THE UNIVERSITY OF ILLINOIS LIBRARY 54 ^ N48t ° r ^ are m the ascendency, the molecules are The wave-length of the blue rav ICt\ o 0000169 inch. ^ ' ' o. 00043 n millimetre, or 4 Introductory Outlines drawn more or less closely together, and the substance assumes the solid state. If the repellent forces have the upper hand the material takes the gaseous condition, while the liquid state may be regarded as resulting from a certain balance of these opposite forces. Changes which matter undergoes by the action of these forces are physical changes, they do not affect the inherent nature and properties of the substance, which properties, as already stated, reside in the molecules themselves. In each of the three states of matter, viz., solid, liquid, or gaseous, the molecules are conceived as being in a state of motion ; they are regarded as executing some vibratory movement within the spaces that divide them. In the solid state this movement is the most restricted for the reason that the intermolecular spaces are in this case the smallest. In the gaseous condition the amplitude of vibration of the molecules is very greatly increased ; for the attractive forces being at a minimum, and the intermolecular spaces being greatest, the molecules have a further distance to travel before they strike one another. Such changes in matter, which are merely the result of altera- tions in the motions of the molecules, are likewise purely physical chancres. Molecules may be defined as the smallest particles of matter which can exist in the free state j or as the smallest weight of matter in which the original properties of the matter are retained. Atoms —It is the belief of chemists that most molecules are possessed of a structure. That is to say, they are not simple, single indivisible masses, but themselves consist of aggregations of still smaller particles, which are held together by the opera- tions of some other force. These particles of which molecules are composed are termed atoms, and the force which holds them together is called chemical affinity, or chemical attraction. To the mind of the chemist, such molecules are little systems, con- sisting of a number of atoms which are attracted to each other by this particular force ; in the ordinary movements of the mole- cule the system moves about as a whole. In this respect it bears some analogy, on an infinitely minute scale, to a solar system. The atoms of a molecule are regarded as in a state of motion as respects one another, possibly revolving about one another, while the entire system, or molecule, at the same time performs its in- dependent movements, just as in a solar system the various members perform various movements towards each other, whi e Molecules and Atoms $ at the same time the whole system travels upon its prescribed orbit. In the case of the heavenly bodies the force which regulates the movements of the individual members of the system amongst themselves, is the same force that controls the motion of the united system, namely, gravitation. What is the precise relation, or difference, if any, between the forces which control the movements of molecules, and those which operate between the atoms of the molecule, is not known ; but as the effects produced are different the latter force is distinguished by the name of chemical affinity. Any change which matter undergoes, in which the integrity of the molecules is not destroyed, is regarded as a physical change ; whde any change which arises from an alteration in the structure of the molecule is a chemical change. For example, the molecules of water consist of three separate atoms, one of oxygen and two of hydrogen ; any change which water can be made to undergo in which these three atoms still remain associated together as the molecule, is a physical change. The water may be converted into ice, or it may be changed into steam ; but these alterations still leave the molecules intact, the three atoms still remain united as an unbroken system, and so long as this is the case chemical change has not taken place.. Suppose now the molecules of water are heated to a much higher temperature than that which is necessary to convert the iT'm \T I' 6 ? 1 ’ b 7 PaSS ' ng eleCtric sparks thr <»tgh the steam. It will then be found that a very different kind of change has come over the substance. The steam, after being so heated, no longer condenses to water again when cooled ; it has been changed into a gas which can be bubbled through water and collected in an inverted vessel fiUed with water standing in a pneumatic trough, and if a flame be applied to this gas a sharp explosion takes place. e change in this case is a chemical change, for the integrity of the molecules of water has been destroyed. The two atoms of hydrogen have become detached from the oxygen atom, and the ongmal triune structure of the system is destroyed ,„£T' S a ' e , th f efore defined as th ‘ smallest particles of matter which can take part in a chemical change. CHAPTER II ELEMENTS AND COMPOUNDS There are certain molecules in which all the atoms present are of the same kind, and there are other molecules which are com- posed of atoms which differ from one another. Thus, in the substance sulphur, all the atoms composing the molecules are alike ; while in water, as already mentioned, there are two distinct kinds of atoms in the molecule. Matter, therefore, is divided into two classes, according as to whether its molecules are composed of similar or of dissimilar atoms. Molecules consisting of atoms of the same kind are termed elementary molecules , and substances whose molecules are so constituted are known as elements ; mole- cules, on the other hand, which contain dissimilar elements are called compound molecules , and substances whose molecules are thus composed are distinguished as compounds. Sulphur, therefore, is an element , and water is a compound. It will be evident that in the case of elementary molecules, whatever processes they may be subjected to, only one kind of matter can be obtained from them ; while in the case of compounds, the molecules consisting of dissimilar atoms, as many different kinds of matter can be obtained as there are different atoms present. By appropriate means the atoms of hydrogen and oxygen in water molecules can be separated, and two totally different kinds of matter, namely, hydrogen and oxygen, can be obtained from this compound. At the present time there are about seventy substances known to chemists which are believed to be elements. In the history of the science it has frequently happened that substances which were considered to be elements have proved, when subjected to new methods of investigation, to be in reality compound bodies : thus, prior to the year 1783, water was thought to be an elementary substance, it was indeed regarded as the very type of an element, until Cavendish and Lavoisier proved that it was composed of 6 Elements and Compounds two entirely different kinds of matter. In the year 1807 Sir Humphrey Davy showed that the substances known as potash and soda, which were believed to be elements, were in reality compound bodies, and he succeeded in separating the constituent atoms in the molecules of these substances, and in obtaining from them two essentially different kinds of matter. It is therefore quite possible, perhaps even probable, that some at least of the forms of matter which are now held to be elements, may yet prove to be compound bodies. The number of compounds is practically infinite. The elements are very unequally distributed in nature, and are of very different degrees of importance to mankind. Some are a solutely essential to life as it is constituted, while others might be blotted out of creation without, so far as is known, their absence being appreciated. The following thirty elements include all the most important (for the complete list see page 21) : Aluminium. Antimony. Arsenic. Bismuth. Bromine. Calcium. Carbon. Chlorine. Copper. Fluorine. Gold. Hydrogen. Iodine. Iron. Lead. Magnesium. Manganese. Mercury. Nickel. Nitrogen. Oxygen. Phosphorus. Platinum. Potassium. Silicon. Silver. Sodium. Sulphur. Tin. Zinc. On account of certain properties common to a large number of the elements, and more or less absent in others, properties which are for the most part physical in character, the elements are lvided into two classes, known as metals and non-metals. The metals generally are opaque, and their smoothed surfaces reflect light to a high degree, thus giving them the appearance known as metallic lustre. They also conduct heat and electricity. Gold silver, copper, iron, are metals ; sulphur, bromine, oxygen, phos- phorus, are non-metals. These two classes, however, gradually merge into one another, and certain elements are sometimes placed in one division and sometimes in the other, depending upon w ether the distinction is based more upon their physical or their chemical properties : thus, the element arsenic possesses 8 Introductory Outlines many of the physical properties of a metal, but in its chemical relations it is more allied to the non-metals : such elements as these are often distinguished by the name metalloids. By general consent the following fifteen elements are regarded as including all the non-metals and metalloids : — Arsenic. Boron. Bromine. Carbon. Chlorine. Fluorine. Hydrogen. Iodine. Nitrogen. Oxygen. Phosphorus. Selenium. Silicon. Sulphur. Tellurium. The number of atoms which compose the various elementary molecules is not the same in all cases : thus in the elements sodium, potassium, cadmium, mercury, and zinc, the molecules consist of only one atom. The molecules of these substances are single particles of matter. The terms molecule and atom , there- fore, as applied to these elements, are synonymous. Such mole- cules as these are called mono-atomic molecules. In many cases elementary molecules consist of two atoms ; such is the case with the elements hydrogen, bromine, chlorine, oxygen, nitrogen, and others. Elementary molecules of this twin or dual nature are known as di-atomic molecules. Only one instance is known in which an elementary molecule consists of a trio of atoms, namely, the molecule of ozone, which is an aggregation of three oxygen atoms. This molecule is said to be tn-atomic. In two cases, namely, arsenic and phosphorus, the molecules are composed of a quartette of atoms, and these elements, therefore, are said to form tetr-atomic molecules. In a large number of instances the atomic constitution of the molecule of the elements is not known. These terms, mono -atomic, di-atomic , &c., are applied exclu- sively to molecules of elements , and are not used in reference to compounds, where the molecules are composed of dissimilar Mechanical Mixtures.— When molecules of different kinds of matter are brought together, one of two results may follow : either they will merely mingle together without losing their identity, that is to say, the atoms composing the individual molecules will still remain associated together as before, or the atoms in the molecules of one kind will attach themselves to certain atoms present in molecules of another kind to form still different molecules ; in other Mechanical Mixtures words, there will be a redistribution of the atoms, whereby diffe- rent systems or molecules are produced. In the first case the result is said to be a simple or mechani- cal mixture, in the second it is the formation of a chemical compound. In a simple mixture, the ingredients can be again separated by purely mechanical methods ; and as the properties of a substance are the properties of the molecules of that substance, it follows that if the integrity of the molecules is not broken, the properties of a mechanical mixture will be those of the ingredients. For example, oxygen is a colourless gas without taste or smell $. hydrogen also is a colourless gas without taste or smell : when these two gases are mixed together, the mixture is gaseous, is colourless, and tasteless, and, being only a mixture, the molecules of one gas can be readily sifted away from the other. Again, charcoal is a black solid, insoluble in water ; sulphur is a yellow solid, also insoluble in water ; nitre is a white solid, readily dissolved by water: when these three substances are finely powdered and mixed together, the result is a mechanical mixture which is solid, and which is dark grey or nearly black in colour. If this mixture be placed in water, the nitre is dissolved away and the charcoal and sulphur are left. When, however, the integrity of the molecules is disturbed when, by bringing together molecules of different substances a rearrangement of the atoms takes place, resulting in the formation of new molecules, then it is said that chemical action has taken place. Chemical action, therefore, always results in the formation of new molecules,— new molecules which are endowed with their own special properties, differing often in the most remarkable and quite inexplicable manner from those of the original molecules. ne or two examples may be quoted in order to illustrate this extraordinary modifying effect of chemical action. The two colourless gases, oxygen and hydrogen, when simply mixed to- gether, give rise, as already mentioned, to a colourless, gaseous mixture, in which the dual molecules of hydrogen and the simi- larly constituted oxygen molecules move about freely amongst each other. By suitable means chemical action may be made to take place between these two Elements, whereby a complete rearrangement of the atoms takes place, resulting in the formation of molecules of water- molecules in which, as has been already io Introductory Outlines mentioned, one atom of oxygen is associated with two atoms of hydrogen. The product of the chemical action is therefore water, while both the forms of matter of which it is composed are gaseous. The air we breathe, and which is necessary to life, consists of a simple mixture of two colourless gases, viz., oxygen and nitrogen : when chemical action takes place between these substances, a brown-coloured gas is produced in which no animal or vegetable life could exist for many minutes, on account of its suffocating nature. Common salt, which is a white solid substance, and not only harmless but even a necessary article of food, contains two atoms in its molecules — one an atom of chlorine, which is a yellow gas, intensely suffocating and poisonous ; and the other an atom of sodium, a soft, silver-like metal, which takes fire in contact with water. Why it is that a molecule, consisting of an atom of chlorine and an atom of sodium held together by chemical affinity, should be endowed with properties so totally different from those of the contained elements, is altogether unknown; and, similarly, it is quite impossible to predicate from the properties of any compound what are the particular elements of which it is composed. Thus, sugar is a white crystalline solid, soluble in water, and possessing a sweet taste, but no one would have ventured to predict that the molecules of this substance were composed of atoms of carbon (i.e., charcoal), a black, tasteless, insoluble solid; hydrogen , , a colourless, tasteless gas ; and oxygen , another colourless, tasteless gas. . . Chemical Affinity.— When molecules, consisting of two atoms, say A B, come in contact with molecules consisting of other two atoms, C D, and a chemical change takes place resulting in the formation of new molecules, A C and B D ; the question naturally arises, Why does the atom A leave the atom B and attach itself to C? In other words, what determines the rearrangement of the atoms into new molecules ? At present no exact answer can be given to this question. Chemists express the fact by saying that the chemical affinity existing between A and C is greater than that exerted by B upon A. This remarkable selective power possessed by the atoms of different elements lies at the root of all chemical phenomena, and it differs between the various elements to an extraordinary degree. Chemical Affinity u For example, the atom of chlorine possesses a very powerful chemical affinity for the atom of hydrogen : when hydrogen mole- cu es, which consist of two atoms, are mixed with chlorine mole- cules, which are also aggregations of two atoms, at first a simple mechanical mixture is obtained, the two different kinds of mole- cules move amongst each other without undergoing change. On ffir ry tl! ma m Pr0V0Catl0n, however ’ the affinit y of the hydrogen atoms or the chlorine atoms can be caused to exert itself; by merely momentarily exposing the mixture to sunlight a complete redistri- bution of the atoms suddenly takes place with explosive violence and new molecules are formed, each containing one atom of hydrogen and one atom of chlorine. Again an atom of nitrogen is capable of associating itself in chemical union with three atoms of the element chlorine, forming a compound whose molecules therefore contain four atoms The chemical affinity between the atoms of chlorine and nitrogen is the astern is, so to speak, in a state of such unstable equilibrium, that the very slightest causes are sufficient to instantly separate the atoms in the most violently explosive manner, and so break up the compound molecule into separate molecules of ch orine and nitrogen. In this case the affinity between one hlonne atom and another chlorine atom is greater than that between chlorine and nitrogen, consequently the redistribution example .* 1 tS “ ° f ° PP ° Slte order t0 that of the former As a rule, those elements which the more closely resemble each other in their chemical habits have the least affinity for each other rar affinity usuaiiy exists — k- « Chemical Aetion.-The actual process of redistribution of the atoms that takes place when molecules of different kinds of matter are brought together, is called chemical action. In many cases chemical action takes place when the substances are merely brought together, while in others it is necessary to expose the bodies to the influence of some external energy: thus chemical action is brought about in a great number of instances by the of hVht ha rh the / UbStanCeS - In some «** «he influence of light has the effect of causing chemical action to take place ■ for example, when the gases chlorine and hydrogen are mfngled together, no chemical action takes place in the dark, but on 1 2 Introductory Outlines exposing the mixture to light the hydrogen and chlorine combine, and form the compound hydrochloric acid. It is upon the effect of light in causing chemical action to take place that the art of photography depends. Chemical action may sometimes be induced by the influence of pressure : thus when the two gases, hydrochloric acid and phos- phoretted hydrogen, are subjected to increased pressure, they combine together to form a crystalline solid compound known as phosphonium chloride. In the same way, by very great mechanical pressure, a mixture of powdered lead and sulphur can be caused to combine together, when they form the compound, lead sulphide. There are also a number of chemical actions that are only able to proceed in the presence of small quantities (often extremely small) of a third substance, which itself remains unchanged at the conclusion of the action. These cases are generally included under the name of catalytic actions : in some of them the modus operandi of the third substance can be traced (see Oxygen, Modes of Formation ; also Chlorine, Deacon’s Process), while in others it is not understood ; thus it is found that a number of chemical actions are quite unable to take place if the materials are abso- lutely dry : for example, the element chlorine has a powerful affinity for the metal sodium, and when these substances are brought together under ordinary conditions, chemical action in- stantly takes place, and the compound known as sodium chloride (common salt) is produced. If, however, every trace of moisture be perfectly removed from both the sodium and the chlorine, no action between these elements takes place when they are brought together, and so long as they remain in this state of perfect dryness no chemical change takes place. The admission into the mixture of the minutest trace of the vapour of water, however, at once induces chemical action between the chlorine and the sodium, but the exact part that the trace of moisture plays in producing this effect, is not known with certainty. (See also foot-note, page 86.) A few interesting cases are known in which chemical action is brought about by the vibration caused by a loud sound or note : for example, the molecules of the gas acetylene consist of two atoms of carbon associated with two of hydrogen ; when a quantity of this gas is exposed to the report produced by the detonation of mercury fulminate, the mere shock of the explosion causes a re- distribution of the atoms whereby solid carbon is deposited and Chemical Action 13 hydrogen set free. We may suppose that the particular vibration produced by the detonation of the fulminate exercises a disturbing effect upon the motions of the atoms constituting the molecules of acetylene, and thereby causes them to swing beyond the sphere of their mutual attractions, and thus the system undergoes disruption and rearrangement. F All known instances of chemical action can be referred to one place" 66 m ° des ’ m whlch the rearrangement of the atoms can take ( 1 .) By the direct union of two molecules to form a more comflex molecule. Thus, if CO and C1C1 represent two mole- cules between which chemical action takes place according to «oms m CO e c,a y Umte ‘° form " m ° IeCU,e C ° Dtaini "S the f °" ( 2 .) By an exchange of atoms taking place between different molecules In its simplest form this is illustrated in the action rinT jT another to form a compound. Thus, if HH and C1C1 stand for two elementary molecules between which chemical action takes place, the result is the formation of the two molecules HC1 HC1. Such a process as this, in which a com pound substance is produced directly from the elements which compose it, is termed synthesis. thJ he Sa . me m ° de ° f chemicaI action ma y also be exemplified by the exact opposite to this process, namely, the resolution of a compound into its constituent elements. Thus, if OHH OHH represent two molecules of the same compound, when chemical action takes place it will result in the formation of the three elementary molecules OO, HH, and HH. Such a process as is 'analysis & C ° mp ° Und ,S reSoIved in ‘° i's elements^ is known (3.) By a rearrangement of the atoms contained in a molecule There are a number of instances of chemical change, in which the molecules of the substance do not undergo any alfemtion in thet Composition— that is to say, no atoms leave the molecule, nor are « h° THe molecuIe stin consists of the same atoms r!n H h .t Change “ “ d ' d before ’ but the ch emical action has caused them to assume new relative positions, or different relative motions with respect to each other. For example, the substances nown to chemists as ammonium cyanate and urea are two totally different and distinct kinds of matter. These molecules, however^ Introductory Outlines each contain the same atoms and in the same number ; they each consist of aggregations of one atom of carbon, one atom of oxygen, two atoms of nitrogen, and four atoms of hydrogen. When am- monium cyanate is gently warmed, the eight atoms composing the molecules undergo this process of rearrangement, and the substance is changed into urea. CHAPTER III CHEMICAL NOMENCLATURE Jf H JLt? meS which u have been given to the various elementary forms of matter are not based upon any scientific system. The names of some have their origin in mythology. Others have received names several bLr^eferencel S ° me Chara e teristic P r °Perty, while those of everai bear reference to some special circumstance connected with heir discovery It has been the custom in modern t0 dh- tmguish metals from non-metals by applying- to tho f ’ ending in the letters «*, and “ more recent discovery all have names with this ternina ion The common metals, however, which have been known sW T bmes such as gold silver, tin, and copper “keep their old ntm Z discovery 6 taught ^ ^ ^ ° f their names with the^erminh, metals in many of their physical properties, but in their chemical relations they are so closely similar to the non-metal sulphur that they are by general consent classed among the non-metals thVare examples of those elements which are distinguished as ^ metalloids On this account selenium is by some chemists termed selenion. ' = s si composing them, thus— elements The compound formed by the chemical union of- Hydrogen with sulphur is called hydrogen sulphide. Sodium „ chlorine „ sodium chloride Copper „ oxygen „ copper oxide. Calcium „ fluorine „ calcium fluoride. Potassium „ iodine „ potassium iodide Introductory Outlines It continually happens, however, that the same two elements combine together in more than one proportion, giving rise to as many different compounds, in which case it becomes necessary to so modify the names that each of the compounds may be dis- tinguished. This is accomplished by the use of certain terminal letters or of certain prefixes ; for example, the element phos- phorus combines with chlorine in two proportions, forming two different compounds— in one the molecules contain one atom of phosphorus united to three atoms of chlorine, in the other the molecules consist of one atom of phosphorus associated with five of chlorine. These two compounds may be distinguished in the following ways : — X atom of phosphorus with 3 atoms of chfcrine forms or — i atom of phosphorus with 3 atoms of chlorine forms pentaMoride. i »» »> »» b ” M 9 The latter method of distinction is the more general, thus— r atom of su!phur with a atoms of oxygen forms sulphur dtadjfa j atom of carbon with ? atom’ of oxygen forms -bon monoxide. Occasionally the prefixes sub and . proto are employed to denote these differences of composition, but their use is more limited and is becoming out of vogue. When more than two compounds are formed by the union of the same two elements, the additional prefixes hypo , under, and per , over, are sometimes use 1 P in a considerable number of instances the systematic names of familiar compounds give way to the vulgar or common names by which they are known, thus— r Ammonia Common) Hydrochloric acid names 1 Sulphuretted hydrogen l Water Hydrogen nitride 'i Hydrogen chloride I Systematic Hydrogen sulphide j names. Hydrogen monoxidej Binary compounds that are formed by the union of elements with oxygen are called the oxides of those elements. Certain of these oxides are capable of entering into combination with water, giving rise to substances known as adds: such oxides are distinguished Chemical Nomenclature \y as acid-forming oxides , or acidic oxides. They are also some- times termed anhydrides. All the non-metallic elements, except hydrogen, form oxides of this order, and the acids derived from them are known as the oxy-acids. Certain other oxides also unite with water, but give rise to com- pounds known as hydroxides. When such oxides,- which are all derived from the metallic elements, are brought into contact with acids, chemical union takes place, and a compound termed a salt is formed. Such oxides are distinguished as salt-forming , or basic oxides. There are also oxides which are neither acidic nor basic. The names of oxy-acids are derived from the name of the particular oxide from which they are formed, thus — Carbon dioxide gives carbonic acid. Silicon dioxide „ silicic acid. When the same element forms two acid-forming oxides, the terminals ic and ous are applied to the acids to denote respectively the one with the greater and the less proportion of oxygen, thus — Sulphur Irioxide gives sulphurcV acid. Sulphur dioxide gives sulphur^zzi" acid. Nitrogen fenloxide gives nitric acid. Nitrogen trioxide gives nitr ous acid. When more than two such acids are known, the additional prefixes hypo or per are made use of. Thus /sulphuric acid denotes an acid containing the highest quantity of oxygen, while nypomtrons acid stands for an acid containing less oxygen than is present in nitrous acid. There is a class of binary compounds formed by the combination of a large number of the elements with sulphur ; these are known as sulphides. Certain of these sulphides are also capable of forming acids which are analogous in their constitution to oxy-acids, but in which the oxygen atoms are substituted by atoms of sulphur. hese acids are known as thio acids (sometimes sulpho acids), and the same system of nomenclature is adopted to distinguish these : thus we have thio-arseni^ acid, thio-arsenzV acid, denoting sulphur^ 6 7 thG aC1<1 Wkh thC SmalICr and thC larger P r °P° rtion of It was at one time believed that all acids contained oxygen, that indeed this element was essential to an acid. The name oxygen j g Introductory Outlines indicates this belief, the word signifying “the acid -P r0 ^ cer '" This view is now seen to have been incorrect, for many acids are known in which oxygen is not one of the constituents^ Thus the elements fluorine, chlorine, bromine, and iodine, which constitute the so-called Halogen group of elements, each c0 ^mes w, hydrogen, giving rise respectively to hydrofluoric, hydrochloric, hydriodic, and hydrobrornic acids. All known acids contain hydrogen as one of their constituents. As already stated, when chemical action takes place between an acid and a base, a salt is formed. Oxy-acids in this way give rise to oxy-salts, thio-acids to thio-salts, and halogen acids to haloid ^The latter salts being binary compounds, their names are given according to the system already explained, such for example as calcium fluoride, sodium chloride, potassium bromide, silver iodide. In the case of the oxy-salts and thio-salts, the names are made up from the names of the acid and of the metal contained in he base, with the addition of certain distinctive suffixes : thus if the acid be one whose name carries the terminal ous its salts wi 1 distinguished by the suffix ite, while the names of the salts derived from acids whose names end in ic are terminated by the letters ate. Nitroar acid and potassium oxide give potassium nitnfc. , , -j .. sulph ite. Sulphunwtf acid „ >> nitMfe . ff* • j ” ” ” sulpha*. Sulphurz^ acid » ” ” The formation of a salt by the action of an acid upon a base is due to the redistribution of the atoms composing the m0 ' , ' c “ les the two compounds, in such a manner that some or hydrogen atoms in the acid molecules, exchange places with certai metallic atoms from the molecules of the base. Acids which con- tain only one atom of hydrogen so capable of becoming exchang for a metal, are termed mono-basic acids ; those with two, fee, four such hydrogen atoms are distinguished respectively as di-basic, tri-basic , and tetra-basic acids. , ps If the whole of the displaceable hydrogen in an acid becomes replaced by the base, the salt formed is known as a normal salt. On the other hand, when only a portion of the hydrogen atoms is displaced by the base, the salt is distinguished as an acid salt Thus sulphuric acid contains two atoms of hydrogen in its molecule (associated with one of sulphur and four of oxygen) ,. Chemical Nomenclature 19 both the hydrogen atoms are exchanged for potassium, the salt obtained is normal potassium sulphate , and when only one is so replaced the salt is known as acid potassium sulphate. By the term acid salt , therefore, must be understood a salt in which one or more of the hydrogen atoms of the original acid are still left in the molecule.* A third class of salts is formed by the association of one or more molecules of normal salt, with one or more additional mole- cules of the base : these are known as basic salts. Thus, carbonic acid and the base lead oxide, form such a salt known as basic lead carbonate. * Some chemists prefer to regard the acids themselves as the hydrogen salts ; accordingly they apply to nitric acid, sulphuric acid, nitrous acid, sulphurous acid, &c., the names hydrogen nitrate, hydrogen sulphate, hydrogen nitrite, hydrogen sulphite, &c., respectively. CHAPTER IV CHEMICAL SYMBOLS CHEMISTS are agreed in adopting certain symbols to denote the atoms of the various elementary forms of matter. The table opposite contains the names of the elements at and in the second column are given the symbols winch are em- ployed to represent their atoms. The names of the rare elements ar !n P rPumber ta of C instances the atomic symbol is the initial 1 letter of the ordinary name of the element : thus Boron, B , Carbon, C , Fluorine F; Hydrogen, H ; Oxygen, O ; Sulphur, S. When more than one element has the same initial, other the fi TZ Tenets of the name, or the first and another that >s pro- minently heard in pronouncing the word are employed, as Bromine Br • Cobalt, Co ; Chlorine, Cl ; Platinum, Pt. In some cases letters taken from the Latin names for the elements are used, such as Antimony {Stibium), Sb ; Gold C Aurum ), Au ; Silver {Argentum), Ag ; Lead {Plumbum), Pb ; and Iron {Ferrum), Fe. These symbols are not intended to be employed as mere short- hand signs to be substituted as abbreviations for the full names hand sign Jn e case they denote one atom of the element The symbol H stands for one atom of hydrogen, the stunbol O stands for one atom of oxygen ; Cl means one atom of cUorine, and Ag represents one atom of silver. No other use of th Tt e hr b Tt '^“mentioned (page 8) that the molecules^ of th different elements are composed of different --hers of morns Chemical Symbols 21 Atomic Weights. Atomic Weights. 1 234 1 234 Name. Atomic Symbols. Approxi- mate Values. More Exact * Values. Name. Atomic Symbols. Approxi- mate Values. More Exact * Values. Aluminium . . . Antimony ( Stibium Argon .... Arsenic .... Barium .... Beryllium . Bismuth . . . Boron .... Bromine . . . Cadmium . . . Caesium Calcium . . . Carbon .... Cerium .... Chlorine . . . Chromium . . . Cobalt .... Copper {Cuprum). Didymium . Erbium .... Fluorine. . . . Gallium .... Germanium . . Gold (Aurum) Hydrogen . . . Indium .... Iodine .... Iridium . . . . Iron (Ferrum ) . Lanthanum Lead {Plumbum ) . Lithium .... Magnesium . . Manganese . . Mercury {Hydr- ) argyrum . . j ( A1 Sb A As Ba Be Bi B Br Cd Cs Ca C Ce Cl Cr Co Cu Di Er F Ga Ge Au H In I Ir Fe La Pb j Li Mg! Mm Hg 27 120 ? 75 137 9 207 5 11 80 112 133 40 12 141*2 35*5 52 59 G3 145 166 19 70 72 197 1 113 127 192*5 56 138*5 207 7 24 55 200 1 27.04 119.6 74-9 136.86 9.08 10.9 79.76 hi . 7 132.7 39-91 11.97 35-37 52.45 58.6 63.18 19.06 69 . 86 I 196.8 XI 3-4 126.54 55-88 206.39 7.01 23-94 54-8 199.8 Molybdenum . Nickel . . . . Niobium . . . Nitrogen . . . j Osmium. Oxygen .... Palladium . . . Phosphorus . . Platinum . . . Potassium {Kal- ) ium) . . . j Rhodium . Rubidium . Ruthenium. . . Samarium . Scandium . Selenium . . . Silicon .... Silver {Argentum) Sodium {Natrium) Strontium . . . Sulphur .... Tantalum . . . Tellurium . . . Thallium . Thorium . . . Tin {Stannum) . Titanium . Tungsten . . . Uranium . Vanadium . Ytterbium . . . Yttrium Zinc Zirconium . . . Mo Ni Nb N Os O Pd P Pt K Rh Rb Ru Sm Sc Se Si A g Na Sr S Ta Te T1 Th Sn Ti W U V Yb Y Zn Zr 96 59 93*7 14 191 16 106 31 195 39 104 85 103*5 150 44 79 28 108 23 87*3 32 182 125 203*7 232 118 48 184 239 8 51*1 173 89*6 65 90*4 95-9 58.6 14.01 15.96 106.2 30.96 194-3 39-03 104. 1 85.2 ;;; ' 1 43-97 78.87 28.3 107.66 22.995 31.98 xx 7.35 64 .’ 88 represented by the symbol 0 3 , while the tetr-atomic character of the phosphorus molecule is expressed in the symbol P 4 . The composition of compound molecules is expressed by placing the symbols of the atoms which compose such molecules in juxta- position : thus a molecule consisting of one atom of sodium (symbol * Where no values are given in this column, those in the third column may be regarded as accurate. 22 Introductory Outlines Na) and one atom of chlorine (symbol Cl) is represented by the united symbols of these two elements, NaCl ; a compound con- sisting of one atom of carbon and one atom of oxygen by the symbols of these two atoms, CO. Such arrangements of symbols representing molecules are termed molecular formula, or, simply, formula. When the molecule contains more than one atom of any parti- cular element, this fact is indicated by the use of numerals placed immediately after the symbol to be multiplied : thus, a molecule of water consists of two atoms of hydrogen and one atom of oxygen, th formula for water is therefore H 2 0. One molecule of ammonia, consisting of an atom of nitrogen with three atoms of hydrogen, is represented by the formula NH 3 ; and a molecule of sulphuric acid, which is an aggregation of two atoms of hydrogen, one atom of sulphur, and four atoms of oxygen, has the formula HoS0 4 . . . It is sometimes necessary to represent the presence in a mole- cule of certain groups of atoms, groups which seem to hold together, and often to function as a single atom. This is accomplished by the use of brackets : thus (NH 4 ) 2 S0 4 is the formula for a molecule containing one atom of sulphur, four atoms of oxygen, eight atoms of hydrogen, and two atoms of nitrogen ; the nitrogen and hydrogen atoms being present as two groups, in each of which one nitrogen atom is associated with four hydrogen atoms. Such groups of atoms are termed compound radicals. When it is required to indicate more than one molecule of the same substance, numerals are placed immediately in front of the formula : thus 2H 2 0 signifies two molecules of water, and 3NH, expresses three molecules of ammonia. By means of these symbols and formulae, chemists are enabled to represent, in a concise manner, the various chemical changes which it is the province of chemistry to examine. Such changes are usually termed chemical reactions , and they are represented in the form of equations in which the symbols and formulae of the reacting substances as they are before the change are placed on the left, and those of the substances which result from the change upon the right, thus — H 2 + Cl 2 = 2HC1 HgCl 2 + 2KI = Hgl 2 + 2KC1. The sign 4- has a different significance as used on the left side Chemical Symbols 23 of the equation to that which it bears upon the right. On the left hand it implies that chemical action takes place between the substances, while on the opposite side it has the simple algebraic meaning. Thus, the second of the above equations is understood to mean, that when the compounds, mercuric chloride and potassium iodide, are brought together in such a way that chemical action results, a redistribution of the atoms will take place, resulting in the formation of mercury iodide and also potassium chloride. As further illustrations of the use of chemical symbols, the following three examples may be given as exemplifying the three modes of chemical action mentioned on page 13 : — (1) NH 3 + HC1 = NH 4 C1. Ammonia combines with hydrochloric acid, and gives ammonium chloride. (2) H 2 S0 4 + Na 2 C0 3 = Na 2 S0 4 + C0 2 + H 2 0. Sulphuric acid combines with normal sodium carbonate, and gives normal sodium sulphate, carbon dioxide, and water. (3) (CN)0(NH 4 ) = (NH 2 ) 2 CO. Ammonium cyanate is converted into urea. In all cases where the nature of the chemical change is under- stood, it is capable of expression by such equations, and as matter is indestructible, every atom present in the interacting molecules upon the left of the expression, reappears on the right hand side in some fresh association of atoms.* * See also Chemical Notation, chapter vii. ) CHAPTER V THE ATOMIC THEORY The atomic view as to the constitution of matter, briefly sketched out in Chapter I., forms a part of what is to-day known as the atoniic theory. When chemical changes were carefully studied from a quantita- tive standpoint, four laws were discovered in obedience to which chemical action takes place. These laws are distinguished as the laws of chemical combination. Three of these generalisations refer to quantitative relations as respects weight ; while one expresses quantitative relations with regard to volume , and only relates to matter in the gaseous state. I. Law of Constant Proportion.— The same compound always contains the same elements combined together in the same proportion by weight j or expressed in other words, The weights of the con- stituent elements of every compound bear an unalterable ratio to each other , and to the weight of the compound formed. II. Law of Multiple Proportions.— When the same two elements combine together to form more than one compound, the different weights of one of the elements which unite with a constant weight of the other , bear a simple ratio to one another ; or this law may be stated thus : When one element unites with another in two or more different proportions by weight, these proportions are simple multiples of a common factor. III. Law of Reciprocal Proportions, or Law of Equivalent Proportions. — The weights of different elements which combine separately with one and the same weight of another element, are either the same as, or are simple multiples of, the weights of these different elements which combine with each other ; or in other words, The relative proportions by weight in which the elements, A, B, C, D, &c., combine with a constant weight of another element, X, are the same for their combinations with any other element, Y. 24 The Atomic Theory 25 IV. Law of Gaseous Volumes, or The Law of Gay-Lussae. — When chemical action takes place between gases, either elements or compounds , the volume of the gaseous product bears a simple relation to the volumes of the reacting gases. These four laws are the foundations upon which the whole superstructure of modern chemistry rests. (1.) The Law of Constant Proportions.— When two sub- stances are mingled together, and remain as a mere mechanical mixture, they may obviously be present in any proportion, and it was at one time thought that when two substances entered into chemical combination with each another, they could do so also in any proportion, and that the composition of the resulting com- pound would vary from this cause. This belief was finally disproved, and the law of constant proportions definitely estab- lished by Proust in the year 1806. The same compound, therefore, however made, and from whatever source obtained, is always found to contain the same elements united together in the same proportion by weight. Thus, common salt, or, to adopt its systematic name, sodium chloride, which is a compound of the two elements sodium and chlorine, may be made by bringing the metal sodium into contact with chlorine gas, when the two elements unite and form this compound. It can also be made by the action of hydrochloric acid upon the metal sodium, or by adding hydrochloric acid to sodium carbonate, and by a variety of other chemical reactions. When the sodium chloride obtained by any or all of these processes is analysed, it is invariably found to contain the elements chlorine and sodium in the proportion by weight of 1 : 0.6479, or, expressed centesimally — Sodium . . 39.32 Chlorine . . 60.68 100.00 and when this is compared with the sodium chloride as found in nature, obtained either from the salt-mines of Cheshire, or the celebrated mines in Galicia, or by evaporating sea-water, it is found that the composition of the compound in all cases is exactly the same. In the same way the compound water, consisting of the two elements hydrogen and oxygen, whether it be prepared synthetically by causing the two elements to unite directly, or obtained from any natural source, as rain, or spring, or river, is 26 Introductory Outlines found to contain its constituent elements hydrogen and oxygen in the ratio by weight of i : 8, or, Hydrogen . . 11.12 Oxygen . . 88.88 100.00 If in the formation of sodium chloride by the direct combination of its constituent elements, an excess of either one or other be present beyond the proportions 39.32 per cent, of sodium and 60.68 per cent, of chlorine, that excess will simply remain unacted upon. If eight parts by weight of hydrogen and eight parts by weight of oxygen be brought together under conditions that will cause chemical action, the eight parts of oxygen will unite with one part of hydrogen, and the other seven parts of hydrogen merely remain unchanged. This fact, that elements are only capable of uniting with each other in certain definite proportions, marks one of the most characteristic differences between chemical affinity and those other forces, such as gravitation, that are usually distinguished as physical forces, for although there are many instances known in which the extent to which a chemical action may proceed (that is, the particular proportion of the reacting bodies which will undergo the permutation that results in the formation of different mole- cules) is influenced by the mass of the acting substances, it never governs the proportion in which the elements combine in these compounds. . , It follows from the law of constant composition that the sum of the weights of the products of a chemical action will be equal to that of the interacting bodies ; and upon the validity of this law depend all processes of quantitative analyses. (2.) The Law of Multiple Proportions was first recognised by Dalton, who investigated certain cases where the same two elements combine together in different proportions, giving rise to as many totally distinct compounds. These proportions, however, were always found to be constant for each compound so produced, so that this law formed no contradiction to the law of constant composition. The simple numerical relation existing between the numbers representing the composition of such compounds will be evident from the following examples. The two* compounds of * In Dalton’s day these two substances were the only known compounds of carbon with hydrogen. The Atomic Theory 27 carbon with hydrogen, known as marsh gas and ethylene , are found to contain these elements in the proportions — Marsh gas . . i part by weight of hydrogen with 3 parts of carbon. Ethylene . . 1 „ M m 6 The two compounds of carbon with oxygen contain these ele- ments in the proportion — Carbon monoxide . 1 part of carbon with 1.334 parts of oxygen by weight Carbon dioxide .1 ,, f| 2.667 The elements nitrogen and oxygen form as many as five different compounds, in which the two elements are present in the propor- tions — Nitrous oxide . . 1 part of nitrogen with o. 571 parts of oxygen by weight Nitric oxide. . . 1 ,, m x I43 ^ s Nitrogen trioxide .1 ,, M 1.714 Nitrogen peroxide 1 ,, M 2.286 Nitrogen pentoxide 1 ,, )t 2.857 The relative proportions of carbon combining with a constant weight of hydrogen in the two first compounds are as 1 : 2. Those of oxygen uniting with a constant weight of carbon in the second example are also as 1 : 2, while in the nitrogen series the relative proportions of oxygen in combination with a constant weight of nitrogen are as I : 2 : 3 : 4 : 5. (3.) Law of Reciprocal Proportions.— Known also as the law of proportionality, or the law of equivalent proportions. When the weights of various elements, which were capable of uniting separately with a given mass of another element, were compared together, it was seen that these weights bore a simple relation to the proportions in which these elements combined amongst them- selves. For example, the elements chlorine and hydrogen each separately combine with the same weight of phosphorus, the pro- portions being — v Phosphorus : chlorine = i : 3.43 Phosphorus : hydrogen = 1 : 0.097 The elements chlorine and hydrogen can combine together, and they do so in the proportion — Chlorine : hydrogen = 35.5 : 1 but 35 : 1 = 3*43 •' 0.097 2 g Introductory Outlines Therefore the proportions by weight in which chlorine and hydrogen separately combine with phosphorus, is a measure of t e proportion in which they will unite together. Again, the two elements carbon and sulphur each separate y combine with the same weight of oxygen, the proportion being- Oxygen : carbon = i : 0.375 Oxygen : sulphur =1:1 But the elements carbon and sulphur themselves unite together, and in the proportion— Carbon : sulphur = 0.1875 : 1 but 0.1875 : 1 = °*375 : 2 Therefore the proportion by weight in which carbon and sulphur separately unite with the same mass of oxygen, is a simple multiple of that in which these two elements combine together. These remarkable numerical relations will be rendered still more evident, by comparing the proportions in which the members of a series of elements combine with a constant weight of various other elements . thus — Hydrogen. Sodium. Potassium. Silver. Mercury. 0 02817 0.6479 i-° 2 3*04 2.816 unite separately with i part. It will be seen that the proportion in which these numbers stand to each other is, as — 1 23 : 39 '• io 7 : 100 : 3S ' S Let us now compare these proportions with those in which the same elements unite with a constant weight of the element bromine — dercury. Bromine. 1.25 unite with 1 part. Hydrogen. Sodium. Potassium. Silver. 0.0125 0.2875 0.4875 1*34 or as — 1 23 : 39 : 107 80 Each of these five elements in like manner combines with oxygen, and the weights which are found to unite with a constant mass of oxygen are — Hydrogen. Sodium. Potassium. Silver. Mercury. 2.875 4.87s 13-38 13-5 unite with i part. 0.125 again as — T 23 39 107 The Atomic Theory 29 The same relation will appear in the case of the combination of these five elements Vith a constant weight of sulphur — Hydrogen. Sodium. Potassium. Silver. Mercury. Sulphur. 0.0625 1.4375 2.4375 6.69 6.25 unite with i part, or as — 1 : 23 : 39 : 107 : 100 • 16 It is thus evident that the proportions in which the members of such a series combine with a constant weight of one element, is the same as that in which they unite with a constant mass of another element. One part by weight of hydrogen combines with 35.5 parts of chlorine, 80 parts of bromine, 8 parts of oxygen, and 16 parts of sulphur — that is to say, these proportions of these four elements satisfy the chemical affinity of 1 part of hydrogen ; they are therefore said to be equivalent. Twenty-three parts of sodium is likewise equivalent to 35.5 parts of chlorine, 80 parts of bromine, 8 parts of oxygen, and 16 parts of sulphur, and by the same reasoning it is also equivalent to 1 part of hydrogen, 39 parts of potassium, 107 parts of silver, and 100 parts of mercury. These numbers, therefore, are known as the equivalent weights of the elements, or their combining proportions , and the combining weight of an element may therefore be defined as the smallest weight of that element which will combine with 1 part by weight of hydrogen. This law of proportionality, or reciprocal proportion, was dis- covered by Richter, but it was left for Dalton to trace the connec- tion between these three generalisations. Dalton adopted and adapted an ancient theory concerning the ultimate constitution of matter which was expounded by certain of the early Greek philo- sophers. The exponents of this theory held that matter is built up of vast numbers of minute indivisible particles, in opposition to the antagonistic theory believed by others, namely, that matter was absolutely homogeneous and capable of infinite subdivision. Dalton embraced the ancient doctrine of atoms, and extended it into the scientific theory which is to-day known as Dalton’s atomic theory, and is accepted as a fundamental creed by modern chemists. According to this theory, matter consists of aggregations of minute particles, or atoms, which are indivisible. Dalton con- ceived that chemical combination takes place between atoms— that is to say, when chemical action takes place between two elements, it is due to the union of their atoms ; the atoms coming into juxtaposition with each other under the influence of chemical 30 Introductory Outlines affinity, are held together by the operation of this force. He furthei assumed that the atoms of the various elements possessed different relative weights, and that the relations existing between these weights, was the same as that between the weights in which experi- ment had shown the elements to be capable of combining together- In other words, he said that the numbers representing the combin- ing proportion of the elements expressed also the relative weights of the atoms. Let us now see how this theory satisfies and explains the first three laws of chemical combination. (i.) The Law of Constant Composition.— It has already been shown, p. 25, that the compound sodium chloride, wheresoever and howsoever obtained, contains the elements chlorine and sodium in the proportion — Chlorine : sodium = 1 : 06479. These numbers have been shown on p. 28 to represent the com- bining proportions— Chlorine : sodium = 35*5 : 2 3 - Now the atomic theory states, that sodium chloride is formed by the union of atoms of chlorine with atoms of sodium, and that the relative weights of these atoms is expressed by the combining weights of the elements, namely, 35-5 and 2 3 - lf > therefore, sodium is to combine with chlorine, since atoms are indivisible masses, it follows that the compound produced by the union of one atom ol each of these two elements must always have the same composi- tion. _ . . ... (2.) The Law of Multiple Proportions.— The ratio in which oxygen combines with hydrogen to form the compound water, is seen on p. 27 to be as 8 : 1. This number 8, therefore, we will for the present argument regard as the relative weight of the atom of oxygen.* . Oxygen combines with carbon as already mentioned, forming two different compounds ; in the first, the elements are present in the proportion — Carbon : oxygen = 1 : 1-334 = 6:8 That is to say, in the proportion of one atom of carbon to one atom * For reasons which will be explained later, chemists now , r f “ d 1 '“f “ 16 as representing (in round numbers) the relat.ve weight of the atom oxygen. The Atomic Theory 31 of oxygen. According to the theory, if the atom of carbon unites with more oxygen than one atom, it must at least be with two atoms. It may be with three or with four, but as the compound must be formed by the accretion of these indivisible atoms, the increment of oxygen must take place by multiples of 8. When the second compound is examined it is found to contain its constituent elements in the proportion — Carbon : oxygen = 1 : 2.667 = 6 : 16 That is to say, in the proportion of one atom of carbon to two atoms of oxygen. This information respecting the composition of these two compounds is conveyed both in their names and their formulae. The first is termed carbon monoxide, and its formula is expressed by the symbol CO ; while the second is distinguished as carbon dioxide, and has the symbol CO* The difference in the composition of the five compounds that nitrogen forms by union with oxygen, will be made evident by the aid of this theory. The proportion of nitrogen to oxygen in these compounds is — (1.) Nitrogen : oxygen = 1 : 0.571 = 14 : 8 (2.) Nitrogen : oxygen = 1 : 1.143 = 14 : 16 (3.) Nitrogen : oxygen = 1 : 1.7 14 = 14 ; 24 (4-) Nitrogen : oxygen = 1 : 2.268 =14 : 32 (5.) Nitrogen : oxygen = 1 : 2.857 = 14 : 40 And it will be seen that the increase in the proportion of oxygen in the compounds, takes place by the regular addition of a weight of t at element equal to 8, which at the present stage of the argument we are regarding as representing the relative weight of the atom of oxygen. (3 ) The Law of Reciprocal Proportions.— If the illustrations given on p. 27 of the operation of this law be examined in the light of the atomic theory, their explanation will be evident : thus, the relative proportions in which hydrogen and chlorine separately combine with phosphorus is 0.097 : 3-43, and the ratio between these numbers is as 1 : 35.5, which is the proportion in which these two elements are known to unite together to form hydrochloric acid, t hese numbers, however, represent the relative weights of the atoms of these elements, therefore hydrochloric acid may be sup- posed to be formed by the union of one atom of hydrogen with one atom of chlorine. 32 Introductory Outlines Again, the relative weights of carbon and sulphur which sepa- rately combine with a constant weight of oxygen, are — carbon 0.375, sulphur 1, and the ratio between these numbers is as 6 : 16. Carbon and sulphur, however, unite together in the relative proportion — Carbon : sulphur = 0.1875 : 1 = 6 : 32 Therefore the compound they produce, may be supposed to consist of one atom of carbon, having the relative weight 6, and two atoms of sulphur, each with the relative weight 16. CHAPTER VI ATOMIC WEIGHTS In the third column of the table on page 21, the numbers are given, which are at the present time generally accepted by chemists as representing the approximate atomic weights of the elements. These numbers depart, in many instances, from those arrived at by Dalton’s methods ; thus, the relative weights of carbon, oxygen, nitrogen, and sulphur, which were found to be equivalent to one part of hydrogen, are carbon = 6,* oxygen = 8, nitrogen = 4.66, sulphur = 16 ; while the figures given as the approximate atomic weights of these elements in the table, are carbon = 12, oxygen = 16, nitrogen = 14, sulphur = 32. We must now discuss some of the chief reasons for these departures. In the two compounds of carbon and hydrogen known to Dalton, namely, marsh gas and ethylene, the proportions of carbon to hydrogen are— In ethylene . . . Carbon : hydrogen = 6:1. In marsh gas . . Carbon : hydrogen = 6:2. Dalton therefore concluded that ethylene was a compound con- taining 1 atom of carbon, united with 1 atom of hydrogen, and to which, therefore, he gave the formula CH ; and that marsh gas consisted of 1 atom of carbon combined with 2 atoms of hydrogen, and which he accordingly represented by the formula CH 2 . There was, however, nothing to prove that the weight of carbon was constant in the two compounds, for it will be obvious that the same ratio between the weight of carbon and hydrogen will still be maintained by assuming that the hydrogen is constant, and that the carbon varies, thus — In marsh gas . . Hydrogen : carbon 1:1:3. In ethylene . . Hydrogen : carbon : : 1 : 3 x 2. * These are the numbers which Dalton ought to have obtained had his methods of determination been more exact. The figures he actually found for the combining weights of these four elements were respectively 5, 7, 5, 13. 34 Introductory Outlines That is to say, the ratios are not disturbed by the assumption that in marsh gas we have i atom of hydrogen combined with i atom of carbon, having the relative combining weight of 3, and in ethylene 1 atom of hydrogen united with 2 atoms of carbon. It will be evident, however, that if we could gain any exact information as to the actual number of atoms which are present in these various molecules, this difficulty would no longer exist. For example, suppose it were possible to ascertain that in the molecule of marsh gas there were 4 atoms of hydrogen, then as the relative weights of hydrogen and carbon in this compound are as 1 : 3, the weight of the carbon atom would obviously have to be raised from 3 to 1 2 ; and if it could be determined that in the ethylene molecule there were also 4 atoms of hydrogen, then seeing that the ratio of hydrogen to carbon in this substance is as 1 : 6, we should conclude that it contained 2 atoms of carbon, of the relative weight not less than 12, and the composition of the two compounds would be expressed by the formulae, marsh gas CHfc ethylene C 2 H 4 . Again, the relative weights of hydrogen and oxygen in water are as 1 : 8. If the molecule of water contains only 1 atom of hydrogen, then we conclude that 8 represents the relative weight of the oxygen atom, and the formula for water will be HO. But suppose it to be discovered that there are 2 atoms of hydrogen in a molecule of this compound, then it becomes necessary, in order to retain the ratio between the weight of these constituents (a ratio ascertained by analyses), to double the number assigned to the oxygen atom, and to regard its weight as 16, as compared with 1 atom of hydrogen, and the formula for water in this case would be H 2 0 . The compound ammonia contains the elements hydrogen and nitrogen in the ratio — Hydrogen : nitrogen : : 1 : 4.66. If the molecule of ammonia contains only 1 atom of hydrogen, then 4.66 represents the relative weight of the nitrogen atom, and the formula will be NH; but if it should be found that there are 3 atoms of hydrogen in this molecule, then again the relative weight assigned to the nitrogen must be trebled in order to pre- serve the ratio, and it will have to be raised from 4-66 to 14 (m round numbers), and the formula for ammonia will be NH3* From these considerations it will be evident, that it is of the Atomic Weights 35 highest importance to gain accurate knowledge as to the actual number of atoms which are contained in the molecules of matter- in other words, to learn the true atomic composition or structure of molecules ; and it may be said that this problem has occupied the minds of chemists from the time that Dalton published his atomic weights, in the year 1808, down to the present time. There is no single method of general application, by means of which chemists are able to determine the atomic weight of an element, but they are guided by a number of independent considerations, some of which are chemical in their character, while others are of a physical nature ; and that particular number which is in accord with the most of these considerations, or with what are judged to be the most important of them, is accepted as the true atomic weight. The chief methods employed for determining atomic weights may be arranged under the following four heads r. Purely chemical methods. 2. Methods based upon volumetric relations. 3. Method based upon the specific heats of the elements. 4. Method based upon the isomorphism of compounds. 1. As an illustration of the chemical processes from which atomic weights may be deduced, the following examples may be given, namely, the case of the two elements oxygen and carbon. Oxygen combines, as already stated, with hydrogen in the proportion — TT , Hydrogen : oxygen =1:8. When water is acted upon by the element sodium, the compound is decomposed and hydrogen is evolved ; and it is found that if 18 grammes of water are so acted on, 1 gramme of hydrogen is evolved, and 40 grammes of a compound are formed, which contains sodium, together with all the oxygen originally in the 18 grammes of water, and some hydrogen. This compound, under suitable conditions, can be acted upon by metallic zinc, and when these 40 grammes are so acted on, 1 gramme of hydrogen is again evolved, and 72.5 grammes are obtained of a compound containing no hydrogen, but sodium and zinc combined with all the oxygen originally contained in the 18 grammes of water. It will be evident, therefore, that the hydrogen contained in water can be expelled in two equal moieties ; there must, therefore, be two atoms of hydrogen in this compound. By no known 36 Introductory Outlines process can the oxygen be withdrawn from water in two stages ; thus, if 18 grammes of water are acted upon by chlorine, under the conditions in which chemical action can take place, 73 grammes of a compound containing only chlorine and hydrogen are found, and the whole of the oxygen is thrown out of combination, and evolved as gas. It is therefore concluded, that water contains in its mole- cule, 2 atoms of hydrogen and 1 atom of oxygen, and as they are combined in the relative proportion of 1 : 8, the atomic weight of oxygen cannot be less than 16. No compounds have been found in which a smaller weight of oxygen, relative to one atom of hydrogen, than is represented by the number 16 (approximately), is known to take part in a chemical change. The compound marsh gas contains hydrogen and. carbon in the proportion by weight of 1:3. By acting on this compound with chlorine, it is possible to remove the hydrogen from it in four separate portions. By the first action of chlorine upon 16 grammes of marsh gas, 1 gramme of hydrogen is removed in combination with 35.5 grammes of chlorine, and a compound containing carbon, hydrogen, and chlorine, in the ratio 12 : 3 : 35.5, formed. By the successive action of chlorine, three other moieties of hydrogen can be thus withdrawn, each being in combination with its equivalent (35.5 parts) of chlorine. The second and third com- pounds that are formed contain carbon, hydrogen, and chlorine in the ratios 12 : 2 : (35.5 x 2) and 12 : 1 : (35.5 >< 3 )- The compound produced by the fourth action of chlorine, whic withdraws the fourth portion of hydrogen, contains only carbon and chlorine, in the ratio 12 : (35.5 x 4)- From the fact that the hydrogen contained in marsh gas can thus be removed m four separate portions, the molecule must contain four hydrogen atoms, and therefore the atomic weight of carbon must be at least 12. No compounds of carbon are known in which a smaller weight of carbon, relative to one atom of hydrogen, than is represented by the number 12, takes part in a chemical change. The definition of atomic weight, furnished by considerations of a chemical nature, may be thus stated : the atomic weight of an element, is the number which represents how many times heavier the smallest mass of that element capable of taking part in a chemical change is, than the smallest weight of hydrogen which can so function. Atomic Weights 37 The choice of hydrogen as the unit of atomic weights is a purely arbitrary selection ; but since atomic weight values can only be determined relatively, it j becomes necessary to select some one element and to assign to its atom some j P artlcu lar number to serve as a standard. As hydrogen is the lightest of all elements, Dalton originally adopted it, and arbitrarily fixed unity as the number which should stand for its atomic weight. The disadvantages of this ! particular unit are twofold : in the first place the number of elements that form hydrogen compounds that are suitable for atomic weight determinations is very n = being held together in a manner resembling that by which the . numbers i o _a cosmical system are bound together. The lines simply denote that the atoms are held to each other by the attractive force which we call chemical affinity. Valency 6 1 Compounds of this order, in which one of the elements has still unsatisfied affinities, are called unsaturated compounds. In its power to satisfy the affinities of an element, a divalent atom is equal to two monovalent atoms : thus, when the affinities of the tetravalent carbon atom are saturated with oxygen, the mole- cule contains two atoms of oxygen, which may be symbolically expressed thus, O = C = O, in which the four affinities of the carbon (represented by the four lines) are satisfied by the two divalent atoms of oxygen. Carbon, however, combines with a smaller proportion of oxygen, forming the compound carbon mon- oxide, CO. The carbon atom in this case is divalent, as expressed by the formula C = O, and this substance is also an unsaturated compound. The number of divalent atoms with which an element can unite, cannot, however, be taken as a safe criterion or measure of the valency of that element in cases where that number is greater than i ; for example, in such a compound as calcium oxide, CaO, we regard the two affinities of the divalent atom of oxygen as being satisfied by two affinities possessed by the calcium, and express this belief in the formula Ca = O, and regard the calcium as divalent. In the same way, in carbon monoxide, CO, the carbon being united with one atom of the divalent element oxygen, is itself divalent in this compound; but in the case of carbon dioxide, where the carbon atom is united with two atoms of divalent oxygen, we are not justified in asserting that the atoms are united, as represented by the formula O = C = O, in which the four affinities of carbon are represented as saturated with oxygen. There exists no posi- tive proof that the carbon is not divalent in this compound, and that the molecule does not consist of three divalent atoms united, C as shown in the formula /\. From the fact, however, that O O carbon forms a compound with four atoms of hydrogen, and another with four atoms of chlorine, we know that this element is tetravalent, and therefore we believe that in carbon dioxide it is also tetravalent. Again, as measured by its compound with hydrogen, sulphur is divalent ; while with chlorine it forms SCI*. But sulphur unites with oxygen, forming che two compounds, sulphur dioxide S0 2 and sulphur trioxide S0 3 . If it be assumed that in these molecules the whole of the oxygen affinities are satisfied with sulphur, then the 62 Introductory On times symbolic representation of these oxides will be O = S = O, and O = S = O, the sulphur being in one case tetravalent, and in the O other hexavalent. There is, however, no positive proof that the affinities of one oxygen atom are not partially satisfied by union with another oxygen atom, and that the valency of the sulphur is higher than either two or four, as seen in the alternative formulae, S S0 2 / > S0 3 O — O ?\ , S = O ; or 0/ S / \ 0 — 0—0 Although there are no known compounds in which an atom of sulphur is united with six monovalent elements, sulphur is regarded by many chemists as capable of fulfilling the functions of a hexa- valent element. It will be evident from these considerations, that in many cases the valency of an element is a variable quantity, de- pending partly upon the particular atoms with which it unites. It is also found that it is dependent in many instances upon tem- perature and upon pressure. Thus, between a certain limited range of temperature, one atom of phosphorus combines with five atoms of chlorine in the compound PC1 6 , but above that limit two atoms of chlorine leave the molecule, and the phosphorus becomes trivalent. Again, if phosphoretted hydrogen, PH S , be mixed with hydro- chloric acid, HC1, and the mixed gases be subjected to increased pressure, the gases combine-and form a solid crystalline compound known as phosphonium chloride, PH 4 C1, in which the phosphorus atom, being united with five monovalent atoms, is pentavalent. When the pressure is released, an atom of hydrogen and an atom of chlorine leave the molecule, and the phosphorus returns to its trivalent condition. A compound, in whose molecules there is an atom which for the time being is not functioning in its highest recognised valency, often exhibits a readiness to unite with additional atoms to form new compounds : thus, ammonia combines eagerly with hydro- chloric acid, forming ammonium chloride — NH 3 + HC1 = NH 4 C1. Carbon monoxide unites directly with chlorine to form carbonyl CO + Cl 2 = COCl 2 . chloride — Valency 63 Carbon monoxide also combines with an additional atom of oxygen, and gives carbon dioxide, thus — In this last action it will be seen that the molecule of carbon monoxide, in being converted into the dioxide, takes up one atom of oxygen ; but as the molecule of oxygen is the smallest isolated particle, it follows that the two atoms contained in such a molecule must first separate, and each one then furnishes the requisite additional oxygen for one molecule of carbon monoxide. In the union of carbon monoxide with chlorine, and of ammonia with hydrochloric acid, are we to suppose that the same action takes place? That is to say, do the two atoms in the molecule of chlorine separate from each other and unite with carbon, thereby satisfying its tetrad valency, in the manner here expressed ? — And in the case of ammonia and hydrochloric acid, do the hydrogen and chlorine atoms part, and each unite with the nitrogen atom, thereby raising it from the trivalent to the penta- valent condition ? thus — Or are we to suppose that the two molecules, without losing their integrity, become held together as independent molecules, by virtue of the unsatisfied affinities of the carbon, or the nitrogen, as the case may be, in which case the compounds might be repre- sented thus — 2CO + 0 2 = 2C0 2 . ci — : — ci + co = Cl H : \/ H — : — C1 + H — N — H = H — N — H. H H Cl j C =0 Cl H — Cl H — N — H H This question would be settled by determining the vapour density of the compound. If, for instance, we were to find the vapour density of ammonium chloride to be 26.75, then the com- 64 Introductory Outlines pound having the composition NH 4 C 1 would have the normal molecular volume, that is, its molecule would occupy two unit volumes,* and the conclusion would be that the vapour consisted of single molecules of the composition represented by the formula NH 4 C 1 . But ammonium chloride at ordinary temperatures is a solid, and when heated to the temperature necessary to convert it into vapour its molecules break up into separated molecules of the two original gases — ammonia, NH 3 , and hydrochloric acid, HC1.+ So that we are unable to gain any information in this direction as to the mode in which the atoms are disposed in the compound. When the two gases are brought together under ordinary condi- tions, they combine with the evolution of considerable heat, owing to loss of energy ; this is taken as evidence that true chemical action, in the sense of atomic rearrangement, has resulted, hence it is believed that in this compound the nitrogen is united with the five monovalent atoms, and consequently is pentavalent. In the case of carbonyl chloride, COCl 2 , the vapour-density can be ascertained, this compound existing in the gaseous condition at the ordinary temperature. Its vapour-density, determined by experiment, is found to be 50.6. This number, divided into the molecular weight of the compound having the composition COCl 2 , gives practically the number 2 as the molecular volume of the compound. Hence we conclude that these four atoms constitute a single molecule. There is a certain class of combinations, in which molecules of different compounds unite, that do not so readily admit of explanation, because in neither of the molecules is there any atom functioning in a lower state of valency than that which it is known to be capable of. For example, the monovalent elements, fluorine and hydrogen, form the compound hydrofluoric acid, HF ; fluorine also combines with the monovalent element potassium, forming potassium fluoride, KF. Both of these com- pounds come under the head of saturated compounds, in the sense that neither of them contains an atom which is known to be capable of exercising a higher valency than it exhibits in these compounds. Nevertheless these two molecules unite together and form a definite chemical compound, known as hydrogen-potassium fluoride. Again, the divalent element zinc combines with two atoms of See p. 42. f See Dissociation, p. 86. Valency 65 ' the monad element chlorine, forming zinc chloride, ZnCl 2 ; the two monovalent elements, sodium and chlorine, also combine, giving the compound sodium chloride, NaCl. Both of these substances must be regarded as saturated compounds, and yet they unite with each other, forming a distinct chemical compound, known as sodium zinc chloride. Such compounds as these are known as double salts , and examples might be multiplied almost indefinitely. A similar union of molecules, where the recognised valency of the atoms is all satisfied, is seen in a large number of compounds containing water for crystallisation ; * for example, the divalent element copper, in combination with two atoms of chlorine, forms cupric chloride, CuCl 2 . The divalent element oxygen, in combination with two hydrogen atoms, forms water, H 2 0. When cupric chloride crystallises from aqueous solution, each molecule of the chloride unites to itself two molecules of water, which is therefore termed water of crystallisation. In chemical notation, it is usual to represent compounds of this order, by placing the formulas of the different molecules that have entered into union, in juxtaposition, with a comma between ; accordingly, the examples here quoted would be indicated thus — Hydrogen potassium fluoride . HF,KF. Sodium zinc chloride .... ZnCl 2 ,NaCl. Crystallised cupric chloride . . . CuC1 2 ,2H 2 0. Combinations of this order are by no means confined to the union of two kinds of molecules, as the following examples will serve to show : — Platinum sodium chloride . . PtCl 4 ,2NaCl,6H 2 0. Mercuric potassium chloride . 2HgCl2,KCl,2H 2 0. At the present time our knowledge of the nature of the union between these various molecules is too imperfect to admit of any precise explanation ; such compounds are frequently distinguished as molecular combinations. It must be remembered that our ideas of valency are based mainly upon the consideration of matter in the gaseous state ; at present we have little certain knowledge as to the valency of elements in liquid and solid compounds. Most of the compounds belonging to the class we are now discuss’ng are solid, and See p. 192. E 66 Introductory Outlines it is quite conceivable that the valency of an element might increase, as the compounds in which it functions pass from the gaseous to the liquid and solid There is also another consideration that must not be overlooked. The unit of measure that has been adopted for estimating valency, namely, i mono- valent atom, is probably only an extremely rough and crude measure, which is incapable of appreciating smaller differences of combining capacity that may, and most probably do, exist. Its use may be compared to the adoption of a single unit, say i gramme, for the estimation of mass, or weight ; when, if a given quantity of matter has a weight equal to i gramme but less than 2 grammes, its weight would be i ; if greater than 2 grammes but less than 3, then its weight would be 2 — a method of estimating which tacitly assumes that no intermediate weights of matter between the various multiples of the selected unit are possible. There is no evidence to show that the com- bining capacity of an element is exactly expressed by simple multiples of a monovalent atom. For example, in the simplest form of combination, such as that between hydrogen and chlorine— where the molecule contains 1 atom of each element —1 hydrogen atom unites with 1 chlorine atom, that is to say, with a mass of chlorine weighing 35.5 times its own weight; and we say that the mutual affinities of these atoms are satisfied. But for anything we know to the con- trary, an atom of hydrogen may have an affinity for chlorine which would enable it to unite with a mass of chlorine weighing 40 or 45 or 50 times its own weight, but not a mass weighing 71 (35- S x 2) times its own. But since a mass of chlorine 35.5 times the weight of a hydrogen atom is the smallest quantity that is ever known to take part in a chemical change, is the chemically indivisible mass we call an atom, it follows that as the hydrogen atom has not sufficient combining capacity to unite with 2 atoms, it is compelled to be satisfied with 1 i t might still, however, retain a residual combining capacity. Or the residual combining capacity may be lodged in the chlorine atom, which may be conceived as being able to unite with a greater weight of hydrogen than is represented by 1 atom, but not so much as that of 2 atoms. Each of the elements fluorine, chlorine, bromine, and iodine unites with 1 atom of hydrogen, and we represent their compounds in a similar manner, thus : — H-F; H - Cl ; H-Br; H-I; but we make an enormous assumption if we suppose that in each of these compounds the mutual affinities of the atoms is equally satisfied. For example, the fluorine compound exhibits a tendency to unite itself to other compounds of fluorine (and to a much more marked degree than is seen in the case of hydrochloric acid), resulting in the formation of such double fluorides as the following Hydrogen potassium fluoride .... Hydrogen bismuth fluoride . . • • BiF 3 ,3HF. Hydrogen silicon fluoride (Hydro-fluo-silicic acid) SiF 4 ,2HF. And there is reason for believing that the molecules of hydrofluoric acid Valency 6 7 HF,HF a * ble ° f Unit ‘ ng ara0ngSt themse1 "*. forming the double molecule Assuming .he residual combining power to reside in the fluorine atom, and representing this by means of dotted lines, we may express the composition of tnese compounds thus— F-F - H I H — F---F-H, H — F - F — Bi - F---F — H ; NTT ■ ,- F \ - F H - F< \si< See Hydrofluoric acid. CHAPTER IX general properties of gases Under the head of the general properties of gases it will be con- venient to consider the following subjects 1. The relatioh of gases to heat. 2. The relation of gases to pressure. 3. The liquefaction of gases. 4. Diffusion of gases. 5. The kinetic theory of gases. The Relation of Gases to Heat. -The fact that substances expand when heated, and again contract upon being cooled, was observed in very early times. The fact also that all substances do not undergo the same alterations in volume when subjected to the same changes of temperature has been long known ; but it was not until the beginning of the nineteenth century that it was proved by Charles and Gay-Lussac that all gases expanded and contracted equally when exposed to the same alterations of temperature. This law is generally known as the Law of Charles, and may e thus stated : When a gas is heated , the pressure being constant, it increases in volume to the same extent whatever the gas may be. The increase in bulk suffered by 1 volume of a gas in being heated from 0° to i° is termed the coefficient of expansion, and if the law of Charles is true all gases will have the same coefficient. Modern research has shown that the law of Charles is not abso- lutely true, and the extent to which gases deviate from the strict expression will be seen from the coefficients of expansion given in the following table : — * The study of these subjects belongs more especially to the scl ““ physics or chemico-physics. For fuller information on these L included within the scope of this book students are referred to special treatises on physics Relation of Gases to Heat 69 Air . Hydrogen . . . . .003667 Carbon monoxide . . . .003667 Nitrogen . . . . .003668. Nitrous oxide . . . .003676 Carbon dioxide . . .003688 Cyanogen . . .003829 Sulphur dioxide . .003845 It will be noticed that the first four gases have almost the same coefficient of expansion : these gases are all very difficult of lique- faction, and it will be seen that the coefficient rapidly rises in the case of the other gases, which are easily liquefied. For purposes of ordinary calculation it is usual to adopt the coefficient of expansion of air, as applicable to all gases. It will be obvious that since the volume of a gas is affected by alterations of temperature, it becomes necessary, when measuring the volume of a gas, to have regard to the particular temperature at which the measurement is made, and in order to compare volumetric measures they must be all referred to some standard temperature. This standard temperature is by general consent o° C. Taking the fraction .003665 therefore for the coefficient— 1 volume of a gas at o° becomes 1 + .003665 volumes at i° 1 » „ o° „ 1 + .003665 x 2 „ 2 0 or 1 » „ o° „ 1 + .003665 / „ t° Therefore the volume at t° equals the volume at o° multiplied by 1 + .003665 t. Let v be the volume at t\ and v 0 the volume at o°, then — v — v 0 (\ + .003665 /) and conversely the volume at o° equals the volume at t° divided by 1 + .003665 / — 1 + .003665 t The vulgar fraction equivalent to .003665 is 2 tt?. 273 volumes at o° become 273 + / at f. What is known as the absolute temperature of a substance is the number of degrees above - 273 0 C. Taking this point as the zero, the absolute temperature of melting ice, for example, will be 273 0 . Charles’ law, therefore, may be thus stated : The volume of any 70 Introductory Outlines gas , under constant pressure, is proportional to the absolute tern- perature. The Relation of Gases to Pressure. — The effect of increase of pressure upon a gas is to diminish its volume. The law which connects the volume occupied by a gas, with the pressure to which it is subjected was discovered by Robert Boyle (1661), and is known as Boyle’s Law. It may be thus stated : The volume occupied by a given weight of any gas is inversely as the pressure. The general truth of this law may readily be illustrated by subjecting a gas to varying pressures, and it will be seen that when the pressure is doubled the volume of gas is reduced to one-half, and so on. Just as in the case of the law of Charles, modern investigations have shown that the law of Boyle is not a mathematical truth. It is found not to be absolutely true of any gas, for with the exception of hydrogen, all gases are more compressible than is demanded by the law. Hydrogen deviates from the law in an opposite sense, in that it requires a higher pressure than the law would indicate, in order to reduce a volume of it to a given point. These deviations from Boyle’s law are explained by the operation of two causes : first, the attraction exerted by gaseous particles upon each other ; second, the fact that increased pressure diminishes the space between the molecules, and not the actual space occupied by the molecules of a gas. When the former cause predominates, the gas deviates from the law by being more compressible ; in the case of hydrogen, the second cause operates more powerfully. (See Kinetic Theory of Gases.) For ordinary purposes of calculation the law of Boyle may be regarded as true. As the volume of a given weight of gas is so intimately related to the pressure, and as the atmospheric pressure is variable, it becomes necessary, in all quantitative manipulation with gases, to know the actual pressure under which the gas is at the time of measurement, and to refer the volume to a standard pressure. The pressure that has been adopted as the standard is that of a column of mercury 760 mm. in height. (See Atmosphere.) If v equals the volume of gas measured at p pressure, and v 0 the volume at the standard pressure, then In practice it is most usual to make both correction for tempe- Liquefaction of Gases yi rature and pressure together; then v 0 being the volume at the standard temperature and pressure, we get * . = Z jL* i + .003665 1 760 The Liquefaction of Gases.— Under certain conditions of tem- perature and pressure, the law of Charles and the law of Boyle both completely break down. According to the law of Charles, 100 c.c. of a gas at o° C. should occupy 96.4 c.c. if the tem- perature were lowered to - io°. If 100 c.c. of the gas sulphur dioxide at o° C. be confined in a glass tube standing in mercury, and the gas be cooled to — io° by surrounding the tube with a freezing mixture, it will be found that the volume of gas, instead of occupying 96.4 c.c., has been reduced to a few cubic .centi- metres only, and that the surface of the mercury in the tube is wet owing to the pressure of a minute layer of a colourless liquid upon it. In this case the law of Charles has broken down, and the sul- phur dioxide has passed from the gaseous to the liquid state. Similarly, according to the law of Boyle, 100 c.c. of a gas measured at the standard pressure should occupy 25 c.c. when exposed to a pressure of four additional atmospheres. If 100 cc. of the gas sulphur dioxide be enclosed in one limb of a long U-tube, as shown in Fig. 1, the other limb being filled with air, and the two gases be simultaneously exposed to increased pressure by raising the mercury reservoir, it will be seen that at first the gases in both tubes are compressed equally. As the pressure approaches three atmospheres, however, the mercury will be seen * The student should familiarise himself with the method of calculating the changes of volume suffered by gases, by changes of temperature and pressure, by working out a number of examples such as the following : — 1. If 30 litres of gas are cooled from 25° to o°, what is the diminution in volume, the pressure being constant? A ns. 2.51 litres. 2. If a litre of air at o° degrees weighs 1.293 grammes when the barometer 72 Introductory Outlines to rise much more rapidly in the tube containing the sulphur dioxide, and when the mercury reservoir has been raised to such a height that the gases are subjected to four atmospheres, the sulphur dioxide will have completely broken down, and will be entirely con- verted into a few drops of liquid, which appear upon the surface of the mercury. The air meantime, in the other limb, will be found to occupy 25 c.c., as that gas at that pressure obeys Boyle’s law almost absolutely. We see, therefore, that at a certain temperature and at a certain pressure the gas sulphur dioxide begins rapidly to depart from the laws of Charles and Boyle, and ultimately passes into the liquid condition. All known gases, with the one exception of hydrogen, when exposed to certain conditions of temperature and pressure, condi- tions which are special for each different gas, will pass from the gaseous to the liquid state : and as the point at which liquefaction takes place is approached, the departures from Boyle’s law become more and more pronounced. The first substance, recognised as being under ordinary condi- tions a true gas, that was transformed into the liquid condition, was chlorine, which was liquefied in the year 1806 by Northmore. The true nature of this liquid was not understood until Faraday inves- tigated the subject. In his earlier experiments, Fara- day’s method consisted in sealing into a bent glass tube (Fig. 2) sub- stances which, when heated, would yield the gas ; the substances being contained in one limb of the tube, and the empty limb being immersed in ice. The pressure exerted by the gas thus generated in a con- fined space, was sufficient to cause a portion of it to condense to is at 760 mm. , what will be the weight of a litre of air at 27 , the barometer standing at the same height? Ans. 1.177 grammes. 3. What will be the weight of a litre of air at 42 0 when the barometer stands at 735 mm. ? Ans. 1.084 grammes. 4. Air at a temperature of 15° is enclosed in a vessel and heated to 93 0 . Compare the pressure of the enclosed air with that of the atmosphere. Ans. As 61 : 48. 5. What will be the volume, at the standard temperature and pressure, of 500 c.c. of hydrogen, measured at 20°, and under a pressure of 800 mm. ? Ans. .490 c.c. 73 Liquefaction of Gases the liquid state, and the liquid collected in the cooled limb. In this way Faraday liquefied such gases as chlorine, sulphur dioxide, ammonia, cyanogen. In his later experiments, Faraday compressed the gas by means of a small compression pump, and at the same time applied a low degree of cold, and by so doing he succeeded in liquefying carbon dioxide, hydrochloric acid, nitrous oxide, and other gases. There were a number of gases, however, which F ara- day found it impossible to liquefy, such as hydrogen, oxygen, nitro- gen, marsh gas, nitric oxide, carbon monoxide, &c. It became the custom to call these permanent gases, and this term was applied to them until the year 1877. In that year it was proved by Pictet, and independently by Cail- letet, that under sufficiently strong pressure, and a sufficiently low degree of cold, the so-called permanent gases could in the same way be reduced to the liquid condition. Pictet’s method was in principle the same as that employed by Faraday, the difference being, that with the machinery at his disposal, he was able to employ enormously increased pressure, and a greater degree of cold. For the liquefaction of oxygen, a quantity of potassium chlorate was heated in a strong wrought iron retort, to which was connected a long horizontal copper tube of great strength and small bore. At the extreme end of this tube there was a pressure gauge capable of indicating pressures up to 800 atmospheres, and a stop- cock. The tube was cooled by being contained in a wider tube, through which a constant stream of liquid carbon dioxide, at a tem- perature of — j2o° to — 140°, was caused to flow. The machinery employed to maintain this flow of liquefied car- bon dioxide was somewhat elaborate, consisting of condensing and exhaust pumps for liquefying and rapidly evaporating sulphur dioxide, and similar condensing and exhaust pumps for liquefying and rapidly evaporating carbon dioxide : the sulphur dioxide being merely the refrigerating agent used to assist the liquefaction of the carbon dioxide. This machinery was driven by two eight- horse-power engines. As the potassium chlorate was heated and oxygen evolved, the internal pressure in the retort and copper tube rapidly rose, and its amount was indicated by the gauge. When the stopcock upon the end of the tube was opened, liquid oxygen was forcibly driven out in the form of a jet. In the method employed by Cailletet, the pressure to which the gas is subjected is obtained by purely mechanical means. The 74 Introductory Outlines gas to be liquefied is introduced into a glass tube (Fig. 3), the narrow end of which consists of a strong capillary tube. The tube carries a metal collar, which enables it to be secured in position in the strong steel bottle (Fig. 4), by means of a nut, E' (Fig. 5), which screws into the mouth. The bottle, which is partially filled with mercury, is connected, by means of a flexible copper tube of fine bore, with a small hydraulic pump, by means of which water is forced into the steel bottle. The water so driven in, forces the Fig. 3. Fig. 4. Fig. 5. mercury up into the glass tube T, and thereby compresses the contained gas. In this way a pressure of several hundred atmos- pheres may be applied to the gas. In his earlier experiments, Cailletet depended almost entirely for the refrigeration he required, upon the fact, that when a gas is allowed suddenly to expand, it undergoes a great reduction in temperature. This method of cooling may be termed internal refrigeration. In the case of oxygen, the gas was first subjected to a pressure of 300 to 400 Liquefaction of Gases 75 atmospheres, and was then allowed suddenly to expand by a rapid release of the pressure. The result of the sudden expansion was to momentarily lower the temperature of the gas to such a point that the tube was filled with a fog, or mist, consisting of liquid particles of oxygen. This principle, namely, the self-cooling of a gas by its own sudden expansion, has recently been applied for the liquefaction of oxygen in large quantities. When oxygen under considerable pressure, say 120 atmospheres, is allowed to escape from a fine orifice at the end of a long pipe, the issuing gas suddenly expands, and thereby its temperature is greatly lowered. If this self-cooled gas is made to sweep over the pipe from which it is escaping, it will cool the pipe, and therefore lower the temperature of the remaining gas be- fore it issues. In this way the cooling effect becomes cu- mulative, the initial temperature of the gas before it es- capes being con- tinually brought lower and lower, until at last the point is reached at which the oxygen is liquefied.* If the oxygen be first cooled to about — 8o° by means of solid carbon di- oxide, then in a few minutes , by the fur- ther cooling due to its own expansion, the temperature will fall below the boiling-point of oxygen, and the liquefied gas be obtained. The apparatus for the purpose is shown in Fig. 6.+ Oxygen Under a pressure of 120 to 140 atmospheres is passed through a * Linde, The Engineer , Oct. 4, 1895. + Designed by Dewar; constructed by Messrs. Lennox, Reynolds & Fyfe. Fig. 6. j6 Introductory Outlines series of spirals of fine copper pipe contained in the chamber C, which is encased in a non-conducting jacket of cork-dust. The gas enters by the pipe O (seen in the enlarged section), and passes through the spiral S S, which is immersed in a mixture of alcohol and solid carbon dioxide (the liquid carbon dioxide from the reservoir being admitted into the alcohol through the valve W, which is regu- lated by the screw B). The oxygen thus cooled passes through the double spiral pipe D D, which ultimately extends through the bottom of the chamber, and terminates in a stirrup, U, the short end of which is closed. In the bend of this stirrup there is a fine hole, which can be closed or opened at will by the pointed end of the rod V, connected to the screw A. On opening this valve, the oxygen, already cooled to about — 8o°, escapes from the hole under a pressure of 120 to 140 atmospheres. It instantly expands and is thereby cooled still lower. This cold gas is prevented from escaping at once into the atmosphere by the glass tube G, but is compelled to rush upwards (as shown by the arrows), and, sweep- ing past the double spiral, D D, cools this pipe, and therefore the succeeding portions of issuing oxygen. In a few minutes the temperature of this pipe is thereby brought so low, that the further cooling of the gas by its expansion causes the liquefaction of a portion of it, and a fine spray of liquid is seen to spurt out from the hole. This spray quickly increases in quantity, and rapidly collects as a clear liquid in the glass tube G. This tube is double-walled, the space between the walls being perfectly vacuous. In such a vessel the liquid oxygen may be kept for a considerable time, eva- porating only very slowly in spite of its extremely low boiling-point, as it has been found that such a vacuous envelope forms the most perfect non- conductor. Oxygen can also be liquefied by the low tempera- ture obtainable by the rapid evaporation of liquid ethylene; and, similarly, by the rapid evaporation of liquid oxygen itself, such a low temperature can be reached that almost all known gases have by this means been reduced to the liquid state (see Hydrogen, p. 157). Thus, if a quantity of liquid oxygen in the glass tube O (Fig. 7), which is provided with a vacuous envelope, V, be caused to boil rapidly by putting the pipe P in connection with an exhaust pump, the temperature can be lowered to - 200°, when air itself becomes liquefied without the application of pressure, and drops of liquid ,-o N Fig. 7. Critical Temperature of Gases 77 air quickly collect upon the walls of the inner empty tube, N, which is freely open to the atmosphere. In this way considerable quan- tities of liquefied air can be collected in a few minutes. The Critical Point. — As far back as the year 1869, it was shown by Andrews that when liquid carbon dioxide was heated to a particular temperature, it passed from the liquid to the gaseous state, and that no additional pressure was able to condense it again so long as the temperature remained at or above that point. This particular temperature is called the critical point , or the critical temperature of the gas. In the case of carbon dioxide this critical temperature is 31.9 0 , and in order that this gas may be liquefied by pressure, it is an essential condition that the temperature be below that point ; above 32 0 no pressure is capable of bringing about liquefaction. All gases have a critical temperature, which is special for each gas, and until the temperature of the gas be lowered to that point, liquefaction is impossible. The critical temperatures of the different gases vary through a very wide range, as will be seen from the following examples : — Nitrogen -146.0° Carbon monoxide . . -140.0° Oxygen — 1 1 8.8° Marsh gas .... — 81. 8° Ethylene + io.i° Carbon dioxide ... 31 .9° Nitrous oxide . . . 35 - 4 * Acetylene ..... 37 -o' Hydrochloric acid . . 52 - 3 ‘ Ammonia 13°.°' Chlorine 141.0' Sulphur dioxide . . . 155 - 4 ' The gases in this list, from ethylene downwards, all have their critical temperatures so high, that there is no difficulty in cooling them below these points. These are the gases which were first reduced to the liquid state. The first four upon the list have very low critical temperatures ; these are the very gases which for so long resisted all attempts to liquefy them, and which were on that account called permanent gases. We now know that the failure to obtain them in the liquid state, was owing to the fact that the relation between the critical temperature and the point of liquefaction was not fully realised. Just as carbon dioxide cannot be liquefied unless its temperature be brought down to 3 1. 9 0 , so oxygen resists liquefaction under the highest possible pressures, until its temperature be lowered to - 118.8 0 , the critical temperature of oxygen. 7 8 Introductory Outlines The critical temperature of a gas is sometimes spoken of as the absolute boiling point. Critical Pressure.— The particular pressure that is required to liquefy a gas at its critical temperature, is called the critical pressure. Thus the pressure necessary to liquefy oxygen, when the temperature has been lowered to - 118.8 0 , is 50 atmospheres ; while that required to condense chlorine at its critical point, viz., + 14 1 °, is 84 atmospheres. Taken at their respective critical points, therefore, chlorine is a more difficultly liquefiable gas than oxygen, although at o° chlorine is condensed by a pressure of only 6 atmospheres ; o°, however, is 14 1° below the critical point of chlorine, and it is more than probable that if it were possible to cool oxygen to a temperature 141 0 below its critical point, that is, to - 259.8°, it, in like manner, would be capable of liquefac- tion by very slight pressure. Diffusion of Gases. — If a jar filled with hydrogen be placed mouth to mouth with a jar of air, the hydrogen being uppermost, it will be found that after the lapse of a few minutes some of the hydrogen will have passed into the bottom jar containing air, and some of the air will have made its way up into the hydrogen jar. The light gas hydrogen does not, as might have been supposed, remain floating upon the air, which is 14.44 times as heavy, but gradually escapes into the lower jar ; and the heavier gas finds its way, in opposition to gravitation, into the upper jar. This process goes on until there is a uniform mixture of air and hydrogen in both jars, and the gases never separate again according to their densities. This transmigration of gases will take place even through tubes of considerable length : thus, if two soda-water bottles be filled one with hydrogen, and the other with oxygen, and the two bottles be connected by a piece of glass tube a metre in length, the system being held in a vertical position with the light hydrogen upper- most, it will be found after an hour or two that the two gases have become mixed. Some of the hydrogen will have descended through the long tube into the lower bottle, and in like manner a portion of the oxygen, although nearly sixteen times as heavy as hydrogen, will have travelled up into the top bottle. That the gases have so mixed may be readily shown by applying a lighted taper to the mouth of each bottle, the detonation which then takes place proving that the bottles contained a mixture of oxygen and hydrogen. This passage of one gas into another is called the diffusion of gases. It was observed by Graham that when the 79 Diffusion of Gases two gases were separated from each other by a thin porous septum, such, for instance, as a piece of unglazed porcelain (so- called “biscuit”), or plaster of Paris, the pressure of the gas on the two sides of the porous partition did not remain the same during the process of diffusion : that is to say, one gas made its way through the partition faster than the other, and it was noticed that the lighter the gas, the more rapidly was it able to transpire or diffuse through the porous medium. This fact, viz., that a light gas diffuses more rapidly than a heavier one, may be observed in a variety of ways.* The apparatus seen in Fig. 8 is a modified form of Graham’s diffusiometer. It consists of a long glass tube with an enlargement or bulb near to one end. Into the short neck of this bulb there is fastened a thin diaphragm of stucco, or other porous material. If the apparatus be filled with hydrogen by dis- placement, the short neck being closed by a cork, and the long limb be immersed in water, it will be seen, upon the withdrawal of the cork, that the water rapidly rises in the long tube. The hydrogen diffusing out through the diaphragm so much more rapidly than air can make its way in, a diminution in pressure * See Experiments Nos. 350-359, Newth’s " Chemical Lecture Experiments,” new ed. 80 Introductory Outlines within the apparatus results, and this causes the water to ascend in the tube. The same phenomenon may be seen even more strikingly by means of the apparatus, Fig. 9, which consists of a tall glass U-tube, upon the end of one limb of which there is fastened, by means of a cork, a porous cylindrical pot, such as is used in an ordinary Bunsen battery. The U-tube is half filled with coloured water. Under ordinary circumstances air is continually diffusing through the porous pot, but as it passes at an equal rate in both directions, there is no disturbance of the pressure, and consequently the coloured water remains level in the two limbs. If now a beaker containing hydrogen be brought over the apparatus, as seen in the figure, the hydrogen will stream through the porous pot so much more rapidly than the air in the pot can make its way out, that there will be an increase in the total amount of gas inside the apparatus, which will be instantly rendered evident by the change of level of the liquid in the U-tube, the water being forcibly driven down the tube which carries the porous pot. Upon removing the beaker the reverse operation will at once take place ; the hydrogen inside the apparatus now rapidly diffuses out, and much more quickly than air can pass in, consequently a reduction of pressure within the apparatus results, which is indicated by a disturbance of the level of the water in the tube, in the opposite direction to that which occurred at first. The Law of Gaseous Diffusion. — Graham established the law according to which the diffusion of gases is regulated, and it may be thus stated : The relative velocities of diffusion of any two gases are mversely as the square roots of their densities. The density of hydrogen being 1, that of air is 14.44, the velocity of the diffusion of hydrogen therefore, as compared with that of air, will be in the ratio of ^14.44 to V14.44 = 3-8, J 1 = 1. Therefore hydrogen diffuses 3.8 times faster than air ; or 3.8 volumes of hydrogen will pass out through a porous septum, while only 1 volume of air can enter. If d = the density of a gas, air being unity, and v = the volume of the gas which diffuses in the same time as 1 volume of air, then The following table gives in the last column, the results obtained by Graham, which will be seen to accord very closely with the cal- culated numbers demanded by the law of diffusion : — Diffusion of Gases 8i Name of Gas. Density of Gas compared with Air =d. Volume of Gas which Diffused in the same Time as one Volume of Air. Hydrogen .... 0.06926 3-7794 3-83 Marsh gas 0-559 1-337 5 1*344 Carbon monoxide 0.9678 1.0165 1. 0149 Nitrogen .... 0.9713 1.0147 1. 0143 Oxygen .... 1.1056 0.9510 0.9487 Sulphuretted hydrogen 1. 1912 0.9162 l o-95 Carbon dioxide . 1.5290 0. 8087 0.812 Sulphur dioxide 2.247 0.6671 0. 68 The property of diffusion is sometimes made use of in order to separate gases, having different densities, from gaseous mixtures. This process of separation by diffusion is known as atmolysis. The principle may readily be illustrated by causing a mixture of oxygen and hydrogen, in proportion to form an explosive mixture, to slowly traverse tubes made of porous material, such as ordinary tobacco pipes. Two such pipes may be arranged as shown in Fig. io, and the gaseous mixture passed through in the direction indicated by the arrow. On collecting the issuing gas over water in a pneu- matic trough, it will be found to have so far lost the hydrogen, by diffu- sion through the tube, that a glowing splint of wood when introduced into it, will be re- ignited. From the rate of dif- fusion of ozone, in a mix- ture of ozone and oxygen, Soret was able to calcu- late the density of this allotropic form of oxy- gen, and so confirm the result he had previously obtained by other methods (see Ozone). Attempts have been made to utilise this principle in order to obtain oxygen from the air. The relative densities of oxygen and 82 Introductory Outlines nitrogen are as 16 to 14 ; the rate of diffusion, therefore, of nitrogen is slightly greater than that of oxygen. Effusion is the term applied by Graham to the passage of gases through a fine opening in a very thin wall, and he found that it followed the same law as diffusion. Bunsen utilised this principle for determining the density, and therefore the molecular weights, of certain gases. The method, in essence, is as follows : — A straight glass eudiometer is so constructed, that a gas contained in it can be put into communication with the outer air through a minute pin-hole in a thin platinum plate. The gas is confined in the tube, which is placed in a cylindrical mercury trough, by means of a stop-cock at the top. When the tube is depressed in the mercury, and the cock opened, the gas escapes through the minute perforation in the platinum plate, and its rate of effu- sion is determined by the time occupied by a glass float, placed in the tube, in rising a graduated distance within the eudiometer. The flow of gases through capillary tubes is called transpiration of gases. In this case the friction between the gas and the tubes becomes a factor in the movement, so that this phenomenon is not governed by the same law as gaseous diffusion. The Kinetic Theory of Gases.— The term kinetic signifies motion, and as applied to this theory it expresses the modem views of physicists concerning matter in the gaseous state, and serves to harmonise and explain the physical laws relating to the properties of gases. Matter in the state of gas or vapour, is regarded as an aggregation of molecules in which the attractive forces which tend to hold them together, are reduced to a minimum, and in which the spaces that separate them are at a maximum. These molecules are in a state of rapid motion, each one moving in a straight line until it strikes some other molecule, or rebounds from the walls of the containing vessel, when it continues its move- ment in another direction until it is once more diverted by another encounter. As they constantly encounter and rebound from each other, it will be evident that at any given instant some will be moving with a greater speed than others ; the majority, however, will have an average velocity. In these encounters no loss of energy results so long as the temperature remains constant, but any change of temperature results in a change in the velocity of movement of the molecules, the speed being increased with increased heat. The actual volume of the molecules is very small as compared with the space occupied by the mass ; the space The Kinetic Theory 83 between the molecules, therefore, in which they pass to and fro, is relatively very great. As the molecules are constantly colliding and rebounding, the distances between them, as well as their speed, will be sometimes greater and sometimes less ; but there will be an average distance, which is known as the mean free path of the molecule. The pressure exerted by a gas, or its elastic force, is the combined effect of the bombardment of its molecules against the containing vessel ; in other words, the pressure of a gas is proportional to the sum of the products obtained by multiplying the mass of each molecule by half the square of its velocity. It will be obvious that if the space within which a given mass of gas is confined be reduced, the number of impacts of the molecules against the walls of the containing vessel, in a given time, will be increased, and therefore the pressure it exerts, or its elastic force, will also be increased. If the space be reduced to one-half the original, the number of these impacts will be doubled, or in other words, the number of impacts in a given time is inversely as the volume. This statement is simply the law of Boyle stated in the language of the kinetic theory. When a given mass of gas contained in a confined space is heated, the pressure it exerts, or its elastic force, is increased. But as the number of molecules present has not been increased by raising the temperature of the gas (provided no chemical decom- position of the gas is brought about by the change of temperature), the increased pressure can only have resulted from the greater frequency, and greater energy, of the impacts of the molecules against the walls of the vessel, owing to their greater velocity. Two equal' volumes of different gases under the same conditions of temperature and pressure, exert the same elastic force upon the containing vessels, that is to say, the kinetic energy in each volume is the same. According to Avogadro’s hypothesis, equal volumes of all gases under the same conditions of temperature and pressure, contain an equal number of molecules, however much the weight of these molecules may vary ; therefore the average kinetic energy of each individual molecule will be the same. It follows from this that the mean velocities of different molecules must vary, and the calculated numbers representing the actual velocities of movement of the molecules of different gases, show that these rates are pro- portional to the inverse square roots of their respective densities. But according to the law of gaseous diffusion (Graham’s law), the 84 Introductory Outlines relative rapidity of diffusion of gases is inversely proportional to the square roots of their densities, hence by purely mathematical processes, based upon the kinetic theory of gases, the law of gaseous diffusion is proved to be true. The deviations from the laws of Boyle and Charles, already referred to,* are also explained by the dynamical theory of gases, from considerations of the following order : — 1. That the molecules themselves are not mathematical points, but occupy a space ; in other words, the space occupied by the actual particles of matter is not infinitely small as compared with the entire volume of the gas, i.e., the bulk of the particle plus the intermolecular spaces. 2. That the impact of the molecules against each other and against the containing envelope occupies time ; or, in other words, the time occupied by the impacts is not infinitely small compared with the time elapsing between the impacts. 3. That the molecules themselves are not entirely without attrac- tion for each other ; that is to say, although the attractive force between the molecules which holds them together in the liquid and solid states of matter, is at a minimum in the case of gases, it is not entirely absent. See page 7a CHAPTER X DISSOCIA TION Dissociation is the term employed to denote a special class of chemical decompositions. When potassium chlorate is heated it breaks up into potassium chloride and oxygen, thus — 2KC10 3 = 2KC1 + 30 2 , and when calcium carbonate (chalk) is heated it breaks up into calcium oxide (lime) and carbon dioxide — CaC0 3 = CaO +C0 2 . In the first case the oxygen is incapable of reuniting with the potassium chloride, but in the second, the carbon dioxide can recombine with the lime and reproduce calcium carbonate ; there- fore both the following expressions are possible — CaCO s = CaO + C0 2 , and CaO + C0 2 = CaC0 3 . Reactions of this order are known as reversible reactions, and the breaking up of calcium carbonate by the action of heat is termed dissociation, while that of the potassium chlorate under similar circumstances is simple decomposition. When ammonia is passed through a tube heated to a dull red heat, the gas is decomposed into nitrogen and hydrogen — 2NH 3 = N 2 + 3H 2 , and the two gases pass out of the heated tube as separated gases, and do not recombine again.* But when steam is strongly heated it is dissociated into oxygen * Nitrogen and hydrogen can be caused to unite under suitable conditions. See Ammonia. 85 86 Introductory Outlines nnd hydrogen, find as these separated gases pass away from the heated region they reunite, forming molecules of water vapour. Such a reversible reaction may be thus expressed — 2 H 2 0 2 H 2 + 0 2 . Again, when the gases ammonia and hydrochloric acid are brought together at the ordinary temperature, they unite to form solid ammonium chloride, and when ammonium chloride is heated it dissociates into its two generators,* hence we have the expression — NH 3 + HC 1 ^NH 4 C 1 . The corresponding compound containing phosphorus in the place of nitrogen, dissociates at a temperature as low as -20°, hence when phosphoretted hydrogen and hydrochloric acid are mixed at ordinary temperatures no combination takes place, the separate molecules are in the same relation to one another as those of ammonia and hydrochloric acid at a high temperature. When, however, the mixture of gases is cooled below - 20, union takes place and crystals of phosphonium chloride are formed, which at once begin to dissociate into the original gases as the temperature again rises. The change, as before, may be represented as a reversible one — PH 3 + HC 1 ^ PH 4 C 1 . In such cases of dissociation as that of calcium carbonate, where one of the products is gaseous and the other solid, no difficulty exists in separating the simpler compounds that result from the decomposition ; but where the products are entirely gaseous, special methods have to be adopted to withdraw the one from the other, while they still exist as separate molecules, and before they reunite again. One such method, which is well adapted for the quali- tative illustration of dissociation, is based on the law of gaseous diffusion. If when ammonium chloride is heated it is dissociated into ammonia, NH 3 , and hydrochloric acid, HC 1 , these two gases, having the relative densities of 8.5 and 18.25, will diffuse through a porous medium at very different rates. According to the law of diffusion, these rates will be inversely as the square roots of the densities of the gases ; if therefore the conditions are so arranged *, Baker has shown (May 1894) that when absolutely dry, these gases do not combine ; and also, that when aqueous vapour is entirely absent, ammonium chloride does not undergo this dissociation. Dissociation 87 that the heating of the ammonium chloride takes place in the neighbourhood of a porous diaphragm, more of the light ammonia gas will diffuse through in a given time, than of the heavier hydro- chloric acid, so that a partial separation of these gases will be effected. Fig. 1 1 shows a convenient arrangement for carrying out the experiment. A fragment of ammonium chloride is heated in a short glass tube, through which passes the stem of an ordinary clay tobacco pipe. As the dissociation takes place, both of the gaseous products begin to diffuse into the interior of the porous clay pipe, but owing to their greater rate of diffusion, a larger number of am- monia molecules will pass in, than of hydrochloric acid, in the same time ; consequently, when the gases pass away from the heated region and once more recombine, there will be a surplus of am- monia molecules within the porous pipe, and for the same reason an excess of hydrochloric acid molecules outside. If the gaseous contents of the porous tube be driven out by means of a stream of air from an ordinary bellows, the presence of the free ammonia may be recognised by allowing the air to impinge upon a piece of paper, coloured yellow with turmeric, which is instantly turned brown by ammonia. The excess of hydrochloric acid within the glass tube may also be proved, by placing a piece of blue litmus paper in the tube before heating the compound, and it will be reddened by the free hydrochloric acid. In all cases of dissociation, we may imagine two opposing forces in operation, one being the external force supplying the energy which tends to bring about the disruption of the molecules, and the other being the force of the chemical affinity existing between the disunited portions of the molecule, which tends to bring about their reunion. When these forces are equally balanced, the same number of molecules are dissociated as are recombined in a given 88 Introductory Outlines unit of time, and the system is said to be in a state of equilibrium, If by any means the balance between the two opposing forces is disturbed, by augmenting or lessening either one or the other of them, the equilibrium of the system will also be disturbed and a new condition of equilibrium will be set up, in which again an equal number of molecules undergo dissociation and combination in a given time, but in which the ratio of the number of united and dis- united molecules is different from that which obtained under the former condition of equilibrium. The relation between these two forces may be most readily disturbed, by either a change of tempe- rature or pressure. Thus, in the case of nitrogen peroxide, N 2 0 4 , when this gas is at a temperature of 26.7°, 20 per cent, of it is dissociated into molecules having the composition N 0 2 ; and so long as this temperature is maintained this ratio of the weight of the dissociated molecules to the total weight of the system (known as the fraction of dissociation) still subsists. When the temperature of the gas is raised to 60.2°, the state of equilibrium existing at the lower temperature is disturbed, and the system gradually assumes a new condition of equilibrium, where once more the actual number of molecules undergoing dissociation and recombination in a given unit of time is the same, but where the percentage of dissociated molecules in the gaseous mixture is now 52.04. It might at first be supposed that when such a gas is heated, and a temperature is reached at which the molecules are dissociated, that they would all dissociate, and that the process once begun would rapidly proceed until the decomposition was complete ; instead of which, we find a definite fraction of dissociation corresponding to a particular temperature. This may be explained on the basis of the kinetic molecular theory. Let us imagine the gas nitrogen per- oxide to be at a temperature below that at which dissociation begins, when all the molecules will have the composition N 2 0 4 . The molecules of the gas are in a state of rapid movement, and the rapidity of their movement is increased by rise of temperature. But the molecules in a given volume of the gas do not all move at the same velocity, and therefore they have not all the same temperature. On account of the infinite complications in their movements caused by their impacts against one another, some will be moving at a speed considerably greater than that of the average, and will have a temperature proportionally higher, while others again will have a velocity and a temperature below the average. Dissociation 89 The observed temperature of the gas, therefore, is not that of the molecules having the highest or the lowest velocity and tempera- ture, but is the average or mean temperature between, possibly, a very wide range. On the application of heat to the gas, the observed or mean temperature rises, but the velocity of some of the molecules, and consequently their temperature, may have been thereby raised to the point at which dissociation takes place, and they consequently separate into the simpler molecules. Let us suppose that the observed temperature of the nitrogen peroxide is 26.7°, and that it is maintained at this point. Although this temperature may be below the dissociation temperature of the molecules, it must be remembered that it only represents the mean temperature, and that while some of the molecules have a lower, some also have a higher temperature. As already mentioned, at the temperature of 26.7°, 20 per cent, of the molecules are dissociated ; that is to say, at any given instant one-fifth of the total number of molecules reach a velocity which causes them to break down into the simpler NO a molecules, which themselves then take up independent movements. If, in the process of their movements, two of these disunited mole- cules come into contact with each other at a moment when their velocities are lower than that at which they dissociated, they at once reunite, so that at the same instant some are uniting and others are dissociating, and, the two processes going on equally, the percentage of disunited molecules at any moment is the same, although the actual molecules which are dissociated at one point of time may not be the identical ones that are in this state at another time. Let us now suppose the gas to be heated until the registered (/.«., the mean) temperature reaches 60.2°, and that it be maintained at this point. At this higher temperature a much larger proportion of the molecules will acquire a velocity at which they are unable to hold together, namely, 52.04 per cent. ; but the remainder, amounting to nearly one-half, still are at a temperature below that at which dissociation takes place. Under these altered conditions a greater number of disunions and reunions takes place during a given interval of time, but the numbers are equal, and therefore the equilibrium exists. If once more the gas be further heated, until the indicated temperature is 140°, then it is found that the whole of the N 2 0 4 molecules have dissociated into NO a molecules ; that is to say, when the mean temperature has reached 140°, then even those molecules that are moving with the slowest go Introductory Outlines speed, have reached the temperature of dissociation. It will be evident that the rate at which the fraction of dissociation in- creases, as the temperature of a gas is gradually raised, will be greatest when the mean temperature approaches the real dissocia- tion temperature of the gas, for the temperature of the greater number of the molecules will be coincident with, or very closely approximating to, that point. The vapour density of nitrogen peroxide, if it could be ascertained when all the gaseous molecules had the composition N 2 0 4 , would be 46 ; while that of the gas, when entirely dissociated into N0 2 molecules, is 23. At temperatures between these extremes, the gas, consisting of mixtures of both molecules, will have a density lying between these figures, thus at 27.6° and 60.2° the density is 38.3 and 30.1 (see Nitrogen Peroxide, and also Phosphorus Penta- chloride). The effect of increased pressure upon a gas being to diminish the mean free path of the molecules, and thereby increase the number of molecules in a given space, the number of impacts between the molecules in a given time will be increased. If, therefore, while the nitrogen peroxide is maintained at a constant temperature, say 62.2°, the pressure be increased, the dissociated molecules, having shorter distances to travel, and making more frequent impacts in a given time, will unite more quickly than others are being disunited, and a fresh condition of equilibrium will be established for any particular pressure. The case of phosphonium chloride already mentioned, may be referred to as an illustration. This compound is completely dis- sociated into molecules of phosphoretted hydrogen, PH 3 , and hydrochloric acid, below a temperature of o°. If, while at this temperature, it be subjected to pressure, the dissociated molecules are caused to unite, and at a pressure of thirteen atmospheres the union is complete, the whole of the disunited molecules having combined to form molecules of phosphonium chloride, PH 4 C 1 . CHAPTER XI ELECTROLYSIS If a strip of pure zinc, and a strip of platinum, be together dipped into a vessel containing dilute sulphuric acid, neither metal is affected by the acid, so long as the metals do not touch each other. If the ends of the strips outside the liquid be joined by means of a metal wire, the zinc gradually dissolves in the acid, and bubbles of hydrogen are disengaged from the liquid in contact with the surface of the platinum plate (which itself is otherwise unaffected by the acid), and at the same time an electric current passes through the wire. So long as the chemical action of the sulphuric acid upon the zinc proceeds, so long will the electric current con- tinue to pass ; in other words, chemical energy will be transformed into electrical energy. If the wire be severed, the electric current can no longer pass, and the chemical action at once stops. Such an arrangement constitutes a galvanic or voltaic element, or cell, and a series of such cells forms a galvanic battery. The zinc plate, or the end of a wire that may be connected to it, is termed the negative pole of the battery, while the end of a wire attached to the platinum plate is the positive pole. Other arrange- ments can be employed for generating a galvanic current, but in all cases the electrical energy is derived from chemical action. If the two poles of a battery are connected together by placing them both in contact with various different substances, it is seen that in some cases the electric current passes, and in others not. For instance, if the poles are joined by placing them both in contact with a bar of sulphur, no current passes, whereas when connected by a rod of graphite the current freely passes. Substances which behave in this respect like the sulphur, are said to be non-con- ductors of electricity, while those that allow the current to pass, are distinguished as conductors. Substances capable of conducting electricity are of two kinds, namely, those which are merely heated, and those which undergo a chemical change, in consequence. All QI 92 Introductory Outlines the metals, and a few of the non-metals, belong to the first of these classes ; while the second includes a large number of compound substances, which are either in the liquid state, or in solution in some solvent. Thus, if the poles of a battery are immersed in pure water, practically no current passes, because this liquid is a non-con- ductor ; but if a quantity of hydrochloric acid (HC1) be dissolved in the water, the solution at once becomes a conductor, and it is seen that gas is disengaged from the liquid upon the surface of each wire. Upon examination it is found that the gas evolved at the negative pole is hydrogen, while that from the positive pole is chlorine : the hydrochloric acid, therefore, is separated into its elements by the passage of an electric current through its aqueous solution. Such a process of decomposition is termed electrolysis ; and the conducting liquid is known as an electrolyte. The poles that are introduced into the electrolyte are called electrodes , the negative electrode being sometimes termed the cathode , and the positive electrode the anode. In a great number of instances, the electrolytic decomposition is accompanied by certain secondary reactions, caused by the action of the primary products of the decomposition upon either the electrolyte or the solvent ; for example, when a solution of sodium chloride (NaCl) is electrolysed, the primary products are sodium and chlorine, the latter appearing at the anode and the sodium making its appearance at the cathode. The metal sodium, however, in contact with the water in the neighbourhood of the cathode, at once exerts chemical action upon the liquid, with the liberation ot its equivalent of hydrogen, according to the equation — Na + H 2 0 = NaHO + H. In the same way, when an aqueous solution of copper sulphate (CuS0 4 ) is submitted to electrolysis, the primary products are copper, Cu, and the group S0 4 . The copper is liberated at the cathode, and is deposited as a metallic film upon the electrode* * This is the essence of the process of electro-plating. The metal to be de- posited, whether it be gold, silver, or nickel, &c. , in the form of a suitable salt (usually a double cyanide) in aqueous solution, forms the electrolyte. The object to be plated is made the cathode, that is, it is suspended in the liquid and is connected to the negative electrode of a suitable battery. The anode consists of a strip of the metal to be deposited. Thus in silver plating, a strip of silver is employed, and in this way the acidic radical that is liberated at the anode dissolves the metal, and thereby prevents the weakening of the solution, which would otherwise result from the gradual deposition of silver upon the cathode. Electrolysis 93 The group consisting of S 0 4 passes to the anode, where it under- goes decomposition in the presence of the water, whereby ulti- mately oxygen is evolved and sulphuric acid produced — S 0 4 + H 2 0 = h 2 so 4 + O. The primary products of electrolysis are termed the ions. Those ions that appear at the anode (positive electrode) are those which are negatively electrified, or which convey negative electricity ,* such as the elements fluorine, chlorine, bromine, iodine, and a number of acidic groups, or radicals, such as the S 0 4 group already mentioned. Inasmuch as the negative ions appear at the anode, they are sometimes spoken of as anions. Those ions, such as hydrogen and the metals, which travel to the cathode (negative electrode) are those that are positively electri- fied, or in other words, which convey positive electricity : positive ions, therefore, are distinguished as cathions. Faraday’S Law. — When the same quantity of electricity is passed through different electrolytes, the ratio between the quan- tities of the liberated products of the electrolysis, is the same as that between their chemical equivalents. Thus, if the two electrolytes, hydrochloric acid and dilute sul- phuric acid, be introduced into the same electric circuit, hydrogen and chlorine are evolved in the one case, and hydrogen and oxygen in the other. If the gases be all collected in separate measuring vessels, it will be seen (1) that the hydrogen and chlorine evolved from the hydrochloric acid are equal in volume ; (2) that the volume of hydrogen collected from the other electrolyte is the same, while that of the oxygen is equal to only one-half this amount. Knowing the relative weights of equal volumes of these three gases to be hydrogen, oxygen, chlorine, as 1, 16, 35.5, we see that they must have been liberated in the proportions by weight of — Hydrogen = 1 Oxygen = 8 Chlorine = 35.5. Similarly, if the same quantity of electricity be passed through aqueous solutions of hydrochloric acid (HC 1 ), silver nitrate (AgN O s ), copper sulphate (CuS 0 4 ), and gold chloride (AuC 1 3 ), by the time that 1 gramme of hydrogen has been liberated from the hydro- chloric acid, there will be deposited upon the cathodes of the other electrolytic cells, 108 grammes of silver, 31.7 grammes of copper, and 65.6 grammes of gold. These numbers, which are the electro- 94 Introductory Outlines chemical equivalents, are identical with the chemical equivalents of those elements, the chemical equivalent of an element being its atomic weight divided by its valency. H. O. Cl. Ag. Cu. Au. Atomic weights . . i 16 35.5 108 63.5 197 Valency ....12 1 1 2 3 Regarding the quantity of electricity required to liberate 1 gramme of hydrogen as the unit, we may say that 16 grammes of oxygen require 2 units of electricity for its liberation, 108 grammes of silver 1 unit, 63.5 grammes of copper 2 units, and 197 grammes of gold 3 units ; or, in other words, the number of units of electricity required to liberate a gramme-atom, is identical with the number representing the valency of that atom in the particular electrolyte employed. Some metals, such as copper, mercury, tin, &c., are capable of functioning with different degrees of valency. Thus copper is divalent in copper sulphate and in cupric chloride, but mono- valent in cuprous chloride. If, therefore, 1 unit of electricity be passed through aqueous solutions of each of these copper chlorides, in the case of cupric chloride 6 -^ = 3 l -7 grammes of copper will be deposited, while in the cuprous chloride = 63.5 grammes are formed. The modem theory now generally held, to explain the pheno- mena of electrolysis, is known as the theory of electrolytic dissocia- tion. The passage of electricity through conductors of the two classes above mentioned, that is, through conductors such as metals, and those which are electrolytes, may be compared with the two ways by which heat is transmitted, namely, by conduction and convection. When a bar of metal is heated at one end, the heat travels along the bar, the metal remaining stationary ; but when water is contained in a tube which is heated at its lower end, the heated particles of water travel along the tube, conveying the heat to the other extremity. In a similar manner, when electricity passes through a metallic conductor, the electricity travels through, or along, the metal, which itself does not move ; but when it is passed through an electrolyte, it is conveyed, or transported, through the liquid by the moving ions. One set of ions charged with negative Electrolysis 95 electricity travels towards the anode, while another set conveying positive electricity moves towards the cathode. In the earlier stages of the development of the present theory, it was supposed that the electrolyte was only separated into its ions as the electric current was passed into it, that the electricity was the prime cause of the dissociation of the electrolyte, hence the expression electrolytic decomposition , still commonly used. If this were in truth the case, it ought to be made manifest by the fact, that the electric current would have to do work in effecting such decompositions ; but exact experiment goes to show that electricity is conducted with equal freedom by electrolytes as by metals. The theory proposed by Arrhenius (1887) is that a certain proportion of the molecules of the electrolyte are in a dissociated condition at all times. When, for example, sodium chloride is dissolved in water, some of the molecules, owing to their collisions, become separated into the ions, sodium and chlorine, much in the same way as a certain proportion of the molecules of a gaseous compound may be dissociated, and that these convey the electricity as soon as the electrodes are introduced into the solution. From a number of other considerations, it is now believed that in such a saline solution the greater proportion of the compound is in the dissociated or disunited condition ; the proportion depend- ing largely upon the state of dilution. The more dilute the solution, and the more complete is the dissociation. At first it might appear contrary to established ideas, that in the case, for example, of such a compound as sodium chloride, the sodium and chlorine in the free state should be capable of existence in the same liquid ; a liquid, moreover, upon which one of the elements, namely sodium, is under ordinary circumstances capable of exerting a chemical action. According to the electrolytic dissociation theory, however, the disunited constituents of the sodium chloride exist as separate atoms, having enormous electrical charges, the sodium with positive and the chlorine with negative electricity. Whenever the chlorine atoms lose their electrical charge, they unite together, forming the chlorine molecule Cl 2 , which then possesses the properties which are usually associated with this element ; and in the same way when the positively electrified sodium atoms give up their charge they likewise unite, forming sodium molecules, which are endowed with the ordinary properties belonging to that element. When the electrodes of a galvanic battery are placed into such a solution of sodium chloride, the negatively charged chlorine atoms travel to 9 6 Introductory Outlines the anode, and there discharge their electricity, and in consequence, the chlorine atoms at that point unite, and molecules of ordinary chlorine escape as gas from the liquid. The sodium atoms with their positive charge travel to the cathode, where in like manner they are discharged, and at once unite to form sodium molecules, having the ordinary properties of sodium, and consequently at that point this element exerts its chemical action upon the water and liberates hydrogen. This theory, that electrolytes in dilute solution are dissociated into their ions, is in harmony with, and derives support from, the laws which regulate the influence of substances in solution, upon osmotic pressure (page 136), upon the lowering of the vapour-pres- sure (page 1 1 8), and upon the lowering of the solidifying point of the solvent (page 121). Dilute solutions of electrolytes (strong acids, bases, and salts) exhibit deviations from these laws, much in the same way that gases which undergo dissociation, depart from the ordinary gaseous laws. It is found that in the case of dilute solutions of electrolytes, the osmotic pressure, the lowering of the vapour pressure, and the lowering of the freezing-point of the solvent, instead of being proportional to the number of molecules of the dissolved substance, are proportional to the number of dis- sociated ions j each ion behaving as a separate molecule. CHAPTER XII CLASSIFICATION OF THE ELEMENTS It has already been mentioned (page 7), that the elements may be classified under the two subdivisions, metals and non-metals. Further classifications have from time to time been in use, based upon other properties, such, for example, as the valency of the elements. Classified according to their valency, the elements fall into six subdivisions, consisting of mono-, di-, tri-, tetra-, penta-, and hexa- valent elements. This system of classification has now largely fallen into disuse, owing partly to the difficulties arising out of the variability of valency so often exhibited, but more especially to the more recent development of another system, known as the natural classification of the elements , or the periodic system , which practi- cally absorbs and includes the older method. Certain remarkable numerical relations have long been observed to exist among the atomic weights of elements that closely re- semble one another in their chemical habits. In such groups or families it is frequently seen that the atomic weight of one member, is approximately the arithmetic mean of the atomic weights of those immediately before and after it, when they are arranged in order of their atomic weights. This will be seen from the following examples : — Li. Na. 7 23 K. Rb. 39 85 P. As. 31 75 s. Se. 32 79 K. 39 7 + 39 = 23 2 J Cs. 39±T33 = 85.8 133 Sb. 120 31 + 120 2 - 75-5 Te. 32 + 125 - 7 S 125 97 G 98 Introductory Outlines If the elements in these various families are so arranged, as to bring out the differences between their atomic weights, the striking fact will be observed that the increase in the atomic weights in each group takes place by practically the same increment. In the following table the elements belonging to the same group are placed in vertical columns, the differences between the various atomic weights being placed between them : — F = 19 N = 14 ° II o\ Na = 23 Mg = 24 Difference . 16.5 Diff. . 17 Diff. . 16 Diff. . 16 Diff. . 16 Cl = 35-5 P = 3 1 S = 32 K = 39 Ca = 40 Difference . 44.5 Diff. . 44 Diff. . 47 Diff. . 46.2 Diff. . 47.3 03 i -i II 8 ? As = 75 Se = 79 Rb = 85.2 Sr = 87.3 Difference . 47 Diff. . 45 Diff. . 47 Diff. . 47-8 Diff. . 49-7 I = 127 Sb = 120 Te = 125 Cs = 133 Ba = 137 It will be seen that in each group the difference between the first and second number is about 16, while between all the others the increase in weight takes place by a number which approximates to 16 x 3. These numerical relations between the atomic weights of elements of the same family, and between the various groups, is obviously not a chance one, and chemists were led by it to believe that the properties of the elements were in some way related to their atomic weights. Newlands (1864) was the first to point out, that if the elements are tabulated in the order of increasing atomic weights, the properties belonging to each of the first seven elements reap- peared in the second seven, and he applied to this relation the name of the law of octaves. A more elaborated and systematic representation of Newlands’ law of octaves was afterwards deve- loped by Mendelejeff (1869), and which is now generally known as Mendelejeff s periodic law. If the fourteen elements with lowest atomic weights, after hydrogen, be arranged in order of increasing atomic weights in two horizontal rows of seven, some of these relations will be recognised — Li =7 Be =9 B = 1 1 C — 12 N=i 4 0 =i 6 F =19 Na = 23 Mg = 24 Al = 27 Si = 28 P =31 S =32 C 1 = 35 S- The Periodic Classification 99 In traversing the upper row from lithium to fluorine, we meet with certain characteristic properties belonging to each member, and also a certain gradation in those properties that are common. Coming to the second row, or octave, many of the characteristic properties of the members of the first row again appear, and the same regular modulation is met with in passing along the series : thus lithium resembles sodium, carbon corresponds to silicon, fluorine to chlorine, and so on. These resemblances are seen both in the physical as well as the chemical properties of the elements, thus lithium and sodium are both soft white metals, and are strongly electro-positive. Fluorine and chlorine are both pungent corrosive gases, and are intensely electro-negative. Tak- ing their power of combining with chlorine and with hydrogen as indicative of their valency, we see that the change in this respect, as the two series are traversed, is the same in each, thus — Li Cl BeCl 2 BC1 3 CC1 4 CH 4 NH 3 OH 2 FH NaCl MgCl 2 (A1C1 3 ) 2 SiCl 4 SiH 4 PH 3 SH 2 C1H The gradation in properties exhibited by the elements in a series, is also seen in their power of combining with oxygen, which will be more clearly brought out if the formulas of the compounds be so written as to indicate the relative proportions of oxygen with which two atoms of each element unites, thus — Na 2 0*(Mg 2 0 2 ) A1 2 0 3 (Si 2 0 4 ) P 2 0 6 (S 2 O e ) Ct 2 O r + Regarding, then, the seven elements of the first row as a period, we find that the various properties exhibited by the several members are met with again in those of the second period. Not only do the properties of the elements themselves reappear, but also those possessed by the various compounds they form ; thus lithium chloride (LiCl) and sodium chloride (NaCl) strongly re- semble one another. The oxides of beryllium and magnesium (BeO and MgO) have similar properties. The compounds of fluo- rine and chlorine with hydrogen (HF and HC1) closely resemble each other, and so on. This periodic reappearance of similar properties, exhibited by the MgO Si0 2 so 3 * See footnote on page 487. t Perchloric oxide is not known in the free state. I oo Introductory Outlines elements and their compounds as the atomic weights of the former gradually increase, is thus stated by Mendelejeff in his law of periodicity. The properties of the elements , as well as the proper- ties of their compounds, form a periodic function of the atomic weights of the eleme 7 its. When the tabulation of the elements according to this system is continued (after the completion of the second period with chlorine), it will be seen, that beginning with potassium, seventeen elements have to be arranged before we meet with the reappearance of those properties that belong to the first ; that is to say, there are two sets of seven each, and three elements over, which in the following table are placed within brackets : — K. Ca. Sc. Ti. V. Cr. Mn. (Fe. Co. Ni.) 39 40 44 48 5 1 52 55 (5 6 59 59) Cu. Zn. Ga. Ge. As. Se. Br. 63-5 65 70 72 75 79 80 This constitutes what is known as a long period, in contradis- tinction to the two first, which are distinguished as short periods, ki certain respects, however, the last seven elements in this long period exhibit resemblances to the first seven ; that is to say, the properties displayed by the members of the first period, which is known as the typical period, reappear twice over in the long period. The three elements within the brackets are termed by Mendelejeff transitional elements . Continuing the arrangement from bromine, another long period occurs, again containing three transitional elements : — Rb. Sr. Y. Zr. Nb. Mo. — (Ru. Rh. 00 yi to 87-3 89.6 90.4 93-7 96 ? ( io 3 - 5 I04 Ag. Cd. In. Sn. Sb. Te. I. I08 1 1 2 113 Il 8 120 125 127 It will be seen that a gap is left where the seventh member of the first part of this period should be, an element which would correspond, in this period, with manganese in the period above. This element is at present unknown. The remaining elements belong to three other long periods, in which, however, the number of gaps is very considerable, thus — The Periodic Classification 101 Cs. ‘ Ba. La. Ce. (- -) 133 137 13S.5 MI Yb. 173 Ta. W. — (Os. Ir. Pt.) 182 184 (191 I92.5 I95) Au. Hg. TL Pb. Bi. 197 200 203.7 20 7 207.5 Th. 232 239-8 Ur. <~ -) Those elements that fall in the first seven places of the long periods, are termed the even series , while the last seven are dis- tinguished as the odd series j arranging them, therefore, in such a manner as to bring the odd and even series into columns, we get the table on page 102. In this manner the elements are arranged in eight groups, the eighth containing the transitional elements that come between the even and odd series of the long periods. In each of the remaining seven groups, the elements belonging to the even series of their lespective long periods, are placed to the left, while those belonging to the odd series are arranged on the right hand side of each vertical column. In this way the groups are divided into the subdivisions, A and B, in which the resemblance between the members is most pronounced. Thus in Group II., although there are certain properties common to all the members, there is a much closer similarity existing between the elements calcium, strontium, and barium than between zinc and calcium, or cadmium and barium.* The elements in the two short periods, have been placed in that subdivision or family, with the members of which they exhibit the closest resemblance. Thus, in Group I. lithium and sodium are more allied to potassium, rubidium, and caesium, than to copper, silver, and gold; while in Group VII. * This, however, is by no means uniformly the case ; thus the element copper (Group I.) in many of its chemical attributes is much more closely allied to mercury (Group II.) than to silver; and silver again more strongly resembles thallium (Group III.) than either copper or gold, with which it is associated in this system of classification. 102 Introductory Outlmes Group VIII. 1 | Fe, Co, Ni '■a PL, rG Pi G pi 1 1 1 £ 6 1 1 1 o Cf 0^ Group VII. (I Uh 1 U pq G 1 1 1 r i 1 1 offi- c£* Group VI. l a [< o . i cn 0 1 ° l< 1 cl i l v As Nb Sb 1 1 PQ J r i >Q OK Pi Pi H (* .1 V 1 w H3 v o .O 'c o 'G O 'C • 4) CL to G JO X j Oxygen compounds 1 Hydrogen compounds 1 The Periodic Classification 103 fluorine and chlorine are placed in the same family with bromine and iodine, with which they exhibit a close similarity. In the eighth group, containing the transitional elements, the families consist of the horizontal and not the vertical rows ; that is to say, the closest resemblance is between the three transitional elements in each series, elements whose atomic weights, instead of exhibiting a regular increase, as in the other families, have almost the same value, such as Fe = 56 ; Co = 59 ; Ni = 59. A glance at the table shows that in the last three long periods there are a large number of gaps. It is possible that these gaps may represent elements which yet await discovery. This supposi- tion gains considerable support from the fact, that at the time Mendelejeff first formulated the periodic law, there were three such gaps in the first long period, which have since been filled up by the subsequent discovery of three new elements ; these will be referred to later. It is noteworthy, however, that all the elements belong- ing to the last three periods, together make a total which is almost exactly the number required for a single complete long period, including three transitional elements ; and it is quite pos- sible that future investigations may necessitate an alteration in the accepted atomic weights of some of these elements, and con- sequently a change in their positions in the system. The periodic recurrence of some of the chemical properties, is indicated in the lowest horizontal column, where the general formulae of the oxygen compounds, and the hydrides, are given ; R standing for one atom of any element in the group. As explained on page 99, these formulae are so written as to show the relative amount of oxygen to two atoms of element, in order to establish the true relation between the different groups. For example, the oxides of the elements of Group I. contain two atoms of the element to one of oxygen, as Li 2 0 ; but those of the second group only con- tain one atom of the element, as CaO : hence the general formula is doubled, R 2 0 2 . It will be seen, therefore, that the proportion of oxygen relative to two atoms of the element regularly increases from the first group to the eighth. The oxides of the members of the first group are strongly basic in character, and in general this basic nature gradually diminishes as we traverse the series, giving place to acidic characteristics, which are strongly marked in the seventh group. The periodic reappearance of the physical properties of the elements is seen in such points as their electrical characters, their I0 4 Introductory Outlines malleability, ductility, melting-points, &c., all of which are in harmony with the periodic law ; but in none is it more strikingly seen than in their atomic volumes in the solid state. The atomic volumes of the elements, are the relative volumes occupied by quantities proportional to their atomic weights, or by gramme- atoms ; and they are obtained by dividing the atomic weights of the elements by their specific gravities. In the case of gases, as has been already explained on page 39, the specific gravity is the density referred to hydrogen as the unit : the atomic volume, therefore, of such a gas as oxygen is — 16 = atomic weight _ 16 = density The specific gravities of solids (and also liquids) are referred to water as the unit, and as 1 cubic centimetre of water weighs 1 gramme, the specific gravity of a solid or liquid, expresses the weight in grammes of 1 cubic centimetre of the substance. Dividing the atomic weight, expressed in grammes, by the weight in grammes of 1 cubic centimetre (i.e., the specific gravity), the atomic volume will be represented in cubic centimetres. It must be remembered that the atomic volumes do not express the relative volumes that are actually occupied by the atoms, they represent in reality the relative volume of the atoms plus the unknown volumes of the spaces that separate them. The following table gives the specific gravities, and the calculated atomic volumes, of the first and the middle elements of the two short and two long periods. 1 Specific Gravity. Atomic Weights. Atomic Volumes. 1st Period < 2nd Period < 3rd Period j 4th Period j ( Lithium [ Carbon ( Sodium 1 Silicon ! [ Potassium j 1 Iron (Cobalt-Nickel) Rubidium , (Ruthenium-Rhodium) Palladium . Caesium 0.59 3 ° 0.97 2.4 0.865 7-79 1.52 11.4 1.88 7 12 2 3 28.3 39 56 ' 85 106.2 *33 r ! 23.6 n.8 1 45 i 7.2 i 56.0 { 9 70 1 From the figures in the last column it will be seen, that beginning The Periodic Classification 105 with lithium, 1 1.9, the atomic volume falls as the middle element of the period, namely carbon, is reached ; after which it again rises and reaches a maximum with the first member of the second period, namely sodium. In this period the same gradual fall in atomic volume is again noticed until the middle element (silicon) is reached, when the value of this function of the elements once more rises, and a second maximum is attained with the first member (potassium) of the third period. The two next are long periods, and the atomic volumes steadily decrease until the middle three (transi- tional) elements, after which they gradually increase again to a maximum in rubidium, the starting-point of the fourth period. In the fourth period the same thing once more occurs, the minimum atomic volumes being those of the middle or transition elements, after which a maximum is again reached in caesium. This periodicity of the atomic volumes may be graphically represented by a curve, where the ordinates represent atomic volumes, and the abscissas atomic weights. This curve, which was first constructed by Lothar Meyer, is known as Lothar Meyer’s curve (page 106), and a comparison of it with Mendelejeffs table is most instructive. The divisions indicated by the Roman numerals correspond to the different periods : Groups I. and 1 1. being the two short periods, III. and IV. the two complete long periods, while V., VI., and VII. correspond to the fragmentary portions of the last three periods. The transitional elements of periods III., IV., and VI. are all to be found at the minima of the large hollows ; separating the even series (situated on the descending portion of the curve), from the odd series which lie on the ascending slope. The elements belong- ing to the different groups in Mendelejeffs table, are seen to occupy the same relative positions upon the different portions of this curve. Thus in Group I. the elements Li, Na, K, Rb, Cs, are all found upon the maxima of the curve, and Cu, A g, and Au at those points at the minima where the electro-negative properties reappear. The halogen elements (chlorine, bromine, iodine) are seen in similar positions upon the ascending, and the alkaline earths (beryllium, magnesium, calcium, strontium, barium) on the descending portions. When the periodic law was first formulated by Mendelejeff (1869), there were a number of instances in which the system did not harmonise with the then accepted atomic weights of the elements. The discoverer boldly asserted that the atomic weights, io6 Introductory Outlines Atomic foygrhts The Periodic Classification 107 and not the system, were at fault, and in every such case the care- ful reinvestigation of the atomic weights by numerous chemists, has proved the correctness of the assertion. One or two instances may be quoted. The element indium had assigned to it the atomic weight 76. Its combining proportion is 38, and being regarded as a divalent element, its oxide was believed to have the formula InO. Having an atomic weight = 76, indium would occupy a place between As = 75 and Se = 79 ; but in the system (see table on page 102), there is no room for an element with such an atomic weight ; and, moreover, if indium be a divalent element having this atomic weight, it should come between Zn = 65 and Sr = 87 in Group II., where again there is no room. Mendelejeff made the assumption that the oxide of indium had the formula ln 2 0 3 , believing the element to be an analogue of aluminium (Group III.). If this be the true composition of the oxide, the atomic weight of the element would be 38 x 3 = 114, and indium would then take its place in Group III., between the elements cadmium =112 and Sn = 118, in the odd series of the second long period. Bunsen afterwards determined the specific heat of indium by means of his ice calorimeter, and found it to be 0.057 : — Mean atomic heat 6.4 Specific heat . . 0.057 1 12.3 = atomic weight (see page 45). Hence 114 and not 76 is the accepted (approximate) atomic weight of indium. Again, the element beryllium (formerly known as glucinum) has a combining proportion of 4.6. Its chloride was believed to have the composition BeCl 3 , and its oxide to be a sesquioxide having the formula Be 2 0 3 . The atomic weight assigned to the element, therefore, was 13.8. With this atomic weight beryllium would take its place between carbon = 12 and nitrogen = 14; but according to the periodic classification there is no room for such an element, and moreover, in such a position it would be among elements with which it has no properties in common. On the supposition that the oxide of beryllium has the formula BeO, that is, that the element is divalent, its atomic weight would have to be lowered from 13.8 to 9.1 in order to maintain the same ratio between the weights of metal and oxygen in the compound. On this assumption, beryllium would fall into the second place in the first series, between lithium = 7 and boron = 11. and in the same group as magnesium and zinc. 1Q 8 Introductory Outlines When the specific heat of beryllium was determined, it gave the value 0.45, and this number divided into the atomic heat constant, 6-4 j gave 14 as the atomic weight. In spite of this evidence in favour of the higher value as the atomic weight of beryllium, Mendelejeff still regarded the lower number as correct, and it was suggested that possibly beryllium, like carbon and boron (elements also of very low atomic weight), had an abnormally low specific heat at ordinary temperatures. This was found to be the case (see page 46), and at 500° the specific heat of beryllium was found to be 0.6206. This divided into 6.4 gives the value 10 as the atomic weight, which indicates that 9.1 and not 13.8 is in reality the atomic weight of beryllium. Not only has the periodic law been of service in bringing about the correction of a number of doubtful atomic weights, but by means of it, its originator was enabled to predict with considerable certainty the existence of hitherto undiscovered elements, and even to predicate many of the properties of these elements. As already mentioned, at the time when the periodic law was first formulated, there were three gaps in the system in the first long period, namely, No. 3 in the even series (now occupied by scandium), and Nos. 3 and 4 in the odd series (now filled by gallium and germanium). To the unknown elements which were destined to occupy these positions, Mendelejeff gave the names eka-boron , eka-aluminium , , and eka-silicon (the prefix eka being the Sanscrit numeral one), and from the known properties of the neighbouring elements of the series (horizontal rows in the table, page 102), and also of those situated nearest in the same family (vertical columns), he predicted some of the prominent properties that would pro- bably be possessed by these elements. Thus in the case of eka- aluminium, from the known properties of aluminium and indium, the neighbouring elements in the same family, and from zinc, the contiguous element in the same series (the 5th place in the series being unoccupied), Mendelejeff deduced the following properties for the unknown element that he called eka-aluminium : — Predicted Properties of Eka-Aluminium (1871). (1.) Should have an atomic weight about 69. (2.) Will have a low melting-point. (3.) Its specific gravity should be about 5.9. (4.) Will not be acted upon by the air. The Periodic Classification YO9 (5.) Will decompose water at a red heat. (6.) Will give an oxide E/ 2 0 3 , a chloride El 2 Cl 6 , and sulphate ^/ 2 (S 0 4 ) 3 . (7.) Will form a potassium alum, which will probably be more soluble and less easily crystallisable than the corresponding alumi- nium alum. (8.) The oxide should be more easily reducible to the metal than alumina. The metal will probably be more volatile than alumi- nium, and therefore its discovery by means of the spectroscope may be expected. In the year 1875 Lecoq de Boisbaudran discovered a new element in a certain specimen of zinc blende (zinc sulphide), the individuality of which he first recognised by the spectroscope, the spectrum being characterised by a brilliant violet line. This element he named gallium. The properties of this metal, as they were subsequently observed, showed that it was, in fact, the pre- dicted eka-aluminium of Mendelejeff, as will at once be seen by a comparison of the following facts. Properties of Gallium ( discovered 1875). (1.) Atomic weight = 69.9. (2.) Melting-point, 30. 15 0 . (3.) Specific gravity, 5.93. (4.) Only slightly oxidised at a red heat. (5.) Decomposes water at high temperatures. (6.) Gallium oxide, Ga 2 O a . Gallium chloride, Ga 2 Cl 6 . Gallium sulphate, Ga 2 (S 0 4 ) 3 . (7.) Forms a well-defined alum. (8.) Is easily obtained by the electrolysis of alkaline solutions. In a similar manner the properties of eka-boron and eka-silicon were predicted, and the subsequent discovery of scandmtn (Nilson. 1879), and germanium (Winkler, 1886), whose properties were found to closely accord with these hypothetical elements, formed an additional demonstration of the truth of the periodic law. No satisfactory theory has yet been offered, to explain the law of periodicitv. CHAPTER XIII GENERAL PROPERTIES OF LIQUIDS Under this head the following subjects will be considered 1. The passage of liquids into vapours or gases. 2. The passage of liquids into solids. 3. Solution. 1. The Passage of Liquids into Gases. Evaporation and Boiling. — Just as in the gaseous condition, so in the liquid state, the molecules are in a state of motion : in the liquid state, however, the mean kinetic energy of the molecules is unable to overcome the force of their mutual attraction. Some of the molecules have a smaller kinetic energy (that is, a lower temperature), and others a greater kinetic energy, than the average ; and when in the course of their movements the latter strike the surface of the liquid and break through it, they continue their movements in the space above, as gaseous molecules. If the space into which they wander be unlimited, that is, if the liquid be freely exposed to the air, these molecules escape away altogether, and consequently the liquid diminishes in quantity. This process is known as evaporation , and as the molecules which so leave the liquid are those having the highest temperature, it follows that the temperature of the liquid, which is the average temperature of the molecules, will fall. The more completely the molecules that so escape from the surface of a liquid are prevented from falling back, that is, the more rapidly they are swept away from the immediate neighbourhood of the liquid, the more quickly will this escape of molecules take place, and therefore the greater will be the fall of temperature that results from evaporation. Thus, if a quantity of liquid, say water, be exposed in a dish so that a current of air is blown across the sur- face, the rate of evaporation is increased, and the temperature con- sequently falls lower than if the water be merely placed in a still atmosphere : similarly, if the water be placed in a vacuum the rate no Evaporation t t t of evaporation is increased, because the molecules that escape from the surface of the liquid are not impeded in their motions by collisions with the molecules of air. This fall of temperature resulting from evaporation, may be readily seen by enveloping the bulb of a thermometer in a piece of thin muslin, and moistening it with water. If such a thermometer be placed by the side of a naked thermometer, it will be seen that the mercury will fall lower in the one that is moistened, and the difference will be still more marked if the instruments are placed in a draught, whereby the evaporation of the water from the muslin is accele- rated. If the space above the liquid be limited, molecules still continue to escape from the surface ; but a state of equilibrium is soon established, when as many are thrown back again by rebounding from one an- other and from the walls of the containing vessel, as leave the surface in a given time. Under these conditions the enclosed space is said to be satu- rated with the vapour of the liquid. The number of molecules which escape from the surface, depends upon the temperature, and Fig. 12. is independent of the pres- sure, for if the volume of a saturated vapour be forcibly diminished, it merely results in the condensation of a portion of the vapour ; and if expanded, a corresponding vaporisation of an additional quantity of the liquid, the pressure remaining always constant. The number of molecules that re-enter the liquid, is determined by the number and the velocity of those that exist as gaseous molecules in a unit volume. But the pressure exerted by a gas is caused by the number and velocity of the molecules in a given volume, hence the 1 1 - / ntroductory Outlines condition of equilibrium is set up, when the vapour above the liquid exerts a definite pressure, which pressure will be constant for any given temperature. The pressure exerted by a vapour under these conditions is termed the vapour tension of the liquid. The fact that the vapour given off from a liquid exerts pressure, may readily be experimentally illustrated by means of the apparatus seen in Fig. 12. Three glass tubes, A, B, and c, about one metre long, are com- pletely filled with mercury and inverted in a trough of the same liquid. The mercury will sink to the same level in each tube, the length of the mercury column representing the atmospheric pres- sure at the time. Into two of these barometer tubes, B and c, a few drops of water are introduced, when it will be found that the mercury is depressed, as indicated in B, below the level at which it previously stood. This depression of the mercury column, repre- sents the tension of the vapour of the water for the particular temperature at which the experiment is made. If tube C be sur- rounded by a wider glass tube, through which steam from a small boiler is passed, it will be noticed that as the temperature of the water in the tube rises, the mercury is more and more depressed, thus showing that the tension of the vapour increases with rise of temperature. As soon as the steam circulates freely and is escaping at the bottom of the wide tube, in other words, as soon as the temperature of the enclosed water in tube c reaches ioo°, i-e ., the temperature of the steam surrounding it, the mercury in the tube will be depressed to the level of that in the trough. The tension of the vapour within the tube, under these circumstances, is therefore equal to the atmospheric pressure. If, instead of introducing water into the barometer tube, ether were employed, and a stream of vapour from boiling ether were passed through the outer tube, it would be seen that when the ether within the tube reached the temperature of the vapour from the boiling ether, namely, 35 0 , the mercury would again be depressed to the level of that in the trough ; that is, the tension of the ether vapour would then be equal to the pressure of the atmosphere. We see, therefore, that when water is heated to its boiling-point, viz., ioo°, the tension of its vapour is equal to the atmospheric pressure; and when ether is heated to its boiling-point, viz., 35 0 , the pressure exerted by its vapour is equal to the pressure of the atmosphere. The boiling-point of a liquid may therefore be defined as the temperature at which the vapour pressure is equal to the pressure of the atmosphere. As soon as this point is passed, the kinetic Boiling-Points of Liquids 1 1 3 energy of the molecules has been so much augmented by the supply of external heat, that it is able to overcome the force of their mutual attractions, and, consequently, the molecules freely pass away from the surface of the liquid. As will be seen from the illustrations given, namely, water and ether, the temperatures at which the vapours of different liquids exert a pressure equal to that of the atmosphere are widely diffe- rent. This fact will be still more evident from the following table, giving the temperatures at which the vapour pressure of various liquids is equal to the standard atmospheric pressure : — Liquid oxygen ..... - 1S1 0 Liquid nitrous oxide ... - 92 0 Liquid sulphur dioxide Ethyl chloride + ii° Carbon disulphide .... 47 ° Water IOO° Aniline ..... 182° Mercury ...... 358 ° Since the boiling-point of a liquid is that temperature at which its vapour tension is equal to the atmospheric pressure, it will be evident that, if the latter increases or decreases, the temperature necessary to produce an equal vapour pressure must also rise or fall ; in other words, the boiling-point of a liquid is dependent upon the pressure. If a quantity of water, no warmer than the hand, be placed beneath the receiver of an air-pump, which is then quickly exhausted, the water will be seen to enter into violent ebullition. It does this, when the pressure within the receiver is reduced to the point at which it is equal to the tension of aqueous vapour at the temperature taken. For this reason water boils at a lower temperature in high altitudes than at the sea-level ; and as the vapour tension of water at various temperatures has been experimentally determined, we can, by ascertaining the boiling-point of water at any particular altitude, calculate the atmospheric pressure, and consequently the height above the sea-level. Many liquids when heated, especially in glass vessels that have been carefully cleansed, may be raised several degrees above the boiling-point without ebullition taking place. The liquid under these circumstances assumes a pulsating movement, which con- tinues for a short time, when a burst of vapour is suddenly evolved H 1 14 Introductory Outlines with violence, and the temperature at once drops to the boiling- point. The liquid then becomes quiescent, and again, as the temperature rises, the pulsating movement begins, ending once more in an explosive evolution of vapour. This successive boiling, or bumping , is sometimes sufficiently violent to cause the fracture of the vessel. In order to experimentally ascertain the boiling-point of a liquid, tiie thermometer, for this reason, is not immersed in the liquid , but is suspended in the vapour, the temperature of which remains constant throughout these irregularities in the boiling. Latent Heat cf Vaporisation.— When a liquid is heated, its temperature rises, as indicated by the thermometer, until a certain point is reached (t'.;e boiling-point of the liquid), when the con- tinued application cf heat causes no further rise of temperature. Thermometers placed in the liquid, and in the vapour, indicate the same temperature and remain constant, and all further applica- tion of heat is unapprec iated by these instruments, and disappears in changing the liquid into vapour. The heat which in this way is absorbed during the vaporisation of a liquid, is spoken of as the latent heat of vaporisations and the same amount of heat which thus disappears during the conversion of a liquid into a vapour, is again rendered sensible when the vapour passes back into the liquid state. The heat which is thus said to become latent , is in reality con- verted into kinetic energy; it is expended in imparting to the molecules the kinetic energy necessary to overcome the attractive forces operating between them while in the liquid state ; in other words, it is doing the work of overcoming cohesion (internal work), and also the external pressure on the vapour (external work). In order that a liquid may pass into a vapour it is necessary that heat be absorbed. We have seen (page iio) that a liquid under- going spontaneous evaporation becomes colder (that is, heat is absorbed by the molecules that are converted into the gaseous state), and also that the more rapidly the liquid can be made to pass into the vaporous condition, without supplying external heat, the lower will its temperature fall. Upon this fact depend a number of methods for the artificial produc- tion of low degrees of cold. For example, ether boils at 35 } but if a small quantity of ether be placed in a glass flask standing Latent Heat of Vaporisation II 5 upon a wooden block, upon which a few drops of water have been poured, and a current of air from a bellows be briskly blown through the ether (Fig. 13), the temperature of the ether will fall so rapidly that in a few moments the flask will be frozen to the block. By the rapid evaporation of liquids with lower boiling- points, the extreme degrees of cold necessary for the liquefaction of such gases as oxygen, carbon monoxide, air, &c., are obtained. Thus, liquid methyl chloride boils at - 23 0 ; by causing it to rapidly vaporise, its temperature can be reduced to - 70°. Liquid ethy- lene in the same way falls to a temperature of - 120°, and liquid oxygen by rapid evaporation gives a temperature as low as -210°. The temperature of water, in like manner, may be so lowered by its own rapid evaporation, as to cause it to freeze. We have already seen that by reducing the pressure, the boiling-point of a liquid is lowered ; if, therefore, a quantity of water be placed in a vacuum, and methods be adopted to remove the water vapour as rapidly as it is formed, the water will enter into rapid ebullition. The i6 Introductory Outlines evaporation will therefore proceed so rapidly, and consequently absorb heat so quickly, that the temperature of the boiling liquid will quickly fall to o° when it passes into the solid state. The instrument known as Carre's freezing machine depends upon this principle. The water to be frozen is placed in the glass bottle C (Fig. 14), which is in connection with a metal reservoir R, half filled with strong sulphuric acid. This in its turn is connected by b with an air-pump P, worked by the lever M, to which is also attached a connecting rod /, so that a stirrer within the reservoir is kept constantly in motion. As soon as the apparatus is exhausted to a pressure of two or three millimetres, the water begins rapidly to boil, and as the sulphuric acid absorbs the water vapour as rapidly as it is given off, the temperature quickly falls and the water freezes. Fig. 15 illustrates another method by which the same result may be obtained. A tall glass vessel is exhausted by means of an ordinary air-pump, and water is allowed slowly to enter from a stoppered funnel, upon the end of which is secured a short string. At the same time strong sulphuric acid is admitted by the second funnel, and caused to flow down a glass rod, round which is wound a spiral of asbestos thread. The acid at once absorbs the aqueous vapour from the evaporating water, the tempera- ture of which, therefore, falls below the freezing-point, and it solidifies as it flows over the string into the form of an icicle. Just as diminution in pressure lowers the boiling-point of a liquid, so increased pres- sure raises the boiling-point. If water be heated in a closed iron vessel, as in a high-pressure steam boiler, the pressure caused by its own vapour raises the boiling-point many degrees above ioo°. There is a definite temperature, how- ever, for every liquid, beyond which the liquid state is impossible, whatever may be the pressure ; that is to say, the liquid when heated beyond this fixed point passes into the gaseous state, how- ever great the pressure may be. This temperature is the critical Vap, mr~ Pressures of Solutions I 17 temperature (see page 77). If a liquid be heated in a sealed and strong glass tube, as the critical temperature is approached, the surface of the liquid gradually becomes ill-defined, and finally the tube is completely occupied by transparent vapour. On again cooling, as soon as the critical point is passed, the contents of the tube again separate into two distinct layers consisting of liquid and gas. Vapour-Pressures of Solutions.— The boiling-point of a liquid is modified by the presence in the liquid of dissolved substances. If the substance in the solution be less volatile than the liquid, the boiling-point is raised. Thus, while the boiling-point of pure water (under the normal atmospheric pressure) is ioo°, the tem- perature at which saturated aqueous solutions of salts boil, is considerably higher, thus : — Containing Grammes of Water Saturated with Salt in roo Grammes Boiling-point. of Water. Sodium chloride . . 41.2 108.4° Potassium nitrate . . . 335.1 115.9 0 Potassium carbonate 205.0 133.0 0 Calcium chloride . . 325.0 T79.5 0 The temperature of the steam of these boiling solutions, as ascertained by suspending a thermometer in the vapour, appears to be the same as that from pure water, as the thermometer in all cases indicates ioo°. In reality, however, the temperature is higher, although not so high as that of the boiling liquid. The reason that the thermometer indicates ioo° in all cases is because the water vapour continually condenses upon the bulb of the instrument, covering it with a film of pure water, which boiling off from the bulb indicates only the boiling-point of the pure liquid. By special arrangements this condensation may be pre- vented, when it has been shown (Magnus) that the temperature of the vapour, from such boiling solutions, rises as the solutions become more concentrated — that is, as the temperature of the boiling liquids rise. It has been already explained that the boil- ing-point of a liquid is that temperature at which the vapour tension is equal to the atmospheric pressure ; since, then, the presence of dissolved substances raises the boiling - point, it therefore lowers the vapour-pressure, for (in the case of aqueous solutions) when the temperature has reached ioo° the vapour- T , g Introductory Outlines pressure is still below that of the atmosphere, for the liquid does not enter into ebullition at that temperature. By measuring vapour-pressures of solutions at a constant temperature, instead of measuring the temperature at a constant pressure {i.e., the boiling-point), the following general laws have been estab- lished : — 1. The relation between the quantity of a substance in solution and the diminution of the vapour-pressure below that of the pure solvent , is the same at all temperatures. 2. The diminution of the vapour-pressure of a liquid, by a dissolved substance , is proportional to the amount of the substance in solution {provided the substance itself exerts no appreciable vapour-pressure at the temperature of the experiment). 3. The molecular lowering of vapour -pressure by chemi- cally similar substances is constant ; that is to say , solu- tions containing one molecular weight in grammes {one gramme-molecule) of such substances in equal volumes of the solvent , give rise to the same diminution of vapour - pressure. 4. The relative lowering of vapour-pressure is propor- tional to the ratio of the number of molecules of the dis- solved substance, to the total number of molecules in the solution, i.e., the sum of the number of molecules of the dissolved substance and of the solvent * Upon these considerations it becomes possible, by means of the lowering of the vapour-pressure, to find the molecular weight of a substance that is capable of being dissolved in a volatile liquid. The Passage of Liquids into Solids.— Most liquids, when cooled to some specific temperature, pass into the solid state ; the temperature at which this change takes t is termed the solidifying point. Generally speaking, the temperature at which a liquid solidifies is the same as that at which the solid again melts ; but as the solidification of a liquid is subject to disturbances from causes that do not affect the melting-point, this is not always the case. Thus, water may be cooled many degrees below o° if it be previously freed from dissolved air, and be kept perfectly still. This super- cooling of water may readily be illustrated by means of the * Except in the case of electrolytes. See page 96. Solidifying Points of Liquids 1 19 apparatus represented in Fig. 16. This consists of a thermometer whose bulb is enclosed in a larger bulb containing water, which before the bulb is sealed at a, is briskly boiled to expel all the air. When the instrument is immersed in a freezing mixture the tem- perature of the water may be lowered to - 1 5 0 without congeala- tion taking place, but on the slightest agitation it at once solidifies and the temperature rises to o°. It is on account of this property of water to suspend its solidification, that in determining the lower fixed point of a thermometer, the temperature of melting ice, and not that of freezing water, is made use of. Many other liquids exhibit suspended solidification to a very high degree ; thus glycerine may be cooled to — 30° or - 40° with- out solidifying, but if a crystal of solid glycerine be placed in the liquid the entire mass freezes, and does not again melt until a temperature of 15.5 0 is reached. Change of Volume on Solidification.— Most liquids, in the act of solidifying, contract ; that is to say, the solid occupies a smaller volume than the liquid. Consequently the solid is specifically denser, and sinks in the liquid. Thus 100 volumes of liquid phosphorus at 44 0 (the melting-point) when solidified, occupy only 96.7 volumes. Water expands upon solidification, hence ice is relatively lighter than water, and floats upon the liquid. The reverse change of volume accompanies the change of state in the opposite direction. Effect of Pressure upon the Solidifying Point of Liquids. — In the case of liquids that contract upon solidification, increased pressure raises the point of solidification, and consequently raises the melting-point of the solid. The effect, however, is extremely small : thus the solidifying point (and melting-point) of spermaceti under the standard atmospheric pressure is 47.7 0 , while under a pressure of 1 56 atmospheres it is raised to 50.9 0 . With liquids that expand on solidification, increased pressure has the opposite effect, and lowers the solidifying point. Thus, water under great pressure may be cooled below o° and still remain liquid ; and in the same way, ice may be liquefied by increased pressure without altering its temperature. In the case of water it has been found that an increased pressure of n atmospheres, lowers the soli- difying point by 0.0074;*° > hence under a pressure of 135 atmos- pheres, the freezing-point of water (and the melting-point of ice) is lowered i°. This lowering of the melting-point of ice under pres- sure may be illustrated by the experiment represented in Fig. 17. Over a block of ice is slung a fine steel wire, to which are hung a 1 20 Introductory Outlines number of weights. The pressure thus exerted upon the ice, by lowering the melting-point, causes the ice to liquefy immediately beneath the wire, which therefore gradually cuts its way through the block. But as the wire passes through the mass, each layer of water behind it, again resolidifies, being no longer subject to the increased pressure ; hence, although the wire cuts its way com- pletely through the ice, the block still remains intact. Latent Heat Of Fusion. — When a liquid, at a temperatuie above its solidifying point, is cooled, a thermometer placed in the liquid indicates its loss of heat until solidi- fication begins. At this point the temperature remains constant until solidification is complete, when the thermometer again begins to fall. And again, when a solid, at a tem- perature below its melting-point, is heated, its temperature rises until the melting begins, but no further rise of temperature takes place by the application of heat, until lique- faction is complete. The sensible heat that so disappears during fusion is spoken of as the latent heat of fusion. Just as in the pas- sage of liquids into gases, this so- called latent heat represents heat that has ceased to be heat , but which is converted into kinetic energy that is taken up by the molecules : when the liquid passes back into the solid state, this energy is again transformed into sensible heat. The fact that heat is thus changed into energy, and so rendered insensible to the thermometer, may be seen by adding boiling water to powdered ice. A thermometer placed in ice indicates the tem- perature o°, and although boiling water is poured upon it, so long as any ice remains unmelted no rise of temperature of the mixture results, the heat contained in the boiling water being expended in doing the work of liquefying the ice, and converting it into water at o°. When such an experiment is made more exactly, it is found that j kilogramme of water at 80.25°, when mixed with 1 kilogramme of Fig. 17. r 2 1 Solidifying Points of Liquids ice at o°, gives 2 kilogrammes of water at o°. That is to say, the amount of heat contained in a kilogramme of water at 80.25°, is exactly capable of transforming an equal weight of ice at o° into water at o°. As the heat required to raise the temperature of 1 kilogramme of water from o° to i° is the unit of heat, or major calorie , we say that the latent heat of fusion of ice is 80.25 thermal units, or calories. During the solidification of a liquid, the latent heat of fusion is again given out. The solidification, therefore, only takes place gradually, for the heat evolved by the congelation of one portion, is taken up by the neighbouring particles, whose solidification is thereby retarded until this heat is dissipated. In the case of super- cooled liquids and super-saturated saline solutions, the solidifica- tion takes place more suddenly, and the evolution of the latent heat is therefore manifest by a rise of temperature. Effect of Substances in Solution upon the Solidifying Point of a Liquid. — It has long been known, that a lower degree of cold is necessary to freeze salt water than fresh ; and also that the water obtained by remelting ice from frozen sea water, is so little salt as to. be drinkable. Careful experiments have shown that when an aqueous solution of a salt is frozen, pure ice alone separates out, provided the solution is sufficiently dilute to prevent the dissolved salt from crystallising out on account of the reduction in tempera- ture. This also holds in the case of all other solvents that are capable of being solidified, the pure solidified solvent alone sepa- rating when the solution is frozen. The effect of dissolved sub- stances in lowering the solidifying point of the solvent was first discovered by Blagden (1788), who formulated the law that the depression of the freezing-point of aqueous solutions of the same substance, was proportional to the strength of the solution. By referring the lowering of the solidifying point to quantities of the dissolved substances that were in molecular proportions, instead of to equal weights, modern physicists have established the following general law : Solutions containing in equal volumes of the solvent , quantities of dissolved substances proportional to their molecular weights , have the same point of solidification* The relations thus established between the molecular weight of a compound and its influence in lowering the solidifying point of a solvent, furnishes a method for the determination of the molecular weight of a substance. * Except in the case of electrolytes. See page 96. CHAPTER XIV SOLUTION A solution may be defined as a homogeneous mixture of either a gas, a liquid, or a solid, with a liquid, this liquid being termed the solvent .* Substances that are capable of forming such homogeneous mix- tures with a solvent, are said to be soluble in that liquid. The solution of matter in its three states will be treated separately. i. Solution of Gases in Liquids.— When a gas is dissolved by a liquid, the liquid is said to absorb the gas, and although it is held that most liquids are capable of absorbing most gases to a greater or less degree, most of the investigations in this direction have been made with the two liquids, water and alcohol, by Bunsen. The quantity of a gas which a liquid is capable of absorbing depends upon four factors— (i) the specific nature of the liquid ; (2) the nature of the gas ; (3) the temperature of the liquid ; (4) the pressure. (1.) The influence of the solvent may be seen by a comparison of the quantities of the same gas which equal volumes of water and of alcohol are capable of dissolving, thus— 100 volumes of water at o° dissolve 179.6 volumes of carbon dioxide, while 100 ,, alcohol „ 43 2 -9 »* ” (2.) The various quantities of different gases which the same liquid will absorb are found to extend over a very wide range, thus — 100 volumes of water at o° dissolve 4- ™4 volumes of oxygen, while too 114800.0 „ ammonia. * Mixtures of gases are sometimes regarded as solutions, one gas being said to be dissolved in the other. Gases also are sometimes spoken of as dissolving liquids and solids, when liquid and solid substances directly vaporise into them. 123 Henry's Law (3.) The volume of any gas which a liquid can absorb, diminishes with rise of temperature. This will be seen from the following table, where the volumes of different gases are given which 100 volumes of water will absorb at various temperatures. Temperature. Carbon Dioxide. Nitrous Oxide. Oxygen. | 1 Nitrogen. 0 179.6 130- 5 4. 11 2.03 s 144.9 109.3 3.62 1.79 10 118.4 91.9 3- 2 S I.60 20 1 90.I 67.0 2.83 1.40 It was at one time believed that the solvent power ofwater for hydrogen was the same at all temperatures between o° and 25 0 . Recent experiments have shown, however, that there is no excep- tion to the general law in this case ; thus it has been found that 100 volumes of water — At o° dissolve 2.15 volumes of hydrogen. At 5 0 „ 2.06 „ „ At io° „ 1.98 „ u At 20 ,, I.84 55 55 When a solution of a gas in water is heated, the gas being less soluble at the higher temperature is expelled, and in most cases the whole of the gas is driven off at the boiling temperature. This, however, is not invariably the case ; for example, the solution of hydrochloric acid in water, when boiled, will distil, without further evolution of gas, when a solution of definite strength is reached (see Hydrochloric Acid). (4.) The influence of pressure upon the volume of a given gas which a liquid can absorb, was discovered by Henry (1803), and is known as Henry’s law, namely, The volume of the gas absorbed by a liquid is directly proportional to the pressure of the gas. If the pressure be doubled, the same volume of liquid will dissolve twice the volume of the gas, the volume in each case being measured at o° and 760 mm. But since, according to Boyle’s law, the volume of a gas is inversely as the pressure, this law may be thus stated : A. given volume of a liquid will absorb the same volume of a gas at all pressures. 124 Introductory Outlines Thus, if 100 volumes of water at o* dissolve 2.03 volumes of nitrogen, under the standard atmospheric pressure (the volume of the gas being measured at o° and 760 mm.), under twice this pressure, z. in No - &) heat is evolved, and , ?!, Ts Va Ues h *' and ^ 3 ) We have as the net result The number of heat-units, therefore, which expresses the heat of formation of hydrochloric acid, is the heat produced by the union of two atoms of hydrogen with two atoms of chlorine, minus the heat molecule ““ deCOmposition of one hydrogen and one chlorine Compounds such as hydrochloric acid, in the formation of which biTwhirWV, °* 5e ’ are l j rmed exot ^ er,mc compounds, the reaction y which they are produced being an exothermic change ; com- pounds, on the other hand, whose heats of formation are expressed by a negative sign, that is, in whose formation heat disappears are distinguished as endothermic compounds, and the reactions by which they are formed are endothermic reaction. Thus C + S 2 — CS 2 - 19,600 cal. signifies that in the formation of carbon disulphide, heat is absorbed and the compound is therefore an endothermic compound. mentsTmed em T aldetermmati0nS are made by means of instru - tiTnLciX? °T et t rS - TheSe are 0f S reat variety, although to take place Z “ me> The chemical reaction is caused t k r ? U ! ld 1 er such circumstances, that the whole of the heat that is liberated shall be communicated to a known volume of water at a known temperature * water, have ir hi C iL d : te r inati0aS ° f the thermaI value op chemical changes, rimnle r •” m m ° nly a limited number ° f comparatively cuTate thelhe™'! P< f lble ’ h ° Wever ’ a few known data, to cal ^ tly measumd "rT ° f ' cha "^ cannot be y measured. This depends upon the fundamental principle * For descriptions of the various calorimeters, see Treatises on Physics. T ^g Introductory Outlines of thermo-chemistry, which is itself the corollary of the law of the conservation of energy, and which was first experimentally proved by Hess (1840). This principle, which is sometimes termed the law of constant heat consummation , or the law of equivalence of heat and chemical change, may be thus stated : The amount of heat that is liberated, or absorbed, during a chemical process, is de- pendent solely upon the initial and final states of the system, and is independent of the intermediate stages. The following examples will serve to explain the application of the principle :— 1. Let us suppose it is desired to find the heat of formation of carbon monoxide, the data at our disposal being (1) the heat pro- duced when carbon unites with oxygen to form carbon dioxide ; and (2) the heat formed by the combustion of carbon monoxide to carbon dioxide. The thermal equations are— (1) C + 0 2 = C 0 2 + 97,ooo cal. (2) 2 CO + 0 2 = %C 0 2 + 136,000 cal. Halving the second equation, in order to get the heat produced in the formation of 44 grammes of carbon dioxide {i.e., the same weight as in the first), we may represent the equation as— CO + 0 = CO 2, + 68,000 cal* The difference between the two values 97,000 and 68,000 will be the heat of formation of carbon monoxide, therefore we get the equation — C + O = CO + 29,000 cal. 2. The compound, methane (marsh gas), CH 4 , cannot be forme by the direct union of its elements, but its heat of formation can be calculated by the application of this principle. The data in this case are the ascertained heats of formation of carbon dioxide * It must be remembered that this equation does not express the whole truth • as it here stands it would imply that 68,000 calories represent the heat formed by the simple chemical union of 28 grammes of carbon monoxide with of ox P yg en. m reality this number is half the sum o two values namely, the heat of combination of 56 grammes of carbon monoxide with 32 grammes of oxygen, minus the heat absorbed by the decom^sitionof 02 grammes of oxygen molecules into their constituent atoms. The oxygen atom does not exist alone, and whenever free oxygen takes part in ^chemical change, the molecules of the element are first separated into their atoms. T hermo- Chemistry i and of water, and the heat produced by the combustion of marsh gas, the thermal equations being — (1) 0 + 0 2 = C 0 2 + 97,000 cal. (2) 2 H 2 + 0 2 = 2 H 2 0 + 136,800 cal. (3) CH± + W 2 = C 0 2 + 2 H 2 0 + 212,000 cal. The difference between the thermal value of the last process, and the sum of the first and second, represents the heat of formation of marsh gas — 97,000 + 136,800 - 212,000 = 21,800, hence we get the expression— C + 2 H 2 = CH \ + 21,800 cal. PART II THE STUDY OF FOUR TYPICAL ELEMENTS HYDROGEN— OXYGEN— NITROGEN— CARBON AND THEIR MORE IMPORTANT COMPOUNDS CHAPTER I HYDROGEN Symbol, H. Atomic weight = t. Molecular weight = 2. Density = 1. History.— The existence of hydrogen as an individual sub- stance was first established by Cavendish (1766), who applied to it the name inflammable air . He obtained the gas by acting upon certain metals, as iron, tin, and zinc, with either sulphuric or hydro- chloric acid. Occurrence.— In the free state hydrogen occurs only m small quantities upon the earth. It is evolved with other volcanic gases, and is present in the gases which escape from petroleum wells. It is evolved also during the fermentation and decomposition of certain organic compounds, and is therefore present in the breath and the intestinal gases of animals. Hydrogen has also been found in many specimens of meteoric iron, where it is present as occluded gas. . . Hydrogen in the uncombined state exists in enormous ma upon the sun, and is present in certain stars and nebulas. The so-called prominences which are seen projecting from the suns disk to a distance of many thousands of miles, and which were first observed during solar eclipses, consist of vast masses of in- candescent hydrogen. In combination with other elements hydrogen is extremely Hydroge?i 1 5 j abundant : its commonest compound is water, which consists of one part by weight of this element combined with eight parts of oxygen. In combination with chlorine, as hydrochloric acid, with carbon as marsh gas, and with sulphur as sulphuretted hydrogen, this element also occurs in large quantities. All known acids contain hydrogen as one of their constituents, and it is present in almost all organic compounds. Modes Of Formation. -(1.) Hydrogen may be obtained from Fig. 25. water by the action of various metals upon that compound under certain conditions. The metals sodium and potassium will decom- pose water at the ordinary temperatures ; when, therefore, a frag- ment of either of these metals is thrown upon water, the latter is decomposed, and hydrogen set free : — H 2 0 + Na = NaHO + H. The metals being lighter than water float upon its surface, and owing to the heat of the reaction, melt and roll about upon the liquid as molten globules. With potassium, the heat developed is 1 52 Inorganic Chemistry sufficiently great to cause the hydrogen to inflame, and it burns with a flame coloured violet by the vapour of the metal. The hydroxide of the metal, which is the second product of the action, dissolves in the excess of water, rendering the liquid alkaline. The alkalinity of the solution may be made evident by the addition of a reddened solution of litmus, which will be turned blue by the In order to collect the hydrogen evolved by the action of sodium upon water, the metal is placed in a short piece of lead tube closed at one end, which causes it to sink in the liquid, and an inverted glass cylinder filled with water is placed over it, as shown in Fig. 25. The evolved hydrogen then rises as a stream of bubbles into the cylinder and displaces the water.* (2.) Water may be readily decomposed at the boiling-point, by means of zinc, if the metal be previously coated with a thin film of copper by immersion in a dilute solution of copper sul- phate When this copper -coated zinc (known as zinc- copper couple) is heated in a small flask filled with water, and provided with a delivery tube, the oxygen of the water combines with the zinc forming zinc oxide, and hydrogen is evolved, which may be collected over water at the pneumatic trough : *— Zn + H 2 0 = ZnO + H 2 . * For detailed description of these experiments see News’s “Chemical Lecture Experiments," p. 2. Hydrogen 1 5 3 ( 3 .) At a still higher temperature, water in the state of steam can be readily decomposed by the metal magnesium, magnesium oxide being formed and hydrogen liberated : — Mg + H a O = MgO + H 2 . For this purpose the magnesium is strongly heated in a glass bulb (Fig. 26 ), while steam from a small boiler is passed over it. As the temperature of the metal approaches a red heat it bursts into flame, and the issuing hydrogen may be ignited as it escapes from the end of the tube. (4-) If iron be heated to bright redness, and steam be passed over it, the water is decomposed, the oxygen uniting with the iron to form an oxide known as triferric tetroxide, or magnetic oxide oj iron , thus — 3Fe + 4H 2 0 = Fe 3 0 4 + 4H 2 . This method is employed on a large scale for the preparation of hydrogen for commercial purposes. Iron borings or turnings are packed into an iron tube, which is strongly heated in a furnace, and steam from a boiler is passed through the tube. ( 5 .) For laboratory purposes hydrogen is most conveniently pre- pared by the action of dilute sulphuric acid upon zinc : — » Zn + H 2 S0 4 = ZnS0 4 + H 2 . For this purpose granulated zinc (i.e., zinc which has been melted ‘54 Inorganic Chemistry and poured into water) is placed in a two-necked WoulPs bottle (Fig. 2 7), and a quantity of sulphuric acid, previously diluted with six times its volume of water, is introduced by means of the funnel. A brisk action sets in, and hydrogen is rapidly disengaged. After the lapse of a few minutes, the air within the apparatus will be swept out by the hydrogen, when the gas may be collected over water in the pneumatic trough. The hydrogen so obtained is never absolutely pure : it is liable to contain traces of arsenuretted hydrogen, sulphuretted hydrogen, phosphoretted hydro- gen, oxides of nitrogen, and nitrogen. The nitrogen is derived from the air, which finds its way through joints in the apparatus, and also from air dissolved in the acid. There is no known process for removing this impurity. The other gases are due to impurities in the zinc and the sulphuric acid, and can be removed, if required, by passing the hydrogen through a series of tubes containing absorbents (see p. 187). Absolutely pure sulphuric acid, even when diluted with water, has no action upon perfectly pure zinc. Scrap iron may be substituted for zinc, but the hydrogen so obtained is much less pure, and is accompanied by compounds of carbon (derived from the carbon in the iron), which impart to the gas an unpleasant smell ; the reaction in this case is the following Fe + H 2 S0 4 = FeS0 4 + H 2 . Hydrochloric acid can be employed in place of sulphuric acid with either zinc or iron, the reaction then being : — Zn + 2HC1 = ZnCl 2 + H 2 . (6.) Hydrogen in a high degree of purity is conveniently prepared in small quantity by the electrolysis of water acidulated with sul- phuric acid (see p. 183). (7.) Hydrogen is disengaged when certain metals, such as zinc, iron, magnesium, and aluminium, are boiled with an aqueous solu- tion of potassium or sodium hydroxide. Thus, in the case of zinc, when this metal in the form of filings is boiled with a solution of potassium hydroxide, hydrogen is evolved, and a compound of zinc, potassium, and oxygen remains in solution, namely, potassium zinc oxide, thus : — ♦ 2KHO + Zn = H 2 + ZnK 2 0 2 . (8.) Hydrogen is also obtained by heating alkaline oxalates, or Hydrogen i$5 formates, with either potassium or sodium hydroxide, with the simultaneous formation of an alkaline carbonate ; thus with sodium oxalate : — Na 2 C 2 0 4 + 2NaHO = H 2 + 2Na 2 CO a . Properties. — Hydrogen is a colourless gas, and has neither taste nor smell. It is the lightest known substance, being 14.435 times lighter than air. Its specific gravity is 0.0693 (air = 1). One litre of the gas at o° C., and under a pressure of 760 mm. of mercury (/. S ° fer 35 il has been Penetrated by man t must be remembered, however, that the greatest depth to 7 which is after a !l| Xa ? med ’ When COm P ared with the diameter of the earth IS after all only, as it were, a mere scratch. 159 i6o Inorganic Chemistry Average Composition of the Earth's Crust. Oxygen . Silicon . Aluminium Iron Calcium . Magnesium Sodium . Potassium 44.0 to 48.7 22.8 9-9 9-9 6.6 2.7 2.4 1-7 36.2 6.1 2.4 0.9 0.1 2.5 3-i 100.00 100.00 Modes Of Formation.— (1.) Oxygen may readily be obtained by a slight modification of Priestley’s original method, namely, y heating mercuric oxide in a glass tube, by means of a Bunsen flame The red oxide of mercury first darkens in colour, an decomposed by the action of the heat into mercury and oxygen, thus — T _ 2HgO = 2Hg + 0 2 . The evolved oxygen may be collected over water in *e pneumatic trough, while the mercury condenses in the form of “ etal ' C globules upon the cooler parts of the tube. This method of obtaining oxygen is never employed when any quantity of the gas is required— it is chiefly of histone interest. S (2 ) For experimental purposes, oxygen is best prepared from potassium chlorate. When this salt is heated it melts, and at about 400" decomposes with brisk effervescence due to the evolution of oxygen, while potassium chloride remains : — KC10 3 = KC1 + 30. If the potassium chlorate be previously mixed with about one- fourth of its weight of manganese dioxide, it gives up the whole of [roxygen at a temperature considerably below the melting-point of the^salt and at a greatly accelerated rate. When, therefore, the oxygen U not required to be perfectly pure, a mixture of these two * The mechanism of this J and £- resoived into KCIO3 and O. Oxygen 161 substances is usually employed. The mixture may be conveniently heated in a “ Florence ” flask, supported in the position shown in the figure, and gently heated with a Bunsen flame. The gas is washed by being passed through water, and then collected either at the pneumatic trough or in a gas-holder. The manganese dioxide is found at the end of the reaction to be unchanged : the part it plays in the decomposition belongs to a class of phenomena to which the name catalysis is applied ; the manganese dioxide, in this instance, being the catalytic agent. It was at one time supposed that by its mere presence, itself under- going no change, the manganese dioxide enabled the potassium chlorate to give up its oxygen more readily and at a lower tempera- ture ; but the accumulated evidence which has been collected by the study of an increasing number of similar cases of catalytic action, leads to the conclusion that the manganese dioxide is here playing a more distinctly chemical part in the reaction. So far as is known, in all phenomena of this order, the catalytic agent is a substance which possesses a certain degree of chemical affinity for one of the constituents of the body to be decomposed, and the influence of this attraction is a necessary factor in determining the splitting up of the compound. Owing, however, to certain condi- tions which are present, such, for example, as the particular temperature at which the reaction is conducted, the catalytic agent is unable to actually combine with the constituent for which it has this affinity, or if it combines, the combination it forms is unable to exist, and is instantly resolved again : hence the catalytic agent comes out of the action in the same state as it was at the commencement. In the case before us, it is believed that a cycle of changes takes place,* in which the power possessed by manganese to enter into * M'Leod. L 162 Inorganic Chemistry higher states of oxidation, results first in the formation of potassium permanganate, KMn0 4 ; with the simultaneous production of chlorine and oxygen, thus — (1) 2Mn0 2 + 2KC10 3 = 2KMn0 4 + Cl 2 + 0 2 . The potassium permanganate then passes into potassium man- ganate, K 2 Mn0 4 , with evolution of oxygen, and partial reformation of manganese dioxide, thus — ( 2 ) 2KMn0 4 = K 2 Mn0 4 + Mn0 2 + 0 2 . And this is decomposed, by the chlorine evolved by the first re- action, into potassium chloride, manganese dioxide, and oxygen, thus — ( 3 ) K 2 Mn0 4 + Cl 2 = 2KC1 + Mn0 2 + 0 2 . ( 3 .) When manganese dioxide itself is heated to bright redness, it parts with one-third of its oxygen, and is converted into trimanganic tetroxide. 3Mn0 2 = Mn 3 0 4 ( 4 .) Other peroxides, when heated, similarly yield a portion of the oxygen they contain. One of these, namely, barium peroxide, is now largely employed for the preparation of oxygen upon a manu- facturing scale. -This method, known as BrinI process, from the name of the inventor, is based upon the fact, that when barium oxide (BaO) is heated in contact with air, it unites with an additional atom of oxygen, forming barium peroxide, thus— BaO + O = Ba0 2 . And that when this substance is still further heated, it again parts with the additional oxygen and is reconverted into the monoxide— Ba0 2 = BaO + O. The process, therefore, is only an indirect method of obtaining oxygen from the air, the same quantity of barium monoxide being employed over and over again. In practice it was found, that instead of effecting the two reactions by altering the temperature, which involved loss of time and considerable expense, the same result could be obtained by altering the pressure and keeping the temperature constant. If the monoxide be heated to the lower temperature, at which the first reaction takes place, and air be passed over it at the ordinary atmospheric pressure, atmospheric oxygen is taken up and barium peroxide is formed. If the pressure Oxygen 163 be then slightly reduced by suitable exhaust pumps, the peroxide immediately gives up one atom of oxygen without any further application of heat, and is retransformed into the monoxide. In this way, by alternately sending air through the heated retorts containing the oxide, and then exhausting the retorts, a continuous process is obtained without change of temperature. The modus ofitrandi of the process will be seen from Fig 31 164 Inorganic Chemistry which represents the general arrangement of the apparatus. A number of retorts, R, consisting of long narrow iron pipes, are arranged vertically in rows in the furnace, where they are heated by means of “ producer-gas ” (i.e. carbon monoxide with atmospheric nitrogen, obtained by the regulated combustion of coke). By means of the pump P, air is drawn in at the “ air intake,” and forced through purifiers in order to withdraw atmospheric carbon dioxide ; the complete removal of this impurity being essential to the successful carrying out of the operation. The purifiers are so arranged, that any of them can be thrown out of use at will. By means of automatic gear the purified air is sent through pipe J to the distributing valve X, from which it passes by the pipe Y into the retorts, being made to pass down through one row, and up through the other. The oxygen is then absorbed, and the accumulating nitrogen escapes by the relief valve W. When the absorption of oxygen by the barium monoxide in the retorts has continued for ten or fifteen minutes, the automatic reversing gear comes into operation. The relief valve W is thereby closed, communication with the purifiers is cut off, and the action of the pumps at once causes a reduction of pressure within the retorts. When the pres- sure falls to about 660 mm. (26 inches, or about 13 lbs. on the square inch), the peroxide gives up oxygen, and is reduced to the monoxide. The oxygen is drawn away by the pipe J and is passed on to a gas-holder. The first portions of gas that are drawn out of the retorts, will obviously be mixed with the atmospheric nitrogen which was there present ; in order that this shall be got rid of, the automatic gear is so arranged, that communication with the pipe leading to the gas-holder is not opened until a few seconds aftei; the reversing gear is in operation, and the first portions of gas that are pumped out, are made to escape into the air by a snifting valve S, which is automatically opened and closed. (5.) Oxygen may be obtained by heating manganese dioxide with sulphuric acid, the dioxide parting with the half of its oxygen, and a sulphate of the lower oxide being formed — Mn 0 2 + H 2 S 0 4 = MnS 0 4 + H 2 0 + O. (6.) Similarly, potassium dichromate (a salt containing chromium trioxide, Cr 0 3 ), when heated with sulphuric acid, yields oxygen ; the chromium at the same time being reduced to a lower state of oxi- Oxygen 165 dation, viz., Cr 2 0 3 , in which condition it unites with sulphuric acid, forming chromium sulphate — K 2 Cr 2 O r + 4H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4H 2 0 + 30. During the reaction, the red colour of the dichromate changes to the deep olive-green colour possessed by chromium sulphate. (7.) Many other highly oxidised salts yield oxygen when acted upon by sulphuric acid ; thus, with potassium permanganate, the following action takes place : — K 2 Mn 2 0 8 + 3H 2 S0 4 = K 2 S0 4 + 2MnS0 4 + 3H 2 0 + 50. (8.) If hydrogen peroxide be added to dilute sulphuric acid, and the mixture dropped upon a solution of potassium permanganate contained in a suitable generating flask, a rapid evolution of oxygen takes place at the ordinary temperature, thus — K 2 Mn 2 0 8 + 3H 2 S0 4 + 5H 2 0 2 = K 2 S0 4 + 2MnS0 4 + 8H 2 0 + 50 2 . (9.) When strong sulphuric acid is dropped upon fragments of brick or pumice-stone, contained in an earthenware or platinum retort, and maintained at a bright red heat, the acid is decomposed into water, sulphur dioxide, and oxygen— h 2 so 4 = H s O + so 2 + O. The products of the decomposition are passed through water, which absorbs the sulphur dioxide, and also arrests any undecomposed sulphuric acid, and the oxygen is collected over water. When this process is used on a large scale, the sulphur dioxide is absorbed by being passed through a tower filled with coke, and down which a stream of water is allowed to trickle, and the solution so obtained can be utilised in the manufacture of sulphuric acid. (10.) Oxygen can be obtained from bleaching-powder by methods which afford interesting instances of catalytic action.* The composition of bleaching-powder is expressed by the formula Ca(OCl)Cl. If this substance be mixed with water, and a small quantity of precipitated cobalt oxide added, and the mixture gently warmed, oxygen is rapidly evolved. The cobalt oxide, CoO, is the catalytic agent; it is able to combine with more oxygen to form * Experiments 35, 36, 37, 152, “ Chemical Lecture Experiments,” new ed. x 66 Inorganic Chemistry Co 2 0 3 , but this compound is reduced as fast as it is formed, and the oxygen is evolved as gas — (i.) Ca(OCl)Cl + 2CoO = Co 2 0 3 + CaCl* ( 2 .) Co 2 O s = 2CoO + O. A solution of calcium hypochlorite, which may be obtained from bleaching-powder (see Bleaching-powder), behaves in the same way ; and, as in the above reaction, nickel oxide may be substi- tuted for cobalt — Ca(OCl ) 2 = CaCl 2 + 0 2 . (ii.) A similar instance of catalysis, by which oxygen maybe obtained, is seen when a stream of chlorine gas is passed through boiling milk of lime, to which a small quantity of the oxide of cobalt or nickel has been added — CaH 2 0 2 + Cl 2 = CaCl 2 + H 2 0 + O. A reaction of the same order takes place, when the milk of lime is replaced by either potassium or sodium hydroxide — 2NaHO + Cl 2 = 2NaCl + H 2 0 + O. ( i 2 .) When a mixture of steam and chlorine gas is heated to bright redness, the steam is decomposed, the hydrogen combining with the chlorine to form hydrogen chloride (hydrochloric acid), and the oxygen is set free — H 2 0 + Cl 2 = 2HC1 + O. In order to prepare oxygen by this reaction, chlorine gas is caused to bubble through water which is briskly boiling in a glass flask, F (Fig. 32 ). The mixture of chlorine and steam is then passed through a porcelain tube filled with fragments of porcelain, and maintained at a bright red heat in a furnace. The issuing gases are passed through a Woulf’s bottle, containing a solution of sodium hydroxide, in order to absorb the hydrochloric acid, and the oxygen is collected at the pneumatic trough. ( 13 .) Oxygen is formed on a large scale in nature by the decom- position of atmospheric carbon dioxide by the green leaves of plants, under the influence of light. The carbon dioxide is decom- Oxygen 167 posed into carbon, which is assimilated by the plant, and into oxygen which is thrown into the atmosphere. It has been esti- mated that 1 square metre of green leaf is able, under the influ- ence of sunlight, to decompose more than 1 litre of carbon dioxide per hour. ( 14 .) Of the many other methods by which it has been proposed, from time to time, to manufacture oxygen on a large scale, may be mentioned one, known as the Tessie du Motay process, from the name of the inventor. This method consists in the alternate for- mation and decomposition of sodium manganate. The process consists of two operations, which are carried out at different tem- peratures. When a current of air is passed over a moderately heated mixture of manganese dioxide and sodium hydroxide, sodium manganate is formed — 2Mn0 2 + 4NaHO + 0 2 = 2H 2 0 + 2Na 2 Mn0 4 . And if this sodium manganate be heated to bright redness, and a current of steam at the same time passed over it, the manganate is reduced to dimanganic trioxide, sodium hydroxide is reformed, and oxygen evolved, thus — 2Na 2 Mn0 4 + 2H 2 0 = Mn 2 0 3 + 4NaHO + 30. On again passing air over the residue, after allowing the tempera- 1 68 Inorganic Chemistry ture of the mass to fall to that at which the first reaction was conducted, sodium manganate is once more reformed — Mn 2 0 3 + 4NaHO + 30 = 2H 2 0 + 2Na 2 Mn0 4 . Properties. — Oxygen is a colourless gas, having no taste or smell. It is slightly heavier than air, its specific gravity being 1.1056 (air =1). One litre of the gas, at the standard temperature and pressure, weighs 1.43028 grammes. Oxygen is slightly soluble in water. 1 c.c. of water at o° C. dissolves 0.0489 c.c. of oxygen measured at o° C. and 760 mm. pressure. The solubility of oxygen in water, diminishes as the temperature rises in accordance with the interpolation formula (Winkler) : — £■=0.0489 - .ooi34i3^+.oooo283/ 2 -.oooooo29534/ 3 . Fish are dependent upon the dissolved oxygen in water for their supply of this gas for respiration. Oxygen is also soluble in molten silver, which is capable of absorbing about twenty times its own volume of this gas (see Silver). Oxygen is endowed with very powerful chemical affinities. Even at the ordinary temperature it is able to combine with such elements as phosphorus, sodium, potassium, and iron. All the chemical phenomena exhibited by the atmosphere, are due to the presence in it of free oxygen, the atmosphere being practically oxygen diluted with four times its volume of nitrogen. Thus, when a piece of bright metallic sodium is exposed to the air, the surface becomes instantly tarnished, and coated over with a film of oxide : when iron rusts, it in the same way is being acted upon by the oxygen of the air, forming oxide of iron : in these cases the metals are said to become oxidised. If the metal be obtained in a sufficiently finely divided condition before being exposed to the air, or to pure oxygen, this process of oxidation may proceed so rapidly, that the heat developed by the combination will cause the metal to burn. When the process of oxidation is accompanied by light and heat, the phenomenon is known as combustion , the oxygen being spoken of as the supporter of combustion : bodies which burn in the air, therefore, are simply undergoing rapid combination with oxygen. It will obviously follow, that bodies which are capable of burning in the air, will burn with greatly increased rapidity and brilliancy, when their combustion is carried on in pure or undiluted oxygen. A glowing chip of wood, or a taper with a spark still upon the Oxygen 169 wick, when plunged into pure oxygen, will be instantly rekindled. Such substances as sulphur, charcoal, phosphorus, which readily bum in air, when burnt in pure oxygen, carry on their combustion with greatly increased brilliancy. Many substances which are not usually regarded as combustible bodies will burn in oxygen, if their temperature be raised sufficiently high to initiate the combustion ; thus a steel watch-spring, or a bundle of steel wires, if strongly heated at one end, will burn in oxygen, throwing out brilliant scintillations. This experiment is most readily shown by project- ing a spirit-lamp flame upon the ends of a bundle of steel wire, by means of a stream of oxygen, as shown in Fig. 33. As soon as the ends of the wire are sufficiently heated, and begin to burn, the lamp may be withdrawn and the wire held in the issuing stream brilliancy 1 * “ ^ C ° minUe itS combustion with great It is a remarkable fact, and one which has not yet received anv satisfactory explanation, that these instances of combustion in oxygen will not take place if both the gas and the material be absolutely dry. It has been shown that phosphorus, sealed up in a tube with oxygen which has been absolutely freed from aqueous vapour, may even be distilled in the gas without any combination taking place. The presence of the minutest trace of moisture however is sufficient to allow the action to proceed, but the exact’ way in which this operates in causing the effect is at present not known with certainty. See also p. 12. Oxygen is the only gas which is capable of supporting respira- tion : an animal placed in any gas or gaseous mixture containing no free oxygen rapidly dies. Undiluted oxygen may be breathed * See also Experiments 48 to 52, " Chemical Lecture Experiments." j j 0 Inorganic Chemistry with impunity for a short time, but its continued inhalation soon produces febrile symptoms. The inhalation of oxygen is occa- sionally had recourse to in cases of asphyxiation, or under circumstances of great bodily prostration, where the necessary oxygenation of the blood cannot take place on account of the enfeebled action of the lungs. Compressed oxygen acts upon the animal economy as a poison : an animal placed in oxygen gas under a pressure of only a feu atmospheres quickly dies. . , During the respiration of man, air is drawn into the lungs, and is there deprived of 4 to 5 per cent, of its oxygen, and gains 3 to 4 per cent, of carbon dioxide.' The oxygen that is withdrawn from the inhaled air by means of the lungs, is absorbed by the blood. The power to absorb this oxygen is believed to reside m a 1 crystal- line substance, contained in the corpuscles of the blood, called hemoglobin, with which it enters into feeble chemical union, form- ing the substance known as oxyhemoglobin. This substance is red, and imparts to arterial blood its well-known colour. During its circulation in the system, the oxyhaemoglobin parts with its oxygen, and is reconverted into the purple-coloured hemoglobin Under normal conditions, the whole of the oxyhemoglobin is not so reduced, for venous blood is found still to contain it to some extent. The amount of carbon dioxide exhaled is diminished during sleep, and to a still greater extent dunng hibernation. Oxygen can be liquefied at very low temperatures by the appli- cation of moderate pressure (see Liquefaction of Gases), first liquefied in 1877 by Cailletet, and independently by Pictet. Its critital temperature is - 118.8°, at which point a pressure of 50 atmospheres is required to bring about its liquefaction. 5 Liquid oxygen is a pale steel-blue, mobile liquid, which bolls at -181°. Its specific gravity at - 181° is 1. 124. The liquid expan s. when warmed, much more rapidly than gases do for the same increment of temperature, and its density diminishes in proportion, thus — At - 181" „ - 139 ' ,, " 134 ' •» “ I2 9* density = 1.124. „ = -877. „ = .806. „ = - 755 - Allotropy 171 Isomerism— Polymerism— Allotropy. Isomerism.— It is frequently found that two different compounds have the same composition ; that is, their molecules are composed of the same number of the same atoms, and yet the substances have different properties. Such compounds are said to be isomeric , the one is an isomer of the other, and the phenomenon is called isomerism. Cases of isomerism are so numerous among the compounds of carbon (i.e. , in the realm of organic chemistry, see Carbon, p. 259), that it has been found convenient to classify them. The term isomerism, therefore, is frequently restricted to cases in which the compounds have the same percentage composition, the same molecular weight, and belong to the same chemical type or class of substances. Thus, the two compounds dimethyl benzene and ethyl benzene, are both expressed by the formula C 8 H 10 . The molecules in each case contain 8 atoms of carbon and 10 atoms of hydrogen, they therefore have the same molecular weight and the same percentage composition ; and as they both belong to the same type, or family they are said to be isomeric with each other. The difference in the properties of these compounds is due to a difference in the arrangement of the atoms within the molecules, and this difference is expressed in their formulae in the following manner : — Dimethyl benzene, C 6 H 4 (CH 3 )2. Ethyl benzene, C 6 H 5 (C 2 H 5 ). * Different compounds having the same molecular weight and the same per- centage composition, but which do not belong to the same family of compounds are distinguished as metamers. Thus, the two compounds, acetone and allyl alcohol are each expressed by the formula C 3 H 6 0 . They have the same molecular weight and the same percentage composition, but belong to two widely different types of compounds ; they are therefore called metameric compounds. The difference between them is again due to a difference in molecular structure, and they are distinguished by formulae which convey this difference, thus : — J Acetone, CO(CH 3 ) 2 . Allyl alcohol, C 3 H 5 (HO). Polymerism. This term is employed to denote those cases in which dif- ferent compounds belonging to the same family have the same percentage composition but differ in molecular weight ; that is to say, their molecules are composed of the same elements, which are pr .sent in the same proportion ■ but they do not contain the same actual numbers of the various atoms, and therefore CH ?£? W r g ^ S - ,T hUS ’ thC com P oun ds, ethylene, C 2 H 4 ; propylene, C 3 H 6 , butylene, C 4 H S , belong to the same family, and have each the same percentage composition, but they differ in molecular weight. These sub- stances are said to be polymers of one another. Allotropy may be regarded as a special case of polymerism. In its widest sense the term is sometimes used to denote polymerism in general, buT ff is usuaUy restricted to those instances of polymerism which are exhibited by elementary bodies only. Many of the elements are capable, under special 172 Inorganic Chemistry conditions, of assuming such totally different habits and properties, that they appear to be entirely different substances. Thus, the element sulphur, as usually seen, is a primrose-yellow, opaque, solid substance, extremely brittle, and readily dissolved by carbon disulphide. Under certain circumstances it may be made to appear a totally different thing ; it is then a translucent amber- coloured substance, soft and elastic like indiarubber, and insoluble in carbon disulphide ; it is still sulphur, and sulphur only. Phosphorus, again, as usually known, is a nearly colourless, translucent, wax-like solid, which melts at a temperature only slightly above that of the hand, and which takes fire a few degrees higher ; it is also extremely poisonous. Under special influences phosphorus can be made to assume the following properties : — A dark reddish- brown powder, resembling chocolate, which may be heated to 250° without taking fire, and which is non-poisonous. The substance is still phosphorus, and phosphorus only. This property possessed by certain of the elements of appearing in more than one form, of assuming, as it were, an alias, is called allotrofy ; the more uncommon form being spoken of as the allotropic modifica- tion, or the allotrope of the other. From a study of the best known instances of this phenomenon, it is believed that allotropy, in all cases, is due to a difference in the number of atoms of the element that are contained in the molecule. In the case of ozone, which is the allotrope of oxygen, this is known to be the case. Ordinary oxygen molecules consist of two atoms, while the molecule of ozone is an aggregation of three oxygen atoms. OZONE. Molecular symbol, 0 3 . Molecular weight, 47.88. Density, 23.94. History.— When an electrical machine is in operation, a peculiar and characteristic smell is noticed in its vicinity. The same smell is sometimes observed in and about buildings, or other objects, when struck by lightning. In 1785 it was observed by Van Marum that when electric sparks were passed in oxygen, the oxygen acquired this peculiar smell. Schonbein (1840) showed that the oxygen obtained by the electrolysis of water, also contained this substance having a smell, and he gave to it the name ozone, signifying a smell. Schonbein made a careful study of the substance, and found that it might be obtained by various other methods. The more recent work of Andrews, Soret, and Brodie has brought our knowledge of the constitution of ozone to its present state. Occurrence. — Ozone is present in the atmosphere in extremely small quantities (see Atmospheric Ozone). Modes of Formation.— (1.) Mixed with an excess of oxygen, ozone is best obtained by exposing pure dry oxygen to the influence of the silent electric discharge. This may be effected by means of the instrument shown in Fig, 34, known as “ Siemens’ ozone tube. Ozone 173 1 1 consists of two concentric glass tubes, A and B. Tube A is coated upon its inner surface with tinfoil, which is brought into metallic contact with the binding screw D, as shown in the figure. Tube B is coated upon the outer surface, also with tinfoil, which is in metallic connection with binding screw C. These two surfaces of tinfoil are connected by means of their respective binding screws Fig. 34. lo the terminals of a Ruhmkorf coil, and the slow stream of oxygen w ich is admitted at E, and which passes along the annular space between the two tubes, is there exposed to the action of the silent electric discharge. A small portion of the oxygen so passing, becomes converted into the allotropic modification, and the mixture Fig. 35. end'offhe a^arTs* ^ “ “* be « opposite For general purposes of illustration, a very simple arrangement ‘? a ybe Subs “ fo / the afa ove. It consists, as shown in fig. 35 of a straight length Of narrow glass tube having a piece of platinum wire down the inside, which passes out through the walls of the tube near to one end, and is there sealed to the glass. A second 174 Inorganic Chemistry platinum wire is coiled round the outside of the tube, and these two wires are connected to the induction coil. On passing a slow stream of oxygen through the tube, the issuing gas will be found to be highly charged with ozone. (2.) Ozone is also formed when an electric current is passed through water acidulated with sulphuric acid. Thus, in the ordinary electrolysis of water, the oxygen evolved from the positive electrode is found to contain ozone in sufficient quantity to be readily detected, both by its odour, and by other tests. (3.) During many processes of slow oxidation at ordinary tempera- tures, ozone is formed in varying quantities. Thus, when phos- phorus is exposed to the air, an appreciable amount of ozone is formed. One or two short sticks of freshly scraped phosphorus are for this purpose put into a stoppered bottle containing air, and allowed to remain for a short time, when the air will be found to contain ozone. (4.) Ozone is also found during the combustion of ether upon the surface of red-hot platinum. When a spiral of platinum wire is warmed in a gas-flame, and while hot is suspended over a small quantity of ether contained in a beaker, the mixture of ether vapour and air undergoes combustion upon the surface of the platinum, which continues in an incandescent state so long as any ether remains. During this process of combustion, a considerable quantity of ozone is formed. (See also Peroxide of Hydrogen.) (5.) Ozone is formed during the liberation of oxygen in a number of the reactions by which that gas is obtained ; thus from manga- nese dioxide and sulphuric acid the oxygen that is evolved contains sufficient ozone to answer to the ordinary test. In the same way, by the action of sulphuric acid upon barium peroxide or potassium permanganate, this allotrope is present with the ordinary oxygen that is evolved, Properties.— As prepared by any of the methods described, ozone is always mixed with a large excess of unaltered oxygen, probably never less than about 80 per cent, of the latter gas being present. Even in this state of dilution it has a strong and rather unpleasant smell, which rapidly induces headache. When inhaled it irritates the mucous membranes, and is rather suggestive of dilute chlorine. Ozone is a most powerful oxidising substance ; it attacks and rapidly destroys organic matter : on this account ozonised oxygen cannot be passed through the ordinary caoutchouc tubes, as these Ozone 175 are immediately destroyed by it. It bleaches vegetable colours, and most metals are at once acted upon by it. Even metals like ™ erC . Ur> ’’ wh,ch are entirely unaltered by ordinary oxygen, are attacked by ozone. Its action upon mercury is so marked in its result, that the presence of exceedingly small traces of ozone can be detected by it ; the mercury is seen to lose its condition of perfect liquidity, and adheres to the surface of the glass vessel containing it leaving “tails” upon the glass. Ozone converts lead sulphide • bS - 'j-j ead sulphate ( pbs °4)> and liberates iodine from potas- sium iodide — r 2KI + H 2 0 + 0 3 = 0 2 + I 2 + 2KHO. This property is generally made use of for detecting the presence of ozone, advantage being taken of the fact, that iodine, when set rWn kT* C0 . mbmatl0n ,n the Presence of starch, gives rise to a deep blue-coloured compound, the reaction being one of extreme delicacy. In order to apply this test for ozone, strips of paper are dipped in an emulsion of starch to which a small quantity of potas- sium iodide has been added. These papers may be dried and preserved, and are usually spoken of as ozone test papers. When one of these papers is moistened with water, and placed in air containing ozone, the iodine is liberated from the potassium iodide bW hvT t ■ PreS ? Ce ° f S * arCh ’ * he paper instan “y becomes blue by the formation of the coloured compound of starch. It will e o vious that this method of testing for ozone can only be relied upon, when there is no other substance present which is able to decompose potassium iodide ; for example, when testing for ozone m the atmosphere, the presence °f oxides of nitrogen or peroxide are lfobfom h ^ f* Capable ° f liberatin ? iodine, and are able to be present in the air), would materially vitiate the result (see also Atmospheric Ozone). The above decomposition of potassium iodide by ozone, may be made use of as a test for MkT I" an ° ther way> which > abb ough less delicate, is also less ’itini^ ° 6 Vltlat j d by tbe presence of Other substances. Blue Sv Ea” T. T u iDt ° Water Which has been rendered very been ad C ^/ nd ‘° Wh ‘ Ch a SmaI1 \ uantity of p o‘ aa sium iodide has r aadded - Tbe papers may be dried and preserved. On oistemng one of these papers with water and exposing it to one, the iodide is decomposed as in the former case, and the potassium hydroxide which is formed, being a powerfully alkaline substance, converts the colour of the litmus from “bfuf 176 Inorganic Chemistry When heated to a temperature of about 250°, ozone is retrans- formed into ordinary oxygen ; if, therefore, the ozonised gas obtained by means of the Siemens’ ozone tube, be passed through a glass tube heated by means of a Bunsen flame, the whole of the ozone will be decomposed, and the issuing gas will therefore be found to be without action upon the ozone test papers. Ozone is also decomposed by certain metallic oxides, such as those of manganese, copper, and silver. The action appears to be one of alternate reduction and oxidation, the metallic oxide remain- ing unaltered at the conclusion, thus— Ag 2 0 + 0 3 = Ag 2 + 20 * Ag 2 + 03 = Ag 2 0 + 0 2 . The oxidising power of ozone is due to the instability of the mole- cule, and the readiness with which it loses an atom of oxygen, leaving a molecule of ordinary oxygen, thus— 0 3 = O + O* The oxygen molecule is comparatively inert, but the liberated atom in its nascent state, is endowed with great chemical activity. No change of volume accompanies these processes of oxidation by ozone, as the volume of the oxygen molecule (0 2 ) is the same as that of the ozone molecule (0 3 ), the third atom of oxygen being that which enters into new combination with the oxidised substance. Ozone is soluble to a slight extent in water, imparting to the solution its own peculiar smell. 1000 c.c. of water dissolve about 4.5 c.c. of ozone. Under the influence of extreme cold, ozone condenses to liquid having an intense blue colour. So deep is the colour, that a layer of it 2 mm. in thickness is opaque. This liquid is obtained by passing ozonised oxygen through a tube which is cooled by being immersed in boilirfg liquid oxygen, which has a temperature of - 181 0 . At this temperature the ozone liquefies, but the oxygen with which it was mixed passes on. Liquid ozone boils at — 106. It is described by Olszweski and Dewer as an extremely explosive substance. Constitution of Ozone.— The fundamental difference between ordinary oxygen and its allotrope ozone, lies in the fact that the molecule of the latter contains three atoms, while that of ordinary Ozone '77 oxygen consists of only two. Ozone, therefore, is a polymer of oxygen ; Us molecule is more condensed, three atoms occupying two unit volumes. This conclusion as to the constitution of ozone has been arrived at from the consideration of a number of experi- mental facts. (i.) When oxygen is subjected to the action of the electric dis- charge, it is found to undergo a diminution in volume.* This was shown by Andrews and Tait, by means of the tube seen in Fig 36 The tube was filled with dry oxygen, which was prevented from escaping by means of the sulphuric acid contained in the bent por- tion of the narrow tube, which served as a manometer. When the silent discharge was passed through the oxygen, a contraction in the volume took place, indicated by a disturbance of the level of the acid in the syphon. When the tube was afterwards heated to about 300° C. and allowed to cool, the gas was found to have returned to its original volume, and to be devoid of ozone. This could be repeated inde- finitely, the gas contracting when ozonised, and re- expanding when the ozone was converted by heat into ordinary oxygen. As only a very small propor- tion of the oxygen was converted into ozone, this experiment alone afforded no clue as to the rela- tion between the change of volume and the extent to which this conversion took place. (2.) A small sealed glass bulb, containing a solu- tion of potassium iodide, was placed in the tube before the experiment. The oxygen was ozonised, and the usual contraction noticed. The bulb was then broken, and on coming in contact with the ozone present, the potassium iodide was decomposed, iodine being liberated. ’ No further contraction, however, followed ; and, further, when the tube was subsequently heated to 300° and cooled, the gas suffered no increase in volume. By carefully estimating the amount of iodine that was liberated by the ozone, the actual amount of oxygen which had caused this liberation could be determined, according- to the equation — 5 2K1 + H 2 0 + O = I 2 + 2 KHO, * "Chemical Lecture Experiments,” new ed., Nos. 63, 64. M lyS Inorganic Chemistry and it was found, that the volume of oxygen so used up, was exactly equal to the contraction which first resulted on the ozonisation of the oxygen. . , , These facts proved that when potassium iodide was oxidised by ozone, a certain volume of ordinary oxygen was liberated, which was equal to the volume of ozone ; and a certain volume was used up, which was equal to the original contraction. These facts were explained by the supposition, that ozone was represented by the molecular symbol 0 3 ; and its action upon potassium iodide may be expressed as follows— 2 K 1 + H 2 0 + 0 3 = 0 2 + I 2 + 2KHO. (3.) To prove the correctness of this supposition, however, it was necessary to learn the exact relation between these two volumes. This Soret did, by making use of the property possessed by turpen- tine (and other essential oils) of absorbing ozone without decom- posing it ; and he found, that the diminution in volume which took place, by' absorbing ozone from ozonised oxygen, was exactly twice as great as the increase in volume that resulted when the same volume of ozonised oxygen was heated. (4.) If the molecule of ozone be correctly represented by 0 3 , its density will be 24, as against 16 for oxygen ; and its rate of diffu- sion will be proportionately slower in accordance with the law of gaseous diffusion (see Diffusion of Gases, p. 81). Soret found that this was actually the case, and from his experiments the number 24 for the density of ozone receives conclusive confirmation. CHAPTER III COMPOUNDS OF HYDROGEN WITH OXYGEN There are two oxides of hydrogen known, viz. Hydrogen monoxide, or water . „ n Hydrogen dioxide . * • h 2 o 2 Formula, H 2 0. WATER. Molecular weight = 17.96. and he had also observed that ’ WaS pr ° duced ’ the water so obtained was some- times acid. Cavendish showed, that the water was actually the product of the chemical union of hydrogen with oxygen, and he also discovered that the acidity which this water sometimes pos- sessed, was due to the presence of small quantities of nitric acid ; and he traced the formation of this acid to the accidental pre- sence of nitrogen (from the at- mosphere) with which the gases were sometimes contaminated. Cavendish filled a graduated bell-jar with a mixture of hydro- gen and oxygen, in the propor- Fig. 37. w"h°e f Zn V °'r eS ,° f the , former t0 one of °xyg en - and he attached the bell-jar a stout glass vessel, resembling the pear-shaned apparatus shown in Fig. 37> which was perfect, /dry^nd renZed I79 T 8o Inorganic Chemistry vacuous. On opening the stop-cocks, gas entered the exhausted tube, which is furnished at the top with two platinum wires. The cocks were again closed and an electric spark passed through the mixed gases, thereby causing their explosion, when the interior surface of the previously dry glass vessel was found to be dimmed with a film of moisture. On again opening the stop-cocks, more gas was drawn into the upper vessel, the same volume passing in as originally entered the evacuated apparatus. This showed that the two gases in their combination with each other had entirely disappeared. By repeatedly filling the vessel with the mixed gases, and causing them to unite in this way, Cavendish succeeded in collecting sufficient of the water to identify the liquid, and prove that it was in reality pure water. The more exact volumetric proportion in which oxygen and hydrogen combine to form water, has been determined by modem eudiometric methods, which have been developed from Cavendish’s experiment. Accurately measured volumes of the two gases are introduced into a long graduated glass tube, standing in the mer- curial trough, and provided with two platinum wires, by means of which an electric spark can be passed. The gases are caused to unite by means of the spark, and the contraction in volume is carefully observed. Fig. 38 shows the apparatus for this purpose. The long glass tube A having a millimetre scale graduated upon it, and having two platinum wires sealed into the glass near the upper and closed end, is completely filled with mercury, and inverted in the trough of the same liquid : this tube is known as a eudiometer A quantity of pure oxygen is then introduced into the tube, and the volume occupied by the gas carefully read off upon the gradua- tions. Seeing that the volume occupied by a given mass of gas is dependent both upon the temperature and the pressure, each of these factors has to be taken into account in the process of this experiment. The temperature is ascertained by the attached thermometer T. The pressure under which the gas is, will be the atmospheric pressure at the time (ascertained by the barometer B placed near the apparatus) minus the pressure of a column of mercury, equal to the height of the mercury within the eudiometer above the level of that in the trough. This height is obtained in millimetres, by carefully reading upon the graduated scale the level of the mercury in the trough, and the top of the column in the tube, and the number of millimetres so obtained is deducted from the barometric reading. These observations are made by means of Water 1 8 1 a telescope placed at such a convenient distance, that the heat of the body may not introduce disturbances. The data obtained, give the volume of gas at a particular tem- perature, and under a pressure less than that of the atmosphere. By the process of calculation explained under the general pro- perties of gases (p. 69), this is reduced to the standard temperature and pressure, viz., o° and 760 mm. A quantity of hydrogen is then introduced into the eudiometer, considerably in excess of that required for complete combination with the oxygen, and the volume again ascertained with the above precautions and corrections. The difference between the first and second reading will give the volume of hydrogen which has been added. The eudiometer is then firmly held down against a pad of caout- chouc upon the bottom of the trough, and the gases fired by an electric spark from a Ruhmkorff coil. A bright flash of light passes down the tube, and on releasing it from the indiarubber bed, mercury enters to fill the space previously occupied by the gases which have combined. 82 Inorganic Chemistry On allowing the instrument to once more acquire the tempera- ture of the surrounding atmosphere, the residual volume is read off, and corrected for temperature and pressure. The following data have now been obtained : — (i.) The volume of oxygen, corrected for temperature and pressure. (2.) The volume of mixed oxygen and hydrogen, corrected for temperature and pressure. (3.) The volume of residual hydrogen, corrected for tempera- ture and pressure. A concrete example will explain how the result is deduced from these observations : — Corrected volume of oxygen used 45*35 Corrected volume after the addition of hydrogen . . 256.05 Corrected volume of residual hydrogen .... 120. 10 256.05 — 45.35 = 210.70 = total volume of hydrogen employed. 210.70 — 120. 10 = 90.60 = volume of hydrogen which has combined with 45.35 volumes of oxygen. • •• 43*35 : 1 : : 90*60 : 1.997. . •. One volume of oxygen has combined with 1.997 volumes of hydrogen to form water.* The volume composition of water may be shown by analytical processes, as well as the synthetical method described above. This decom- position of water is most conveniently effected by means of an electric cur- rent. If the two terminals from a gal- vanic battery are connected to two pieces of platinum wire or foil, and these are dipped into acidulated water, bubbles of gas make their appearance upon each of the wires. If these two strips of platinum be so arranged in a bottle, that all the gas evolved escapes by a delivery-tube (Fig. 39), it will be found that the gas explodes violently on the application to it * In accurate experiments the volume occupied by the minute quantity of water formed, has to be taken into account, and a number of other corrections have to be made that are not mentioned in this outline description of the process. Water 183 of a lighted taper, showing it to be a mixture of oxygen and hy- drogen. By modifying the apparatus in such a way that the gas from each platinum plate shall be collected in sepa- rate tubes, so arranged that the volumes of the gases can be measured, it is found that twice as much hydrogen is evolved, in a given time, as oxygen. A conve- nient form of voltameter is seen in Fig. 40, where the two measuring tubes are suspended over the platinum plates con- tained in a glass basin. The electrode, which is connected with the negative terminal of the battery, is the one from which the larger volume of gas, viz., the hydrogen, is evolved, while the oxy- gen is liberated at the positive plate. When the volumes of the gases are carefully measured, it is found that they are not exactly in the proportion of two of hydrogen to one of oxygen, but that the oxygen is in deficit of this propor- tion. This is due, in the first place, to the greater solubility of oxygen in water than of hydrogen ; and, secondly, to the formation of a certain quantity of ozone during the electrolysis, whereby there is a shrinking of volume in the proportion of three to two. The “electrolysis of water,” as this process is usually called, is not the simple phenomenon that at first sight it might appear to be. In the first place, pure water is not an electrolyte, and it is necessary either to acidulate il’ or to render it alkaline by the addition of sodium or potassium hydroxides. I he first action of the electric current is (in the case of water acidulated with sulphuric acid) to decompose the sulphuric acid into H 2 , which appears at the negative electrode, and S0 4 , which is liberated at the anode. S0 4 , however, breaks down into SO 3 +O, the oxygen being liberated, and the sulphur trioxide at once uniting with a molecule of water present, to regenerate sulphuric acid, i he changes may be thus represented At the At the Cathode. Anode. h 2 so 4 = H 2 + (SO 3 + O). so 3 + h 2 o = h 2 so 4 . Indirectly, therefore, the water is decomposed. As already stated, some of Fig. 40. 184 Inorganic Chemistry the nascent oxygen is converted into ozone, some also unites with water to form hydrogen peroxide, H 2 0 2 , and probably a still larger quantity is employed in oxidising the sulphuric acid to persulphuric, which is always formed in solution at the anode. Berthelot considers that the variation of the proportion of hydrogen and oxygen evolved during the electrolysis of dilute sulphuric acid, from the proportion of two volumes of hydrogen to one volume of oxygen, is mainly, if not entirely, due to the formation of persulphuric acid. The Volumetric Composition of Steam.— When a mixture of oxygen and hydrogen is exploded in a eudiometer, we have seen that a certain contraction of volume follows, due to the formation of water by the uniting gases. The oxygen and hydrogen that have entered into combination have disappeared as gases, the volume of the resultant water being practically negligable. It is important to know what relation exists between the volume of the uniting gases, and the volume of the product of their combina- tion when in a state of vapour ; that is to say, what volume of steam is produced by the union of one volume of oxygen with two volumes of hydrogen ; in other words, whether there is any molecular contraction in the formation of steam. To ascertain this, the mixed gases, in the exact proportions to form water, must be made to combine under such circumstances that the product shall remain in a state of gas or vapour, so that its volume, and that of the mixed gases, may be measured under comparable conditions. For this purpose a mixture of oxygen and hydrogen, obtained by the electrolysis of acidulated water, is in- troduced into the closed limb of the U-shaped eudiometer shown in Fig. 41* This tube is graduated into three equal divisions, indicated by the broad black bands, and is furnished with two platinum wires at the closed end. It is also surrounded by an outer tube, so that a stream of vapour from some liquid, boiling above the boiling-point of water, can be made to circulate. The most convenient liquid for the purpose is amyl alcohol, which boils at 130° In this way the eudiometer, and the contained gases, will be maintained at a constant temperature, high enough to keep the water formed by their combination, in the state of vapour. The amyl alcohol is briskly boiled in the flask, and its vapour is led into the tube surrounding the eudiometer. The temperature of the mixed gases is thereby raised to 130°, and they occupy the three divisions of the tube when the mercury in the open limb is at the same level, that is, when the gases are under atmospheric * See Experiments Nos. 7 A and 75- "Chemical Lecture Experiments,” new ed. Water 185 pressure. The amyl alcohol vapour leaves the apparatus by the glass tube at the bottom, and is conveyed away and condensed. An electric spark is then passed through the gases by means of the induction coil. (In order to prevent the mercury from being forcibly ejected from the open limb of the U-tube at the moment of explosion, an additional quantity of mercury is poured in, and the open end is closed by the thumb when the spark is passed.) On bringing the enclosed gas again to the atmospheric pressure, by adjusting the level of the mercury until it is once more at the same Fig. 41. height in each limb, it will be found that the mercury in the eudio- meter is now standing at the second band ; that is to say, the three *r S r O SaS 0r ' gmall . y P resent have now become two volumes of am. This condensation is expressed in the molecular equation— 0 2 + 2H 2 = 2H 2 0. The Gravimetric Composition of Water.— Having learned the composition of water by volume, and knowing als 0 g that the X 86 Inorganic Chemistry relative weights of equal volumes of oxygen and hydrogen are as 1 5.96 : 1, the composition by weight can readily be calculated, thus 1 volume of oxygen = 1 5 - 9 ^ 2 volumes of hydrogen = 2.00 17.96 17.96 parts by weight of water are composed of 2.00 parts by weight of hydrogen, and 1 5.96 parts of oxygen, or, expressed centesi- mally, we have — Oxygen . . * .88.86 Hydrogen . . • II * I 4 100.00 The composition of water by weight has been experimentally determined with great care by a number of chemists. Fig. 42. The apparatus shown in Fig. 42 represents the method employed by Dumas (1843). When copper oxide is heated in a stream of hydrogen, the copper oxide is deprived of its oxygen, which unites with the hydrogen to form water— CuO + H 2 = Cu + H 2 0 . Dumas’ method is based upon this reaction. A weighed quantity of perfectly dry copper oxide was heated in the bulb A, in a current of hydrogen generated from zinc and sulphuric acid in the bottle H, and rendered absolutely pure and dry by its passage through a Water 187 series of tubes containing absorbents. The water formed by the union of the hydrogen with the oxygen of the copper oxide, was collected in the second bulb, B, previously weighed ; and the un- condensed aqueous vapour which was carried forward in the stream of hydrogen, was arrested in the weighed tubes which follow. The increase in weight of the bulb B and the weighed tubes, gave the total weight of water produced ; while the loss of weight suffered by the copper oxide, gave the weight of oxygen contained in that water. The difference between these two weights is the weight of the hydrogen that entered into combination with the oxygen. As a mean of many experiments it was found that in the forma- tion of 236.36 grammes of water, the oxygen given up by the copper oxide was 210.04 grammes. 236.36 - 210.04 = 26.32, therefore 236.36 grammes of water were made up of Hydrogen = 26.32 Oxygen = 210.04 236.36 The ratio of hydrogen to oxygen is therefore as 2 : 15.96. Hydrogen prepared from zinc and sulphuric acid is liable to contain traces of (1.) Sulphuretted hydrogen. This is absorbed in the first tube containing broken glass moistened with a solution of lead nitrate ( 2 .) Arsenuretted hydrogen ( absorbed in the second tube, ’filled with glass (3.) Phosphoretted hydrogen \ moistened with silver sulphate. , ^absorbed in the third tube, containing in one limb 4 - Sulphur dioxide I pumice moistened with a solution of potassium (5.) Carbon dioxide j hydroxide, and in the other fragments of solid l potassium hydroxide. Tubes 4 5, 6, and 7, containing solid potassium hydroxide and phosphorus pentoxide (the two latter being placed in a freezing mixture), are for the pur- pose of withdrawing every trace of aqueous vapour. Tube 8 was weighed before and after the experiment in order to test the absolute dryness of the hydrogen at entered the bulb. In order to get rid of dissolved air, the dilute sulphuric acid used was previously boiled. Tubes 9, 10, n were weighed both before and after the experiment ; while tube 12, which was not weighed, was placed at the nd to prevent any absorption of atmospheric moisture by the weighed tubes Since the time of Dumas this subject has been reinvestigated by other experimenters, who have introduced various modifications into the process • wl ?J a Vlew findin S the weight of hydrogen directly, and of eliminating many of the possible sources of error arising from the presence of impurities in 188 Inorganic Chemistry the hydrogen, the hydrogen has been absorbed by palladium. The metal so charged with hydrogen, can be weighed before and after the experiment, and the actual weight of hydrogen used, directly ascertained. Most recently the matter has been investigated by Scott and Rayleigh, and the results obtained show only the slightest departure from the numbers obtained by Dumas. Properties of Water.— Pure water is a tasteless and odourless liquid. When seen in moderate quantities it appears to be colour- less, but when viewed through a stratum of considerable thickness Fig. 43- ii presents a beautiful greenish-blue colour. This colour may be seen by filling a horizontal tube about 1 5 feet long with the purest water, and passing a strong beam of light through it. It may also be perceived by directing a ray of light through a tall cylinder of water in the manner shown in the figure, and causing it to be reflected up through the water from the surface of a layer of mercury at the bottom ; the immerging ray, being then reflected upon a screen, shows the characteristic colour of the water. By intercepting the Water 189 ray by a hand mirror at A, the white light can be thrown upon the screen, as a contrast to the greenish-blue tint. Aitkin has recently shown, that the presence of extremely finely divided suspended matters in water will give to the liquid the appear- ance of a blue colour. Thus, in tanks where water is being softened by the addition of milk of lime, after the bulk of the precipitated chalk has settled, and only the finest particles still remain suspended in the liquid, it is often noticed that the water appears to have a rich blue colour. The wonderful blue colour of the waters of many of the Swiss lakes is probably due in part to this optical phenomenon, as well as to the intrinsic colour of the water. When a mass of pure snow, such as falls in high mountainous regions, is broken open in such a way that the light is reflected from side to side of the small crevice, the true greenish-blue colour of the water is very manifest. 7 Water is compressible to only a very slight extent ; thus, under an additional pressure of one atmosphere, 1000 volumes of water become 999.95 volumes. Small as this compressibility is, it exerts an important influence upon the distribution of land and water upon the earth. It has been calculated, that owing to this compression, where the ocean has a depth of six miles, its silrface is lower by 620 feet than it would be if water were absolutely non-compressible • and calculated from the average depth of the sea, its average level is depressed 1 1. 6 feet - Th e effect of this depression of the sea-level is that 2,000,000 square miles of land are now uncovered, which would otherwise be submerged beneath the ocean. Water is an extremely bad conductor of heat. A quantity of water contained in a tube held obliquely, may be boiled by the application of heat to the upper layers, without appreciably affecting the temperature of the water at the bottom ; a fragment of weighted ice sunk to the bottom will remain for a long time unmelted, while the water a few inches above it is vigorously boiling. This low conductivity for heat is shared in common by all liquids that are not metallic. Indeed, Guthrie has shown, that water conducts heat better than any other substance which is liquid at the ordinary temperature, with the exception of mercury. Steam.— Under a pressure of 760 mm., water boils at ioo° (seep. 1 12), and is converted into a colourless and invisible gas, or vapour. The visible effect that is observed when steam is allowed to issue into the atmosphere, is due to the condensation of the steam in the form of minute drops of water. What is popularly 190 Inorganic Chemistry called steam is in reality, therefore, not steam, but an aggrega- tion of small particles of liquid water. The invisibility of steam is readily demonstrated by boiling a small quantity of water in a capacious flask ; as the steam issues from the neck it condenses in contact with the cool air and presents the familiar appear- ance, but within the flask it will be perfectly transparent and invisible. Ice. — At a temperature of o° water solidifies to a transparent crystalline mass. In the act of solidification, the water expands by nearly ^th of its volume, 10 volumes of water become 10.908 volumes of ice : solid water, therefore, is specifically lighter than liquid water, and floats upon its surface. Water in this respect is anomalous, for in the case of most other substances, the solid form is denser than the liquid. The disruptive force exerted by water at the moment of freezing, is the cause of the bursting of pipes and other vessels containing water during winter ; and it is also an important factor in the economy of nature, in the disintegration of rocks and of soil. Under certain conditions, water may be cooled many degrees below o° without solidification taking place. Thus, if a small quantity of water contained in a vacuous tube be care- fully cooled without being subjected to vibration, its temperature may be lowered to -15 0 without it solidifying; a slight shock, however, at once causes it to pass into the solid state, when its temperature instantly rises to o° (see p. 118). Although the exact temperature at which water freezes is liable to uncertainty from this cause, the point at which ice melts is, under ordinary cir- cumstances, constant, viz., o°. Under increased pressure ice will melt at temperatures below o° ; thus Mousson found that, under a pressure of 13,000 atmospheres, ice melted at — 18 0 . The melting-point of ice is lowered by about 0.0074° by each additional atmosphere of pressure (see p. 119). Between the temperatures of +4 0 and ioo°, water follows the ordinary laws that govern the expansion and contraction of liquids due to change of temperature ; if water be cooled from 100 , it gradually contracts until the temperature reaches 4 0 . Between this point and o° it forms a remarkable exception to the general law, for, when cooled below 4 0 , it slowly expands instead of con- tracting, and continues expanding until o° is reached, when it solidifies. At 4 0 , therefore, water expands whether it be heated or cooled ; consequently, at this point it is denser than at any other temperature. This temperature is known as its point of Water 19 (The most accurate observations fix the exact maximum density, point at 3-945°.) The following table shows the change of volume suffered bv water on being heated from o° to 8 ° 7 1. 000000 volumes at o° becomes °-9999 I 5 „ +2 0 o-99987o „ 4 ° 0.999900 „ 6° „ 1. 000000 „ 8° One cubic centimetre of water, measured at its point of maxi- mum density and at 760 mm., is the unit of weight of the metrical system, and is called a gramme. It is also at this temperature, that water is taken as the unit for comparison of the densities of other liquids and of solids • thus, when it is stated that the density, or specific gravity, of ciamond is 3.5, it is meant that diamond is 3.5 times as heavy density^ 1131 ^ ° f Wa ‘ er measured at its P oint of maximum The fact that water has a point of maximum density remote from Its ireezmg-pomt, is one of far-reaching consequences in the opera- tions of nature. ^ of rim 3 " water ’ such as a lake > » ^posed to the influence a cold wind, the superficial layer of water is cooled, and thereby becoming specifically denser, it sinks to the bottom and exposes a fresh surface. This in its turn has its temperature lowered, and in like manner falls to the bottom. A circulation of the water in this way is set up, until the entire mass reaches a temperature of 4 0 At this point the further cooling of the surface-layer causes expan- sion instead of contraction, and the colder water becoming speci- fically lighter now floats upon the top, where it remains until it “ ngea ! 5 ' , Ifwa ! er continued to contract as its temperature was whPl C< h H 6 °7 i ’ the c,rculat0r y motion w °n'd Continue until the whole body of the water was cooled to 0°, when solidification of the entire mass would take place. The reason that certain very deep waters seldom or never freeze, is because the duration of the cold ' , not long enough to bring the temperature of the entire mass upon th W e a surface Wn *° 4 ’ UntU ** * effeCted ’ n ° ice «*» ^rm The Solvent Power of Water-Water is possessed of more ig 2 Inorganic Chemistry general solvent powers than any other liquid ; that is to say, a larger number of substances are dissolved by water than by any other liquid. The solvent action of water upon gases, liquids, and solids, in so far as it is shared by other liquids, has been dealt with under the General Properties of Liquids (Part I., chap. xiii.). Water of Crystallisation. — When solid substances are dis- solved in water, and the water afterwards evaporated, the dissolved substance is frequently deposited in definite crystalline shapes. Many salts owe their crystalline nature to the fact, that a certain number of molecules of water have solidified along with molecules of the salt, each molecule of the salt being associated with a defi- nite number of molecules of solid water. The water molecules must be regarded as having entered into a feeble chemical union with the salt molecule, but a union which is of a somewhat diffe- rent order from that which holds together the atoms of oxygen and hydrogen in the water molecules, or the atoms composing the salt in the salt molecule (see p. 65). Thus, copper sulphate crystallises associated with five molecules of water, CuS0 4 ,5H 2 0 ; magnesium sulphate with seven, MgS0 4 ,7H 2 0. Water so associated with crystals is known as water of crystallisation, and the compound is called a hydrate. - Many salts are capable of crystallising with more than one defi- nite number of molecules of water, depending upon the temperature at which the crystallisation takes place : thus sodium carbonate, crystallised at the ordinary temperature, has the Na 2 CO 3 ,10H 2 O ; while at temperatures between 30 and 5° the sMt that is deposited contains 7 molecules of water, Na 2 C03,7H 2 0. Sodium chloride, crystallised from solution at -7 , has the compo- sition NaCl,2H 2 0 ; while the crystals that are deposited at 23 contain 10 molecules of water, NaCl,10H 2 O. In such cases as these, the particular crystalline form of the salt differs with the different degrees of hydration. Salts containing water of crystallisation, which are deposited from solution at temperatures below 0°, are sometimes termed cryohydrates. Many crystalline salts, when exposed to the air, lose elt her some or all of their water of crystallisation, and in so doing, lo particular geometric form. Thus, the salt, Na 2 CO 3 ,10H 2 O (ordinary washing soda), when freely exposed, gradually loses its crystalline form and falls down to a soft white powder, which consists of smaU crystals of another form, having the composition Na^Oaj^O. This process is known as efflorescence, the crystals being said Water 193 effloresce. Other crystals undergo exactly the reverse change ; they combine with moisture from the air, and pass into other crystalline forms containing more water, of crystallisation, or in some cases they absorb sufficient moisture to cause them to liquefy. Such substances are said to deliquesce. This property of certain salts is made use of for withdrawing traces of water from either liquids or gases. Thus, such a liquid as ether may be freed from dissolved water, by adding to it copper sulphate containing one molecule of water of crystallisation, CuS0 4 ,H 2 0 ; this compound takes up water and passes into CuS0 4 ,5H 2 0, and thereby has the effect of drying the ether. Gases in the same way are frequently dried by being passed through tubes containing calcium chloride from which the water of crystallisation has been removed. This substance absorbs water with avidity, passing into the hydrated salt CaCl 2 ,6H 2 0. The^ characteristic colours of certain salts are in many cases dependent upon the amount of water of crystallisation they contain. Thus cobalt chloride, CoC 1 2 ,6H 2 0, is a pink salt. If it be gently heated to 120 it loses its water, and becomes CoCl 2 , which has a rich blue colour. Solutions of this salt have been employed for the so-called sympathetic inks. The faint colour of the pink salt renders words written upon paper with its dilute solution prac- tically invisible ; but on warming the paper, and thereby expelling the water from the salt, the written characters appear in a blue colour, which again disappears as the salt is allowed to rehydrate itself by exposure to the air. One of the most striking examples of this change of colour resulting from varying proportions of water of crystallisation, is seen in the salt magnesium platino-cyanide, which crystallises under ordinary circumstances as a bright scarlet salt with seven molecules of water, MgPt(CN) 4 ,7H 2 0. When this salt is heated to about w° it loses two molecules of water, and is converted into a canary- yellow salt Mg p t (c N )4 ,5 H 2 °. If the temperature be raised To too the yellow salt becomes white by the loss of three more mole- cules, the composition of the white salt being MgPt(CN)„,2H,0. a d h< h 3 *,° Utl0 “ 0f the salt is carefully evaporated to dryness in a dish and gently warmed, these colour changes will be rendered evident ; and upon exposing the dried and white residue to the air, Trn y „ g T a breath ^ mt0 the dish > ‘he salt rehydrates itself, and of water mt ° CnmSOn com P° und having seven molecules Many salts can have their combined water withdrawn by power- ! Inorganic Chemistry ful dehydrating agents ; thus, if a crystal of copper sulphate (“ blue vitriol,” CuS0 4 ,5H 2 0) be immersed in strong sulphuric acid, the acid abstracts four out of the five molecules from the salt, leaving the nearly white salt CuS0 4 ,H 2 0 ; or when alcohol is added to a solution of cobalt chloride, or to crystals of the salt, CoC1 2 ,6H 2 0, the alcohol abstracts water, and the solution becomes blue. When salts containing water of crystallisation are heated, it frequently happens that a portion of the water is more easily parted with than the remainder. Thus copper sulphate, CuS0 4 ,5H 2 0, when heated to ioo°, parts with four molecules of water, leaving the salt CuS0 4 ,H 2 0 ; and in order to drive off this one remaining mole- cule, the temperature must be raised above 200 °. Zinc sulphate (or white vitriol), ZnS0 4 ,7H 2 0, in like manner loses six molecules^ of water at ioo°, but retains the seventh until a temperature of 240 ° is reached. In order, therefore, to distinguish between the water that is more firmly held and that which is readily parted with, the term water of constitution is frequently applied to the former, and the fact is sometimes expressed in notation in the following manner CuS0 4 H 2 0,4H 2 0 ; ZnS0 4 H 2 0,6H 2 0. Natural Waters. — On account of the great solvent powers of water, this compound is never found upon the earth in a state of absolute purity ; even rain, as it falls in regions far removed from the dirty atmosphere of towns, not only dissolves the gases of the atmosphere, but also small quantities of those suspended matters which are always present in the air. As soon as the rain reaches the earth, the water at once exerts its solvent action upon the mineral matter constituting the portion of the earth’s crust over which it flows, and through which it percolates, and the liquid is rapidly rendered less and less pure as it travels on its course to lake or ocean. Natural waters may be broadly divided into two classes, based upon the amount of dissolved impurities they contain. If the sub- stances in solution are present in excessive quantities, or to such an extent as to be perceptible to the taste, the water is said to be a mineral water; while, on the other hand, waters that are not so rich in dissolved impurities are known as fresh waters. Mineral Waters.— The most exaggerated examples of mineral waters are to be found in sea water, and in the waters of certain lakes, which, having no outlet, are fulfilling the purpose of enormous Natural Waters l 9S evaporating basins, in which the waters that flow into them are undergoing evaporation and therefore concentration; such for example as the salt lakes of Egypt, the Elton lake in Russia’ and the Dead Sea. In waters of this description the total quantity of dissolved solid matter is very considerable, and, as in the case of the Dead Sea, is often deposited in crystalline masses round the shores of the lake. The following table gives the total amount of dissolved saline matter contained in 1000 grammes of certain of these waters Irish Sea 86 Mediterranean Sea 4o o J?, ead Sea 22S.S7 Elton Lake 271 .43 nf rtlV v‘?^ XamP ! e ° f 3 Sea Water ’ the com P os 'tion of the water Br,tlsh Channel m ay be quoted ; 1000 grammes of this water con 13,1 n — Sodium chloride . 27.059 3.666 . . 2.296 • • 1.406 • . 0.766 Magnesium chloride Magnesium sulphate Calcium sulphate . Potassium chloride . Calcium carbonate . Magnesium bromide 0-033 * • 0.029 Water .... 35.255 964.745 1000.000 Passing from these highly concentrated mineral waters, we find a large number of spring waters which are classed as mineral, not because the total quantity of foreign matter in solution is excessive ut rather because they contain an abnormally large proportion of sullar 01 ^ S m Sta r S - ThUS ’ ' arge ^amities* of* magnesium sulphate, and chloride, are found in such springs as those at Epsom and Fnednchshall. Others are found to contain consider- Zr nhUeS °I f, dium SUlphate and s0dium carbonate ; while hose known as chalybeate waters, contain ferrous carbonate in solu- . 1 >,1 Pnng waters that are charged with unusual quantities of u h S T’ ^ ' lkeWiSe P ' aCed in the cate £ory of mineral waters such as the waters of Apollinaris and Seltzer, containing large quantities of carbon dioxide ; and the sulphur springs at Harrogate Inorganic Chemistry 196 and Aachen, which hold in solution sulphuretted hydrogen as well as alkaline sulphides. Fresh Waters.— The purest form of natural water is rain water. The average weight of solid matter dissolved in rain water, collected in the country and in perfectly clean vessels upon which it exerts no solvent action, is found to be 0.0295 parts in 1000 parts of water. Collected in or near towns, rain water always contains a larger amount of dissolved impurities, such as nitrates, sulphates, ammo- niacal salts, and often considerable quantities of sulphuric acid : it is the acid nature of the rain that causes so much damage to stone buildings. , . The nature and extent of the contamination that rain water suffers after it has fallen, must obviously depend very largely upon geographical and geological circumstances, and therefore there are no special features that are distinctly characteristic of waters from rivers, lakes, or springs. Thus, the total solid impurity in 1000 parts of water from the river Dee at Aberdeen is 0.057, while that contained in the Thames is 0.30 parts. c , The water of Loch Katrine only contains 0.032 parts of solid matter dissolved in 1000 parts, while that of Elton lake contains a* m The^ame wide differences are also seen in spring waters from different geological strata. Spring waters from granite and gneiss rocks, contain on an average 0.059 parts of dissolved solid matter in 1000 parts, while those from magnesian limestone average as much as 0.665 parts. As a broad general rule, river waters are found to contain less solid matter in solution than spring waters, and these in their turn less than deep well waters. Thus, com- paring waters from different sources, and selecting only such samples as are known to be free from pollution from either sewage matter or other abnormal impurities, it will be seen that, with regard to the dissolved solid matter they contain, they fall in the following order — Total Solid Impurity Dissolved in 1000 Parts of Unpolluted Waters. Rain water (average of 39 samples) ■ -0295 Rivers and lakes (average of 195 samples) ■ -°9 7 Spring waters (average of 198 samples) . • -2820 Deep well waters (average of 157 samples) . 437» Natural Waters 197 Hardness of Water.— Certain of the salts that are very fre- quently present as impurities in natural waters, give to these waters the property that is known as hardness. The chief com- pounds that produce this effect, are the salts of lime and magnesia. he term hardness is applied to such waters, on account of the difficulty of obtaining a lather, with soap, in the ordinary process of washing. Pure soap may be regarded as a mixture of the sodium salts of certain fatty acids (oleic, stearic, palmitic, &c.) which are soluble in pure water. In the presence of salts of lime and magnesium the soap is decomposed ; and an insoluble curdy precipitate is formed, by the union of the fatty acid of the soap with the lime and magnesium of the salts. Until the whole of the hardening salts have in this way been thrown out of solution, no lather can be obtained, and the soap is useless as a cleansing agent • but as soon as this point is reached, the addition of any further quantity of soap at once raises a lather on the water, and the soap is capable of acting as a detergent. This process of precipitating the salts of lime and magnesium is known as softening, and in this instance the water is softened at the expense of the soap. Hard waters often become less hard after being boiled for a short time, and this hardness which is so removed is termed the temporary hardness. The degree of hardness which the water still possesses after prolonged boiling, is distinguished by the term permanent hardness. The diminution of the total hardness of a water by boiling, is due to the fact that the soluble acid carbonates o lime and magnesium are decomposed during this process, into water, carbon dioxide (which escapes as gas), and the practically insoluble normal carbonates of these metals ; thus, in the case of the lime salt — CaH 2 (C0 3 ) 2 = H 2 0 + C0 2 + CaC0 3 . When such a water is boiled, the calcium carbonate is thrown down as a white precipitate, which gradually collects upon the bottom o the containing vessel. The “ furring ” of kettles, and the forma- tion of calcareous deposits in boilers, is largely due to this cause. n the case of waters that are highly charged with calcium car- bonate, held m solution by dissolved carbonic acid, this deposition of calcium carbonate may even take place at the ordinary tempe- rature, owing to the diffusion of the dissolved carbon dioxide into f„n T . • f th ’ S Way that those remarl «*ble, and often beauti- fully fantastic, formations, known as stalactites, have been produced 198 Inorganic Chemistry in certain subterranean caves. Water charged with the soluble calcium carbonate, in slowly dropping from the roof of such a cave, loses a portion of its dissolved carbon dioxide, and, in consequence, deposits a certain amount of the calcium carbonate which was in solution. Each drop, as it slowly forms, adds its little share of calcium carbonate to the deposit, which thereby gradually grows, much as an icicle grows, as a dependent mass called a stalactite. Whether the water that drops from the stalactite has deposited the whole of its calcium carbonate, will depend largely upon the time occupied by each drop in gathering and dropping ; if, as often happens, the whole has not been precipitated, the remainder is deposited upon the floor of the cave, and a growing column of calcium carbonate, called a stalagmite, gradually rises from the ground until it ultimately meets the stalactite. Clark’s Process for Softening Water.— Waters whose hard- ness is due to the presence of the carbonates of lime and magne- sium, can be deprived of their hardness by the addition to them of lime. 5 The amount of hardness is first estimated, and such an amount of milk of lime is then added as is demanded by the follow- ing equation : — CaH 2 (C0 3 ) 2 + CaO = H 2 0 + 2CaC0 3 . In this way, the soluble lime salt is converted into the insoluble normal carbonate, which settles to the bottom of the tank. The salts, which are mainly instrumental in causing the per- manent hardness, are the sulphates of lime and magnesium. The degree of hardness and its particular order, that is, whether tempo- rary or permanent, will obviously be determined entirely by the particular geological formation from which the water is derived. Potable Waters. — Undoubtedly the most important use to which water is put, is its employment as an article of food to man, and since it has been proved beyond dispute that many virulent diseases, such as cholera, typhoid fever, and others are propagate through the medium of drinking-water, it becomes a matter of the greatest sanitary importance that the waters supplied for this pur- pose should be as pure as possible. Excepting in very rare in- stances, where poisonous mineral matters accidentally gain access to drinking-water (as for example, in the case of certain waters which are capable of attacking, and to a slight extent dissolving, the lead of the pipes through which they may be passed), the solid Natural Waters 199 matters that are usually found in waters are not injurious to health. he living germs or bacilli, through whose agency zymotic diseases are caused, cannot be detected in a sample of water by any direct chemical or microscopical analysis. A specimen of pure distilled water might be artificially contaminated with such organisms so as to constitute it a most virulent poison, and still chemical analysis would fail to detect the danger, and the water would be pronounced pure. Chemical analysis can, however, reveal the presence of excrementitious matter, and also of the characteristic products re- sulting from its decomposition : it can with certainty detect in the water the evidence of recent contamination with sewage matters and it can also, with considerable precision, trace the evidences of its having been so contaminated at an earlier stage of its history. It cannot, however, distinguish between pollution with healthy and with infected excreta, and therefore it is necessary to regard *with the greatest suspicion any water to which sewage has at any time gained access. Waters that are made use of for drinking purposes may be classified in the following order : ! i. Spring water. 2. Deep well water. 3. Mountain rivers and lakes. Suspicious -f Stored ra i n water. ( 5. Surface water from cultivated land. Dangerous j ^* ver water to which sewage gains access. ( 7. Shallow well water. HYDROGEN PEROXIDE. Formula, H 2 0 «j. Occurrence.—' This compound is occasionally found in small quantities in the atmosphere, and also in dew and rain. Mode s of Formation.— ( 1.) Hydrogen peroxide is produced in small quantities during the burning of hydrogen in the air. If j jet of burning hydrogen be caused to impinge upon the surface of water, the temperature of which is not allowed to rise above 20° peroxid^* ^ ^ ^ t0 C ° ntain hydrog e " See “Chemical Lecture Experiments,” new ed.. p. 74. 200 Inorganic Chemistry ( 2 .) This compound is also produced by the decomposition of barium peroxide by carbonic acid. For this purpose a stream of carbon dioxide is passed through ice cold water, into which from time to time small quantities of barium peroxide are stirred. Barium carbonate is precipitated, and a dilute aqueous solution of hydrogen peroxide is obtained — Ba0 2 + H 2 C0 3 = BaCO s 4- H 2 02 - ( 3 .) Barium peroxide may be decomposed by either hydrochloric, sulphuric, silicofluoric, or phosphoric acid. Whichever acid be employed, the barium peroxide, previously mixed with a small quantity of water, is added gradually to the acid ; which, in the case of either hydrochloric or sulphuric acid, should be diluted with from five to ten times its volume of water. The temperature of the mixture is not allowed to rise above 20 °. Thus, in the case of hydrochloric acid — Ba0 2 4 - 2HC1 - BaCl 2 4- H 2 0 2 , the soluble barium chloride is removed by the addition of sulphuric acid, whereby barium sulphate is precipitated, and hydrochloric acid formed — BaCl 2 4 - H 2 S0 4 = BaS0 4 4- 2HC1. The hydrochloric acid may be removed by adding a solution of silver sulphate, which precipitates silver chloride, leaving sulphuric acid in solution — 2HC1 4 - Ag 2 S0 4 = 2AgCl 4- H 2 S0 4 . And, lastly, the free sulphuric acid is withdrawn by the addition of barium carbonate — H 2 S0 4 4- BaC0 3 = BaS0 4 + H 2 6 4- C0 2 . When sulphuric is employed for the decomposition of barium peroxide, the crystallised, or hydrated peroxide (Ba0 2 ,8H 2 0), is most advantageous for the purpose. This salt, made into a paste with water, is gradually added to the diluted and cooled acid, until the acid is nearly but not quite neutralised. The slight excess of acid is removed by the addition of the exact quantity of barium Hydrogen Peroxide 20 1 hydroxide (baryta-water) necessary to neutralise it, and the insoluble barium sulphate is removed by filtration. On a large scale silico- fluoric acid, or phosphoric acid, is usually employed, preferably the latter as it is found that small quantities of free phosphoric acid in hydrogen peroxide greatly retard its decomposition. (4.) Hydrogen peroxide is also readily obtained by decomposing potassium peroxide by means of tartaric acid. The potassium peroxide is added to a cooled strong aqueous solution of tartaric acid, when potassium tartrate separates out, and an aqueous solu- tion of hydrogen peroxide is obtained. (5.) When small quantities of hydrogen peroxide are required for the purpose of illustrating its properties, it is most conveniently obtained by adding sodium peroxide to dilute and well-cooled hydrochloric acid, whereby sodium chloride and hydrogen per- oxide are formed, both of which remain in solution — Na 2 0 2 + 2HC1 = 2NaCl + H 2 0 2 . (6.) Hydrogen peroxide is formed in considerable quantity when ozone is passed through ether floating upon water. Probably a peroxidised compound of ether is first produced, which is then decomposed by the water. This production of hydrogen peroxide may readily be demonstrated by placing a small quantity of water and ether in a beaker, and suspending into the vapour a spiral of platinum wire which has been gently heated. The combustion of the ether vapour upon the wire, whereby the latter is maintained at a red heat, is attended with the formation of ozone, and this acting upon the ether, as already described, results in the pro- duction of hydrogen peroxide, which may be detected in solution in the water. (7-) In small quantities, hydrogen peroxide is produced when moist ether is exposed to the action of oxygen, under the prolonged influence of sunlight. Properties.— The dilute aqueous solution of hydrogen peroxide obtained by the foregoing methods, is concentrated by evaporation over sulphuric acid in vacuo. In the pure condition it is a colour- less and odourless, syrupy liquid, having an extremely bitter and metallic taste. The specific gravity of the liquid is 1.4532 The substance is extremely unstable, giving up some of its oxygen even at temperatures as low as -20°, and decomposing with explosive violence when heated to 100°. Hydrogen peroxide bleaches 202 Inorganic Chemistry organic colours, but less rapidly than chlorine. When placed upon the skin it destroys the colour, and gives rise to an irritating blister. When diluted with water, and especially if rendered acid, the compound is far more stable, and in this condition may be preserved at the ordinary temperature for a considerable length of time. When such an aqueous solution is strongly cooled, it deposits ice, and in this way, by the removal of the frozen water, the solu- tion may be concentrated ; hydrogen peroxide itself remains liquid at - 30 °. When heated the solution is decomposed into water and oxygen — H 2 0 2 = H 2 0 + O. Owing to the readiness with which hydrogen peroxide gives up the half of its oxygen, and is converted into water, its properties are generally those of a powerful oxidising agent. It liberates iodine from potassium iodide ; it converts sulphurous acid into sulphuric acid, and oxidises lead sulphide into lead sulphate. Its action upon lead sulphide is made use of, in restoring something of the original brilliancy to oil paintings that have become dis- coloured. The “ white-lead ” used in oil paints is gradually con- verted into lead sulphide when such paintings are exposed to air, especially the air of towns, which is liable to contain small quantities of sulphuretted hydrogen. Lead sulphide being black, the picture slowly assumes a uniformly dark colour, until it is finally quite black. When such a discoloured picture is washed over with dilute hydrogen peroxide, the black sulphide is oxidised into the white lead sulphate— PbS + 4H 2 0 2 = 4H 2 0 + PbS0 4 . This compound is employed for bleaching articles that would suffer injury by the use of other bleaching agents, such as ivory, feathers, and even the teeth. Hydrogen peroxide is also capable of oxidising hydrogen, when that element in the nascent condition is brought in contact with this compound. Thus, if a dilute acidulated solution of hydrogen peroxide be subjected to electrolysis, oxygen will be evolved from the positive electrode, but no gas will be disengaged at the negative pole. The hydrogen, in the presence of the hydrogen peroxide, is oxidised into water— H,0 2 + H 2 = 2H a O. 203 Hydrogen Peroxide Hydrogen peroxide, in many of its reactions, appears to act as a deoxidising agent ; thus, manganese dioxide in contact with this substance is reduced to manganous oxide — Mn0 2 + H 2 0 2 = MnO + 0 2 + H 2 0. Similarly silver oxide is reduced to metallic silver with the evolution of oxygen — Ag 2 0 + H 2 0 2 = A g 2 + 0 2 + H 2 0. In like manner, when ozone is acted upon by hydrogen per- oxide, a reaction takes place exactly analogous to that with silver oxide, which will be the more obvious if the formula for ozone be written 0 2 0 instead of O s , thus— 0 2 0 + H 2 0 2 = 0 2 + 0 2 + H 2 0. Although, in a sense, these reactions may be regarded as reduc- ing, or deoxidising, actions, in essence they are not different from those which have been given as illustrative of the oxidising power of hydrogen peroxide. It will be seen that they all depend upon the readiness with which the compound parts with an atom of oxygen, but that in these latter cases the oxygen that is so given up is engaged in oxidising another atom of oxygen, contained in the other compound. Thus, in the case of silver oxide, its atom of oxygen is oxidised by the liberated oxygen from the hydrogen peroxide, and converted into the complete molecule of oxygen. By these reactions Brodie first demonstrated the dual, or di- atomic, character of the molecule of oxygen. When hydrogen peroxide is added to a dilute acidulated solution of potassium dichromate, a deep azure-blue solution is obtained (see Chromium), which affords a delicate test for this com- pound. To apply the test, the dilute hydrogen peroxide is shaken up with ether, and being soluble in this liquid, the ethereal layer which rises to the surface will contain nearly the whole of the peroxide ; a few drops of acidulated potassium dichromate are then added, and the mixture again shaken, when the ethereal liquid will separate as a blue layer. In this way, the presence of 0.00025 grammes of hydrogen peroxide in 20 c.c. of water can be detected. Hydrogen peroxide is decomposed by contact with many sub- stances which themselves do not combine with the oxygen ; thus 204 Inorganic Chemistry charcoal, finely divided palladium, platinum, mercury, and notably silver, when brought into hydrogen peroxide, determine its decom- position into water and oxygen, the rapidity of the action being increased if the liquid be made alkaline. The action is doubtless catalytic, although in all cases the exact modus operandi is not clearly understood. In the case of silver it is believed that silver oxide (perhaps peroxide) is first formed, and then decomposed, A g 2 + H 2 0 2 = H 2 0 + Ag 2 0 Ag 2 0 + H 2 0 2 = H 2 0 + 0 2 + Ag 2 . When hydrogen peroxide is added to solutions of the hydroxides of barium, strontium, or calcium, the peroxide of the metal is precipitated — Ba(HO) 2 + H 2 0 2 = 2H 2 0 + BaO* The ^compound is deposited in crystals having the composition Ba0 2 ,8H 2 0. ....... With the hydroxides of the alkali metals, the peroxide (which is soluble in water) may be precipitated by the addition of alcohol ; when in the case of sodium peroxide, crystals are obtained of Na 2 0 2 ,8H 2 0. Hydrogen peroxide is a useful antiseptic: it possesses the ad- vantages of being free from smell, without poisonous or injurious action upon the system, and of leaving as a residue, after having . furnished its available oxygen, only water. CHAPTER IV NITROGEN Symbol, N. Atomic weight = 14.01. Molecular weight = 28.02. History. — Nitrogen was discovered by Rutherford in 1772. He showed that when an animal is placed in a confined volume of air for some time, and the air afterwards treated with caustic potash, to absorb from it the carbon dioxide (“fixed air’ 5 ), there still remained a gas which was incapable of supporting either respira- tion or combustion. He called the gas mephitic air. Scheele was the first to recognise that this gas was a constituent of the air. Lavoisier applied the name azote to the gas, to denote its inability to support life. The name nitrogen , signifying the nitre-producer, was suggested by Chaptal, from the fact that the gas was a con- stituent of nitre. Occurrence. — In the free state nitrogen is present in the atmos- phere, of which it forms about four-fifths. Certain nebulas have been shown, by spectroscopic observation, to contain nitrogen in the uncombined condition. In combination, nitrogen is found in ammonia, in nitre (potassium nitrate), and in a great number of animal and vegetable compounds. Modes Of Formation.— ( 1.) Nitrogen is very readily obtained from the atmosphere, by the abstraction of the oxygen with which it is there mixed.* This is conveniently done by burning a piece of phosphorus in air, confined over water. The phosphorus in burning combines with the oxygen, forming dense white fumes of phosphorous pentoxide, which gradually dissolve in the water, and nitrogen remains in the vessel. The nitrogen obtained in this way is never quite pure, for the phosphorus becomes extinguished before the oxygen is entirely removed ; and also the gas will contain atmospheric carbon dioxide. (2.) Nitrogen in a purer state can be prepared from the atmos- * Experiments 254, 255, “ Chemical Lecture Experiments," new ed. 205 206 Inorganic Chemistry phere, by passing a stream of pure air over metallic copper contained in a combustion tube, and heated to redness in a furnace. The air is contained in a gas-holder, and is passed through two U -tubes, the first containing potassium hydroxide (caustic potash), in order to absorb the carbon dioxide ; and the second filled with fragments of pumice moistened with sulphuric acid, in order to arrest the aqueous vapour. The purified air, on passing over the heated copper, is deprived of the whole of its oxygen, cupric oxide, CuO, being formed, while the nitrogen passes on and may be collected. This gas contains small quantities of argon (p. 648). (3.) Oxygen is rapidly absorbed by a solution of cuprous chloride in hydrochloric acid ; a ready method, therefore, of obtaining nitrogen from the air, is to place a quantity of this solution in a stoppered bottle, and shake it up with the contained air. The colourless cuprous chloride solution quickly absorbs the oxygen, becoming dark in colour, and being converted into cupric chloride, the nitrogen of the air remaining in the bottle — Cu 2 Cl 2 + 2HC1 + O = H 2 0 + 2CuCl2. (4.) Nitrogen is obtained by heating a strong solution of ammo- nium nitrite in a flask, the salt splitting up into water and nitrogen — NH 4 N0 2 = 2H 2 0 + N 2 . In practice it is found more convenient to employ a mixture of ammonium chloride and sodium nitrite — NH 4 C1 + NaN0 2 = NaCl + 2H 2 0 + N 2 . (5.) By heating a mixture of ammonium nitrate and ammonium chloride, a mixture of nitrogen and chlorine is evolved ; the latter gas may be absorbed, by passing the mixture through either milk of lime, or a solution of sodium hydroxide — 2NH 4 N0 3 + NH 4 C1 = 5N + Cl + 6H 2 0. (6.) Nitrogen is also evolved when ammonium chromate, or a mixture of potassium dichromate and ammonium chloride, is heated — (NH 4 ) 2 Cr 2 0 7 = Cr 2 0 3 + 4H 2 0 + N 2 , or — K 2 Cr 2 0 7 + 2NH 4 C1 = Cr 2 O s + 2KC1 + 4H 2 0 4- N 2 . Nitrogen 207 (7.) When ammonia is acted upon by chlorine, it is decomposed the chlorine combining with the hydrogen to form hydrochloric acid, and the nitrogen being liberated — 2NH 3 + 3C1 2 = 6HC1 + N 2 . If the chlorine be passed into a strong solution of ammonia, the hydrochloric acid which is produced combines with the excess of ammonia, forming ammonium chloride ; thus— 8 NH 3 + 3C1 2 = 6NH 4 C1 + N* U ul', e J Ch ! 0rine ; after being washed b y P assin S through water, is bubbled through strong aqueous ammonia contained in a Woulf’s bottle. As each bubble of chlorine enters into the ammonia, the combination is attended by a feeble yellowish flash of light and a rapid stream of nitrogen is evolved. The nitrogen, which carries with it dense white fumes of ammonium chloride, should be scrubbed by being passed through a second bottle, filled with fragments of broken glass moistened with water, and it can then be collected over water in the ordinary way, as shown in Fig. 44.* I n prepar- ing nitrogen by this reaction, it is very necessary that the ammonia should be in considerable excess, otherwise there is liable to be formed the dangerously explosive compound of nitrogen and chlo- rine. See Nitrogen Trichloride. * Experiment 261 . 208 Inorganic Chemistry Properties.— Nitrogen is a colourless gas without taste 01 smell. It is slightly lighter than air, its specific gravity being 0.972 (air = 1). One litre of the gas at o° C. and 760 mm. weighs 14 criths, or 1.256 grammes. Nitrogen is only very slightly soluble in water, its coefficient of absorption at o° C. being 0.020346. Nitrogen will not burn, neither will it support the combustion of ordinary combustibles. It is not poisonous, but is incapable of sup- porting respiration. Nitrogen is one of the most chemi- cally inert substances known, com- bining directly, and with difficulty, with only a very few elements. Under the influence of the high temperature of the electric spark it can be made to unite directly with oxygen (see p. 210). Certain metals also com- bine directly with it, forming nitrides. Thus, when lithium or magnesium are heated in nitrogen, they form respectively NLi 3 and N 2 Mg 3 . This reaction may be conveniently shown by means of the apparatus seen in Fig. 45. A small quantity of powdered magnesium is placed in a hard glass tube, which is connected to a long narrow tube dipping into water, and a Fig. 45. stream of nitrogen is passed through. When the air is all displaced, the passage of the nitrogen is stopped, and the magnesium strongly heated. At a red heat the nitrogen will be rapidly absorbed, and the water will be seen to rise in the long tube. The critical temperature of nitrogen is - 146°, and when cooled to this point, a pressure of 35 atmospheres causes its liquefaction. Under ordinary atmospheric pressure, the liquid boils at —193 5 the gas, therefore, can be liquefied by the cold obtained by the rapid evaporation of liquid oxygen (see p. 76). CHAPTER V OXIDES AND OXY-ACIDS OF NITROGEN Nitrogen combines with oxygen, forming five oxides (i.) Nitrous oxide (hyponitrous anhydride) ( 2 .) Nitric oxide. . ' (3-) Nitrogen trioxide (nitrous anhydride) (40 Nitrogen peroxide . (5-) Nitrogen pentoxide (nitric anhydride) N 2 0. NO. N 2 0 3 . N 0 2 and N 2 0 4 . n 2 o 6 . Three oxy-acids of nitrogen three oxides, Nos. 1 , 3 , 5 : are known, corresponding to the Hyponitrous acid Nitrous acid . Nitric acid HNO. HNOj, HN0 3 . ~nnmP relat !° n in which these three acids stand to the ponding oxides may be seen by the following formate corre Hyponitrous anhydride JJ | O. Hyponitrous acid £ } O. Nitrous anhydride . j Q Nitrous add NOj Q Nitric anhydride . JJg» [ 0 . Nitric acid . J. Q 2 > H y NITRIC ACID. Formula, HN0 3 . Molecular weight = 62. 88 2 to Inorganic Chemistry (1776) its true nature was not known ; he showed that oxygen was one of its constituents, but as to its other components he was un- certain. Its exact composition was determined by Cavendish. Modes of Formation.— (1.) When an electric spark is passed through a detonating mixture of oxygen and hydrogen with which a certain quantity of air, or nitrogen, is mixed, the water that is produced by the union of the oxygen and hydrogen is found to contain nitric acid. This fact was first observed by Cavendish in the course of his investigations on the composition of water, when, owing to the accidental admixture of air with the mixed gases, oxygen and hydrogen, he found that the water resulting from the union was sometimes acid. The direct union of nitrogen and oxygen may be brought about by allowing a series of electric sparks to pass between platinum wires, in a confined volume of air, contained in a glass globe, as shown in Fig. 46. In a short time the air in the globe will become distinctly reddish in colour, owing to the formation of nitrogen peroxide. The rapidity of the formation of the red fumes will be greatly increased, by compressing the air within the globe by means of a small compression pump, as indicated in the figure. If a small quantity of water be introduced, and the contents of the globe shaken up, the red gas will be seen to dissolve in the water, which will then acquire an acid reaction, owing to the forma- tion of nitric acid. Similarly, when a jet of hydrogen is allowed to burn in air to which additional oxygen has been added, considerable quantities of nitrogen peroxide are formed. The hydrogen may be burnt from a jet, surrounded by a glass tube, as shown in Fig. 4 7 , into which oxygen can be passed by means of the small bent tube at the bottom. On holding a clean dry cylinder over the flame, sufficient of the products of combustion will collect in a few seconds to show the presence of nitrogen peroxide. (2.) Nitric acid is formed when nitrogenous animal matter under- goes slow oxidation in the air, in the presence of water and an alkali, the nitric acid combining with the alkali to form a nitrate. In this way nitrates are found in the soil, and from the soil often find their way into shallow well waters of towns. In hot and rain- less countries these nitrates are sometimes found as crystalline deposits on the surface of the soil, as in Chili and India. (See Potassium Nitrate.) (3.) Nitric acid is prepared by acting upon potassium nitrate Nitric Acid 2 ji 0 nitre-saltpetre ) with sulphuric acid. The nitre is placed in a glass retort, together with an equal weight of sulphuric acid, and the “ ^ The nitric add readily distils Uer, and insist, f n eC i ln a COOled receiver - The residue in ‘he retort consists of hydrogen potassium sulphate— kno 3 + h 2 so 4 = i and “"‘ains nitrogen perox.de in solut.on, which imparts to it a yellowish-red colour. To purify it, it is again distilled with an equal volume of sulphunc acid ; and the redistilled acid is deprived of the las traces b bb. S e r d pe r ide ° f by causing a stream of d "a" o bubble through .t while slightly warm. Nitric acid so prepared fIl C N‘t am aS M UCh 35 99-8 Per Cent of anhydrous acid, HNO,. Drltess 'T aCld 1S an artlCle 0f commerciaI manufacture. In this process potassium nitrate is replaced by the sodium salt, as being emnl e T r mat ’ and the Proportion of acid to sodium nitrate employed ,s arranged in accordance with the equation- 2NaNO s + H 2 S0 4 = Na 2 S0 4 + 2HNO,. 2 1 2 Inorganic Chemistry It will be seen that the whole of the hydrogen of the sulphuric acid has been replaced by the alkali metal, derived from two mole- cules of the nitrate, and that two molecules of nitric acid result. The reaction takes place in two stages ; in the first we have — NaN0 3 + H 2 S0 4 = NaHS0 4 + HNOg. And then, as the temperature is raised, the hydrogen sodium sulphate reacts upon a second molecule of the nitrate, thus — NaNOg + NaHS0 4 = Na 2 S0 4 + HNO s . The temperature necessary to effect this second stage, however, causes the decomposition of a certain quantity of the nitric acid itself, thus — * 2HN0 3 = H 2 0 + 2N0 2 + O. The retorts usually employed for the manufacture of this acid are large cast-iron cylinders, which are sometimes lined, either entirely or in part, with fireclay, and which are built into a furnace in such a manner as to allow of their being heated as uniformly as possible. The ends of the cylinders are closed by slabs of Yorkshire flag, securely cemented to the iron. The charge of sodium nitrate (Chili saltpetre) and sulphuric acid is introduced through a hole in one end, which is afterwards plugged up, and the vapours are carried off through an earthenware pipe ( c , Fig. 48), cemented through a hole in the other end, and connected to a series of earthenware pots, t>, in the manner shown in the figure. The last of these jars is connected with a tower, filled with coke, down which water is caused to percolate, and any peroxide of nitrogen which escapes condensation with the acid in the jars is thereby absorbed. Properties. — Nitric acid is a colourless liquid having a specific gravity of 1.53. It fumes strongly in the air, and has a peculiar and choking smell. It is extremely hygroscopic, absorbing moisture from the air with great readiness. Nitric acid is an intensely corrosive liquid : the strongest acid, when brought in contact with the skin, causes painful wounds, while in more dilute conditions it stains the skin, and other organic materials, a bright yellow colour. A quantity of strong nitric acid thrown upon sawdust causes it to burst into flame. When nitric acid is distilled it first begins to boil at 86°, at the same time it is partially decomposed into water, nitrogen peroxide, and oxygen ; the distillate, therefore, Nitric Acid 213 gradually becomes weaker, and the boiling-point gradually rises. This continues until a certain point is reached, when both the temperature of the boiling liquid and the strength of the distillate SXSTTh V“ the 0ther hand - a weak acid b ' -JMled. the distillate gradually increases in strength, until when the same point is reached, the boiling liquid has again the same temperature. 214 Inorganic Chemistry This constant boiling-point is 120.5 0 , and the distillate which comes over at that temperature contains 68 per cent, of HN 0 3 . Whatever the strength of the acid, therefore, on being boiled it loses either nitric acid or water until the strength reaches 68 per cent., and this liquid boils at 120° C. The specific gravity of this acid at 15 0 is 1.414. It was formerly supposed that the acid of this strength constituted a definite hydrate, but Roscoe has shown that the strength of the acid is purely a function of the pressure, for by varying the pressure under which the distillation is conducted, acids of various compositions can be caused to distil at a constant temperature. When nitric acid is mixed with water there is a rise in tempera- ture and a contraction in volume, the maximum effect being pro- duced when the mixture is made in the proportion of three molecules of water with one molecule of acid. Nitric acid is a powerful oxidising agent, on account of the readi- ness with which it parts with oxygen. Elements such as sulphur and phosphorus are oxidised into sulphuric and phosphoric acids ; arsenious oxide into arsenic acid ; and many protosalts are con- verted into persalts. It attacks a large number of metals, forming in many cases the nitrate. Its action upon metals is often of a complicated nature, and depends not only upon the particular metal, but also upon the strength of the acid, the temperature, and the presence of the saline products of the reaction ; thus, when nitric acid acts upon copper, the following reaction takes place — 3 Cu + 8HNO3 = 3 Cu(N 0 3 ) 2 + 4 H 2 0 + 2 NO. It is found, however, that as the amount of copper nitrate accu- mulates, the nitric oxide which is evolved is mixed more and more largely with nitrous oxide, N 2 0 , and even with nitrogen. Again, when dilute nitric acid acts upon zinc, nitrous oxide is produced, according to the following equation — 4 Zn + IOHNO3 = 4 Zn(NO a ) 2 + 5 H 2 0 + N 2 0 . When, however, strong nitric acid is employed, ammonia is formed, which combines with the excess of acid — 4 Zn + 9 HN 0 3 = 4 Zn(NO s ) 2 + 3 HX> + NH 3 . In some cases, as with copper and silver, the presence of nitrons acid (either as an impurity in the nitric acid, or as a first product of its attack upon the metal) is believed to be a necessary condition of the action. Nitric Acid 215 Owing to the strong oxidising properties of nitric acid, hydro- gen is rarely isolated by the action of metals upon this acid, the hydrogen which is displaced from the acid being converted into water. With magnesium, however, free hydrogen is evolved. The chief reactions of nitric acid may be broadly divided into three classes : — (1.) With metallic oxides its behaviour is in common with other acids. It exchanges its hydrogen for an equivalent quantity of the metal, forming a nitrate, with the elimination of water, e.g . — Ag 2 0 + 2HN0 3 = 2 AgN 0 3 + H 2 0. (2.) Reactions in which it acts as an oxidising agent ; as an example, its action upon iodine, which is converted into iodic acid, may be cited — 1 + 3HN0 3 =HI0 3 + H 2 0 + NO + 2N0 2 . (3.) Actions in which hydrogen in an organic compound is replaced by the elements N0 2 , with the elimination of H a O, no gas being evolved. The conversion of cotton- wool, or cellulose, Ci 2 H 2 o0 10 , into gun-cotton, or nitro-cellulose, C 12 H 14 O 10 (NO 2 ) 6 , is an illustration of this class of reactions — C 12 H 20 O 10 + 6 HNO 3 = 6H 2 0 + C 12 H 14 O 10 (NO 2 ) 6 . Nitric acid is without action upon the so-called noble metals, gold and platinum. Commercial nitric acid, which is of a reddish colour, is liable to contain many impurities : chlorine and iodic acid, derived from the Chili saltpetre ; iron, sulphuric acid, and sodium sulphate, carried mechanically over from the retorts ; and nitrogen peroxide, from the decomposition of the acid. From these it is purified by redistillation. Nitric acid is a monobasic acid ; the salts of which, known as the nitrates, are for the most part readily soluble in water, and crystallise in well-defined forms. They are all decomposed at a high temperature, evolving oxygen and nitrogen peroxide, or oxy- gen and nitrogen, leaving an oxide of the metal. The presence of a nitrate in solution is easily recognised by the following characteristic test. A solution of ferrous sulphate is first added to the solution containing the nitrate, and concentrated sul- 216 Inorganic Chemistry phuric acid is then cautiously poured down the side of the test- tube, held in a sloping position, so as to fall to the bottom without mixing with the solution. The sulphuric acid acting upon the nitrate, liberates nitric acid ; this is reduced by the ferrous sulphate to nitric oxide, which, dissolving in the ferrous sulphate, forms a brown-coloured solution at the point where the two layers of liquid meet. (See Nitric Oxide.) When nitric acid is added to hydrochloric acid, a mixture is obtained which is known by the name of aqua regia. This name was applied to it by the alchemists on account of its power of dis- solving gold. Aqua regia is used in the laboratory for dissolving gold, platinum, and certain ores. Its solvent power depends upon the free chlorine which is evolved from the mixture — HNO s + 3HC1 = 2H 2 0 + NOC1 + Cl,. NITROGEN PENTOXIDE {Nitric Anhydride). Formula, N a 0 5 . Molecular weight = 107.8. Modes of Formation. — (i.) By withdrawing from nitric acid the elements of water, by means of phosphorus pentoxide — 2HN0 3 + p 2 o 6 = 2HP0 3 + n 2 o 6 . For this purpose the strongest nitric acid is cautiously added to phosphorus pentoxide in a cooled retort, in the proportion de- manded by the equation ; the mixture being made as far as possible without rise of temperature. The pasty mass is then gently heated, when the nitrogen pentoxide distils over, and, if collected in a well- cooled receiver, at once crystallises. (2.) The method adopted by Deville, who discovered this com- pound (1849), was by passing dry chlorine over dry silver nitrate contained in a U-tube, which was kept at the desired temperature by being immersed in a water-bath. The following equation ex- presses the final result of the action — 2AgN 0 3 + Cl 2 = 2AgCl + N 2 0 6 + O. Properties. — Nitrogen pentoxide is a white solid substance, crystallising in brilliant prismatic crystals, which melt at 30° with partial decomposition. Between 45 0 and 50° it undergoes rapid decomposition, evolving brown fumes. It is a very unstable Nitrogen Peroxide 217 compound ; when suddenly heated it decomposes with explosive violence, and even at ordinary temperatures decomposition slowly takes place. It absorbs moisture rapidly, and when thrown into water it dissolves with the evolution of great heat N 2 0 6 + H 2 0 = 2HN0 3 . When nitrogen pentoxide is gradually mixed with nitric acid a compound is formed having the composition 2N„0 5 ,H 2 0 • which separates, on cooling, as a definite crystalline hydrate NITROGEN PEROXIDE. Formula, N0 2 and N 2 0 4 . Molecular weight = 45-9= and 91.84. Density = 22. 96 and 45. 92. Modes of Formation.— ( 1 .) This compound may be prepared by mixing one volume of oxygen with two volumes of nitric oxide and passing the red gas so obtained through a tube surrounded by a freezing mixture— 2NO + 0 2 = 2N0 2 . ( 2 .) rhe nitrates of certain metals, when heated, are decomposed mto nitrogen peroxide oxygen, and an oxide of the metal ; thus, if dry lead nitrate be heated in a retort, and the gaseous products of decomposition are conducted into a U-tube placed in a freezing mixture, the nitrogen peroxide collects in the tube Pb(N 0 3 ) 2 PbO + N 2 0 4 + O. (3-) When arsenious oxide is gently warmed with nitric acid a mixture of n.tnc oxide, NO, and peroxide, NO, is evolved, and if t^a b ue° U l S iafod tUr o PaSSed thr ° Ugh “ C °° led tUbe ’ h condenses to a blue liquid. On passing a stream of oxygen through this hquid it loses its blue colour, and is converte?intoTS^ liquid, which consists of nitrogen peroxide. Properties.— At low temperatures nitrogen peroxide is a colour- less crystalline compound. It melts at - 9 0 , but requires a tern- above hs^ IT* 3 ~ 3 °° f ° S ° Hdify At a tem P er ature slightly above its melting-point the hquid begins to acquire a pale yellowish nh winch rapidly deepens, until at the ordinary temperature it is a full orange colour. The liquid boils at M -, and gives a vapour 2 1 6 Inorganic Chemistry having a reddish brown colour. The colour of the vapour also becomes deeper as its temperature is raised, until at 40° it is a dark chocolate brown, and almost opaque. On allowing the vapour to cool the reverse changes take place. This change of colour, as the temperature rises, is accompanied by a steady change in the density of the gas, as will be seen from the table : Temperature. Density, H = 1. Percentage of N 0 2 Molecules. 26 . 7 ° 38.3 20.00 60. 2° 30.I 52.04 100.1° 24 - 3 ' 89.23 i 35 -o° 23.1 98.69 140.0° 22.96 100.00 The density required by the formula N 2 0 4 is 45-9 2 > whlle that demanded by the formula N 0 2 is 22.96 ; hence as the temperature rises a process of dissociation goes on, in which N 2 0 4 molecules are broken down into molecules of the simpler composition. At 140° this process is complete, and the gas is entirely re- solved into N0 2 . It is believed that at low temperatures, nitrogen peroxide has the composition represented by the formula N 2 0 4 , but that dissociation begins to take place even during the state of liquidity, as indicated by the gradual change of colour ; and there- fore at temperatures between the boiling-point of the liquid, viz., 22 0 , and 140°, the gas consists of mixtures of molecules of N 0 2 and N 2 0 4 . The calculated percentage of N0 2 molecules, which the gas contains at the temperatures at which the above densities are taken, are given in the third column. Nitrogen peroxide is decomposed by water. At low tempera- tures, and with small quantities of water, nitric and nitrous acids are the products of the action, thus — N 2 0 4 + H 2 0 = HN0 3 + HNO2. At the ordinary temperature, and with an excess of water, the following reaction takes place — 3N0 2 + H 2 0 = 2HN0 3 + NO. Gaseous nitrogen peroxide is incapable of supporting the com- bustion of a taper. Phosphorus, when strongly burning and plunged into the gas, continues its combustion with brilliancy, Nitrous Acid 219 the temperature of the burning phosphorus being sufficiently high to effect the decomposition of the gas. Nitrogen peroxide is a suffocating a^id highly poisonous gas, and even when largely diluted with air rapidly prodaces headache and sickness. Nitrogen peroxide unites directly with certain metals, giving rise to a re- markable series of compounds, to which the name nitro-metals, or metallic nitroxyls, may be given (Sabatier and Senderens).* Thus, when the vapour of nitrogen peroxide is passed over metallic copper (obtained by the reduction of copper oxide in a stream of hydrogen), the gas is rapidly absorbed by the metal with considerable rise of temperature, and a solid brown compound is formed. This substance is the copper-nitroxyl, and its composition is ex- pressed by the formula Cu 2 N 0 2 . Copper-nitroxyl is a fairly stable compound, and is unacted upon by dry air. It is decomposed by water and by nitric acid, hence in its preparation care must be taken to free the nitrogen peroxide from these substances. At a temperature of about 90° copper-nitroxyl is decomposed into copper and nitrogen peroxide. If, therefore, a quantity of the compound be sealed up in a bent glass tube, and the empty limb of the tube be immersed in a freezing mixture while the compound is gently warmed, the nitrogen peroxide which is evolved will be condensed in the cold portion of the tube. Similar compounds are formed with the metals cobalt, nickel, and iron. Nitrous Acid, HN0 2 . — This substance is not known in the pure state. Even in dilute aqueous solution it rapidly decomposes into nitric acid, nitric oxide, and water — 3HN0 2 = HN0 3 + 2NO + H 2 0. The solution of this acid sometimes acts as a reducing agent, taking up oxygen from such highly oxidised compounds as per- manganates, or chromates, and passing into nitric acid — hno 2 + O = hno 3 . Under other conditions it exerts an oxidising action, as when it bleaches indigo, or liberates iodine from potassium iodide, being itself reduced to nitric oxide and water, with the elimination of oxygen — 2HN0 2 = 2NO + H 2 0 + O. The salts of nitrous acid, viz., the nitrites , are stable compounds. 1 he nitrites of the alkalies are best prepared by carefully heating * Bulletin de la SociiU Chimique , September 1893. 220 Inorganic Chemistry the nitrates ; thus, when potassium nitrate is fused, it parts with oxygen, and is transformed into potassium nitrite — kno 3 = KN 0 2 + O. At a higher temperature the nitrite is also decomposed. Nitrites are decomposed by dilute acids, evolving brown vapours, and in this way are at once distinguished from nitrates. Nitrogen Trioxide. — There is considerable doubt as to the existence of this compound. It has been usually stated that it is formed by the action of nitric acid upon arsenious oxide, according to the equation — As 4 0 6 + 4HN0 3 = 2As 2 0 5 + 2H 2 0 + 2N 2 0 3 . It has, however, been shown by the determination of the vapour density, that in the gaseous state the compound N 2 0 3 does not exist, but that the gas is a mixture of molecules of NO and N0 2 . It will be seen that a mixture contain- ing equal volumes of these two gases will have a composition represented by the formula N 2 0 3 , therefore the above reaction may be regarded as taking place thus — As 4 0 6 + 4HN0 3 = 2As 2 0 5 + 2H 2 0 + 2NO + 2N0 2 . Simultaneously with this reaction the following decomposition also goes forward — As 4 0 6 + 8HN0 3 = 2As 2 0 5 + 4H 2 0 + 8N0 2 . The result, therefore, of the action of nitric acid upon arsenious oxide is a mixture of nitric oxide and peroxide in varying proportions. When this mixture is strongly cooled, it condenses to a blue liquid, believed by some to be the true compound N 2 0 3 . Others regard it as merely a solu- tion of the difficultly liquefiable gas, NO, in liquid nitrogen peroxide, N0 2 . If the two oxides are in a state of combination, it would appear to be at best a feeble union, for it has been shown that at temperatures as lov as -90° the liquid slowly evolves NO, while at this temperature no nitrogen peroxide is given off. NITRIC OXIDE. Formula, NO. Molecular weight = 29.96. Density = 14.96. History. — Nitric oxide was first obtained by Van Helmont. Priestley, however, was the first to investigate this gas, which he termed nitrous air , and which was employed by him in his analysis of air. Modes of Formation. — (i.) This gas is obtained by the action of nitric acid of specific gravity 1.2 upon copper or mercury. In Nitric Oxide 221 practice, copper is always employed * The action may be repre sented thus— H 3Cu 4 - 8 HNO 3 = 3 Cu(N 0 3 ) 2 4 4H 2 0 + 2NO. The gas obtained by this method is always liable to contain nitrous oxide, and even free nitrogen ; the amount of these im- purities rapidly increasing if the temperature be allowed to rise and still more so as the amount of copper nitrate in solution increases. ( 2 .) Pure nitric oxide is readily obtained by the action of nitric acid upon ferrous sulphate. The reaction is best applied by gene- rating the nitric acid from potassium nitrate and sulphuric acid in the presence of ferrous sulphate. A mixture of the two salts in the proportion of about one part of nitre to four of ferrous sulphate is introduced into a flask, with a small quantity of water. Strong sulphuric acid is dropped upon the mixture by means of a drop- ping funnel, and the mixture gently warmed, when a steady stream of pure nitric oxide is evolved 2KN0 3 + 5H 2 S0 4 + 6FeS0 4 = 2HKS0 4 + 3Fe 2 (S0 4 ) 3 + 4H 2 0 + 2NO. A precisely similar result may be obtained by the reduction of potassium nitrate by means of ferrous chloride in the presence of hydrochloric acid, thus — KNOq + 3FeCl 2 + 4HC1 = 3FeCl 3 + KC1 + 2H 2 0 + NO. Properties.— Nitric oxide is a colourless gas, having a specific gravity of 1 . 039 . When brought into the air, it combines with the atmospheric oxygen, forming red brown vapours, consisting of a mixture of nitrogen tri- and per-oxides, the combination being attended with a rise of temperature. The formation of these red fumes m contact with oxygen, is characteristic of this gas, therebv distinguishing it from all other gases. This property of nitric oxide renders it impossible to ascertain whether this gas has any smell, or is possessed of any toxicological action. Nitric oxide is only very sparingly soluble in water. It is the most stable of all the oxides of nitrogen, being able to stand a dull red heat with- out decomposition. It is not a supporter of combustion. A lighted taper, or a burning piece of sulphur, when introduced into the gas, * Experiment 314 , “ Chemical Lecture Experiments,” new ed 222 Inorganic Chemistry arc extinguished. If the temperature of the burning substance is sufficiently high to decompose the gas, combustion then continues at the expense of the liberated oxygen : thus, if a piece of phos- phorus, which is freely burning in the air, be plunged into this gas, it continues its combustion with great brilliancy ; if, however, the phosphorus be only feebly burning when thrust into the gas, it is at once extinguished. A mixture of carbon disulphide vapour and nitric oxide, obtained by allowing a few drops of the liquid to fall into a cylinder of the gas, bums, when inflamed, with an intensely vivid bluish flame, which is especially rich in the violet or actinic rays, and has on this account been sometimes employed by photo- graphers to illuminate dark interiors. Nitric oxide is soluble in a solution of ferrous sulphate, forming a dark brown solution, containing an unstable compound of ferrous sulphate and nitric oxide, 2 FeS 0 4 ,N 0 . This compound is readily decomposed by heat, nitric oxide being evolved. By means of this reaction, nitric oxide may be separated from other gases. Nitric oxide is a difficultly liquefiable gas, its critical temperature being - 93.5 : at this temperature a pressure of 71.2 atmospheres is required to liquefy it. The composition of nitric oxide may be proved, by heating a spiral of iron wire by means of an electric current, in a measured volume of the gas (as shown in Fig. 49).* As the metal becomes red hot the gas is gradually decomposed, and the oxygen combines with the iron to form ferric oxide. The residual nitrogen will be found to occupy one-half the original volume. Two vols. of nitric oxide, weighing 29.96 Contain 1 vol. of nitrogen, weighing 14.00 1 5. 96= weight of 1 vol. of oxygen. Therefore we learn that two volumes of nitric oxide consist of one volume of nitrogen and one volume of oxygen united without condensation. Experiment 321. Nitrous Oxide 223 NITROUS OXIDE ( Hyponitrous anhydride , Laughing gas). Formula, N 2 0 . Molecular weight = 43.96. Density = 21.98. History.— This gas was discovered by Priestley, and called by him dephlogisticated nitrous air. Modes of Formation.— (1.) Nitrous oxide is formed by the reduction of nitric acid by certain metals, as zinc or copper under special conditions (see Nitric Acid). These reactions, however are never made use of for the preparation of the gas for experi- mental purposes. (2.) The most convenient method for obtaining this compound is by the decomposition of ammonium nitrate. A quantity of the dry salt is gently/heated in a flask fitted with a cork and delivery- tube. The salt rapidly melts and splits up into nitrous oxide and water — , NH 4 N 0 3 = 2H 2 0 + N 2 0. The heat should be carefully regulated, or the decomposition is liable to become violent, in which case nitric oxide is also evolved Nitrous oxide being rather soluble in cold water, the gas should be collected either over mercury, or over hot water. Wien the gas is to be used for anaesthetic purposes, it should be purified y being passed first through a solution of ferrous sulphate to absorb any nitric oxide and afterwards through caustic soda, to remove any chlorine which may have been derived from the presence of ammonium chloride in the nitrate. Properties.— Nitrous oxide is a colourless gas, having a faint and not unpleasant smell, and a peculiar sweetish taste. Its jjpecific gravity is 1.52. The gas is somewhat soluble in water its coefficient of absorption at o° being 1.3052. The solubility rapidly decreases as the temperature rises, as will be seen by the follow- ing table (Carius) : — 1 c.c. Water at 760 mm. Dissolves At o° „ io° jj 20° „ 25 ° c.c. N 2 0 at o° C. and 760 mm. . I.3052 . *0.9196 - O.670O • O.5962 The loss of gas during its collection over water in the pneumatic trough, arising from its solubility in that liquid, is therefore greatly 224 / norgartic Chemistry lessened by using warm water. Nitrous oxide is much more readily decomposed than nitric oxide, a red hot splint of wood is instantly rekindled, and bursts into flame when plunged into the gas. Phosphorus burns in it with a brilliancy, scarcely perceptibly less dazzling than in pure oxygen. If a piece of sulphur, which is only feebly burning, be thrust into a jar of this gas, the sulphur is extinguished, the temperature of the flame not being sufficiently high to decompose the gas. When, however, the sulphur is allowed to get into active combustion before being placed in the gas, the combustion continues with greatly increased brilliancy. In all cases of combustion in nitrous oxide, the combustion is simply the union of the burning body with oxygen, the nitrogen being eliminated. From its behaviour towards combustibles, nitrous oxide might readily be mistaken for oxygen ; it can, how- ever, be easily distinguished from that gas by the fact that when added to nitric oxide it does not produce red vapours, whereas when oxygen is mixed with nitric oxide these coloured gases are instantly formed. When equal volumes of nitrous oxide and hydrogen are mixed in a eudiometer, and an electric spark passed through the mix- ture, the gases combine with explosion, water being produced and nitrogen set free ; the volume of nitrogen so resulting being equal to that of the nitrous oxide employed. This compound, therefore, contains its own volume of nitrogen, and half its own volume of oxygen. Nitrous oxide, when inhaled, exerts a remarkable action upon the animal organism. This fact was first observed by Davy. If breathed for a short time, the gas induces a condition of hysterical excitement, often accompanied by boisterous laughter, hence the name laughing gas. If the inhalation be continued, this is followed by a condition of complete insensibility, and ultimately by death. On account of the ease with which the state of insensibility can be brought about, this gas is extensively employed as an anaesthetic, especially in dentistry. Nitrous oxide is a gas which is moderately easily liquefied ; at o° C. a pressure of thirty atmospheres is required to effect its liquefaction. Liquid nitrous oxide is colourless and mobile ; it boils at — 92 and when dropped upon the skin produces painful blisters. When thrown upon water, a quantity of the water is at once con- verted into ice : mercury poured into a tube containing a small quantity of the liquid is instantly frozen. An ignited fragment of Nitrous Oxide 225 charcoal thrown upon the liquid floats upon the surface, at the same time burning with brilliancy. If the liquid be mixed with carbon disulphide, and placed in vacuo, the temperature falls to -140. By strongly cooling the liquid, contained in a sealed tube, Faraday succeeded in solidifying it; this may also be effected by the rapid evaporation of the liquid. The solid melts at -99°, and if placed upon the hand causes a painful blister; in this respect it differs from solid carbon dioxide, which gasifies without previous liquefaction. Hyponitrous Acid, NHO.— This substance has not yet been isolated, being only known in its salts and in aqueous solution. When a solution of potassium nitrate, or nitrite, is acted upon by sodium amalgam (an alloy of sodium and mercury), the salt is reduced by the nascent hydrogen, evolved by the action of the amalgam upon water, and the potassium salt of hyponitrous acid is left in solution— KN0 3 + 2H 2 = 2H s O + KNO. The solution, which is alkaline, owing to the presence of sodium hydroxide is then made neutral by the addition of acetic acid, and silver nitrate added.’ A yellow precipitate is thrown down, consisting of silver hyponitrite, AgNO. When a solution of potassium hyponitrite is acidified and then heated, the hyponitrous acid, which may be regarded as liberated by the acid, is broken up into nitrous oxide and water — v 2HNO = N 2 0 + H 2 0. Nitrosyl Chloride, * NOC1.— This compound maybe obtained by the direct combination of nitric oxide with chlorine — 2NO + Cl 2 = 2NOC1. It is also formed by the action of phosphorus pentachloride upon potassium nitrite, thus — PC1 3 + KN0 2 = NOC1 + POCI3 + KC1. Nitrosyl chloride is formed together with chlorine when a and hydrochloric acids is gently heated- mixture of nitric HNO3 + 3HC1 NOC1 + Cl 2 + 2H 2 0. Nitrosyl chloride is also readily prepared by the action of nitrosyl hydrogen sulphate upon dry sodium chloride, thus— (N0)HS0 4 + NaCl = NOC1 + NaHS0 4 . Properties. — Nitrosyl chloride is an orange-yellow denses when passed through a tube immersed in a gas, which easily con- freezing mixture, to an * Tilden has shown that this is the only oxy-chloride of nitrogen that exists. P 226 Inorganic Chemistry orange-yellow liquid, which boils at about - 8°. It is decomposed by water into nitrous acid and hydrochloric acid — NOC1 + H 2 0 = HN0 2 + HC1. In a similar manner it is decomposed by metallic oxides and hydroxides, thus — NOC1 + 2KHO = KN0 2 + KC1 + H 2 0. Nitrosyl chloride has no action upon gold and platinum, but it attacks mercury with the formation of mercurous chloride and the liberation of nitric oxide — 2NOC1 + Hg., = Hg 2 Cl2 + 2NO. CHAPTER VI THE ATMOSPHERE Thb atm °s pher e is the name applied to the gaseous mixture wHch envelops the earth, and which is commonly called the air The older chemists used the word air much as in modem times the word gas is employed; thus they spoke of inflammable air, dephlogisticated air, alkaline air, and so on. The air consists of a mixture of gases, the two chief ingredients being nitrogen and oxygen. Lavoisier was the first to clearly prove that oxygen was a constituent of the air, although Robert Boyle and others before him had shown, that air was absorbed by metals in the process of forming a calx, and that the metal gained weight as the calx formed. When the fact that the air was com- posed of oxygen and nitrogen became established, various devices were adopted to determine the proportion of oxygen in it Priestley s method was by means of nitric oxide. It depended upon the fact that when nitric oxide is mixed with air, it combines With the oxygen, forming brown fumes which dissolve in the water A contraction in volume therefore takes place, from which the volume of oxygen may be calculated. This method yielded results which seemed to show that there was considerable variation in the proportion of oxygen present in different samples of air, and the denenT 5 ! 5 u wholesom eness, or goodness, of the air was dependent upon the quantity of oxygen which it contained. Hence arose the term eudiometry , signifying to measure the goodness Cavendish, on the other hand, as the result of a large number of expenments; made by him, came to the conclusion that there was no difference in the samples of air that he experimented upon. bince the time of Cavendish, eudiometric analysis has been brought to a state of great perfection and accuracy by Bunsen from the’ and °‘ herS - The “nclusio/to be Z2 t7e atmo f res< ? rches of ‘hese chemists is, that although posittln the er V e ' rta, y S a remar hable uniformity of com- position, there do exist perceptible, though very slight, variations 228 Inorganic Chemistry in the amount of oxygen present at different places and at different times. Samples of air collected from all parts of the globe, from mid ocean, from high mountain peak, American prairie, and i crowded cities, show a variation in the proportion of oxygen rang- j ing from 20.99 to 20.86. Angus Smith has shown, that in foggy | The Atmosphere 229 weather the oxygen in the air in towns, sometimes falls as low as 20.82. Samples of air taken from crowded theatres, have been found to contain as little as 20.28, while in many mines the amount averages as low as 20.26. The mean proportions of oxygen and nitrogen in the atmosphere may be given as — The composition of the atmosphere by weight was determined by Dumas and Boussingault (1841). In their method, air which was freed from carbon dioxide and moisture, was slowly drawn through a glass tube containing a known weight of metallic copper, heated to redness. The oxygen combined with the copper, forming copper oxide, which was afterwards weighed, and the nitrogen passed into a vacuous flask, and was also weighed. The apparatus as em- ployed by Dumas is seen in Fig. 50. B is a glass flask having a capacity of 10 to 15 litres, which was exhausted and then weighed. It was then attached, as shown, to the tube T, containing a known weight of metallic copper, and which was also exhausted. The bulbs L contained a solution of potassium hydroxide, and the tubes f solid potash, for the removal of atmospheric carbon dioxide. The bulbs O contained strong sulphuric acid, and the tubes t were filled with pumice moistened with the same acid, by means of which the moisture was withdrawn from the air. When the copper was heated and the cocks partially opened, air, free from carbon dioxide and moisture, was slowly drawn over the heated metal, which was thereby converted into the oxide. At the conclusion of the experiment the globe and the tube T were reweighed. The nitrogen remaining in tube T was then pumped out and the tube once more weighed. The difference between the two last weigh- ings of the tube, added to the gain in weight suffered by the globe, gave the nitrogen ; while the difference between the original and final weights of the tube gave the increase of weight suffered by the copper, that is, the amount of oxygen. The result of numerous experiments gave the mean composition — Oxygen 23 parts by weight. Nltro S en 77 „ Oxygen . Nitrogen . 20.96 parts by volume. . 79.04 „ 100.00 100 230 Inorganic Chemistry A more modern method for estimating the amounts of oxygen and nitrogen in the air, based upon the same principle, namely, the absorption of the oxygen by heated metallic copper, is illustrated in Fig. 51 (known as Jolly’s apparatus). The sample of air to be examined is allowed to enter the glass globe A (whose capacity is about 100 c.c., and which has been previously exhausted) by means of the three-way cock b. (The air is first dried, by being drawn through tubes filled with pumice moistened with sulphuric acid, on Fig. 51. its way into the apparatus.) The bulb is then surrounded by the metal jacket B, which is filled with broken ice, and when the tem- perature has fallen to o° the bulb is put into communication with the tube d by means of the three-way cock. The tube g is then raised or lowered, so as to bring the mercury in d to a fixed point in the tube at m , and the tension of the enclosed air is ascertained by the graduated scale behind tube g. The ice-jacket is then removed, and the spiral of copper wire within the bulb is heated 231 The Atmosphere to redness by the passage through it of an electric current. The copper combines under these conditions with the oxygen, form- ing copper oxide, thereby reducing the volume of the contained gas. The globe is again cooled, and the tube g lowered to such a position, that when communication is once more made between the globe and tube d ’ the mercury shall stand at the same point m. From the observed tension of the gas before and after the experiment, the volume relations of the two constituents can be calculated. Thus, suppose the tension of the enclosed air to be 720.25 mm., and that of the residual nitrogen 569.28 mm., then for 1 volume of air the reduction would be — 569.28 720.25 = .7904 vols. Therefore in 100 volumes the composition would be— Nitrogen * = 79.04 Oxygen = 20.96 100.00 Besides oxygen and nitrogen, the air contains variable quantities of the following gases : carbon dioxide, aqueous vapour, ammonia, ozone, nitric acid. With the exception of aqueous vapour, these substances are present only in relatively small proportions, and with all of them the amount is liable to considerable variation. Especially is this the case with the aqueous vapour, as the amount of this constituent present at any time is largely influenced by the temperature. The average composition of normal air may be taken as follows : — Nitrogen* Vols. per tooo. . 779.0600 Oxygen . . 206.5940 Aqueous vapour . . I4.OOOO Carbon dioxide O.3360 Ammonia 0.0080 Ozone . O.OOI5 Nitric acid . 0.0005 1000.0000 Aqueous Vapour. — For any given temperature there is a maximum amount of aqueous vapour, which a given volume of air * The small percentage of argon present is here included with the nitrogen. 232 Inorganic Chemistry is capable of taking up : under these conditions the air is said to be saturated with moisture at the particular temperature. Thus 1 cubic metre of air is saturated with moisture at the various tempera- tures stated, when it has taken up the following weights of water : — At o° . 4.871 grammes. I At 20° . 17* 1 57 grammes. „ io° . 9.362 „ I „ 30° . 30.095 „ When air, saturated with moisture at say 20°, is cooled to io°, the excess of water beyond 9.362 (the maximum for io°) is deposited either as mist or rain. The temperature at which air thus begins to deposit moisture is called the dew-point. The deposition of moisture from the air caused by the lowering of the temperature is a matter of everyday observation. A glass vessel containing iced water, becomes bedewed with moisture upon the outside, as the air in its immediate vicinity is cooled. When a season of severe frost is suddenly followed by a warm wind, highly charged with aqueous vapour, it is not unusual to see condensed moisture collecting upon, and streaming down, the cold surface of walls. For the same reason, after the sun has set, and the heat from the ground has radiated, leaving the ground colder than the atmos- phere, the temperature of the air is lowered, and it begins to deposit its aqueous vapour in the form of dew. The amount of aqueous vapour in the air, or the humidity of the air, is estimated by meteorologists by means of an instrument called the wet and dry bulb thermometer. Carbon Dioxide. — The proportion of this gas present in the air is also liable to considerable variation, although not through such a wide range as the aqueous vapour. The processes of respiration, combustion, and putrefaction are attended by the evolution of carbon dioxide, hence the amount of this gas present in closed inhabited places, is greater than that in the open air ; in badly ventilated and crowded rooms, the proportion sometimes rises to three parts in 1000 vols. Frankland has found that at high elevations the amount of carbon dioxide in the air is often, although not invariably, considerably above the normal. At Chamounix (3cxx> feet) the amount of carbon dioxide was 0.63 per 1000 vols. ,, Grands Mulets (11,000 feet) ,, ,, ,, i.n .. .. ,, Mont Blanc (15.732 feet) ,, ,, .. 0.61 ,, This fact is probably due to the absence, in these high regions, The Atmosphere 233 of the vegetation which is one of the chief natural causes operating to remove atmospheric carbonic dioxide (see Oxygen, page 166). The amount of carbon dioxide is slightly higher during the night, and often rises considerably during foggy weather. Thorpe has shown that near the surface of the sea, the amount of carbon dioxide in the air is slightly less, being on an average 0.300 vols. per 1000. Ammonia in the atmosphere is derived from the decomposition of nitrogenous organic matter. Although present in relatively very small quantities, it varies in amount very considerably. From the experiments of Angus Smith, 1000 grains of air from various sources were found to contain the following amounts of ammonia : London Glasgow Manchester . 0.05 grammes. 0.06 The proportion of ammonia appears to be higher during the night than in the daytime, and immediately after heavy rain the amount is perceptibly diminished. Rain water always contains ammonia, although the amount varies greatly with changing atmospheric and climatic conditions. Lawes and Gilbert, Angus Smith, and others, have made a large number of estimations of the amount of ammonia in rain water at various places and seasons, and under many different conditions. Nitric Acid, present in the form of nitrates and nitrites, is produced in the atmosphere by the direct union of oxygen and nitrogen whenever a lightning flash passes through the air (see Nitric Acid). Rain which falls during or immediately after a thunderstorm is found to contain nitrates and nitrites. These two nitrogenous compounds, ammonia and nitric acid, although present only in such small proportion in the atmosphere! iulfil a most important function in the economy of nature. From the experiments of Lawes and Gilbert and others, it has been shown that most plants are unable to draw upon the free nitrogen of the atmosphere, for the supply of that element which they require for the development of their structure and fruit.* Although they are surrounded by, and bathed in, nitrogen, they cannot assimilate it. Plants that are growing in unmanured soil, therefore, derive their * Leguminous plants, such as clovers, vetches, beans, root-nodules or tubercles, are exceptions. peas, which develop 234 Inorganic Chemistry nitrogen from the ammonia and nitric acid which are present in the air, and which are washed into the ground by the rain. It has been found that a plant grown under such experimental conditions as to exclude the possibility of its obtaining supplies of these nitro- genous compounds, will yield upon analysis exactly the same amount of nitrogen as was originally contained in the seed from which it grew. Ozone. — The causes which operate in the formation of this sub- stance in the air are at present imperfectly known : it is supposed that its occurrence is related to the development of electricity in the atmosphere. On account of the powerful oxidising character of ozone, its presence can never be detected in the air, where much organic matter of an oxidisable nature is present, as is the case in the air of such places as malarial swamps, dwelling-houses, and large towns. The amount of ozone in pure country air has been found to vary with the time of year, reaching a maximum in the spring-time, and gradually falling towards winter. Thorpe has found that in sea air the amount of ozone is practically constant during all seasons. The usual method which is available for the detection and estimation of ozone in the air is extremely crude. It consists in exposing ozone test-papers (see Ozone) to the air for a certain time, and comparing the colour that is produced with a standard scale of tints : moreover, other substances than ozone, which may be present in the atmosphere, will also liberate iodine from potassium iodide, and these are therefore measured as ozone. Besides the higher oxides of nitrogen, which, as we have seen, are formed in the atmosphere, and which liberate iodine from potassium iodide, it has been shown that peroxide of hydrogen is also present. The state of our knowledge at present, therefore, respecting the exact amount of atmospheric ozone, and its variation, is far from satisfactory ; it is, indeed, quite possible that many of the effects which have been attri- buted to ozone, are in reality due to peroxide of hydrogen. Thus it has been shown by Schonbein that this compound is formed during the evaporation of water, and this statement probably derives con- firmation from the fact that its presence may be detected in rain water. The salubrity of the air of the sea-shore, where large areas of wet sand and stones offer the most perfect conditions for the rapid evaporation of water, and, consequently, for the formation of peroxide of hydrogen, may therefore be attributable as much to the presence of this substance as to the proverbial ozone. 235 The Atmosphere The various gases of which the air is composed, are not com- bined, but are merely mingled together. The remarkable con- stancy of its composition, as regards the oxygen and nitrogen, led chemists at one time to suppose, that these gases were in chemical union with each other in the atmosphere ; but a number of facts which have since been learnt respecting these gases, prove with- out doubt that this is not the case, and that the air is simply a mechanical mixture. This evidence may be briefly summed up as follows : — (i.) When oxygen and nitrogen are mixed together in the pro- portion in which they occur in air, the resulting mixture behaves in all respects like ordinary air, and the mixing of the gases is not attended by any volumetric or thermal disturbance, such as would be expected to accompany the chemical union of two elements. (2.) The degree to which air is capable of refracting light, is found to be the mean of the refractive power of oxygen and nitrogen. Were these gases chemically combined, the compound should behave in this respect as other compound gases, where it is found that the refractive index is always either greater or less than the mean of that of the constituents. (3.) According to a fundamental law of chemical science, the composition of a chemical compound is constant. Such a thing as variability in the composition of a compound is unknown. The proportion of oxygen and nitrogen, as we have seen, does vary in the air, although through only small limits, hence they cannot be united to form a compound. (4.) The proportion by weight in which oxygen and nitrogen are present in air, bears no simple relation to the atomic weights of these elements. (5.) When air is dissolved in water, the oxygen and nitrogen dissolve as from a simple mixture of these gases, in accordance to the law of partial pressures (see page 127). (6.) The oxygen and nitrogen can be partially separated, by taking advantage of the different rates of diffusion of these two gases (see Diffusion of Gases, page 80). The various gases of the atmosphere are maintained in a state of uniform admixture, in spite of their widely different densities, by the operation of two causes : first, air currents, which effect the rapid removal of large masses of air from place to place ; and, second, their own molecular movements, which bring about the phenomena of gaseous diffusion. 236 Inorganic Chemistry Suspended Impurities in the Atmosphere. — Besides the gaseous constituents of the air, there is always present a certain quantity of suspended matter, both liquid and solid. The exist- ence of this suspended matter in the air, can be rendered evident from the fact, that these minute particles are capable of reflecting light ; if, therefore, a strong beam of light be passed through a darkened room, the track of the beam is distinctly visible, on account of its being reflected from innumerable particles floating about in the air, many of them appearing quite large. Pasteur has shown, that this suspended matter can be removed by filtration through cotton wool.* Tyndall also has shown, that in undis- turbed air the suspended matter settles in the course of a few hours, leaving the air almost entirely free from this impurity. For this purpose the floor of a large oblong glass box was smeared over with glycerine. The box, after being hermetically closed, was then allowed to stand for twenty-four hours, during which time the suspended matter subsided, and adhered to the glycerine. When a beam of light is allowed to pass through air that has been thus freed from suspended matter, there being nothing present to reflect the light, the beam cannot be seen \ its track will be evident in the air of the room as it enters and leaves the box, but within the box it will be invisible (as repre- sented in Fig. 52). To air in which a beam of light is in this way invisible, Tyndall has applied the term “ optically pure.” The suspended matters are partly mineral and partly organic. Of the mineral matters, sodium chloride and certain sulphates are present in greatest quantity. These are thrown into the air in the sea-spray, and as the small globules of water evaporatq See Experiments 334 to 341, “ Chemical Lecture Experiments,” new ed. The Atmosphere 237 they leave minute residua) particles of saline matter, which being driven by the wind, remain floating in the atmosphere It is only very rarely, even at far inland places in Europe, that spec- troscopic examination fails to detect the presence of sodium com- pounds in the air. In the air of islands, such as England, it is never absent. Sulphates are also produced by the oxidation and combustion of sulphuretted compounds ; the amount of these, there- fore, is greatly increased in the neighbourhood of towns The organic suspended matter of the air has of late years been made the subject of extended research. Pasteur has shown that amongst these organic substances are the germs and organisms which produce fermentation, putrefaction, and disease. Putrescible substances, such as milk, urine, flesh, &c., if themselves carefully freed from al such germs, may be preserved unchanged, for ap- parently any length of time, in air that has been deprived of all suspended matter. It is highly probable that the salubrity or otherwise of different places, is associated with the nature and amount of the organic matter in the air, and it is certain that these organisms play a most important part in relation to the life and health of man. The feelings of lassitude and headache, which result from the prolonged breathing of the air of rooms containing many people, are brought about more by the poisonous effects of e organic emanations evolved during respiration, than by any diminution in the supply of oxygen, or increase in the proportion of carbon dioxide in the air. The well-known and unpleasant smell that is perceived on first entering a crowded room, is also due to the same cause, and it has been shown that the moisture which condenses from such an atmosphere upon a cold object if preserved for a short time, rapidly becomes putrescent, owing’to the decomposition of this organic matter. The presence of suspended matter in the air, appears to exert a remarkable influence upon the formation and character of fogs bonofafo a 1 th0Se COnditions whi ch result in the forma- tion of a fog in ordinary air, are incapable of producing that effect m air that has been freed from suspended matter. It would Lpt, that the suspended particles act as innumerable points or nuclei Ts tie “nr *. 6 dep , osltio , n of moisture, much in the same way from Crys ! alI,sat,on of a s alt, from its solution, is known to start the liquid. minUte particles of forel & n matter that may be floating in The height to which the atmosphere extends has been variously 238 Inorganic Chemistry estimated. From observations of the flight of meteorites, it ap- pears that even at a height of seventy to seventy-five miles the air still has a sensible degree of density. The air being elastic, and subject to the law of gravitation, its density, which is greatest at the earth’s surface, rapidly diminishes as the altitude increases ; thus, at about three and a half miles the density is only one-half, and at seven miles one-third, of that which obtains at the sea-level. From a consideration of the physical properties of gases, there is every reason to believe, that in an extremely attenuated condition the atmosphere extends far into space, and it has been calculated, that the pressure exerted by our atmosphere upon the surface of the moon, is equal to about 1 mm. of mercury. The density of the atmosphere varies at different points of the earth’s surface, and at the same point at different times. The pressure exerted by the atmosphere is measured by the height of a column of mercury which it is capable of supporting, the instrument employed for the purpose being called the barometer. At the sea-level in the latitude of London, the average weight of the atmosphere is equal to that of a column of mercury 760 mm. at o°, and this is taken as the standard pressure of the atmos- phere. CHAPTER VII COMPOUNDS OF NITROGEN AND HYDROGEN of nitrogen with hydro?en have be - p- Ammonia Hydrazine Hydrazoic acid NH, N 2 H 4 or (NH 2 ) 2 N 3 H or HNo. AMMONIA. Formula, NH 3 . Molecular weight = 17. Density = 8.5. History -Salts of ammonia, and also the aqueous solution, were known to the alchemists. It was termed by Glauber, sfiritus vola- tile salts armomaa, being obtained by the action of an alkali upon sal-armomacum Subsequently, when ammonia was obtained by the destructive distillation of such refuse as hoofs and horns of animals, the name spirits of hartshorn was applied to it The actual discovery of gaseous ammonia was made by Priestley (1774), when he collected the gas, evolved by the action of lime upon sabammoniac, by means of his mercurial pneumatic trough. Priestley named the gas alkaline air. Occurrence. -In combination as carbonate of ammonia it is pre- sent in small quantities in the air, derived by the decay of nitro- found^ a ” lma and ve?etable matter - A s nitrate and nitrite it is found in rain water. It is evolved, along with boric acid, from the fumaroles of Tuscany (see Boric Acid), and is found as chloride and sulphate in the vicinity of active volcanoes dufed b 0f formation— (i.) Ammonia can be synthetically pro- duced, by submitting a mixture of nitrogen and hydrogen to the influence of the silent electric discharge (Donkin) The amoum ammonia so obtained, however, is extremely small and can best l S ”«roul P a aSSi r g d the gaS6S ’ 32 th6y iSS " 6 ^ one ’ hrough a cylinder containing a small quantity of NesslePs 239 240 Inorganic Chemistry solution.* In a short time the solution will begin to show a yellowish brown colour, indicating the presence of traces of ammonia. (2.) Ammonia may be prepared by gently heating any of its salts, with either of the caustic alkalies, potash or soda, or with slaked lime. The salt most commonly employed is the chloride. When this is mixed with an excess of slaked lime, and the mixture gently heated in a flask, ammonia is evolved, and calcium chloride and water are formed — 2NHjCl + CaH 2 0 2 = CaCl 2 + 2 H 2 0 + 2 NH 3 . The gas may be dried by being passed through a cylinder con- taining lumps of quicklime, t and may then be collected either by upward displacement, or in the mercurial trough. On account of its extreme solubility it cannot be collected over water. (3.) Ammonia is formed by the action of nascent hydrogen upon salts of nitrous and nitric acid, thus — NaN 0 3 + 4 H 2 = NaHO + 2 H a O + NH 3 . This method is often made use of, in the quantitative estimation of nitrates in drinking water. (4.) When nitrogenous organic matter is subjected to destruc- tive distillation, that is, strongly heated out of contact with air, ammonia is formed : hence when coal, which usually contains about 2 per cent, of nitrogen, is distilled in the process of the manu- facture of ordinary illuminating gas, one of the products of the decomposition is ammonia. The “ ammoniacal liquor” of the gas works, is the source of all ammonia salts at the present day. The liquor is boiled with milk of lime, and the ammonia thus expelled is absorbed by sulphuric acid. The ammonium sulphate so obtained is evaporated to dryness, and purified by recrystallisation. Properties. — Ammonia is a colourless gas, having a powerfully pungent smell, and a strong caustic taste. It is lighter than air, its density being 0.589 (air = 1). Ammonia possesses the property of alkalinity in a very high degree ; it turns red litmus blue, and yellow turmeric brown. The gas is unable to support combustion, * A solution of mercuric iodide in potassium iodide, rendered alkaline with potassium hydroxide. f The usual desiccating agents, namely, sulphuric acid, or phosphorus pentoxide, are inadmissible in the case of ammonia, as this gas at once unites with such compounds. A mtnonia 24 1 and is irrespirable. Under ordinary conditions ammonia is not combustible, but if the air be heated, or if the amount of oxygen be increased, the gas will then burn with a flame of a characteristic yellow-ochre colour. This behaviour of ammonia as regards com- bustibility, is most conveniently illustrated by means of the ap- paratus shown in Fig. 53. A stream of the gas, obtained by gently heating a quantity of the strong aqueous solution in a small flask, is delivered through a tube which is surrounded by a wider glass tube. Through the cork which carries this tube a second tube passes, through which a supply of oxygen can be passed. On applying a lighted taper to the jet of am- monia as it issues from the tube, it will be noticed that the gas burns in the heated air round the flame of the taper, but is unable to continue burning when the taper is withdrawn. If now a gentle stream of oxygen be admitted into the annular space Fig. 53. between the two tubes, the am- monia readily ignites, and continues to burn with its character- istic flame. On cutting off the supply of oxygen, the flame of the burning ammonia languishes and dies out. Ammonia is extremely soluble in water ; 1 c.c. of water at o° C., and at the standard pressure, dissolves 1148 c.c. of ammonia,’ measured at o° C. and 760 mm. The solubility rapidly decreases as the temperature rises, as will be seen by the following table 1 c.c. of Water at 760 mm. Dissolves At o° . „ 8 ° . „ 16 „ 30° . „ 5o° . Grammes, NH 3 . • O.875 • . O.713 . . O.582 . • 0*403 • . 0.229 • c.c. at o* C. and 760 mm. . II48 • 923 • 764 . 529 306 242 Inorganic Chemistry When a solution of ammonia is heated, the gas is rapidly evolved, and at the boiling temperature the whole of it is given up. The great solubility of this gas in water may be shown by filling a large bolt-head flask with ammonia by displacement, the flask being closed by means of a cork through which a long tube passes, as shown in Fig. 54. On removing the cork from the end of the tube, water slowly rises until it reaches the top, and as soon as the first drops enter the globe the absorption proceeds with great rapidity, the water being forced up the tube in the form of a fountain, which continues until the flask is filled. Commercial liquor ammonice is prepared by passing ammonia gas into water ; the strongest solution has a specific gravity of 0.882 at 1 5 0 , and contains 35 per cent, of ammonia. During the process of solution heat is liberated, and when the gas is again expelled, the same amount of heat is reab- sorbed. If a rapid stream of air be driven through a quantity of strong ammonia solu- tion, contained in a glass flask, the am- monia gas is quickly expelled ; and if the flask be placed upon a wooden block, as seen in Fig. 55, upon which a few drops of Fig. 54. water have been poured, it will be found that after a few moments the flask will have become firmly frozen to the block. By the rapid evaporation of ammonia in this way, it is possible to lower the temperature to -40° C. Ammonia is an easily liquefiable gas ; thus at 15.5 0 it requires a pressure of 6.9 atmospheres, and at o° only 4.2 atmospheres, in order to liquefy it. The gas was first liquefied by Faraday (1823), by heat- ing in one limb of a closed and bent glass tube (see Fig. 2), a quantity of a compound of ammonia with silver chloride, the other limb of the tube being immersed in a freezing mixture- The experiment may be made in a tube constructed as seen in Fig. 56. The wide limb is nearly filled with dry precipitated silver chloride, which has been saturated with ammonia gas. This compound melts at about 38°, and at a somewhat higher temperature it gives up its ammonia. If the narrow limb of the tube be immersed in a freezing mixture while the compound is being heated, the com- A mmonia 243 bined influence of the cold, and the pressure exerted by the evolved ammonia, will cause the gas to liquefy and collect in the cold portion of the tube. On removing the tube from the freezing mixture, and allowing the other end to cool, the liquid ammonia will boil off, and be reabsorbed by the silver chloride, reforming the original compound. Liquid ammonia is easily obtained in larger quantity, by passing the gas through a glass tube immersed in a bath of solid carbonic acid and ether. Liquid ammonia is a colourless, mobile, and highly refracting liquid, boiling at -33.7 0 , and having a specific gravity at o° of 0.6234. When cooled below -75 0 it solidifies to a mass of white crystals. Liquid ammonia dissolves the metals sodium and potassium, the solution in each case being of an intense blue colour. On the evaporation of the liquid, the metal is deposited unchanged. 0 - Fig. 55. Fig. 56. During the evaporation of liquid ammonia, boiling as it does at so low a temperature as - 33.7 0 , a rapid absorption of heat takes place, and as this substance/ is so easily obtained it was one of the earliest liquids employed for the artificial production of ice. Various ice-making machines have been invented by M. Carr£, in which the reduction of temperature required is obtained by the evapora- tion of liquid ammonia. Ammonia is decomposed into its elements at a temperature below a red heat. In this decomposition, two volumes of ammonia give one volume of nitrogen, and three volumes of hydrogen. The gaseous products, therefore, obtained by passing ammonia through a red hot tube, are inflammable. In the same way, when electric sparks are passed through ammonia, the gas is resolved into its constituents. By performing this experiment upon a measured volume of ammonia, confined in a eudiometer over mercury, it will 244 Inorganic Chemistry be found, that after the passage of the sparks for a short time, and the readjustment of the levels of mercury, the original volume of the gas has been doubled. The fact that the hydrogen and nitrogen are present in ammonia in the proportion of three volumes of hydrogen to one of nitrogen, can be shown by taking advantage of the fact that ammonia is decomposed by chlorine, the latter combining with the hydrogen to form hydrochloric acid, and the nitrogen being set free. This is effected by means of the apparatus shown in Fig. 57. The long glass tube, divided into three equal divisions, is filled with chlorine and closed by a cork carrying a small dropping funnel. A few cubic cen- timetres of strong aqueous ammonia are poured into the funnel and allowed to enter the tube drop by drop. As the first two or three drops fall into the chlorine, it will be seen that the combination is attended with a feeble flash of light, and fumes of ammonium chloride are formed. When the reaction is complete, the whole of the chlorine will have combined with hydrogen, derived from the ammonia, to form hydrochloric acid, and this in its turn will combine with the excess of ammonia added, forming ammonium chloride. This substance dissolves in the water. A small quantity of dilute sulphuric acid is next introduced by means of the dropping funnel in order to absorb the remaining excess of Fig. 57. ammonia. The atmospheric pressure is then once more restored by attaching to the funnel a bent tube, dipping into a beaker of water as shown in the figure, and when the water is allowed to enter, it will be found to flow into the tube until it reaches the second gradua- tion. The gas which is left, and which occupies one of the divisions of the tube, is found on examination to be nitrogen. This one measure of nitrogen, therefore, has been eliminated from that amount of ammonia which has been decomposed by the chlorine, with which the tube was originally filled. Now chlorine combines with its own volume of hydrogen, therefore the volume Hydrazine 245 of hydrogen which was in combination with the one measure of nitrogen, is equal to the volume of chlorine contained in the tube, that is to say, it was three measures. We have, therefore, one measure of nitrogen and three measures of hydrogen, or, in other words, ammonia is a combination of nitrogen and hydrogen in the proportion of one volume of nitrogen to three volumes of hydrogen Ammonia combines directly with acids forming salts, known as ammonium salts, in which the nitrogen functions as a pentad element ; thus with hydrochloric and sulphuric acids it forms respec- tively ammonium chloride and ammonium sulphate— NH 3 + HC1 = (NH 4 )C1. 2NH 3 4 - H 2 S0 4 = (NH 4 ) 2 S0 4 . (The ammonium salts will be described with the compounds of the alkali metals.) Hydrazine, NH 2 *NH 2 or N 2 H 4 .— This compound was first prepared by Curtius (1887). It is obtained by heating together in a sealed tube to a temperature of 170°, hydrazine hydrate, N 2 H 4 ,H 2 0, and barium monoxide. Under these circumstances the barium oxide takes up the water from the hydrazine hydrate, according to the equation— BaO + N 2 H 4 H 2 0 = Ba(HO) 2 + N 2 H 4 . When the tube is opened, the gaseous hydrazine, which is under considerable pressure, rushes out of the tube, forming dense fumes in contact with the atmospheric moisture, with which it combines with great readiness Hydrazine Hydrate, N 2 H 4 H 2 0 .-The compound formed by the combina- NHH SO m W “ h Wa ‘ er “ ° btained by dis,illin g hydrazine sulphate, in » > 7 ' , a “ T e0US SOlmion 0f P otass ium hydroxide (caustic potash , ™ of sllver ; 11 ls a colourless, fuming, powerfully corrosive liquid winch bods at 118.5". It attacks glass, corks, and indiarubber, and can 0“?; be prepared in vessels of silver or platinum which are screwed together at their onTnT' 1 ', he ha ’° gen acids il forms two of salts, if which either “id we hfv?- 6 hal ° gen add ^ PreSent 1 thUS with hydrochloric Hydrazine monohydrochloride . . . N H Hri Hydrazine dihydrochloride .... N a H 4> 2 HCl Hydrazoic Acid, HN 3 .— Discovered by Curtius (1800) The c 0r ii u ■ prepared by boiling benzoylazo-imide whh sodium SLxTde wh ” sfd “ benzoate and sodium hydrazoate are formed, thus- nyar ° X,de ' when sodlun ’ ^ /N C 6 H 5 CO - N || + 2 NaHO : ! \ N LL^gS° d ^ t (° btai “cd by heating sodium in dry g ’ 1S heated t0 200 »n a stream of nitrous oxide * 2NH 2 N a + N 2 0=NaN 3 + NaH0 + NH 3 . /N : QH 5 CO-ONa + Na-N || +H 2 0 . * See Experiment new ed ‘ 246 Inorganic Chemistry The sodium hydrazoate so obtained is then gently warmed with dilute sulphuric acid, when sodium sulphate and hydrazoic acid are formed, thus — 2NaN 3 + H. 2 S0 4 = Na^SC^ + 2HN 3 . Properties. — This compound is a colourless gas, possessing a most un- pleasant and powerfully penetrating odour. If inhaled, even when largely diluted with air, it exerts an irritating action upon the mucous membrane. As its name denotes it is an acid substance, and in many of its properties it strongly resembles the halogen acids. The gas is extremely soluble in water, and forms a strongly acid liquid which smells of the gas. This solution when boiled, finally assumes a definite strength, and yields on distillation an aqueous acid of constant composition, in this respect resembling aqueous hydrochloric acid, q.v. In its constitution this acid may be compared with hydrocyanic acid, and with the halogen acids — H(N 3 ); H(CN) ; H(C1) ; H(Br), in which the radical cyanogen (CN), or the halogen elements, Cl and Br, are replaced by the group consisting of three nitrogen atoms. When a solution of hydrazoic acid is added to a solution of silver nitrate, a white precipitate of silver hydrazoate is formed, strongly resembling silver cyanide and silver chloride. This silver salt, however, is not acted upon by light in the way the chloride is, and it differs also in being extremely explosive. A minute quantity of the compound, when touched with a hot wire, detonates violently. When gaseous hydrazoic acid is mixed with gaseous ammonia, dense white fumes are formed, consisting of ammonium hydrazoate. These two hydrides of nitrogen, apparently so similar, but in reality so widely different, unite to form the ammonium salt, just as gaseous hydrochloric acid and ammonia combine to form ammonium chloride, thus — NH 3 + H(N 3 ) = NH 4 (N 3 ). NH 3 + HC1 = NH 4 C1. The alkaline hydride of nitrogen, ammonia, combines with the acid hydride of nitrogen, hydrazoic acid, and forms the salt ammonium hydrazoate NH 4 N 3 or N 4 H 4 . HYDROXYLAMINE. Formula, NH 2 (OH). Discovered by Lossen in 1865 . Modes of Formation. — (1.) By the action of nascent hydrogen upon nitric oxide, nitric acid, or certain nitrates — 2NO + 3H 2 = 2NH 2 (OH). HN0 3 + 3H 2 - 2H 2 0 + NH 2 (OH). The nascent hydrogen is evolved from tin and hydrochloric acid, and a stream of nitric oxide passed through the mixture. The Hydroxylamine 247 hydrochloride of hydroxylamine is thus obtained. This is purified by first passing a stream of sulphuretted hydrogen through the solution. The tin is thus precipitated as sulphide, and is removed by filtration. The filtered solution is then evaporated to dryness, and the hydrochloride of hydroxylamine is dissolved out of the residue by means of absolute alcohol, in which solvent ammonium chloride is only very slightly soluble. The alcohol is then distilled off, and the residue is converted into the sulphate by being treated with the requisite quantity of sulphuric acid. Hydroxylamine itself, in aqueous solution, is obtained from the sulphate by the addition of baryta-water. ( 2 .) By boiling potassium hydroxylamine-disulphonate with water for several hours — 2N(0H)(S0 3 K) 2 + 4H 2 0 = (NH 3 0H) 2 S0 4 + 2K 2 S0 4 + H 2 S0 4 . The potassium sulphate is removed by crystallisation. Properties. — Hydroxylamine is known only in aqueous solution. The solution is colourless, and has an alkaline character. When the solution is distilled the substance is partially decomposed. The solution is a powerful reducing agent : it precipitates gold and mercury from their solutions, and reduces cupric salts, throwing down the red cuprous oxide on being boiled. Hydroxylamine is a base, and may be regarded as ammonia, in which one of the hydrogen atoms has been replaced by the monad group hydroxyl (OH). Its salts, like those of ammonia, are formed by direct union with an acid, without the elimination of water. NH 2 OH + HC1 = NH 3 OHCl (or NH 2 OH,HCl). 2NH 2 OH + H 2 S0 4 = (NH 3 OH) 2 S0 4 (or 2NH 2 OH,H 2 S0 4 ). The salts of hydroxylamine all decompose on the application of heat, with a more or less sudden and violent evolution of gas ; thus the nitrate breaks up with almost explosive violence into nitric oxide and water — NH 2 0H*HN0 3 = 2NO + 2H 2 0. AMMON-SULPHONATES. These compounds may be regarded as derived from ammonia, by the gradual replacement of the hydrogen by the group S0 3 H or S0 2 0II. Ammon-sulphonic acid . . . NH 2 (S0 3 H). Ammon-disulphonic acid . . NH(S0 3 H) 2 . Ammon-trisulphonic acid . N(SO s H) s . 248 Inorganic Chemistry Potassium ammon-trisulphonate is precipitated as a crystalline salt when excess of a solution of potassium sulphite is added to a solution of potassium nitrite — 3K a SO a + KNO a + 2H a O = 4KHO + N(SO a K) a . Prolonged boiling with water converts it first into the ammon-disulphonate — N(SO a K) a + H 2 0 = NH (S0 8 K) 2 + HKS0 4 , and finally into ammon-sulphonate — NH(SO a K) 2 + H 2 0 = NH 2 (SO a K) + HKS0 4 . Ammon-sulphonic acid is a stable crystalline body ; the other two acids are only known in their salts. When an ice-cold solution of sodium nitrite is added to hydrogen sodium sulphite, a compound is obtained which may be regarded as derived from ammon-trisulphonic acid by the replacement of one of the groups, SO,,H. by hydroxyl, OH — NaN0 2 + 2NaHSO a = N(OH)(SO a Na) 2 + NaHO. On the addition of a saturated solution of potassium chloride, in the cold, the sodium salt is converted into the potassium salt, which slowly crystallises from the solution, with two molecules of water, N(OH)(SO a K) 2 ,2H 2 0. This potassium hydroxylamine disulphonate is an unstable compound, and on boiling with water the two SO a K groups are replaced by hydrogen, forming first potassium hydroxylamine monosulphonate, NH(OH)SO a K; and finally hydroxylamine, NH 2 OH. Compounds of Nitrogen with the Halogen Elements. Nitrogen Trichloride, NCl a . — This compound was discovered by Dulong (1811). Its true composition was proved by (\attermann (1888). Mode of Formation. — Nitrogen trichloride is obtained by the action of chlorine upon ammonium chloride — NH 4 C1 + 3C1 2 = NC1 3 + 4HC1. When a solution of ammonium chloride is electrolysed, the chlorine, which is evolved at the positive electrode, acts upon the ammonium chloride, forming the trichloride of nitrogen.* Properties. — Nitrogen trichloride is a thin oily liquid, of a pale yellow colour, and having a specific gravity of 1.65. It is very volatile, and has an unpleasant pungent smell, the vapour being extremely irritating to the eyes. It is the most dangerously explosive compound known, and when suddenly heated, or brought into contact with grease, turpentine, or phosphorus it at once explodes. It also explodes on exposure to sunlight. At a temperature of 71 0 it may be distilled, but the operation is one of the utmost danger. Nitrogen trichloride is decomposed by ammonia, forming ammonium chloride and tree See “Chemical Lecture Experiments,” new ed. , No. 301. 249 Nitrogen Iodide nitrogen ; hence in the preparation of nitrogen by the action of chlorine upon ammonia, the presence of an excess of ammonia prevents the formation of this dangerous compound. Nitrogen Tribromide, NBr 3 . — When potassium bromide is added to nitro- gen trichloride beneath water, a red, oily, highly explosive substance Is obtained, believed to be the tribromide of nitrogen. Nitrogen Iodide, NHI 2 . — When strong aqueous ammonia is added to powdered iodine, a brown-coloured powder is formed which has violently explosive properties.* Also when alcoholic solutions of iodine and of ammonia are mixed, a brown and highly explosive compound is produced. The com- pound has generally been called iodide of nitrogen, or nitrogen iodide, but its exact composition has been a matter of dispute, and it is still an unsettled point whether the compound obtained by the different methods is identical or not. Courtois, who tirst prepared the substance, believed it to have the composi- tion NI 3 , and this view was held by Gay-Lussac and others. Gladstone and others considered that the substance contained one atom of hydrogen, and that the formula NHI 2 expressed the composition. The recent investigations of Szuhay (1893) prove indubitably that the compound obtained by the addition of an excess of aqueous ammonia to a concentrated solution of iodine in potassium iodide, has the composition NHI 3 . Whether under different experimental conditions other compounds are capable of formation, in which either more or less of the hydrogen has been substituted by iodine, such as NI 3 or NH 2 I, still remains to be proved. The reaction by which the com- pound is produced may be thus expressed — 3 NH 3 + 2 I 2 = NHI 2 + 2 NH 4 1 . Properties. — Nitrogen iodide is a dark chocolate-brown powder. When moist it may be handled without much risk of explosion, although it has been known to explode even under water. When dry, the substance is extremely explosive, the shock caused by the tread of a fly upon it, is more than sufficient to explode it ; even falling dust particles will sometimes cause it to explode. To illustrate its explosive nature, a small quantity of the wet substance should be allowed to dry spontaneously upon a wooden block : in a few hours it will have become sufficiently dry to readily explode if lightly touched by a feather. When nitrogen iodide is placed in dilute aqueous ammonia, and exposed to bright light, it is decomposed, and bubbles of nitrogen are seen escaping from the compound. ) Experiment 302. CHAPTER VIII CARBON Symbol, C. Atomic weight = 11.97. Occurrence. — This element is capable of assuming three allo- tropic forms, and it occurs free in nature in each of these modifica- tions, viz., diamond, graphite, and charcoal. In combination with oxygen, carbon occurs in carbon dioxide, a gas which is present in the air, being a constant product of com- bustion and respiration. In combination with hydrogen it occurs as marsh gas. Carbon is a constituent of all the natural car- bonates, such as limestone, dolomite, &c., which form an important fraction of the earth’s crust, and it is also an essential constituent of all organic substances. Diamond. Occurrence. — This substance has been known and prized from the remotest antiquity. It is found in various parts of India, mostly in river gravels and superficial deposits, in Brazil, South Africa, Australia, and various parts of the United States. The diamond has also recently been obtained from extra-terrestrial sources. In a meteorite which fell in Russia on September 22, 1886, carbon was found, partly as amorphous, and partly as ada- mantine carbon. The diamond form of carbon is found of various colours ; some- times it is dark grey, or even black, stones of these colours being known as carbonado and bort. The former of these is extremely hard, and is of great value for use in rock-boring and drilling instruments. Bort is used in the crushed condition by lapidaries for grinding and polishing. Occasionally the diamond is found coloured blue, or red, or green, by traces of foreign materials. Some of these coloured Carbon 25 stones are of great value as gems: the well-known “Hope” diamond, a stone weighing 44^ carats, has a fine sapphire colour. The origin of the diamond is unknown, although many theories have been put forward to explain its formation. Newton’s famous suggestion that diamond was “ an unctuous substance coagulated,” was based upon its remarkably high refractive index. The cellular structure which is sometimes to be seen in the ash that is left when the diamond is burnt, seems to indicate that it is of vegetable origin. Modes of Formation. — Innumerable attempts have been made to effect the crystallisation of carbon in the adamantine form ; but while it is readily possible to convert this variety of carbon into its allotropes graphite and charcoal, the transformation of these back again to the diamond, is a problem that is beset with the greatest difficulties. Moissan has recently shown * that the carbon which is capable of being dissolved in molten iron, and which is usually deposited in the graphitic form on cooling, can, under certain conditions, be caused to take up the adamantine form. By raising the temperature of the iron to about 3000° by means of an electric furnace, and then suddenly cooling the molten mass by plunging the crucible into water or molten lead, until the cooled and solidified surface is at a dull red heat, an enormous pressure is brought to bear upon the interior and still liquid portion. Under these circumstances, a part of the carbon which is deposited by the slowly cooling mass, was found by Moissan to be in the adamantine form. On dissolving the iron in hydrochloric acid, amongst the carbonaceous residue were found fragments having a specific gravity between 3.0 and 3.5, and sufficiently hard to scratch ruby. Some of the fragments were the black or carbonado variety, while others were transparent. On combustion in oxygen, Moissan proved that these were really carbon in the diamond form. Properties. — The diamond in its purest condition is a colourless crystalline substance. Its crystalline forms belong to the cubic system, and appear to some extent to be characteristic of the locality in which the element occurs. It is extremely hard, and moderately brittle. When struck with a hammer, the diamond not only splits along its cleavage-planes, but also in other directions, with a conchoidal fracture. It does not conduct electricity. The specific gravity of diamond varies slightly in different specimens, Comptes Rendus de /' Acadtmie des Sciences, vol. cxvi. p. 218. 252 Inorganic Chemistry the mean being about 3.5. Its refractive index is higher than that of any other substance, and it is this property which gives its peculiar beauty and brilliancy to the cut stone. The value of diamond as a gem, depends largely upon its colour- lessness, except in the case of those rare instances where the colour is quite definite and also pleasing, such as distinct red, blue, or green. When diamond is strongly heated it becomes black, and in- creases in bulk, being converted into a substance having the properties of graphite. Lavoisier (1772) was the first to show that the diamond was a combustible body, and that.it yielded carbon dioxide. Davy (1814) showed that carbon dioxide was the only product of its combustion, and proved that diamond was pure carbon. The combustion of diamond in oxygen may readily be accomplished by means of the apparatus shown in Fig. 58. A fragment of diamond is supported upon a small gutter of platinum foil, which bridges across two stout copper wires, A. These wires pass through a cork in a perforated glass plate, and are lowered into a cylinder of oxygen. By the passage of an electric current, the little platinum boat can be strongly heated, when the diamond will become ignited, and continue to burn brilliantly in the oxygen, with the formation of carbon dioxide. The ash, which is always left after a diamond has been burnt, varies from 0.2 to 0.05 per cent, of the stone. It is found usually to contain ferric oxide and silica. Graphite. Occurrence. — This second allotrope of carbon is much more plentiful in nature than the first. It is found in large quantities in Siberia, Ceylon, and various parts of India. In England the chief source of graphite has been the mines at Borrowdale, in Cumberland ; this supply is now practically exhausted. Enor- mous quantities of very pure graphite are now obtained from the Eureka Black-Lead Mines, in California. Graphite also occurs in many specimens of meteoric iron. Mode of Formation. — Molten iron, especially when it contains Fig. 58 . Carbon 253 silicon, is capable of dissolving a considerable amount of carbon, which, on cooling, is deposited in the form of black shining crystals of graphite. Occasionally considerable quantities of graphite are found deposited in this way in iron-smelting furnaces, to which the name “ kish ” has been applied. Properties. — Graphite is a soft, shiny, greyish-black substance, which is smooth and soapy to the touch. It is usually found in compact laminated masses, but sometimes crystallised in six-sided plates. Its specific gravity varies in different specimens, aver- aging about 2.5. Graphite is a good conductor of both heat and electricity. When strongly heated in oxygen, graphite takes fire and burns, forming carbon dioxide, and leaving an ash consisting of silica, alumina, and oxide of iron. Graphite has been found by Regnault to contain, usually, traces of hydrogen. Graphite is employed for the manufacture of ordinary lead pencils ; for, on account of its soft- ness, it leaves a black mark upon paper when drawn across it. For the purposes of the pencil manufacture the natural graphite is ground to powder, and carefully washed free from gritty matter. It is then mixed with the finest washed clay, and the pasty mass is forced by hydraulic pressure through perforated plates. The name “ graphite,” from the Greek to write , is given to this substance on account of its use for this purpose. It was formerly supposed that this material contained lead, hence the names black-lead and plumbago. Graphite is largely employed, on account of its refractoriness, for the manufacture of the so-called plumbago crucibles, which consist of fireclay, mixed with finely-ground graphite. Other uses to which graphite is put, are for glazing or polishing gunpowder, especially the larger grained varieties ; as a lubricant for machinery, where oil is inadmissible on account of high tem- perature ; for electrotyping processes, and also as a coating for ironwork, to prevent rusting. Amorphous Carbon. This non-crystalline form of carbon, may be obtained by the decomposition of a great variety of carbon compounds, by the process known as destructive distillation. The carbon so obtained differs very much as regards its purity, according to the particular organic compound used for its preparation. The commonest forms 254 Inorganic Chemistry of amorphous carbon to be met with, are lampblack, or soot ; gas carbon ; coke ; charcoal ; animal charcoal, or bone-black. None of these substances is pure carbon ; animal charcoal, for example, usually containing only about io per cent, of carbon. Lampblack. — This substance is manufactured by burning sub- stances rich in carbon, and which burn with a smoky flame, (as turpentine, petroleum, or tar,) with a limited supply of air. The smoke is passed into chambers in which are suspended coarse blankets, upon which the soot collects. This lampblack always contains hydrogen, in the form of hydrocarbons. If the soot be heated to redness in a stream of chlorine, this hydrogen can be removed, and pure amorphous carbon will be left. Lampblack is used for printers’ ink, and for black paint. Gas Carbon. — This form of carbon is obtained by the destruc- tive distillation of coal, in the manufacture of illuminating gas. It remains in the retorts as an extremely hard deposit, lining the roof and sides. It is a very pure carbon, coming second to puri- fied lampblack. Its specific gravity is about 2.35. Gas carbon is a good conductor of electricity, and is extensively used for the manufacture of carbon rods for the arc light. Coke. — This substance differs from gas carbon, although it also is obtained in the process of coal distilling. It contains all the inorganic matter which constitutes the ash of the coal, and also small quantities of hydrogen, nitrogen, and oxygen. The average amount of carbon in coke is about 91 per cent. Charcoal. — The purest form of charcoal is obtained by the carbonisation of pure white sugar, and the subsequent ignition of the charcoal in a stream of chlorine gas. Charcoal so obtained has a specific gravity of 1.57. Charcoal in a much less pure con- dition is manufactured from wood. The methods by which the carbonisation of wood is carried out are, broadly, of two kinds : first, those in which access of air is permitted to the burning material ; and, second, those in which air is excluded. The first of these, and the most ancient, is generally carried on in more primitive parts, where wood is plentiful. The wood is piled into mounds or stacks, which are built with some care. They are set on fire in the interior, by means of a lighted bundle of brushwood, which is introduced through a vertical opening or chimney, left for this purpose in the centre of the mound during its construction. The outside of the heap is covered with brushwood, and finally with turf, in order to regulate the access of air to the Charcoal 255 interior, and therefore to control the rate of combustion of the wood. The object of the charcoal-burner is to carbonise the wood as slowly as possible. In this process there is great liability to loss, by the too rapid combustion of the wood ; and, in addition, it pos- sesses the disadvantage that the secondary products, such as the pyroligneous aqid, tar, &c., are entirely lost. Various modifications have been introduced into the method of coaling wood, as the process is termed, with a view to collect these products. In the second general process of carbonising wood, the material is heated in ovens or retorts from the outside, no air being admitted to the wood. The operation is very similar to that employed in the destructive distillation of coal, in the manufacture of coal gas. In these methods all the volatile and condensable products are collected ; among these are water, pyroligneous acid, wood spirit, acetone, and fatty oils. The non-condensable products consist mainly of such gases as hydrogen, carbon monoxide, carbon di- oxide, marsh gas, and acetylene. Animal Charcoal. — Bone-black is obtained by the carbonisa- tion of bones in iron retorts. This variety of charcoal is the least pure of all the ordinary forms of amorphous carbon. Bone contains only about 30 per cent, of organic matter, the other 70 per cent, consisting chiefly of calcium phosphate, asso- ciated with small quantities of magnesium phosphate and calcium carbonate. It will be obvious, therefore, that as the carbon is derived from the organic matter, the amount of it in carbonised bones must be small. The average composition of animal char- coal is found to be — Carbon Calcium phosphate . Other saline substances . 10.0 . 88.0 2.0 100.0 Although containing relatively so small an amount of carbon, animal charcoal possesses many of the valuable properties of charcoal in a highly marked degree, owing to the fact that it con- tains its carbon disseminated throughout an extremely porous mass of calcium phosphate. Properties of Chareoal. — Charcoal varies very considerably in its properties, depending upon the particular wood from which it 256 Inorganic Chemistry is obtained, and the method by which it is prepared. Thus, char- coal obtained at 300° is a soft, brownish-black, very friable material, having an igniting point as low as 380°. On the other hand, charcoal prepared at very high temperatures is black and com- paratively dense, and requires to be strongly heated in order to ignite it. Under ordinary circumstances, charcoal burns in air without the formation of a flame, or the production of smoke. At high tem- peratures, however, the combustion of charcoal is seen to be attended by a flame. This is probably accounted for by the fact, that as the temperature at which the combustion of carbon takes place is raised above 700°, the amount of carbon monoxide which is formed increases, and the carbon dioxide decreases.* When charcoal is thrown upon water it floats, on account of the air which is enclosed within its pores. The specific gravity of charcoal when thus filled with air, varies from 0.106 (charcoal made from the ash) to 0.203 (charcoal from the birch). If the air be withdrawn from charcoal it sinks in water, the average specific gravity of charcoal itself being 1.5. Ordinary charcoal is a bad conductor of electricity, but its con- ductivity is greatly increased by strongly heating the charcoal in closed vessels. Charcoal has the power of absorbing gases and vapours to a remarkable extent : this power, which is exhibited to a different degree by the various kinds of charcoal, is due to the porosity of the material, whereby it exposes a very large surface ; and it belongs to a class of phenomena known as surface action. If a fragment of charcoal, recently strongly heated to expel the air from its pores, be passed up into a cylinder of ammonia gas, standing in a trough of mercury, the ammonia will be gradually absorbed by the charcoal, and the mercury will ascend in the cylinder. Saussure found, that recently heated beech-wood char- coal was capable of absorbing ninety times its own volume of ammonia gas ; while Hunter, by employing charcoal made from cocoa-nut shell, found that 17 1.7 volumes of ammonia were absorbed by one volume of charcoal. The results of both of these experi- ments show, that those gases are absorbed in the largest quantities which are the most readily liquefiable. The gas so held by the charcoal, is in a highly condensed condition upon the surface of Ernst [Cfiemischcs Repertorium, vol. xvii. p. 2). Coal 25 7 the porous mass. Probably in the case of easily liquefied gases such as ammonia, sulphur dioxide, and others, the gases are par- tially liquefied upon the surface of the charcoal. In this condensed state, the gas is more chemically active than under ordinary condi- tions, and charcoal is therefore able to induce many striking com- binations to take place. Thus, if charcoal be allowed to absorb chlorine, and dry hydrogen be then passed over it, the chlorine is capable of combining with the hydrogen even in the dark, with the formation of hydrochloric acid. This chemical activity of gases, when absorbed by charcoal, is strikingly exemplified in the case of sulphuretted hydrogen. If a quantity of powdered charcoal, which has been saturated with sulphuretted hydrogen, be brought into oxygen, the rapid combination of the two gases is attended with the development of so much heat, that the charcoal bursts into active combustion. In the same way a mixture of air, with 10 or 15 per cent, of sulphuretted hydrogen, may be passed rapidly through a tube, about a metre in length, filled with char- coal, without a trace of sulphuretted hydrogen escaping at the end.* Owing to this property, charcoal is largely employed to absorb noxious gases, the atmospheric oxygen which is con- densed in the pores of the charcoal, oxidising these offensive and injurious compounds : thus sewer ventilators are often trapped with a layer of charcoal, which effectually arrests all bad-smelling gases. Charcoal also has the power of absorbing colouring-matters from solution : thus, if water which has been tinted with an organic colouring-matter, be shaken up with powdered charcoal, and filtered, the solution will be found to be entirely decolourised. The variety of charcoal which possesses this property in the highest degree, is animal charcoal, or bone-black, and this substance is largely em- ployed in many manufacturing processes, such as sugar refining, in order to remove all colouring-matter from the liquid. Charcoal under ordinary conditions is unacted upon by the air but when the temperature is raised, it enters into active combus- tion, forming carbon dioxide. In an extremely divided condition, however, carbon is capable of combining spontaneously with the oxygen of the air, and with so much energy as to take fire. Coal.— The carbonaceous minerals that are included under the name coal , are an impure form of carbon, containing compounds * “Chemical Lecture Experiments,” 394-396, new ed. R 258 Inorganic Chemistry of carbon with hydrogen and oxygen. Coal is the final result of a series of decomposition changes, which have been undergone by vegetable matter of the remote past, the process having extended over long geological periods. During this prolonged process, a por- tion of the carbon and hydrogen is eliminated as marsh gas ; and large quantities of this gas are found associated with, and occluded in, coal. Broadly speaking, the numerous varieties of coal may be divided into soft or bituminous , and hard or anthracitic. The former are employed for the manufacture of coal gas, and for ordinary domestic purposes ; they burn with a smoky flame, and evolve large quantities of gases and volatile vapours, on combus- tion or distillation. Anthracite coal is much harder, ignites with more difficulty, and burns with the formation of very little flame or smoke. It contains a higher percentage of carbon, and gives out great heat on combustion, and is employed largely as a steam- coal. The following table shows the average composition of coals from various sources, and the general difference between coals of the two main classes — Locality. Carbon. Hydrogen. Oxygen. Nitrogen. Sulphur. Ash. Water. Coke. C /3 ( N orthumberland 81.41 5-83 7.90 2.05 O.74 2.0 7 i -35 66.70 Bitu- linou Wales 83.78 4-79 4 - IS 0.98 1-43 4.9I 72.60 5 | [Staffordshire 78.57 5-29 12.88 1.84 0-39 I.03 11.29 57-21 1 2*j |S. Wales . 92.56 3-33 2-53 i-58 co ] 'll Pennsylvania 90-45 2.43 2-45 4.6 CHAPTER IX CARBON COMPOUNDS The compounds of the element carbon are so numerous, that it has been found convenient to constitute the study of these sub- stances a separate branch of chemistry. In the early history of the science, it was believed that there were a large number of substances which could only be obtained as the product of living organisms. They were known to be elaborated by the action of life, or, as it was termed, the vital force , and it was believed, that owing to some inherent specific quality belonging to this vital force, the substances produced by its action were distinct from such substances as could be prepared by any laboratory processes. To denote this distinction, the term organic was applied to those things which were known to be the products of living organisms, and other compounds were distinguished as inorganic substances. This distinction received its deathblow in 1828, when Wohler produced, by purely laboratory processes, one of the most typical of all organic compounds, namely, urea. The names “ organic ” and “ inorganic ” chemistry are still retained, but their old significance is entirely gone, as no distinction is to-day recognised between products elabo- rated by the action of life, and those which can be synthetically produced. Speaking broadly, organic chemistry may be defined as the chemistry of the carbon compounds. This definition, however, includes such compounds as the oxides of carbon and all the carbonates, such as chalk, limestone, dolomite, &c., compounds which are more conveniently classed as inorganic. A more exact definition is the following: — The chemistry of those compounds of carbon which contain in the molecule an atom of carbon, directly associated with either hydrogen , nitrogen , or another carbon atom. This definition excludes all the carbon compounds which are by general consent regarded as belonging to the inorganic division of the science. 260 Inorganic Chemistry Three organic compounds will be briefly studied in the following chapters, namely, methane (marsh gas), CH 4 ; ethene (ethylene), C 2 H 4 ; and acetylene, C 2 H 2 . These three compounds play an important part in our ordinary illuminating flames, and in coal gas. Compounds of Carbon with Oxygen. Two oxides of carbon are known, both of which are colourless gases, viz. : — Carbon monoxide CO. Carbon dioxide CO<*. CARBON MONOXIDE. Formula, CO. Molecular weight = 27.96. Density = 13.98. Modes of Formation. — (i.) Carbon monoxide is formed when carbon dioxide is passed over charcoal heated to bright redness — C0 2 + C = 2CO. The same result is obtained, when a slow stream of air or oxygen is passed over red-hot charcoal contained in a tube. The first action of the air, on coming in contact with the carbon, is to form carbon dioxide, which, passing over the remainder of the heated material, is deprived of a portion of its oxygen, according to the above equation. This operation goes on in an ordinary fire-grate ,* the air, on first gaining access to the burning coal or coke, causes the complete oxidation of a portion of the carbon, to carbon dioxide ; and as this gas passes through the mass of red-hot carbon, it is reduced to the lower oxide, which either escapes with the other products of combustion, or becomes ignited and bums with a lambent bluish flame, such as may frequently be noticed upon the top of a “ clear ” fire. ( 2 .) When steam is passed over strongly heated carbon, a mixture of carbon monoxide and hydrogen is produced. This mixture, known as water gas, is employed in many manufacturing processes as a gaseous fuel — H 2 0 + C = CO + H* Carbon Monoxide 261 (3.) Carbon monoxide is also formed by the action of carbon dioxide upon red-hot iron — 4CO a + 3Fe = Fe 3 0 4 + 4CO. (4.) Or by strongly heating either carbon or iron with a car- bonate, such as calcium carbonate, which is capable of yielding carbon dioxide, thus — CaCO a + C = CaO + 2CO. 4CaCO a + 3Fe = Fe 3 0 4 + 4CaO + 4CO. (5.) Carbon monoxide is most conveniently prepared, by the decomposition of certain organic compounds by means of sulphuric acid. Thus, when formic acid, or a formate, is acted upon by sul- phuric acid, the sulphuric acid withdraws the elements of water from the molecule of formic acid, and leaves carbon monoxide — H-COOH - H 2 0 = CO. (6.) By a similar decomposition, oxalic acid yields a mixture of carbon monoxide and carbon dioxide in equal volumes — c 2 h 2 o 4 - H a O = co 2 + CO. The carbon dioxide is readily removed from the mixture, bypassing the gases through a solution of sodium hydroxide (caustic soda), in which carbon dioxide is absorbed with the formation of sodium carbonate. (7.) The method usually employed when carbon monoxide is required for experimental purposes, consists in heating a mixture of 1 part by weight of crystallised potassium ferrocyanide (yellow prussiate of potash) with ten parts of strong sulphuric acid in a capacious flask, when the following reaction takes place — K 4 FeC 6 N„ + 6H 2 S0 4 + 6H 2 0 = 2K 2 S0 4 + FeS0 4 + 3(NH 4 ) 2 S0 4 + 6CO. The six molecules of water required by the reaction, are derived partly from the acid employed, and partly from the salt, which contains three molecules of water of crystallisation.* Properties. — Carbon monoxide is a colourless, tasteless gas, * "Chemical Lecture Experiments,’’ new ed., 435-439. 262 Inorganic Chemistry having a faint smell. It is only slightly soluble in water, its co- efficient of absorption at o" being 0.03287. It bums in the air with a characteristic pale blue flame, forming carbon dioxide— 2CO + O a = 2CO2. When mixed with half its own volume of oxygen, and inflamed, the mixture explodes with some violence.* If the two gases be confined in a eudiometer standing over mercury, and be rendered absolutely free from aqueous vapour by powerful desiccating agents, no explosion will take place upon the passage of an electric spark through the mixture. And in the same way if carbon monoxide, which has been deprived of all aqueous vapour, be burned from a jet in the air, and the jet be lowered into a cylinder containing air which has been similarly dried, the flame will be extinguished. Carbon monoxide is an extremely poisonous gas : very small quantities present in the air, rapidly give rise to headache and giddiness, and if inhaled for a length of time, or if taken into the lungs in a less dilute condition, insensibility and death quickly follow. The deaths that have resulted from the use of unventi- lated fires — either of charcoal or coke, or in some cases of coal gas —in dwelling-rooms, have been due to the escape of this poisonous gas into the air. The extremely deadly nature of the after-dam $ resulting from a colliery explosion, is due to the presence of carbon monoxide in the carbon dioxide which is formed as a product of the combustion. The poisonous action of this gas is due to its absorption by the blood, with the formation of a bright red compound, to which the name carboxy-hamo- globin is applied. Blood so charged appears to be unable to fulfil its function of absorbing and distributing oxygen throughout the system. This carboxy- haemoglobin gives a characteristic absorption spectrum, which furnishes a ready method of detection, in cases of poisoning from the inhalation of carbon- monoxide. Carbon monoxide is one of the most difficultly liquefiable gases, its critical temperature being — 140°. At high temperatures this gas is a powerful reducing agent, uniting with another atom of oxygen to form carbon dioxide. This fact is made use of in many metallurgical processes, for re- ducing the oxides of the metals to the metallic state. * The rate at which the combustion is propagated throughout a mixture of carbon monoxide and oxygen, is much slower than through hydrogen and oxygen. Bunsen has estimated it at less than 1 metre per second. Carbon Monoxide 263 Carbon monoxide is absorbed at ordinary temperatures by a solution of cuprous chloride, forming the compound COCu 2 Cl 2 . At a temperature of boiling water, carbon monoxide is slowly absorbed by solid potassium hydroxide, with the formation of potassium formate — KHO + CO = H COOK. Carbon monoxide unites directly with chlorine, under the in- fluence of sunlight, forming the compound known as phosgene gas, or carbonyl chloride — . CO + Cl 2 = COCl* If the two gases are mixed in equal volumes, and kept in the dark, no action takes place, but on exposure to sunlight they com- bine, and the yellowish colour due to the chlorine will disappear. On opening the vessel in moist air, clouds of hydrochloric acid are formed, owing to the decomposition of carbonyl chloride by the moisture, according to the equation— COCl 2 + H 2 0 = CO a + 2HC1. Carbonyl chloride may be readily condensed to a liquid, its boiling-point being +8°. Carbon monoxide unites directly with certain metals, giving rise to compounds which possess some very remarkable properties, and to which the name metallic carbonyls has been applied by their discoverer.* When carbon monoxide is allowed to stream slowly over metallic nickel (obtained by the reduction of nickel oxide in a stream of hydrogen), the gas is absorbed by the finely-divided metal, forming a compound having the composition Ni(CO) 4 . If the issuing gas be passed through a cooled tube, the nickel carbonyl condenses to a colourless, mobile, highly refracting liquid, having a specific gravity at o° of 1.356, and boiling at 43° under a pressure of 751 mm.t Nickel carbonyl vapour burns with a luminous flame, which produces a black deposit of metallic nickel when a cold porcelain dish is depressed upon the flame. The gas is decomposed into nickel and carbon monoxide if passed through a hot glass tube, * Mond, 1890. f See ' 1 Chemical Lecture Experiments,” new ed. , 446-448. 1 2 ^4 Inorganic Chemistry the nickel being deposited as a bright metallic mirror upon the glass — Ni(CO) 4 = 4 CO + Ni. A similar compound of carbon monoxide and iron has also been obtained having the composition Fe(CO) s . Iron carbonyl is a pale- yellow, viscous liquid, boiling at 102. 8° under a pressure of 749 mm. Its specific gravity at 18 is 1.4664. When heated to 180° the vapour is decomposed, iron being deposited and carbon monoxide being evolved. This compound has been found in iron cylinders in which the so-called water gas (a mixture of H and CO) has been stored under pressure for a length of time ; it is also said to be present in minute quantities in coal gas. CARBON DIOXIDE. Formula, CO a . Molecular weight = 43.92. Density = 21.96. History. —Van Helmont, in the seventeenth century, was the first to distinguish between this gas and ordinary air : he observed that it was formed during the processes of combustion and fermentation, and he applied to it the name gas sylvestre. Black showed that this gas was a constituent of what in his day were known as the mild alkalis (alkaline carbonates), and on account of its being so combined, or fixed, in these substances, he named the gas fixed air. Lavoisier first proved its true chemical composition to be that of an oxide of carbon. Occurrence.— Carbon dioxide is a constant constituent of the atmosphere, being present to the extent of about 3 volumes in 10,000 volumes of air. It is also found in solution in all spring- water, which is sometimes so highly charged with this gas under pressure, that the water is effervescent, or “ sparkling,” from the escape of the gas. Carbon dioxide is evolved in large quantities from vents and fissures in the earth in volcanic districts. The well-known Poison Valley in Java, which is an old volcanic crater, and the Grotto del Cane near Naples, owe their peculiar pro- perties to the discharge into them of large quantities of carbon dioxide from such subterranean sources. Modes of Formation. — (i.) Carbon dioxide* is produced when carbon is burnt with a free supply of air or oxygen — C + o 2 = C 0 2 . * Experiments on carbon dioxide, Nos. 400-434, “Chemical Lecture Ex- periments,” new ed. Carbon Dioxide 265 If an insufficient supply of oxygen be employed, carbon mon- oxide is produced at the same time. (2.) When limestone, or chalk, is strongly heated, as in the process of burning lime, carbon dioxide is evolved in large quantities— CaC0 3 = CaO + COj. (3.) In the ordinary processes of fermentation, and during the decay of many organic substances, carbon dioxide is also formed. Thus, when sugar undergoes alcoholic fermentation by means of yeast, the sugar is converted into alcohol and carbon dioxide — Ci 2 H 22 O n + H 2 0 = 4C 2 H 6 0 + 4C02. (4./ Carbon dioxide is formed during the process of respiration ; also by the combustion of all ordinary fuels, and of any compound containing carbon, such as candles, oils, gas, &c. (5.) For experimental purposes, carbon dioxide is most readily obtained by the decomposition of a carbonate by means of a stronger acid. The effervescence that results from the action of tartaric acid upon sodium carbonate, in an ordinary Seidlitz powder, is due to the disengagement of this gas. The most convenient carbonate for the preparation of this gas is calcium carbonate, in one of its many naturally occurring forms, such as marble, lime- stone, or chalk. Fragments of marble are for this purpose placed in a two-necked bottle (Fig. 59), with a quantity of water, and 1 266 Inorganic Chemistry strong hydrochloric acid is added by means of the funnel-tube. A rapid effervescence takes place owing to the elimination of the gas, and a solution of calcium chloride remains in the bottle— CaCO s + 2HC1 = CaCl 2 + H 2 0 + CO * If sulphuric acid be substituted for hydrochloric acid, the frag- ments of marble rapidly become coated with a crust of insoluble calcium sulphate, which soon prevents the further action of the acid, and therefore puts an end to the reaction : by employing finely powdered chalk, however, instead of lumps of calcium car- Fig. 6o. Fig. 6i. bonate, this difficulty is obviated. This gas is largely manufac- tured from these materials. Properties. — Carbon dioxide is a colourless gas, having a feeble acid taste, and a faint and pleasantly pungent smell. It is incap- able of supporting either combustion or respiration ; a burning taper is instantly extinguished, and an animal speedily dies when introduced into this gas. Although carbon dioxide is not such a poisonous compound as the monoxide, it nevertheless does exert Carbon Dioxide 267 a direct poisonous effect upon the system, and death caused by this gas is not merely due to the absence of oxygen. The pro- longed inhalation of air containing only a very slightly increased amount of carbon dioxide, has a distinctly lowering effect upon the vitality. Carbon dioxide is a heavy gas, being about one and a half times heavier than air. On this account it may readily be col- lected by displacement. By virtue of its great density it may be poured from one vessel to another, much in the same way as an ordinary liquid : thus, if a large bell jar be filled with the gas by displacement, a beaker-full may be drawn up, as water from a well (Fig. 60). If the gas so drawn up be poured into a similar beaker, suspended from the arm of a balance, and counterpoised, the weight of the gas will be evident by the disturbance of the equilibrium of the system. If a soap bubble be allowed to fall into a large jar filled with carbon dioxide, it will be seen to float upon the surface of the dense gas (Fig. 61). The power of carbon dioxide to extinguish flame is so great, that a taper will not bum in air in which this gas is present to the extent of 2.5 per cent., and in which the oxygen is reduced to 18.5 per cent. For this reason a comparatively small quantity of carbon dioxide, brought into the air surrounding a burn- ing body, is capable of extinguishing the flame. This property has been put to valuable service in the construction of numerous con- trivances for extinguishing fire, such as the “ extincteur.” This is a metal vessel containing carbon dioxide under pressure, the gas having been generated within the closed apparatus by the action of dilute sulphuric acid upon sodium carbonate. A stream of the gas, projected judiciously upon a moderate conflagration in a dwelling, readily extinguishes the fire. This property may be illustrated by inflaming a quantity of turpentine in a dish, and pouring upon the flames a quantity of carbon dioxide contained in a large bell jar (Fig. 62), when it will instantly extinguish the conflagration. Although carbon dioxide is incapable of supporting combus- tion in the ordinary sense, certain metals are capable of burn- ing in this gas. Thus, a fragment of potassium when heated in this gas, burns brightly, forming potassium carbonate with the deposition of carbon — 2 K 2 + 3 CO a = 2 K 2 C 0 3 + C. 268 Inorganic Chemistry When carbon dioxide is passed into a solution of calcium hydroxide (lime water) a turbidity at once results, owing to the precipitation of insoluble calcium carbonate, or chalk — CaH 2 0 2 + CO a = CaCO s + H 2 0 . This reaction furnishes the readiest means for the detection of carbon dioxide. Thus, if the gas obtained by any of the modes of formation described, be passed into clear lime water, the formation of this white precipitate of chalk, is proof that the gas is carbon dioxide. By this test it may readily be shown that carbon dioxide is a product of respiration, by merely causing the exhaled breath Fig. 62. to bubble through a quantity of lime water, which will quickly be rendered turbid. Carbon dioxide is moderately soluble in water. At the ordinary temperature, water dissolves about its own volume of this gas. The coefficient of absorption at o° is 1.7967, the solubility de- creasing with rise of temperature in accordance with the ’inter- polation formula — c = 1.7967 - 0.07761/ + 0.00 1 6424/* Carbon dioxide shows a slight departure from Hemys law (see page 123), when the pressures are greater than that of the atmosphere. Thus, when the pressure is doubled, the amount dis- Carbon Dioxide 269 solved is slightly more than doubled. The solubility of carbon dioxide in water, and its increased solubility under pressure, is illustrated in the ordinary aerated waters. Water under a pres- sure of several atmospheres is saturated with the gas, and upon the release of this pressure by the withdrawal of the cork, the excess of gas, over and above that which the water can dissolve at the ordinary pressure, escapes with the familiar effervescence. In a similar manner the natural aerated waters have thus become charged with carbon dioxide, under subterranean pressure, and when such waters come to the surface, the dissolved gas begins to make its escape. The solution of carbon dioxide in water is feebly acid, turning blue litmus to a port-wine red colour, characteristically different from the scarlet red given by stronger acids. This acid may be regarded as the true carbonic acid — C 0 2 4- H 2 0 = H 2 C 0 3 . A recently-made sample of aerated water is seen to effervesce more briskly, and give off the dissolved gas more rapidly, than specimens that have been long preserved. In process of time the dissolved carbon dioxide gradually combines with the water, with the formation of carbonic acid, an unstable compound which slowly decomposes into carbon dioxide and water, especially at a slight elevation of temperature. Many of the naturally occurring aerated waters, such as Apollinaris, when opened, exhibit scarcely any effervescence, but give off carbon dioxide gradually. Such waters have in all probability been exposed to pressure for a great length of time, and their dissolved carbon dioxide has almost entirely combined to form carbonic acid. When such a water is gently warmed, a rapid stream of gas is evolved. When carbon dioxide is strongly heated, as by the passage of electric sparks, it is partially dissociated into carbon monoxide and oxygen. This decomposition is never complete; for when the amount of these two gases in the mixture reaches a certain proportion, they reunite to form carbon dioxide, and a point of equilibrium is reached, when as many molecules are united as are dissociated in the same time. Liquid Carbon Dioxide. — Carbon dioxide is easily liquefied. At - 5 0 it requires a pressure of 30.8 atmospheres ; at + 5 0 , 40.4 atmospheres; while at + 15 0 a pressure of 52.1 atmospheres is required. 270 Inorganic Chemistry Faraday first liquefied this gas, by introducing into a strong bent glass tube a quantity of sulphuric acid, and a few lumps of ammo- nium carbonate, which were prevented from touching the acid by means of a plug of platinum foil. The tube was then hermetically sealed, and the acid allowed gently to come in contact with the carbonate, which was at once decomposed with the formation of ammonium sulphate and carbon dioxide. By the internal pres- sure exerted by the evolved gas, aided by the application of cold to one end of the bent tube, the gas condensed to a colourless liquid. Large quantities of this liquefied gas were obtained by Thilorier by a precisely similar method, the experiment being performed in strong wrought-iron vessels. Liquid carbon dioxide is to-day manufactured on a large scale, by pumping the gas into steel cylinders by means of powerful compression pumps. The enormous volumes of carbon dioxide evolved in the process of brewing, and which until quite recently were allowed to escape into the atmosphere, are now utilised fot this purpose. The gas, as it is evolved from the fermenting vats, is washed and purified, and pumped into steel bottles for the market. In this form the gas is largely employed by manufac- turers of aerated waters. Liquid carbon dioxide is a colourless and extremely mobile liquid, which floats upon water without mixing. It boils at - yS°.2 under atmospheric pressure. When heated, liquid carbon dioxide expands at a more rapid rate than a gas, its coefficient of expansion being greater than that of any known substance. Its rapid change of volume is seen by the following figures : — 95 volumes at - io° become ioo „ „ o „ I°6 a a + a 1 14 „ „ +20° The critical temperature of carbon dioxide is 31.9 0 . If the liquid be heated to this point, it passes into the gaseous state without any change of volume. The line of demarcation, between the liquid and gas in the tube, gradually fades away, and the tube appears filled with gas. Above this temperature no additional pressure is able to liquefy the gas. On once more cooling the tube, when the Carbon Dioxide 271 critical point is passed the liquid again appears, and the dividing line between it and the gas is once more sharply defined. Solid Carbon Dioxide. — When liquid carbon dioxide is allowed to escape into the air, the absorption of heat due to its rapid eva- poration causes a portion of the liquid to solidify. This solid is most conveniently collected, by allowing the jet of liquid to stream into a round metal box (Fig. 63), in which it is caused to revolve by being made to impinge upon the curved tongue of metal. The box is furnished with hollow wooden handles, through which the gas makes its escape. Large quantities of the frozen carbon dioxide can in this way be collected in a few minutes. Solid carbon dioxide is a soft, white, snow-like substance. When exposed to the air it quickly passes into gas, without going through the intermediate state of liquidity. Fig. 63. Solid carbon dioxide is readily soluble in ether, and this solution constitutes one of the most convenient sources of cold. A large number of gases can readily be liquefied by being passed through tubes immersed in this freezing mixture. When this ethereal solution is rapidly evaporated, its temperature can be lowered to - IIO°. “ Carbonic acid snow,” as this substance is sometimes termed, is now an article of commerce, the compound being sent into the market in this form, to avoid the cost of the carriage of the necessarily heavy steel bottles containing the liquid. Composition of Carbon Dioxide. — When carbon bums in '*72 Inorganic Chemistry oxygen, the oxygen undergoes no change in volume in being trans- formed into carbon dioxide. The volume of carbon dioxide produced, is the same as that of the oxygen which is required for its produc- tion. This may be shown by means of the apparatus, Fig. 64. The bulb of the U-tube is filled with oxygen, and the stopper, which carries a small bone -ash crucible, upon which a fragment of charcoal is placed, is lowered into position. The charcoal is ignited by means of a thin loop of platinum wire, as shown in the figure, which can be heated by an electric current. As the carbon burns, the heat causes a temporary expansion of the included gas ; but after the combustion is complete, and the apparatus has cooled, the level of mercury will be found to be undisturbed. Carbon dioxide, therefore, contains its own volume of oxygen. From this experiment the composition of carbon dioxide by weight can be deduced. One litre of carbon dioxide weighs 21.96 criths ; deducting from this the weight of 1 litre of oxygen, viz., 15.96 criths, we get 6 as a remainder. Six parts by weight of carbon, therefore, combine with 1 5.96 parts by weight of oxygen to form 21.96 parts of carbon dioxide : expressing this proportion atomic- ally, the proportion of carbon to oxygen is 12 to 31.92. The gravimetric composition of carbon dioxide may be directly determined, by the combustion of a known weight of pure carbon in a stream of oxygen gas, and absorbing and weighing the carbon dioxide that is formed. This was done with great care and accuracy by Dumas and Stas, in the experiments by which they determined the atomic weight of carbon. Fig. 65 represents the apparatus employed for this purpose. A weighed quantity of diamond, contained in a small platinum boat, was introduced into a porcelain tube, which could be strongly heated in a furnace. The oxygen for its combustion was contained in a glass bottle, from which it could be expelled by allowing water to enter through the funnel. As it was necessary that the oxygen should be absolutely free from any carbonic dioxide, the water used in the little gas- holder contained potassium hydroxide in solution. The oxygen was then passed through the tubes A, B, C in order to deprive it of carbon dioxide and moisture, and lastly through a small desic- cating tube, d , which was weighed before and after the experiment. The pure dry oxygen then entered the strongly heated tube, and the carbon there burnt away to carbon dioxide, leaving a minute quantity of ash, which was carefully weighed at the conclusion of the experiment. A small layer of copper oxide was placed in the Carbon Dioxide 273 indicated in the figure, in order to oxidise any tube, in the position traces of carbon monoxide, which were liable to be formed, into the dioxide. The pro- duct of the combustion was carried forward by the stream of oxygen, through a series of tubes ; d' is a small weighed desiccating tube, the weight of which, if the diamond used contained no hydrogen, should remain unchanged. It then passes through the bulbs F and G, where the carbon di- oxide is entirely absorbed. To arrest aqueous vapour which would be carried away from the solution in these bulbs by the escaping oxygen, the gas is passed through H, contain- ing fragments of solid potas- sium hydroxide ; this tube is weighed along with the potash bulbs. K is a guard tube, containing fragments of solid potassium hydroxide, in order to prevent atmospheric carbon dioxide, and moisture, from gaining access to the weighed portions of the apparatus. The weight of the diamond, minus the weight of the ash which was left, gave the actual weight of the carbon burnt ; the increase in weight of the tubes gave the weight of the carbon dioxide which was pro- duced, and this weight, minus the weight of carbon used, gave the weight of oxygen that was consumed. As a mean of 9 Fig. 65. 274 Inorganic Chemistry a number of experiments, Dumas and Stas found that 80 parts of oxygen by weight, combined with 29.99 parts of carbon. From a knowledge of the density of carbon dioxide, and the volume of oxygen it contains, we know that the molecule of this gas contains two atoms ; therefore, by the simple equation — 80 : 31.92 : : 29.99 : H-97, 11.97 parts of carbon combine with 31.92 parts of oxygen, and the number 1 1.97 is therefore the atomic weight of carbon as determined by these chemists. The Carbonates. — Although carbonic acid, H 2 C 0 3 , is a very unstable compound, the salts it forms are stable. Being a dibasic acid, it is capable of forming salts in which either one or both of the hydrogen atoms have been replaced by an equivalent of a metal ; thus in the case of sodium we have — (1) Disodium carbonate (normal sodium carbonate) . NajCC^. (2) Hydrogen sodium carbonate (bicarbonate of soda) . HNaC 0 3 . Similarly, with the divalent metal calcium, it is possible to form— (1) Normal calcium carbonate CaC 0 3 , and — (2) Hydrogen calcium carbonate (bicarbonate of lime) . CaH 2 (C 0 3 )> »> !-26 ,, carbon dioxide. »» >* ». 2.27 ,, air. ** >» »i 2.30 ,, nitrogen. •• »» >» 5 * 11 m carbon monoxide. That the atmospheric oxygen effects the result by a direct oxidising action, and is not acting merely as nitrogen does, is proved by the fact, that mixtures of oxygen and nitrogen containing a higher proportion of oxygen than is Fig. 84. See “ Chemical Lecture Experiments,” new ed., 498-502. f Ibid., 506. The Bunsen Flame 305 present in air, destroy the luminosity more rapidly than is effected by air. Thus, when mixtures containing nitrogen and oxygen in the proportion of 3 to 1, 2 to 1, 1 to 1 by volume are employed, the volumes of the mixtures required to destroy the luminosity of one volume of coal gas, are respectively 2.02, 1.49, and 1. 00. It has been shown that when coal gas is diluted with nitrogen, a higher temperature is necessary to effect its decomposition ; hence the action of the atmospheric nitrogen in causing the loss of luminosity of a gas flame, is in part due to the higher temperature that is required for the formation of acetylene, which, as already mentioned, is the first step in the decomposition and con- densation of the hydrocarbons in the gas. As already mentioned, the luminosity of a flame is very much influenced by alterations of temperature ; and just as the non-luminosity of the outer mantle of an ordinary flame, is partly due to the cooling action of the air which is dragged into the flame from the outside, so the want of luminosity of the Bunsen flame, is in part due to the cooling influence of the large volume of air that is drawn up into the interior of the flame. That the gases which are drawn into a flame reduce the luminosity by virtue of their cooling action, is borne out by the fact, that the higher the specific heat of the diluent (and therefore the greater its power to abstract heat from the flame) the less of it is required to effect the destruction of the luminosity ; thus, as already men- tioned, less carbon dioxide than nitrogen is necessary to render a flame non- luminous : the specific heat of nitrogen is 0.2370, while that of carbon dioxide is 0.3307. The specific heat of oxygen is also slightly greater than that of nitrogen, being 0.2405 ; but the cooling effect of dilution with this gas, is enormously overpowered by the increased temperature due to its oxidising action upon the combustible materials of the flame. Experiments made upon the actual temperatures of various regions of a Bunsen flame, rendered non-luminous by admixture with different gases, the results of whifch are seen in the following table (Lewes), show the cooling effect of these diluents upon the flame : — Temperature of Flame from Bunsen Burner , burning 6 cubic feet of Coal Gas per Hour. Region in Flame. Luminous Flame. Flame rendered Non- luminous by Air. Nitrogen. Carbon Dioxide. \ inch above burner .... 1^ inch above burner .... Tip of inner cone .... Centre of outer cone .... Tip of outer cone .... Side of outer cone, level with tip of ) inner cone . . . , ) Degrees. 135 421 9 i 3 1328 728 1236 Degrees. 54 175 1090 *533 ii 75 1333 Degrees. 30 III 444 ' 999 1151 1236 Degrees. 35 70 393 770 95 i 970 I U 30 6 Inorganic Chemistry In the case of air, it will be seen that the first effect is to cool the flame • but in the upper region, where the oxidising action of the oxygen is felt the h^tinof^h raP1<3 y n T t0 a maximum at a Point about half way between bv thT , 'T r r CT C ° neS - In thC flames rendered non luminous y the two inert gases, the highest temperature is only reached at the outer oX'a'tXhe'e am ° Unt ° f ** Combustion is Gained from the On account of the wide range of temperature exhibited by the various f h Un$e " flame * h constitutes a most valuable analytical instru- ment for by the judicious use of the different parts of the flame, it is often possible to detect the presence of several flame-colouring substances in a mixture. Thus, if a mixture of sodium and potassium salts be introduced upon platinum wire into the cooler region of the flame near its base, the more volatile potassium compound will impart its characteristic violet tint to the flame, before the sodium salt is volatilised sufficiently to mask the colour by gives to the flame - In this way many “ If a piece of copper wire be held horizontally across a Bunsen flame, so as to cut the inner cone, it will be seen that the wire in contact with the edges of the flame becomes coated with copper oxide, while the portion in the centre remains bright. On moving the wire so as to bring the oxidised portion into the inner region, the oxide will be reduced, the metal once more becoming bri .^ t ’. 7 . he outer ar ea of a flame, where oxygen is in excess, is called the oxidising flame ; while the inner region, in which heated and unburnt hydro- gen or hydrocarbons exist, is spoken of as the reducing flame. These regions exist m ail ordinary flames. The oxidising action of the outer flame of a candle, for example, is illustrated in the behaviour of the wick. So Ion? as the wick remains in the inner region of the flame it is not burnt: and in the early days of candles, as the tallow gradually consumed, the wick remained standing straight up, and by degrees extended into the luminous area of the flame where, owing to the deposition of soot upon it, it frequently developed a cauliflower-like accretion, which greatly impaired the luminosity of the flame, and which necessitated the use of snuffers. In the modem candle owing to a method of plaiting the wick, it is caused to bend over (as shown m Fig. 79), and so thrusts its point into the oxidising region, where it is continually burnt away. PART III THE SYSTEMATIC STUDY OF THE ELEMENTS, BASED UPON THE PERIODIC CLASSIFICA- TION. CHAPTER I THE ELEMENTS OF GROUP VII. (. FAMILY B.) Fluorine, F . . 19.06 I Bromine, Br . . . 79.76 Chlorine, Cl . . 35-37 I Iodine, I 126.54 The first to be discovered, and the most important element of the group, is chlorine, which is a constituent of sea salt (sodium chloride). The term halogen , signifying sea salt producer, has been applied to this family of elements, on account of the close resemblance between their sodium salts and sea salt. This family exhibits, in a marked manner, many of the features which are found to exist in most chemical families of elements. In their general behaviour they strongly resemble one another, and readily displace each other in combinations without producing any very marked change upon the character of the compounds. They each unite with hydrogen, giving rise respectively to hydro- fluoric acid, HF ; hydrochloric acid, HC1 ; hydrobromic acid, HBr; hydriodic acid, HI. These hydrogen compounds are all colourless gases, which fume strongly in the air ; they are extremely soluble in water, and are strongly acid in character. In combination with potassium and with sodium, the halogens form a series of compounds, which are similarly constituted, and which closely resemble each other in their habits. Their similarity of composition is expressed in the follow- ing formulae — Compounds with potassium, KF, KC1, KBr, KI. Compounds with sodium, NaF, NaCl, NaBr, Nal. 307 308 Inorganic Chemistry The physical properties of the elements, exhibit a regular grada- tion with increasing atomic weight ; thus, fluorine and chlorine are gases, bromine is liquid, while iodine is solid, at ordinary tempera- tures. In their chemical activity they also show the same gradual change ; thus, in the case of their combination with hydrogen, when fluorine and hydrogen are brought together, combination instantly takes place with explosion, even in the dark. Chlorine and hydrogen do not combine in the dark, but in diffused daylight they unite slowly, and in direct sunlight their combination takes place suddenly with explosion. Bromine vapour and hydrogen do not combine even in direct sunlight, but a mixture of the two gases ignites in contact with a flame, yielding hydrobromic acid, while iodine vapour and hydro- gen require to be strongly heated in contact with spongy platinum to effect their combination. This difference in the activity of the halogens towards hydrogen, is seen by a comparison of the heats of formation of their hydrogen compounds, thus — H + Cl = HC1 + 22,000 cal. (Thomsen). H + Br = HBr + 8,440 „ * H + I = HI — 6,040 „ t The heat of formation of hydrofluoric acid has not yet been determined, but there can be no doubt that it is considerably greater than that of hydrochloric acid. Although a strong resemblance exists between all the members of the halogen family, the element fluorine, which is the typical member (see p. 100), stands marked off from the others in many of its attributes. Thus fluorine exhibits a great tendency to form double salts which have no counterpart among the compounds of the other elements of the family, and at temperatures below 32 °, the molecule of hydrofluoric acid consists of two atoms of hydrogen and two of fluorine, having the composition HgFjj. FLUORINE. Symbol, F. Atomic weight = 19.06. History. — This element, the first of the halogen series, was the most recent to be discovered, it having baffled all attempts to * This value refers to bromine in the liquid state, f Iodine as solid. Fluorine 309 isolate it until the year 1886, when Moissan succeeded in solving the problem. Occurrence. — Fluorine occurs in considerable quantities in com- bination with calcium in the mineral fiuor spar (CaF 2 ), which is found in cubical crystals. On account of the occurrence of this mineral in large quantities in Derbyshire it is frequently termed Derbyshire spar. It is a constituent also of cryolite, Na 3 AlF 6 , fluor- apatite, , 3 P 2 0 8 Ca 3 ,CaF 2 , and many others. In small quantities fluorine is found in bones, in the enamel of teeth, and also in certain mineral waters. Mode of Formation.— When an electric current is passed into an aqueous solution of hydrochloric acid, the acid is decomposed into its elements, chlorine being liberated at the positive electrode, while hydrogen is evolved at the negative. When aqueous hydro- fluoric acid is treated in the same way, the water only is decom- posed, oxygen and hydrogen being liberated. Davy found that the more nearly the acid approached the anhydrous condition, the less easily did it conduct electricity ; and that in the perfectly pure state, that is, entirely free from water, hydrofluoric acid was a non- conductor. Moissan’s recent success in the isolation of fluorine, depends upon the discovery that a solution of the acid potassium fluoride, HF,KF, in anhydrous hydrofluoric acid is an electrolyte, and that by the passage of an electric current through this solution, fluorine is disengaged at the anode, or positive electrode, and hydrogen is evolved at the cathode. The primary decomposition taking place is the breaking up of the acid potassium fluoride — HF,KF = F 2 + H + K. The atom of potassium, in contact with the free hydrofluoric acid present, is then converted into potassium fluoride with the elimina- tion of an equivalent of hydrogen — K + HF = KF + H. And the normal potassium fluoride then unites with a molecule of the acid, to regenerate the acid salt — KF + HF = HF,KF. The reaction is performed in a U-tube made of an alloy of platinum and iridium, a material which is less acted upon by the 3 TO Inorganic Chemistry liberated fluorine than platinum alone. The apparatus has two side-tubes (Fig. 85), which can be either closed with a screw cap, c, or connected to platinum delivery tubes by means of the union D. The two limbs of the tube are closed by means of stoppers made of fluor spar, shown in section at s, and which can be securely screwed into the tube. These serve to insulate the electrodes, which are constructed of the same platinum-iridium alloy. The anhydrous hydrofluoric acid is introduced into the apparatus, and about 25 per cent, of its weight of the acid potassium fluoride is added, which really dissolves in the liquid. The tube is immersed in a bath of methyl chloride (m, Fig. 86), which boils at -23; the supply being continuously replenished from the reservoir B, while the vapour is drawn away by the pipe c. On passing a current from 20 to 25 Grove’s cells through the apparatus, fluorine is evolved at the positive electrode, and hydrogen is liberated at the negative.* Properties.— Fluorine is, of all known elements, the most chemically active. It is on account of its intense chemical affinities * Fluorine has recently been obtained by Brauner (June 1894) by heating potassium fluorplumbate, 3KF,HF,PbF 4 . At 200° this salt gives off hydro- fluoric acid, HF, and when heated to 23o°-25o° fluorine is evolved. Fluorine 3 ” that it so long resisted all attempts to isolate it, as when libe- rated from combination it instantly combined with the materials of the vessels in which the reactions were made. It is impossible to collect this gas by any of the usual methods, for it decomposes water, and instantly combines with mercury. It also attacks glass, so that it can only be collected by displacement of air, in vessels of platinum. Fluorine appears to be a colourless gas, judged by the small volume of it that can be viewed at once. The smell of the gas is very characteristic, it is irritating to the mucous membranes, and is not unlike the odour of the mixture of chlorine and chlorine Fig. 86. peroxide, evolved from potassium chlorate and hydrochloric acid. Whether the smell actually perceived is the true smell of fluorine is doubtful, for when fluorine comes into contact with the moisture in the nostrils, water is decomposed, with the formation of ozonised oxygen and hydrofluoric acid. Fluorine not only decomposes potassium iodide, with liberation of iodine, but also displaces chlorine from sodium chloride. It combines directly with a large number of elements with in- tense energy ; in contact with hydrogen it instantly explodes. 312 Inorganic Chemistry Iodine, sulphur, and phosphorus first melt, and then take fire in fluorine. Crystals of silicon when brought into the gas snontane unon b nfl fl ame ’ and bUm WUh brillianc >'' AI1 of ‘he mftal’s are acted upon by fluorine ; some, when finely divided, undergoing spontane- ous inflammation when thrown into the gas. Eve/gold and plati- num are attacked by fluorine, especially if gently warned : itsacion opon the latter metal being seen by the corrosion of the apparSus an Or 3 c ally lhe P T' tIVe e,eCtrode em P'oyed in its preparation. ’ ofte„Tflame d mP ° U ^ flUOrine Whh “ d When fluorine is cooled to a temperature about - 18c (i e a fpw bofling S th b e e o W tHe 'TTr ° f b ° iling 0xygen > obtain ’ ed »r- y f ” Und t e f Sl,ghtly reduced Pressure) it condenses the liquid state* Liquid fluorine is a mobile yellow liquid resembling liquid chlorine. It is without action upon siKcon, phosphorus, sulphur, or glass ; it can therefore be produced and contained in glass vessels. Even at this low temperature, how- ever, fluorine attacks hydrogen and hydrocarbons. HYDROFLUORIC ACID. Formula, H 2 F 2 and HF. Molecular weight, 40 and 20. Density, 20 and 10. Modes ° f Formation.— (1.) Hydrofluoric acid is produced when s;s« c " a ™ <*-" •-'> " »«■ »«» . CaF 2 + H 2 S 0 4 = CaS 0 4 + 2 HF. This method is employed for the commercial preparation of aqueous solutions of hydrofluoric acid. The mixture of fluor spar and sulphuric acid is gently warmed in a leaden retort, and the gaseous acid passed into water contained in leaden bottles. This aqueous acid is sent into the market in gutta-percha bottles. (2.) The anhydrous acid is prepared by heating hydrogen potas- sium fluoride (aad potassium fluoride) in a platinum retort. The ouble fluoride of potassium and hydrogen splits up into norma! potassium fluoride, and hydrofluoric acid HF,KF = KF + HF. For this purpose the perfectly dry double fluoride is placed in a platinum retort, which is screwed to a platinum condensing arrangement, as seen in Fig. 87. The wooden trough through which the long tube passes is filled with a freezing mixture, and the platinum bottle is also surrounded by a similar mixture. * Moissan, May 1897. Hydrofluoric A cid 3 1 3 Properties. — Anhydrous hydrofluoric acid is a colourless, limpid, strongly fuming liquid, which boils at 19.5 0 . It has a powerful affinity for water, and can only be preserved in perfectly stoppered platinum vessels, which are kept in a cool place. The acid at once attacks gutta-percha. Gore found that the anhydrous acid was without action upon glass. Pure hydrofluoric acid is an extremely dangerous substance to manipulate ; its vapour, even when diluted with air, has a most irritating and injurious effect upon the respiratory organs, and if inhaled in the pure state causes death. A single drop of the liquid upon the skin causes the most painful ulcerated sores, accompanied by distressing aching pains through- out the whole body. The metals potassium and sodium dissolve in pure hydrofluoric acid, with the formation of fluorides and evolution of hydrogen. The vapour density of hydrofluoric acid at a temperature of 32 0 , corresponds to the. formula H 2 F 2 . As the temperature is raised, the density of the vapour falls steadily, until at 88° it is 10, a density which is normal for the formula HF. The composition of this com- pound, therefore, as ordinarily seen, is probably represented by the formula H 2 F 2 , the molecules of which undergo dissociation when the temperature rises, until at 88° the substance consists entirely of the simpler molecules HF. Gaseous hydrofluoric acid rapidly attacks glass, and it is largely employed for etching purposes, both for obtaining designs upon 314 Inorganic Chemistry glass, and for the purpose of etching graduations upon glass mea- suring instruments. The object to be etched is first coated with wax, and the design or other marks cut upon the wax by means of a pointed steel tool. In this way the surface of the glass is laid bare in parts, and on exposing the object to the action of the acid, either as gas or aqueous solution, the glass is rapidly eaten into, where the surface has been exposed. Its action upon glass is due to the readiness with which it attacks silicates, the fluorine combining with the silicon to form silicon tetrafluoride — Si0 2 + 4HF = 2H 2 0 + SiF 4 . Crystallised silicon, when gently heated, takes fire in gaseous hydrofluoric acid, giving silicon fluoride and hydrogen. Hydrofluoric acid is extremely soluble in water, forming a strongly acid corrosive liquid, which readily dissolves many of the metals with evolution of hydrogen — Fe + 2HF = FeF 2 + Hj. Silver and copper are also dissolved by this acid. CHLORINE. Symbol, Cl. Atomic weight, 35.37. Molecular weight, 70.74. History. — Chlorine was discovered by Scheele (1774), but was regarded by him as a compound substance. He applied to it the name of dephlogisticated marine acid air , having obtained it by the action of hydrochloric acid upon ores of manganese. The belief that chlorine was a compound of oxygen and hydrochloric acid, was generally held until Davy’s time, and gave rise to the name of oxymuriatic acid. The elementary nature of chlorine was proved by Davy (1810), who gave to it the name chlorine , in allusion to the greenish-yellow colour of the gas. Occurrence.— In the uncombined condition chlorine does not occur in nature. In combination with metals, as chlorides, chlorine is very abundant, the commonest chloride being sodium chloride (common salt). Many of the salts found in the Stassfurt deposits, consist largely of chlorides (see Alkali Metals). Chlorides of the alkali metals are also found in animal secretions and in certain plants. Chlorine occurs in combination with hydrogen, as hydrochloric acid, in volcanic gases, and also in the gastric juice. Modes of Formation. — (1.) When hydrochloric acid is poured upon manganese dioxide, and the mixture kept cool, a dark-brown Chlorine 3 1 5 solution is obtained, which rapidly decomposes at a slight rise of temperature with the evolution of chlorine. It has not yet been clearly established whether this brown solu- tion consists of the compound MnCl 4 or Mn 2 Cl 6 , formed according to one of the equations — Mn0 2 + 4HC1 = MnCl 4 + 2H 2 0. or — 2MnO a + 8HC1 = Mn 2 Cl 6 + Cl 2 + 4H 2 0. When this dark-brown solution is gently warmed, the higher chloride breaks up into manganous chloride (MnCl 2 ) and chlorine ; the complete reaction being expressed by the equation — MnO a + 4HC1 = 2H 2 0 + MnCl 2 + Cl 2 . The experiment is conveniently carried out in the apparatus seen in Fig. 88. The mixture of manganese dioxide and hydrochloric acid is gently heated in a large flask, and the gas, after being passed through water in the Woulf’s bottle, may be collected by downward displacement, as shown in the figure.* ( 2 .) Instead of employing hydrochloric acid, the materials from which this compound is prepared, namely, sodium chloride and sulphuric acid, may be used. Thus, if a mixture of sodium See Experiment 154, “ Chemical Lecture Experiments,” new ed. 3 J 6 Inorganic Chemistry chloride, manganese dioxide, and sulphuric acid be gently warmed chlorine is readily evolved — 2NaCl + MnO a + 2H 2 S0 4 = Na 2 S0 4 + MnS0 4 + 2H a O + Cl 2 . It will be seen that by this reaction the whole of the chlorine con- tained in the reacting compounds is evolved as gas, while in the former case a part of it remains in combination with the man- ganese. (3.) Many other highly oxygenised compounds, when acted upon by hydrochloric acid, evolve chlorine ; thus, when crystals of potas- sium bichromate are drenched with hydrochloric acid, and the mixture heated, a rapid stream of chlorine takes place, thus K 2 Cr 2 0 7 + 14HC1 = 2KC1 + Cr 2 Cl 6 + 7H 2 0 + 3C1 2 . (4.) When crystals of potassium chlorate are similarly treated, a mixture of chlorine and chlorine peroxide is evolved, even without the application of heat — 4KC10 3 -1- 12HC1 = 4KC1 + 6H 2 0 + 3C10 2 + 9C1. (5.) Red lead (Pb 3 0 4 ), when treated with hydrochloric acid, reacts in a manner similar to manganese dioxide and many other peroxides. In the case of lead, however, there is no intermediate chloride formed — Pb 3 0 4 + 8HC1 = 3PbCl 2 + 4H 2 0 + Cl* (6.) Manufacturing Processes — Deacon’s Process. —This method for the preparation of chlorine, is by the oxidation of the hydrogen in hydrochloric acid by atmospheric oxygen. It will be seen that in the foregoing methods the oxidation of this hydrogen is carried on at the expense of the oxygen contained in either the metallic peroxide, or the highly oxygenated salt used ; in the Deacon process atmospheric oxygen is made use of. When a mixture of gaseous hydrochloric acid and oxygen is heated, a slight decomposition takes place ; but if these gases be heated in the presence of a third substance, which acts as a catalytic agent, the decomposition of the hydrochloric acid is much more readily effected. The catalytic agent employed in the Deacon process is cuprous chloride (Cu 2 Cl 2 ). This substance is capable of taking up Chlorine 3 1 7 an additional quantity of chlorine, and of being converted into cupric chloride (CuCl 2 ), thus — Cu 2 Cl 2 + Cl 2 = 2 CuC1 2 . If, therefore, a mixture of hydrochloric acid and oxygen be passed over fragments of pumice impregnated with cuprous chloride, contained in a tube which is heated to dull redness, the hydrochloric acid will be decomposed. We may suppose that the affinity of oxygen for the hydrogen in hydrochloric acid, is un- able to overcome the affinity existing between the hydrogen and chlorine, but that the additional pull exerted upon the molecules of hydrochloric acid by the cuprous chloride, is sufficient to dis- turb the equilibrium and rupture the molecule — q H | Cl Cu Cl 0 H i Cl 2 2 ‘ The result of the action being H a O + 2CuCla. At the temperature at which the reaction is carried on, however, the compound CuCl 2 cannot exist ; two molecules of it are con- verted into one of Cu 2 Cl 2 , and a molecule of chlorine is evolved. The final result, therefore, of the reaction may be thus expressed — O + 2HC1 + Cu 2 Cl 2 = H 2 0 4 - Cl 2 + Cu 2 Cl 2 . In reality the action is rather more complex, there being an intermediate compound formed by the combination of cuprous chloride with oxygen. This 3 1 8 Inorganic Chemistry oxychloride of copper then acts upon the hydrochloric acid as seen in the following equations : — (1) Cu 2 Cl 2 + O = CujjOCl* ( 2 ) CuaOClj + 2HC1 = 2CuCl2 + H 2 0. (3) ^uClj = CujjC^ -f- Clj. This reaction may be made on a small scale by means of the apparatus shown in Fig. 89. Hydrochloric acid is generated from salt and sulphuric acid in the flask, and a stream of the gas passed through the WoulPs bottle, into which also enters a stream of oxygen. The mixed gases are then passed through the bulb-tube, containing fragments of pumice which have previously been soaked Fig. 90. in a solution of cupric chloride, and dried. On heating the bulb by means of a Bunsen flame, chlorine will issue from the end of the tube. When chlorine is manufactured on an industrial scale by the Deacon process, the mixture of hydrochloric acid and air (in the proportion of four volumes of the latter to one volume of hydro- chloric acid) is drawn by means of a Root’s blower first through iron pipes, which are heated to a temperature of about 500°, and then the hot gases pass on through the decomposer. This consists of a cylinder of cast iron, containing masses of broken brick, or burnt clay, impregnated with cupric chloride, and so arranged that the gases are drawn through the mass. Chlorine 319 The gas leaving the decomposer, consists of a mixture of chlorine, undecomposed hydrochloric acid, and atmospheric nitrogen and oxygen. By passing them through water, the hydrochloric acid is removed, and the chlorine is usually converted at once into bleach- ing-powder. The process by which chlorine is usually made on a manufactur- ing scale, is by the action of hydrochloric acid upon manganese dioxide. The best ore for the purpose is pyrolusite. The process is conducted in stills made of thick slabs of stone, usually “York- shire flag,” which are fitted and luted together, and securely bound by cast-iron clamps. Fig. 90 shows such a chlorine still, repre- sented as cut across the centre. The charge of manganese is placed upon the false bottom a , and the acid is run in through the funnel tube b , which, dipping into a small pot, does not allow the gas to escape. As the action begins to slacken, steam is cautiously blown in from time to time. The chlorine escapes by the pipe g , and passes from thence into the main h. The reaction that goes on in the still, is the same as that given above in the first mode of formation, except that as pyrolusite is not pure Mn 0 2 , small quantities of other compounds are formed. The following analysis, by Black, of still-liquor from stone stills, shows the general composition of this substance : — (7.) The Weldon Process, although indirectly a method for making chlorine, is in reality a process for recovering the man- ganese contained in the still-liquors as manganous chloride, and of reconverting it into available manganese dioxide. The manganese so recovered, however, is again utilised for the preparation of chlorine by the decomposition of a further quantity of hydrochloric acid. The essence of the process is the following The still- liquor is mixed with ground chalk, or limestone dust, in large tanks or wells, and the mixture thoroughly stirred by agitators. One of these wells, A, is shown in the diagrammatic figure. By this opera- MnCl 2 . 10.5700 0.6200 0.4551 6.6220 81.7329 HC1 (undecomposed) h 2 o . 100.0000 320 Inorganic Chemistry tion the free acid is neutralised, and the iron precipitated as hydrated oxide. The neutral liquor, consisting of manganous chloride and calcium chloride, is then pumped into large tanks, where it is allowed to settle ; one of these “ settlers,” B, is shown in the figure. By means of a pipe upon a swivel-joint,/, the clear liquid from the settler can be drawn off without disturbing the sediment, and run into the oxidiser C. The oxidiser is merely a flat-bottomed iron cylinder, open at the top. Milk of lime from the tank E, where lime and water are stirred together, is pumped into the oxidiser as required. The milk of lime is added in quantity more than sufficient to precipitate the manganese as manganous hydroxide, MnHgOg. Into this mixture, which consists of manganous hydroxide and calcium hydroxide (milk of lime) in suspension, and to a smaller extent in solution in the calcium chloride which is present, a stream of compressed air is forced by means of the pipe h, which Chlorine 321 passes to the bottom of the oxidiser, where it ends in perforated branches. During this process the manganese becomes oxidised, and is converted mainly into calcium manganite, a compound of manganese dioxide with calcium oxide, CaO,Mn 0 2 , or CaMnO s . By a further addition of the neutral liquor from tank B, and by raising the temperature within the oxidiser by injecting steam, a portion of the calcium manganite is converted into a compound having the composition Ca 0 , 2 Mn 0 2 . When the operation is complete, the contents of the oxidiser are run out into a series of tanks called mud settlers , of which one is shown at D in the figure. The product here settles as a thin black mud, known as the Weldon mudj and this is ultimately drawn from the settlers, and run direct into chlorine stills, where it is at once treated with hydrochloric acid for the preparation of chlorine. The Weldon stills are similar to the ordinary chlorine stills, but are much larger, and usually octagonal in shape. Properties. — Chlorine is a greenish-yellow coloured gas, with a strong suffocating smell. It is quite irrespirable, and if inhaled in the pure state causes death. Even when largely diluted with air it is extremely disagreeable and injurious, as it acts rapidly upon the mucous membranes of the nose and throat, causing irritation and inflammation, which usually result in severe catarrh. A few bubbles of chlorine allowed to escape and diffuse into thd air of a room, give to the air a distinct and rather pleasant smell. Chlorine is an extremely heavy gas, being 35^37 _ 2.45 times heavier than 1 4*44 air. One litre of the gas, measured under the standard conditions of temperature and pressure, weighs 3.168 grammes. The density of chlorine, taken at all temperatures, does not exactly agree with that which is required for the molecular formula Cl 2 . At tempe- ratures above 1200° the density is markedly less than theory demands, showing that partial dissociation of the chlorine mole- cules into single atoms has taken place. (Compare Bromine and Iodine.) On account of its heaviness, chlorine is readily collected by dis- placement ; it cannot be collected over mercury, as it attacks that metal, and in water it is considerably soluble. It may, however, be collected over a strong brine, as it is much less soluble in this solution than in water. Chlorine is not inflammable, but it supports the combustion of many burning bodies. It is possessed of such extremely powerful x 322 / norganic Ckem istiy chemical affinities that it acts upon a large number of substances at ordinary temperatures, and in many cases the combination is sufficiently energetic to result in the inflammation of the bodies. Phosphorus, when introduced into chlorine, first melts and then spontaneously inflames, burning with a somewhat feeble light to form phosphorus trichloride (PC1 3 ) and phosphorus pentachloride (PC1 6 ). The elements arsenic and antimony, when finely powdered and dusted into a vessel of chlorine, at once take fire and burn, forming their respective chlorides. Many metals, when finely divided, or in the form of thin leaf, such as ordinary Dutch metal, instantly take fire when brought into chlorine. If a quantity of sodium be heated in a deflagrating spoon until it begins to burn in the air, and be then plunged into chlorine, the sodium continues to burn in the gas with dazzling brilliancy, forming sodium chloride. Although under ordinary circumstances chlorine unites with metals with great readiness, it has been shown that this action will not take place if the chlorine be absolutely dry. Thus, if chlorine which has been completely freed from aqueous vapour be passed into a tube containing bright metallic sodium, and the tube sealed, the sodium not only remains bright and unaffected by the gas, but may even be melted in the atmosphere of chlorine without any action taking place. Similarly, dry chlorine, when allowed to enter a flask filled with Dutch metal, has no action upon it ; but upon the introduction of the smallest trace of moisture, the metal at once takes fire.* These facts are of the same order as those mentioned in connection with oxygen, and have not yet received any satis- factory explanation. Chlorine is not capable of direct combination with carbon ; ordi- nary combustibles, therefore, which consist of hydrocarbons, burn in chlorine by virtue of the combination of their hydrogen with the gas, and they burn with a lurid smoky flame, owing to the elimina- nation of their carbon in the form of soot. A burning taper or ordinary gas flame, when introduced into chlorine, burns in this manner, emitting a dense smoke and forming fumes of hydro- chloric acid. Chlorine has a most powerful affinity for hydrogen ; a jet of hydrogen bums freely in chlorine, with the formation of hydro- chloric acid. A mixture of hydrogen and chlorine unites with explosion on the application of a flame. This combination takes place also under the influence of light (see Hydrochloric Acid). * See Experiments 174, 175, “ Chemical Lecture Experiments,” new cd. Chlorine 323 The affinity shown by chlorine for hydrogen is seen in its action upon many of the compounds of hydrogen and carbon. If one volume of ethylene (C 2 H 4 ) be mixed with two volumes of chlorine, and the mixture ignited, the carbon is instantly thrown out of com- bination as a black smoke, while the hydrogen unites with the chlorine, forming a cloud of hydrochloric acid. Similarly, if a liquid hydrocarbon, such as turpentine (C 10 H 16 ), be poured upon a piece of filter paper and the paper be thrust into a jar of chlorine, instant inflammation takes place, with deposition of a large quantity of carbon. Chlorine possesses strong bleaching properties, which depend upon its power of combining with hydrogen, for it is an essential condition that water shall be present. ' The chlorine unites with the hydrogen of the water, and the liberated oxygen oxidises the colouring matter. If chlorine be bubbled into liquids coloured with any vegetable colouring matter, or if a dyed rag be dipped into chlorine water, the colour will be rapidly discharged. Ordinary writing-ink (which usually consists of a compound of iron with tannic and gallic acids) is readily bleached by chlorine ; while printer’s ink, which consists mainly of carbon, in the form of lamp- black, is not acted upon by this gas. If, therefore, a piece of printed paper be brushed over with writing-ink so as to completely obli- terate the print, and the blackened paper be immersed in chlorine water, the writing-ink will be rapidly bleached away, leaving the print unchanged. The bleaching power of chlorine constitutes its most valuable property from an industrial point of view ; the chlorine for this purpose is combined with lime to form the substance known as / bleaching-powder. (See Calcium Compounds.) Chlorine is soluble in water to a considerable extent. One volume of water at io° absorbs 3.0361 volumes of chlorine measured at o° and under 760 mm. pressure. This solution, known as chlorine water, has the same colour as the gas, and smells strongly of chlorine. If exposed to the air, the chlorine rapidly diffuses out of the solution. Chlorine water cannot be preserved for any length of time, as it slowly undergoes de- composition, the chlorine combining with the hydrogen of the water, forming hydrochloric acid, which remains in solution, and the oxygen being liberated, thus — H 2 0 + Cl 2 = 2 HC 1 + O. 324 Inorganic Chemistry This action, which proceeds slowly under ordinary conditions, is greatly accelerated by the influence of light, and if exposed to direct sunlight the decomposition is very rapid. If chlorine water be cooled to within one or two degrees of the freezing-point of water, or if chlorine be passed into ice-cold water, a solid crystalline compound of chlorine with water is deposited. This substance is termed chlorine hydrate , and has a composition expressed by the formula Cl 2 ,10H 2 O. The compound is very un- stable, and when exposed to the air it melts and rapidly gives off chlorine. If the crystals are quickly freed from adhering water, and are then sealed up in a glass tube, they may be heated to a temperature of 38° before being decomposed. Faraday made use of this compound in order to obtain liquefied chlorine. A quantity of the hydrate was sealed up in one limb of a bent tube and was gently warmed, the compound dissociated into water and chlorine, and the internal pressure caused the condensa- tion of the chlorine to the liquid condition. Liquid Chlorine. — Under the ordinary atmospheric pressure, chlorine may be liquefied by lowering its temperature to — 34 0 . At a temperature of o° the pressure required to effect its lique- faction is equal to six atmospheres. When, therefore, the liquid is obtained by heating the crystalline hydrate, as in Faraday’s method, one limb of the tube should be cooled by being placed in ice. The critical temperature of chlorine is 14 1°, and the pressure required to effect its liquefaction at that point, or its critical pressure, is 84 atmospheres. (See Liquefaction of Gases.) Liquid chlorine has a bright golden yellow colour, entirely free from the greenish tint possessed by the gas. Its specific gravity is 1.33, and it boils at — 33-6°- When cooled to a temperature of - 102, it freezes to a yellow crystalline mass. Liquid chlorine is now an article of commerce. It is contained in iron bottles lined with lead, and is largely exported in this form, for use in the extraction of gold, to parts of the world where the carriage of the plant and materials necessary for generating large quantities of chlorine, would be attended with great difficulties. Hydrochloric Acid 325 HYDROCHLORIC ACID ( Hydrogen Chloride ). Formula, HC 1 . Molecular weight = 36.37. Density = 18.185. History. — In solution in water, this compound was known to the early alchemists, and the mixture of this solution with nitric acid constituted the valued liquid known as aqua regia. The preparation of hydrochloric acid from common salt is associated with the name of Glauber (1650), who obtained it by the action of sulphuric acid upon sodium chloride (common salt). Gaseous hydrochloric acid was first collected and examined by Priestley, who collected it over mercury, in the mercurial pneumatic trough invented by him. He named the gas marine acid air. Occurrence. — Gaseous hydrochloric acid is evolved in consider- able quantities from volcanoes during active eruption. Modes of Formation. — ( 1.) Hydrochloric acid may be syn- thetically produced directly from its elements ; thus, this compound is formed when a jet of hydrogen is caused to burn in an atmos- phere of chlorine. If a mixture of chlorine and hydrogen be ignited, the union takes place instantaneously with explosion, and hydrochloric acid is produced. The union of hydrogen with chlorine will also take place under the influence of light ; thus, if a / 326 Inorganic Chemistry mixture of these two gases be exposed to even diffused daylight for a few hours, the greenish colour imparted to the mixture by the chlorine will gradually disappear, and on examination it is found that the tube contains hydrochloric acid. This combination, which is only gradual when the mixture is exposed to diffused daylight, becomes explosively sudden if the mixed gases are exposed to direct sunlight, or any artificial light which is rich in rays of high refrangibihty the so-called actinic rays. If, therefore, a glass vessel be filled with a mixture of these gases in equal volumes, and the mixture be placed in bright sunshine, a violent explosion will result, and hydrochloric acid will ■ be produced. This phenomenon is best illustrated by U filling small thin glass bulbs with a mixture of the two f gases obtained by the electrolysis of aqueous hydrochloric || ac ^* The bulbs when filled can be hermetically sealed _ before the blowpipe without causing the combination of the gases ,* and if kept in the dark may be preserved indefinitely. On exposing one of these bulbs to the light of burning magnesium, the combination of the two gases instantly takes place, with a sharp explosion, which shatters the bulb to powder. The bulb should therefore be screened, as shown in Fig. 92. p> The rays of light which are capable of causing this combination are those which compose the blue and violet end of the spectrum ; if these particular rays are absorbed V from the light by means of ruby glass, the mixture of gases may be exposed to the red light so obtained with- out any action taking place.t The combination of chlorine with hydrogen is not attended by any alteration in volume ; one volume of Fig. 93. chlorine combines with one volume of hydrogen, and the resultant hydrochloric acid occupies two volumes. This may be readily proved by filling a stout glass tube, provided with a stop-cock at each end, with a mixture of the two gases in exactly equal volumes, and causing them to combine either by the influence of light, or by the passage of an electric spark by means of the platinum wires sealed into the tube (Fig. 93). On opening one of the stop-cocks under mercury, it will be seen that * See Experiment 178, “ Chemical Lecture Experiments,” new ed. t Ibid., p. 93. 327 Hydrochloric Acid no mercury is drawn in, neither does any gas pass out from the tube, thus showing that the union has taken place without any alteration in the volume. If one of the cocks be now opened beneath water, the hydrochloric acid which has resulted from the union of the hydrogen and chlorine, being extremely soluble m water, the liquid will rush up into the tube and completely fill it, showing that no free hydrogen or chlorine remains in the tube. (2.) For all ordinary purposes, hydrochloric acid is always obtained by the action of sulphuric acid upon sodium chloride. For labo- ratory uses the apparatus seen in Fig. 94 may be conveniently employed. Sulphuric acid, previously diluted with rather less than its own volume of water, is placed in the flask, and a quantity of common salt is added. On the application of a gentle heat a steady stream of gas is evolved, which may be dried by being passed through the tubulated bottle, containing pumice moistened with strong sulphuric acid. The gas is then collected either over mercury, or by displacement. The reaction which takes place is expressed by the equation — NaCl + H 2 S0 4 = NaHS0 4 + HC1. If strong sulphuric acid be employed along with an excess of salt, both of the atoms of hydrogen can be displaced from the acid; and instead of the hydrogen sodium sulphate, there is formed the normal sodium sulphate— 2NaCl + H 2 S0 4 = Na 2 S0 4 4- 2HC1. 328 Inorganic Chemistry A much higher temperature is necessary in order to complete the reaction indicated by this equation. Properties.— Hydrochloric acid is a colourless gas with a choking, pungent odour. In contact with the moist air it forms dense fumes, consisting of minute globules of a solution of the gas in the atmospheric aqueous vapour. Hydrochloric acid does not burn, neither does it support ordinary combustion. It is heavier than air, its specific gravity being— 18.185 14.44 1.25 (air = 1). Hence the gas is readily collected by displacement. One litre of the gas weighs 18.185 criths. Hydrochloric acid is extremely soluble in water ; 1 volume of water at o° and under a pressure of 760 mm. is capable of dis- solving 503 volumes of gaseous hydrochloric acid, measured at o and 760 mm. As the temperature rises the solubility diminishes as seen by the following table : — Temperature. Coefficient of Absorption. The solubility of hydrochloric acid may be illustrated by com- plMely fining a large globular flask with the gas, by displacement, the flask being provided with a long tube passing through the cork, as seen in Fig. 95. On opening this tube beneath water, the gas begins to dissolve, and the liquid rises slowly in the tube until it reaches the top. As soon as the first few drops enter the globe they rapidly absorb the gas, thereby causing a partial vacuum in the vessel, so that the water is driven up the tube with consider- able force, forming a fountain, which continues until the globe is nearly filled with liquid. If the water in the dish is rendered blue by the addition of litmus solution, the acid nature of the solution of the gas will be evident by the reddening of the liquid as it enters the globe. When a weak aqueous solution of hydrochloric acid is boiled it loses water and becomes stronger ; while, on the other hand’ if a strong solution be heated, it loses gas and becomes weaker’ until in both cases an acid containing 20.24 per cent, of HC1 is produced which boils at no 0 . This strength of add corresponds 329 Hydrochloric Acid to a composition expressed by the formula HC1 + 8H z O, and it was at one time supposed to represent a definite compound. Roscoe and Dittmar have shown, however, that, as with nitric acid, the composition of the liquid which boils at a constant temperature is simply a function of the pressure. The strongest aqueous solution of hydrochloric acid at 15 C. has a specific gravity of 1.212, and contains 42.9 per cent, of HC1. Hydrochloric acid gas is readily lique- fied by pressure. At a temperature of io° a pressure of 40 atmospheres will effect its liquefaction. If the temperature be lowered to -16 0 , the same result is ob- tained by a pressure of 20 atmospheres. The critical temperature of hydrochloric acid is 52.3 0 . Condensed hydrochloric acid is a colour- less liquid. Gore has shown that this liquefied acid is without action on most of the metals which are readily dissolved by the aqueous acid. The composition of hydrochloric acid may be experimentally proved by a num- ber of methods. It may be shown synthetically by the volumetric experiment referred to above (page 326). The volumetric proportion of hydrogen contained in the gas may be shown by means of sodium amalgam. The sodium in the amalgam, acts upon the hydrochloric acid, combining with the chlorine, and liberating the hydrogen — Na + HC1 = NaCl + H. For this purpose gaseous hydrochloric acid is introduced into one limb of the U-shaped eudiometer (Fig. 96), and its volume indicated by means of a ring upon the tube, the mercury being level in both limbs. A second ring marks exactly half the volume. A quantity of liquid sodium amalgam is then poured into the open limb until it is completely filled, and on being closed by the thumb the tube can be inverted so as to decant the gas into this limb. After being bubbled once or twice through the amalgam, the gas is again returned to its former place ; and by drawing mercury from the 33 ° Inorganic Chemistry branch tube, the levels in each limb can be again adjusted, when it will be found that the gas remaining in the tube occupies the space exactly down to the upper ring, that is to say, two volumes of hydrochloric acid contain one volume of hydrogen. That the gas is hydrogen can be shown by again filling up the open limb with mercury, and driving the gas out of the stop-cock, where it can be inflamed as it escapes. Fig. 96. Fig. 97. Fig. 98. The fact that hydrochloric acid contains the same volume 0/ ch lorine as of hydrogen, may also be demonstrated by collecting the mixed gases, evolved by the electrolysis of the aqueous acid, in a long tube provided with a stoppered funnel, as shown in Fig. 97. The gases may be collected over a saturated solution of salt in wat er, and the tube filled to the lower ring. On allowing a solution of potassium iodide to enter by means of the funnel, the chlorine is 33 T Hydrochloric Acid absorbed with the liberation of iodine, which partially dissolves and partly separates as a solid. When the absorption of the chlorine is complete, the water will have risen to the second band placed half way up the tube, showing that one-half of the gaseous mixture consists of chlorine. The former experiment proved that hydrochloric acid contained half its volume of hydrogen, therefore the two elements, in uniting to form this compound, do so in equal volumes and without any contraction in volume. When aqueous hydrochloric acid is subjected to electrolysis, the hydrochloric acid is decomposed, hydrogen being evolved at the negative electrode and chlorine at the positive. At first the liberated chlorine is dissolved in the solution ; but after the liquid has become saturated with the gas, the whole of the chlorine is liberated. By conducting this decomposition in the apparatus seen in Fig. 98 , and continuing the passage of the electric current until the liquid in one limb is saturated with chlorine before closing the stop-cocks, it will be seen, when the gases are collected in the tubes, that they are evolved in equal volumes. The Manufacture of Hydrochloric Acid. — The aqueous solution of hydrochloric acid is an object of commercial manu- facture, which is carried out on an enormous scale. It is obtained by the decomposition of common salt by means of sulphuric acid, according to the reaction — 2 NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. Formerly hydrochloric acid was a waste product obtained in the manufacture of sodium carbonate by the method known as the Leblanc process ; the first stage in this process being the conversion of sodium chloride into sodium sulphate, by the action upon it of sulphuric acid. The hydrochloric acid evolved as gas in this reaction was allowed to escape into the atmosphere. The nuisance caused by this acid gas being thrown into the air, ultimately resulted in the "Alkali Act,” which compelled manufacturers to absorb this waste acid. Since that time, the Leblanc process for the manufacture of sodium carbonate has had a formidable rival in another method, known as the ammonia-soda process (see Sodium Compounds), which would probably have completely driven the older method out of the field, but for the commercial value of the hydrochloric acid which is obtained as a secondary product in the Leblanc process. The hydrochloric acid, therefore, which formerly was thrown away as a waste product, is now the salvation of the process, and the utmost care is taken to prevent any of it from escaping, not now by com- pulsion of the Alkali Act, so much as from purely economic reasons. The charge of salt and sulphuric acid is heated in an enormous hemispherical cast-iron pan, built into a brickwork chamber, so 332 Inorganic Chemistry that it can be heated by a fire beneath, and so that the evolved gas can be conveyed away by brick or earthenware flues. The gas evolved by the reaction, is led into towers which are filled with coke or bricks, and down which water is made to percolate ; the water being caused to flow equally over the mass, by means of special distributing contrivances. As the gaseous hydrochloric acid passes up the towers and meets the descending stream of water it is entirely dissolved, and the aqueous acid becomes nearly saturated as it reaches the bottom of the tower. In works where the condensers, or towers, are not of great height, it is usual either to cool the gas before admitting it into the towers, or to pass it through a series of jars, resembling gigantic Woulfs bottles (Fig. 99). The water in these bottles is made to flow steadily from one to the other by the side pipes c, c (in the direction from left to right), while the gas passes through the system in the opposite direction! In this way a constantly changing surface of water is exposed to the gas, and a very strong solution is obtained. Commercial hydrochloric acid is generally yellow in colour, owing to the presence of iron as an impurity ; and it is always liable to contain sulphuric acid, free chlorine, arsenic, and some- times sulphur dioxide. This aqueous solution of hydrochloric acid is also known under the names of “spirits of salt,” and muriatic acid. Chlorine Monoxide 333 OXIDES AND OXYACIDS OF CHLORINE. The elements oxygen and chlorine have never been made to unite together directly : two compounds, however, of these elements can be obtained by indirect methods ; these are— Chlorine monoxide (hypochlorous anhydride) . ci 2 o. Chlorine peroxide . C10 2 . Three oxyacids are known, viz. : — Hypochlorous acid . HCIO. . HCIO 3 . Chloric acid Perchloric acid . HC10 4 . CHLORINE MONOXIDE ( Hypochlorous anhydride ). Formula, C1 2 0. Molecular weight = 86 . 70 . Density = 43-35- Mode of Formation. — This compound is obtained by passing dry chlorine over dry precipitated mercuric oxide contained in a glass tube, the temperature of which is not allowed to rise. The chlorine combines with the mercuric oxide, forming mercuric oxy- chloride, and chlorine monoxide is liberated — 2 HgO + 2C1 2 = HgO,HgCl 2 + C1 2 0. Properties. — At ordinary temperatures chlorine monoxide is a pale yellow gas, without the greenish tint possessed by chlorine. Its smell strongly suggests chlorine, but is readily distinguishable from it. It is a very unstable compound, decomposing with more or less violence with moderate rise of temperature. When strongly cooled it is condensed to an orange-yellow coloured liquid, which boils at about - 20 °. This liquid is extremely unstable, exploding with great violence on the gentlest application of heat, and some- times on merely being poured from one vessel to another. When exposed to direct sunlight it also explodes with violence. Gaseous chlorine monoxide is considerably soluble in water, one volume dissolving about 100 volumes of the gas, forming hypo- chlorous acid — C1,0 4- H a O = 2HC10. 334 / norganic Chemistry CHLORINE PEROXIDE. Formula, CI 0 2 . Molecular weight = 67.29. Density = 33.645 Modes of Formation. -( 1 .) By the action of sulphuric acid upon potassium chlorate — 3 KCIO 3 + 2H 2 S0 4 = KC10 4 + 2HKSO, + H 2 0 + 2C10 2 . F inely powdered potassium chlorate is added little by little to concentrated sulphuric acid in a small retort. The salt dissolves with the formation of a reddish liquid, and if the temperature is not allowed to rise, no gas is evolved. On very cautiously warm- ing the retort by means of warm water, taking care not to heat the glass above the level of the liquid in the retort, the chlorine peroxide is evolved. ( 2 .) A mixture of chlorine peroxide and carbon dioxide, in equal volumes, is obtained by heating a mixture of powdered potassium chlorate and oxalic acid to a temperature of 70 ° in a water-bath 2KC10 3 + 2H 2 C 2 0 4 = K 2 C 2 0 4 + 2H 2 0 + 2C0 2 + 20102. ( 3 .) Chlorine peroxide, mixed with chlorine, is evolved by the action of hydrochloric acid upon potassium chlorate— 4KC10 3 + 12HC1 = 4KC1 + 6H 2 0 + 9C1 + 3010,. This mixture of gases was formerly supposed to be a definite compound of oxygen and chlorine, and received the name of euchlorine. Properties.— Chlorine peroxide is a heavy gas, with a deep yellow colour. It has an intensely unpleasant smell, and if in- haled, even when largely diluted with air, produces headache. The gas attacks mercury, and is soluble in water, so that it can only be collected by displacement. Chlorine peroxide is an ex- tremely unstable compound, it is gradually resolved into its ele- ments by the influence of light ; the passage of an electric spark, or the introduction into it of a hot wire, causes it to decompose with violent explosion. It is a powerful oxidising compound ; a piece of phosphorus introduced into the gas takes fire spontane- ously. If a jet of sulphuretted hydrogen be lowered into a jar of 335 Hypochlorous Acid chlorine peroxide, the sulphuretted hydrogen ignites spontaneously, and continues burning in the gas. Its oxidising action upon organic matter, may be shown by liberating the gas in the presence of such a substance as sugar, by adding a drop of sulphuric acid to a mixture of powdered sugar and potassium chlorate. The chlorine peroxide, liberated by the action of the acid upon the chlorate, ignites the mixture, and the entire mass then bursts into flame. When chlorine peroxide is strongly cooled, it condenses to dark red liquid, which is even more explosive than the gas. HYPOCHLOROUS ACID. Formula, HCIO. Modes Of Formation.— (I.) As already mentioned, this acid is formed when chlorine monoxide is dissolved in water. (2.) It may readily be obtained in dilute solution, by passing an excess of chlorine through water in which precipitated mercuric oxide is suspended — HgO + H 2 0 + 2C1 2 = HgCl 2 + 2HC10. On distilling the liquid, the -dilute acid passes over as a colourless distillate. . (3.) In dilute solution, hypochlorous acid may be obtained by the decomposition of a hypochlorite by a very dilute mineral acid, and subsequent distillation of the mixture; thus, if to a solution of calcium hypochlorite (obtained by treating bleaching-powder with water and filtering the solution) very dilute nitric acid be added and the solution distilled, a dilute colourless acid is obtained— Ca(C10) 2 + 2HN0 3 = Ca(N 0 3 ) 2 + 2HC10. (4.) This compound is also formed, when a stream of chlorine is passed through water containing precipitated calcium carbonate in suspension — CaC0 3 + H 2 0 + 2C1 2 = CaCl 2 + C0 2 + 2HC10. Properties. — Pure hypochlorous acid, free from water, has never been obtained. The acid produced by the solution in water of chlorine monoxide, has a pale straw-yellow colour, and a very 336 Inorganic Chemistry characteristic chlorous smell. Dilute solutions of this acid are moderately stable, while more concentrated solutions readily undergo spontaneous decomposition. Hypochlorous acid is a powerful oxidising and bleaching agent as it readily gives up its oxygen, and is resolved into hydrochloric acid — HCIO = HC1 + O. As an oxidising agent it is twice as effective as an equivalent quantity of chlorine in chlorine water, for two atoms of chlorine are here necessary for the liberation of one atom of oxygen— Cl 2 + H 2 0 = 2HC1 + O. Hypochlorous acid is decomposed by hydrochloric acid, with the evolution of chlorine — HCIO + HC1 = H 2 0 + CL, It is also decomposed by silver oxide, oxygen being liberated Ag 2 0 + 2HC10 = 2AgCl + H 2 0 + 0 2 . The salts of hypochlorous acid may be obtained by the action of the acid upon the hydroxides of the metals ; thus HCIO + KHO = KCIO + H 2 0. The most important salt of this acid is bleaching-powder (see Calcium Salts). CHLORIC ACID. Formula, HC10 3 . Mode of Formation.— This compound is best obtained, by ecomposing barium chlorate with an exact equivalent of sulphuric acid, previously diluted with water— Ba(C10 3 ) 2 + H 2 S0 4 = BaS0 4 + 2HC10 3 . The clear liquid is decanted from the precipitated barium sul- phate, and is then concentrated by evaporation in vacuo. The strongest acid that can be obtained still contains 80 per Perchloric Acid 337 cent, of water. Attempts to concentrate it further, result in its decomposition into free chlorine and oxygen, with the formation of perchloric acid and water. Properties. — The strong aqueous acid has powerful oxidising properties ; many organic substances, as wood or paper, are so rapidly oxidised by it that when the acid is dropped upon them they are frequently inflamed. The acid even in dilute solution has strong bleaching powers. The salts of chloric acid are far more stable than the acid, and some of them are of considerable technical importance. The chlorates are all soluble in water, and all yield oxygen on being heated. Chloric acid is a monobasic acid ; the chlorates, there- fore, have the general formula M'C10 3 and M"(C 10 3 ) 2 , where M' and M" stand for monovalent and divalent metals respectively. Of all the chlorates, potassium chlorate, KC10 3 , is by far the most important. (See Potassium Compounds.) PERCHLORIC ACID. Formula, HC10 4 . Mode Of Formation. — Perchloric acid is best prepared, by the action of strong sulphuric acid upon potassium perchlorate — 2KC10 4 + H 2 S0 4 = K 2 S0 4 + 2HC10 4 . Pure and dry potassium perchlorate is mixed with four times its weight of concentrated sulphuric acid, and the mixture gently dis- tilled in a small retort. The distillate at first consists of perchloric acid ; but as the operation proceeds, a portion of the perchloric acid is decomposed into lower oxides of chlorine, and water, and the latter, combining with the first portions of the distillate, forms a white crystalline compound, having the composition HC10 4 ,H 2 0. This body, when gently heated, gives off perchloric acid ; it may, therefore, be employed for the preparation of the acid in a state of purity. Properties. — Perchloric acid is a colourless, volatile, and strongly fuming liquid, having a specific gravity of 1.782 at 15 0 . It is an extremely powerful oxidising substance ; a drop of the liquid allowed to fall upon paper, wood, or charcoal is instantly decom- posed, sometimes with a violent explosion. In contact with the Y 33 ^ Inorganic Chemistry skin it produces most painful wounds ; when allowed to drop into water it produces a hissing sound, owing to the energy of the combination. Perchloric acid cannot be preserved, as it slowly decomposes even in the dark, and often explodes spontaneously. The salts of this acid are the perchlorates, of which the most important is potassium perchlorate ; they are all soluble in water. Constitution of the Oxides and Oxyacids of Chlorine.— On the assump- tion that chlorine is a monovalent element, the constitution of these compounds may be thus represented : — Chlorine monoxide, Cl - O - Cl. I Hypochlorous acid, Cl - O - H. Chlorine peroxide, Cl- O-O- . I Chloric acid, Cl- O - O- O-H. Perchloric acid, Cl -0-0— 0-0- H. It is possible, however, that in some of these compounds the chlorine functions as a trivalent element, and that these compounds have a constitution similar to the oxides and oxyacids of nitrogen, thus — Chlorine monoxide, Cl - O - CL .'O Chlorine peroxide, — Cl< Nitrogen monoxide, N - O - N. SO Nitrogen peroxide, — N<^ | . \o Hypochlorous acid, Cl - O - H. Hyponitrous acid, N - O - H. /O < '• No Chloric acid, H-0-Cl<^ | . \o Nitric acid, H - O - N<^ | . No o Perchloric acid, H - O - Cl\^/0. There are several facts which point to the belief that not only chlorine, but also bromine and iodine, are capable of fulfilling the functions of a trivalent element. The existence, for example, of such a compound as trichloride of iodine, IC1 3 , is difficult to explain on any other assumption than that iodine is here a trivalent element. Indeed, from a consideration of the salts of periodic acid, some chemists are in favour of assigning to iodine even a still higher valency, and of regarding it as a heptad element in these compounds (see Periodates, page 356). The constitution of such molecules as those of hydrofluoric acid at low temperatures, namely, H 2 F 2 , and of the acid fluoride of potassium, HF,KF, is readily understood if we regard the fluorine as functioning in these compounds as a trivalent element, thus — H - F = F - H, and H - F = F — K. Bromine 339 BROMINE. Symbol, Br. Atomic weight = 79.76. Molecular weight = 159.52. Vapour density = 79.76. History. — This element was discovered by Balard (1826), in the mother-liquor obtained after the crystallisation of salt from con- centrated sea-water. He applied the name bromine (signifying a stench) to the element, in allusion to its unpleasant smell. Occurrence. — Bromine is never found in the uncombined state in nature. In combination chiefly with the metals potassium, sodium, and magnesium, it occurs in small quantities in all sea- water, and more abundantly in many mineral waters and salt springs. The saline deposits of Stassfurt contain notable quantities of bromides, and the main supply of bromine for the market is manufactured from this source. Modes of Formation. — (1.) Bromine may be obtained from a bromide by displacement with chlorine. If to a solution of mag- nesium bromide, chlorine water is added, the chlorine combines with the magnesium and the bromine is liberated — MgBr 2 + Cl 2 = MgCl 2 + Br 2 . On distilling the liquid the bromine is driven off, and can be collected in a well-cooled receiver. The addition of any excess of chlorine results in the formation of bromide of chlorine, and is therefore to be avoided. (2.) Bromine is readily obtained from potassium bromide by the action of manganese dioxide and sulphuric acid, a reaction exactly analogous to that by which chlorine is obtained from sodium chloride — 2KBr + Mn0 2 + 2H 2 S0 4 = MnS0 4 + K 2 S0 4 + 2H 2 0 + Br 2 . The mixture is gently distilled from a retort into a receiver kept cold by means of ice. (3.) Manufacturing Methods. — Practically all the bromine that is required at the present day, is manufactured from crude carnallite obtained at Stassfurt (see Alkali Metals). This salt contains bromine combined with magnesium, the magnesium bromide forming about 1 per cent, of the magnesium chloride in the crude substance. The final mother liquors from the manufac- 340 Inorganic Chemistry ture of potassium chloride, and which were formerly run to waste, are found to contain about .25 per cent, of bromine as magnesium bromide, and these liquors are now utilised for the manufacture of bromine. The bromine is liberated from its combination with magnesium, by means of chlorine. In some processes, the mother liquor is mixed with manganese dioxide and sulphuric acid in a stone vessel resembling an ordinary chlorine still. The magnesium chloride in the liquor, is acted upon by the manganese dioxide and sulphuric acid with the evolution of chlorine, and this decomposes the bromide present, displacing the bromine — MgCl 2 + Mn0 2 + 2H 2 S0 4 = MnS0 4 + MgS0 4 + 2H 2 0 + Cl 2 . MgBr 2 + Cl 2 = MgCl 2 + Br 2 . The bromine that is driven out, is condensed by means of a worm condenser. Bromine 341 Instead of the chlorine being generated within the mother liquor, it is now more usually produced in a separate chlorine still, and passed into the liquor. Fig. 100 shows in diagrammatic form the method employed. The hot mother liquor is admitted by the pipe A into the tower T, which is filled with earthenware balls, between which the liquid percolates. It leaves the tower by the pipe B, and flows into the tank W, which is provided with shelves in such a way that the liquid must circulate through it in the direction indicated by the arrows. The exit-pipe from this tank, empties into a waste, placed at such a height that the tank is always nearly full. The liquid in the tank is kept at, or near, the boiling-point, by means of a current of steam blown in through S. Chlorine from a still is admitted by the pipe L, and passing into the tower by the pipe B, travels in an opposite direction to the current of liquid. As the chlorine passes up the tower, it meets the descending mother liquor, and decomposes the magnesium bromide contained in it with the liberation of bromine. The bromine vapour leaves the tower by the pipe C, and is conveyed to a worm (Fig. 101), where it is condensed. Any bromine which dissolves in the water in the tower, is again expelled from solution by the steam as the liquid tra- verses the tank W, and is swept up into the tower by the current of chlorine. The condensed bromine, as it leaves the worm, is collected in a tubulated bottle, and any vapour which escapes condensation 342 Inorganic Chemistry is arrested by the vessel F, Fig. ioi. This tube is filled with iron borings, kept moist by the constant dropping of water upon them, and any bromine, or bromide of chlorine, is there converted into iron compounds, which are dissolved by the water, and flow away into the receiver. The bromine is purified by redistillation. Properties.— Bromine is a heavy but mobile liquid, of a deep reddish-brown colour. Except in extremely thin layers it is opaque. It is the only non-metallic element which is liquid at the ordinary temperature. Bromine boils at 59 0 , but being a very volatile liquid it gives off vapour rapidly at the ordinary temperature. A drop of bromine allowed to fall into a flask, immediately evaporates and fills the vessel with a dark red-brown vapour. The specific gravity of the liquid at 0° is 3.188. At - 7 0 bromine solidifies to a crystal- line mass. Bromine has a powerful and disagreeable smell. When the vapour, largely diluted with air, is inhaled, it suggests chlorine by its smell and by its action upon the mucous membrane of the throat and nose ; it has in addition, however, a most irritating action upon the eyes. It is very poisonous, and the liquid exerts a corrosive action upon the skin ; it produces a yellow colour when brought in contact with starch. The vapour density of bromine, taken at moderately high tem- peratures, gradually becomes less than is demanded by the formula Br 2 , showing that dissociation takes place. In the case of bromine this is more marked than with chlorine. Bromine is soluble in water, imparting its own colour to the solution which is known as bromine water. 100 grammes of water at o dissolve 3*60 grammes of bromine. The solubility steadily diminishes as the temperature rises : at 20° it is 3.208, and at 30° it is 3.126. When bromine water is cooled to o° it deposits a crystalline hydrate similar in composition to the hydrate of chlorine, Br 2 ,10H 2 O. Bromine resembles chlorine in its chemical attributes ; it com- bines directly with metals and many other elements, although with less energy than is exhibited by chlorine. A fragment of arsenic, for example, when dropped upon bromine, ignites and bums upon the surface of the liquid. Like chlorine, it has bleaching properties, due to its power of combining with hydrogen. Hydrobromic Acid 343 HYDROBROMIC ACID. Formula, HBr. Molecular weight = 80.76. Density = 40.38. Modes of Formation.— ( 1.) Hydrobromic acid can be obtained by the direct union of its elements. Bromine vapour and hydrogen, when mixed, do not combine under the influence of light ; neither does such a mixture explode when a light; is applied to it. The mixture, however, may be caused to burn, when hydrobromic acid is formed ; or, if the mixed gases be passed through a red-hot tube, the same result follows. A simple method of preparing hydro- bromic acid synthetically, consists in passing a mixture of hydrogen and bromine vapour over a spiral of platinum wire, maintained at a red heat by means of an electric current.* (2.) The best method for the preparation of gaseous hydrobromic acid, consists in dropping bromine upon red phosphorus which has been moistened with a small quantity of water, when tribasic phosphoric acid is formed, and hydrobromic acid is liberated — P + 4H 2 0 + 5Br = H 3 P0 4 + 5HBr. We may suppose that in this reaction the bromides of phosphorus are formed and simultaneously decomposed, the action of water upon these compounds being thus expressed — PBr 3 + 3H 2 0 = H 3 P0 3 + 3 HBr. PBr 6 + 4H 2 0 = H 3 P0 4 + 5HBr. (3.) Hydrobromic acid may be obtained by the action of phos- phoric acid upon potassium bromide — 3KBr + H 3 P0 4 = K 3 P0 4 + 3HBr. (4.) If sulphuric acid be employed (as in the formation of hydro- chloric acid from a chloride), free bromine is simultaneously pro- duced, owing to the reduction of a portion of the sulphuric acid by the hydrobromic acid which is first evolved, thus — H 2 S0 4 + 2HBr = S0 2 + 2H 2 0 + Br 2 . (5.) A dilute aqueous solution of hydrobromic acid may also be * See “Chemical Lecture Experiments,” new ed. , No. 225. 344 Inorganic Chemistry obtained by passing a stream of sulphuretted hydrogen through bromine water — (6.) Hydrobromic acid is readily obtained, by the action of bromine upon certain hydrocarbons, such as turpentine or melted paraffin. The action is one of substitution, one atom of bromine replacing one atom of hydrogen in the compound, and the hydrogen so displaced combining with a second bromine atom to form hydro- bromic acid. Thus, if the hydrocarbon be represented by the general formula, C n H2n + 2, the action of bromine will be repre- sented thus — Properties. Hydrobromic acid is a colourless, pungent-smel- ling gas, which fumes strongly in the air. It is extremely soluble in water, forming an acid liquid strongly resembling aqueous hydrochloric acid. When boiled, this solution loses either acid or water, until it reaches a degree of concentration at which it contains 48 per cent, of hydrobromic acid. The acid of this strength then continues to boil unchanged at 126°. As with hydrochloric acid, the strength of the liquid which boils at a constant temperature, depends upon the pressure. Hydrobromic acid is decomposed by chlorine, with the liberation of bromine — In its chemical behaviour, hydrobromic acid closely resembles hydrochloric acid, and this resemblance is extended to the bromides. All bromides are soluble in water, except mercurous bromide, silver bromide, and lead bromide, the latter being slightly soluble! No oxides of bromine corresponding with the oxides of chlorine have as yet been obtained ; two oxyacids, however, are known, viz. — SH 2 + Br 2 = S + 2HBr. C n H2n + 2 + Br 2 = C n H2n + iBr + HBr. 2HBr + Cl 2 = 2HC1 + Br 2 . OXYACIDS OF BROMINE. Hypobromous acid Bromic acid . . HBrO. . HBrO,. Hypobromous Acid 345 HYPOBROMOUS ACID. Formula, HBrO. Mode Of Formation.— An aqueous solution of hypobromous acid may be obtained, by shaking together admixture of bromine water and precipitated mercuric oxide, the reaction being ana- logous to that by which hypochlorous acid is prepared — HgO + H 2 0 + 2Br 2 = HgBr 2 + 2HBrO. Properties. — Hypobromous acid is an unstable compound ; it breaks up on distillation into oxygen and bromine. By heating to 40° in vacuo, however, it can be distilled without decomposition. The aqueous liquid so obtained has a pale yellow colour. It readily gives up its oxygen, and is a strong bleaching agent ; when heated to about 6o° it decomposes. BROMIC ACID. Formula, HBr0 3 . Modes of Formation.— (1.) This acid is only known in aqueous solution ; in this form it may be obtained by the action of bromine upon silver bromate in the presence of water 5AgBrO s + 3Br 2 + 3H 2 0 = 5AgBr + 6HBr0 3 . The insoluble silver bromide separates out, and the aqueous acid can be decanted from the precipitate. (2.) A solution of this acid, mixed with hydrochloric acid, is also formed when chlorine is passed through bromine water— Br 2 4- 5C1 2 4- 6H 2 0 = 10HC1 + 2HBr0 3 . (3.) The decomposition of barium bromate by the requisite weight of sulphuric acid, affords the best method for the preparation of a pure aqueous solution of bromic acid Ba(Br0 3 ) 2 4- H 2 S0 4 = BaS0 4 + 2HBr0 3 . Properties. — Bromic acid is an unstable, strongly acid sub- stance, closely resembling chloric acid. The aqueous solution may be concentrated in vacuo until it contains about 5° P er cent. 346 Inorganic Chemistry of bromic acid, representing a composition of i molecule of the acid to 7 of water. Beyond this degree of concentration, or if heated to ioo°, the acid decomposes into bromine, oxygen, and I he bromates are formed by reactions similar to those by which the chlorates are produced ; thus, by adding bromine to a solution of potassium hydroxide, a mixture of potassium bromide and bromate is obtained — 6KHO + 3Br 2 = 5KBr + KBr0 3 + 3H 2 0. And the two salts can be separated by crystallisation, owing to the greater solubility of the bromide. The bromates decompose on being heated, some with the liberation of oxygen and formation of bromide— KBr0 3 = KBr + 30, but without the intermediate production of a perbromate. Others give off their bromine as well as a part of the oxygen they contain leaving the metal in combination with oxygen — Mg(Br0 3 ) 2 = MgO + Br 2 + 50. IODINE. Symbol, I. Atomic weight = 126.54. Molecular weight = 253.08. Vapour density = 126.54. History.— Iodine was discovered by Courtois (1812), who ob- served that a beautiful violet vapour was evolved, during his endeavours to prepare nitre from liquors obtained by lixiviating the ashes of burnt seaweed. The substance was subsequently investigated by Gay-Lussac. Occurrence. Like all the other members of this group of elements, iodine is never found in nature in the uncombined con- dition. In combination it occurs associated principally with potassium, sodium, magnesium, and calcium, as iodides and lodates. Iodine is a widely distributed element, although not occurring in more than small quantities in any particular source. Thus it is found in small quantities in sea-water and in both marine plants and animals. The amount of iodine in seaweed, varies with diffe- Iodine 347 rent plants ; generally speaking, those from greater depths, contain more than weeds which grow in comparatively shallow waters. Dry Weeds. Drift weed Cut weed ( Laminaria digitata (stem) 1 Laminaria stenophylla • j Fucus serratus . I Ascophyllum nodosum Per Cent, of Iodine. • 0.4535 . 0.4777 . 0.0856 . 0.0572 Iodine is also found in small quantities in many mineral waters and medicinal springs. In small quantities iodine is present in the natural sodium nitrate of Chili and Peru, known as Chili saltpetre, and at the present day this constitutes the most abundant source of this element. Mode of Formation.— Iodine may be readily obtained by a precisely similar reaction to that by which both bromine and chlorine are produced ; thus, if potassium iodide be mixed with manganese dioxide and sulphuric acid, and the mixture gently heated in a retort, iodine distils over and condenses in the form of greyish black crystals — 2KI + Mn0 2 + 2H 2 S0 4 = K 2 S0 4 + MnS0 4 + 2H 2 0 + I 2 . Manufacturing Processes.— On an industrial scale iodine is obtained from two sources, namely, from seaweed and from caliche (Chili saltpetre). (1.) From seaweed. The weeds chiefly employed, are the Lami- naria digitata and Laminaria stenophylla. The weed is burnt in shallow pits, care being taken to avoid too high a temperature ; the maximum yield of iodine being obtained if the ash is not allowed to fuse. This ash is technically known as kelp, and if the weed is properly burnt, it should yield a kelp containing from 25 to 30 lbs. of iodine per ton. The kelpers, however, usually lose about half the iodine on account of burning the weed at too high a tempera- ture, thereby fusing the ash into a hard slag, instead of obtaining a porous residue. An improved process of carbonising the weed, was introduced by Stanford (1863), in which it was heated in large retorts, whereby the volatile products of the distillation, consisting largely of tar and ammoniacal liquor, could be collected. The kelp obtained by this method is in a very porous condition, and contains the whole of the iodine originally present in the weed. 34 ° Inorganic Chemistry A still more recent process for extracting the iodine from sea- weed, and at the same time obtaining other useful materials, has since been discovered by Stanford. The weed is boiled with sodium carbonate and filtered : the residue consists of a substance called algulosc. Hydrochloric acid is added to the filtered liquid which precipitates a compound known as alginic acid , and this is again separated by filtration. The liquor is neutralised will sodium hydrate, evaporated to dryness and carbonised. The residue, which is known as “kelp substitute,” contains all the iodine, as well as the potash salts, and should yield about to lbs of iodine per ton. [The alginic acid obtained in this process, is purified and converted into the sodium salt, which constitutes the commercial “ algin,” a material of a gelatin- apphca^nsT n0US ^ ** be6n P Ut t0 a " umber of useful The kelp obtained by either of these methods, is lixiviated with water in large iron vats, whereby all the soluble salts are extracted This aqueous liquid is concentrated in large open boiling pans, and the less soluble salts, viz., the alkaline sulphates, carbonates and chlorides, are allowed to crystallise. The mother liquor is then mixed with sulphuric acid and allowed to stand. The sul- phuric acid decomposes any sulphides and sulphites which may be present, with the separation of sulphur ; it also converts the bromides and iodides into the corresponding sulphates, with the liberation of hydrobromic and hydriodic acids which remain in solution, while the alkaline sulphates are deposited from the liquid and are technically known as plate sulphate. The liquor is then transferred to the iodine still , which is an iron pot furnished with a leaden cover into which two exit-pipes are fixed (Fig. 102 ). These are connected to a series (usually ten in each row) of large earthenware jars or aludels. A gentle heat is applied, and manganese dioxide is introduced from time to time through the opening. The iodine is evolved according to the following equation — 6 2H1 + Mn0 2 + H 2 S0 4 = MnS0 4 + 2H 2 0 + I 2 , and condenses in the jars. These vessels are also furnished with a tubulus upon their under side, so that the water which is evolved during the distillation can drain out, and run off down the trough in which the jars are resting. Iodine 349 ( 2 .) From Chili satlletre. The crude sodium nitrate of Chili and Peru, known as caliche , contains small quantities of iodine, chiefly as sodium iodate. Although the amount of iodine in caliche is only very small, averaging about 0.2 per cent., in view of the enormous quantity of nitrate that is turned out, the aggre- gate amount of iodine is very great. This iodine is now extracted, and the supply of this element that is now manufactured from this source, is more than the total consumption of iodine in the whole world. The process is based upon the fact, that when a solution Fig. X02. of hydrogen sodium sulphite (sodium bisulphite) is added to a solution of an iodate, iodine is precipitated, thus— 2NaI0 3 4 - 5NaHS0 3 = 3NaHS0 4 4- 2Na 2 S0 4 4- H 2 0 4- I 2 . The final mother liquor from the sodium nitrate, or caliche, in which all the iodate has concentrated, contains as much as 22 per cent, of this salt. This liquor is mixed with the requisite proportion of the hydrogen sodium sulphite solution, in large lead- lined vats, and the precipitated iodine allowed to settle. It is then washed and pressed into blocks, and is found to contain from 80 to 85 per cent, of iodine. This impure product is then distilled at a gentle heat from iron retorts, the vapour being con- densed in a series of earthenware receivers much as in the older method. Properties. — Iodine is a bluish-black shining solid, somewhat resembling graphite in lustre and general outward appearance. It 350 Inorganic Chemistry crystallises in large brilliant plates, which have a specific gravity of 4.95. When heated to 107” iodine melts, and gives off vapour having a beautiful violet colour. Its boiling-point is about 175 0 . Iodine vaporises slowly at ordinary temperatures, and sublimes from one part to another of a bottle in which a small quantity of Jt is contained. The smell of iodine vapour is somewhat irritating and unpleasant, recalling the smell of moderately diluted chlorine. When iodine vapour is heated, it passes from a violet colour to a deep indigo blue* This change in the colour is accompanied by a diminution of the vapour density. Up to a temperature of 700° the density of iodine corresponds to the formula I 2 : as the temperature is raised the density gradually diminishes, until at 1468 it is reduced to less than two-thirds. At this point, 73.1 per cent, of the iodine molecules have become dissociated into single atoms. Iodine is slightly soluble in water, 1 gramme of iodine being dissolved by 5.524 litres of water at io°. This dilute solution, however, has a perceptible brown colour. Iodine is freely soluble in aqueous potassium iodide solution, in alcohol, ether, and aqueous hydnodic acid ; in all these solvents it dissolves to a dark reddish brown solution. In chloroform, carbon disulphide, and many liquid hydrocarbons, iodine is also soluble, but in these solvents it dissolves to a deep violet solution, resembling the colour of the vapour. When iodine is brought in contact with starch, it forms an intense blue colour. This reaction is so extremely delicate, that it is capable of revealing the minutest trace of iodine. The exact nature of this blue compound is not known. The colour disappears when the liquid is heated to about 8o°, but returns on cooling ; continued boiling destroys it permanently. In its chemical relations iodine resembles chlorine and bromine, but with a lesser degree of energy. Both these elements are capable of displacing iodirre from its combinations with electro- positive elements, thus — KI + Br = KBr + I. KI + Cl = KC 1 + I. Iodine combines with many elements, both metals and non- metals forming iodides. Phosphorus, when brought in contact with iodine, at once melts and inflames ; antimony powder dropped * “Chemical Lecture Experiments,” new ed., No. 231. 35i Hydriodic Acid into iodine vapour also spontaneously inflames. When mercury and iodine are gently heated, energetic combination takes place, and mercuric iodide is formed. HYDRIODIC ACID. Formula, HI. Molecular weight = 127. 54. Density=255.o8. Modes of Formation. — (1.) Hydriodic acid can be obtained synthetically, by passing a mixture of hydrogen and iodine vapour over strongly heated, finely divided platinum. (2.) It is also obtained by the action of phosphoric acid upon sodium or potassium iodide. (See Hydrobromic Acid.) As in the case of the corresponding bromine compound, sul- phuric acid cannot be employed, as by its action upon the iodide, iodine and sulphur dioxide are liberated, thus — 2KI + 3H 2 S0 4 = 2HKS0 4 + 2H 2 0 + S0 2 + I 2 . (3.) Hydriodic acid is produced by the action of sulphuretted hydrogen upon iodine (p. 371). At the ordinary temperature, and in the absence of water, these two substances do not react, hydriodic acid being an endother- mic compound (p. 147) ; but if the iodine be suspended in water and sul- phuretted hydrogen passed through, the heat of solution of the hydriodic acid supplies the neces- sary energy to en- able the action to proceed. When, however, the solu- tion reaches a sp. g. of 1.56 the action ceases, because, as Naumann has shown, the heat produced by the solution of the product is insufficient to carry on the process beyond this degree of concentration. (4.) Hydriodic acid is most readily prepared, by the action of phosphorus upon iodine in the presence of water — P + 51 + 4H 2 0 = H 3 P0 4 + 5HI. 352 Inorganic Chemistry The red phosphorus and iodine for this reaction may be placed in a dry flask, and water gradually dropped upon the mixture, when hydriodic acid is rapidly evolved. The gas is allowed to pass through a U-tube containing red phosphorus, in order to arrest any iodine vapour which may accompany it. Hydriodic acid may be collected over mercury or by displacement, as shown in Fig. 103. Properties. Hydriodic acid is a colourless, pungent-smelling gas, which fumes strongly on coming into the air. The gas is readily decomposed by heat into hydrogen and iodine. Thus, if a heated wire be thrust into the gas, or if a spiral of platinum wire be heated in the gas by means of an electric current, the violet vapour of iodine at once makes its appearance. When mixed with chlorine, hydriodic acid is at once decomposed, with the liberation of iodine, thus — 2 HI + Cl 2 = 2 HC 1 + I 2 . Hydriodic acid is one of the most readily liquefied gases; at o°, and under a pressure of four atmos- pheres, it condenses to a colourless liquid. The gas is extremely soluble in water. An aqueous solution of it is readily pro- duced, by allowing the gas, obtained by the method of preparation above de- scribed, to pass into water. In order to prevent the water from being drawn back into the generating flask, it is convenient to pass the gas through a retort arranged in the position seen in Fig. 104. Should there be any back rush of water, owing to the intermission of the evolution of gas in the apparatus, the liquid in the beaker will be drawn up into the retort and there lodge, leaving the end of the neck open to the air. A saturated aqueous solution of hydriodic acid at o° has a specific gravity of 2. At the ordinary pressure the strongest acid that can be obtained by distillation has a specific gravity of 1.67, and contains 57.7 per cent, of hydriodic acid. This solution boils at 1 2 7 0 . As in the case of the corresponding bromine and chlorine compounds, the particular strength of acid which has a constant boiling-point, is a function of the pressure. Fig. 104. Iodic Acid 353 Aqueous hydriodic acid, when freshly prepared, is colourless ; but it rapidly turns brown, owing to the oxidation of the compound, and the solution of the liberated iodine in the acid — 4HI + 0 2 = 2H 2 0 + 2 I 2 . OXIDE AND OXY ACIDS OF IODINE. One compound of iodine with oxygen is known, and two oxy- acids, viz. : — Iodine pentoxide .... I 2 0 6 . Iodic acid HIO s . Periodic acid .... HI0 4 . IODINE PENTOXIDE {Iodic Anhydride). Formula, I 2 0 5 . Mode of Formation.— When iodic acid is heated to 170 °, it loses water and is converted into the pentoxide — 2HI0 3 = H 2 0 + I 2 0 6 . Properties. — Iodine pentoxide is a white crystalline solid body. It is soluble in water, combining with a molecule of the water to form iodic acid. Iodine pentoxide is more stable than any of the oxides of the other halogens ; but, when heated to a temperature of 300 °, it decomposes into its elements. IODIC ACID. Formula, HIO 3 . Modes Of Formation. — ( 1 .) Iodic acid can be prepared by adding to a solution of barium iodate the requisite amount of sulphuric acid demanded by the equation — Ba(IO s ) 2 + H 2 S0 4 = BaS0 4 + 2HI0 3 . The aqueous solution of iodic acid is decanted from the preci- pitated barium sulphate, and may be concentrated at ioo° without being decomposed. . ( 2 .) When chlorine is passed through water in which powdered z 354 Inorganic Chemistry iodine is suspended, a mixture of iodic acid and hydrochloric acid is produced — 3H 2 0 + I + 5C1 = 5HC1 + HIO s . The hydrochloric acid may be removed by the addition of preci- pitated silver oxide to the solution, and separating the precipitated silver chloride by filtration. (3.) Iodic acid is most conveniently prepared by heating iodine with nitric acid, whereby the iodine is oxidised, and a mixture of oxides of nitrogen is evolved as dense red vapours — 3HN0 3 + I = H 1 0 3 + H 2 0 + N 2 0 3 + NO* Properties. — Iodic acid is a white crystalline solid, soluble in water. The aqueous solution shows an acid reaction with litmus, but the colour is ultimately discharged by the bleaching action of the compound. Iodic acid does not form any blue colour with starch ; being, however, an oxidising substance, it readily gives up oxygen to such reducing agents as sulphur dioxide, sulphuretted hydrogen, or hydriodic acid, with the liberation of iodine, thus — 2HI0 3 + 4H 2 0 + 5S0 2 = 5H 2 S0 4 + I* 2HIO a + 5H 2 S = 5S + 6H 2 0 + l 2 . H10 3 + 5HI = 3H 2 0 + 3I 2 . If, therefore, a small quantity of sulphurous acid be added to a dilute solution of iodic acid, previously mixed with starch, the blue iodide of starch will be formed. This reaction affords an excellent illustration of the time required for certain chemical changes to go forward. It is readily possible to obtain an interval of 30 to 60 seconds between the addition of the sulphurous acid and appear- ance of any visible result, when at the expiration of that time the whole mass of the liquid suddenly turns blue.* Iodates. — When iodine is dissolved in potassium hydroxide, a mixture of potassium iodide and iodate is produced, by an analogous reaction to that which takes place with either bromine or chlorine — 6KHO + 3I 2 = 5KI + KIO s + 3H 2 0. With the exception of the iodates of the alkali metals, the iodates are for the most part insoluble in water. On being heated they * See Experiment 246, “Chemical Lecture Experiments,” new ed. Periodic Acid 355 behave in a similar manner to the bromates ; some being decom- posed into an iodide and oxygen, while others leave a metallic oxide and evolve iodine as well as oxygen. The alkaline iodates are capable of uniting with iodic acid, forming salts which are termed acid and di-acid iodates, thus — Normal potassium iodate . . . KI0 3 . Acid potassium iodate . . . KI0 3 ,HI0 3 . Di-acid potassium iodate . . . KI0 3 ,2HI0 3 . PERIODIC ACID. Formula, HI0 4 ,2H 2 0 or H 5 I0 6 . Modes of Formation. — (i.) The compound represented by the formula HI 0 4 has never been obtained; when aqueous solutions of periodic acid are evaporated, the compound which crystallises out has the composition HI0 4 ,2H 2 0, or H 6 10 6 . It may be obtained by boiling silver periodate with water, when an insoluble basic silver salt is produced — 2AgI0 4 + 4H 2 0 = Ag 2 H 3 I0 6 + HI0 4 ,2H 2 0. The silver periodate is prepared by passing chlorine into an aqueous solu- tion of sodium iodate and sodium hydroxide, when the sparingly soluble disodium periodate separates out — NaI0 3 + 3NaHO + Cl 2 = 2NaCl + Na 2 H 3 I0 6 . This sodium salt is then dissolved in nitric acid, and silver nitrate added, whereby AgI0 4 is formed, which crystallises out on concentration — ( 2Na 2 H 3 I0 6 + 2HN0 3 = 2NaNO a + 4H 2 0 + 2NaI0 4 . X 2NaI0 4 + 2AgN0 3 = 2NaN0 3 + 2AgI0 4 . ( 2 .) Periodic acid is also formed by the addition of iodine to an aqueous solution of perchloric acid — 2 HCIO 4 + 2H 2 0 + I 2 = Cl 2 + 2HI0 4 ,2H 2 0. Properties. — The acid having the composition HI0 4 ,2H 2 0 is a colourless, crystalline, deliquescent substance. It melts at 1 33 0 , and at 150 ° is decomposed into iodine pentoxide, water, and oxygen — 2H 6 I0 6 = I 2 0 6 + 5H 2 0 + 0 2 . The acid cannot be converted into HI0 4 by heat, for oxygen is evolved as soon as water begins to be given off. 356 Inorganic Chemistry The Periodates constitute a numerous class of salts, many of them being of a highly complex composition. On the assumption that iodine is monovalent in these compounds, their classification is somewhat difficult, and they must be represented as associations of molecules of salts of the unknown mono-basic periodic acid, HI0 4 , with metallic oxide and water in various proportions thus, the silver periodate in the foregoing equation, Ag 2 H 3 I0 6 , would be expressed by the formula, 2AgI0 4 ,Ag 2 0,2H 2 0. The classification of these compounds is much simplified, if we regard iodine as here functioning as a heptavalent element. On this assumption the perio- dates may be considered as the salts of various hypothetical acids, which are all derived from the compound I(HO ) 7 (itself hypothetical) by the withdrawal of varying quantities of water. Thus, by the successive removal of one mole- cule of water the following three acids would be formed— I(HO ) 7 - H 2 0 = IO(HO ) 5 . IO(HO ) 5 - H a O = I0 2 (H0 ) 3 . I0 2 (H0 ) 3 - H 2 0 = IOj(HO) . From these three acids the following salts derived — • H 5 IO„. (I.) • H 3 IO s . ( 2 .) • hio 4 . ( 3 .) may be regarded as being ( 1 .) Na 2 H 3 IOg ; Ag 2 H 3 I0 6 ; Ag 5 IO e ; Ba 5 (I0 6 ) 2 . ( 2 .) Ag 3 I0 5 ; Pb (10 5 ) 2 . ( 3 .) KI0 4 ; Agl0 4 . By the abstraction of one molecule of water from two molecules of these acids, still more complex acids would be derived, thus— IO(HO ) 4 ioIhoIs “ H 2 ° = o or h 8 i 2 o u . IO(HO ) 4 I0 2 (HO ) 2 I0 2 (H0 ) 3 “ H 2° = O or H 4 r 20 9 . io 2 (HO ) 2 And from these two acids the following periodates may be regarded as being derived — ( 4 .) Zn 4 I 2 O n ; Ba 4 I 2 O u . (5-) Ag 4 I 2 0 9 ; Ca2l 2 0 9 ; Ba2l 2 0 9 . (4.) (SO Compounds of the Halogens with each Other. Chlorine unites both with bromine and with iodine, and the two latter elements combine with each other. ( 1 .) Chlorine and Bromine . — Bromine monochloride. This substance is obtained as a reddish-yellow liquid, when chlorine gas is passed into bromine. The compound is believed to have the composition BrCl. ( 2 .) Chlorine and Iodine .— Iodine monochloride, IC1. When dry chlorine is passed over iodine, the latter rapidly melts, forming a dark reddish-browr. Iodine Trichloride 357 liquid, strongly resembling bromine in appearance. The liquid solidifies to a mass of red prismatic crystals, which melt at 25°. It -is decomposed by water into iodic and hydrochloric acids, and iodine is liberated— 5IC1 + 3H a O = HI0 3 + 5HC1 + 2I 2 . Iodine trichloride , IC1 3 . This compound is formed by passing an excess of chlorine over iodine, or by passing chlorine through iodine monochloride. It is also formed when hydriodic acid is acted upon by an excess of chlorine— HI + 2C1 2 = HC1 + IC1 3 . Iodine trichloride is a yellow solid substance, crystallising in long brilliant needle-shaped crystals, which sublime at the ordinary temperature. When gently warmed it melts, at the same time dissociating into chlorine and the monochloride ; on cooling, re-union takes place with the reformation of IC1 3 . (3.) Bromine and Iodine. — Two compounds of these elements are believed to exist, viz. , a crystalline solid, and a deep-coloured liquid. Their composition is probably expressed by the formulae, IBr and IBr 3 . CHAPTER II THE ELEMENTS OF GROUP VI. (. FAMILY B.) Oxygen, O . 15.96 I Selenium, Se . 78.87 Sulphur, S . . 31.98 | Tellurium, Te . 125 The relation in which oxygen, the typical element, stands to the remaining members of the family is very similar to that between fluorine and the other halogens. All the elements of this family unite with hydrogen, forming compounds of the type RH 2 — OH 2 , SH 2 , SeHg, TeH 2 ; but the hydride of oxygen stands apart from the others in many of its attributes. Thus at ordinary temperatures it is a colourless and odourless liquid, while the remaining compounds are all foetid- smelling and poisonous gases. Sulphur, selenium, and tellurium each combine with oxygen, forming respectively S0 3 , Se0 3 , and Te0 3 , while none of these elements in a divalent capacity forms a similar compound ; that is to say, no such combinations are known as OS 3 , or OSe 3 , although amongst themselves they unite, forming SeS 2 and TeS 2 . Sulphur, selenium, and tellurium also unite with oxygen, forming dioxides, S0 2 , SeOg, and Te0 2 , in which these elements are pos- sibly tetravalent, in which case the constitution of the compounds will be represented thus, 0 = S = 0; 0 = Se = 0. We may, however, consider them as functioning in a divalent in which case we may look upon ozone as being the corresponding /O oxygen compound, 00^ | . \o 358 Sulphur 359 All the elements of this family combine with chlorine, producing compounds having the following composition Oxygen. Sulphur. Selenium. o 2 ci S 2 C1 2 Se 2 Cl 2 OCl 2 SC1 2 SC1 4 SeCl 4 Oxygen again differs from the other members, by alone forming a compound of the type, R a CL This element also shows no ten- dency to function with a higher atomicity than that of a divalent , while the others unite with four atoms of the halogen, thereby exhibiting their tetravalent nature. The members of this family pass by a regular gradation from the strongly electro-negative, gaseous, non-metal oxygen, to t e feebly negative and slightly basic element tellurium, which possesses many of the properties of a true metal. Selenium and tellurium are both elements which lie very close to that ill-defined boundary between the metals and non-metals, and are on this account some- times termed metalloids. In tellurous oxide, TeO a , we have a compound which is both an acid-forming and a salt-forming oxide, its acidic and basic properties being nearly equally balanced. Thus, it replaces hydrogen in sulphuric acid, forming tellurium sulphate, Te(S0 4 ) 2 ; and it also unites with water, forming tellurous acid, H 2 Te0 3 , corresponding to sulphurous acid, H 2 S0 3 . Of the four elements of this family, oxygen is by far the most abundant, both in combination and in the free state ; sulphur is more plentiful than the other two, and tellurium occurs in the smallest quantity. The element oxygen has already been treated in Part II., p 1 59* Tellurium. TeCl 2 TeCl 4 SULPHUR. Symbol, S. Atomic weight = 31.98. Molecular weight = 63.96. Occurrence —In the free state this element occurs chiefly in volcanic districts. In Italy and Sicily large quantities of native sulphur are found, which have long been the most important European sources of this substance. Large deposits are to be met with in Transylvania and in Iceland, and it also occurs in beds, often of great thickness, in parts of China, India, California, and 360 / norganic Chemistry the Yellowstone district of the Rocky Mountains. These natural deposits are sometimes found stratified with beds of clay or rock, but they often occur as what are known as “living” beds, in which the sulphur is continuously being formed as the result of chemical decompositions which are at present at work. Such a “living” sulphur bed is known as a solfatara, and, as in the case of the Ice- land deposits, they are usually found associated with geysers, fumaroles, and other signs of volcanic action. In combination with hydrogen, sulphur occurs as sulphuretted hydrogen. Enormous quantities of sulphur are found combined with various metals, constituting the important class of substances known as sulphides ; as, for example, galena, , or lead sulphide, PbS ; zuic blende , or zinc sulphide, ZnS ; pyrites , or iron sulphide, FeS 2 ; copper pyrites , or copper iron sulphide, Cu 2 Fe 2 S 4 ; stibnite, or antimony sulphide, Sb 2 S 3 ; cinnabar , or mercury sulphide, HgS.* In combination with metals and oxygen, sulphur occurs in sulphates, such as gypsum, CaS0 4 ; heavy spar , BaS0 4 ; kieserite , MgS0 4 . Modes of Formation.— (i.) Sulphur is formed when sulphu- retted hydrogen is brought in contact with sulphur dioxide ; the two gases mutually decompose one another, with the formation of water and the precipitation of sulphur— 2H 2 S + S0 2 = 2H 2 0 + 3S. ( 2 .) It is also produced when sulphuretted hydrogen is burnt with an insufficient supply of air— H 2 S + O = H 2 0 + S. This reaction probably takes place in two stages, a portion of the sulphuretted hydrogen burning to sulphur dioxide, and this then reacting upon a further quantity of sulphuretted hydrogen, thus — (a) H 2 S + 30 = H 2 0 4- S0 2 . (b) 2H 2 S + S0 2 = 2H 2 0 + 3S. It is supposed that some of the free sulphur found in volcanic regions, has been produced by this action of these two gases upon one another. Extraction of Sulphur from Native Sulphur.— Natural sulphur is always more or less mixed with earthy or mineral Sulphur 361 matters, from which it is necessary to free*it. This is usually effected by melting the sulphur and allowing it to flow away from the accompanying impurities. The crude sulphur rock is stacked in brick kilns having a doping floor, and the mass ignited by introducing through openings in the heap, burning faggots of brushwood. The heat produced by the combustion of a part of the sulphur, causes the remainder to melt and collect upon the sloping floor of the kiln, from which it can be drawn off into rough moulds. The loss of sulphur by this method is very considerable, usually not more than two-thirds of the total amount contained in the rock being obtained. (3.) Sulphur may be obtained by heating certain metallic sul- phides ; thus when iron pyrites is heated it yields one-third of its sulphur — 3FeS 2 = Fe 3 S 4 J r S 2 . If the pyrites be roasted in kilns, the whole of the sulphur is obtained, partly as free sulphur, and partly as sulphur dioxide, thus — 3FeS 2 + 50 2 = Fe 3 0 4 + 3S0 2 + 3S. This method was at one time rather extensively employed for the preparation of sulphur on a manufacturing scale, but has now practically gone out of use, the pyrites being usually roasted with excess of air, whereby the whole of the sulphur is converted into sulphur dioxide for use in the manufacture of sulphuric acid. By a similar process, sulphur is obtained as a bye-product during the roasting of copper pyrites, in the first stage of the operation of copper-smelting. (4.) Large quantities of sulphur are now extracted from the vat- waste, or alkali-waste , obtained in the manufacture of sodium carbonate by the Leblanc process. This material consists largely of an insoluble oxy-sulphide of lime, a compound containing calcium sulphide (CaS) and calcium oxide (CaO) in varying proportions. Either in the lixiviating tanks themselves, or in special vats, a current of air is blown through the compound, whereby the calcium sulphide it contains is ultimately converted into a mixture of calcium hydrosulphide (CaH 2 S 2 ), thiosulphate (CaS 2 0 3 ), and polysulphide (CaS 6 ), according to the following equations— (1.) 2CaS + 2H 2 0 = CaH 2 S 2 + CaH 2 O a . 3^2 Inorganic Chemistry This reaction goes forward in several stages, in the course of which a quantity of sulphur is set free ; this is then acted upon by the calcium hydroxide, with the formation of calcium polysulphide and calcium thiosulphate, thus— ( 2 .) 3CaH 2 0 2 + 12S = 2CaS 6 + CaS 2 0 3 + 3H 2 0. The materials are alternately oxidised and lixiviated several times, and the liquor is then treated with excess of hydrochloric acid, at a temperature of about 6 o°, which decomposes the various sulphur compounds according to the following equations — (a.) CaH 2 S 2 + 2HC1 = CaCl 2 + 2H 2 S. (b.) CaS 6 + 2HC1 = CaCl 2 + H 2 S + 4S. ( c . .) CaS 2 0 3 + 2HC1 = CaCl 2 + S0 2 + S + H 2 0 The best results are obtained, when the sulphur compounds are present in such proportions that the S0 2 evolved by reaction c is sufficient to decompose the whole of the SH 2 produced by the other two reactions, so that neither gas escapes— S0 2 + 2H 2 S = 2H 2 0 + 3S. ( 5 .) Sulphur is also obtained from the spent oxide of iron which has been used in the “ purifiers ” employed upon gas-works. Coal gas contains considerable quantities of sulphuretted hydrogen, which are removed from the gas by passing it through hydrated ferric oxide (Fe 2 H 6 0 6 ), which absorbs the whole of the sulphuretted hydrogen, thus — Fe 2 H 6 0 6 + 3H 2 S = 2FeS + S + 6H 2 0. When the compound has lost its power to absorb sulphuretted hydrogen, the material is thrown out of the purifiers and exposed to air and moisture, when the iron becomes reconverted into the hydrated oxide, and the sulphur is set free— 2FeS + 30 + 3H 2 0 = Fe 2 H 6 0 6 + 2S. This revivified material is then employed for the purification of a further quantity of gas. It is found that after a certain number of revivifying operations, the substance begins to lose its power of absorbing any additional sulphuretted hydrogen, and as it then Sulphur 363 contains nearly half its weight of sulphur .t .becomes a valuable source of this element. The sulphur is obtained from ■t by distd- lation, or the material may be roasted ,n special kilns whereby the sulphur is converted into sulphur dioxide, and employed for the manufacture of sulphuric acid. , Purification.— The crude sulphur obtained by the forego g methods U purified by distillation, the process being earned out m the arrangement shown in Fig. 105. The sulphur is rs me in an iron pot, d, and the liquid substance drawn off at will y Fig. 105. means of the pipe F into the retort B. The sulphur is there boiled by means of the fire, and the vapour allowed to issue into the large brickwork chamber G. As the vapour enters the chamber U condenses upon the walls and floor in the form of a light powdery deposit, consisting of minute crystals, and constituting *e>^ of sulphur of commerce. As the process continues, and the brick- work becomes hot, this soft powder melts and collects upon the floor as an amber-coloured liquid, which is run out from time to 364 I tiorganic Chemistry form it is known as roll sulphur iatteT f^a^ the Chamber, therefore, is famished with a valve, S, at the , 0 „’ be'reheved. 6 PreSSUrC deVe ’° ped * ‘ he — ' of combustion may Properties. —Sulphur, as ordinarily seen, is a pale-yellow brittle crystalline solid It ,s insoluble in water, but readily dissolves hi carbon disulphide, and to a greater or less degree in turpentine enzene, chloroform, sulphur chloride, and many other solvents’ t is a non-conductor of electricity, and an extremely bad con- ductor of heat. A piece of sulphur on being very gently warmed, Fig. 106. ulen b> I ’l emg graspe f in the hand > may be heard to crack by the unequal heating, and will ultimately fall to pieces. At a tem- perature of ,14.5 sulphur melts to a clear amber-coloured and liauid'its^rnl rn ° ' ’ ° n raiSmg the temperature of this mohnit l , raP,<3 y aarkens ’ and at the same time it loses its ahnost black at a . temperature of about 230" the mass appears most Mack, ^ nd f 50 V1SC0US that il can no Icmger be poured subTtanr e e Ve h S ? $ ! em P erature is still further raised, the ahhou^h b H retaining its dark colour, again becomes liquid, although ,t does not regain its original limpidity. At 448° the vaoour On 3 ^ “• COn ^ e « edinto a yellowish-brown coloured vapour. On allowing the boiling sulphur to cool, it passes through the same changes in reverse order until it solidifies. When the vapour of sulphur is heated ,0 ,000", it is converted Sulphur 365 into a true gas, and has a density of 32, one lifre of the gas weigh- ing 32 criths. Sulphur is known to exist in four allotropic modifications, two of which are crystalline and two amorphous. (a) Rhombic Sulphur . — Of the two crystalline varieties this is the more stable. Sulphur, therefore, that occurs native is found crystallised in this form. It may be obtained by allowing a solu- tion of sulphur in carbon disulphide to slowly evaporate. Fig. 106 represents two large rhombic octahedral crystals of sulphur obtained in this way. Octahedral crystals of sulphur can also be obtained under certain conditions, when melted sulphur is allowed to crystallise. Sulphur Fig. 107. in the liquid condition exhibits the phenomenon of suspended solidification to a very high degree, and if the liquid be carefully cooled out of contact with air, the temperature may fall to 90° before solidification takes place. If into the liquid in this state a crystal of the rhombic variety be dropped, the sulphur begins to solidify in crystals of that form. If the superfused sulphur be con- tained in a hermetically closed flask, the liquid frequently deposits octahedral crystals, and by allowing the mass to partially solidify, and quickly inverting the flask, the crystals may be seen upon the bottom of the vessel. The specific gravity of this form of sulphur is 2.05. (/ 3 ) “ Prismatic ” Sulphur. — When melted sulphur is allowed 3^6 Inorganic Chemistry to cool under ordinary conditions, such as in a crucible or beaker, it crystallises in the form of prismatic needles, belong- ing to the monoclinic, or monosymmetric system. By allowing the mass to partially solidify, and pouring off the still liquid por- tion, these crystals will be seen lining the inside of the beaker as long translucent prisms. Fig. 107 shows such a mass cf crystals. Prismatic crystals of sulphur are also obtained, when this element is crystallised from a hot solution in oil of tur- pentine. The specific gravity of this form of sulphur is less than that of the octahedral variety, being 1.98. At ordinary temperatures this modification is unstable, and in the course of a day or two the crystals lose their translucent appearance, owing to their becoming broken down into a number of smaller crystals of the rhombic variety, and present the opaque yellow appearance of ordinary roll sulphur. This change from the prismatic to the octahedral variety, which takes place more quickly when the crystals are scratched, or subjected to vibration is attended with evolution of heat. When monoclinic sulphur is thrown into carbon disulphide, its transformation into the stable modification takes place rapidly, and in this way, by means of a thermopile, the heat evolved by the change may be rendered evident. As carbon disulphide, however, at once exerts its solvent action upon the rhombic sulphur the moment it is formed, the reduction of temperature resulting from this cause, would com- pletely over-balance and mask the more feeble heat effect produced by the passage of the sulphur from the unstable to the stable form. In order, therefore, to render evident the heat resulting from the change of crystalline form, the carbon disulphide must be pre- viously allowed to dissolve as much sulphur as it can take up. If a small quantity of carbon disulphide, so saturated with sulphur, be placed in a corked flask, and stood upon the face of a thermo- electric pile * in connection with a galvanometer, and a quantity of prismatic crystals of sulphur be quickly thrown into the liquid, a sensible deflection of the galvanometer needle will be seen in the direction caused by heat. Although under ordinary conditions, monoclinic sulphur is un- stable and passes into rhombic form, at temperatures between The thermo-electric pile is a delicate physical instrument employed for registering shght changes of temperature ; for descriptions of the apparatus the student must consult text-books on physics. Sulphur 367 100* and 1 14°, it appears to be the more stable variety, for at this temperature rhombic sulphur passes into the monoclinic variety. (y) Plastic Sulphur . — When sulphur which has been heated until it reaches the viscous condition, is suddenly plunged into water, or when boiling sulphur is poured in a thin stream into water, the substance solidifies to a tough elastic material some- what resembling indiarubber. The sulphur in this form is known as plastic sulphur. This variety is best obtained by distilling a quantity of ordinary sulphur from a glass retort (Fig. 108), and allowing the distilled liquid to flow in a fine stream into cold water placed for its reception. As the liquid sulphur falls into the water, it congeals to the plastic condition as a continuous thread, which winds itself in a regular manner into beauti- ful coils of a delicate trans- lucent amber colour. The specific gravity of plastic sul- phur is 1.95, and it is not soluble in carbon disulphide. At ordinary temperatures this allotrope of sulphur is gra- dually transformed into the stable rhombic variety ; in the course of a few days it loses its transparency and elasticity, and becomes con- verted into the ordinary lemon yellow brittle condition of common sulphur. This change takes place more quickly if the plastic material be stretched and worked between the fingers, and still more readily by heating it for a few moments to ioo°, and allowing it again to cool. (S) White Amorphous Sulphur.— When sulphur is heated, and the vapour condensed upon a cool surface, as in the formation of ordinary flowers of sulphur, although the greater portion of the sulphur is sublimed in the rhombic octahedral form, the sublimate contains a small amount of sulphur in the form of an amorphous powder, which is almost milk-white in colour. This modification is best obtained by treating flowers of sulphur, which usually contains as much as 5 or 6 per cent, of amorphous sulphur, with carbon disulphide, whereby the rhombic variety is dissolved, and the white amorphous substance, which is insoluble 3^8 Inorganic Chemistry in that liquid, is left behind. By filtering the liquid and washing the residue with carbon disulphide until the whole of the soluble sulphur is removed, the amorphous powder may be obtained in a state of purity. This amorphous substance is also produced in small quantity, by the action of light upon a solution of sulphur in carbon disulphide. Thus, if a perfectly clear solution of sulphur in this liquid be placed for even a few minutes in the path of a beam of electric light, the solution will be seen to become rapidly turbid, owing to the forma- tion of this insoluble modification. This milk-white amorphous modification is stable at the ordinary temperature, and therefore does not pass spontaneously into the rhombic variety. When heated to a temperature of ioo°, it quickly becomes yellow in colour, and is then readily soluble in carbon disulphide, having been transformed at that temperature into the ordinary stable form. Milk Of Sulphur.— This substance is a medicinal preparation, obtained by precipitating sulphur from a polysulphide of lime by means of hydrochloric acid. Flowers of sulphur and milk of lime are boiled together for some time, and after settling, the clea? reddish liquid containing the calcium polysulphides is decanted off; and hydrochloric acid added to it ; calcium chloride is formed^ and sulphur in a fine state of subdivision is precipitated — CaS* + 2HC1 = CaCl 2 + H 2 S + (x- 1)S. The product so obtained is pale yellow in colour, and consists of ordinary sulphur often contaminated with considerable quantities of calcium sulphate, derived from sulphuric acid present in the hydrochloric acid employed in the precipitation. When sulphur in any of its modifications is heated in the air, it takes fire and burns with a pale blue flame, giving rise to sulphur dioxide ; when burnt in oxygen a small quantity of sulphur tri- oxide is at the same time produced. Finely divided sulphur, when exposed to air and moisture, under- goes slow oxidation even at ordinary temperatures, with the forma- tion of sulphuric acid. Thus, if flowers of sulphur be moistened with water and freely exposed to the air, in a short time the water will be distinctly acid. On this account, sulphur that is used for pyrotechnic purposes, is thoroughly washed and dried, and pre- served in warm dry places. Hydrogen Sulphide 369 Sulphur combines directly with many metals under the influence of heat, forming sulphides ; the union in many cases being accom- panied by vivid combustion. Thus, a strip of copper, when intro- duced into sulphur vapour, burns brilliantly with the formation of copper sulphide ; and a red-hot bar of iron, when pressed against a roll of sulphur, bums in the vapour which is generated, and the molten sulphide falls in scintillating masses through the air — Fe + S = FeS. Heated with sodium or potassium, the alkaline sulphides are formed with deflagration — K 2 + S = K 2 S. COMPOUNDS OF SULPHUR WITH HYDROGEN. Two compounds of these elements are known, namely — Hydrogen sulphide or sulphuretted hydrogen . H 2 S. Hydrogen persulphide H 2 S 2 . HYDROGEN SULPHIDE. Formula, H 2 S. Molecular weight = 34. Density = 17. Occurrence. — This gas is evolved in volcanic regions, and is met with in solution in sulphur mineral waters. Modes Of Formation. — ( 1 .) Sulphuretted hydrogen may be formed by the direct union of its elements, by passing a mixture of hydrogen and the vapour of sulphur through a strongly heated tube. In small quantity it is produced when hydrogen is passed into boiling sulphur, or over certain heated metallic sulphides. ( 2 .) Sulphuretted hydrogen is most readily obtained by the action of either hydrochloric or sulphuric acid upon ferrous sul- phide, thus — FeS + 2HC1 = FeCl 2 + H 2 S. FeS + H 2 S0 4 = FeS0 4 + H 2 S. The ferrous sulphide in broken fragments is placed in a two- necked bottle, similar to the apparatus, Fig. 27 , employed for the preparation of hydrogen, and the dilute acid poured upon it. The gas is rapidly evolved without the application of heat. The gas obtained by this method always contains free hydrogen, owing to the presence of uncombined iron in the ferrous sulphide. 2 A 37° Inorganic Chemistry (3.) Pure sulphuretted hydrogen may be obtained by heating antimony trisulphide (grey antimony ore) with strong hydrochloric acid, when antimony trichloride is produced and hydrogen sulphide evolved — Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. (4.) Sulphuretted hydrogen is produced during the putrefaction of organic substances containing sulphur, the offensive smell of a decomposing egg being due to the presence of this gas. It is also produced during the destructive distillation of coal, by the direct union of hydrogen with the sulphur contained in the pyrites, hence coal gas always contains sulphuretted hydrogen amongst its impurities. Properties. — Sulphuretted hydrogen is a colourless gas having a somewhat sickly sweetish taste, and an extremely offensive odour. It acts as a powerful poison when inhaled in the pure state, and even when very largely diluted with air it gives rise to dizziness and headache. Its poisonous effects are more marked upon some animals than others : thus, a bird was found to die in an atmosphere containing only iVtttj of sulphuretted hydrogen, while it required an amount equal to to poison a hare ; and again, cold-blooded animals are in no way affected by inhaling these proportions of the gas. Sulphuretted hydrogen is moderately soluble in water ; at ordinary temperatures water dissolves about three times its own volume of the gas. In collecting it over water, therefore, consider- able loss results unless the water be warm. The coefficient of absorption by water at o° is 4.3706. The aqueous solution gives an acid reaction with litmus, and possesses the taste and smell of the gas. It quickly decomposes on exposure to air, the hydrogen of the sulphuretted hydrogen combines with oxygen, and the liquid becomes turbid by the preci- pitation of sulphur. Hydrogen sulphide is an inflammable gas, burning with a bluish flame, and producing sulphur dioxide and water — 2H 2 S + 30 2 = 2S0 2 + 2H 2 0. If mixed with oxygen in the proportion demanded by this equa- tion, viz., two volumes of sulphuretted hydrogen and three volumes of oxygen, and ignited, the mixture explodes with violence. When the gas is burned with an insufficient supply of air or oxygen for its complete combustion, the sulphur is deposited. Sulphuretted hydrogen is decomposed by the halogens, with the Hydrogen Sulphide 37 1 deposition of sulphur, and the formation of the hydrogen compound of the halogen element, thus — H 2 S + Cl 2 = 2HC1 + S. Fluorine, chlorine, and bromine are capable of bringing about this decomposition at ordinary temperatures ; in the case of iodine, the reaction is attended with absorption of heat, which may be supplied by passing the mixture of iodine vapour and sulphuretted hydrogen through a hot tube, «or by causing the action to take place in the presence of water. In the latter case the heat of solu- tion of the hydriodic acid, determines the reaction. When passed into sulphuric acid, reduction of the acid takes place with the precipitation of sulphur — H 2 S0 4 4- H 2 S = S0 2 + 2H 2 0 + S. Sulphuretted hydrogen, therefore, cannot be dried by means of sulphuric acid. The gas acts upon many metals with the formation of sulphides ; thus, when potassium is heated in a stream of hydrogen sulphide it readily bums and produces potassium hydro-sulphide, H 2 S + K = KHS + H. Such metals as tin, lead, silver, &c., are rapidly tarnished in contact with this gas. On this account articles of silver, when exposed to the air of towns, quickly become covered with a film of sulphide, which first appears yellowish-brown, and gradually becomes black. The discoloration of a silver spoon, when introduced into an egg which is partially decomposed, is due to the same cause. Sulphuretted hydrogen also acts upon metallic salts, combining with the metal to form a sulphide. The “ white-lead ” employed in ordinary paint is gradually blackened on prolonged exposure to the air by the formation of lead sulphide. Hydrogen sulphide is rapidly absorbed by lime, with the forma- tion of calcium hydrosulphide — CaH 2 0 2 + 2H 2 S = CaH 2 S 2 + 2H 2 0. It is also absorbed by calcium sulphide, yielding the same compound. This reaction is employed in the method known as Chances process, for utilising the sulphur of the vat waste of the alkali manufacture. This consists in passing lime kiln gases through a series of vessels containing the waste mixed with water. In the first vessels the carbon dioxide is absorbed, and sulphuretted 372 Inorganic Chemistry hydrogen evolved. This, passing into the later vessels, is absorbed by the vat waste, forming calcium hydrosulphide, which in its turn is decomposed by carbon dioxide, with the evolution of twice the volume of sulphuretted hydrogen for a given volume of carbon dioxide, as in the first reaction — (1) CaS + C0 2 + H 2 0 = CaC0 3 + H 2 S. ( 2 ) CaS + H 2 S = CaH 2 S 2 . ( 3 ) CaH 2 S 2 + C0 2 + H 2 0 = CaC0 3 + 2H 2 S. The sulphuretted hydrogen, mixed with atmospheric nitrogen and a small quantity of carbon dioxide, is sufficiently rich to burn, yielding sulphur dioxide, which can then be employed for the manufacture of sulphuric acid. Sulphuretted hydrogen is also decomposed by ferric hydroxide, with the formation of ferrous sulphide and water, and the deposi- tion of sulphur, as described on page 362 . This action takes place with the evolution of considerable heat, the temperature rising high enough to ignite a mixture of sulphuretted hydrogen and oxygen.* Sulphuretted hydrogen is a valuable laboratory reagent, on account of the general behaviour of certain classes of sulphides. Thus, the sulphides of certain metals, being insoluble in dilute acids, are precipitated from acid solutions ; for example — CuS0 4 + H 2 S = CuS + h 2 so 4 . CdCl 2 + H 2 S = CdS + 2HC1. Others are soluble in acids, but insoluble in alkaline liquids, and are therefore precipitated by sulphuretted hydrogen in the presence of ammonia, or by the addition of ammonium sulphide, thus — ZnS0 4 + (NH 4 ) 2 S = ZnS + (NH^SO* A third group of metals yield sulphides that are soluble in water, and therefore are not separated either in acid or alkaline solutions. Many of the metallic sulphides are also possessed of charac- teristic colours, which readily serve for their identification. Thus, arsenious sulphide is pale yellow, and cadmium sulphide golden yellow. Antimonious sulphide has a bright red colour, while zinc sulphide is white. This behaviour of metals towards sulphuretted hydrogen is the basis upon which certain methods of qualitative analyses are founded. * “Chemical Lecture Experiments,” new ed., Nos. 577, 578. Hydrogen Persulphide 373 HYDROGEN PERSULPHIDE. Formula, H 2 S 2 . Modes of Formation.— (i.) This substance, which stands in the same relation to hydrogen sulphide, as hydrogen peroxide does to water, may be obtained by slowly pouring a solution of calcium disulphide into hydrochloric acid, in the cold ; the liquids being rapidly stirred during the process of mixing, and the acid being kept in considerable excess — CaS 2 + 2HC1 = CaCl 2 + H 2 S 2 . The hydrogen persulphide separates out as a heavy, pale yellow, oily compound, which settles to the bottom of the liquid. The calcium disulphide is prepared by boiling together one part of lime with about twenty parts of water, and one part of flowers of sulphur. The yellow liquid that is obtained will contain more or less of the higher sulphides of lime, and in proportion as these are present there will be a precipitation of sulphur with the hydrogen persulphide, thus — CaS 6 + 2HC1 = CaCl 2 4* H 2 S 2 + 3S. (2.) Hydrogen persulphide may also be obtained from a com- pound that is produced by the action of sulphuretted hydrogen upon strychnine, in the presence of oxygen— 2C 21 H 22 N 2 0 2 + 6SH 2 + 30 = 3H 2 0 + 2C 21 H 22 N 2 0 2 ,3H 2 S 2 . This substance, on treatment with an acid, yields the persulphide. Properties. — Hydrogen persulphide is an oily liquid having a specific gravity of 1.73. It has a pungent smell, accompanied by the odour of sulphuretted hydrogen, due probably to the partial decomposition of the compound, and its vapour is irritating to the eyes. It is an unstable substance, decomposing at ordinary tem- peratures into sulphur and sulphuretted hydrogen : when heated, this decomposition takes place rapidly. It readily dissolves sul- phur, and on this account, and the readiness with which the compound decomposes, it is extremely difficult to obtain it in a state of purity, and so to determine its exact composition. It is insoluble in water, but dissolves readily in carbon disulphide and 374 Inorganic Chemistry ether: its solution in the former liquid is more stable than the liquid substance itself. Hydrogen persulphide burns with a blue flame, yielding sulphur dioxide and water. v Hydrogen persulphide possesses feeble bleaching properties, and like its oxygen analogue it is decomposed by certain metallic oxides, with the evolution of sulphuretted hydrogen. COMPOUNDS OF SULPHUR WITH CHLORINE. Two of these compounds exist at ordinary temperatures, while a third is only known at temperatures below - 22 0 . 1. Disulphur dichloride or sulphothionyl chloride S 2 C1 2 . 2. Sulphur dichloride SCI 3. Sulphur tetrachloride SC 1 4 . Disulphur Diehloride, S 2 Cl 2 .-This substance is obtained by passing dry chlorine over the surface of heated sulphur, con- tained in a retort ; the compound, which distils away as fast as it is formed, condenses in the receiver as a yellow liquid- s' Cl 2 = S 2 C 1 2 . Properties.— The redistilled liquid is an amber-coloured fuming substance with a disagreeable penetrating odour, the vapour of which irritates the eyes. Its specific gravity is 1.709, and it boils at 138. i°. In contact with water it gradually decomposes into hydrochloric acid and sulphur dioxide, with the precipitation of sulphur. The action takes place in two stages, thiosulphuric acid being formed as an intermediate product, thus — (a.) 2 S 2 C 1 2 4 - 3 H 2 0 = 4 HC 1 + S 2 + H 2 S 2 0 3 . OO H 2 S 2 0 3 = H 2 S 0 3 + S. Disulphur dichloride dissolves sulphur with great readiness, and the solution so obtained is largely employed in the process of vulcanising indiarubber. This compound is the most stable of the three chlorides of sulphur. From the fact that it contains chlorine and sulphur in 375 0 Aides and Oxy acids of Sulphur the. proportion ot one atom of each element, it is sometimes called sulphur monochloride ; but as its vapour density (67.4) shows that it contains two atoms of each element in the molecule, the use of the word monochloride is calculated to mislead. The name sul- phothionyl chloride indicates its analogy to thionyl chloride, SOCl 2 , from which it may be regarded as being derived, by the replace- ment of the oxygen by an atom of sulphur, O ./Cl. \ Cl s = s /Cl Cl Sulphur Dichloride, SC1 2 .— This compound is obtained by passing a stream of dry chlorine into disulphur dichloride at a temperature not above o°. When the maximum amount of chlorine is absorbed, the liquid assumes a dark reddish-brown colour. Excess of chlorine is removed by passing a stream of carbon dioxide through the liquid. Sulphur dichloride rapidly dissociates with rise of temperature into free chlorine and disulphur dichloride; at +20° this decomposition amounts to 6.5 per cent., at 50° 24.59 per cent., and at ioo° 80.85 per cent. On boiling the compound, therefore, chlorine is evolved, and the disulphur dichloride remains behind. In contact with water it is decomposed in the same manner as the more stable compound. Sulphur Tetrachloride, SC1 4 .— This compound only exists at temperatures below - 22 0 , and is produced by saturating sulphur dichloride with chlorine at that temperature. It dissociates very rapidly as the temperature rises ; thus, at 7 0 above the temperature at which it is formed, viz., at - 15 , this decom- position amounts to 58.05 per cent. At -2 0 88.07 per cent, of the compound dissociates, while at +6.2 the percentage rises to 97.57. . The compound is decomposed by water with violence into sulphur dioxide and hydrochloric acid — SC1 4 + 2H a O = S0 2 + 4HC1. Compounds of Sulphur with Bromine and Iodine have been obtained, corresponding to S 2 C1 2 . S 2 Br 2 as a red-coloured liquid, boiling with partial decomposition at 200° ; and S 2 I 2 as a dark-grey crystalline solid, which melts at a temperature about 6o°. OXIDES AND OXYACIDS OF SULPHUR . Four oxides of sulphur are known, namely — (1.) Sulphur dioxide (sulphurous anhydride) . S 0 2 . (2.) Sulphur trioxide (sulphuric anhydride) . S 0 3 . (3.) Sulphur sesquioxide . . . - S 2 0 3 . (4.) Persulphuric anhydride .... S 2 0 7 . Three of these oxides, namely, Nos. 1, 2, 4, give rise respectively 376 Inorganic Chemistry to the three acids, sulphurous, sulphuric, and persulphuric, besides which several others are known— Hyposulphurous acid Sulphurous acid Sulphuric acid . Persulphuric acid Thiosulphuric acid t . Pyrosulphuric acid sulphuric acid) . h 2 so 2 h 2 so 3 h 2 so 4 hso 4 * H 2 S 2 03 (Nordhausen (uc „ j- rl 2 i> 2 U 7 H0 ? s HO) 5 * HO Uo Ho; ba HO? o n HO HO (so O ] bU 2 - H °[ SO Hs | sv* ho.so 2 ho.so 2 o. Besides these acids, there is a series known under the general name of the polythionic acids. They may be regarded as being derived from dithionic acid, which is the first of the series, by the absorption into the molecule of various quantities of sulphur. Four of these acids are believed to exist, viz. : — I Dithionic acid (sometimes hyposulphuric acid) . c a Ued } h 2 s 2 0 6 ho.so 2 ho.so 2 Trithionic acid . h 2 s 3 o 6 ho.so 2 ho.so 2 Tetrathionic acid ho.so 2 ho.so 2 Pentathionic acid h 2 s 6 0 6 ho.so 2 ho.so 2 SULPHUR DIOXIDE. S. S* | ^3. Formula, S 0 2 . Molecular weight = 63.92. Density = 31.96. Occurrence. —This compound is met with in the gaseous emanations from volcanoes, and in solution in certain volcanic springs. It is also present in the air of towns, being derived from the combustion of the sulphur compounds present in coal. Modes of Formation.— ( 1.) Sulphur dioxide is formed when sulphur burns in air or oxygen — S + o 2 = SCX. * By some chemists— H 2 S 2 0 8 , jjO SO 2 } ° 2 ‘ f This acid is sometimes incorrectly called hyposulphurous acid , its sodium salt being known as sodiumhyposulphite: the so-called ‘ ‘ hypo ” of the photographers. 377 Sulphur Dioxide At the same time small quantities of sulphur trioxide are formed, which render the gas obtained by this combustion more or less foggy. (2.) Sulphur dioxide may also be obtained by heating sulphur with a metallic peroxide, such as manganese dioxide, thus — S 2 + Mn0 2 = S0 2 + MnS. (3.) It is obtained when such metallic sulphides as copper pyrites or iron pyrites are roasted in a current of air, the metal being con- verted into oxide, thus — 2FeS 2 + 110 = Fe 2 0 3 + 4S0 2 . (4.) The most convenient laboratory process for preparing sul- phur dioxide consists in heating sulphuric acid with copper, the final products of the reaction being copper sulphate, water, and sulphur dioxide — Cu + 2H 2 S0 4 = CuS0 4 + S0 2 + 2H 2 0. The metals mercury or silver may be substituted for copper, but in practice the latter metal is usually employed. (5.) Sulphur dioxide is also formed when sulphuric acid is heated with sulphur, the oxidation of the sulphur and the reduction of the sulphuric acid going on simultaneously — S + 2H 2 S0 4 = 2H 2 0 + 3S0 2 . (6.) The reduction of sulphuric acid may be brought about by means of carbon ; thus, if sulphuric acid be heated with carbon, the latter is oxidised to carbon dioxide, and the acid is reduced to sulphur dioxide — C + 2H 2 S0 4 = 2H 2 0 + 2S0 2 + C0 2 . This method is employed on a large scale for the preparation of alkaline sulphites. The carbon dioxide which accompanies the sul- phur dioxide, not being soluble to any extent in water containing sulphurous acid, is not in any way detrimental. (7.) Sulphur dioxide is formed by the decomposition of a sulphite by dilute sulphuric acid, thus — Na 2 SO s + H 2 S0 4 = Na 2 S0 4 + H a O -I- SO a . 378 / norganic Chemistry Properties. — Sulphur dioxide is a colourless gas, having the well-known suffocating smell usually associated with burning sulphur. The gas will not burn in the air, nor will it support the combustion of ordinary combustibles : a taper introduced into the gas is instantly extinguished. Sulphur dioxide is more than twice as heavy as air, its specific gravity being 2.21 1 (air=i). On this account it is readily collected by displacement : it cannot be collected over water on account of its solubility in that liquid, but may be collected over mercury. The solubility of sulphur dioxide in water at various temperatures is seen by the following figures 1 vol. of water at o° dissolves 79.789 vols. SOg. » » 2 °° „ 39-374 „ >j » 4° >» 18.766 „ The solution is strongly acid, and is regarded as sulphurous acid, the gas having entered into chemical union with the water— so 2 + H 2 0 = h 2 so 3 . On cooling a saturated solution of sulphur dioxide to o°, a solid crystalline hydrate is deposited, having the composition H 2 S0 3 , 8H 2 0* When the solution is boiled, the whole of the sulphur dioxide is expelled. Sulphur dioxide is an easily liquefied gas. At o° a pressure of 1.53 atmospheres is sufficient to condense it, while at ordinary pressures it may be liquefied by a cold of - io°. Its critical temperature is 155.4 0 . To obtain liquid sulphur dioxide, the gas, as evolved from the action of sulphuric acid upon copper, is dried by being passed through a bottle containing sulphuric acid, and is then passed through a gas-condensing tube, Fig. 109, immersed in a freez- ing mixture. The gas at once condenses in the bulb of the apparatus, as a colourless, transparent, mobile liquid, which boils at - 8°. When the liquid is cooled to - 76° it solidifies to a trans- parent, ice-like mass. Liquid sulphur dioxide is largely employed as a refrigerating agent, low temperatures being obtained by its rapid evaporation * Several cryohydrates of sulphurous acid have been obtained, H 2 S0 3 ,6H 2 0 H 2 SO 3 ,10H 2 O ; H 2 S0 3 ,14H 2 0. 379 Sulphur Dioxide under reduced pressure. The liquid dissolves phosphorus, iodine, sulphur, and many resins. When thrown upon water, a portion of the liquid dissolves, and, owing to the reduction of temperature caused by the rapid evaporation of the remainder, a quantity of the water is frozen. The ice so produced contains a large pro- portion of the solid hydrate H 2 S0 3 ,8H 2 0. Although sulphur dioxide is incapable of supporting the com- bustion of ordinary combustibles, many metals will take fire and burn when heated in the gas. Thus, when finely divided iron is heated in a stream of sulphur dioxide it burns, forming sulphide and oxide of the metal. It also unites with many metallic peroxides, and often with so much energy as to give rise to light and heat. Thus, when passed over peroxide of lead, the mass glows spontaneously in the gas, and lead sulphate is produced — Pb0 2 + S0 2 = PbS0 4 . Sulphur dioxide is decomposed by the influence of strong light 380 Inorganic Chemistry If a concentrated beam of electric light be passed through a vessel filled with gaseous sulphur dioxide, the gas at first will appear perfectly transparent and clear ; but in the course of a few minutes the track of the beam will become more and more visible as it traverses the gas, owing to the formation of thin clouds of sulphur trioxide and sulphur, until the atmosphere of the vessel appears to be filled with fog (Fig. iio) — 3S0 2 = 2S0 3 + S. After the lapse of a short time, if the vessel be removed from Fig. ho. the strong light, the atmosphere will once more become clear, owing to the reformation of sulphur dioxide. Sulphur dioxide possesses powerful bleaching properties, when in the presence of water. Its bleaching action is due to its absorption of oxygen from water, and consequent liberation of hydrogen, thus — S0 2 4 - 2H 2 0 = H 2 S0 4 + H 2 . The hydrogen so set free reduces the colouring-matter, with the formation of colourless compounds : the action in this case being the reverse to that which takes place with chlorine. In some instances, the bleaching is due to the formation of a colourless compound, by the direct combination of sulphur dioxide with the colouring-matter, as the original colour may often be restored by treatment with dilute sulphuric acid, or by weak alkaline solutions. Thus, by passing sulphur dioxide into an infusion of rose leaves, the red colour of the liquid is quickly discharged, but on the Sulphur Dioxide 381 addition of a small quantity of sulphuric acid the colour is restored. Sulphur dioxide is employed in bleaching materials that would be injured by exposure to chlorine, such as straw, silk, wool, sponge, &c., and the familiar yellow colour which gradually comes over a sponge or a piece of bleached flannel when it is washed with soap, is an illustration of the power of alkalies to restore the original colour to materials that have been bleached by this substance. In the presence of water, sulphur dioxide converts chlorine into hydrochloric acid, and on this account is employed as an “anti- chlor ” — S0 2 + 2H 2 0 + Cl 2 = 2HC1 + H 2 S0 4 . In the same way it acts upon iodine, with the formation of hydriodic acid — S0 2 + 2H 2 0 + I 2 = 2 HI + H 2 S0 4 . In the case of iodine, however, this reaction only takes place when a certain degree of dilution is maintained, for in a more concentrated condition sulphuric acid is reduced by hydriodic acid, into sulphur dioxide, according to the reverse equation, thus— H 2 S0 4 + 2HI = I 2 + 2H 2 0 + S0 2 . It has been shown* that aqueous sulphurous acid can only be completely oxidised by iodine, as indicated in the former equation, when the proportion of sulphur dioxide does not exceed 0.05 per cent. : when the amount exceeds this proportion the second reaction comes into operation. Sulphur dioxide brought into contact with iodic acid, or an iodate, is oxidised into sulphuric acid and liberates iodine, thus — 2HI0 3 + 4H 2 0 + 5S0 2 = 5H 2 S0 4 + I 2 . This reaction is made use of, as a method for the detection of the presence of sulphur dioxide. Paper which has been moistened with a solution of potassium iodate and starch, on exposure to sulphur dioxide is at once turned blue, owing to the liberated iodine combining with the starch. * Bunsen. 3 82 Inorganic Chemistry 1 he composition of sulphur dioxide may be determined by the combustion of sulphur in a measured volume of oxygen, in the apparatus employed for showing the volume composition of carbon dioxide (Fig. 64). After the fragment of sulphur has burnt, and the apparatus has been allowed to cool, it will be seen that there is no alteration in the volume of the contained gas ; the sulphur dioxide produced, occupying the same volume as the oxygen used in its formation. Sulphur dioxide, in other words, contains its own volume of oxygen. One molecule, therefore, of sulphur dioxide contains one molecule of oxygen, weighing 31.92. But the molecular weight of sulphur dioxide is 63.92 ; therefore 63-92 - 31-92 = 32 = the weight of sulphur contained in the molecule of the gas. Sulphur dioxide, therefore, contains in the molecule one atom of sulphur, combined with two atoms of oxygen, hence its composition is expressed by the formula S 0 2 . Sulphurous Acid and Sulphites. -Sulphurous acid is only known in solution and in its cryohydrates. The solution smells of sulphur dioxide, and gradually undergoes decomposition by absorption of oxygen. The acid is dibasic, having two atoms of hydrogen replaceable by metals : it is therefore capable of form- ing two series of salts, according to whether one or both of the hydrogen atoms are replaced. Thus, by its action upon potassium hydroxide, when the acid is in excess the so-called acid potassium sulphite, or hydrogen potassium sulphite, is obtained— KHO + H 2 S 0 3 = H 2 0 + HKSOg. Whereas, if the metallic hydroxide be in excess, the normal potassium sulphite is formed— 2 KHO + H 2 S 0 3 = 2 H 2 0 + K 2 SO s . The alkaline sulphites are readily soluble in water, all other normal sulphites being either difficult of solution or insoluble. SULPHUR TRIOXIDE. Formula, SO s . Molecular weight = 79.88. Vapour density = 39.93. Modes of Formation. — ( 1 .) This compound is produced, when a mixture of sulphur dioxide and oxygen is passed over heated spongy platinum or platinised asbestos— S0 2 + O = so 3 . Sulphur Trioxide 383 On leading the product through a well-cooled receiver, the sulphur trioxide condenses in white silky needles. This method has been successfully employed on a commercial scale. The mixture of sulphur dioxide and oxygen, is obtained by allowing ordinary strong sulphuric acid to drop into earthenware retorts heated to bright redness, whereby it is almost entirely broken up into these two gases and water, thus — H 2 S0 4 = S0 2 + O + h 2 o. The gases are then deprived of the water, by passage first through a condenser, and then through a leaden tower containing coke moistened with sulphuric acid, and are finally passed over heated platinised asbestos contained in glazed earthenware pipes. (2.) Sulphur trioxide is most conveniently obtained by gently heating pyrosulphuric acid in a glass retort. The trioxide distils over and may be collected in a well-cooled receiver — H 2 S 2 0 7 = H 2 S0 4 + S0 3 . (3.) It may also be obtained by heating sodium pyrosulphate to bright redness — Na 2 S 2 0 7 = Na 2 S0 4 + SO s . The sodium pyrosulphate is produced when hydrogen sodium sulphate (so-called bisulphate of soda ) is heated to about 3 0O °’ thus — 2HNaS0 4 = H 2 0 + Na 2 S 2 O r . And on account of this origin it is sometimes termed anhydrous sodium bisulphate. (4.) Sulphur trioxide can also be produced by the action of phosphorus pentoxide upon sulphuric acid. This most powerful dehydrating substance withdraws from the sulphuric acid the elements of water, when gently heated, thus — H 2 S0 4 + P 2 O s = 2HP0 3 + SO3. The trioxide is distilled from the mixture, and the metaphosphoric acid remains in the retort. Properties. — Sulphur trioxide is a white, silky-looking, crystal- line substance, which melts at 14.8° and boils at 46°. It is very volatile, and gives off dense white fumes in contact with air, owing 3^4 Inorganic Chemistry to the combination of its vapour with atmospheric moisture to form sulphuric acid. It combines with water with great energy to form sulphuric acid ; a fragment of the compound dropped into water dissolves with a hissing sound resembling the quenching of red-hot iron — S0 3 + h 2 o = H 2 S 0 4 . When brought in contact with the skin, or other organic matter containing hydrogen and oxygen, it abstracts these elements and produces a burnt or charred effect upon the substance. Sulphur trioxide unites directly with barium oxide, BaO, and if the baryta be dry, the mass becomes incandescent owing to the heat of the union, and barium sulphate is formed — BaO + S0 3 = BaS0 4 . When the vapour of sulphur trioxide is passed through a red-hot tube, it is broken down into sulphur dioxide and oxygen. When the trioxide is heated, it melts to a colourless liquid, which exhibits a remarkably high rate of expansion by heat ; between 25 0 and 45 0 its mean coefficient of expansion is 0.0027, nearly three- fourths of the expansion coefficient of a gas. Sulphur Sesquloxide, S 2 Og. — A solution of this compound in fuming sulphuric acid was obtained early in the century by heating flowers of sulphur with Nordhausen sulphuric acid, whereby a blue solution was obtained. The substance may be prepared by the gradual addition of dry flowers of sulphur to melted sulphur trioxide, at a temperature just above its melting-point, when a malachite-green crystalline solid separates out. The compound is unstable at ordinary temperatures, being resolved into sulphur dioxide and sulphur, the decomposition taking place rapidly upon gently warming — 2S 2 0 3 = S + 3S0 2 . If the sesquioxide be sealed up in a bent glass tube and gently warmed, the sulphur dioxide may be obtained liquid in one limb of the tube. PERSULPHURIC ANHYDRIDE. Formula, S 2 0 7 . This compound is formed when a mixture of dry sulphur dioxide and oxygen is subjected to the silent electric discharge, in an ozone Per sulphuric Acid 385 tube, or by treating sulphur trioxide and oxygen in the same manner. At the end of some hours a small quantity of a viscous liquid collects upon the walls of the glass vessel, which when cooled, solidifies in the form of long transparent needle-shaped crystals, resembling sulphur trioxide in appearance. It is a very unstable substance, and can only be preserved a short time even at low temperatures. It is soluble in water, with the formation of persulphuric acid, but the solution rapidly undergoes decomposition into oxygen and sulphuric acid — S 2 O r + H 2 0 = 2HS0 4 . 2HS0 4 + H 2 0 = 2H 2 S0 4 + O. When very gently warmed, persulphuric anhydride rapidly breaks up into sulphur trioxide and oxygen — S 2 O 7 = 2 SO 3 -f* O. The readiness with which it gives up oxygen, constitutes this com- pound a powerful oxidising agent, and affords the clue to most of its reactions. Persulphuric Acid and Persulphates. — When dilute sulphuric acid is sub- jected to electrolysis (as in the process commonly spoken of as the electrolysis of water), appreciable quantities of persulphuric acid are found in solution at the positive electrode, or anode. The acid itself has never been obtained in a state of purity, its aqueous solution rapidly undergoing decomposition, as already mentioned. In solution this compound displays all the oxidising properties of the oxide. The potassium salt may be obtained by the electrolysis of a saturated solu- tion of hydrogen potassium sulphate in a divided cell, the action being due to the oxidation of the hydrogen potassium sulphate by the nascent oxygen developed at the anode, thus — 2HKS0 4 + O = H a O + 2KS0 4 . The potassium persulphate, being a sparingly soluble salt, crystallises out, and may be freed from the acid sulphate by recrystallisation. The ammonium salt, NH 4 S0 4 , and the barium salt, Ba(S0 4 ) 2 ,4H 2 0, have also been obtained. Barium persulphate is soluble in water, being much more readily dissolved than the potassium salt ; thus, ioo parts of water at o° dissolve 1.77 parts of potassium persulphate and 52.2 parts of the barium salt. On this account solutions of persulphates give no precipitate with barium chloride, whereby they are distinguished from sulphates : if the mixture be warmed, however, the persulphate is decomposed into a sulphate, with evolution of chlorine, thus — 2KS0 4 + BaCl 2 = Ba(S0 4 ) a + 2KC1 = BaS0 4 + K,S0 4 + Cl 3 . 2 B 3^6 Inorganic Chemistry In the solid condition the persulphates are stable salts, but their aqueous solutions gradually evolve oxygen and become converted into sulphates— 2KS0 4 + H 2 0 = 2HKS0 4 + O. The reactions exhibited by solutions of persulphates are those of strong oxidising agents. Hyposulphuroua Acid,* H 2 S0 2 . — This compound is obtained by the reduc- tion of sulphurous acid by means of nascent hydrogen. Thus, when zinc is acted upon by an aqueous solution of sulphurous acid, no hydrogen is evolved, as the nascent gas combines with oxygen of the acid to form water— HjjSOa + H 2 = H 2 0 + H 2 SO s . The solution so obtained, has a yellowish colour, and possesses powerful reducing and bleaching properties. Sodium hyposulphite (HNaS0 2 ) may be obtained by the action of zinc upon a cooled concentrated solution of hydrogen sodium sulphite (HNaS0 3 ), air being carefully excluded, a double sulphite of sodium and zinc being at the wme time produced, thus — 3HNaSO a + Zn = HNaSO a + NaaSO^nSOg + H 2 0. The greater part of the double sodium-zinc sulphite is deposited as crystals, the rest is removed by adding to the mother liquor about four times its volume of alcohol, in a closed flask. The double salt, being less soluble in alcohol than the hyposulphite, is first deposited, and the clear liquid, after being poured off and corked up, is cooled, when it solidifies to a mass of crystals of nearly pure sodium hyposulphite. The crystals in the wet condition are rapidly oxidised on exposure to air ; but if quickly pressed between blotting-paper and dried in vacuo, the dry salt is not acted upon by atmospheric oxygen. The acid is obtained from the sodium salt, by the action upon it of oxalic acid. Sodium hyposulphite is also formed when a solution of hydrogen sodium sulphite is subjected to electrolysis, the nascent hydrogen developed at the negative electrode reducing the sulphite to hyposulphite, by the abstraction of one atom of oxygen. Properties.— Hyposulphurous acid is an extremely unstable, yellow-coloured liquid, which rapidly decomposes, first into thiosulphuric acid and water, and finally into sulphur dioxide, water, and sulphur — 2H 2 S0 2 = H 2 S 2 0 3 + H 2 0. H 2 S 2 O 3 — S 0 2 + h 2 o + s. The acid reduces salts of silver or mercury, with precipitation of the metal, thus — HgCla + H a O + H 2 S0 2 = Hg + 2HC1 + HaSO,. * This compound (sometimes called hydrosulphurous acid) must not be con- founded with thiosulphuric acid, which is often incorrectly called hypo- sulphurous acid (page 396 ). Sulphuric Acid 387 The sodium salt possesses the same bleaching and reducing powers as the acid ; and when wet, or in solution, it rapidly absorbs oxygen from the air and is converted into hydrogen sodium sulphite — HNaSOa + C = HNaSO a . SULPHURIC ACID. Formula, H 2 S0 4 . Modes of Formation. — (1.) This acid is formed when sulphur trioxide is dissolved in water — so 3 + H 2 0 = h 2 so 4 . (2.) It is also formed by the direct union of sulphur dioxide with hydrogen peroxide — S0 2 + H 2 Oa = H 2 S0 4 . (3.) An aqueous solution of sulphur dioxide gradually absorbs oxygen, and is converted into sulphuric acid — H 2 S0 3 + O = H 2 S0 4 . (4.) Manufacture of Sulphuric Acid.— Sulphur dioxide is un- able to absorb an additional atom of oxygen, and so pass into sulphur trioxide, without the aid of some third substance, which can act as a catalytic agent, or a carrier of oxygen. The material which is employed for this purpose, in the process by which sul- phuric acid is manufactured, is one of the oxides of nitrogen, which is capable of giving up oxygen to the sulphur dioxide, and of again taking up oxygen from the air. Thus, nitrogen peroxide (N0 2 ), by the loss of one atom of oxygen, is reduced to nitric oxide, N O ; which in its turn combines with atmospheric oxygen, and is re- converted into nitrogen peroxide. Therefore, when sulphur dioxide and oxygen are mixed with nitrogen peroxide in the presence of steam, a series of reactions takes place, the final result of which is that the oxygen is caused to combine with the sulphur dioxide and water, with the formation of sulphuric acid — S0 2 + O + H 2 0 = H 2 S0 4 . The nitrogen peroxide at the end of the reaction is unchanged, and is able to react in the same series of changes over and over 388 Inorganic Chemistry again, thus transforming, theoretically, an unlimited, and, in practice, a relatively large quantity of sulphur dioxide into sul- phuric acid. The series of changes that gives rise to the ultimate product is the following : — The sulphur dioxide, nitrogen peroxide, and water give rise, in the first place, to the formation of nitro-sulphonic acid and a molecule of nitric oxide— (i.) 2S0 2 + 3N0 2 4- H 2 0 = 2H(N0)S0 4 + NO. Nitro-sulphonic acid (sometimes called nitro- sulphuric acid \ and nitrosyl sulphate ) may be regarded as sulphuric acid in which one of the hydrogen atoms is replaced by the group (NO), thus, /OH S0 2 <^q^ n q in which case the nitrogen is linked to the sulphur by the intervention of oxygen : or it may be considered as derived from sulphuric acid by the replacement of one of the groups (HO) by the group N0 2 , /OH j \no 2 , when the nitrogen is directly attached to the sulphur. The substance is a white crystalline compound, which in the presence of water is instantly decomposed into sulphuric acid and a mixture of nitric oxide and nitrogen peroxide, thus — (2.) 2S0 2 (HO)(N0 2 ) + H 2 0 = 2H 2 S0 4 4- NO 4- N0 2 . The nitric oxide in this and the former reaction, on coming in contact with the atmospheric oxygen, is at once reconverted into nitrogen peroxide — (3.) NO + O = N0 2 . In the process of the manufacture, the crystalline compound S0 2 (H0)(N0 2 ) (known as chamber crystals) is not actually isolated, unless from accidental causes the supply of water is in deficit, the production of these crystals being regarded as an indication that the process is not being well carried out. The formation of sulphuric acid by these reactions, with the intermediate production of the chamber crystals, may be carried out on a small scale by means of the apparatus shown in Fig. 111. A large flask, F, is fitted with a cork, through which pass five tubes : three of these are connected to separate two-necked bottles containing sulphuric acid, through which can be delivered respec- Sulphuric Acid 389 lively, nitric oxide, sulphur dioxide, and oxygen. The fourth tube is attached to a flask in which water may be boiled, and through which oxygen can be passed, and the fifth tube (not shown in the figure) serves as an exit. A quantity of oxygen is first passed into the large flask through the drying-bottle D, and sufficient nitric oxide is then allowed to enter, to form deep red vapours ; at the same time sulphur dioxide is passed in through the bottle S. In order to introduce a small quantity of moisture, oxygen is allowed to enter through the flask of boiling water, and in a few moments large white crystals begin to form all over the interior Fig. hi. of the flask, and rapidly spread until the whole surface is covered. In order to show the second reaction in the cycle, the gaseous contents of the flask may be swept out by means of a rapid stream of oxygen, passed in through the drying-bottle D ; and when the atmosphere within the apparatus is colourless, a quantity of steam is driven in from the small flask. The chamber crystals will be seen to dissolve with effervescence, and the flask once more becomes filled with brown fumes. The nitric oxide evolved by the decomposition of the nitrosyl sulphate, coming in contact with the oxygen within the flask, at once regenerates nitrogen peroxide, in accordance with equation No. 3. The solution formed in the flask will be found to yield a pre- 39 ° Inorganic Chemistry cipitate of barium sulphate, on the addition to it of a soluble barium salt. On a manufacturing scale, the combination of the reacting gases and vapours which gives rise to the sulphuric acid, takes place in large leaden chambers, usually about ioo feet long, 25 feet wide, and 20 feet high, having therefore a capacity of 50,000 cubic feet ; several of such chambers being placed in series. Into these chambers there is delivered sulphur dioxide, air, oxides of nitrogen, and steam. The plant employed for the manufacture of sulphuric acid con- sists broadly of four parts. 1. Apparatus for generating sulphur dioxide. 2. Apparatus for producing oxides of nitrogen. 3. Appa- ratus for absorbing oxides of nitrogen from the gases leaving the chambers. 4. The chambers in which the reactions are made. (1.) Pyrites Burners . — The sulphur dioxide is obtained either by burning native sulphur, or roasting the “spent oxide” of the gas works (see Sulphur), or by roasting pyrites, the latter being the most general method. The pyrites burner, Fig. 112, B, is essen- tially a small furnace or kiln in which the ore is heated, and in which the admission of air can be duly regulated ; as not only is it necessary to admit sufficient air to completely bum the whole of the sulphur, and so prevent any volatilisation of it in an unburnt condition, but also to supply the requisite volume of oxygen for the requirements of the reactions which are to go on within the cham- ber. Too large a volume of air must be avoided, in order not to unduly dilute the chamber gases. (2.) If no loss of nitrogen peroxide took place during the cycle of changes, the same quantity of this gas would convert an infinite amount of sulphur dioxide and water into sulphuric acid ; but in practice, owing to leakage, defective absorption, and the reduction of a certain percentage of this compound into nitrous oxide, it is necessary to constantly replenish the supply. This is usually done by generating a small quantity of nitric acid (by the action of sulphuric acid upon nitre) in earthenware pots, which are usually placed in an enlarged part of the flue of the pyrites burner, known as the “nitre oven,” and which is provided with a door for the introduction of the pots, Fig. 112, N. The heated gases playing upon these pots, promotes the evolution of the nitric acid, which in contact with sulphur dioxide is at once decomposed according to the equation — 2HNO3 + SOj = H 3 S0 4 + 2NO a . 39i Sulphuric Acid It is found that to make up for the loss of nitrogen peroxide, about three to four parts of nitre are required for every ioo parts of sulphur, burnt as pyrites. (3.) The apparatus for the absorption of the nitrogen peroxide from the gases that are drawn from the chamber at the end of the series, is known as the “Gay-Lussac Tower,” Fig. 112, T. This consists of a square leaden tower filled with fragments of coke, and down which there is caused to slowly percolate, a stream of cold strong sulphuric acid, the acid being evenly spread over the mass of coke by a special distributing arrangement. The nitrogen peroxide is absorbed by the acid, with the formation of nitro- sulphonic acid, S0 2 (H0)(N0 2 ). In order to make use of the absorbed nitroxygen compound, the acid which flows from the Gay-Lussac tower is pumped to the top of another very similar tower, situated between the “burners” and the first of the cham- bers, and known as the “ Glover Tower,” G. The hot gases from the burners, consisting of sulphur dioxide, nitrogen, and oxygen, together with the small quantities of nitrogen peroxide from the nitre pots, are made to pass up this tower on their way to the first chamber, and meeting with the descending stream of nitro-sul- phonic acid as it percolates through the mass of coke with which the tower is filled, denitrification of the latter takes place, thus— 2S0 2 (H0)(N0 2 ) + S0 2 + 2H 2 0 = 2N0 + 3S0 2 (H0)(H0), or 3H 2 S0 4 . The nitric oxide thus evolved, in presence of the atmospheric oxygen, is converted into nitrogen peroxide, and swept along with the other gases into the chambers. In practice, it is usual to deliver down the Glover tower, besides the nitro-sulphonic acid, a quantity of “chamber acid” from a separate tank. The effect of the heated gases upon this dilute acid, is to remove a portion of the water from it, thereby effecting its partial concentration, and furnishing the water demanded by the above equation. It will be seen, therefore, that there is a scrubber tower at each end of the series of chambers, the “ Gay-Lussac ” at the exit, where nitrogen peroxide is absorbed ; and the “Glover” at the commencement, where the dissolved nitrogen compound is again liberated and returned to the chambers. (4.) The chambers are made of sheet lead, connected together 392 Inorganic Chemistry by fusing the edges by means of an oxyhydrogen flame, without the intervention of solder, as the presence of another metal gives rise to the rapid corrosion of the lead on account of galvanic action being set up ; this method of joining the lead is known as autogenous soldering. The enormous leaden chamber is supported in a framework of wood, to which the lead is secured by bands of the same metal, and the whole is usually supported on iron or brick pillars. The general arrangement of a modern sulphuric acid works is seen in Fig. 112. The gases from the double row of pyrites burners B, are led through the Glover tower G, where they effect the denitrification of the nitro-sulphonic acid, as already explained. From this tower the/ are delivered into the series of chambers, where they meet with the necessary supply of steam. The acid collects upon the floor of the chambers, and samples are constantly drawn off by means of an arrangement known as a drip pipe, which, acting in a manner similar to a rain gauge, indi- cates the progress of the processes going on within. The gases after being drawn through the entire series of chambers, by means of the draught caused by the tall chimney, are finally passed up the Gay-Lussac tower T, where all the nitrogen peroxide is absorbed, and returned to the chambers through the intervention of the Glover tower G, as above described. The acid which collects in the chambers, is usually not permitted to reach a higher specific gravity than about 1.6, when it contains about 68 per cent, of sulphuric acid ; for if the strength be allowed to exceed this, the acid not only begins to dissolve the nitrogen peroxide in the chamber, but exerts a corrosive action upon the lead of which the chamber is constructed. It is therefore with- drawn, and the first stage in the further concentration is effected either by the action of the Glover tower, or by evaporation in shallow leaden pans. In order to bring up the strength of the acid to that of “oil of vitriol,” that is, to about 98 per cent., the acid from the Glover tower, or the leaden pans, is heated in either glass or platinum stills. Sulphuric acid, unless specially purified, is liable to contain a number of impurities, such as lead sulphate, derived from the action of the acid upon the chamber ; arsenic, from the pyrites employed ; oxides of nitrogen, and sulphur dioxide. From most of the impurities, except the arsenic, the acid may be purified Sulphuric Acid 393 by the addition of ammonium sulphate, and subsequent redis- tillation — Fig. i 12. B.— Double row of pyrites burners, placed back to back : one being shown open. N. — Hearth where the nitre pots are placed : one shown as open. G- — Glover Tower, with two tanks at top : one for the nitro-sulphuric acid de- rived from the Gay-Lussac tower, the other for the “chamber acid." These acids are forced up from the leaden vessels E, called “eggs.” C. — Leaden chamber, of which there are three shown in the figure. P . — Pipe conveying the gases from the third chamber to the Gay-Lussac tower. T. — Gay-Lussac Tower. The tanks at the top of this and the Glover tower are enclosed in wooden sheds. Arsenic may be removed by boiling the acid with hydrochloric acid or sodium chloride, when it passes away as arsenious chloride. Properties. — Sulphuric acid is a perfectly colourless, heavy, 394 Inorganic Chemistry oily liquid. The acid obtained by distillation, always contains about 2 per cent, of water ; stronger than this it cannot be prepared by evaporation or distillation. When, however, acid of this strength is cooled to o°, colourless crystals of pure sulphuric acid, containing ioo per cent. H 2 S 0 4 , are deposited. The crystals melt at 10.5 0 , and remain liquid at temperatures much below this point. The specific gravity of the pure acid is 1.854 at o°. When boiled, it gives off sulphur trioxide until the amount of water in it rises to 1.5 per cent., when it distils unchanged at a temperature of 338°. Sulphuric acid has a powerful affinity for water, and absorbs moisture from the atmosphere with great readiness. On this account it constitutes one of the most valuable desiccating agents, and is constantly made use of for depriving gases, upon which it exerts no chemical action, of aqueous vapour. Owing to its strong affinity for water, it decomposes many organic substances contain- ing hydrogen and oxygen, withdrawing from the compounds these elements in the proportion to yield water : its action upon formic acid, oxalic acid (see Carbon Monoxide), and alcohol (see Ethy- lene) are examples of this action. When the acid is poured upon such substances as wood or sugar, the elements composing water are withdrawn, and the carbon is liberated, with the result that the compounds are blackened or charred. When sulphuric acid is mixed with water, considerable heat is disengaged, the temperature often rising to the boiling-point of water, and at the same time a diminution in volume takes place. The maximum contraction is obtained upon mixing the materials in the proportion of one molecule of acid to two molecules of water. The diminution in volume in this case amounts to 8 per cent., and the composition of the acid produced, corresponds to the formula H 2 S 0 4 , 2 H 2 0 . Sulphuric acid combines with water in various proportions, form- ing a number of hydrates, and cryohydrates, of a more or less definite character. The best known hydrates are those represented by the formulas H^O^O and HjSO^HjO. These compounds may be regarded as respectively tetrabasic and hexabasic sulphuric acid, and their relation to the ordinary dibasic acid may be expressed by the following formulae — H 2 S 0 4 ... or S0 2 (H0)2. H 4 S 0 6 or H 2 S 0 4 ,H 2 0 „ SO(HO) 4 . H fl S0 6 „ HjSO^HaO,, S(HO) 6 . 395 Pyrosulphuric Acid Salts of each of these acids are known — Hydrogen potassium sulphate . . HKSOa -Derived from H 2 S0 4 . Normal potassium sulphate . . k 2 so 4 Barium sulphate BaS0 4 J Tetrabasic lead sulphate . . . Pb 2 S0 5 „ „ H 4 S0 6 . Hexabasic mercuric sulphate ) (Turpeth mineral) f Hg 3 S0 6 „ „ HeSOe. Most sulphates are soluble in water : those of lead, calcium, and strontium are only very sparingly soluble, whilst barium sulphate is insoluble both in water and acids. The presence of sulphuric acid or a sulphate, may therefore be readily detected by the addition of a soluble barium salt, which causes the immediate precipitation of white barium sulphate, insoluble in hydrochloric acid. PYROSULPHURIC ACID {Nordhausen Acid; Fuming Sulphuric Acid). Formula, H 2 S 2 0 7 or ho!8 2 } °* Modes of Formation. — (i.) This acid may be obtained by dis- solving sulphur trioxide in ordinary sulphuric acid — H 2 S0 4 + SO s = H 2 S 2 0 7 . On cooling the solution to o°, the pyrosulphuric acid separates out in the form of large colourless crystals. ( 2 .) Pyrosulphuric acid is manufactured by the distillation of ferrous sulphate in clay retorts, mounted in series in a large “ galley ” furnace. The first action of heat upon crystallised ferrous sulphate (green vitriol) is to expel six molecules of water of crystal- lisation, leaving the salt of the composition FeSO^H^O. When this substance is further heated it is decomposed finally into ferric oxide, with the formation of sulphur trioxide, water, and sulphur dioxide, thus — 2FeS0 4 ,H 2 0 = Fe 2 0 3 + S0 3 4- S0 2 + 2H 2 0. The decomposition takes place in two stages, the sulphur dioxide and water being evolved in the first part of the process with the formation of ferric sulphate, which is afterwards broken up in the manner shown in the following equation — ( 1 .) 6FeS0 4 ,H 2 0 = Fe 2 (S0 4 ) 3 + 2Fe 2 0 3 + 3SO a + 6H a O. ( 2 .) Fe 2 (S0 4 ) 3 = Fe 2 0 3 + 3SO,. 39^ Inorganic Chemistry The sulphur tnoxide is condensed in receivers, containing either a small quantity of water, or a charge of sulphuric acid. ( 3 ) Pyrosulphuric acid may also be obtained by decomposing sodium pyrosulphate (Na 2 S 2 0 7 ), either by heating it to a high temperature (see Sulphur Trioxide, page 383), or by acting upon it NajSjOj + H 2 S 0 4 = 2 HNaS 0 4 + S 0 3 . The sulphur trioxide obtained, is dissolved in sulphuric acid, as in the former methods ; and the hydrogen sodium sulphate, when gently heated to about 300°, is reconverted into pyrosulphate by the loss of a molecule of water (page 383). Properties. Pyrosulphuric acid is a colourless, strongly fuming liquid, having a specific gravity of 1.88. When cooled, it solidifies to a crystalline mass, which melts at 35 0 . The compound may be regarded as consisting of one molecule of sulphuric acid plus a molecule of sulphur trioxide, H 2 S 0 4 ,S 0 3 ; or, as being derived from two molecules of sulphuric acid, by the withdrawal of one molecule of water, thus — SO. < O-H O-H + H- 0 \ JJ.Q = H 2 0 + SO. so 2 . Pyrosulphuric acid forms a stable series of salts, of which the sodium compound already mentioned is a typical example. These salts are sometimes spoken of as the disulphates , and are analogous to the dichromates ( q.v.\ Two other definite compounds of sulphur trioxide and sulphuric acid are known to exist, both of which are fuming acids. The composition of these substances is expressed by the formulae — H2S0 4 ,3S0 8) or H 2 S 4 0 13 ; and 3H2S0 4 ,S0 3 , or HeS 4 0 16 . THIOSULPHURIC ACID. Formula, HgSgOs. This acid has never been obtained in the free state, as it decom- poses almost as soon as liberated from its salts, into sulphur dioxide and water, with precipitation of sulphur — h,s 2 o 3 = SO a + H 2 0 + S. 397 Thiosulphuric Acid The thiosulphates, however, are stable and important salts, the sodium salt being largely used in photography under the name of hyposulphite of soda, or “hypo.” Modes of Formation of Thiosulphates.— (i.) These salts may be obtained by digesting flowers of sulphur with solutions of the sulphites, thus — Na 2 SO s + S = Na 2 S 2 0 3 . (2.) Sodium thiosulphate is also formed, when sulphur dioxide is passed into a solution of sodium sulphide. The reaction may be regarded as taking place in three steps, in which sodium sulphite and sulphuretted hydrogen are the first products. The latter com- pound is then acted upon by sulphur dioxide with the precipitation of sulphur, thus — S0 2 + H 2 0 + Na 2 S = Na 2 S0 3 + H 2 S. S0 2 + 2H 2 S = 2H 2 0 + 3S. And the sulphur reacts with the already formed sulphite, as indi- cated in the equation given above. (3.) When sulphur is boiled with sodium hydroxide, or with milk of lime, mixtures of sulphides and thiosulphates are obtained in both cases — 6NaHO + 4S = Na 2 S 2 0 3 + 2Na 2 S + 3H 2 0. 3Ca(HO) 2 + 12S = CaS 2 0 3 + 2CaS 6 + 3H 2 0. The sodium sulphide can be converted into thiosulphate by the reactions given above. Calcium pentasulphide, on exposure to air, absorbs oxygen and forms a further quantity of thiosulphate with precipitation of sulphur — CaS 6 + 30 = 3S + CaS 2 0 3 . The thiosulphates are decomposed by most acids, with the libera- tion of sulphur dioxide, and precipitation of sulphur. They show a great tendency to form double salts, many of which are soluble in water; thus sodium thiosulphate, in contact with either silver chloride, bromide, or iodide, forms the soluble double sodium-silver thiosulphate, NaAgS 2 O s — Na 2 S 2 0 3 4- AgCl = NaCl + NaAgS.O*. 39 « Inorganic Chemistry The employment of sodium thiosulphate in photography, for fi *' ng i ne i gat . 1VCS ° r S,lver P rmts > depends upon this property. Thiosulphunc acid may be regarded as being derived from sul- phuric acid by the replacement of one of the (HO) or hydroxyl groups, by an equivalent of (HS) or hydrosulphyl— HO HO SO, HS HO | SO2. Dithionic Acid, H 2 S 2 0 6 or }.-This compound is prepared, by passing a stream of sulphur dioxide through water in which manganese dioxide is suspended, whereby manganese dithionate is formed; while It the verted* into" V*** ^ I* * ^ U ?° n b * dioide, and con! verted into manganous sulphate, thus — 2SO a + Mn0 2 = MnS 2 O e . MnS 2 O e + Mn0 2 = 2MnS0 4 . On the addition of barium hydroxide to the solution, barium dithionate barium sulphate, and manganous hydrate are formed— ’ MnS 2°6 + Ba(HO) 2 = BaS 2 O e + Mn(HO)j. Barium dithionate, being soluble, is separated by filtration, and upon evaporation separates out m crystals of the composition BaS 2 0 fi ,2H 9 0 Upon the addition of dilute sulphuric acid in amount demanded' by the equation * BaS 2 0 6 + H2S0 4 = BaS0 4 + H 2 S 2 0 6 , 1S ° btained - The so ^tion may be concentrated in vacuo until h v SP ! ClfiC , gr v aVlty ° f I<347 - Further concentration results in its de- composition into sulphuric acid and sulphur dioxide— H 2 S 2 0 6 = S0 2 + HjjSCV Dithionic acid forms well-defined crystaUine salts, which on heating decom- pose into sulphates with evolution of sulphur dioxide Dithionic acid was formerly called hyfomlfhuric acid, and its salts are stiU sometimes referred to as hypo sulphates. Trithiomc Acid, H 2 S 3 O e or no sOg } S * — ‘ The potassium salt of this acid may be obtained, by passing sulphur dioxide through a strong solution of potassium thiosulphate — 3SO a + 2K 2 S 2 0 3 = S + 2K 2 S 3 0 6 . It is also formed when a solution of potassium silver thiosulphate is boiled - KO AgS AgS KO }so 2 }so, Pentathionic Acid 399 The sodium salt may be obtained, by the addition of iodine to a mixture of sodium sulphite and thiosulphate — 22 } S ° a + Na2S ° s + Ia = 2NaI + NaO’SOo } S - The acid itself is obtained by the addition of fluosilicic acid to a solution of the potassium salt, when insoluble potassium fluosilicate is precipitated. Both the acid itself, and its salts, are readily decomposed into sulphur dioxide, sulphur, and either sulphuric acid or a sulphate, thus — k 2 s 3 o 6 = k 2 so 4 + so 2 + S. When acted upon by sodium amalgam, sodium trithionate is converted back again into its generators, sodium sulphite and thiosulphate, thus — NaO-S0 2 } s + Na a = Na 2 SO s + NaS } S0 * HO # SO ) Tetrathionic Acid, H 2 S 4 0 6 or H q. S q 2 \ S 2*— Tbe sodium salt is obtained by the action of iodine upon sodium thiosulphate — NaO'SOo ) o 2NaS-Na0-S0 2 + I 2 = 2NaI + Na0 -S0 2 f ^ The barium salt, from which the acid itself is most readily obtained, is pre- pared by the gradual addition of iodine to barium thiosulphate in water — 2BaS 2 0 3 + Ij — Bal 2 + BaS 4 0 3 . The barium tetrathionate is separated by the addition of alcohol, which dis- solves the iodide and excess of iodine, leaving the tetrathionate. By the addition of dilute sulphuric acid to an aqueous solution of this salt, in amount demanded by the equation — BaS 4 O e + H2S0 4 = H 2 S 4 O s + BaS0 4 , a dilute aqueous solution of the acid may be obtained. The dilute acid may be boiled without decomposition ; but when concentrated, it readily passes into sulphuric acid, sulphur dioxide, and sulphur. Sodium amalgam decomposes the sodium salt into two molecules of thio- sulphate, reversing the reaction by which it is produced. Pentathionic Acid, H 2 S 5 0 6 or ho's 0 2 } S 3* — ' This acid is P re P ared passing sulphuretted hydrogen into a strong aqueous solution of sulphur dioxide — 5S0 2 + 5H 2 S = H 2 S 6 0 6 + 5S + 4H a O. or— 5H2S0 3 + 5H 2 S = H 2 S 5 0 6 + 5S + 9H 2 0. The solution contains, however, more or less of the other thionic acids, but as the passage of sulphuretted hydrogen is continued, these are gradually 400 Inorganic Chemistry decomposed, and ultimately the pentathionic acid also, so that the final products of the action of excess of this gas will be sulphur and water— H aS 6 O s + 6HjS = 6H a O + 10S. The solution obtained by the first action, may be concentrated by cautious evaporation in vacuo, until a specific gravity of 1.46 is obtained, when on partial saturation with potassium hydroxide and filtration, a solution is obtained which on spontaneous evaporation deposits crystals of potassium pentathionate, having the composition K 2 S 5 0 6 ,3H 2 0. On heating, the salt splits up into potassium sulphate, sulphur dioxide, and sulphur. OXYCHLORIDES OF SULPHUR. Four of these compounds are known, all of which may be re- garded as being derived from the oxyacids by the replacement of hydroxyl (HO) by its equivalent of chlorine. Sulphurous °Moride Cl } S0 corres P on =c!} so *+h 8 }so* Properties. — Sulphuryl chloride is a colourless liquid, which fumes in moist air, and has a specific gravity of 1.66. It boils at 70°, and is decomposed by water with formation of sulphuric acid and hydrochloric acid — g } S0 2 + 2H a O = 2HC1 + } SO* Sulphuric Chlorhydrate, S0 2 C1(H0). This compound is the first pro- duct of the replacement of the (HO) groups in sulphuric acid by chlorine, and is formed by the direct combination of sulphur trioxide and hydrochloric acid — S0 3 + HC1 = HCISO3 or S0 2 C1(H0). It may be obtained by distilling sulphuric acid with phosphorus oxychloride 2 HO } S ° 2 + P0Cl3 = 2 C1° } S ° 2 + HC1 + hp 0 3* Or by passing dry gaseous hydrochloric acid into melted pyrosulphuric acid — HjjSjA + 2HC1 = H a O + 2HCISO3. Properties. — Sulphuric chlorhydrate is a colourless fuming liquid, having a specific gravity of 1.76, and boiling at i49°-i5i°, with partial dissociation into its generators, sulphur trioxide and hydrochloric acid. In contact with water it is decomposed with considerable violence, with formation of sulphuric and hydrochloric acids — q° } S °2 + H 2 0 = HC1 + } SO* ciso ) Disulphuryl Chloride ( pyrosulphuric chloride ), Qg© 2 \ O or S 2 0 5 C1 2 . This substance is obtained by the action of sulphur trioxide, or sulphuric chlor- hydrate, upon phosphorus pentachloride — 2SO3 + PCI5 = POCI3 + s 2 o 5 ci* 2S0 2 Cl(HO) + PC1 5 = POCI3 + 2HC1 + S 2 0 5 C1* It is also produced by the action of sulphur trioxide upon sulphur dichloride — 5S0 3 + S 2 C1 2 = S 2 0 5 C1 2 + 5SO* Or by the action of sulphur trioxide upon sulphuric chloride— Cl Cl so, + so, = gig 402 Inorganic Chemistry Properties. -Pyrosulphurlc chloride is a heavy, oily, fuming liquid, resem JS? a P r ,^ Ph w C h r id m a H PP “ ?* has a ^ gravity o', ,8^ ..I k V V, Whe " mlMd w,lh *»«r it Slowly decomposes into sulphuric XMXlSt 5howing a “ differen “ 1 * his »p« ^ S a 0 6 Cl 2 + 3H 2 0 = 2H 2 S0 4 + 2HCL CARBON DISULPHIDE. Formula, CS 2 . Molecular weight = 76. Vapour density = 38. History.— This compound was accidentally produced by Lam- padius (1796) when heating a mixture of charcoal and pyrites Mode of Formation.-Carbon disulphide is prepared by passing the vapour of sulphur over red-hot charcoal, when the two elements unite and form the volatile product, which is condensed in vessels surrounded with cold water— C + S 2 = CS* The product is always contaminated with free sulphur, which volatilises, and is also accompanied by considerable quantities of sulphuretted hydrogen, formed by the action of sulphur upon the hydrogen contained in the charcoal. When carbon disulphide is prepared on a manufacturing scale, the charcoal is heated in a vertical cast-iron or earthenware retort,’ C, Fig. 113, having an elliptical section, and provided with three openings. The retort is built into a suitable furnace, whereby it can be uniformly heated to redness. A quantity of sulphur, con- tained in the pot S, kept liquid by the heat of the furnace, is allowed to enter at intervals through the pipe B. As the vapour comes in contact with the red-hot charcoal, combination ensues, and the carbon disulphide escapes through the pipe D, which is inclined to the retort so as to allow condensed sulphur to run back. Sulphur which escapes condensation in this pipe, collects, for the most part, in the vessel E, which is closed by water seals as seen in the figure. The volatile compounds are then passed through a Liebig’s condenser about 30 ft. long, and the crude disulphide so condensed is collected in a receiver. Any vapour of carbon disulphide which is carried on by the sulphuretted hydrogen, is absorbed by passing the gas through a scrubber containing oil ; and finally the Carbon Disulphide 403 sulphuretted hydrogen is absorbed in a lime purifier, similar to those employed for the purification of coal gas. The ashes are withdrawn from the retort through the wide tube B ; and the fresh charcoal is introduced through the opening A. In order to prevent the escape of the unpleasant and injurious vapours from A, during the addition of fresh charcoal, the opening A' is put into communication with the chimney of the furnace. The sulphur which flows back into the retort from D, is conveyed by means of the pipe f nearly to the bottom of the mass of heated char- coal, so that its vapour shall once more be made to pass over the carbon. The crude product is purified by distillation, and subsequent agitation with mercury. Properties. — Carbon disul- phide is a colourless, mobile, and highly refracting liquid. When perfectly pure it possesses a sweet- ish, and not unpleasant, ethereal smell, but as usually met with the odour is decidedly foetid. Its specific gravity at o° is 1.292, and it boils at 46°. The vapour Fig. 113. of carbon disulphide has a very low igniting-point (see page 291). It burns with a blue flame, which, when fed with oxygen, emits a dazzling blue light. When carbon disulphide vapour is mixed with three times its volume of oxygen, and a light applied, the mixture explodes with violence ; the products of the combustion being carbon dioxide and sulphur dioxide — CS 2 + 30 2 = C 0 2 + 2SO a . The vapour of carbon disulphide, when constantly inhaled in small quantities, has an injurious effect upon the health, and if breathed in large quantities is a powerful poison. When heated to a bright red heat, carbon disulphide vapour is decomposed into its constituent elements : on this account, in the manufacture of this compound, care is taken that the temperature does not rise too high. The vapour of carbon disulphide is decomposed by potassium, 404 Inorganic Chemistry which, when sulphide, and heated, burns in the vapour, forming potassium liberating carbon — CS 2 + 2K 2 = C + 2K 2 S. When passed over heated slaked lime, carbon disulphide vapour is converted into carbon dioxide and sulphuretted hydrogen— CS 2 + 2CaH 2 0 2 = 2CaO + C0 2 + 2H 2 S. Ttus reaction is made use of for converting the carbon disul- phide, which is always present in coal gas, into the two easily removed substances, carbon dioxide and sulphuretted hydrogen When a mixture of carbon disulphide vapour and sulphuretted hydrogen, is passed over heated copper, marsh gas is formed 4Cu + CS 2 + 2H 2 S = CH 4 + 4CuS. Carbon disulphide is soluble to a minute extent in water ; i volume of water dissolves .001 volume of this liquid, and the solution possesses the taste and the smell of the disulphide. It mixes in all proportions with alcohol, ether, the hydrocarbons of the benzene family, and most essential oils. It also dissolves sulphur, phosphorus, iodine, bromine, caoutchouc, and most fats ; and is largely used in the arts, both as a solvent for caoutchouc, and in extracting essential oils, spices, and perfumes. Thioearbonic Acid.— Carbon disulphide is the sulphur ana- logue of carbon dioxide, CS 2 ; C0 2 . Like the oxygen compound, it forms a feeble acid, which has received the name thioearbonic acid, H 2 CS 3 ; carbonic acid, H 2 C0 3 . The thiocarbonates are produced by reactions analogous to those by which carbonates are formed. Thus, when carbon disul- phide is brought into contact with potassium sulphide, potassium thiocarbonate is obtained — CS 2 + K 2 S = K 2 CS 3 , which may be compared with the action of carbon dioxide upon potassium oxide — co 2 + k 2 o = k 2 co 3 . Thiocarbonates are likewise formed by the action of carbon disulphide upon metallic hydroxides — 3CS 2 + 6KH0 = 2K 2 CS 3 + K 2 CO a + 3H 2 0. Selenium 405 The acid itself is obtained as £ yellow oil, having an unpleasant odour, by the decomposition of a thiocarbonate by dilute hydro- chloric acid. A large number of compounds are known, in which divalent sulphur replaces oxygen, and which therefore stand in the same relation to the oxygen compounds, as thiocarbonic acid stands to carbonic acid ; for example — Other Compounds of Carbon and Sulphur. — When carbon disulphide is exposed to the influence of light, there is gradually formed upon the glass vessel containing it, a brown deposit, which is believed to be carbon mono- sulphide, CS ; the sulphur analogue of carbon monoxide. When electric sparks from carbon poles are passed through the vapour of carbon disulphide, or when the electric arc is produced in the vapour, an offensive-smelling liquid is obtained, which exerts a most irritating and tear-producing effect upon the eyes. This liquid has been shown to have the composition C 3 S 2 . * SELENIUM. Symbol, Se. Atomic weight = 78.87. Molecular weight = 157.74. History. — This element was discovered by Berzelius (1817), who gave it the name selenium (signifying the moon) on account of its close analogy with the previously discovered element tellurium (signifying the earth). Occurrence. — Selenium is occasionally met with associated with native sulphur, probably as a selenide of sulphur. In a few minerals of considerable rarity, selenium is met with in the form of selenides of such metals as mercury, lead, silver. It occurs in very small quantities in a large number of metallic sulphides. Modes of Formation. — (1.) When pyrites containing selenium is employed in the manufacture of sulphuric acid, the selenium is oxidised by the atmos- pheric oxygen into selenium dioxide, which is carried forward with the sulphur dioxide. Selenium dioxide, being a solid substance, is partly deposited in the flues, and in the Glover tower, and partly carried forward into the chambers, where it forms a red-coloured deposit. To obtain the selenium, either the flue dust or the chamber deposit, is first boiled with dilute sulphuric acid, and either nitric acid or potassium chlorate added, in order to oxidise it completely into selenic acid, H 2 Se 0 4 . The solution is then boiled with strong hydrochloric acid, whereby it is reduced to selenious acid, H 2 Se 0 3 , when a stream of sulphur dioxide is passed through it which precipitates the selenium as a red powder— Tl > ; H 2 Se 0 3 4- 2 S 0 2 + H 2 0 — Se + 2 H 2 S 0 4 . Von Lengyel, 1894. 4 06 I norganic Chemistry (2.) A second method for the preparation of selenium from the chamber deposit, consists in digesting the substance with potassium cyanide, whereby it is converted into soluble potassium selenocyanide, SeK(CN). On the addition of hydrochloric acid to this solution, the element is precipitated as a red amorphous powder, and hydrocyanic acid and potassium chloride go into solution — SeK(CN) + HC 1 = Se 4- KC 1 + H(CN). Properties.— Selenium is known in various allotropic modifications. 1. Soluble in carbon disulphide. — a. Brick-red amorphous powder, obtained by precipitation with acids, or reduction of selenious acid, in the cold, by sulphur dioxide. P- Black crystalline powder, obtained by reduction of hot selenious acid by sulphur dioxide. 7. Dark red translucent monoclinic crystals, specific gravity 4.5, deposited from solution in carbon disulphide. d. Black, shining, brittle amorphous mass, having a conchoidal fracture, and a specific gravity of 4.3, obtained by rapidly cooling melted selenium. 2. Insoluble in carbon disulphide. —Black, metallic-looking crystalline mass, having a granular fracture. Obtained by quickly cooling melted selenium to 210 0 and keeping it for some time at that temperature, when the mass solidifies with rise of temperature to 217 0 . This insoluble variety, sometimes called metallic selenium, is also formed as a deposit of minute black crystals, when concentrated solutions of sodium or potassium selenide are exposed to the air, This modification has a specific gravity of 4.5, and melts at 217 0 . Selenium boils at 68o°, forming a dark red vapour which condenses in the form of flowers of selenium, having a scarlet-red colour. At high temperatures the vapour of selenium, like that of sulphur, becomes a true gas ; thus at 1420°, the vapour density is found to be 81.5, approaching very closely to the normal density demanded by the molecule Se^. “Metallic” selenium conducts electricity, and the element exhibits the remarkable property of having its conductivity increased by light ; the con- ductivity of selenium when exposed to diffused daylight being about twice as great as when in the dark. This alteration in the electrical resistance with varying intensities of light, is a property of selenium that was made use of, in the construction of an instrument known as the photophone, but it has not as yet been put to any practical use. When selenium is heated in the air, it burns with a blue flame, with the formation of selenium dioxide, and at the same time emits a powerful and characteristic smell resembling rotten horse-radish. When selenium is heated in a tube, filled with an indifferent gas, it sublimes in the form of a red deposit ; but when heated in hydrogen, the sublimate is in the form of black shining crystals. The formation of these crystals is due to the fact, that selenium combines with the hydrogen, and the hydrogen selenide is again decomposed by the heat. Hydrogen Selenide ( selenuretted hydrogen ), H 2 Se ; molecular weight = 81. Hydrogen selenide is formed when selenium is heated in hydrogen. This compound is also obtained, by the action of dilute hydrochloric or sul- phuric acid, upon either potassium selenide or ferrous selenide — FeSe + HjS 0 4 = FeSQ 4 4 HjSe. Selenium Dioxide 407 Properties. — Hydrogen selenide is a colourless gas, strongly resembling sulphuretted hydrogen, both in its smell, and in its chemical behaviour. It is readily soluble in water, and when passed through metallic solutions, precipi- tates insoluble selenides of most of the heavy metals. Hydrogen selenide burns with a blue flame, with the production of water and selenium dioxide. Its smell, although resembling that of its sulphur analogue, is more unpleasant, and its effects upon the system are more persistent and injurious. A single small bubble inhaled through the nostril, produces temporary paralysis of the olfac- tory nerves, accompanied by inflammation of the mucous membrane. No compound of selenium corresponding to hydrogen disulphide is known. Compounds with Halogens. Diselenium Dichloride, Se 2 Cl 2 , is obtained by passing chlorine over selenium, or by passing gaseous hydrochloric acid through a solution of selenium in nitric acid. > Properties.— Selenium chloride is a brown oily liquid, in which selenwm itself is readily soluble, and from which the element is deposited in the form which is insoluble in carbon disulphide. It is slowly decomposed by water, thus — 286303 + 3H a O = H 2 Se0 3 + 3Se + 4HC1. Corresponding bromine and iodine compounds are known, Se 2 Br 2 , and ^Selenium Tetrachloride, SeCl 4 , is prepared either by the action of chlorine upon selenium chloride — Se 2 Cl 2 + 3C1 2 = 2SeCl 4 , or by heating a mixture of selenium dioxide and phosphorus pentachloride— 3Se0 2 + 3PC1 5 = 3SeCl 4 + P 2 O s + POCl a . Properties.— Selenium tetrachloride is a white, crystalline, volatile com- pound ; which may be sublimed without decomposition and without fusion. When the vapour is heated above 200 °, it begins to dissociate into selenium and chlorine. It dissolves in water, with decomposition into hydrochloric and selenious acids — SeCl 4 + 3H a O = 4HC1 + H 2 Se0 3 . Corresponding bromine and iodine compounds are known, SeBr 4 and Sel 4 . Oxides and Oxyacids of Selenium. Only one oxide of selenium is known, namely, selenium dioxide, Se0 2 , although a second oxide of unknown composition is believed to exist, and to constitute the peculiar smelling substance which is always formed when selenium is burnt in the air. Selenium Dioxide is prepared by burning selenium in a stream of oxygen in a glass tube ; the element burns in the gas with a blue flame, and the oxide condenses upon the distant portions of the tube, as a white crystalline deposit. 408 Inorganic Chemistry solves .n water and gives rise to selenious acid liquidity. It dis- The following oxyacids of selenium are known— Splenic acid Ci H S?° 3 ' corres P° ndin S t0 sulphurous acid, H,S0 3 . , C d ' corres P° nd ‘ng to sulphuric acid, H s SO,. Selenosulphunc HO lg 0 corresponding ( thiosulphuric HO 1 ’ * acid, "" t acid, HSe HS fh S “ 3 A f ld ’ , H 2 Se °s- is obtained as a white crystalline compound when irid d !wh ,S - d,SS °l V ! d m hot water> and the solut *on allowed to cool ’ The ^ U:^ C it a ^ mS b ° th aC, ’ d and n ° rmal $eleniteS ’ corresponding to moW„u f # U L l° rms a senes of so-called superacid salts, containing a molecule of the acid salt, combined with a molecule of acid, thus— g H KSe 0 3 , HgSeOj. Selenic Acid, H 2 Se 0 4 .— This acid is best prepared, by the addition of 7 SelCn !! e T PendGd “ Water ’ ^en insoluble silver tomiSe is formed and selenic acid is left in solution Ag 2 Se 0 3 + H a O + Br 2 = 2 AgBr + H 2 Se 0 4 . srlldc so !“ tion , ms ;) , | 5 * eva P or uted by heating, until it contains 94 per cent, of selenic acid, and still further evaporated in vacuo, until it reaches 97 4 per cent when its specific gravity is 2.637. When heated .0 280° it dec?mpTsL“nto selenium dioxide, water, and selenium. Properties. —Selenic acid in its most concentrated condition, is a colourless ^ g w T d h WhlCh miX6S With water with the development of con- siderable heat. It dissolves iron and zinc with evolution of hydrogen ; and when heated, dissolves copper with formation of selenious acid. The selenates closely resemble the sulphates. Barium selenate, like the sulphate is quite insoluble in water, but differs from that compound in being converted by boiling hydrochloric acid into barium selenite, which is soluble & Selenium also forms a compound with oxygen and chlorine, selenium oxv- SOCLj de ’ ^ ^ 6nyl Chl ° ride ’ Se0CI 2- corresponding with thionyl chloride, TELLURIUM. Symbol, Te. Atomic weight,* 125 (?). Occurrence.— In the free state, small quantities of this element have been found as crystals, consisting of almost pure tellurium. In combination it is * Various numbers have been obtained, by different observers, for the atomic weight of tellurium. Some of these numbers are higher than the atomic weight of iodine, which would make it impossible to give to tellurium a posi- tion between antimony (atomic weight = 120) and iodine (atomic weight = 126.54) as demanded by the periodic law. Brauner, who has spent many years investigating this point, considers that hitherto pure tellurium has never been obtained. Tellurous Acid 409 met with in a few rare minerals, such as tellurite [ Te0 2 ), and, more commonly, tetradymite (Bi 2 Te 3 ). Some specimens of pyrites contain small quantities of this element, hence it is found in the deposit from the vitriol chambers, from which selenium is obtained. Mode of Formation. — Tellurium is obtained from bismuth telluride, Bi 2 Te 3 , by fusion with an intimate mixture of sodium carbonate and carbon. The mass on treatment with water, yields a solution containing a mixture of sodium telluride and sodium sulphide, which on exposure to the air deposits tellurium as a grey powder. The element is purified by distillation in a stream of hydrogen. Properties.— Tellurium is a bluish- white, silver-like solid, possessing metallic lustre. It conducts heat and electricity, although badly, and is very brittle. Its specific gravity is 6.26, and it melts at 452 0 . When melted tellurium is slowly cooled, it forms rhombohedral crystals. When heated in the air, it burns with a blue flame, and forms tellurium dioxide, Te0 2 . When heated in a sealed tube with hydrogen, tellurium sublimes in the form of brilliant prismatic crystals. Hydrogen Telluride ( Telluretted Hydrogen ), H 2 Te. — When tellurium is heated in hydrogen, the elements combine, forming hydrogen telluride, which exhibits the same phenomenon as is shown by selenuretted hydrogen, of being decomposed by heat, and depositing the element as a crystalline sublimate. Hydrogen telluride is obtained by the action of hydrochloric acid upon zinc telluride — ZnTe + 2HC1 = ZnCl 2 + H 2 Te. Properties. — Hydrogen telluride is a most offensive smelling, and highly poisonous gas. It behaves like sulphuretted hydrogen in precipitating metals from solutions. It is soluble in water, and the solution gradually absorbs oxygen and deposits tellurium. Compounds with the Halogens. Two chlorides of tellurium are known, namely, tellurium dichloride, TeCl 2 , and tellurium tetrachloride, TeCl 4 . It will be noticed that the composition of the dichloride is not analogous with the lower chloride of either selenium {Se 2 Cl 2 ) or sulphur (S 2 C1 2 ). Two bromides, TeBr 2 and TeBr 4 , and corresponding iodides are known. Oxides and Oxyacids of Tellurium. Two oxides of tellurium are known with certainty, namely, tellurium dioxide, Te0 2 , and tellurium trioxide, Te0 3 , which give rise respectively to the two acids, tellurous acid, H 2 Te0 3 , and telluric acid, H 2 Te0 4 . Tellurous Acid is obtained, by pouring a solution of tellurium in nitric acid into an excess of water. The acid is precipitated as a white amorphous powder. When strongly heated, it is converted into the dioxide and water. Tellurous acid, like sulphurous acid, is dibasic, and gives rise to both acid and normal salts : thus, with potassium it forms hydrogen potassium tellurite, 4io Inorganic Chemistry HK TeOj, and dipotassium tellurite, KjTeOj. It also forms super-acid salts such as — Quadracid potassium tellurite . . . HKTeO,, H 2 Te0 3 . Potassium tetratellurite .... K 2 Te0 3 , 3Te0 2 . Telluric Acid is prepared by fusing either tellurium, or tellurium dioxide, with a mixture of potassium nitrate and carbonate— Te a + K 3 C0 3 + 2KNO a = 2K 2 Te0 4 + N 2 + CO. 1 he fused mass, after solution in water, is mixed with a solution of barium chloride, which precipitates barium tellurate ; this is then decomposed by the addition of the exact amount of sulphuric acid, and after filtration, the clear solution deposits crystals of telluric acid, H 2 Te0 4 ,2H 2 0. When these crystals are heated to i6o°, the water is expelled, and the anhydrous acid in the form of a white powder is left. On strongly heating, telluric acid decomposes into water and tellurium trioxide, which at a higher temperature splits up into the dioxide and oxygen. Like tellurous acid, telluric acid forms not only normal and acid salts, but a number of more complex superacid salts — Normal potassium tellurate Hydrogen potassium tellurate Quadracid potassium tellurate Potassium tetratellurate K 2 Te0 4 ,5H 2 0. HKTe0 4 . K 2 Te0 4 ,H 2 Te0 4 ,3H 2 0. K 2 Te0 4 ,3H 2 Te0 4 , H 2 0. CHAPTER III THE ELEMENTS OF GROUP V. ( FAMILY B .) Nitrogen, N . . 14.01 I Antimony, Sb . . 119*6 Phosphorus, P . . 30* 9 6 Bismuth, Bi . . 207.5 Arsenic, As. . .74.9 ' In this family of elements we have a gradual transition from the non-metals to the metals. Nitrogen and phosphorus may be con- sidered as typical non-metallic elements, both as regards their physical and chemical properties. The third member, arsenic, begins to exhibit metalline properties ; its specific gravity is more than three times as high as that of phosphorus, and it possesses considerable metallic lustre : arsenic is called a metalloid on this account. Antimony is still more metallic in its character, possess- ing most of the physical attributes of a true metal, while in bismuth all non-metallic properties cease altogether to exist. All these elements form more than one compound with oxygen, of which the following may be compared — N 2 O s ; (P 2 0 3 ) 2 ; (As 2 0 3 ) 2 ; Sb 2 0 3 ; Bi 2 0 3 . N 2 0 4 ; P 2 0 4 ; — Sb 2 0 4 ; Bi 2 0 4 . N 2 0 5 ; P 2 0 6 5 As 2 0 5 ; Sb 2 0 6 ; Bi 2 0 6 . The oxides (which in the case of nitrogen and phosphorus are strongly acidic in their nature, combining with water to form acids) gradually become less and less acidic and more basic as the series is traversed. Thus, nitrogen pentoxide, N 2 0 5 , unites violently with water to form nitric acid, which with bases yields nitrates. Antimony pent- oxide is insoluble in water, and no antimonic acid has been isolated, although its salts, the antimonates, are known. The oxides of antimony, on the other hand, begin to exhibit basic properties, and unite with acids, forming salts in which the antimony functions as the base. 411 412 Inorganic Chemistry In the case of the last element, the acidic nature of the oxides is to arn monl; n0b,smuth com P oun ds being known, corresponding to antimonates or arsenates, while these oxides unite with acids in the capacity of bases, giving rise to bismuth salts. Four of the elements of this group unite with hydrogen, forming similarly constituted compounds, NH 3 , PH 3 , AsH 3 , SbH g The stability of these compounds gradually decreases a's we pass mi^Hf ° g r t0 an ' lm ° n> ' Antimon y hydride has never been ob- tained free from other gases, while no similar bismuth compound is known. Ammonia is alkaline and strongly basic, and unites readily with acids to form ammonium salts. Phosphorus hydride has no .^ ‘ n , e cba , racter > and is only feebly basic. It combines, however, with the halogen acids to form phosphonium chloride, bromide and iodide, PH 4 C 1 , PH 4 Br, PH 4 I, analogous to ammonium ch onde, bromide, and iodide. The hydrides of arsenic and antimony exhibit no basic character. All the elements of this group unite with chlorine, giving rise to the compounds NCIj, PC 1 3 , AsC 1 3 , SbCl„ BiCI l 3> which also exhibit a gradation in their properties : thus, nitrogen trichloride is an extremely unstable liquid, exploding with extra- ordinary violence upon very slight causes, while the analogous bismuth compound is a perfectly stable solid. The boiling-points of these compounds show a gradual increase with the increasing atomic weight of the element : thus, nitrogen chloride boils at 71 0 , phosphorus trichloride at 78°, arsenic tri- chloride at 130.2 0 , and antimony trichloride at 200°. The elements arsenic, antimony, and bismuth are isomorphous, and their corresponding compounds are also isomorphous. The first member of this family, namely, nitrogen, has been already treated in Part II., as one of the four typical elements studied in that section of the book. It occupies a position in relation to the other members of the family, very similar to that of oxygen towards sulphur, selenium, and tellurium. PHOSPHORUS. Symbol, P. Atomic weight = 30. 96. Vapour density = 61. 92. Molecular weight = 123.94. History.— Phosphorus was first discovered by the alchemist Brand, of Hamburg (1669), who obtained it by distilling a mixture 413 Phosphorus of sand with urine which had been evaporated to a thick syrup. The process, however, was kept secret. Robert Boyle (1680) dis- covered the process of obtaining this element, but the method was not published till after his death. Until the year I77 1 ? when Scheele published a method by which phosphorus could be ob- tained from bone ash, this element was looked upon as a rare chemical curiosity. The nam z phosphorus was not first coined for this element : it had been in previous use to denote various sub- stances known at that time, which had the property of glowing in the dark. To distinguish the element it was called Brand's phos- phorus , or English phosphorus. Occurrence. — Phosphorus has never been found in nature in the free state. In combination with oxygen and metals, as phosphates, it is very widely distributed, especially as calcium phosphate. The following are some of the commonest natural phosphates — Sombrerite, or estramadurite . Ca 3 (P0 4 ) 2 . Apatite 3Ca 3 (P0 4 )2,CaCl 2 . Wavellite 2A1 2 (P0 4 ) 2 ,A1 2 (H0) 6 ,9H 2 0. Calcium phosphate is present in all fertile soils, being derived from the disintegration of rocks : the presence of phosphates in soil has been shown to be essential to the growth of plants. F rom the vegetable it passes into the animal kingdom, where it is chiefly present in the urine, brain, and bones. Bones contain about 60 per cent, of calcium phosphate, to which they entirely owe their rigidity. Mode of Formation. — Manufacture. The chief source of phosphorus is bone ash, a material obtained by burning bones, and which consists of nearly pure calcium phosphate, Ca 3 (P0 4 ) 2 . Other varieties of calcium phosphate, such as sombrerite and apatite, are also employed, as well as phosphates of other metals, such as the Redonda phosphates, which consist of phosphates of iron and alumina. The bone ash, in fine powder, is first decom- posed by means of sulphuric acid, specific gravity 1.5 to 1.6. This operation is performed in large circular wooden vessels, resem- bling a brewer’s “ mash tun,” provided with an agitator, and into which high pressure steam can be driven. Finely-ground bone ash and sulphuric acid, in charges of a few cwts. at a time, are alter- nately stirred into the decomposer, until from four to five tons of 414 Inorganic Chemistry ^^ en : ntro f uced - with sufficiem add to c ° nvert ° f the hme mt0 calclum sulphate, according to the equa- Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 P0 4 . The contents of the decomposer are next run out into filtering tanks and the phosphoric acid is then concentrated to a syrup in arge lead-lined pans through which steam-coils pass, the liquor being constantly agitated by a mechanical stirrer q The concentrated liquor is next mixed, either with sawdust or Ttelvd 3 - S H h S K U n T* 1 ’ ° r C ° ke> and the mixture com- pletely dried by being heated in a cast-iron pot, or in a muffle to a Fig. 114. dull red heat. During this process, the tribasic (or orthophosphoric acid), H 3 P0 4 , is converted into metaphosphoric acid, HP0 3 — phosphoric acid by loss of water 3 po 4 = h 2 o + HPO, The charred mixture is then distilled in bottle-shaped retorts of Stourbridge clay, about 3 feet long, and having an internal diameter of 8 inches. A number of these retorts, usually twenty-four, are arranged in two tiers, in a galley furnace, as seen in section in Fig 1 14. The empty retorts are first gradually raised to a bright red heat, and a charge of the mixture is then quickly introduced. Bent 415 Phosphorus pieceb of 2 -inch malleable iron pipe are then luted into the mouths of the retorts, connecting them with the pipes, D D'. These pipes dip into troughs of water, E E', which run along the entire length of the furnace, and in which the phosphorus con- denses. The temperature of the furnace is then raised to a white heat, when decomposition of the metaphosphoric acid commences, and phosphorus begins to distil over. The process is continued for about sixteen hours. The change that goes on is mainly represented by the following equation — 4HP0 3 + 12C = 12CO + 2H 2 + 4P. The crude product, which is usually dark red or black in appear- ance, is first melted under hot water, and thoroughly stirred, in order to allow the greater part of the rougher suspended matters to rise to the surface. The mass is then allowed to resolidify. The exact processes by which phosphorus is further purified on a manu- facturing scale are guarded as trade secrets : one method that has been in use, consists in treating the phosphorus while melted under water, with a mixture of potassium dichromate and sulphuric acid, whereby some of the impurities are oxidised, and others are caused to rise to the surface as a scum, leaving the phosphorus as a clear liquid beneath. Phosphorus usually comes into commerce either in the form of wedges, or as sticks. The operation of casting the phosphorus into sticks is performed beneath water. A quantity of phosphorus beneath a shallow layer of water is placed in the vessel C, Fig. 115 , which is contained in a tank of water through which a steam-coil passes. Connected to the phosphorus reservoir is a glass tube, G, which passes into a second shallow tank of cold water. On open- ing the cock D, the liquid phosphorus flows into the cold glass tube, where it congeals, and it may then be drawn through as a continuous rod of phosphorus, if care be taken not to draw it out faster than it 416 Inorganic Chemistry solidifies. It is the custom to adopt a uniform length and thick- s n u e c S h inCheS ,0nS and * inCh di “ Ni- Properties-When freshly prepared and kept in the dark, Even '!) V translucent . almost colourless, wax-like solid. Even in the dark it soon loses its transparency, and becomes the fif m W that a f n white c film ! while if exposed to the light, the nl h f becomes first y ellow . then brown, and in time the phosphorus assumes a red and even a black colour throughout !ts entire mass. Its specific gravity at i6° is 1.82. At 0“ phos- phorus becomes moderately brittle, and a stick of it may be readily snapped, when its crystalline character will be seen. At it” it becomes soft, and may be cut with a knife, like wax. Phosphorus melts under water at 43-3°, and the liquid exhibits the property of suspended solidification. If the melted material, which has been cooled below its solidifying point, be touched with a fragment of .W^r US UP °, n th t eDd ° f a CapiMar y S lass tube > *e mass instantly congeals with rise of temperature.* Phosphorus contained in a closed vessel, without water, melts at as low a temperature as 30°, t and when heated in air to 34* it takes fire. At a temperature of 269* phosphorus boils, and forms a colourless vapour. Phosphorus is volatile at ordinary temperatures : if a small quantity of phosphorus be sealed in a vacuous tube, and the tube be placed in the dark, the phosphorus will slowly vaporise ; and if one end of the tube be kept slightly cooler than the rest, the phos- phorus will sublime upon that part, in the form of brilliant, colour- less, and highly refracting rhombic crystals, which retain their beauty so long as they are kept in the dark. The density of the vapour of phosphorus is 61.92, giving a molecular weight of 1 23.94, which is four times the atomic weight, showing that the molecule of phosphorus contains four atoms. Even at temperatures as high as 1040 these tetratomic molecules are stable, but it has pi?ce Sh ° Wn that ^ higher tem P eratures , dissociation begins to take On account of its ready inflammability, phosphorus is always preserved under water, which exerts practically no solvent action upon it. It is extremely soluble in carbon disulphide, 1 part of this * See “ Chemical Lecture Experiments,” new ed., Nos. 528, c 2 q . f Readman. J ’ 3 y Phosphorus 4 1 7 liquid dissolving 9.26 parts of phosphorus. On evaporation, the element is deposited in the form of colourless crystals. Phosphorus is also soluble, but to a less extent, in chloroform, benzene, turpen- tine, alcohol, olive oil, and many other solvents. A solution of phosphorus in carbon disulphide, when allowed to evaporate upon a piece of blotting-paper, leaves the element in so finely divided a condition, that its rapid oxidation almost immediately raises the temperature to the ignition point of the phosphorus, when it takes fire. On exposure to moist air in the dark, phosphorus appears faintly luminous, emitting a pale greenish-white light, and at the same time evolving white fumes which possess an unpleasant, garlic-like smell, and are poisonous. These fumes consist mainly of phos- phorous oxide, P 4 0 6 , and the glowing of the phosphorus is the result of its oxidation ; phosphorus does not glow when placed in an inert gas which is perfectly free from admixed oxygen, although the presence of very small traces of free oxygen in such a gas, is sufficient to cause the phosphorescence. At a few de- grees below o°, phosphorus ceases to glow in the air. Although the- glowing is due to oxidation, phosphorus does not appear luminous in pure oxygen at temperatures below about 15 0 . If, therefore, a stick of phosphorus which is glowing in the air, be immersed in a jar of oxygen, its phosphorescence is at once stopped. If, however, the oxygen be slightly rarefied, the phos- phorus again becomes luminous. Similarly, the phosphorescence that is exhibited in air, is stopped if the air be compressed.* The glow of phosphorus is believed to be associated with the formation of ozone, for the presence in the air of traces of such gases and vapours as ethylene, turpentine, or ether, which are known to possess the power of destroying ozone, at once stops the glowing of a stick of phosphorus. Phosphorus is incapable of uniting with oxygen if the gas be perfectly pure and free from aqueous vapour. It has been shown that in oxygen which has been dried by prolonged exposure to the desiccating action of phosphorus pentoxide, phosphorus may not only be melted, but even distilled, without any combination with the oxygen taking place. If water, beneath which is a small quantity of melted phosphorus, be boiled, the phosphorus vaporises with the steam, and renders “ Chemical Lecture Experiments,” new ed., Nos. 530 to 534. 4*8 Inorganic Chemistry the steam luminous : use is made of this property, as a means of detecting free phosphorus, in toxicological analysis. Phosphorus is a powerfully poisonous substance ; in large doses it causes death in a few hours, in smaller quantities it produces stomachic pains and sickness, usually ending in convulsion. Persons constantly exposed to the vapours arising from the hand- ling of phosphorus, either in its manufacture or in the manufacture of matches, are very liable to suffer from caries of the bones of the jaw and nose ; it is believed that this injurious effect is caused by the white fumes which are the product of oxidation, and not by the actual vapour of phosphorus. Red Phosphorus.— When phosphorus is heated to a tempera- ture between 240° and 250°, out of contact with air, it passes into an allotropic modification. The same transformation takes place when phosphorus is heated to 200° with an extremely small pro- portion of iodine. Red phosphorus is manufactured by heating ordinary phosphorus in a cast-iron pot, provided with a cover, through which passes a Fig. 1 16. 419 Red Phosphorus short open pipe, D (Fig. 116). The pot is carefully and uniformly heated to between 240’ and 250°, as indicated by the thermometers C C, which are encased in metal tubes, to prevent the phosphorus from attacking the glass. A small quantity of the phosphorus becomes oxidised by the air within the vessel, but after this atmos- pheric oxygen is used up, no further oxidation takes place. If the temperature be allowed to rise above 260°, the red phosphorus is reconverted into the ordinary modification, and with the evolution of so much heat, that unless the open tube be provided, as a safety- valve, the iron vessel is liable to burst. The material that is obtained at the end of the operation, is in the form of hard, solid lumps, which still contain a certain amount of the unchanged phosphorus mixed with them. It is first ground to powder beneath water and then boiled with a solution of sodium hydroxide ^caustic soda)’ to remove the ordinary phosphorus, and finally washed and dried. Properties. — Red phosphorus, as usually sent into commerce, is a chocolate-red powder, having a specific gravity of 2.25. It is not luminous in the dark, and has no taste or smell. It is not poisonous, and when taken into the system is excreted un- changed. It is not soluble in carbon disulphide, or in any of the solvents which dissolve ordinary phosphorus. Red phosphorus is unaffected by exposure to dry air or oxygen, but in the presence of moisture it is very slowly oxidised. If red phosphorus which has been perfectly freed from ordinary phosphorus, and carefully washed and dried, be exposed to air and moisture, it is found after the lapse of some time to have become acid, owing to slight oxida- tion into phosphoric acid. When heated in contact wither, red phosphorus does not ignite below a temperature of 240°. Red phosphorus may be obtained in the form of rhombohedral crystals by heating the substance under pressure to a temperature of 580°. The chief use of phosphorus is in the manufacture of matches. When ordinary phosphorus is employed, the bundles of wooden splints are first tipped with melted paraffin wax, and afterwards dipped into a paste, made of an emulsion of phosphorus, chlorate of potash, and glue. Matches so made, ignite when rubbed upon any rough surface ; the paraffin (which is sometimes replaced by sulphur) serving to transmit the combustion from the phosphorus to the wood. Since the discovery of red phosphorus, and its non- injurious properties, the old phosphorus match has been largely superseded by the so-called safety matches. In these matches the 4 20 Inorganic Chemistry splints are tipped with a mixture of potassium chlorate, potassium dichromate, red lead, and antimony sulphide, and they are ignited by being rubbed upon a prepared surface, consisting of a mixture of antimony sulphide and red phosphorus. Although these matches will not ignite by ordinary friction upon any but the specially prepared surface, they may be inflamed by being swiftly drawn along a sheet of ground glass, or strip of linoleum. COMPOUNDS OF PHOSPHORUS WITH HYDROGEN. Three compounds of phosphorus and hydrogen are known namely — PH 3 (gaseous) ; P 2 H 4 (liquid) ; and P 4 H 2 (?) (solid). GASEOUS PHOSPHORETTED HYDROGEN {Hydrogen phosphide: Phosphine ). Formula, PH 3 . Molecular weight = 33.96. Density = 16.98. Modes of Formation. — (1.) This substance is formed when red phosphorus is gently heated in a stream of hydrogen. (2.) It may be prepared by bpiling phosphorus with a solution of potassium or sodium hydroxide — 3NaHO + 4P + 3H 2 0 = 3NaH 2 P0 2 + PH 3 . In this reaction a small quantity of the liquid hydride (P-jH^ is simultaneously formed, which imparts to the gas the property of spontaneous inflammability. It also contains a certain quantity of free hydrogen, produced by the action of the caustic alkali upon the sodium hypophosphite, thus — NaH 2 P0 2 + 2NaHO = 2H 2 + Na 3 P0 4 . To obtain the gas by this method, a quantity of a strong solution of caustic soda, and a few fragments of phosphorus, are placed in a flask, fitted as shown in Fig. 117. A stream of coal gas is passed through the apparatus, in order to displace the air, and the solution is gently heated. Phosphoretted hydrogen is readily disengaged, and as each bubble escapes into the air, it bursts into flame, and forms a vortex ring of white smoke of phosphoric acid. If alcoholic potash be substituted for the aqueous solution in this reaction, the liquid phosphoretted hydrogen is dissolved in Phosphoretted Hydrogen 4 21 the alcohol, and the gas which is evolved is therefore not spon- taneously inflammable .* , , • „ . (3.) Phosphoretted hydrogen is also produced by the action ol water upon calcium phosphide 3P 2 Ca 2 + 12H 2 0 = 6Ca(HO) 2 + 4PH 3 + 2P. A secondary reaction, by which liquid phosphoretted hydrogen is formed, goes on simultaneously P 2 Ca 2 + 4H 2 0 = 2Ca(HO) 2 + P 2 H 4 . The gas, therefore, that is evolved is spontaneously inflammable. Fig. i 17. (4.) Pure gaseous phosphoretted hydrogen may be prepared by the action of potassium hydroxide upon phosplionium iodide— PH 4 I + KHO = KI + H 2 0 + PH 3 . Properties. — Gaseous phosphoretted hydrogen, or phosphine, is a colourless gas, having an offensive smell resembling rotting fish It is not spontaneously inflammable, but its ignition tem- perature is below ioo° C. (see page 291 ). The gas burns with a brightly luminous flame, producing water and metaphosphoric acid — PH 3 + 20 2 = HP0 3 + H 2 0. When burnt in oxygen the flame is extremely dazzling. The gas is not acted upon by oxygen at ordinary temperatures and pressures, but if a mixture of these gases be suddenly rarefied, * See “Chemical Lecture Experiments,” new ed., No. 545. 422 Inorganic Chemistry combination at once takes place with explosion. Phosphoretted hydrogen is decomposed by chlorine or bromine, a jet of the gas spontaneously igniting when introduced into chlorine, or the vapour of bromine — PH 3 + 4C1 2 = 3HC1 + PC1 6 . The gas is also decomposed by iodine, but in this case the action is less energetic, and a portion of the hydriodic acid which is produced, unites with the phosphine and forms phosphonium iodide, thus— (i.) PH 3 + 3I 2 = PI 3 + 3HI. ( 2 .) PH 3 + HI = PH 4 I. Phosphine is a highly poisonous gas, and the inhalation of even small quantities of it is attended with injurious effects. It is slightly soluble in water, and imparts its own smell and an un- pleasant taste to the liquid. The solution decomposes after a short time, especially in the light, and deposits red phosphorus. Phosphoretted hydrogen has no action upon either litmus or turmeric paper, but it resembles its nitrogen analogue, ammonia, in combining with hydrochloric, hydrobromic, and hydriodic acids’ forming respectively phosphonium chloride, bromide, and iodide Phosphonium Chloride, PH 4 C1.— When a mixture of phos- phine and gaseous hydrochloric acid is passed through a tube immersed in a freezing mixture, the gases unite and form a white crystalline incrustation upon the tube. If the iube be afterwards sealed up, the compound may be sublimed from one part of the tube to another, when it crystallises in large, brilliant, transparent cubes. If the tube be opened, the compound rapidly dissociates into its two generators. This compound may also be obtained by subjecting a mixture of the two gases to pressure. Under a pres- sure of about eighteen atmospheres at the ordinary temperature, crystals of phosphonium chloride are deposited ; and as the pres- sure is released the crystals gradually dissociate again. Phosphonium Bromide, PH 4 Br. — Phosphoretted hydrogen combines with hydrobromic acid at ordinary temperatures and pres- sures, but as the compound begins to dissociate at the ordinary temperature, the combination is only completely brought about by cooling the gases. Phosphonium bromide may be readily pre- pared by passing the two gases into a flask immersed in a mode- rate freezing mixture. The salt may be obtained in the form of Liquid Phosphoretted Hydrogen 423 large transparent cubical crystals by sublimation in a sealed ^Phosphonium Iodide, PH 4 I.-This compound may be obtained by a method similar to that given for the preparation of the ro mide. It is also produced when phosphoretted hydrogen is passed over iodine, as already mentioned. It is most readily propared by the action of water upon a mixture of phosphorus and iodine. Fo this purpose ten parts of phosphorus are dissolved in carbon . disul- phide in a tubulated retort, to which seventeen parts of ‘°^ ne are gradually added, the retort being kept cold. The carbon disul phide is then distilled off from a water-bath, a stream c ^° dioxide being passed through the apparatus towards the end of the distillation, to assist in expelling the last traces of the disul ph stx parts of water are then gradually introduced from a dropping funnel, when a brisk action takes place, and the phosphonium iodide produced is volatilised, and may be condensed m a i long wide glass tube connected to the retort. Hydnodic acid is at the same time formed — 91 + 13P + 24H 2 0 = 7PH,I + 2HI + 6H 3 P0 4 . The phosphonium iodide condenses in the form of brilliant ''“Liquid Phosphoretted Hydrogen, P S H 4 .-This compound is obtained in small quantities, when phosphorus is boded with a solution of caustic soda. It is obtained in large q^ntdies by he decomposition of calcium phosphide with water, by the react on already mentioned. In order to collect the compound, aq»ant.ty of calcium phosphide is introduced into a flask provided with a dropping funnel and exit tube. After displacing the air from he apparatus by an inert gas, water is gradually introduced from t the funnel, and the products of the reaction, after passing hrough a small empty tube, where water is arrested, are passed through a U-tube immersed in a freezing mixture, where the liquid phos- W”*'" » highly refracting, and spontaneously inflammable liquid. On ex- posure to light it is quickly decomposed into the gaseous and the solid hydrides of phosphorus 5P 2 H t = P 4 H, + 6PHj. 424 Inorganic Chemistry The formation of a spontaneously inflammable gas by the action in Th?3e C a r " aS f ° U " d 3 P-ticaUpphS in the marine appliance known as “ Holmes’ signal.” This con- sists of a tin canister filled with lumps of calcium phosphide A metal tube, Cosed at the bottom with a piece of block tin enters the caniste from be , ow> and a shor( con£ „ f ^ same fived 1 Up °" !, tOP ' When the si « rnal is to be use d> it is securely fixed into a wooden float, Fig. „8. The cone is cut off, and a hole punctured through the bottom of the tube below, and the apparatus thrown into the sea. The phosphoretted hydrogen rrsr-r and burns with a iarge briuiant flame tmm me top of the tin, illuminating a considerable area. Fig. h8. Solid Phosphoretted Hydrogen, P 4 H 2 (?)-The composition of this compound is not known with certainty. It is a yellow powder, obtained, as already mentioned, by the spontaneous decom- position of the liquid compound. COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS. Phosphorus combines with all the halogen elements, forming following compounds — the PF 3 pf 6 PC1 3 PBr 3 p Clfi PBr fi PI 3 . P 2 I 4 . Phosphorus Trifluoride, PF 3 , is obtained by the action of arsenic trmuoride upon phosphorus trichloride — AsF 3 + PC1 3 = PF 3 + AsC1 3 . It is more conveniently prepared by gently heating a mixture of zinc fluoride and phosphorus tribromide — 3ZnF a + 2PBr 3 = 2PF 3 + 3ZnBr 2 . Phosphorus Trichloride 4 2 5 Properties.— Phosphorus trifluoride is a colourless, pungent- smelling gas. It has no action upon glass in the cold, but when heated it forms silicon fluoride and phosphorus. It is moderately soluble in water. Phosphorus trifluoride unites directly with bromine, forming the compound PF 3 Br 2 . Phosphorus Pentafluoride, PF 6 .— This compound is formed when phosphorus burns in fluorine. It is best prepared by the action of arsenic trifluoride upon phosphorus pentachloride— 5AsF 3 + 3PC1 6 = 3PF 6 + 5AsC1 3 . Properties. — Phosphorus pentafluoride is a heavy, colourless gas, which fumes strongly in moist air, being decomposed by water into hydrofluoric and phosphoric acids— PF 6 + 4H 2 0 = 5HF + H 3 P0 4 . Owing to this decomposition it has a pungent and irritating effect upon the mucous membrane. It is not acted upon by oxygen, but unites directly with dry gaseous ammonia, forming a white solid compound having the composition 2PF 6 ,5NH 3 . . Phosphorus pentafluoride is an extremely stable compound, being capable of withstanding a very high temperature without dissocia- tion. On this account it is of special interest, as affording an example of a compound in which phosphorus is united to five monovalent atoms to form a stable substance. The corresponding chlorine and bromine compounds readily dissociate, when heated, into compounds containing trivalent phosphorus and the free halogen. . Phosphorus Trichloride, PC1 3 .— This compound is prepared by passing dry chlorine over red phosphorus, gently heated in a tubulated retort. The two elements readily combine, and the volatile trichloride, mixed with more or less of the pentachloride, distils off, and is collected in a well-cooled receiver. The product is freed from the higher chloride by redistillation over ordinary phosphorus. Properties. — Phosphorus trichloride is a colourless, mobile liquid, which boils at 75-95°- It has a pungent smell, and fumes strongly in moist air. Water at once decomposes it into hydro- chloric and phosphorous acids — PC1 3 + 3H a O = H 3 P0 3 + 3HC1. 4 2 ^ Inorganic Chemistry Phosphorus Pentachloride, PC1 6 . — This compound is formed when phosphorus burns in excess of chlorine. It is prepared by the action of chlorine upon the trichloride. Dry chlorine is passed on to the surface of a quantity of the trichloride, contained in a flask which is kept cool. The absorption of the chlorine is attended with considerable rise of temperature, and the contents of the flask rapidly become converted into a dry, pale yellow solid. Phosphorus pentachloride is conveniently obtained by passing chlorine through a solution of phosphorus in carbon disulphide, the solution being kept cold. Properties.— Phosphorus pentachloride is a yellowish- white, crystalline solid, having a pungent and irritating odour. It fumes strongly in contact with moist air, being decomposed by moisture into hydrochloric acid and phosphorus oxychloride — PC1 6 + H 2 0 = 2 HC 1 + POCl 3 . With excess of water, both phosphorus oxychloride and phos- phorus pentachloride dissolve with evolution of heat, forming hydrochloric and phosphoric acids — POCI 3 + 3H 2 0 = H 3 P0 4 + 3HC1. PC1 6 + 4H 2 0 = H 3 P0 4 + 5HC1. Phosphorus pentachloride readily sublimes, without melting, at a temperature below that of boiling water. It can only be melted by being heated under pressure, to a temperature of 148°. As the vapour of phosphorus pentachloride is heated, the com- pound dissociates into phosphorus trichloride and free chlorine. At 300 this dissociation is complete, and the vapour consists of equal molecules of the trichloride and chlorine. The gradual breaking down of the molecules of pentachloride, is seen from the following table, which gives the densities of the gas at different temperatures — Temperatures 182° 200° 250° 300° Density . . 72.5 69.2 57.0 51.95 At 300° it consists of molecules of PC1 3 (molecular weight = I 37-°7)> an d molecules of chlorine (molecular weight = 70.74), in equal numbers, which theoretically gives the molecular weight— 137.07 + 70.74 =51-95- 4 Phosphorus P cntachlonde 427 Phosphorus pentachloride is an important chemical reagent, in- asmuch as by ite action upon oxyacids, both inorganic and organic the'(HO) group in the acid can be replaced by chlorine. Thus with sulphuric acid, chlorosulphunc acid is formed HO J SOa + PCI, = } SO, + POCI 3 + HC1. With acetic acid it yields acetyl chloride C1 | CO + POCI 3 + HC1. HO l CO + PCI, = rH CH,i CH i It also effects the replacement of (HO) by chlorine in alcohols. Thus, with ethyl alcohol (spirits of wine) it forms ethyl chloride- Cl CH }CH< : / POCL + HC1. "pho^PenUbromide, PBr 5 . is prepared by adding broaiine to the b DmhosDhorus Tetriodide (phosphorus ii-ioiide), P,I, -This substance is acids. j’ OXY AND THIO COMPOUNDS OF PHOSPHORUS AND THE HALOGENS. The following compounds are known, containing phosphorus combined with the halogens, and either oxygen or sulphur- POF, ; POCI 3 ; P 2 0 3 C1 4 ; POBrClj ; POBr 3 . PSF 3 ; PSCI 3 ; — — PSBr 3 . These compounds may be regarded as derived from the haloid compounds, by the replacement of two atoms of the halogen by an 428 / norgamc Chemistry equivalent of oxygen, or of divalent sulphur • or thev u ■ t-oF r %c?i r r s r M •• ^ S S a derivative of ' 'I , 0mp0Und > P ^ C, <’ ma y be viewed as a denvative of pyrophosphonc acid, P 2 0 3 (H0) 4 . Phosphoryl Fluoride ( phosphorus oxyfl,uoride\ pdf „ , , the action of phosphoryl chloride P 2 0 3 (HO) 4 l hpo 3 „ P0 2 (H0) J p A p 2 o 6 weigh?- h ° , c ^ d {phosphorous anhydride , ), P 4 O 0 ; molecular weight - 221.52. -This oxide is obtained, mixed with a large excess Of the pentoxide, when phosphorus is burned in a tube through Fig. 1 19. which a regulated stream of air is passed. In order to obtain the compound in a state of purity, the following method is employed. A quantity of phosphorus is introduced into a glass tube, bent in the manner indicated in Fig. 119, and fitted into one end of a long wide, Liebig's condenser. Into the end of the condenser nearest to the U-tube, there is introduced a loose plug of glass wool, which serves to arrest the pentoxide. The phosphorus is ignited at the open end of the glass tube, and a stream of air drawn through the apparatus by means of an aspirator. A stream of water, at 60 , is circulated through the condenser, when the easily fusible phosphorous oxide is carried over, and condenses in the' U-tube. which is immersed in a freezing mixture. P Jiosphor us Pentoxide 43 1 Properties. — Phosphorous oxide, as it collects in the cooled tube, is a snow-white crystalline solid, which melts at 22.5 to a colourless liquid. The liquid solidifies at 2 1 0 to a white, waxy-looking mass, consisting of monoclinic prismatic crystals. The liquid boils at 1 73.1 °. It possesses an unpleasant garlic smell, and is highly poisonous. Phosphorous oxide is only very slowly acted upon by cold water, which gradually dissolves it, forming phosphorous acid — P 4 O e + 6 H 2 0 = 4H3PO3. In contact with hot water a violent action takes place, in which spontaneously inflammable phosphoretted hy- drogen is evolved, and a red deposit, consisting of red phosphorus and phosphorus monoxide, is formed. When heated in a sealed tube to a tem- perature of 440°, phosphorous oxide is de- composed into phosphorus tetroxide, and red phosphorus — 2 P 4 0 6 = 3 P 2 0 4 + 2P. , When exposed to air or oxygen, phosphorous oxide is gradually oxidised into phosphorus pentoxide, but when placed in warm oxygen it bursts into flame. When brought into chlorine it also spon- taneously inflames. Phosphorus Tetroxide, P 2 0 4 .— This substance is obtained when phos- phorous oxide is heated in a sealed tube to a temperature of 440°. It forms brilliant transparent crystals, which appear as a sublimate in the tube. This oxide is highly deliquescent, and dissolves in water with evolution of heat. Phosphorus Pentoxide, P 2 0 6 ; molecular weight = 141.82.— This oxide is the main product of the combustion of phosphorus in air or oxygen. It may readily be obtained, by burning a quantity of phosphorus in a small capsule, and covering the whole with a large bell-jar, Fig. 120. The white clouds of phosphorus pent- oxide collect as a soft snow-like substance. Properties. — Phosphorus pentoxide is a white, amorphous, and very voluminous powder. It is without smell, although as usually prepared it frequently possesses a slight garlic odour, owing to the presence of phosphorous oxide. 432 Inorganic Chemistry sea, d vessels. ful desiccating agent known to chemists : prolonged exposure to fro7ga°se, C . Pen '° Xide ’ rem ° VeS * he ' aSt ° f aqueou^" When thrown into water, phosphorus pentoxide is dissolved with LtapZp\°o U ^acid- b,ing thC qUenChing ° f h0t iron ’ and P A + H 2 0 = 2HP0 3 , which gradually passes into orthophosphoric acid— HPO, + H 2 0 = H 3 P0 4 . Phosphorus pentoxide is capable of abstracting the elements oi water from a number of substances, both inorganic andTrganic thus, it converts nitric acid into nitrogen pentoxide— ? 2HN0 3 - H 2 0 = N 2 0 6 . In the same way it withdraws the elements of water from alcohol with the evolution of ethylene— alcohol, c 2 h 6 o - H 2 0 = C 2 H 4 '. Hypophosphorous Acid, H 3 P0 2 .-This acid is prepared by the action of sulphuric acid upon the barium salt y Ba(H 2 P0 2 ) 2 + H 2 S0 4 = BaS0 4 + 2H 3 P0 2 . The solution, after the removal of the barium sulphate by filtra- rion, is gently heated until its temperature rises J , 3 o-, whenlt to'o'be SUffiC ’ entIy COncentra ‘ ed *0 deposit crystals, when cooled The barium hypophosphite is obtained by boiling phosphorus with a solution of barium hydroxide— g pnosphorus 3Ba(HO) 2 + 8P + 6H 2 0 = 2PH S + 3Ba(H 2 P0 2 ) 2 . Properties, Hypophosphorous acid is a white crystalline com pound, which melts at , 7 . 4 ». When strongly heated, ^converiTd Phosphorous Acid 433 into orthophosphoric acid with the evolution of gaseous phos- phoretted hydrogen— 2H 3 P0 2 = H 3 PO, + PH 3 . Hypophosphorous acid acts as a powerful reducing agent, on account of the readiness with which it absorbs oxygen and is con- verted into orthophosphoric acid — H 'i HOI H 1 HCn H IPO + 0 2 = HO IpO ; or HO VP + 0 2 = HO VPO. HOj HO J HOj HOj Hypophosphorous acid is a feeble monobasic acid ; its salts with monovalent metals being represented by the formula MH 2 P0 2 . It is customary to express the basicity of oxyacids by the number of (HO) groups that are contained in the molecule, and as this acid is monobasic its constitution would be expressed by the formula POH 2 (HO). Many of the oxvacids of phosphorus, however, show a tendency to exhibit a lower degree of basicity than is represented by the number of (HO) groups they contain; thus, orthophosphoric acid, PO(HO) 3 , which is tribasic, and forms he salt trisodium phosphate, PO(NaO) 3 , holds the third atom of the metal so loosely, that even such a feeble acid as carbonic acid is capable of expelling it or — PO(NaO) 3 + C0 2 + H 2 0 = PO(HO)(NaO) 2 + HNaCO s . Na 3 P0 4 + C0 2 + H 2 0 = HNa2P0 4 + HNaC0 3 . The weaker acid, phosphorous acid, is also tribasic, P(HO) 3 , and forms trisodium phosphite, P(NaO) 3 , or Na 3 P0 3 . But this salt .s even decomposed by water, into the disodium phosphite, P(HO)(NaO) 2 , or HNasP0 3 . Hypophosphorous add being a still weaker acid, its acidic power is destroyed as soon as one atom of hydrogen is replaced by a base, and its constitution may, in harmony with these facts, be expressed by the formula PH(HO) 2 , or HO | ,P. Phosphorous Acid, H 3 P0 3 , or P(HO) 3 .-As already mentioned this acid is formed when phosphorous oxide is dissolved in cold water. , It is most readily prepared by the action of water upon p os- phorus trichloride — PC1 3 + 3H 2 0 = 3HC1 + P(HO) 3 . The production and decomposition of the phosphorus trichloride 434 Inorganic Chemistry IhroCt T ie t° n Sim , Ul 'f nC0USly ' by P assi "g a stream of chlorine through phosphorus which is melted beneath water. The solution ■s evaporated until its temperature rises to , 80 *, when the liouid ,h “ "■ - »««»»» ohosphoric acid, with evolution of phosphoretted hydrogen- 4 H 3 P0 3 = 3H 3 PO, + PH 3 . Like hypophosphorous acid this compound absorbs oxygen and therefore ,s a powerful reducing agent; silver salts are reduced salts Bv theTh andmercur,c salts ar « reduced to mercurous phosphoric ° f ° Xygen ft " C ° nverted in ‘° h 3 po 3 + o = h 3 po 4 . Although a tribasic acid, its tribasic salts are unstable ; the sodium compound, Na 3 P 0 3 , which is the most stable inorganic salt, is decomposed by water into the dibasic salt Na 3 P 0 3 + H 2 0 = HNa 2 P 0 3 + NaHO. NaCri HO -I NaO f-P + H 2 0 = NaO } P + NaHO NaO-t NaoJ fnn"d ^ P u ACld> HaPO o ° r P °(HO) 3 .-This acid is formed when phosphorus pentoxide is dissolved in boiling water or when the solution of the oxide in cold water is boiled- P 2 0 6 + 3H 2 0 = 2H 3 P0 4 . Orthophosphoric acid is readily obtained by the oxidation of red phosphorus with nitric acid. Copious red fumes are evolved, and phosphoric acid remains in solution. Phosphoric acid is prepared on a large scale by the action of ashi as in the process for the — Ca 3 (P 04) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 P0 4 . The calcium sulphate is removed by filtration, and the solution 435 Pyrophosphoric Acid evaporated to a syrup. Prepared in this way the acid usually contains arsenic. This is removed by first reducing it to arsemous oxide, by means of sulphur dioxide, and after boiling off the excess of sulphur dioxide, precipitating the arsenic as sulphide by means of sulphuretted hydrogen. Properties. — The solution obtained by these methods is either concentrated in vacuo, or heated to a temperature of 140°, and allowed to cool, when the acid is obtained in the form of trans- parent six-sided prismatic crystals, belonging to the rtiombic system. The substance is deliquescent, and melts at 38.6 . Phosphoric acid is tribasic, and forms three series of salts, according as one, two, or three of its hydrogen atoms are replaced by an equivalent of metal. Thus, with the metal sodium the three salts are known— Trisodium phosphate (normal sodium phosphate) Na 3 P0 4 . Hydrogen disodium phosphate Dihydrogen sodium phosphate . • • • 2 a The hydrogen may be replaced by its equivalent of more than one base. Thus, the well-known compound, microcosmic salt, is hydrogen sodium ammonium phosphate, HNa(NH 4 )P0 4 . The salt which is precipitated when magnesium sulphate (in the pre- sence of ammonium chloride and ammonia) is added to a solution of a phosphate, consists of the compound ammonium magnesium phosphate (NH 4 )MgP0 4 . The heavy metals usually only form normal phosphates. Thus, on the addition of silver nitrate to a solution of either of the three sodium salts, the same silver salt is precipitated, namely, tn- argentic phosphate. Na 3 P0 4 + SAgNO, - Ag s P0 4 + 3NaNOj. * HNa 2 P0 4 + 3AgNO s = Ag 3 P0 4 + 2NaN0 3 + HN0 3 . H 2 NaPO, + 3AgNO a a cr p n Pyrophosphoric Acid, H 4 P 2 0 7 , or P 2 0 3 (HO) 4 . This acid is * Hydrogen disodium phosphate, although belonging to that class of com- pounds commonly called acid salts, on account of the fact that it still retains a portion of the replaceable hydrogen of the acid, is strongly alkaline in its action upon litmus : silver nitrate is a neutral compound, hence in this reaction, by mixing an alkaline and a neutral liquid, an acid liquid is obtained, on account of the molecule of nitric acid that is set free. 43^ Inorganic Chemistry derived from orthophosphoric acid, by the withdrawal of one molecule of water from two molecules of the acid. This change is effected by heating the ortho acid to 213 0 — 2H 3 P0 4 ^ H 2 0 = H 4 P 2 0 7 . The formation of this acid from two molecules of orthophos- phoric acid will be made clear by the following formulae— HO^HO HO HO HO HO HO HO ° = P- jo-H h!-o-VC° = h 2 o + o^p^o^p1o. Pyrophosphates are formed when monohydrogen orthophos- phates are heated. Thus, by heating hydrogen disodium ortho- phosphate, sodium pyrophosphate is formed — 2HNa 2 P0 4 - H 2 0 = Na 4 P 2 0 7 . When ammonium magnesium phosphate (see above) is heated in the same way, it loses water and ammonia, and is transformed into magnesium pyrophosphate, thus — 2(NH 4 )MgP0 4 = Mg 2 P 2 0 7 + H 2 0 + 2 NH 3 . Properties. — Pyrophosphoric acid is an opaque white crystal- line solid, readily soluble in water. Its aqueous solution passes slowly into orthophosphoric acid, the change taking place rapidly on boiling ; a solution of this acid therefore cannot be concen- trated by boiling. The pyrophosphates are stable salts, and their solutions may be boiled without change ; by boiling with acids, however, they are converted into orthophosphates. Metaphosphoric Acid, HP0 3 or P0 2 (H0).— This acid is formed when phosphorus pentoxide is allowed to deliquesce. It may be obtained by the abstraction of one molecule of water from one molecule of orthophosphoric acid, which is brought about by heat- ing the tribasic acid to redness — H 3 P0 4 - H 2 0 = HPOg. It is also obtained by strongly heating pyrophosphoric acid — H 4 P,O r - H a O = 2HP0 3 . Metaphosphoric Acid 437 7* f £;;• ttn? rss? ss > S ™>>hW - «■”«» sodium pyrophosphate H 2 Na 2 P 2 0 7 . transDar ent vitreous Properties. — Metaphosphoric acid s a transpare solid (frequently termed glacial phosphoric acid). It 7 fusible, and is usually casf into sticks. At a ^^ ‘e-pera ure mav be volatilised. Metaphosphoric acid is easily soluble in 21, Ld its solution is slowly transformed into otthophosphoric acid- this change takes place rapidly on boiling, and the acid passes directly into the tribasic acid without the intermediate formation of pyrophosphoric acid HP0 3 + h 2 o = H 3 P0 4 . Metaphosphoric acid is monobasic, but it possesses the remark- able property of forming a number of salts which may be rega as derived from several hypothetical polymeric varieties of the acid. Monometaphosphoric acid, HP0 3 , forms ^ 0 "7 e “ ph “P'’ a, “'p N o P03 ' Dimetaphosphoricacid,(HP0 3 ) 2 , dimetaphosphates W, Trimetaphosphoric aciU (HP^b. ;; Hexametaphosphoric acid! (HPCwI. . . hexametaphosphates, Na 6 P 6 0 18 . The three compounds, ortho-, pyro-, and metaphosphoric acids, are readily distinguished from each other by means of silver nitrate, and their action upon albumen Reagent. 1 Orthophosphoric Acid. Pyrophosphoric Acid. Metaphosphoric Acid. Silver nitrate . Albumen . . Canary yellow precipitate of Ag 3 P0 4 No action White crystalline precipitate of Ag 4 P 2 0 7 No action White gelatinous precipitate of AgPO s Coagulates j Orthophosphoric acid is also distinguished by giving a yellow precipitate of ammonium phospho-molybdate, upon the addition o excess of a solution of ammonium molybdate in nitric acid (see Molybdenum). 438 Inorganic Chemistry Compounds of Phosphorus containing: Nitrogen. -By the action of am- monia upon phosphorus pentachlonde, and upon phosphoryl chloride (POCU a number of nitrogen derivatives are obtained. Thus, when gaseous ammonia is passed over phosphorus pentachloride, and the solid mass so obtained is heated in a stream of an inert gas, until the ammonium chloride is driven ofT, a white insoluble powder remains, having the composition represented by the formula PN(NH), to which the name phospham has been given— PCI 5 + 7NH 3 = 5NH 4 C1 + PN(NH). Phosphoryl Triamide, PO(NH 2 ) 3 , is obtained by the action of gaseous ammonia upon phosphoryl chloride— 6 POCI 3 + 6 NH S = PO(NH 2 ) 3 + 3NH 4 C1. 3n ^, en . h , eate . d °; U ° f c ° ntact with air, phosphoryl triamide yields ammonia and phosphoryl nitride , thus — p O(NfT 2 ) 3 = 2NH 3 + PON. Pyrophosphamic Acids.— Three of may be regarded as pyrophosphoric at groups have been replaced respectively 1 Pyrophosphoric acid Pyrophosphamic acid . Pyrophosphodiamic acid Pyrophosphotriamic acid these compounds are known, which id, in which 1 , 2 , and 3 of the (HO) the group (NH 2 ), thus— • • P 2 0 3 (HO) 4 . . . P 2 0 3 (HO) 3 (NH 2 ). • . P 2 0 3 (HO) 2 (NH 2 ) 2 . • . P 2 0 3 (HO)(NH 2 ) 3 . Compounds of Phosphorus with Sulphur.-A number of compounds of phosphorus and sulphur have been obtained by heating together varying pro- known"-^ SU PhUr and rCd P hos P horus - The following compounds are Phosphorus monosulphide . Phosphorus sesquisulphide Phosphorus trisulphide Phosphorus tetrasulphide (?) Phosphorus pentasulphide . Analogous Oxides. P 4 S . . p 4 o. P 4 S 3 . P 2 S 3 . . P 4 0 G . P2S4 . • p 2 o 4 . P2S5 . . p 2 o 5 . Phosphorus Pentasulphide, PA.-This compound is the best-known member of the series. It is prepared by gently heating red phosphorus and fragments of sulphur, in the proportion required by the formula, in a flask. The elements combine with energy, and on cooling, a solid mass is obtained. This solid material is then distilled in a current of carbon dioxide, when the pentasulphide is obtained in the form of yellow crystals. The compound may also be obtained by dissolving ordinary phosphorus and sulphur in the proper proportions in carbon disulphide, and heating the solution in sealed tubes to 210 . On allowing the solution to cool, yellow crystals of the pentasulphide are deposited. Phosphorus pentasulphide is decomposed by water with the formation of orthophosphoric acid and the evolution of sulphuretted hydrogen— PsSj + 8 H 5 O = 2H 3 P0 4 + 5HgS. Arsenic 439 ARSENIC. Symbol, As. Atomic weight = 74-9* Molecular weight = 299.6. Vapour density = 149.8. Occurrence.— Arsenic is found in the free state in nature, usually in the form of small nodules, more rarely as distinct crystals. In combination with sulphur it constitutes the minerals realgar , or ruby sulphur, As 2 S 2 ; and orpiment, As 2 S 3 . In combination with metals, as arsenides, it occurs widely distributed, the commonest ores being arsenical iron , FeAs 2 and Fe 4 As 3 ; kupfermckel , NiAs and NiAs 2 ; and tin white cobalt ', CoAs 2 . With metals and sulphur, it is met with in such minerals as arsenical pyrites , mispickel , ox white mundic , FeS 2 ,FeAs 2 ; cobalt glance, CoS 2 ,CoAs 2 ; nickel glance, NiS 2 , NiAs 2 . Arsenic is present in small quantities in most samples of iron pyrites, hence it finds its way into sulphuric acid manufactured from pyrites. It also occurs in coal smoke, being derived from the pyrites contained in coal, and hence is present in the atmosphere : during the prevalence of yellow fogs the amount of arsenic present is very appreciable. Modes of Formation.— On the small scale, arsenic is obtained by heating a mixture of arsenious oxide, As 4 O e , with powdered charcoal — As 4 0 6 + 6C = 6CO + 4 As. On a larger scale it is usually obtained either from native arsenic or from arsenical pyrites ; the latter substance, when heated, gives up arsenic, and ferrous sulphide is left behind— FeS 2 ,FeAs 2 = 2 As + 2 FeS. The mineral is heated in long narrow horizontal earthenware retorts, into whose mouths are fitted earthenware receivers. The arsenic volatilises, and condenses in these receivers as a compact crystalline solid. It is purified by redistillation. Properties. — Arsenic which has been resublimed, is a brilliant steel -grey metallic-looking substance, forming hexagonal rhombo- hedral crystals, having a specific gravity of 5.62 to 5.96. It is very brittle, and is a good conductor of heat and electricity. Arsenic begins to volatilise at ioo°, and rapidly vaporises at a dark- red heat, passing from the solid to the vaporous states without 440 Inorganic Chemistry liquefying The vapour has a yellow colour and an unpleasant garlic smell. When heated under pressure, arsenic melts at 500“ and on cooling, solidifies to a compact crystalline mass When arsenic is vaporised in a glass tube, in a current of hydrogen, it condenses along the tube in three distinct conditions : that which is deposited nearest to the heated portion of the tube is in the form of rhombohedral crystals ; that which sublimes a little farther along, and condenses at a point where the temperature is about 2 10-220 , consists of a black shining amorphous deposit, while at a still more distant and cooler portion of the tube, a grey crystalline sublimate is formed. These are regarded as allotropic modifica- tions of arsenic. The amorphous variety is also formed, when arsenuretted hydrogen is decomposed by being passed through a heated tube (g.v.). Amorphous arsenic is unacted upon by air at ordinary temperatures, and only slightly oxidised at 8o°. The m-ev crystalline variety is readily oxidised on exposure to air at ordinarv temperatures. , y Amorphous arsenic, when heated out of contact with air to 360° is converted into the rhombohedral variety. Arsenic, like phosphorus, forms tetratomic molecules, its molecular weight as deduced from its vapour density being 74.9 x 4 = 299.6. When heated in oxygen, arsenic burns with a bright bluish-white flame, forming arsenious oxide, As 4 O e . It is oxidised by sulphuric acid, nitric acid, and other oxidising agents. It combines readily with chlorine, and when thrown into this gas in the condition of powder it spontaneously inflames, forming arsenic trichloride. Thrown into bromine, a fragment of arsenic spontaneously in- flames, and burns as it floats about upon the surface of the liquid. Arsenic, in many of its characteristics, resembles the true metals • it is one of those elements lying on the borderland between true metals and non-rftetals, to which the name metalloid is applied. It is capable of forming alloys with metals, and an alloy of this element with lead, is employed for the manufacture of shot. It is found that by the addition of a small proportion of arsenic to lead, the melted metal is more fluid, and therefore more readily assumes the spheroidal form when projected from the shot tower , and on solidification the alloy is considerably harder than pure lead. Arsenuretted Hydroge?i 441 ARSENURETTED HYDROGEN ( Arsenic trihydride. Arsine). Formula, AsH 3 . Molecular weight = 77.9- Density = 38.95. Modes of Formation.— ( 1.) Arsenuretted hydrogen is formed .when soluble arsenic compounds are exposed to the action of nascent hydrogen : thus, when a solution of arsenious oxide is introduced into a mixture from which hydrogen is being generated, such as zinc or iron and dilute hydrochloric or sulphuric acid, arsenic trihydride is obtained, mixed with free hydrogen— As 4 0 6 + 12H 2 = 4 AsH 3 + 6H 2 0. (2.) By the same action of nascent hydrogen, arsenic trihydride is formed when a solution of either arsenious oxide, As 4 O e , or arsenic oxide, As 2 0 6 , is subjected to electrolysis. (3 ) Arsenic trihydride is also formed, when arsenical compounds are in contact with organic matter which is undergoing decom- position During the growth of certain moulds and fungi a small quantity of hydrogen is evolved, which by its action upon the arsenic compound, gives rise to the formation of arsenic trihydride. By this action arsenic trihydride is sometimes formed in dwelling- houses where arsenical wall-papers are employed, and where, from dampness or other causes, mould develops. (4.) Pure arsenic trihydride is prepared by the action of dilute hydrochloric or sulphuric acid upon an alloy of arsenic and zinc As 2 Zn 3 + 3H 2 S0 4 = 2AsH 3 4- 3ZnS0 4 , or by the action of either water or dilute acid upon an alloy of arsenic and sodium, prepared by heating sodium in the impure arsenic trihydride obtained by method No. 1. Properties. — Arsenic trihydride is a colourless, offensive-smell- ing, and highly poisonous gas. Under pressure it condenses to a colourless liquid, which boils at -54.8° and solidifies at -113.5 • The gas burns with a lilac-coloured flame, forming water and white fumes of arsenious oxide— 4AsH 3 + 60 2 = As 4 O s + 6H 2 0. When the supply of air to the flame is limited, as when a cold 44 2 Inorganic Chemistry surface is depressed upon it, water is formed and arsenic is deposited as a shining black amorphous film — 4AsH 3 + 30 2 = As 4 + 6H 2 0. Arsenuretted hydrogen is readily decomposed by heat into its elements : thus, when the gas is passed through a glass tube, which is heated at one point by a Bunsen flame, arsenic in the amor- phous condition is deposited upon the tube immediately beyond the heated spot Even when greatly diluted with hydrogen this reaction takes place, and it therefore affords a delicate test for the presence of exceedingly small quantities of arsenic. This method for the detection of arsenical compounds is known as Marsh's test, and may be carried out by means of the apparatus seen in Fig.’ 12 1 . Hydrogen is generated in the two-necked bottle from zinc and dilute sulphuric acid (which are themselves free from arsenic), and the arsenic in the form of an oxygen or a haloid compound is intro- duced. On igniting the issuing gas, and depres- sing a white porcelain cap- sule into the flame, black stains of amorphous arsenic are produced : and if the tube be heated as shown in the figure the arsenic is deposited as a black film. The corresponding anti- mony compound, SbH 3 (q.v.\ gives rise to a similar deposit of metallic antimony, when treated in the same way ; but the arsenic deposit is readily distinguished by being easily soluble in a solu- tion of calcium hypochlorite. Many metals, such as sodium, or potassium, when heated in arsenuretted hydrogen, form alloys with the arsenic, and hydrogen is set at liberty ; while metallic oxides when similarly treated form metallic arsenides and water. Arsenuretted hydrogen is slightly soluble in water, but the solu- tion on exposure to air deposits arsenic When passed into a solution of silver nitrate, metallic silver is precipitated, and a solution of arsenious oxide (the hypothetical arsenious acid, H 3 AsO a ) is obtained, thus — AsH 3 + 6AgNO a + 3H 2 0 = 3Ag 2 + 6HNO, + H 3 AsO s . A rsenic Chloride 443 When the gas is passed into copper sulphate solution, cuprous arsenide is precipitated — 2AsH 3 + 3CuS0 4 = 3H 2 S0 4 + As 2 Cu 3 . Arsenuretted hydrogen is decomposed by the halogens with energy, forming the haloid compound of arsenic, and the halogen acid — AsH 3 + 3C1 2 = AsC1 3 + 3HC1. Solid Arsenuretted Hydrogen.— When arsenide of potassium or sodium is acted upon by water, a soft brown solid substance separates, which contains equal atomic proportions of arsenic and hydrogen. Its molecular weight is unknown, its composition is therefore expressed by the formula, (AsH) n . COMPOUNDS OF ARSENIC WITH THE HALOGENS. The following compounds are known — AsF 3 ; AsC1 3 ; AsBr 3 ; Asl 3 . Two other compounds with iodine have been described, contain- ing the elements in the proportion represented by the formulae, AsI 2 and As 2 I 5 , the molecular weights of which are unknown. Arsenic Fluoride, AsF 3 ; molecular weight = 13 1.9, is formed when sodium fluoride is heated with arsenic chloride— 3NaF + AsC1 3 = 3NaCl + AsF 3 . It is best obtained by distilling a mixture of arsenious oxide, powdered fluor spar, and sulphuric acid in a leaden retort. The hydrofluoric acid generated by the action of the acid upon the calcium fluoride, reacts upon the arsenious oxide, thus — As 4 O s + 12HF = 4AsF 3 + 6H 2 0. Properties. — Arsenic fluoride is a colourless fuming liquid, boiling at 60.4°. It is rapidly decomposed by water into arsenious oxide and hydrofluoric acid. On this account it forms painful wounds when brought into contact with the skin. Arsenic Chloride, AsCl 3 ; molecular weight = i8i.it, is ob- tained when arsenic burns in chlorine, or when chlorine is passed over fragments of arsenic in a tube. 444 Inorganic Chemistry It is also produced when either arsenic or arsenious sulphide is distilled with mercuric chloride — 2As + 0HgCl 2 = 3Hg 2 Cl 2 + 2AsC1 3 . As 2 S 3 + 3HgCl 2 = 3HgS + 2AsC1 3 . It is readily prepared by the action of hydrochloric acid upon arsenious oxide ; for which purpose sodium chloride, arsenious oxide, and sulphuric acid are gently heated together in a retort connected with a well-cooled receiver— As 4 O fl + 12HC1 = 4AsC1 3 + 6H 2 0. Properties.— Arsenic chloride is a colourless, fuming, and some- what oily liquid, which boils at 130 . 2 0 , and is extremely poisonous. In the presence of excess of water, or when added to warm water, it is decomposed into arsenious oxide and hydrochloric acid. With a small quantity of water a solid crystalline arsenic chlor- hydroxide is formed, As(HO) 2 C1 — AsC1 3 + 2H a O = 2HC1 + As(HO) 2 C1. Arsenious Bromide, AsBr 3 . — This compound is formed by the direct union of arsenic with bromine, and is prepared by adding powdered arsenic to a solution of bromine in carbon disulphide. On evaporation, the compound is deposited in the form of colourless deliquescent crystals, which melt at 20 ° to 25 0 to a straw-coloured liquid. Arsenious Iodide, Asl 3 , is obtained by heating a mixture of arsenic and iodine. It is most conveniently prepared by -digesting a saturated ethereal solution of iodine, with powdered arsenic, in a flask with a reflux condenser. On filtering and cooling, the iodide deposits in the form of lustrous red hexa- gonal crystals. OXIDES AND OXYACIDS OF ARSENIC. Two oxides of arsenic are known, both of which act as anhy- drides — Arsenious oxide As 4 O b . Arsenic oxide (arsenic pentoxide) . . As 2 0 6 . No acid corresponding to arsenious oxide is known in the free state, although the arsenites constitute a class of stable salts. Arsenious Oxide 445 Three arsenic acids, derived from arsenic pentoxide, are known, analogous in constitution to the three phosphoric acids, namely— Ortho-arsenic acid . • H 3 As 0 4 or AsO( H 0 )g. Pyro-arsenic acid . . • H 4 As 2 0 7 or As 2 0 3 (HO) 4 . Metarsenic acid . • • HAsOg or As 0 2 (HO). ARSENIOUS OXIDE. Formula, As 4 0 6 . Molecular weight = 395- 3&- Mode of Formation. — Arsenious oxide is formed when arsenic burns in air or in oxygen, or when arsenic minerals are -roasted in a current of air. On a small scale it may be produced by burning arsenic in a hard glass tube in a stream of oxygen, and allowing the white fumes of arsenious oxide to pass into a glass cylinder (as shown in Fig. 122), where the greater part condenses, while the rest is led into a draught flue. Arsenious oxide is obtained as a secondary product, m the metallurgical process of roasting arsenical ores of nickel, cobalt, tin silver, and others, for the extraction of these metals. It is also obtained as a principal product by roasting arsenical pyrites. The ore is heated either upon the hearth of a rever- beratory furnace, where it is raked over from time to time, or it is introduced by means of a hopper , , into one end of a long clay- lined iron cylinder, placed at an incline of about 1 in 18, and caused slowly to revolve about its longitudinal axis, Fig. 123. The lower end of this cylinder enters a furnace, the upper end is connected to 446 Inorganic Chemistry series of brickwork flues. The ore is delivered into the upper end of the revolving cylinder, and as it gradually gravitates down the incline, it is com- pletely roasted by the furnace flames which pass over it, and finally falls out into a chamber beneath. The fumes of arseni- ous oxide pass through a series of chambers or flues, so ar- ranged as to present an exten- sive condensing surface to the gases, and the crude product, known as arsenical soot , is from time to time collected. This is known as Oxland and Hock- ing’s revolving calciner. Prop epties. — Arsenious oxide, known familiarly as white arsenic , or simply arsenic , is p? known in three modifications (i.) Amorphous. ( 2 .) Octahedral crystals of the regular system. (3-) Prismatic crystals of the trimetric system. Amorphous Arsenious Oxide is a colourless, transparent, vit- reous substance, which is ob- tained when the vapour of the oxide is condensed at a tem- perature only slightly below its vaporising point. On exposure, it gradually becomes opaque, being transformed into the regu- lar octahedral variety. This change takes place from the- outside, and lumps of opaque “white arsenic,” when broken. A rsenites 447 often show a nucleus of the vitreous modification. Amorphous arsenious oxide may be preserved unchanged in sealed tubes. The change from the vitreous to the crystalline form is attended with evolution of heat, and a diminution of specific gravity from 3.738 to 3.689. Amorphous arsenious oxide, when heiated to about 200 , melts, and at a higher temperature vaporises. It is soluble in 108 parts of cold water. Octahedral Arsenious Oxide.— The vitreous form passes spon- taneously into this variety. It is obtained directly, by quickly cooling the vapour of arsenious oxide, or by crystallisation from the aqueous solution of either form of the oxide. Arsenious oxide is also deposited in this form from solution in hydrochloric acid. Octahedral arsenious oxide is less soluble in water than the amor- phous variety, 1 part requiring 355 parts of water for its solution When heated, the crystals vaporise without fusion, but when heated under pressure they melt, and are converted into the vitreous form. Prismatic Arsenious Oxide is obtained by crystallisation from a hot saturated solution of arsenious oxide in potassium hydroxide. Aqueous solutions of arsenious oxide possess a feeble acid re- action, probably due to the formation of unstable arsenious acid. H 3 As 0 3 . The acid has not been isolated, and on concentration the solution deposits crystals of arsenious oxide. Arsenious oxide is a powerful poison : from 2 to 4 grains usually prove fatal. It is possible, however, by the habitual use of it, to so accustom the system to this poison, that doses sufficiently large to cause certain death to one unused to it, may be taken with apparent impunity. The use of arsenic is said to beautify the complexion, and to improve the wind. The men who are em- ployed upon arsenic works, are constantly liable to swallow doses of arsenious oxide which would cause death to one unaccustomed to the occupation. Arsenites. — Three classes of arsenites are known, which may be regarded as being derived from the three hypothetical acids— r Silver ortho-arsenite, Ag 3 AsO a . J Hydrogen copper 'j Ortho-arsenious acid, H 3 As0 3 , or As(HO) 3 -j arsenite (Scheele’s VHCuAs0 3 . I green), J r Calcium pyro-arsenite, Ca 2 As 2 0 5 . I Barium pyro-arsenite, Ba 2 As 2 0 5 . Pyro-arsenious acid, H 4 As 2 0». or As 2 0(H0) 4 a Ammonium \ (NH 4 ) 4 As 2 0 6 . i. pyro-arsenite, 1 44» Inorganic Chemistry { Potassium metarsenite, KAs0 2 . AC metar taSS ’t im } KAs °2.HAsOa. Lead metarsenite, Pb(As0 2 ) 2 . The pigment known as Schweinfurt green , is a double metar- senite and acetate of copper — 3Cu(As0 2 ) 2 ,Cu(C 2 H 3 0 2 ) 2 . All arsenites, except those of the alkali metals, are insoluble in water. When heated, most arsenites are converted into arsenates and arsenic ; and when heated with charcoal the whole of the arsenic is reduced. Arsenic Pentoxide, As 2 0 6 .— This oxide is not formed when arsenic burns in oxygen. It is obtained by the oxidation of arseni- ous oxide by nitric acid, and subsequently heating the arsenic acid so produced, to a dark-red heat— 2H 3 As0 4 = 3H 2 0 + As 2 0 6 . Properties. Arsenic pentoxide is a white deliquescent solid completely soluble in water, with the formation of arsenic acid ’ When strongly heated it breaks up into arsenious oxide and oxygen — 2As 2 0 6 = As 4 0 6 + 20 2 . ARSENIC ACIDS AND ARSENATES. When arsenic pentoxide is dissolved in water, crystals are ob- tained having the composition 2As0(H0) 3 ,II 2 0. At ioo° these melt and lose water, leaving ortho-arsenic acid, H 3 As0 4 . By the withdrawal of water from this acid, both pyro- and metarsenic acids are obtained. Heated between 140 ° and 180 °, two molecules of the “ortho” acid 16se one of water — 2H 3 As0 4 = H 4 As 2 0 7 + H 2 0. And by heating the pyro-arsenic acid so dbtained to 200 °, another molecule of water is expelled, with the formation of metarsenic acid ( compare corresponding acids of phosphorus) H 4 As 2 0 7 = 2HAs0 3 + H 2 0. Pyro- and metarsenic acids are both crystalline solids, which Arsenic Disulphide 449 dissolve in water with the evolution of heat and formation of ortho- lrsenic acid; aqueous solutions of these two acids, therefore, cannot exist. In this respect they differ from the' corresponding phosphorus acids, both of which can be obtained in aqueous solution. „ , . , _ r 11 Each of the three arsenic acids forms salts, of which the following are examples : — Trisodium ortho-arsenate . • • Na 3 As0 4 . Hydrogen disodium ortho-arsenate ■ HNa 2 As0 4 . Dihydrogen sodium ortho-arsenate . H 2 NaAs0 4 . Ammonium magnesium ortho-arsenate . (N H 4 )MgAs0 4 . Sodium pyro-arsenate . . • • Na 4 As 2 0 7 . Sodium metarsenate • NaAs0 3 . The salts of pyro- and metarsenic acids, like the acids them- selves, only exist in the solid state ; when dissolved in water they pass into the ortho-compounds. ■ The arsenates are isomorphous with the corresponding phos- phates. COMPOUNDS OF ARSENIC WITH SULPHUR. Three sulphides of arsenic are known, namely— Arsenic disulphide (. found native as Realgar) Arsenic trisulphide {found native as Orpiment) Arsenic pentasulphide . As 2 S 2 . As 2 S 3 . As 2 S 6 . Arsenic Disulphide, As 2 S 2 , is formed when sulphur and arsenic, or arsenic trisulphide and arsenic, are heated together ; or by heating arsenious oxide and sulphur — As 4 0 6 + 7S = 2As 2 S 2 + 3S0 2 . It is prepared on a large scale by distilling a mixture of iron pyrites and arsenical pyrites — FeS 2 ,FeAs 2 + 2FeS 2 = As 2 S 2 + 4FeS. Properties. — Arsenic disulphide is a red, vitreous, brittle solid, having a specific gravity of 3-5- It is readily fusible, and sublimes unchanged. Heated in air or oxygen, it burns with a blue flame, forming arsenious oxide and sulphur dioxide 2As 2 S 2 + 70 2 = 4S0 2 + As 4 0 6 . 2 F 450 Inorganic Chemistry Arsenic disulphide is employed in pyrotechny. So-called Bengal fire consists of a mixture of realgar, sulphur, and nitre. Arsenic Trisulphide, As 2 S 3 , is obtained by heating sulphur and arsenic in the proportion required by the formula, and sublim- ing the compound. It may readily be produced by passing sulphuretted hydrogen tli rough a solution of arsenious oxide in hydrochloric acid— As 4 0 6 + 6H 2 S = 2As 2 S 3 + 6H 2 0. Properties.— The compound, as obtained by precipitation with sulphuretted hydrogen, is a pure canary-yellow solid, which easily melts, and on again cooling forms a brittle crystalline mass. It volatilises and sublimes unchanged, but when heated in air or oxygen it burns with formation of arsenious oxide and sulphur dioxide. Arsenic trisulphide may be regarded as a thio-anhydride, as it gives rise to a series of salts known as thio-arsenites , or sulpharsenites. Thus, when arsenic trisulphide is brought into a solution of a caustic alkali, such as potassium hydroxide, the sulphide readily dissolves with the formation of an arsenite and thioarsenite, thus — As 2 S 3 + 4KHO = HK 2 As0 3 + HK 2 AsS 3 + H 2 0. Upon the addition of an acid, the salts are decomposed and arsenic tri- sulphide reprecipitated — HK 2 As0 3 + HK 2 AsS 3 + 4HC1 = 4KC1 + 3H 2 0 + AsgS* Thio-arsenites.— These salts may be looked upon as being derived from three hypothetical thio-arsenious acids, corresponding to the oxyacids— Ortho-thio-arsenious acid, H 3 AsS 3 . Potassium ortho-thio-arsenite, K 3 AsS 3 . { Ammonium pyro - thio - arsenite, (NH 4 ) 4 As 2 S 5 . Lead pyro-thio-arsenite, Pb 2 As 2 0 5 . Meta-thio-arsenious acid, HAsS 2 . Potassium meta-thio-arsenite, KAsS 2 . Thio-arsenites of the alkali metals, the metals of the alkaline earths, and o( magnesium, are soluble in water, but decompose on boiling. Their solutions are decomposed by acids, with evolution of sulphuretted hydrogen and pre- cipitation of arsenic trisulphide, thus — 2K 3 AsS 3 + 6HC1 = 6KC1 + 3H£ + A s£ 3 . Arsenic Pentasulphide, As 2 S 6 . — This compound is prepared by adding an acid to a solution of a thio-arsenate, thus — *Na 3 AsS 4 + 6HC1 = 6NaCl + 3H 2 S + As 2 S fr A Htimony 45i Arsenic pentasulphide is a yellow, easily fusible solid. It is readily soluble in caustic alkalies, forming an arsenate and a th.o- arsenate — As 2 S 6 + 4KHO = HK 2 As 0 4 + HK 2 AsS 4 + H 2 0. Arsenic pentasulphide, like the trisulphide, gives rise to a series of salts, known as thio-arsenates. These may be regarded as being derived from the three hypothetical thio-arsenic acids— Ortho-thio-arsenic acid,H 3 AsS 4 . Pyro-thio-arsenic acid,H 4 As 2 S7. Meta-thio-arsenic acid, HAsS 3 . Tripotassium ortho-thio-ar senate, K 3 AsS 4 . Hydrogen disodium ortho -thio- arsenate, HNajjAsS* Magnesium pyro-thio-arsenate, Mg 2 As 2 S 7 . Ammonium meta-thio-arsenate, (NH 4 )AsS 3 . ANTIMONY. Symbol, Sb. Atomic weight = 119.6. Occurrence.— Antimony in the uncombined state is found in small quantities in various parts of the world, and notably in Borneo. In combination with oxygen, as Sb 2 0 3 , it constitutes the mineral antimony bloom, or white antimony j and as Sb 2 0 4 it occurs in antimony ochre. In combination with sulphur, as Sb 2 b 3 , it occurs as the mineral stibnite, or grey antimony ore , which is the most important source of the metal; and with both oxygen and sulphur, as Sb 2 0 3 ,2Sb 2 S 3 , it constitutes the mineral antimony blende, or red antimony. . . It also occurs in combination with sulphur and with metals, in the form of thio-antimonites. . , Modes of Formation.— (1.) Antimony is obtained from the native sulphide by one of the two following methods The broken-up ore is heated in plumbago crucibles, along with scrap iron. As the mass melts, the sulphur combines with the iron, forming a slag of iron sulphide, and the liberated antimony settles out beneath — Sb 2 S 3 + 3Fe = 2Sb + 3FeS ( 2 ) The crude sulphide is first liquated, or melted in such a manner as to separate the sulphide from the rocky matter associated with it. The liquated sulphide is then mixed with about halt its weight of charcoal, in order to prevent the mass from caking, an 45 2 Inorganic Chemistry carefully roasted. During this process, the antimony sulphide is partially converted into antimony trioxide (Sb 2 0 3 ) 2> which passes into flues, and is there condensed, leaving a mixture containing antimony tetroxide (Sb 2 0 4 ), and unchanged sulphide. Most of the arsenic present is also oxidised, and volatilises with the anti- mony trioxide, while sulphur dioxide escapes. The residue, con- sisting of the tetroxide and sulphide (known as antimony ash), is mixed with an additional quantity of charcoal and with sodium carbonate, and heated to redness in a crucible, when the changes represented by the following equations take place — (i.) Sb 2 0 4 + 4C = 4CO + 2Sb. By the action of the carbon upon the sodium carbonate, sodium is liberated, which combines with the sulphur of the trisulphide, forming sodium sulphide and metallic antimony — (2.) Na^Og + 2C = 3CO + 2Na. (3.) Sb 2 S 3 + 3Na2 = 3Na 2 S + 2Sb. The sodium sulphide in its turn unites with a further quantity of antimony sulphide, forming a double sulphide of sodium and anti • mony, which, mixed with the sodium carbonate and charcoal, constitutes the slag. The metal obtained by either process is subsequently refined. Properties. — Antimony is a bright, highly crystalline, and very brittle metal, possessing a bluish-white colour, and having a specific gravity of 6.7 to 6.8. It is unacted upon by air or oxygen at the ordinary temperature, but when heated it burns brilliantly, forming antimony trioxide. The metal melts at 450° ; and when allowed to solidify, its crystalline character is seen by the fern-like appearance of its surface. If a quantity of the molten metal be allowed slowly to cool, and when partially solidified the remaining liquid portion be poured off, the interior of the mass is found to be lined with well-formed rhombohedral crystals, isomoiphous with arsenic. In the act of solidification antimony expands, a property which it imparts to its alloys, thus giving to them the valuable quality of taking very fine and sharp castings. The most important of these alloys are type metal (lead 75, antimony 20, tin 5) ; stereotype metal (lead 1 12, antimony 18, tin 3) ; Britannia metal (tin 140, copper 3, antimony 9). Regarded as a metal, antimony is a bad conductor of heat and electricity. Dilute sulphuric and hydrochloric acids are without action upon Antimony Hydride 453 antimony. The concentrated acids convert it into sulphate and chloride respectively — 2 Sb + 6H 2 S0 4 = 3S0 2 + 6H,0 + Sb 2 (S0 4 ) 3 . Antimony is oxidised by nitric acid, dilute acid converting it into antimony trioxide, or a compound of the oxide with nitrogen pent- oxide, Sb 2 0 3 ,3N 2 0 6 , while strong acid oxidises it chiefly into anti- mony tetroxide and pentoxide. Powdered antimony, when thrown into chlorine, takes hre spontaneously, and forms antimony trichloride. Amorphous Antimony.— Antimony is obtained in an amor- phous form, by the electrolysis of a solution of tartar emetic in antimony trichloride. Properties.— Amorphous antimony presents the appearance of a smooth polished rod of graphite, and has a specific gravity ° { c 78 It always contains a certain quantity of antimony trichloride (from 4 to 12 per cent.); but whether this is in chemical union, or merely mechanically retained by the metal, is not known. Amor- phous antimony is very unstable, and readily passes into t e crystalline modification: a slight blow, even a scratch with a needle, causes it instantly to transform itself into the stable form, with explosive violence ; the temperature at the same moment rising to 250 °, and clouds of the vapour of antimony trichloride being evolved. ANTIMONY HYDRIDE ( A ntimon inret ted Hydrogen ). Symbol, SbH 3 . Modes of Formation.— ( 1 .) This compound has never been obtained in the pure state, as usually prepared it is always mixed with hydrogen. It is formed when a solution of antimony tri- chloride in hydrochloric acid is introduced into a mixture generating hydrogen, such as zinc and sulphuric acid. ( 2 .) It is also found by the action of dilute sulphuric acid upon alloys of antimony and zinc— Sb 2 Zn 3 + 3H 2 S0 4 = 3ZnS0 4 + 2 SbH 3 . Properties.— Antimony hydride is a colourless, offensive- smell- ing, and poisonous gas, closely resembling the corresponding 454 / norganic Chemistry arsenic compound in its general behaviour. It bums with a violet* tinted flame, forming water and antimony trioxide— 4SbH 3 + 60 2 = 6H 2 0 + Sb 4 O fl . When the supply of air is limited, water is formed and antimony is deposited ; when, therefore, a cold object is depressed upon the flame, black stains of metallic antimony are obtained. The gas is easily decomposed by heat, and if passed through a glass tube heated at one point, a black deposit of antimony is formed upon the glass. The antimony so deposited is insoluble in a solution of bleaching powder (see Arsenic Hydride). Antimony hydride is decomposed by the halogen elements, with the formation of the halogen hydride, and the halogen compound of antimony SbH 3 + 3C1 2 = 3HC1 +SbCl 3 . Sulphuretted hydrogen, under the influence of sunshine, converts antimony hydride into antimony trisulphide — 2SbH 3 + 3H 2 S = Sb 2 S 3 + 6Hij. When passed into silver nitrate solution, the antimony is preci- pitated in combination with silver, in this way differing from the arsenic analogue — SbH 3 + 3AgN0 3 = 3HNO s + SbAg 3 . COMPOUNDS OF ANTIMONY WITH THE HALOGENS. The compounds represented by the following formulae are known — SbF 3 ; SbCl 3 ; SbBr 3 : Sbl 3 . SbF 6 ; SbCl 6 . Antimony Trifluoride, SbF 3 , is prepared by dissolving the trioxide in aqueous hydrofluoric acid. From the concentrated solution it is deposited in the form of white deliquescent crystals. It dissolves in water, and is gradually converted into an oxyfluoride. Antimony Pentafluoride, SbF 5 , is obtained when hydrated antimony pent- oxide is dissolved in aqueous hydrofluoric acid. When the solution is evapo- rated the compound remains as an amorphous eum-like residue. Antimony Pentachlonde 455 Both of these fluorides exhibit a great tendency to unite with alkaline fluorides, forming double salts, such as SbF 3 ,2KF; SbF 3 2NH 4 F, in the case of the trifluoride ; and SbF 5 ,KF ; SbF 5 ,2KF, with the pentafluonde. Antimony Trichloride, SbCl 3 , is formed when chlorine is passed over metallic antimony, or antimony trisulphide— 2Sb + 3CI2 — 2SbCl 3 . 2Sb 2 S 3 + 9C1 2 = 4SbCl 3 + 3S 2 C1 2 . It may also be obtained by the action of boiling hydrochloric acid, containing a small quantity of nitric acid, upon either metallic antimony, antimony trioxide, or trisulphide— Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. Properties. — Antimony trichloride is a colourless, deliquescent, crystalline substance, melting at 73.2 0 to an oily liquid, which again solidifies to a soft translucent mass. It is soluble m alcohol and in carbon disulphide, and from the latter may be crystallised. It may be dissolved in a small quantity of water unchanged. Thus, if allowed to deliquesce it liquefies in the water it absorbs, forming a colourless solution, which, upon evaporation over sulphuric acid, again deposits crystals of the trichloride. The addition of larger quantities of water results in the formation of oxychlorides — (1.) SbCl 3 + H 2 0 = 2HC1 + SbOCl. (2.) 4SbCl 3 + 5H 2 0 = 10HC1 + S^C^Cl^ Continued boiling with water removes the whole of the chlorine, forming the trioxide— Sb 4 0 5 Cl 2 + H 2 0 = 2HC1 + Sb 4 Oe. Antimony chloride unites with alkaline chlorides, forming double salts (see Antimony fluoride), such as SbCl 3 ,2NH 4 Cl ; SbCl 3 ,3KCL With potassium bromide it forms the compound SbCl 3 ,3KBr, which, strangely enough, appears to be identical with the double compound of antimony tnbromide with potassium chloride, SbBr 3 ,3KCl. Antimony Pentaehloride, SbCl 5 , is obtained by passing excess of dry chlorine over metallic antimony, or antimony trichloride, m * The mixed product obtained by the action of water upon antimony tri- chloride is known as powder of Algaroth. 45 6 I nor gariic Chemistry a retort, when antimony pentachloride distils over in the excess of chlorine— SbCl 3 + Cl 2 = SbCl 6 . Properties.— Antimony pentachloride is a nearly colourless strongly -fuming liquid. It solidifies, when cooled, to a mass of colourless crystals, which remelt at - 6 °. Under the ordinary atmospheric pressure the pentachloride dissociates, when heated into the trichloride and chlorine, but under reduced pressure it may be boiled and distilled. Thus, under a pressure of 22 mm it boils at 79 0 . By the regulated action of ice cold water, oxychlorides are formed — ( 1 .) SbCl 6 + H 2 0 = SbOCl 3 + 2HC1. ( 2 .) SbOCl 3 + H 2 0 = Sb0 2 Cl + 2HC1. Antimony pentachloride, and also the oxychlorides, are con- verted by hot water into pyro-antimonic acid (analogous to pyro- arsenic and pyro-phosphoric acids) — 2SbCl 6 + 7H 2 0 = H 4 Sb 2 0 7 + 10HC1. 2Sb0 2 Cl + 3H 2 0 = H 4 Sb 2 0 7 + 2HC1. Sulphuretted hydrogen (the sulphur analogue of water) acts upon antimony pentachloride, forming antimony sulphotrichloride corresponding to the oxytrichloride — SbCl 6 + H 2 S = SbSCl 3 + 2HC1. Antimony tribromide, SbBr 3 , and antimony tri-iodide, Sbl 3 , are obtained by adding powdered antimony to solutions of the halogens in carbon disulphide from which liquid the compounds are crystallised : the bromide as colourless deliquescent crystals, and the iodide as hexagonal ruby-red crystals. Both of these compounds are similarly acted upon by water, forming the oxy bromide' SbOBr ; Sb 4 0 6 Br 2 , and the oxyiodides SbOI ; Sb 4 0 5 I 2 . OXIDES AND OXY ACIDS OF ANTIMONY. Three oxides of antimony are known — Antimony trioxide (antimonious oxide) . (Sb 2 0 3 ) 2 or Sb 4 O fl . Antimony tetroxide .... Sb 2 0 4 . Antimony pentoxide .... Sb 2 0 6 . Antimony Tetr oxide 457 No acids ace Unown responding to of the hypothetical metantimomous acid, HSbO* has oeen the composition NaSb0 2 , 3H 2 0. Three acids are known derived from antimony pentoxide, which are analagous to the three arsenic and phosphor, c actds- Orthoantimonic acid Pyroantimonic acid . Metantimonic acid . H 3 Sb0 4 . H 4 Sb 2 0 7 . HSb0 3 . Antimony Trlo*H«. St.O. . — - sodium carbonate to remove the free acid 4SbCl 3 + 6H 2 0= Sb 4 0 6 + 12HC1. Properties.— Antimonious oxide is a white powder which, former are P regions. (See Arsenious Oxide , with potassium antimony tartrate, or tartar emetic 4HK(C 4 H 4 0 6 ) + Sb 4 O e = 4(SbO)K(C 4 H 4 O c ) + 2H 2 0. Antimonious oxide burns in the air, forming the tetroxide- (Sb 2 0 3 ) 2 + 0 2 = 2Sb 2 0 4 . oxlde 2Sb 2 O s = 2Sb 2 0 4 + 0 2 45 8 Inorganic Chemistry potassium tartrate, forming tartar emetic and metantimonic acid thus— HK(C 4 H 4 O 0 ) + Sb 2 0 4 = (Sb0)K(C 4 H 4 0 6 ) + HSb0 3 . Antimony Pentoxide, Sb 2 O t , is obtained by oxidising metallic antimony with nitric acid, and heating the antimonic acid so obtained to a temperature not exceeding 275 0 . Properties — Antimony pentoxide is a straw-coloured powder insoluble m water. When heated to 3 oo”, it gives up oxygen and is converted into the tetroxide. Its feeble acidic character is seen by its formation of an alkaline metantimonate, when fused with an alkaline carbonate — Sb A + Na 2 C 0 3 = C 0 2 + 2NaSb0 3 . caf^h, 1110 a C K dS , and ^* ttaonates — None of 20 7 . Hydrogen potassium pyro- antimonate, H 2 K 2 Sb 2 0 7 . Metantimonates. * Potassium metantimonate, KSbO s . Barium metantimonate, Ba(Sb0 3 ) 2 . 0nly yP es of antimonates are known, and as the salts of the type MobOi. are the best known, the name antimonates was formerly applied to Antimony Pentasulphide 459 COMPOUNDS OF ANTIMONY WITH SULPHUR. Two sulphides of antimony are known, namely— Antimony trisulphide • ^ 2 ^ 3 ‘ Antimony pentasulphide • • • bb 2 b &‘ Antimony Trisulphide, Sb 2 S 3 .-This compound occurs native as the mineral stibnite, or grey antimony ore. It is prepared by heating a mixture of powdered antimony and sulphur tion required by the formula) beneath a layer of fused sodt chloride in a crucible. It is also formed when sulphuretted hydro- gen is passed through a solution of anttmony tnchlor.de, or a solution of tartar emetic— 2SbCl 3 + 3H 2 S = Sb 2 S 3 + 6HC1. Properties.-Antimony trisulphide as it occurs native and as obtained by the direct union of antimony and sulphur, is a grey bllck crystalline substance ; as prepared by preap™ with sulphuretted hydrogen, and subsequently drying at 200 it is a brick-red amorphous powder, which when melted and slow* cooled solidifies in the crystalline form. Antimony sulphide sub- “m’ slchanged when heated in an inert gas, but when ^ m air, sulphur dioxide is evolved, and ant.mon.ous ox.de and tetrox de are formed. Heated with hydrochloric acid, it evolves sulphuretted hydrogen, and forms antimony trichloride— Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. Antimony Pentasulphide, Sb 2 S 5 , is obtained when antimony pentachloride is mixed with water, and sulphuretted hydrogen passed through the liquid— 2SbCl 6 + 5H 2 S = Sb 2 S 6 + 10HC1, 2Sb0 2 Cl + 5H 2 S = Sb 2 S 6 + 4H 2 Q + 2HC1. for the a^mony 3 impounds, as that which is in use for the similarly constituted arsenic and phosphorus compounds— Ortho Pyro Meta Phosphates. . M 3 PO 4 , M4P2O7 . mpo 3 Arsenates. M 3 As0 4 M 4 As 2 0 7 MAsOq Antimonates. M 4 Sb 2 0 7 MSbOg 4-6° Inorganic Chemistry Properties.— Antimony pcntasulphide is a dark, orange-red powder, which on being heated, is decomposed into the trisulphide and free sulphur. i u° th i° f th t SC anUm ° ny sul P hides ma Y be regarded as thio-anhydrides for although no thio-acids derived from them are known, salts have been produced which may be viewed as derivatives of hypothetical thio-acids. When the trisulphide is either fused with caustic potash, or boiled in an aqueous solution potassium thio-antimonite is formed— ’ 2Sb 2 S 3 + 4KHO = 3KSbS 2 + KSb0 2 + 2H 2 0. Similarly, when antimony pentasulphide is dissolved in potassium hydroxide a mixture of antimonate and thio-antimonate is obtained 4Sb2S B + 24KHO = 6 K 3 SbS 4 + SKgSbCh + 12H 2 0. The following are illustrations of the thio-salts of antimony— Sulphide. Sb 2 S 3 Sb2S 5 Hypothetical Acids. ( (Ortho) H 3 SbS 3 . < (Meta) HSbS 2 . l(Pyro) H 4 Sb 2 S 5 . . (Ortho) H 3 SbS 4 . Only ortho-thio-antimona/er are known. Salts. Potassium thio-antimonite, KgSbSg. Silver thio-antimonite, AgSbS* Lead thio-antimonite, Pb^h^. /Potassium thio-antimonate, K^bS* J Sodium thio-antimonate ) I (Schlippe’s salt) j Na 3 SbS 4 . ^ Barium thio-antimonate, Ba 3 (SbS 4 ) ? BISMUTH. Symbol, Bi. Atomic weight = 207 . 5 . Occurrence.— Bismuth occurs most commonly in the uncom- bined condition. It is met with in combination with oxygen, as B i 2 0 3 , in bismuth ochre j and in combination with sulphur as Bi 2 S 3 , in bismuth glance. Mode of Formation.— Bismuth is principally obtained from the native metal, and from ores with which metallic bismuth is associated. The broken-up ore is liquated by being heated in inclined iron pipes, when the bismuth readily melts and drains away. Pure bismuth can be prepared from the crude metal thus ob- tained, by first dissolving it in nitric acid, forming bismuth nitrate. Bismuth Trichloride 461 Bi(N0 3 ) 3 , and then precipitating the basic nitrate by the addition ° f Wate, ~Bi(N0 3 ) 3 + 2H 2 0 = (Bi0)N9 3> H 2 0 + 2HNO, of heat, and the oxide is reduced then by the carbon 2(Bi0)N ° 3 S; + ^^ 2BL + 2H2 °' Properties. -Bismuth allowed °S?sr sstf s =wa ^ ^ » “™£ 1,7'S sr “ ;s -« - «•. — • •< < p»- of bismuth, 2 of lead, 1 of tin, and 1 of cadmium. COMPOUNDS OF BISMUTH WITH THE HALOGENS. Compounds represented by the following formulae are known- BiF ■o;n. BiBr, Bil 3 . BiBr 3 (BiBr 2 ) 2 ? BiCl 3 (BiCl 2 ) 2 Bismuth Trichloride, BiCl s , may be prepared, by pa ssing dry chlorine over powdered bismuth gently heated m a retort. ture of powdered bismuth and mercuric chlonde- Bi + 6HgCl 2 = 3Hg 2 Ct + 2BiCl 3 . 462 Inorganic Chemistry roperties.— Bismuth trichloride is a white, extremely deli- quescent- crystalline compound. Heated in an atmosphere of chlorine, it melts to a yellow liquid. It is decomposed by water with the precipitation of bismuth oxychloride BiCl a + HjO = 2HC1 + BiOCl. Bismuth Dichloride, (BiCy* is obtained by the prolonged heating of mercurous chloride and finely powdered bismuth to 2V in a sealed tube. The mixture melts, and mercury collects at the bottom, and on cooling, the dichloride solidifies as a black extremely deliquescent solid upon the surface of the mercury! When heated above 300 ”, the dichloride is resolved into the tri- chloride and metallic bismuth. The molecular weight of the compound is unknown. Bismuth Tribromide, BiBr 3 , is prepared by gradually adding bromine to powdered bismuth, and slightly warming the mixture for some time. The bromide sublimes in the form of golden yellow, deliquescent crystals, which are decomposed by water, forming oxybromide, BiOBr. msmuth Tri-iodide, Bil„ is prepared by subliming a mixture ofiodine and !,“ h ' ^ he , S “ b ™ at f. 1S aft e™ards finely powdered and again sublimed, and the product finally distilled in a stream of carbon dioxide, when it forms S, K W r‘ h a br ‘ ght “ eta ' lic lustre - Boilin e water decomposes the compound, with formation of bismuth oxyiodide, BiOI. COMPOUNDS OF BISMUTH WITH OXYGEN . F our oxides of bismuth are known, namely— Bismuth dioxide ( Hyftobismuthous oxide ) . . BigOg. » trioxide ( Bismuthous oxide ) . . Bi 2 0 3 . „ tetroxide (. Hypobismuthic oxide ) . . Bi 2 0 . „ pentoxide ( Bismuthic oxide ) . . Bi 2 0 6 . None of these compounds is an acid-forming oxide, although with the exception of the first, they all form hydrated oxides! These hydrated oxides have no acidic properties, and no salts have been obtained in which the acidic portion of the molecule consists of bismuth and oxygen. All the four oxides, when acted upon by acids, yield the same series of salts, in which the bismuth fulfils the functions of a tnvalent element, replacing three atoms of Bismuth Trioxide 4^3 hydrogen. In the case of the dioxide, metallic bismuth is de- posited, thus — 3Bi 2 0 2 + 6H 2 S0 4 = 2Bi 2 (S0 4 ) 3 + 2Bi + 6H 2 0. While with the higher oxides oxygen is evolved— Bi 2 0 6 + 6HN0 3 = 2Bi(N0 3 ) 3 + 20 + 3H 2 0. Bismuth trioxide is the most stable and the most important of the oxides; when heated in air, the remaining three compounds are converted into the trioxide ; the dioxide by oxidation, and the tetroxide and pentoxide by loss of oxygen. The trioxide alone is unchanged on being heated in air or oxygen. Bismuth Dioxide, Bi 2 0,— This oxide is prepared by adding a mixed solu- tion of bismuth trichloride and stannous chloride to an excess of a 10 per cent, solution of caustic potash, air being excluded: potassium stannate is formed, and bismuth dioxide is precipitated— SnCl 2 + 2BiCl 3 + 10KHO = Bi 2 0 2 + 8KC1 + K 2 SnO s + 5H 2 0. properties -The precipitated compound, after being washed in dilute caustic potash, and dried in vacuo, is obtained as a black crystalline powder. When heated in air it smoulders, uniting with oxygen to form the trioxide. When moist it oxidises spontaneously BioOo + O = Bi 2 0 3 . Bismuth Trioxide, Bi 2 0 3 , is formed when the metal is burnt in air or oxygen. It may also be obtained by heating the hydrated oxides, the carbonate, or basic nitrate, thus— Bi 2 0 3 ,H 2 0 = Bi 2 0 3 + H 2 0. Bi 2 0 3 ,C0 2 = Bi 2 0 3 + CO, 2(Bi0)N0 3 ,H 2 0 = Bi 2 0 3 + N 2 0 4 + O + 2H 2 0. Properties.— Bismuth trioxide is a cream-coloured powder, insoluble in and unacted upon by water, and is the only oxide of bismuth which is unchanged when heated in the air or in oxygen. It dissolves in acids, forming salts of bismuth— Bi 2 0 3 + 6HN0 3 = 3H 2 0 + 2Bi(N0 3 ), Bi 2 0 3 + 3H 2 S0 4 = 3H 2 0 + Bi 2 (S0 4 ) 3 . With small quantities of hydrochloric acid it first forms bismuth 4^4 / n organic Chemistry oxychloride, BiOCl, which dissolves in additional acid, yielding the trichloride — Bi 2 0 3 + 2HC1 = H 2 0 + 2BiOCl. BiOCl + 2HC1 = H 2 0 + BiCl 3 . None of these compounds is soluble in water without the presence of excess of the acid. Water alone converts them into insoluble basic salts, and free acid, which in the state of extreme dilution is unable to exert any solvent action. Thus, in the case of the nitrate, when water is added, this compound is decomposed into the basic nitrate and free nitric acid— Bi(N0 3 ) 3 + 2H 2 0 = (BiO)N0 3 ,H 2 0 + 2HNO a . Bismuth trioxide forms three hydrates, represented by the formulae — Bi 2 0 3 ,H 2 0 Bi 2 0 3 ,2H 2 0 Bi 2 0 3 ,3H 2 0 These hydrates have no acid properties, and are incapable of combining with bases to form salts, but themselves play the part of a base, uniting with acids to form bismuth salts. The trihydrate is obtained, by pouring an acid solution of bis- muth nitrate into an excess of strong aqueous ammonia — 2Bi(N0 3 ) 3 + 6NH 3 H 2 0 = 6NH 4 (N0 3 ) + Bi 2 0 3 ,3H 2 0. Heated to ioo°, it is converted by loss of water into the mono- hydrate — Bi 2 0 3 ,3H 2 0 = Bi 2 0 3 ,H 2 0 + 2H 2 0. Bismuth Tetroxide, Bi 2 0 4 , is formed by the action of potassium hypochlorite upon the trioxide, the product being dried at i8o°— Bi 2 O a + KCIO = Bi 2 0 4 + KC1. Properties, — Bismuth tetroxide is a brownish -yellow powder, which readily parts with an atom of oxygen, and passes into the trioxide. Bismuth Pentoxide, Bi 2 0 6 , is prepared by passing chlorine into a nearly boiling solution of caustic potash, in which is sus- pended a quantity of bismuth trioxide — Bi 2 O a + 4KHO + 2C1 2 = 4KC1 + H 2 0 + Bi 2 0 5 ,H 2 0. Bismuth Trisulphide 4^5 Properties. — Bismuth pentoxide is a red powder, which is readily deoxidised into the tetroxide and trioxide by heat. It com- bines with water, forming the hydrate Bi 2 0 6 ,H 2 0, but with excess of water it is gradually deoxidised into hydrates of the tetroxide or trioxide. Bismuth pentoxide is reduced, with evolution of oxygen, by both nitric and sulphuric acids — Bi 2 0 6 + 3H 2 S0 4 = Bi 2 (S0 4 ) 3 + 3H 2 0 + 0 2 . With hydrochloric acid it behaves in the usual manner of peroxides, causing the evolution of chlorine Bi 2 0 6 -I- 10HC1 = 2BiCl 3 + 5H 2 0 + Cl 2 . Bismuth Trisulphide, Bi 2 S 3 .— This compound is the only com- pound of bismuth with sulphur that is known with certainty. It occurs native as the mineral bismuth glance. It is precipitated when sulphuretted hydrogen is passed into a solution of a bismuth salt — 2Bi(N0 3 ) 3 + 3H 2 S = Bi 2 S 3 + 6HN0 3 . It is also obtained by heating together the requisite proportions of bismuth and sulphur. . Properties. — As obtained by precipitation, bismuth sulphide is a dark brown, almost black powder ; the native sulphide forms steel-grey lustrous crystals. It is decomposed, when strongly heated, into its constituent elements. Bismuth sulphide differs from the corresponding anti- mony and arsenic compound, in not being dissolved by alkaline hydrates or sulphides. CHAPTER IV THE ELEMENTS OF GROUP I. (. FAMILY A.) This family comprises the following five elements, known as the alkali metals — Atomic Weights. Melting-points. Lithium (Li) . . . 7.01 . . 180 0 Sodium (Na) . . 22.99 • . • 95-6° Potassium (K) . • 3903 • . 62.5° Rubidium (Rb) . . 85.2 . . • • 38.5° Caesium (Cs) . . • 132.7 . 26.5° The most important, and the most abundant of these elements, are potassium and sodium, which also were the first to be dis- covered, having been isolated by Davy in the year 1807. The element lithium, although widely distributed in nature, is for the most part found only in minute quantities : the element was first isolated by Bunsen in the year 1855. The two remaining elements are still rarer substances, usually met with in very minute quantities accompanying sodium and potassium. Both of these elements were discovered by Bunsen by means of the spectroscope — caesium in i860, and rubidium in the following year. All these elements are soft, silvery-white metals, which may be readily cut with a knife, and which rapidly tarnish in the air. They all decompose water at the ordinary temperature. The mejnbers of this family exhibit that gradation in properties which is met with in all similar families. Thus, their melting-points gradually decrease as their atomic weights rise, as will be seen from the figures given above. Their chemical activity also steadily increases as we pass from lithium to caesium. Thus, in the case of their behaviour in contact with water ; potassium, when thrown upon cold water, decomposes that liquid with sufficient energy to cause the ignition of the hydrogen which is evolved : sodium under the same conditions melts and floats about upon the surface, but the action is not sufficiently energetic to effect the inflammation of the gas, unless the water be previously heated ; while with lithium, even with boiling water, the temperature produced by the reaction 466 The Alkali Metals 467 does not rise to the ignition-point of hydrogen. The same is also seen in the spontaneous oxidation of these elements when they are exposed to the air. Thus, lithium when cut with a knife, although it is soon covered with a film of oxide, nevertheless retains its bright metallic surface for some seconds : sodium tarnishes so much more quickly, that the film of oxide appears almost to follow the knife. When potassium is cut, the bright surface can scarcely be seen, so rapid is the oxidation, and if left exposed, a fragment of the metal soon begins to melt by the heat of its own oxidation, and frequently spontaneously ignites. With rubidium and caesium the oxidation is even more rapid, and a fragment of these metals freely exposed to the air very rapidly takes fire spontaneously. The electro-positive character of these elements gradually in- creases from lithium to caesium, which is the most electro-positive of all the known elements. The term alkali , applied to metals of this family, was originally used (before any distinction was made between potash and soda) to denote the salt obtained by treating the ashes of plants with water. Later on, in order to distinguish between this substance and what became known as the volatile alkali (i.e. } ammonium carbonate), it was termed the fixed alkali. The first distinction between potash and soda, was based upon the erroneous belief that the former was entirely of vegetable origin, while the latter was only to be found in the mineral kingdom ; hence the names vegetable alkali and mineral alkali were used to denote these two sub- stances, both of which were regarded as elementary bodies until 1807, when Davy showed that they contained the two metals, potassium and sodium. The resemblance between the different members of this family, and between their compounds, is very close ; so much so, that in the case of sodium, potassium, rubidium, and caesium, there are scarcely any ordinary chemical reactions by which they can be distinguished. They are all readily identified, however, by means of the spectro- scope. When a minute quantity of a lithium salt is introduced upon a loop of platinum wire, into the non-luminous Bunsen flame, the latter is tinged a brilliant crimson red colour : a potassium salt similarly treated, colours the flame a delicate lilac, while a sodium compound gives a brilliant daffodil -yellow colour. The colour imparted to a flame by rubidium and caesium salts, is indistinguish- able by the eye from that given by potassium compounds ; and, moreover, when any of these are mixed with a sodium salt, the 468 Inorganic Chemistry intense yellow emitted by the latter, completely masks the colours given by the others. By means of the spectroscope, not only are the apparently similar colours given by potassium, rubidium, and caesium readily distinguished, but the presence of any or all of them is easily detected, even when admixed with sodium salts. Spectrum analysis is based upon the fact, that light of different colours has different degrees of refrangibility, and therefore when passed through a prism, the different coloured rays are bent out of their straight course at different angles. Ordinary white light Fig. 124. is composed of rays of all degrees of refrangibility, i.e., rays of all colours : hence, when a beam of such light is passed through a prism, the various coloured rays are separated, and become spread out in the order of their refrangibility, from the least refrangible red at the one extreme, to the deep violet at the other. This familiar « rainbow ” coloured band of light is termed the continuous spectrum. A simple form of spectroscope is seen in Fig. 124. The light is caused to pass through a narrow slit at the end of the fixed tele- scope B. If the prism P be removed, and the telescope A be The Alkali Metals 4^9 moved round so as to be in a continuous line with B, a magnified image of the slit is seen by the observer. When the prism is replaced, and A is moved into such a position that the bent rays fall upon its lens, the continuous spectrum is seen, which is an infinite number of strips of light (corresponding to the image of the slit) of all colours, arranged side by side. If the light to be examined, instead of being ordinary white light, were composed of rays all of one degree of infrangibility (/.*., monochromatic light), there would be produced only a single image of the slit, which would fall in that position corresponding to the particular degree of refrangibility of the light. Such a monochromatic light is pro- duced when a sodium salt is heated in a Bunsen flame ; if, there- fore a salt of this metal be introduced upon a loop of platinum wire into the non-luminous flame G, and the light, after passing through the prism, be observed through A, instead of a continuous spectrum, there will be seen a single image of the slit, falling in the brightest yellow part of the spectrum. When the sodium salt is replaced by a lithium salt, it is seen that two images of the slit are obtained, one in the red and the other in the yellow regions of the spectrum. The light emitted from this element consists of rays of two degrees of refrangibility. We say, therefore, that the spectrum of sodium is one yellow line * and that of lithium consists of one red and one yellow line. In order to distinguish the posi- tions of, for example, the yellow lithium line and that given by sodium, an image of a graduated scale, illuminated by the candle flame F, is also thrown into the telescope A. If salts of sodium and lithium mixed together be introduced into the flame G, then three images of the slit are seen, namely, the yellow line given by the sodium, the yellow line of the lithium, situated slightly nearer the red, and the lithium red line. Potassium, like lithium, gives a light of two degrees of refrangi- bility, forming consequently two images of the slit, one in the deep red and the other in the deep violet ; if, therefore, lithium, sodium, and potassium salts are mixed, and examined by the spectroscope, five lines are seen (Fig. 125), namely, two red (one belonging to lithium and one to potassium), two yellow (one belonging to lithium and one to sodium), and the violet line o potassium. * In reality, when examined by a higher dispersive power, the sodium line is seen to be a group of lines. 470 Inorganic Chemistry When analysed in this manner, the lights emitted by rubidium and caesium compounds, are seen to be totally different from each other, and from potassium. The spectrum of rubidium consists of two prominent lines in the violet (nearer the blue region than that belonging to potassium), two brilliant red lines (very near the potassium red line), and a number of less brilliant lines in the orange, yellow, and green. That of caesium consists of two bril- liant blue lines, two bright red lines (near the lithium red line), and a number of less prominent lines in the yellow and green. It will be seen, therefore, that the three elements potassium, rubidium and caesium may be at once sharply distinguished by this optica! method of analysis, although they so closely resemble one another in their chemical behaviour, as to render it highly probable that the separate existence of the two latter would never have been dis- covered by chcmicnl methods alone. Fro. T25. Indeed, before the discovery of caesium by Bunsen, a rare mineral known as Pollux (now known to contain caesium), was mistaken for a potassium mineral.* The element lithium, the member of the family that belongs to the Typical series , exhibits certain characteristic differences from the other members. This is seen particularly in the case of the carbonate and phosphate of this element. Lithium , carbonate is so little soluble in water, that it is precipitated by the addition of carbonate of either sodium or potassium to a solution of a lithium compound. The phosphates of all the other members are readily soluble in water, while lithium phosphate is almost insoluble, and is precipitated from solutions of a lithium salt by the phosphates of either sodium or potassium. In these two compounds, the car- bonate and phosphate, lithium behaves more like one of the metals of the alkaline earths. * The student should consult special works on spectrum analysis. Potassium 47 1 All the metals of this family are monovalent, and replace each other, atom for atom, in chemical compounds. POTASSIUM. Symbol K. Atomic weight = 39 . 03 * Occurrence. — In combination this element is widely distributed in nature. It forms an essential constituent of many of the com- mon silicates, and rocks, which form the earth’s crust. From these rocks, by processes of disintegration, the potassium com- pounds find their way into the soil, from whence they are absorbed by plants, which can only flourish in a soil that contains com- pounds of potassium. Most of the potassium found in plants is present in combination with organic acids. From the vegetable kingdom, potash passes directly into the bodies of animals. The material known as suint, which is the oily perspiration of the sheep, that accumulates in, and is extracted from the wool, consists of the potassium salt of an organic acid (sudoric acid). In the form of chloride and sulphate, potassium is present in sea water and many mineral springs. As nitrate it is found as a crystallised efflorescence upon the soil, notably in Feru and Chili, where it is associated with sodium nitrate. The largest supplies of potassium compounds are met with in the great same deposits of Stassfurt, where the element is found as chloride (KC ) in sylvine , as a double chloride of potassium and magnesium (KCl,MgCl 2 ,6H 2 0) in carnallite , and as a mixed sulphate in kainite (K 2 S0 4 ,MgS0 4 ,MgCl 2 ,6H 2 0). Modes of Formation.— ( 1 .) The method by which Davy first effected the isolation of potassium was by the electrolysis of potassium hydroxide: the method maybe illustrated by the ex- periment represented in Fig. 126 . A small quantity of potassium hydroxide is gently heated in a platinum capsule, which is con- nected to the positive terminal of a powerful battery. A stout platinum wire, flattened out at one end, is made the negative pole. When this is introduced into the fused potash, a brisk evolution of gas takes place, and minute beads of metallic potassium make their appearance in the liquid, and upon the negative electrode, some of which ignite upon the surface. The decomposition takes place according to the equation — 2KHO - H 2 4- 0 2 + Kg. 47 2 Inorganic Chemistry (2.) Potassium may also be obtained by allowing melted potassium hydroxide to pass over iron turnings heated to whiteness, when the magnetic oxide of iron is formed— 4KHO + 3Fe = Fe 3 0 4 + 2H 2 + 2 ^ This is known as Gay-Lussac and Th^nard’s method. (3.) The method devised by Brunner, and modified by Wohler, Deville, and others, consisted in heating to whiteness an intimate mixture of potassium carbonate and carbon. This mixture was obtained by first igniting in a covered iron pot, crude tartar (hydro- gen potassium tartrate, or cream of tartar), which was thereby decomposed as indicated by the equation— 2HKC 4 H 4 0 6 = K 2 C0 3 + 3C + 5H 2 0 + 4CO. Potassium 473 The charred mass was then introduced into an iron retort ( P , Fig. 127), and strongly heated in the furnace, when the potassium carbonate was reduced by the carbon as follows— K 2 C 0 3 + 2 C = 3 CO + K 2 . As soon as the potassium began to be formed, which was readily Fig. 127. ascertained by the green appearance of the vapour, seen on looking in at the open mouth of the retort, the condenser was attached. This consisted of a flat, shallow iron tray, b (Fig. 128), upon which was fitted the cover a, the two portions being clamped together. The object of this special form of condenser is to cool the potassium as rapidly as possible, for it is found that unless the metal is quickly cooled, it combines with the carbon monoxide, forming a highly explosive compound (be- lieved to have the composition K 6 (CO) 6 ). By the use of this form of condensing apparatus the formation of this compound is reduced to a minimum, but not entirely prevented. (4.) A more recent method by which potassium (and sodium) . is prepared on a manufacturing scale, was devised by Castnei Fig. 128. 474 Inorganic Chemistry (1886). In this process potassium hydroxide is strongly heated with a carbide ot iron, having approximately the composition CFe^. (I his compound is obtained by heating a mixture of pitch and finely divided iron.) The potassium hydroxide, with the powdered carbide of iron, is introduced into large egg-shaped retorts, one of which is repre- sented in Fig. 129. These retorts are placed upon hydraulic lifts, so that they can be lowered away from their covers, to the ground- Fig. 129. level, in order to be discharged at the end of the distillation. The retorts are heated by gaseous fuel, and the metal, as it distils, is passed into long narrow cast-iron condensers, from which it drops into iron pots, and is protected from oxidation by mineral oil. The reaction which takes place may be represented by the equation — 6KHO + 2C = 2K 2 C0 3 + 3H 2 + K 2 . By avoiding any excess of carbon, no carbon monoxide is pro- duced, and hence there is no formation of the explosive compound of this gas with potassium. Potassium Peroxide 47 5 Properties.— Potassium is a lustrous, white metal, which at ordinary temperatures is sufficiently soft to be moulded between the fingers ; at o° it is brittle, and shows a crystalline fracture. The metal is readily crystallised by melting a quantity of it in a vacuous tube, and when it has partially solidified, pouring the still liquid portion to the other end of the tube. Potassium melts at 62.5°, and when boiled gives an emerald-green vapour. The metal is rapidly acted on by ordinary air, its freshly cut surface becoming instantly covered with a film of oxide, which, by absorp- tion of atmospheric moisture and carbon dioxide, passes first into the hydroxide and finally into the carbonate. Potassium is there- fore usually preserved beneath naphtha, or some other liquid devoid of oxygen. , When potassium is volatilised in a vacuous tube, the thin him ol metal which condenses upon the cool portion of the tube, is seen to possess a rich violet-blue colour, when viewed by transmitted light. The density of potassium vapour is about 20 (Dewar and Scott), showing that in the vaporous condition the molecules are mon- atomic. _ . . Potassium dissolves in liquefied ammonia, forming a deep indigo solution (p. 243). When potassium is thrown upon water, that liquid is decomposed with sufficient energy to cause the ignition of the liberated hydrogen (p. 151). When heated in carbon dioxide, potassium takes fire, forming potassium carbonate, and carbon (p. 267). Heated in carbon monoxide, it forms the explosive compound already mentioned. Potassium takes fire spontaneously in contact with the halogens, forming the haloid compounds of the metal. When heated in hydrogen, it absorbs the gas, forming a brittle lustrous substance, which inflames spontaneously in the air. This compound has the composition K 4 H 2 . When heated in nitric oxide, potassium burns, forming a mixture of potassium nitrate and nitrite (Holt and Sims). Oxides of Potassium.— When potassium is heated in ordinary air, it takes fire and burns, giving rise to a mixture of the oxides of the metal. Perfectly dry air or oxygen is without action upon potassium. Potassium Peroxide, K 2 0 4 , may be obtained by melting potas- sium in an atmosphere of nitrogen, and gradually displacing the nitrogen by moderately dry oxygen. It is also produced by heating potassium in nitrous oxide. Potassium peroxide is a yellow powder, which, when strongly 476 Inorganic Chemistry heated, is converted into the dioxide K 2 0 2 and oxygen. When thrown into water, oxygen is evolved, potassium hydroxide and hydrogen peroxide being formed — K 2 0 4 + 2H 2 0 = 2KHO + H 2 0 2 + By the regulated combustion of potassium in nitrous oxide, Holt and Sims have obtained compounds having the composition K 2 0 2 and K 2 0 3 . The latter, in contact with the air, undergoes active oxidation into K 2 0 4 . Potassium Hydroxide ( caustic potash ), KHO, is prepared by adding lin.e to a dilute boiling solution of potassium carbonate, in iron vessels, when calcium carbonate is precipitated and potassium hydroxide remains in solution — K 2 C0 3 4- Ca(HO) 2 = CaC0 3 + 2KHO, the reaction being- complete, when the addition of an acid to a small test sample of the clear liquor, produces no effervescence. This reaction is a reversible one, and if the concentration is beyond a certain limit, the potassium hydroxide reacts upon the calcium carbonate, reforming potassium carbonate. The liquid is therefore constantly maintained at a certain state of dilution during the reaction, at the completion of which the mixture is allowed to settle, and the clear solution is then partially concentrated in iron vessels, and finally in silver, until on cooling, the substance solidifies. It is then usually cast into sticks. Caustic potash is a white brittle solid ; it is extremely deliques- cent, and dissolves in water with evolution of heat, forming a highly caustic liquid. The solid, as well as the solution, readily absorbs carbon dioxide, and is employed in the laboratory for this purpose when it is desired to deprive a gas of the last traces of any admixed carbon dioxide. A hot saturated solution of potassium hydroxide, when cooled, deposits crystals of a hydrate having the composition KH0,2H 2 0. Potassium Fluoride, KF. — This salt is prepared by neutralising aqueous hydrofluoric acid with potassium carbonate, and evaporat- ing the solution in a platinum vessel, when the salt is obtained in the form of deliquescent cubical crystals. Potassium fluoride dis- solves in aqueous hydrofluoric acid with evolution of heat, forming the acid fluoride of potassium, HF,KF, which is obtained as an anhydrous salt when the solution is evaporated to dryness and heated to no°. This salt is not deliquescent. When heated to a Potassium Chlorate 477 dull red heat it decomposes into the normal sal, and hydrofluoric r P rA" Stassfu t. Kri MeCl«6H,0, is based upon the fact, from this double salt, KCl,MgU^on 2 y, . . steam can be drive . . -j and it is found that while strong so ^on of magnesium chloride^ ^ ^ sodium potassium chloride is reaany preS ent in the crude * «, „ p a » — « is drawn off into large iro y potassium and n^gnesium CU The mother liquor from these crystallising tanks is either used mmsm §mmmm Potassium Chlorate, KC10 3 .-When chlorine is passed nto a sofution of potassium hydroxide, a mixture of potassium chlora and chloride is obtained, thus— 6KHO + 3C1, = KClOj + 5KC1 + 3H,0. 478 Inorganic Chemistry Th e two salts in solution may be separated by crystallisation, the chlorate being much less soluble in cold water than the chloride. On the manufacturing scale, potassium chlorate is obtained by passing chlorine into milk of lime, when a mixture of calcium chlorate and chloride is formed 6Ca(HO) 2 + 6C1 2 = Ca(C10 3 ) 2 + 5CaCl 2 + 6H s O. The operation is conducted in cast-iron cylinders connected in series, one of which is shown in section in Fig. 130, furnished with mechanical stirring gear, a, b , b. The shaft, and the pipes conveying the chlorine into, and from the vessel, are connected to it by means of the water-sealed joints c, e, e. The manhole /is a short wide leaden pipe, dipping a few inches into the liquid, which allows of the periodic withdrawal of samples for examination. Several reactions are involved in the final formation of the calcium chlorate ; in the first case calcium hypochlorite is produced, thus — 2Ca(HO) 2 + 2C1 2 = Ca(OCl) 2 + CaCl 2 + 2H 2 0. The calcium hypochlorite then passes into a mixture of chlorate and chloride in accordance with the equation — 3Ca(OCl) 2 = Ca(C10 3 ) 2 + 2CaCl 2 . The second change is brought about by the operation of two Potassium Chlorate 479 causes, namely, rise of temperature, and the presence of excess o chlorine. Heat alone is incapable of converting more than a small proportion of the hypochlorite into chlorate, for the former compound is at the same time decomposed into calcium chloride and free oxygen. The excess of chlorine is believed to act, through the intervention of hypochlorous acid, HOC1, merely as a carrier of oxygen, reducing two molecules of calcium hypochlorite to chloride, and oxidising the third to chlorate, thus— 3Ca(OCl)« + 2C1 2 + 2H 2 0 = CaCl 2 + 4HOC1 + 2Ca(OCl) 2 = 2CaCl 2 + Ca(C10 3 ) 2 + 2C1 2 + 2H s O. The absorption of chlorine by the milk of lime is attended with evolution of heat ; care is taken to prevent the temperature from rising above about 7o°, otherwise loss results by the decomposition of hypochlorite with evolution of oxygen, thus— Ca(OCl) 2 = CaCl 2 + 0 2 . When the formation of calcium chlorate is complete, the liquid is allowed to settle, and is then run into concentrating pans, where the requisite amount of potassium chloride in solution demanded by the following equation, is added Ca(C10 3 ) 2 + 2KC1 = CaCl 2 + 2KC10 3 . The liquid is then concentrated in iron pans and allowed to crystallise, when the moderately soluble potassium chlorate sepa- rates out, leaving the very soluble calcium chloride in solution. The chlorate is afterwards purified by recrystalhsation. Potassium chlorate, although only moderately soluble in water, is much more soluble in a strong solution of calcium chloride, hence there loss (usually about xo per cent.) of chlorate in this process. Pechmey s pro- cess for obviating this, consists in concentrating the liquid, obtained by the chlorination of the lime, to a definite specific gravity, and then cooling u to ra” when about 78 per cent, of the calcium chlonde crystallises out. The mother liquor, containing all the calcium chlorate, and only the comparatively small proportion of calcium chloride, is then treated with potassium chloride as usual. Potassium chlorate crystallises in white tables, belonging to the monosymmetric system, which when of large sire often exhibit fine 4^0 Inorganic Chemistry iridescent colours, ioo parts of water at o° dissolve 3.3 parts of the salt ; while at ioo°, 59 parts are dissolved. Potassium chlorate is used largely in the manufacture of matches, on account of the ease with which it gives up its oxygen : thus, if a small quantity of the finely powdered salt be carefully mixed with an equally small amount of red phosphorus, the friction caused by lightly rubbing it with a spatula, is sufficient to cause the mixture to detonate violently. Similarly, when powdered potassium chlo- rate and sulphur are rubbed together in a mortar, the mixture explodes with violence. Potassium chlorate is also largely em- ployed in pyrotechny, especially in the production of coloured effects, where a fiercely burning mixture is required. Potassium chlorate melts between 360° and 370°, and at a tem- perature about 380° begins to evolve oxygen. Potassium Perchlorate, KC 10 4 .— When the chlorate is heated, it first melts arid begins to give off oxygen ; but it soon begins to partially solidify, owing to the formation of potassium perchlorate, and the evolution of oxygen stops unless a stronger heat be applied. The reaction at this stage is expressed by the equa- tion — 2KCIO3 = KC 10 4 + KC 1 + 0 2 . The perchlorate is separated, by first treating the residue with cold water, which dissolves the greater part of the chloride, and afterwards with warm hydrochloric acid, which decomposes any remaining chlorate. The salt is then purified by crystallisation. Potassium perchlorate is very slightly soluble in cold water, 100 parts of water at o° dissolving only 0.7 parts of the salt ; while at ioo°, 20 parts are dissolved. Potassium Bromide, KBr, and Iodide, KI.— These two salts are obtained by similar methods. When bromine or iodine is added to a solution of potassium hydroxide, the reaction which takes place is exactly analogous to that in the case of chlorine (see Potassium Chlorate, above) — 6 KH 0 + 3 Br 2 = KBrO s + 5 KBr + 3 H 2 0 . If the solution so obtained be evaporated by dryness, and the dry residue ignited, the bromate (or iodate) is decomposed, just as potassium chlorate is decomposed by heat, giving off its oxygen, and being converted into bromide (or iodide)— KBr 0 3 = KBr + 30 . Potassium Carbonate 481 The residue, on being dissolved in water and recrystallised, yields pure potassium bromide (or iodide). These salts are manufactured by decomposing ferrous bromide, Fe 3 Br 8 (or iodide, Fe 3 I 8 ), with potassium carbonate, thus — Fe 3 Br 8 + 4K 2 CO s = Fe 3 0 4 + 8KBr + 4C0 2 . The ferrous bromide is obtained by adding bromine to moistened iron borings (see Manufacture of Bromine). Potassium iodide and bromide both crystallise in cubes, and are both readily soluble in water. These salts are chiefly used for medicinal and photographic purposes. Potassium Sulphate, K 2 S0 4 .— This salt is present in the Stass- furt deposits principally as kainite, K 2 S0 4 ,MgS0 4 ,MgCl 2 ,6H 2 0, and as polyhalite, K 2 S0 4 ,MgS0 4 ,2CaS0 4 ,2H 2 0. When kainite is treated with, small quantities of water, or mother liquors from other processes, the extremely soluble magnesium chloride is removed, leaving the potassium magnesium sulphate ; and on adding to this the requisite amount of potassium chloride, the following change takes place — K 2 S0 4 ,MgS0 4 + 3KC1 = 2K 2 S0 4 + KCl,MgCl 2 . From this solution the potassium sulphate crystallises out. Potassium sulphate is also obtained by the action of sulphuric acid upon the chloride, by a process corresponding exactly to the first stage in the Leblanc soda process (q.v .) — 2KC1 + H 2 S0 4 = K 2 S0 4 + 2HC1. Potassium sulphate forms colourless rhombic crystals, contain- ing no water of crystallisation, therein differing from sodium sulphate, which crystallises with ten molecules of water. Potassium sulphate is largely used for agricultural purposes. Potassium Carbonate, K 2 C0 3 . — This salt was formerly obtained exclusively from the ashes of wood and other land plants, and was known under the name of pot-ashes. The process is still carried on in parts of Canada and the United States. The wood is burned in pits, and the ashes are collected and lixiviated with water (with the addition of a small quantity of lime) in wooden tubs with perforated false bottoms. The liquid which is drawn off is evaporated to dryness, and usually calcined to burn away organic matter. This material, known as American pot-^^s, 2 H 482 Inorganic Chemistry contains varying quantities of caustic potash, on account of the previously added lime. The so-called American pearl-ash is a purer product, obtained by concentrating the liquor from the lixiviating tubs until the less soluble impurities crystallise out, and finally evaporating the mother liquor, containing the potassium carbonate, to dryness, and calcining the residue. Potassium carbonate is also obtained from beet-root molasses, an uncrystallisable residue obtained in the manufacture of beet sugar, carried on chiefly in F ranee. The syrup is fermented with yeast, whereby the sugar it contains is converted into alcohol, and then distilled. The residual liquid, known as vinasse , is evaporated to dryness ; and from the black residue, termed “ vinasse cinder,” the potassium carbonate is extracted. Potassium carbonate is obtained also from suint, which, as already stated, contains considerable quantities of potassium in the form of potassium sudorate. The sheep’s wool is lixiviated with water, and the solution evaporated to dryness. The residue is heated in iron retorts, whereby the organic potassium salts are converted into carbonate, while at the same time ammonia, and an illuminating gas, are evolved. The carbonaceous residue is extracted with water, and the potassium carbonate separated by crystallisation. Since the development of the Stassfurt potash supplies, these sources of potassium carbonate are rapidly sinking into the back- ground, and the bulk of this compound is now manufactured from potassium sulphate by a process exactly similar to the Leblanc soda process (q.v.). Potassium carbonate is not manufactured by a method analogous to the ammonia-soda process (Solvay), on account of the too great solubility of potassium bicarbonate (hydrogen potassium carbonate). Pure potassium carbonate may be obtained by igniting cream of tartar (see page 472), and extracting with water ; or by heating hydrogen potassium carbonate, which gives up water and carbon dioxide, thus — 2HKCO3 = k 2 co 3 + h 2 o + C 0 2 . Potassium carbonate forms long prismatic crystals belonging to the monosymmetric system, and containing three molecules of water, K 2 C 0 3 , 3 H 2 0 . The anhydrous salt is highly deliquescent, and very soluble in water. Hydrogen Potassium Carbonate (< bicarbonate of potash ), Potassium Nitrate 483 HKCOg, is produced by passing carbon dioxide into an aquejus solution of the normal carbonate, thus — K a C0 3 + CO a + H 2 0 = 2HKC0 3 . This salt is much less soluble in water than the normal salt, and is readily purified by crystallisation. Potassium Nitrate {nitre, saltpetre), KN0 3 .— This salt has been known since very early times. It occurs as an efflorescence upon the earth, as a result of the oxidation of organic nitrogenous matter in the presence of the potash in the soil, and is found in the neighbourhood of villages, more especially in hot climates, where urine and other readily decomposable organic matters rich in nitrogen find their way into the surface soil. It has been shown that the process of nitrification which results in the formation of nitre under these circumstances, is due to the action of specific organisms, or microbes, and never takes place in their absence. At various times this natural process has been artificially carried on, by mixing manure and other decomposing refuse, with porous soil, lime, and wood ashes, and exposing the mixture in heaps which were moistened from time to time with drainage from manure. The saltpetre earth , collected from the natural sources, or from the artificial nitre plantations, on lixiviation with water, and subsequent evaporation, yielded crystals of potassium nitrate. At the present time, potassium nitrate is almost exclusively ob- tained from sodium nitrate ( Chili saltpetre ), by treatment with potassium chloride derived from the Stassfurt supplies. The requi- site quantities of the two solutions are run into a tank, and heated by means of steam, when the following double decomposition takes place — NaNO s + KC1 = NaCl + KN0 3 . The greater part of the sodium chloride is at once precipitated, and is removed by canvas filters. The clear liquid is then allowed to crystallise in tanks furnished with stirring gear, in order to cause the formation of small crystals, and the nitre-meal so ob- tained is purified by recrystallisation. Potassium nitrate crystallises usually in rhombic prisms, but it can also be obtained in the form of small rhombohedral crystals, isomorphous with sodium nitrate. The solubility of potassium nitrate rapidly increases with rise of 484 Inorganic Chemistry temperature (see Solubility Curve, p. 131). 100 parts of water at o° dissolve 13.3 parts ; at 50°, 86 parts ; and at ioo°, 247 parts. Nitre melts at 339 0 , and at a higher temperature loses oxygen and is converted into potassium nitrite : on this account it readily oxidises many of the elements when heated in contact with them. Thus, a fragment of charcoal or sulphur thrown upon melted nitre, takes’ fire and burns with great energy ; in the one case with forma- tion of potassium carbonate and carbon dioxide, and in the other of potassium sulphate and sulphur dioxide— 4KNO s + 5 C = 2 K 2 C 0 3 + 3 C 0 2 + 2 N 2 . 2KNO s + 2 S = K 2 S 0 4 + S 0 2 + Njj. Nitre is chiefly used in the manufacture of gunpowder and in pyrotechny. Gunpowder is a mixture of nitre, charcoal, and sulphur. The proportions in which these ingredients are present varies, within small limits, according to the special kind of powder, as will be seen from the following table (Abel and Nobel), giving analyses of various powders manufactured at Waltham Abbey. Fine-grain. Rifle Fine-grain. Rifle Large-grain. Pebble Powder. Potassium nitrate , , sulphate Sulphur . . . • Charcoal . Water 73-55 0.36 10.02 14-59 1.48 75-04 0.14 9-93 14.09 0.80 74-95 0.15 10.27 13-52 1. 11 74.67 0.09 IO.07 14.22 O.95 These proportions are very close to those which would be demanded by the juation 2 KN o 3 + S + 3C = KsS + 3CQ 2 + N a , which was at one time supposed to represent the change which takes place when gunpowder is exploded. In reality the decomposition is much more complex and it has been shown that the solid products consist of mixtures of the following substances in varying proportions, depending upon the particular powder, and the conditions of firing— Potassium nitrate. ,, oxide. Ammonium sesquicarbonate. Carbon. Sulphur. Potassium carbonate. , , sulphate. , , sulphide. „ thiosulphate, thioevanate. Compounds of Potassium with Sulphur 485 While the gases that are evolved consist of — Carbon dioxide. Nitrogen. Carbon monoxide. Sulphuretted hydrogen. Marsh gas. Oxygen. Hydrogen. From the combustion of one gramme of powder, the total weight of solids ranges from 0.55 to 0.58 gramme, and the total weight of the gaseous products from 0.45 to 0.42 gramme. COMPOUNDS OF POTASSIUM WITH SULPHUR. Four sulphides of potassium have been obtained, namely — Potassium monosulphide Potassium trisulphide . Potassium tetrasulphide Potassium pentasulphide . k 2 s . K 2 S 3 . K 2 S 4 • k 2 s 5 Just as potassium decomposes water with evolution of hydrogen and formation of potassium hydroxide, so also, when heated in sulphuretted hydrogen (the sulphur analogue of water) it forms potassium hydrosulphide (the analogue of potassium hydroxide) and liberates hydrogen, thus— H 2 S + K = KHS + H, when potassium hydroxide and hydrosulphide are mixed in equi- molecular proportions, potassium monosulphide and water are formed — KHO + KHS = K 2 S + H 2 0. The liquid, on evaporation in vacuo, deposits reddish prismatic deliquescent crystals having the composition K 2 S,5H 2 0. When potassium carbonate and sulphur are heated together, a mixture of the higher sulphides of potassium with potassium thio- sulphate is obtained, thus — 3K 2 C0 3 + 8S = 2K 2 S 3 + K 2 S 2 0 3 + 3C0 2 . 3K 2 C0 3 + 12S = 2K 2 S 6 + K 2 S 2 0 3 + 3C0 2 . The reddish- Drown solid product was named by the early chemists hefiar suiphuris , or “ liver of sulphur.” 486 Inorganic Chemistry SODIUM. Formula, Na = 22.99. Occurrence. — The most abundant natural compound of sodium is the chloride, which is present in sea-water, and in many salt lakes and springs. Enormous deposits of sodium chloride, or rock-salt , are found in Cheshire, Lancashire, and other parts of the world. As nitrate, this element occurs in large quantities in Chili and Peru ; and in combination with silicic acid it is a con- stituent of many rocks. Modes of Formation.— Sodium was first isolated by Davy, by the electrolysis of sodium hydroxide. On a manufacturing scale it is obtained either by the “Castner” process, as described for the preparation of potassium, or by the still more recent method (also due to Castner) of decomposing fused sodium hydroxide by means of a powerful electric current. The caustic soda is melted in a large iron vessel, through the bottom of which passes vertically the negative electrode. Above this electrode an inverted iron pot is suspended, its mouth dipping into the liquid. The sodium and hydrogen, liberated at the cathode, together rise to the surface under the inverted iron pot. The gas escapes by bubbling beneath the edges of the vessel, while the metal remains floating upon the fused caustic. Properties.— Sodium closely resembles potassium in its general properties. It is a soft, white metal which can be readily moulded by the fingers, and is easily pressed into wire. At - 20° it is hard. The colour of sodium vapour is violet, while the colour exhibited by a thin film of the metal, obtained by sublimation in vacuo, is greenish-blue. The vapour density of sodium is about 1 2 (Dewar and Scott), showing that this metal in the vaporous state is monatomic. Like potassium, sodium dissolves in liquid ammonia, yielding a blue solution. When heated in the air, sodium burns, forming the peroxide, Na 2 0 2 . Perfectly dry air or oxygen is without action upon the metal. When heated in hydrogen, sodium forms the hydride, Na 4 H 2 , analogous to the potassium compound, but not spontaneously in- flammable in air. When this is heated to about 300° in vacuo, the whole of the hydrogen is evolved. Alloy of Sodium and Potassium.— When these two metals are melted together beneath petroleum, an alloy is obtained which is Sodium Hydroxide 4 %7 liquid at ordinary temperatures. When prepared and preserved out of contact with air, the alloy resembles mercury in appearance. This alloy is employed in the construction of thermometers for regis- tering high temperatures, where mercury would be inadmissible. Oxides of Sodium.— Two oxides are said to exist, viz., sodium monoxide, Na 2 0, and sodium dioxide, or peroxide, Na 2 0 2 . Sodium Monoxide, Na 2 0,* is said to be obtained by burning sodium in nitrous oxide, at a temperature not higher than i8o°. Sodium Peroxide, Na 2 0 2 , is obtained by allowing sodium to burn briskly in oxygen. It is a yellowish- white solid, which de- composes in contact with water, with considerable rise of tem- perature and evolution of oxygen — Na 2 O a + H 2 0 = 2NaHO + O. The oxygen which is evolved contains appreciable quantities of ozone. When sodium peroxide is slowly added to water, or to dilute hydrochloric acid in the cold, hydrogen peroxide is formed— Na 2 0 2 + 2H 2 0 = 2NaHO + H 2 0 2 . Sodium peroxide forms a crystalline hydrate of the composition, Na 2 0 2 ,8H 2 0 (p. 204). When heated in either nitrous or nitric oxides, it yields sodium nitrite ; in the former case, with the elimina- tion of nitrogen — Na 2 0 2 + 2N 2 0 = 2NaN0 2 + N* Na 2 0 2 + 2NO = 2NaN0 2 . Sodium Hydroxide {caustic soda), NaHO.— This compound is produced when sodium is brought into contact with water, and also when either sodium monoxide, or peroxide, is dissolved in water. On the large scale, caustic soda is prepared by the action of lime upon a boiling solution of sodium carbonate (see Caustic Potash). The so-called tank liquors (obtained in the manufacture of sodium carbonate by the Leblanc process, q: v.) are heated to the boiling- point, and an excess of lime is stirred into the mixture. The sodium sulphide present in the tank liquor, is oxidised into sul- phate by the combined action of air injected into the mixture, * Doubt has recently been thrown upon the existence of this oxide. Erdmann and Kothner ( Annalen der Chemie, Nov. 1896), have shown that rubidium and potassium do not form oxides of the type R 2 0 . And although lithium, the first member of the series, undoubtedly yields the oxide Li a O, it appears doubtful whether sodium is capable of forming a similar compound. 488 Inorganic Chemistry and of sodium nitrate, which is added for this purpose. The liquor, after being causticised, is decanted, or filtered, from the precipitated calcium carbonate, and is concentrated in large cast- iron hemispherical pans. The decomposition suffered by the sodium nitrate depends upon the temperature and concentration of the liquid: at 300° to 360° the change may be expressed by the equation — NaN0 3 + 2H 2 0 = NaHO + NH 3 + 40. The liberated oxygen oxidises the sulphides to sulphates. Caustic soda is now being manufactured by the electrolysis of brine. Mercury is employed as the cathode, with which the libe- rated sodium forms an amalgam ; and on treating this with water, hydrogen is liberated and a solution of caustic soda is obtained. Sodium hydroxide is a white, strongly caustic, and highly de- liquescent solid. It is soluble in water, with considerable rise of temperature, and a concentrated aqueous solution when cooled to -8°, deposits a crystalline hydrate, having the composition 2NaH0,7H 2 0. Sodium Chloride, NaCl. — Of the compounds of sodium with the halogens, the chloride is the most important. In warm climates, as upon the shores of the Mediterranean, sodium chloride is obtained by the evaporation of sea water in large shallow basins, or pools, constructed upon the sea-shore, and exposed to the sun’s heat. As the brine concentrates in these salterns , the crystals of salt are raked off the liquid, and allowed to drain in heaps at the side of the pools. The mother liquors, known as bittern , were formerly utilised for the extraction of the bromine which they contain. Salt is obtained from salt-beds, where it is found in enormous deposits, either by direct mining operations, when the salt is sufficiently pure, or by first dissolving the material in water, whereby insoluble admixed impurities are removed, and afterwards evaporating the brine so obtained. The latter method is carried out by sinking borings through the upper strata of rock, and sending water down to the salt-beds beneath. The brine is then pumped up, and the salt obtained by evaporation. The first stage of the concentrating process, especially where the brine is not very strong, is in some parts carried on by exposing the liquid to the wind. This is effected by causing the solution to trickle over erections of brushwood known as graduators (Fig. 131)1 which are Sodium Chloride 489 built so that the prevailing winds blow across them. The brine is pumped up into the wooden troughs running along the top, from which it escapes by a number of openings, a , a , a , and flows over the pile of brushwood down into the reservoir upon which the erection is constructed. In this way the solution is made to ex- Fig. 131. pose a large surface to the air, and it quickly reaches a concentra- tion when it contains 20 per cent, of salt in the solution. The liquor is then evaporated in shallow iron pans by means of arti- ficial heat, and as the salt crystallises it is lifted out by means of perforated iron skimmers. Salt obtained in this manner, always contains small quantities of other salts, such as sodium sulphate. 490 Inorganic Chemistry calcium sulphate, calcium and magnesium chlorides. The presence of chlorides of magnesium or calcium, causes the salt to become moist, especially in damp weather. Pure sodium chloride may be prepared by add- ing hydrochloric acid to a strong aqueous solution of salt ; the sodium chloride is thereby precipitated, while the other salts remain in solution. Sodium chloride forms colourless, cubical crystals, which are anhydrous. If deposited at — io° it crys- tallises in monosymmetric prisms, with two molecules of water of crystallisation, o which at the ordinary tem- £ perature lose their water, and break up into minute cubes. Sodium chloride is a necessary article of food, for man and other animals ; it is estimated that about 20 lbs. of salt per head of population is annually used, directly or indirectly, for this purpose. The hydro- chloric acid present in the gastric and other acid fluids of the stomach, is derived from the decomposition of sodium chloride which is taken into the organism. Enormous quantities of sodium chloride are employed in the alkali industry, and all the chlorine that is manufactured is derived primarily from this compound. Sodium Bromide, NaBr, and Sodium Iodide, Nal, are pre- pared by methods similar to those for obtaining the potassium Sodium Carbonate 49 1 compounds. They are both isomorphous with sodium chloride, and when deposited at low temperatures they form monosymmetnc crystals containing two molecules of water. Sodium Carbonate, Na 2 C0 3 .— The preparation of this com- pound is carried on by two methods, and constitutes that important industry, the alkali manufacture. The two processes are known by the names of their respective discoverers, namely, the Leblanc process, and the Solvay process, the latter being also known as the ammonia-soda process. , I. The Leblanc method of manufacture consists essentially ot three processes, namely — (i.) The conversion of sodium chloride into sodium sulphate by the action of sulphuric acid, known as the salt-cake process. Two chemical reactions are involved in the process — NaCl + H 2 S0 4 = NaHS0 4 + HC1. NaCl + NaHS0 4 = Na^C^ + HC1. (2.) The decomposition of sodium sulphate, salt-cake , by means of calcium carbonate (limestone) and coal, at a high temperature, whereby a crude mixture of sodium car- bonate and calcium sulphide is obtained, known as black-ash. This black-ash process takes place in accord- ance with the following equation Na 2 S0 4 + CaC0 3 + 2C = Na 2 C0 3 + CaS + 2C0 2 . The change may be conveniently regarded as taking place in two stages, which proceed simultaneously, according to the equations— Na 2 S0 4 + 2C = Na 2 S + 2C0 2 . Na 2 S + CaC0 3 = CaS 4- Na 2 C0 3 . (3.) The process of extracting and purifying the sodium car- bonate contained in the black-ash. (1.) The Salt-cake Process— The first stage of this process is usually carried on in a large cast-iron pan (D, Fig. 132), built into a furnace in such a manner that it shall be heated as uniformly as possible. The charge of common salt is placed in the covered pan, and the requisite quantity of sulphuric acid is then run in. Hydro- chloric acid is given off in torrents, according to the first of the above equations, and the gas is led away by the pipe E in the arched roof, to the condensing-towers, where it is absorbed by water (see Hydrochloric Acid, page 331). The mixture is heated until it 492 Inorganic Chemistry begins to stiffen into a solid mass, when the damper h is raised, and the mass is raked out of the pan on to the hearth of the roaster , or reverberatory furnace, b. Here it is exposed to the hot gases from the coke fire a , which sweep over it, and ultimately raise its tempera- ture nearly to a red heat, whereby the second of the above reactions is completed. The acid gas, together with the fire gases, leave the roaster by the chimney e, and are also led to condensing - towers, where the hydrochloric acid is ab- sorbed. The mass is from time to time raked, or worked, by means of side - openings, A or “working doors,” in the m roaster, and as soon as the o' operation is completed the salt-cake is withdrawn. The salt-cake so obtained, usually contains from 95 to 96 per cent, of normal sodium sulphate, Na2S0 4 ; the remaining 4 or 5 per cent, consisting of hydrogen sodium sulphate, NaHSO^ un- decomposed sodium chloride, and such impurities as were originally present in the salt. (2.) The Black-ash Process. — The salt-cake is mixed with limestone (or chalk) and coal dust (slack), and heated in a reverberatory furnace, known as the black-ash, or balling , furnace. As the mixture softens with the heat, it requires to be thoroughly mixed together, Sodium Carbonate 493 which, in the older forms of furnace (still used in many places), is accomplished by manual labour. Fig. 133 shows such a furnace in section. The materials are introduced by the hopper k on to the hearth z, where they are exposed to the hot gases from the fire a ; and as the decom- position proceeds, they are raked along to the more strongly-heated front por- tion of the hearth h. During this pro- cess, carbon dioxide is freely evolved, the escaping bubbles of gas giving the semi-fluid mass the appearance of boil- ing. As the temperature rises, and the process approaches completion, the mass thickens, when it is worked up into large balls by means of rakes or paddles. At this stage, carbon mon- oxide begins to be evolved, the bubbles of which, bursting from the doughy material, become ignited and burn upon its surface as small jets of flame, coloured yellow by the soda. As soon as these appear, the ball is quickly withdrawn from the furnace. The for- mation of carbon monoxide, at the high temperature reached at this point in the process, is due to the action of carbon upon the limestone, according to the equation — CaC0 3 + C = CaO + 2CO, excess of these materials being inten- tionally present in the mixture. The effect of the escaping carbon monoxide at this point in the process, in rendering the black-ash light and porous (an impor- tant consideration in view of the next operation), is similar to that of baking-powder when used for cooking purposes. The heated gases from the furnace are made to pass over large evaporating pans, P , where liquors from a subsequent process are con- centrated. 494 Inorganic Chemistry In the more modem forms of black-ash furnace, the mixing and working up of the materials is accomplished mechanically, by means of a revolving hearth. Fig. 134 shows the general arrange- ment of a revolving black-ash furnace. The mixture is placed in the cylinder e , which is made to slowly revolve upon its horizontal axis. The heated gases from the fire a pass through this revolv- ing hearth ; they are then conveyed through a dust-chamber, w, and finally over concentrating-pans. Limestone and two-thirds of the coal are first thrown into the furnace, and heated until the blue flame of burning carbon monoxide makes its appearance, when the salt-cake, along with the rest of the coal, is added, and the process continued until the yellow flames appear upon the surface of the mass. The contents of the cylinder are then thrown out into iron trucks beneath. Black-ash is a mixture of variable composition, containing — Sodium carbonate, Na 2 CO s Calcium sulphide, CaS . Calcium carbonate, CaC0 3 Coke .... Calcium oxide, CaO from 40 to 45 per cent. „ 3° „ 33 >> „ 6 ,, 10 „ » 4 » 7 n » 2 „ 6 ,, And smaller quantities of sodium chloride, sodium sulphate, sodium sulphite, sodium sulphide, sodium thiosulphate, oxides of iron, alumina. (3.) Lixivialion of Black-ash . — The lixiviation of black-ash is carried on in a series of tanks, so arranged that the liquid can be made to pass from one to the other. The action of water upon the black-ash is more than a simple process of dissolving the sodium carbonate from the mixture, for in the presence of water, chemical action takes place between some of the ingredients. Thus, the lime reacts upon sodium carbonate, forming sodium hydroxide, hence the tank liquor always contains caustic soda in varying quantities. Under certain conditions of temperature and dilution, the calcium sulphide also reacts upon the sodium car' bonate, forming sodium sulphide and calcium carbonate, thus — CaS + Na^Og = CaCO s + Na2S. Also by the oxidising influence of atmospheric oxygen, calcium sulphide, CaS, is converted into calcium sulphate, CaSO^ which, Sodium Carbonate 495 in its turn, is acted upon by the sodium carbonate, involving loss of this product — CaS0 4 + Na 2 C0 3 = CaCO s + Na 2 S0 4 . The process of lixiviation is carried on as rapidly as possible, and at temperatures ranging from about 30 ° (for the dilute liquors) to about 6 o° (for those more concentrated) ; for the formation of sodium sulphide diminishes as the concentration of the liquid increases. The tank liquor, after settling, is then either at once concentrated by evaporation, when the soda crystallises out, leav- ing the caustic soda in the mother liquor, or it is submitted to the action of carbon dioxide, whereby both the caustic soda and the sodium sulphide are converted into sodium carbonate, thus — 2NaHO + C0 2 = Na 2 CO s + H 2 0. Na 2 S + C0 2 + H 2 0 = Na 2 C0 3 + H 2 S. The concentration of the tank liquor is accomplished in the shallow pans above-mentioned, by means of the waste heat from the black-ash furnace ; and the product obtained by evaporating the liquid, is usually calcined at a red heat in an ordinary reverberatory furnace. This substance is known as soda-ash, and when dissolved in water, and the solution allowed to crystallise, the so-called soda crystalj are obtained, having the composition Na 2 CO 3 , 10 H 2 O. II. The Ammonia-Soda Process. — This. process is based upon the fact, that hydrogen ammonium carbonate (< bicarbonate of ammonia) is decomposed by a strong solution of sodium chloride, according to the equation — H(NH 4 )C0 3 + NaCl = HNaC0 3 + NH 4 C1. In practice, the brine is first saturated with ammonia gas, and the cooled ammoniacal liquid is then charged with carbon dioxide, under moderate pressure, in carbonating towers. The hydrogen sodium carbonate ( bicarbonate of soda), being much less soluble, separates out, leaving the more soluble am- monium chloride in solution, from which the ammonia is recovered by subsequent treatment with lime. The hydrogen sodium carbonate is converted into normal sodium carbonate by calcination, and the carbon dioxide evolved is again utilised in carbonating a further quantity of ammoniacal brine — 2HNaCO a = Na 2 C0 3 + C0 2 + H a O. 49^ Inorganic Chemistry Sodium carbonate crystallises in large, transparent, monosym- metric crystals, commonly known as “ soda,” or “ washing-soda,” having the composition Na 2 CO 3 , 10 H 2 O. On exposure to the air, the crystals give up water, and become effloresced upon the surface, and finally fall to powder, having the composition Na2C0 3 ,H 2 0. When crystallised from hot solutions, it forms rhombic crystals, containing 7 H 2 0 . The solubility of sodium carbonate in water increases with rise of temperature, reaching a maximum of 32. 5 0 , when 100 parts of water dissolve 59 parts of the salt. Above this temperature the solubility falls, and at ioo° the amount dissolved is 45.4 parts. Hydrogen Sodium Carbonate ( bicarbonate of soda), HNaC 0 3 , may be obtained by the action of carbon dioxide upon the normal carbonate, either in solution, or as crystals— Na 2 CO 3 , 10 H 2 O + C 0 2 = 2 HNaCO s + 9 H 2 0 . The greater part of the bicarbonate of soda of commerce is obtained in the ammonia-soda process above described. This salt is less soluble in water than the normal carbonate. Thus, 100 parts of water at different temperatures dissolves the following quantities of these compounds — IO°. 20*. 30*. 40'. Na 2 C 0 3 . . . 12.6 21.4 38.1 50 parts. HNaC 0 3 . . 8.8 9.8 10.8 11.7 When a solution of hydrogen sodium carbonate is heated, the salt gives off a portion of its carbon dioxide, and on cooling, the solution deposits crystals having the composition Na2C0 3 ,2HNaC0 3 ,2H 2 0, known as sodium sesquicarbonate. On continued boiling, the salt is completely converted into the normal carbonate. Sodium sesquicarbonate occurs as a natural deposit in Egypt, Africa, South America, and elsewhere, known as trona , from which the name natrium is derived. Sodium Sulphate {Glaubers salt), NagSO*, occurs native in the anhydrous condition as the mineral thenardite, and as a double sulphate of sodium and calcium, Na^O^CaSO^ in the mineral G lauberite. It is manufactured in immense quantities in the first ( salt-cake ) process in the alkali manufacture, by the Leblanc method. It is also obtained in large supplies from the Stassfurt deposits. Sodium Nitrate 497 by double decomposition between magnesium sulphate (from kieserite) and sodium chloridej The solution of the mixed salts, when cooled a few degrees below o°, deposits sodium sulphate, and the soluble magnesium chloride remains in solution — 2NaCl + MgS0 4 = Na 2 S0 4 + MgCl 2 . Sodium sulphate is also manufactured by the action of sulphur dioxide and oxygen upon sp.dium chloride. This is known as Hargreaves process. The reaction is expressed by the equation— 2NaCl + S0 2 + O + H 2 0 = Na 2 S0 4 + 2HC1. This process is, in essence, the production of sodium sulphate from sodium chloride, and , the constituents of sulphuric acid, with- out the intermediate manufacture of the acid. The gases from pyrites burners, similar to those used by the “ vitriol ” manufacturer, together with steam, are passed through a series of cast-iron cylinders containing sodium chloride, and maintained at a tem- perature of 500° to 550°. Many days are required for the com- plete conversion of the chloride into sulphate by this process. Sodium sulphate crystallises in colourless prisms belonging to the monosymmetric system, containing ten molecules of water: when exposed to the air the crystals effloresce, and when heated to 33 0 , they melt in their own water of crystallisation (see page 132). When sodium sulphate is heated with sulphuric acid, in the pro- portions required by the following equation, hydrogen sodium sulphate is formed — Na 2 S0 4 + H 2 S0 4 = 2HNaS0 4 . Sodium Nitrate, NaN0 3 , occurs associated with other salts, in Bolivia and Peru, as cubical nitre , or Chili saltpetre . The crude salt is purified by solution in water, and crystallisation. It forms rhombohedral crystals, isomorphous with calcspar. Sodium nitrate is very soluble in water. 100 parts of water dis- solve at o°, 68.8 parts ; at 40°, 102 parts ; and at ioo°, 180 parts, of the salt. When exposed to the air, the salt absorbs moisture, and on this account cannot be employed as a substitute for potassium nitrate in the manufacture of gunpowder, or in pyrotechny. Its chief uses are for the manufacture of nitric acid ; for the manufacture 49 8 Inorganic Chemistry of potassium nitrate by double decomposition with potassium chloride ; and as an ingredient in artificial manures. Sodium Phosphates. — The most important of these compounds is the hydrogen disodium orthophosphate, or common phosphate of soda , HNa 2 P 0 4 . This salt is prepared on a large scale, by adding sodium carbonate to phosphoric acid until the solution is alkaline, and then filtering and evaporating the solution, when large transparent prisms, belonging to the monosymmetric system, are deposited, having the composition HNa2P0 4 ,12H 2 0. Exposed to the air the crystals effloresce, and when heated become an- hydrous. The salt melts at 35 0 . One hundred parts of water at io° dissolve 4.1 parts ; at 50°, 43.3 parts ; and at ioo°, 108.2 parts, of the anhydrous salt. Normal Sodium Orthophosphate, Na 3 P 0 4 , is obtained from hydrogen disodium phosphate, by evaporating a solution of the latter salt with sodium hydroxide, until the liquid crystallises — HNa 2 P 0 4 + NaHO = Na 3 P 0 4 + H 2 0 . This salt contains twelve molecules of water, and forms thin six-sided prisms. Its aqueous solution is strongly alkaline, and absorbs atmospheric carbon dioxide, with the formation of sodium carbonate and hydrogen disodium phosphate, thus — Na 3 P 0 4 + C 0 2 + H 2 0 = HNajjPC^ + HNaC 0 3 . Dihydrogen Sodium Orthophosphate, HgNaPO* is obtained when phosphoric acid is added to ordinary phosphate of soda, until the liquid gives no precipitate with barium chloride. On evapo- rating the solution, the salt crystallises. HNa 2 P 0 4 + H 3 P 0 4 = 2 H 2 NaP 0 4 . The aqueous solution of this salt is acid. Hydrogen Sodium Ammonium Phosphate (microcosmic salt), HNa(NH 4 )P 0 4 , is obtained by adding a strong solution of com- mon sodium phosphate to ammonium chloride — HNa 2 P 0 4 + NH 4 C 1 = NaCl + HNa(NH 4 )P 0 4 . The orthophosphates are readily converted into pyro- and meta- phosphates (see page 436). Lithium 499 LITHIUM. Symbol, Li. Atomic weight = 7.01. Occurrence. — Lithium is only found in combination with other elements. It is a constituent of a few somewhat rare minerals, as fietalite, 30SiO 2 ,4Al 2 O 3 ,Na 2 O,2Li 2 O ; spodumene, 15Si0 2 ,4Al 2 0 3 , 3Li 2 0 ; lepidolite , or lithium mica , 9Si0 2 ,3Al 2 0 3 ,K 2 0,4LiF. By means of the spectroscope, lithium compounds have been detected in sea water, and in most spring and river waters. In a few cases spring waters are met with which contain considerable quantities of lithium salts. Thus, W. A. Miller found as much as 0.372 gramme of lithium chloride in 1 litre of the water of a spring near Redruth in Cornwall. Mode of Formation. — Lithium is obtained by the electrolytic decomposition of the fused chloride. For this purpose the dry salt is heated in a porcelain crucible, when it melts at a low red heat to a mobile liquid. A rod of gas carbon is made the positive electrode ; and a stout iron wire, one end of which is flattened out, is used for the negative pole, upon which the lithium is collected. On passing an electric current through the molten chloride, the metal forms as a bright globule upon the negative electrode. The wire is withdrawn and quickly dipped beneath petiroleum, and the solidified globule of lithium is then cut off with a knife. The reduced metal, in its passage from the crucible to the petroleum, is protected from oxidation by the film of fused chloride which coats it. Properties. — Lithium is a soft, silver-white metal, which soon tarnishes on exposure to the air. It is easily cut with a knife, being softer than lead, but harder than sodium. It may be pressed into wire, and two pieces of the metal may be made to adhere, or welded together, at the ordinary temperature. Lithium is the lightest known solid, its specific gravity being 0.59. Its extreme lightness is illustrated by the fact that the metal floats upon petroleum, a liquid which itself floats upon water. Lithium melts at 180 0 , and at a higher temperature it takes fire and burns with a bright white light. Lithium decomposes water at the ordinary temperature, liberating hydrogen and forming lithium hydroxide, LiHO ; but when a fragment of the metal is thrown upon cold water it does not melt, and even with boiling water the action is not attended by inflammation of the hydrogen. 500 Inorganic Chemistry When strongly heated in nitrogen the two elements unite, with feeble combustion, forming lithium nitride, NLi 3 . Lithium Oxide, Li 2 0, is formed when the metal bums in the air. It is also obtained by heating the nitrate. It dissolves in water, forming lithium hydroxide, Li HO. Lithium Hydroxide is produced by the prolonged boiling of lithium carbonate with milk of lime, the carbonate of this metal, unlike potassium and sodium carbonates, being only very slightly soluble in water. Lithium Carbonate, Li 2 C0 3 , is obtained as a white precipitate when a solution of either potassium, sodium, or ammonium car- bonate is added to a solution of either chloride or nitrate of lithium. The compound is only slightly soluble in cold water, ioo parts of water at 13 0 dissolving 0.77 parts of the carbonate. Lithium Phosphate, Li 3 P0 4 , is precipitated as a crystalline powder, by the addition of hydrogen disodium phosphate to a solution of a lithium salt. In the presence of sodium hydroxide the precipitation is complete, and the formation of this compound is employed as a quantitative method for estimating lithium. The crystals contain 2H 2 0, which they lose when heated. Lithium phosphate is soluble in nitric, hydrochloric, and phosphoric acids, and from the latter solution, on evaporation, the dihydrogen phosphate is deposited (H 2 LiP0 4 ), as deliquescent, and very soluble crystals. The chloride, nitrate, and sulphate of lithium are obtained by dissolving the carbonate in the respective acids. The salts are readily soluble in water. Rubidium and Caesium.* — These two rare elements, which were first dis- covered by Bunsen in the waters of Diirkheim, in the years 1860-61, are met with, associated with sodium and potassium, in certain minerals, such as lepidolites (lithium mica), porphyrites, and in camallite. They are also found in many mineral waters, in the mother liquors from the evaporation of sea water, and in the ashes of plants. Although widely distributed, the quantities present are extremely minute, one of the richest lepidolites in which these metals occur, containing only 0.24 per cent, of rubidium oxide. The rare mineral pollux , a silicate of aluminium and caesium, containing also iron calcium and sodium, is the only known mineral in which either of these two elements occurs as an essential constituent. The analysis of Pisani (1864) gives 34.07 per cent, of caesium oxide in this substance. Rubidium is obtained by heating the carbonate with carbon (the charred tartrate), as in the older method for the preparation of sodium and potassium. * For detailed descriptions of these elements and their compounds, the student is referred to larger works. Ammonium Chloride 501 Caesium cannot be isolated by this reaction, but is obtained by the electro- lysis of the fused cyanide, Cs(CN) (mixed with barium cyanide in order to render it more readily fusible). Rubidium melts at 38*5* caesium at 26.5. Rubidium gives a green vapour, and when sublimed in a vacuous tube yields a thin film of metal, which appears deep blue by transmitted light : when slowly sublimed in this way the metal forms small needle-shaped crystals. The compounds of these metals closely resemble those of potassium, from which they can only be distinguished by the different spectra they give. AMMONIUM SALTS. The monovalent group or radical (NH 4 ) is capable of replacing one atom of hydrogen in acids, thereby giving rise to a series of salts which are closely analogous to, and are isomorphous with, those of potassium. The radical (NH 4 ), to which the name ammonium is given, has never been isolated. When an amalgam of sodium and mercury is thrown into a solution of ammonium chloride, the mercury swells up into a honeycombed or sponge- like mass, which floats upon the surface of the liquid. This so- called ammonium amalgam was at one time thought to be a true amalgam of mercury with the metallic radical ammonium. It is now generally believed to consist of mercury which is simply inflated by the evolution of hydrogen and ammonia gas. When this sponge-like substance is subjected to changes of pressure, it is found to contract and expand in conformity to Boyle’s law : its formation may be represented by the equation — HgxNa y + yNH 4 Cl = yNaCl + xHg + yNH 3 + yH. In the course of a few minutes the inflated mass shrinks down, and ordinary mercury remains at the bottom of the solution, hydrogen and ammonia having been rapidly evolved. The ammonium salts are obtained for the most part from the ammoniacal liquor of the gasworks. This material is treated with lime, and distilled ; and the ammonia so driven off is absorbed in sulphuric or hydrochloric acid, giving rise to ammonium sulphate or chloride. Ammonium Chloride (sal ammoniac\ NH 4 C1. — The product obtained by absorbing ammonia from gas liquor in hydrochloric acid, is purified by sublimation. The crude material is heated in large iron pots, covered with iron dome-shaped vessels, into which the substance sublimes. Ammonium chloride crystallises in 502 Inorganic Chemistry arborescent or fern-like crystals (Fig. 135), consisting of groups of small octahedra belonging to the regular system. 100 parts of water at io° dissolve 32.8 parts, and at ioo°, 77 parts of the salt. On boiling the aqueous solution, dissociation to a small extent takes place, and a portion of the ammonia escapes with the steam ; the solution at the same time becoming slightly acid. Fig. 135. Ammonium Sulphate (NH 4 ) 2 S 0 4 .— The product obtained by the absorption of ammonia obtained from gas liquors, by sulphuric acid, is purified by recrystallisation, when it forms colourless rhombic crystals, isomorphous with potassium sulphate. 100 parts of water at the ordinary temperature dissolve 50 parts of the salt. The chief use of ammonium sulphate is for agricultural purposes, as a manure ; and for this use the crude salt, as first obtained, which is usually more or less coloured with tarry matters, is em- ployed. Ammonium sulphate is also used for the preparation of A mmonium Carbonate 503 ammonia alum, and other ammonium compounds, as well as m th Am^^lum° Carbonates.— Commercial (sal volatile) is obtained, by heating a mixture of a™ 1 " 0 "'™ 1 sulphate and ground chalk to redness in horizontal iron retorts o cylinders^ and conducting the vapours into 1 - where the carbonate condenses as a solid crust, afterwards purified by resublimation, when it is obt ^^ e ^ as a whhT fibrous mass. This substance is a mixture of hydrogen v u/mu , )COo and ammonium carbamate ammonium carbonate, H(JNH 4 )UU 3 , anu * trpat ed fNHi)C 0 9 (NH 2 ), and smells strongly ammomaca . When trea ed lith alcohol the ammonium carbamate is dissolved, leavmg the Ca No 0 p n mallm n lnium Carbonate, (NH 4 ) 2 C 0 3 , is obtained from the commercial compound, by passing ammonia gas into a strong aqueous solution, or by digesting the compound in strong aqueous ammonia. The carbamate present is converted into normal car- bonate by the action of the water, thus (NH 4 )C 0 2 (NH 2 ) + H a O = (NH 4 )C 0 3 (NH 4 ) = (NH 4 ) 2 C 0 3 ; and the ammonia converts the bicarbonate into the normal salt, thUS HfNH^COs + NH 3 = (NH 4 ) 2 C 0 3 . Normal ammonium carbonate on exposure to the air gives off ammonia, and passes back into hydrogen ammonium carb °" a *- When heated to 60“ the salt breaks up into carbon dioxide, an Hydrogen d Ammonium Carbonate, H(NH 4 )C 0 3 , may also be obfled by passing carbon dioxide into a solution of the normal Salt “ (NH 4 ) 2 C 0 3 + C 0 2 + H 2 0 = 2H(NH 4 )CO, It forms large lustrous crystals belonging to the Zombie system which when dry, do not smell of ammonia. 100 parts of water dissolve 12^5 parts of this salt. At ordinary temperatures this solution on exposure to the air slowly gives off carb “^ 0 ^“ t0 becomes alkaline ; and when heated above 36 the b ^ ld beginS J° effervesce, owing to the rapid evolution of carbon dioxide. T 3S with the normal carbonate a double salt, analogous to 5°4 Inorganic Chemistry sodium sesquicarbonate, and having the composition (NHA,CO, 2H(NH 4 )C0 3 , h 2 o. Ammonium Thiocyanate, NH 4 S(CN), is prepared by adding aqueous ammonia to an alcoholic solution of carbon disulphide, and allowing the mixture to stand, when ammonium thiocarbonate is formed, thus — 6NH„ + 3H 2 0 + 3CS 2 = 2(NH 4 ) 2 CS, + (NH 4 ) 2 C0 3 . On heating this solution, the ammonium thiocarbonate is de- composed with evolution of sulphuretted hydrogen — (NH 4 ) 2 CS 8 = 2H 2 S + NH 4 S(CN). Ammonium thiocyanate (known also as ammonium sutyho- cyanate) forms colourless crystals, which are extremely soluble both in water and alcohol. The solution in water is attended with considerable absorption of heat : thus, if 20 grammes of the salt be dissolved in 25 cubic centimetres of water at i8°,the temperature of the liquid falls to - 13 0 . CHAPTER V THE ELEMENTS OF GROUP I. ( FAMILY B.) Copper, Cu 63.18 Silver, Ag 107.66 Gold, Au 196.8 The elements of this family present many striking contrasts to those of the other family belonging to the first group. These three metals are not acted upon by oxygen, or by water, at ordinary temperatures ; they are all found native in the un- combined state, and on this account are amongst the earliest metals known to man. The alkali metals, on the other hand, are instantly oxidised on exposure to air, they decompose water at the ordinary temperature, are never found native, and are amongst the most recently discovered metals. With the exception of sodium and potassium, which are used in a few manufacturing processes, the alkali metals, as such, are of little practical service to mankind, whilst the metals of this family are amongst the most useful of all the metals, and are the three universally adopted for coinage. Many of the compounds of the elements of this family, are similarly constituted to those of the alkali metals : thus, with oxygen and with sulphur we have Cu 2 0, Ag 2 0, Au 2 0, and Cu 2 S, Ag 2 S, Au 2 S, corresponding to Li 2 0 and K 2 S. With the halogens they all form compounds of the type RX. Although the three elements, copper, silver, and gold, fall into the same family, upon the basis of the periodic classification of the elements, they are in many respects widely dissimilar. Thus, silver is consistently monovalent, while copper is divalent, forming compounds of the type CuX 2 , and gold is trivalent, giving compounds AuX 3 . The chlorides, AgCl and Cu 2 Cl 2 , on the other hand, are both insoluble in water, are both soluble in ammonia, and both absorb ammonia. In many of their physical attributes, these metals show a regular 505 506 Inorganic Chemistry gradation in their properties. Thus, as regards malleability and ductility, silver is intermediate between copper and gold, the latter possessing these properties in the highest degree. With respect to their tenacity, silver is again intermediate, copper being the most, and gold the least tenacious of the three. COPPER. Symbol, Cu. Atomic weight = 63.18. Occurrence. — Copper is found in the elementary condition in various parts of the world, notably in the neighbourhood of Lake Superior, where native copper occurs in enormous masses. In combination, copper is a very abundant element, and is widely distributed, the most important of these natural compounds being the following — Ruby ore Copper glance Copper pyrites . Cu s O. . Cu 2 S. . Cu2S,Fe 2 S3. Purple copper ore 3Cu2S.Fe.jS3. Malachite . . CuC 0 3 ,Cu(H 0 ) 2 . Azurite . . 2 CuC 0 3> Cu(H 0 ) 2 . Modes of Formation. — The methods by which copper is obtained from its ores, vary with the nature of the ore. From ores containing no sulphur, such as the carbonates and oxide, the metal may be obtained by a method known as the reducing process , which consists in smelting down the ore in a blast-furnace with coal or coke, when the metal is reduced according to the equation — Cu 2 0 + C = CO + 2Cu. In the case of mixed ores, containing sulphides, the process (known as the English method) consists of six distinct stages — (1.) The ores, which contain on an average 30 per cent, of iron and 13 of copper (the remainder being chiefly sulphur and silica), are first calcined ; usually in a reverberatory furnace, whereby a portion of the sulphur is burnt to sulphur dioxide, and the metals are partially oxidised. (2.) The second step consists in fusing the calcined ore ; when the copper oxides, formed during calcination, react upon a portion of the ferrous sulphide with the formation of cuprous sulphide and ferrous oxide, thus — Cu 2 0 + FeS = Cu 2 S + FeO. 2CuO + 2FeS = Cu 2 S + 2FeO + S. Copper 507 The oxide of iron combines with the silica already present (or which is added in the form of metal slag obtained front ‘he fourt orocess) to form a fusible silicate of iron, or slag, which contains U tk or no copper. This is run off, and a fused regulus remains consist ng of cuprous and ferrous sulphides known as ^-metal and containing from 30 to 35 P« cent. of copper Th>s molten 1 v,;^Vi a comDOsition very similar to copper pynte , « -»« *■ «* « ■ “ k “' d " * ” ,ul “ JJTS fvssi s~5 metal , or white-metal, contains from 60 to ^ .5 P“ ““ w °[ ^ ( 5 .) The fifth operation consists m roast ng oi^ch, as the t^nperature rises, reacts upon another portion of cuprous sulphide, thus 2 Cu 2 0 + Cu 2 S = 6Cu + S0 2 . At the same time any remaining ferrous sulphide is converted into oxide, thus- 3c ^ o + Feg _ 6Cu + FeD + SO, The metallic copper so obtained, presents a blistered appearance, and on this account is known as blister- copper. nroce ss. (f) ) This impure copper is lastly subjected to a r g p r- u <~ r tory furnace, in an oxidising atmosphere. L oxidise • in the metal, such as iron, lead, and arsenic, are the first to oxidise , and the oxides either volatilise, or combine with the siliceous matter 5° 8 Inorganic Chemistry oxide which it still contains, the molten mass is stirred with poles of wood, and a quantity of anthracite is thrown upon the surface to complete the reducing process. Wet Process.— Copper is extracted from the burnt pyrites obtained in enormous quantities in the manufacture of sulphuric acid, which contains about 3 per cent, of copper. Although too poor in copper to be submitted to the smelting process, it is found that when calcined with 12 to 15 per cent, of common salt, the copper is all converted into cupric chloride. On lixiviating the calcined mass with water, the cupric chloride goes into solution, and metallic copper can be precipitated from it by means of scrap-iron. Properties. Copper is a lustrous metal, having a characteristic reddish-brown colour. The peculiar copper-red colour of the metal is best seen, by causing the light to be several times reflected from the surface before reaching the eye. Native copper is occasionally found crystallised in regular octa- hedra, and small crystals of the same form may be artificially obtained, by the slow deposition of the metal from solutions of its salts by processes of reduction. Copper is an extremely tough metal, and admits of being drawn into fine wire, and hammered out into thin leaf. Its ductility and malleability are greatly diminished by admixture with even minute quantities of impurities. When heated nearly to its melting-point, copper becomes sufficiently brittle to be powdered. The specific gravity of pure copper, electrolytically deposited, is 8.945, which by hammering is increased to 8.95. Copper is only slowly acted upon by exposure to dry air at ordinary temperatures ; but in the presence of atmospheric moisture and carbon dioxide, it becomes coated with a greenish basic carbonate. When heated in air or oxygen, it is converted into black cupric oxide, which flakes off the surface in the form of scales. When volatilised in the electric arc, copper gives a vapour having a rich emerald green colour. Copper is readily attacked by nitric acid, either dilute or con- centrated, with the formation of copper nitrate and oxides of nitrogen (page 221). Dilute hydrochloric and sulphuric acids are without action upon copper when air is excluded, but slowly attack it in the presence of air, or in contact with platinum. Cold concentrated sulphuric acid does not act upon copper ; but, when heated, copper sulphate and sulphur dioxide are formed (page 377). 509 Cuprous Oxide Finely divided copper is slowly dissolved by boiling concen- trated hydrochloric acid, with evolution of hydrogen and formation of cuprous chloride : — 2Cu + 2HC1 = Cu 2 Cl 2 + H* In the presence of air, copper is acted upon by a solution of ammonia, the oxide dissolving in the ammonia forming a deep blue solution. Copper is an extremely good electric conductor, being only second to silver in this respect ; it is therefore extensively em- ployed for cables, or leads, for purposes of telegraphy and electric lighting. Copper possesses the property, in a high degree, of being de- posited in a coherent form by the electrolysis of solutions of its salts. On this account it is extensively used in processes of electrotyping. Alloys of Copper. — The most extensive use of copper is in the formation of certain alloys, many of which are of great technical value. The following are among the most important English brass . Copper 2 parts Zinc i part Dutch brass ( Tombac ) . „ 5 » 1 » Muntz metal . » 3 » ,, 1 » Gun metal . . • » 9 » Tin i „ Aluminium bronze . „ 9 Aluminium i „ Oxides Of Copper.— Two oxides of copper are well known, namely, cuprous oxide sub-oxide ), Cu 2 0, and cupric oxide ( copper monoxide ), CuO. Cuprous Oxide, Cu 2 0, occurs native as red copper ore. It is formed when finely divided copper is gently heated in a current of air or when a mixture of cuprous chloride and sodium carbonate is gently heated in a covered crucible. Cu 2 Cl 2 + Na^Og = 2NaCl + C© 2 + Cu 2 0. Cuprous oxide is also obtained when an alkaline solution of a copper salt is reduced by grape sugar. Cuprous oxide is insoluble in water ; it is converted into cuprous chloride by strong hydrochloric acid. Nitric acid converts it intc 5 10 Inorganic Chemistry cupric nitrate with the evolution of oxides of nitrogen. When acted upon by dilute sulphuric acid, it is partly reduced to metallic copper, and partly oxidised into copper sulphate, thus — Cu 2 0 + H 2 S0 4 = CuS0 4 + Cu + H 2 0. When heated with the strong acid it is entirely oxidised, thus— Cu 2 0 + 3H 2 S0 4 = 2CuS0 4 + S0 2 + 3H 2 0. Cuprous oxide fuses at a red heat, and when melted with glass, imparts to the latter a rich ruby-red colour. Cupric Oxide, CuO, occurs as the rather rare mineral, tenorite. It is formed when copper is strongly heated in the air or in oxygen, or by gently igniting either the nitrate, carbonate, or hydroxide.’ It is a black powder, which rapidly absorbs moisture from the air. When heated, it first cakes together and finally fuses, giving up a part of its oxygen, and leaving a residue consisting of Cu0,2Cu 2 0. When heated in a stream of carbon monoxide, marsh gas, or hydrogen, it is reduced to the metallic state. Similarly, when mixed with organic compounds containing carbon and hydrogen, it oxidises these elements to carbon dioxide and water, itself being reduced : on this property depends its use in the ultimate analysis of organic compounds. Cupric Hydroxide, Cu(HO) 2 , is the pale blue precipitate pro- duced when sodium or potassium hydroxide is added in excess to a solution of a copper salt. The compound, when washed, may be dried at ioo° without parting with water ; but if the liquid in which it is precipitated be boiled, the compound blackens, and is con- verted into a hydrate having the composition Cu(HO) 2 ,2CuO. Cuprous hydrate dissolves in ammonia, forming a deep blue liquid, which possesses the property of dissolving cellulose (cotton wool, filter paper, &c.) Salts Of Copper. — Copper forms two series of salts, namely, cuprous and cupric salts. The former, 'which are colourless, readily pass by oxidation into cupric salts, and serve therefore as powerful reducing agents, and are mostly insoluble in water. The cupric salts in the hydrated condition, are either blue or green in colour ; the anhydrous cupric salts are colourless or yellow. The normal salts are mostly soluble in water. Copper salts impart to a non-luminous flame a blue or green colour, and 5i Cupric Chloride on this account are employed in pyrotechny. The soluble salts Cuprous Chloride, Cu 2 Cl 2 , may be obtained by dissolving cuprous oxide in hydrochloric acid. It is more readily prepared by boiling a solution of cupric chloride in hydrochloric acid, with copper turnings or foil. The nascent hydrogen, liberated by the action of the hydrochloric acid upon the copper, reduces the cupric chloride to cuprous chloride. The liquid is then poured into water, which causes the precipitation of the cuprous chloride as a white crystalline powder. A mixture of zinc dust and copper oxide added to strong hydro- chloric acid, also yields cuprous chloride, the nascent hydrogen in this case being derived from the action of the acid upon the zinc, and this causes the reduction of cupric chloride formed by the action of the acid upon the cupric oxide. Cuprous chloride melts when heated, and volatilises without decomposition. It is insoluble in water, but dissolves in hydro- chloric acid, ammonia, and alkaline chlorides. These solutions, on exposure to the air, absorb oxygen, turning first brown, and finally depositing a greenish-blue precipitate of copper oxychloride, CuC1 2 ,3Cu0,4H 2 0. This compound occurs native as the mineral atacamite. Solutions of cuprous chloride also -absorb carbon mon- oxide forming a crystalline compound, believed to have the com- position, C0Cu 2 C1 2 ,2H 2 0. They also absorb acetylene (see page Cuprous bromide, Cu 2 Br 2 ; iodide, Cu 2 I 2 ; and fluoride, Cu 2 F 2 , are also known. Cupric Chloride, CuCl 2 . — This compound is formed when copper is dissolved in nitro-hydrochloric acid, or when cupric oxide, carbonate, or hydroxide, are dissolved in hydrochloric acid. It is also produced when copper is burnt in chlorine. Cupric chloride is readily soluble in water, forming a deep green solution, which, on being largely diluted, turns blue. The salt crystallises in green rhombic prisms, with 2 H 2 0 . When heated, it loses its water, and at a dull red heat is converted into cuprous chloride, with evolution of chlorine (see page 317). Cupric chloride forms three compounds with ammonia. The anhydrous salt absorbs ammonia gas, forming a blue compound, CuC1 2 ,6NH 3 . When ammonia is passed into aqueous cupnc chloride, the solution deposits deep blue quadratic octahedral crystals of the compound, CuC1 2 ,4NH 3 ,H 2 0. Both these sub* 512 Inorganic Chemistry st a nc es , when moderately heated, yield the green compound thus WhlCh at a hi ^ her temperature is decomposed, 6(CuC1 2 ,2NH 3 ) = 3Cu a Cl a + 6NH 4 C1 + 4NH 3 + N 2 . Cupric bromide, CuBr 2 , and fluoride, CuF^ are known, but the iodide is unknown. Cupric Nitrate, Cu(N0 3 ) 2 ,3H 2 0, may be obtained by the action of nitric acid upon cupric oxide, hydroxide, carbonate, or the metal itself. It is deposited from the solution in deep blue deliquescent crystals, soluble in alcohol. When heated to about 65°, the crystals lose nitric acid and water, and are converted into the basic nitrate, Cu(N0 3 ) 2 ,3Cu(H0) 2 . The normal salt, there- fore, cannot be obtained anhydrous. Cupric nitrate is a caustic, powerfully oxidising substance. If the moist salt be rubbed in a mortar with a quantity of tinfoil, the tin is quickly converted into oxide, with considerable rise of temperature. When a solution containing copper nitrate and ammonium nitrate is evaporated, the mixture suddenly deflagrates when a certain degree of concentra- tion is reached. Cupric Sulphate {blue vitriol ), CuS0 4 ,5H 2 0, is the most important of all the copper salts. It is formed when either the metal or the oxide is dissolved in sulphuric acid. On a com- mercial scale, it is obtained from waste copper by first converting the metal into sulphide, by heating it in a furnace, and throwing sulphur upon the red-hot metal. Air is then admitted, and the sulphide is thereby oxidised into sulphate, which is dissolved in water, and crystallised. It is also manufactured from the sulphur ores of copper, by roasting them under such conditions that the iron is for the most part converted into oxide, while the copper is oxidised to sulphate. On lixiviating the roasted mass, the copper sulphate, with a certain amount of ferrous sulphate, is dissolved out. The ores may also be roasted so as to convert both the metals into oxides ; the mass is then treated with “chamber acid,” which dissolves copper oxide, leaving the iron oxide for the most part unacted upon. Cupric and ferrous sulphates cannot be entirely separated by crystallisation, as a solution of these salts deposits a double sulphate of the two metals. If, however, the amount of iron pre- sent is comparatively small, the first crop of crystals obtained, is Copper Sulphides 5 1 3 moderately pure copper sulphate. The copper ^moved from the mother liquors by precipitation upon plates of iron, ‘J? copper so obtained is converted into sulphide, as above describe . Copper sulphate forms large blue asymmetric crystals with 5H..O. At too” it is converted into a blu.sh-wh.te salt, CuS 0 4 ,H 2 0 , and at 220° to 240° it becomes anhydrous. je anhydrous salt is white, and extremely hygroscopic, and is used both for the detection and removal of small quantities of water in organic liquids. , . One hundred parts of water at 10 dissolve 36.6 parts, and a ioo°, 203.3 parts, of the crystallised salt. Several basic sulphates of copper are known : thus, when the normal salt is submitted to prolonged heating it is converted mto an amorphous yellow powder, consisting of CuSO„CuO, which, when thrown into cold water, forms an insoluble green compound, CuS 0 4 , 3 Cu(H 0 ) 2 , and on treatment with boiling water yields CuS 0 4 , 2 Cu(H 0 ) 2 . Copper sulphate forms several compounds with ammonia. Thus, the anhydrous salt readily absorbs ammonia gas, forming the compound, CuS 0 4 , 5 NH ? . When excess of ammonia is added to a solution of copper sulphate, the deep blue solution deposits blue crystals of CuS 0 4 ,H 2 0 , 3- * 5 ° this compound is converted into CuS 0 4 , 2 NH 3 , and at 200 it loses one more molecule of ammonia, leaving CuSO^N H 3 . Cupric Carbonates.— The normal carbonate has not been obtained. The two most important basic carbonates are (1, CuCO„Cu(HO) 2> occurring native as malachite , and obtained when sodium carbonate is added to a solution of copper sulphate (the green deposit which appears upon copper, when exposed to atmos- pheric moisture and carbon dioxide ( verdigris ), is the same com- pound) ; and (2) 2 CuC 0 3 ,Cu(H 0 ) 2 , occurring as the mineral azunt' Sulphides of Copper. — Two sulphides are known, correspond- ing to the two oxides. Cuprous sulphide , Cu 2 S, occurs in nature as copper glance, in the form of grey metallic-looking rhombic crystals. It is produced when copper burns in sulphur vapour, or when an excess of copper filings is heated with sulphur. Cupric Sulphide, CuS, is met with in nature as the mineral indigo- copper. It is obtained when either copper or cuprous sulphide is heated with sulphur to a temperature not beyond 1 14 0 ; so obtained, the compound is blue. As a black precipitate, it is formed when sulphuretted hydrogen is passed into solutions of cupric salts. 5*4 Inorganic Chemistry SILVER. Symbol, Ag. Atomic weight = 107.66. Occurrence. — Silver is found uncombined, occasionally in masses weighing several cwts Such native silver usually contains copper, gold, and other metals. Amongst the more important natural compounds of silver are the following : — Argentite, or silver glance Pyrargyrite, or ruby silver ore Proustite, or light-red silver ore Stephanite .... Polybasite .... Stromeyerite .... Horn silver .... . Ag 2 S. . 3Ag 2 S,Sb 2 S 3 , or Ag 3 SbS 3 . . 3Ag 2 S,As 2 S 3 „ Ag 3 AsS 3 . . 5Ag 2 S,Sb 2 S 3 „ Ag 6 SbS 4 . . 9(Ag 2 S,Cu 2 S),Sb 2 S 3 ,As 2 S 3 . . Ag 2 S,Cu2S. • AgCl. Silver is present also in most ores of lead, notably with galena (lead sulphide) ; argentiferous lead ores constituting one of the main supplies of silver. Modes of Formation. — This element may be obtained from its salts by the electrolysis of their aqueous solutions. The metal is so readily reduced from its compounds, that many organic substances, such as grape sugar, aldehyde, certain tartrates, &c., are capable of effecting its deposition. When a strip of zinc is introduced into silver nitrate solution, the silver is at once de- posited upon the zinc as a crystalline mass. Pure silver for analytical purposes may be prepared by pre- cipitating silver chloride, by the addition of hydrochloric acid to a solution of the nitrate, and reducing the chloride by boiling with sodium hydroxide and sugar, or by means of metallic zinc. In this way the metal is obtained as a fine grey powder. The chloride may also be reduced by fusion with sodium carbonate, when the silver is obtained as a button at the bottom of the crucible. The methods by which silver is obtained from its ores are very varied ; they may, however, be classed under th ree heads, namely — 1 . Processes involving the use of mercury. (Amalgamation processes.) 2 . Processes by means of lead. Silver 515 3. Wet processes. (1.) Amalgamation Processes.— These depend upon the fact that certain compounds of silver are reduced by mercury. The reduced silver then dissolves in the mercury, forming an amalgam, from which the silver is obtained, and the mercury recovered by distillation. The process, as still carried on in Mexico and South America, is the following. The ore is first crushed and then ground to a fine powder with water, and the mud so obtained is mixed with 3 to 5 per cent, of common salt, and spread upon the floor of a circular paved space, the mixing being effected by the treading of mules. After the lapse of a day, mercury is added, together with a quantity of roasted pyrites (known as magistral , and consisting of a crude mixture of cupric and ferric sulphates and oxides), and the materials thoroughly incorporated. Fresh mercury is added from time to time, during the several days required for the completion of the chemical decompositions that take place. The exact nature of these changes is not thoroughly understood, but it is probable that they involve first the formation of copper chlorides, by double decomposition between the copper sulphate and sodium chloride, and the subsequent action of these upon the silver sulphide present in the ore, thus— 2CuC1 2 + Ag 2 S = 2AgCl + Cu 2 Cl 2 + S. Cu 2 Cl 2 + Ag 2 S = 2AgCl + Cu 2 S. The silver chloride dissolves in the sodium chloride present, and is reduced by the mercury, with the production of mercurous chloride ( calomel ), which is ultimately lost in the washing— 2AgCl + 2Hg = Hg 2 Cl 2 + 2Ag. The amalgam is first washed, and freed from adhering particles of mineral, and is then filtered through canvas bags, whereby the excess of mercury is removed. The solid residue, containing the silver, is then submitted to distillation. In other amalgamation processes the ore is first roasted with salt, in order to convert the silver into chloride. The roasted ore is reduced to fine powder with water, and introduced into , revolving casks along with scrap iron, when the chloride is reduced according to the equation — 2AgCl + Fe = 2Ag + FeCl 2 , 5 1 6 Inorganic Chemistry and the reduced silver is then extracted by the addition of mercury, with which it amalgamates. In the modern amalgamation process, the finely crushed ore, with water, is placed in iron pans provided with revolving machinery, which serves the purpose of further grinding, and also of mixing. When the ore is reduced to an almost impalpable powder, mercury is added, and the machinery is kept in operation for a few hours, when the amalgamation is complete ; sometimes common salt and copper sulphate are added, either together or singly. Their pre- sence does not appear to be necessary to the process, except in so far as they aid in keeping the surface of the mercury clean, or “ quick ; ” for in the extremely finely divided condition to which the ore is reduced in this “ pan ” amalgamation process, the silver sulphide is readily acted upon by mercury, with the formation of mercuric sulphide — Ag 2 S + Hg = HgS + 2A g, and the silver so reduced, dissolves in the excess of mercury, from which it is finally separated by distillation. ( 2 .) Processes by Means of Lead. — When silver ores are smelted with lead, or with materials which yield metallic lead ; in other words, when silver ores are smelted with lead ores, an alloy of silver and lead is obtained, from which the silver can be separated. When the argentiferous lead is rich in silver, the alloy is submitted to cupellation , which consists in heating the metal in a reverbera- tory furnace, the hearth of which consists of a movable, oval-shaped, shallow dish, made of bone ash, known as a cupel , or test. The alloy is fed into this cupel from a melting-pot, and a blast of air is projected upon the surface of the molten metal. The lead is thus converted into litharge, and the melted oxide, by the force of the blast, is made to overflow into iron pots. As the oxidation of the lead reaches completion, the thin film of litharge begins to exhibit iridescent interference colours, which presently disappear, leaving the brilliant surface of the melted silver. The sudden appearance of the bright metallic surface is known as the flashing of silver. In the case of argentiferous lead too poor in silver to be directly cupelled, the alloy is submitted to one of two processes of con- centration, namely, the Pat tins on process^ or the Parked s process. The Pattinson process for desilverising lead, depends upon the fact that alloys of silver and lead have a lower melting-point than Silver 517 pure lead, and therefore when argentiferous lead is melted and allowed to cool, the crystals which first form, consist of lead which is nearly or quite pure, and the greater part of the silver is in the still liquid portion. The operation is carried out in a row of iron pots. A quantity of the metal is melted in one pot, and as it cools, the crystals which begin to form are removed by means of a perforated iron ladle, and transferred to the next pot on one side This operation is continued until a definite proportion (either two-thirds or seven-eighths, depending upon the proper- tion of silver) has been removed. The residue is then transferred to the neighbouring pot on the opposite side, and a second charge melted up in the first pot. As the neighbouring pots fill up, they are similarly treated, and in this way an alloy, gradually becoming richer and richer in silver, is passed along in one direction, and purer and purer lead is sent in the opposite way. The rich alloy is then cupelled. The Parked s process depends upon the fact that when zinc is added to a melted alloy of lead and silver, the zinc deprives the lead of the silver, and itself forms an alloy with it. The alloy of zinc and silver rises to the surface, and is the first portion to solidify, and can be removed. The operation is carried out in iron pots. The argentiferous lead is melted, and a quantity of zinc is thoroughly stirred into the molten mass, the amount of zinc depending upon the richness of the lead. As the mixture cools, the first portions to solidify are skimmed off with a ladle, and transferred to another pot. These skimmings, consisting of zinc, silver, and lead, are first liquated ; that is, carefully heated to such a temperature that the adhering lead melts, and flows away from the less fusible zinc silver alloy. The solid alloy is then distilled, and the residue, consisting of silver and lead, is submitted to ' cupellation. (3.) Wet Processes (Ziervogel Process ).— When argentiferous pyrites, or an artificially formed regulus containing sulphides of silver, copper, and iron, is roasted, the sulphides are first converted into sulphates ; and, as the roasting continues, first the iron, then the copper, and lastly, the silver sulphate is converted into oxide. By careful regulation, the process is continued until the whole of the iron and a part of the copper sulphates are decomposed. On lixiviating the roasted mass with water, the silver sulphate, together with the remaining copper sulphate, dissolves. From this solution the silver is precipitated by scrap copper. 5 1 8 Inorganic Chemistry The copper is recovered from the solution by precipitation with iron. The Percy- Patera Process .— In this method the ore is roasted with salt, and the silver chloride so formed is then extracted by means of sodium thiosulphate — Na,S,O s + AgCl = NaCl + NaAgS 2 0 3 . To the solution so obtained, sodium or calcium sulphide is added, which precipitates silver sulphide — 2NaAgS 2 0 3 + Na 2 S = Ag 2 S + 2Na 2 S 2 0 3 . The silver sulphide is then reduced by being roasted in a rever- beratory furnace. Properties.— Silver is a lustrous white metal, which appears yellow when the light is reflected many times from its surface before reaching the eye. It is unacted upon by atmospheric oxy- gen, but quickly becomes tarnished by traces of sulphuretted hydrogen in the air. Silver has the highest conductivity for heat and electricity of all the metals. It is extremely malleable and ductile, being second only to gold. Thin films of silver appear blue by transmitted light. Silver melts at 954 0 , and, when heated by the oxyhydrogen flame, may be readily made to boil, and distil. The pure metal employed by Stas for the determination of the atomic weight, was obtained by distillation in this way. When volatilised in the electric arc, the vapour of silver has a brilliant green colour. Molten silver absorbs as much as twenty-two times its volume of oxygen, which it gives up again (with the exception of 0.7 volume) on solidification. As the mass cools, the oxygen evolved often bursts through the outer crust of solidified metal with considerable violence, ejecting portions of the still liquid silver as irregular excrescences, as seen in Fig. 136. This pheno- menon is known as the “ spitting ” of silver. Small quantities of admixed metals prevent the absorption of oxygen. Silver is readily soluble in nitric acid, forming argentic nitrate. Fig. 136. Silver Oxides 5 ! 9 is Ur fi e 7 lo /1 for clfna^e ank for ornamental purposes. English .WW “S ^LHS ss r. *rr”-- 5W^«» “ “• tains only 800 parts per 1000. Silver mating -for purposes of electro-plating, a solution of silver cyanide Silver-plating, ro p v Whgn a feeble electric current is passed in potassium cyanide * silvered being the negative electrode, through this solution (the ar ic ^ ^ ^ coherent for m is precipitated -ic. ^ S2 sssssssssss: assss „V1.« gj"„ ssu s?“™ are added, and the M t Tbe silvered (after being perfectly id" Siion. in fbout twenty ntinu.es the silver will have formed a brilliant mirror upon the glass.* Oxides Of Silver.— Three oxides are believed to exist, namely— Silver monoxide Silver peroxide Silver suboxide Ag a O. Ag 2 0 2 ? Ag 4 Q ? Silver Monoxide ( argentic oxide), Ag 2 0 , is obtained by adding * Bv the reduction of silver solutions in the presence of certain organic as an allotropic form of silver {American Journal of Science, 1891). 5 20 Inorganic Chemistry sodium or potassium hydroxide to a solution of silver nitrate. A brown precipitate, consisting of hydrated oxide, is obtained, which when heated, is converted into the anhydrous compound. It is also formed when silver chloride is boiled with a strong solution of potassium hydroxide — 2AgCl + 2KHO = 2KC1 + H 2 0 + Ag 2 0. Silver oxide is a black amorphous powder, which when heated to 260°, begins to give off oxygen, and become reduced to metallic silver. It is a powerful oxidising substance, and when rubbed with sulphur, red phosphorus, sulphides of antimony or arsenic, or other readily oxidised substances, it causes them to ignite. Silver oxide, although only very slightly soluble in water (1 part in about 3000), imparts to the solution a distinct metallic taste and an alkaline reaction. It is' reduced by hydrogen at ioo°, with formation of water and metallic silver ; and when brought into contact with peroxide of hydrogen, oxygen is evolved and metallic silver formed (see p. 203). Silver oxide is soluble in strong ammonia, and, on standing, the solution deposits black shining crystals of the so-called fulminating silver. When dry, this compound is extremely explosive, and it often explodes when wet. Fulminating silver is believed to be the nitride, with the composition NAg 3 . Silver Peroxide, Ag 2 0 2 (?). — When a solution of silver nitrate is submitted to electrolysis, a black powder, consisting of small octahedral crystals, is deposited upon the positive electrode. The same compound is obtained when a plate of silver is made the positive electrode in the electrolysis of acidulated water, and also when silver is acted' upon by ozone. The exact composition of the compound has not been placed beyond doubt ; it is believed to be a dioxide. It readily parts with oxygen, and is a still more powerful oxidising agent than the monoxide. It dissolves in aqueous ammonia, with the evolution of nitrogen — 3Ag 2 O a + 2NH S = 3Ag 2 0 + 3H 2 0 + N 2 . Silver Suboxide, Ag 4 0 (?). — The black powder, obtained when silver citrate is reduced in a current of hydrogen at 100 °, and potassium hydroxide is added to the aqueous solution of the residue, is believed to have the composition Ag 4 0. Silver Chloride, AgCl, is obtained as a white, bulky, curdy precipitate when a soluble chloride is added to silver nitrate. It melts at 451 0 to a yellowish liquid, which on cooling, congeals to a tough homy mass (hence the name horn silver , as applied to the Silver Fluoride 5 21 native silver chloride). The precipitated chloride is soluble to a slight extent in strong hydrochloric acid, but readily soluble in alkaline chlorides, in ammonia, and in sodium thiosulphate. Potas- sium cyanide converts silver chloride into silver cyanide, which dissolves in the excess of alkaline cyanide, forming the double cyanide KCN,AgCN. When exposed to the light, silver chloride darkens in colour, assuming first a violet tint, and finally becoming dark brown or black (see Photo-salts, p. 522). Silver chloride absorbs large volumes of ammonia, forming the compound 2AgCl,3NH 3 (see p. 242). Silver Bromide, AgBr, is prepared similarly to the chloride, the precipitated compound having a pale yellow colour. It is less soluble in ammonia than silver chloride ; in dilute ammonia it is nearly insoluble. Silver bromide is decomposed by chlorine, and at a temperature of ioo° by hydrochloric acid. At ordinary temperatures this reaction is reversed, hydrobromic acid convert- ing silver chloride into the bromide. Dry silver bromide does not absorb gaseous ammonia. Silver bromide is extremely sensitive to the action of light, and is the chief silver compound used in dry-plate photography. Silver Iodide, Agl, may be obtained by precipitation from silver nitrate, with a soluble iodide ; or by dissolving silver in strong hydriodic acid. As obtained by precipitation it is an amorphous yellow substance, less soluble in ammonia than either the bromide or chloride. It dissolves in hot hydriodic acid, which on cooling deposits colourless crystals of Agl, HI ; the addition of water to the solution precipitates the normal iodide, Agl. Silver iodide absorbs gaseous ammonia, forming a white compound, 2AgI,NH 3 , which, on free exposure to the air, evolves ammonia, and is recon- verted into the yellow iodide. Silver iodide is the most stable of the three halogen compounds. When either the chloride or bromide is treated with hydriodic acid or potassium iodide, iodine replaces the other halogens, forming silver iodide. Silver Fluoride, AgF.— This compound is markedly different in many respects from the other halogen silver salts. It is obtained * by dissolving silver oxide or carbonate in hydrofluoric acid, and is deposited from the solution in colourless, quadratic pyramids, AgF,H 2 0, or in prisms, AgF,2H 2 0. The salt is extremely deli- quescent, and very soluble in water. When dried in vacuo, the salt AgF,H 2 0 undergoes partial decomposition, leaving a brownish 5 22 Inorganic Chemistry residue. When heated, it is partially decomposed, according to the equation — 2 AgF,H 2 0 = 2Ag + 2 HF + H a O + O. The dry salt absorbs gaseous ammonia in large quantities, more than 800 times its own volume being taken up by the powdered substance. Silver Nitrate, AgNO a , is obtained by dissolving silver in nitric acid. It forms large colourless rhombic tables, which melt at 218 , and resolidify to a white, fibrous, crystalline mass, known as lunar caushc. Below a red heat it gives off oxygen, and forms silver nitrite; and at higher temperatures it is decomposed into metallic silver, and oxides of nitrogen. 100 parts of water at o° dissolve 121.9 parts, and at roo°, mo parts of the crystallised salt ; the solution is neutral. In contact with organic matter, silver nitrate is blackened on exposure to light. Thus, when the skin is touched with a solution of this salt, a few seconds’ exposure to light causes a brown or black stain. Owing to this property, silver nitrate is employed for marking-inks. Silver nitrate absorbs gaseous ammonia, forming the compound AgNO|^8NH s , the ab- sorption being accompanied with considerable rise of temperature. The compound AgN 0 3 , 2 NH 3 is deposited as rhombic prisms when aqueous silver nitrate is saturated with ammonia. Silver Sulphate, Ag 2 S 0 4 , is formed when silver, silver carbo- nate, or silver oxide is dissolved in sulphuric acid. It crystallises in rhombic prisms, isomorphous with sodium sulphate. With aluminium sulphate it forms an alum, in which the monovalent element silver, takes the place of potassium in common alum, Ag 2 S 0 4 ,Al 2 (S 0 4 ) 3 , 24 H 2 0 . Photo-salts. — This name has been applied by Carey Lea, to the coloured compounds formed by the action of light upon silver chloride, bromide, and iodide. The exact composition of the compounds that are formed when these silver salts are exposed to light is not definitely known. The change which they undergo has been attributed (1) to the partial reduction to metallic silver; (2) to the formation of sub-salts, such as Ag 2 Cl, Ag 2 Br, with elimination of chlorine or bromine ; (3) to the formation of oxychloride or oxybromide ; (4) to the production of double compounds of variable composition, of the sub-salt with the normal salt. Gold 523 GOLD. Symbol, Au. Atomic weight = 196.8. Occurrence.— This element occurs in nature almost exclusively in the un combined condition, chiefly in quartz veins and in alluvial deposits formed by the disintegration of auriferous rocks. It is present in small quantities in many specimens of iron pyrites, copper pyrites, and many lead ores, from which it is often profitably extracted. Gold is also met with in the form of an amalgam with mercury, and in combination with the element tellurium in the minerals fietzite (AgAu) 2 Te, and sylvanite (AgAu)Te 2 . Extraction.— Gold is extracted from auriferous quartz by caus- ing the finely-crushed substance to flow, by means of a stream of water, over amalgamated copper plates. The gold particles adhere to the mercury, with which they amalgamate, and the amalgam so obtained is carefully removed and distilled. From alluvial deposits, the native gold is separated by me- chanical washing. Gold is extracted from auriferous pyrites by means of chlorine. The ore is first carefully roasted, and, after being wetted, is exposed to the action of chlorine gas. The gold is thereby converted into the soluble auric chloride, AuC 1 3 , which is extracted by lixiviation, and precipitated by the addition of ferrous sulphate— 2AuC 1 3 + 6FeS0 4 = 2Au + Fe 2 Cl 6 + 2Fe 2 (S0 4 )3. Native gold usually contains silver, from which it may be sepa- rated by passing chlorine over the molten metal, in crucibles glazed with borax. The fused chloride of silver rises to the surface, and is prevented from volatilising by a covering of melted borax. When the operation is complete, the crucible is allowed to cool, when the . gold solidifies, and the still liquid silver chloride is poured off. The Cyanide Process. — Increasing quantities of gold are at the present time extracted by solution in potassium cyanide. The method is specially advantageous in cases where the gold is present in the ore in a very finely divided condition, and it also possesses the advantage over the “ chlorination process,” that the preliminary operation of roasting is obviated. The crushed ore is treated with 5 2 4 T norganic Chemistry a dilute solution of potassium cyanide (containing from 0.25 to 1 per cent, of potassium cyanide), with free exposure to the atmos- phere, since it has been shown that atmospheric oxygen takes a necessary part in the action. The gold is dissolved in the form of a double cyanide, according to the equation — 4Au + 8KCy + 2H 2 0 + 0 2 = 4KHO + 4KAuCy 2 . From this solution the gold is precipitated either by means of metallic zinc (usually in the form of fine turnings) or by electro- lytic deposition. The precipitation by means of zinc takes place according to the equation — 2KAuCy 2 + Zn = K 2 ZnCy 4 + 2Au. The deposit, after being freed as far as possible from zinc, is melted down with a suitable flux, and yields an alloy containing 70 to 80 per cent, of gold. When the gold is precipitated electrolytically, the anodes em- ployed are of. lead foil. These are finally melted down and cupelled, yielding gold of a high degree of purity. Properties. Gold is a soft yellow metal, which, when seen by light many times reflected from its surface, appears red. It is not acted upon by air or oxygen at any temperature, and does not decompose steam. No single acid is capable of attacking it (except selenic acid) ; but it is dissolved by aqua-regia, with for- mation of auric chloride. Gold is the most malleable and ductile of all the metals, and when beaten into very thin leaf, it appears green by transmitted light. Gold is most easily reduced from its combinations. Most metals, when placed in a solution of a gold salt, precipitate the gold, and the most feeble reducing agents bring about the same result. On this account a solution of auric chloride is used for toning photographs. All the compounds of gold, when ignited in the air, are reduced to metallic gold. Gold is readily deposited upon other metals, by the process of electro-gilding, the most suitable solution being that of the double cyanide of gold and potassium, Au(CN) 3 ,KCN. Gold Alloys. — Alloys of gold with copper and with silver are used for coinage, and for ornamental purposes, pure gold being too soft for these purposes. Silver gives the alloy a paler colour than that of pure gold, while copper imparts to it a reddish tinge. The alloy used for English gold coin consists of gold, n 525 Gold Sulphides parts ; copper, i part. The proportion of gold in alloys is usually expressed in parts per 24 (instead of in percentages), these parts being termed carats . Thus pure gold, is said to be 24-carat gold ; \ 8-carat gold contains, therefore, 18 parts of gold, and 6 parts of copper, or silver. Most countries have their own legal standards. In England the legal standard for gold coinage is 22-carats. Compounds of Gold.— Gold forms two series of compounds, namely, aurous, in which the metal is monovalent, and auric , in which it is trivalent. The composition of aurous compounds corresponds to that of the silver compounds. They are very readily decomposed. Thus, aurous chloride cannot exist in the presence of water, being decomposed into auric chloride and metallic gold. For this reason, when aurous oxide, Au 2 0, is acted upon by aqueous hydrochloric acid, it forms auric, and not aurous chloride, thus— 3Au 2 0 + 6HC1 = 2AuCl 3 + 3H 2 0 + 4Au. With iodine, gold forms only aurous iodide, Aul ; therefore, when auric oxide is acted upon by hydriodic acid, aurous iodide and free iodine are formed, thus — _ Au 2 0 3 + 6HI = 2AuI + 2I 2 + 3H 2 0. Auric Chloride, AuC 1 3 , is obtained by dissolving gold in aqua-regia, and evaporating the solution to dryness. When the residue is dissolved in water, the concentrated solution deposits reddish crystals of the composition AuC 1 3 2H 2 0. These lose their water when carefully heated, leaving a brown mass of deliques- cent crystals. Auric chloride forms double chlorides with the alkaline chlorides, and with hydrochloric acid, which may be obtained as crystalline compounds. Thus the compound AuC1 s ,HC 1,3H 2 0 is deposited from a strong solution of gold in aqua-regia. This substance is sometimes termed and the double compounds with metallic chlorides, such as AuCl 3 ,NaCl,2H 2 U and (AuCl 3 ,KCl) 2 H 2 O f are known as chloro-aurates. Auric Oxide, Au 2 0 3 , is obtained as a brown powder, when the hydrated oxide, Au 2 0 3 ,3H 2 0 (or Au(HO) s ), is gently warmed. At 100 it begins to de- compose, and at higher temperatures is completely converted into oxygen and metallic gold. , , , ... 1 . Auric oxide is feebly basic, forming a few unstable salts, m which go d replaces the hydrogen in acids. It is also a feeble acid-forming oxide, and forms salts called aurates , such as potassium aurate, KAu0 2 ,3H 2 0, which may be regarded as being derived from an acid of the composition HAu0 2 . Auric oxide forms a compound with ammonia, known as fulminating gold, the exact composition of which is not known. It explodes with violence when dry, if struck, or gently warmed. , . , Gold Sulphides.— Two sulphides of gold have been obtained, aurous sulphide, Au 2 S, and auro-auric sulphide, Au 2 S,Au 2 S 3 (or AuS). The latter is formed when sulphuretted hydrogen is passed into a cold solution of auric chloride — SAuC 1 8 + 9H a S + 4H a O = 2(Au 3 S,Au a S 3 ) + 24HC1 + H a SO*. CHAPTER VI ELEMENTS OF GROUP II. ( FAMILY A.) Beryllium , Be Magnesium, Mg . Calcium, Ca Atomic Weights. 9.08 Strontium, Sr . 23.94 Barium, Ba Weights. 87.3 . 136.86 Atomic • 39-91 With the exception of the rare element beryllium, these metals were first obtained (although not in the pure state) by Davy, who, soon after his discovery of the metals potassium and sodium, showed that the so-called earths were not elementary bodies as had been supposed, but were compounds of different metals with oxygen. The element beryllium is of later discovery, for although as early as 1798 it had been shown by Vanquelin that the particular “ earth ” in the mineral beryl was different from any other known earth, it was not until 1827 that the metal it contained was iso- lated by Wohler. In a state approaching to purity, beryllium was first prepared by Humpidge, 1885. None of the elements of this family occurs in nature in the un- combined condition; and, with the exception of magnesium, the metals themselves, in their isolated condition, are at present little more than chemical curiosities. In the case of beryllium this is due to the comparative rarity of its compounds ; but with calcium, strontium, and barium, whose compounds are extremely abundant, it is owing partly to the difficulty of isolating the metals in a pure state, and also to the fact that hitherto they have received no useful application. Beryllium and magnesium are white metals, which retain their lustre in the air. Calcium, strontium, and barium have a yellow colour, and on exposure to air become converted into oxide. All these metals form an oxide of the type RO. Beryllium oxide is insoluble in water ; magnesium oxide is very slightly soluble (1 part in 55,000 or 100,000 parts of water), but the solution f Metals of the Alkaline Earths 527 shows a feeble alkaline reaction. The calcium, strontium and barium oxides show increasing solubility, and stronger alkalinity and causticity. On this account these elements are known as the metals of the alkaline earths. These three elements also form peroxides of the type R0 2 - All the monoxides are basic, and combine with acids to form salts of the types RC1 2 , RS0 4 , R(NO s ) 2 . The element beryllium (the typical element) stands apart from the others of this family in many of its chemical relations Thus, the oxide BeO, unlike the corresponding compounds of the other elements, does not combine with water to form the hydroxide. The hydroxide Ee(HO) 2 is soluble in sodium and potassium hydroxide. In this respect beryllium exhibits its resemblance to zinc. The chloride also differs from the other chlorides in being ^In'iS permanence in air, its colour, its high melting-point, the solubility of its sulphate, and the readiness with which its hydroxide is converted by heat into the oxide, beryllium exhibits a close similarity to magnesium. In the solubility of its hydroxi e in potassium hydroxide, and in its inability to decompose water, beryllium also shows a marked resemblance to zinc. The three elements, calcium, strontium, and barium, exhibit a closer resemblance to each other in most of the physical and chemical relations, than to either magnesium or beryllium. They are readily distinguished by their different spectra. Barium salts, when heated in a non-luminous flame, impart to it a green colour. Calcium and strontium, under the same cir- cumstances, each give a red colour; but the red imparted by strontium compounds is more brilliant, and less orange, than that of calcium salts. When the flames are examined by the spectro- scope the most characteristic lines given by barium, are two in the bright green (Baa and Ba/3). These are accompanied by a number of less brilliant lines. The spectrum of strontium consists of four specially prominent lines, one in the bright blue (Sr8), one in t e orange (Sra), and two in the red (Sr/3 and Sry), with others less pronounced; while that of calcium contains one brilliant green line (Ca/3), and one equally brilliant orange line (Caa), besides a large number of less prominent lines. 528 Inorganic Chemistry BERYLLIUM Symbol, Be. Atomic weight — 9.08. Occurrence.— This element occurs principally in the mineral beryl, a double silicate of the composition 3 BeO,Al 2 0 3 , 6 Si 0 2 . The transparent varieties are used as gems, the transparent green beryl being the precious emerald. Phenacite is beryllium silicate, BejSiC^, while chrysoberyl has the compo- sition BeO,ALj 0 3 . Formation.— The clement is obtained by heating sodium in the vapour of beryllium chloride, all air having been previously replaced by hydrogen. The product is afterwards melted beneath fused sodium chloride, when it is obtained as a coherent of solid metal. Properties.— Beryllium is a white metal resembling magnesium. It has a specific gravity of 2.1, and is moderately malleable. It does not readily tarnish in the air at ordinary temperatures, but, when strongly heated becomes coated with a protecting film of oxide. The powdered metal, when heated, lakes fire, and burns with a bright light. It has no action upon water, even at the boiling temperature. Beryllium is. easily dissolved by dilute hydrochloric acid, with evolution of hydrogen. Cold dilute sulphuric acid is without action, but, when heated, slowly dissolves it. Nitric acid slowly attacks it when concentrated and hot. It readily dissolves in potassium hydroxide, with evolution of hydrogen. Beryllium Compounds.— The best known are the oxide (berylla), BeO, a white infusible powder, insoluble in water, soluble in acids; the chloride, BeCl 2 , obtained by heating the oxide with charcoal in a stream of chlorine a white crystalline solid, readily fused and volatilised. Beryllium compounds do not impart any colour to a Bunsen flame. The> are characterised by possessing a sweet taste, hence the name of glucinum originally given to this element. MAGNESIUM. Symbol, Mg. Atomic weight = 23.94. Occurrence. Magnesium is not found in the uncombined state. In combination it is widely distributed, and is extremely abundant. In the mineral dolomite , associated with lime as carbonate, it occurs in mountainous masses. Magnesite, MgCO a ; kieserite, MgS0 4 ,H 2 0 ; carnallite, MgCl 2 , KC1,6H 2 0, are amongst the commoner naturally occurring magne- sium compounds. It is also a constituent of asbestos , meerschaum , serpentine , talc , and a large number of other silicates. As sulphate and chloride it is met with in sea water and many saline springs. Modes of Formation. — Magnesium can be obtained by the Magnesium Oxide 5 29 electrolysis of the fused chloride, or a mixture of magnesium and potassium chlorides. On a larger scale magnesium is prepared by reducing the chloride with sodium. A mixture of anhydrous magnesium chloride (or fused mixed chlorides of magnesium and sodium, or potassium), powdered cryolite, and sodium is thrown into a red-hot crucible, which is quickly closed. A violent reaction takes place, at the conclusion of which the melted mixture is stirred with an iron rod to cause the globules of magnesium to run together. The crude metal is afterwards purified by distillation. Properties. — Magnesium is a silvery -white metal, which does not tarnish in dry air, but becomes coated with a film of oxide when exposed to air and moisture. At a red heat it melts, and at higher temperatures may be distilled. When heated in the air it takes fire, and burns with a dazzling white light, which is extremely rich in the chemically active rays. The flash of light, obtained by projecting a small quantity of magnesium filings into a spirit flame, is used for photographic purposes. Magnesium is only moderately malleable, and is only ductile at high temperatures ; it is readily pressed into the form of wire at a temperature slightly below its melting-point. Magnesium does not decompose water even at the boiling-point ; but when strongly heated in a current of steam, the metal takes fire (p. 152). Magnesium is rapidly dissolved by dilute acids, with brisk evolution of hydrogen, but solutions of caustic alkalies are unacted upon by it (compare Zinc). When heated with aqueous solutions of ammonium salts, hydrogen is evolved, and a double salt of magnesium and ammonium is found in the solution. Magnesium combines directly with nitrogen, when strongly heated in that gas, forming magnesium nitride, N 2 Mg 3 (p. 208). On account of the brilliant light emitted by burning magnesium, it is employed for signalling purposes, and also in pyrotechny. Magnesium Oxide (magnesia), MgO, is found native as the mineral 'periclase. It is formed when magnesium burns in the air, or when magnesium carbonate is submitted to prolonged gentle calcination, when it is obtained as a white, bulky powder, known in commerce as calcined magnesia or magnesia usta. Magnesia is extensively manufactured from the magnesium chloride occurring in the Stassfurt deposits, by first converting the chloride into carbonate, and subjecting this to calcination. Mag- nesia has been obtained in the crystalline form, identical with that of periclase, by heating the amorphous compound in a stream of 5 3° Inorganic Chemistry gaseous hydrochloric acid. It may be fused in the oxyhydrogen flame, and on cooling, it solidifies to a vitreous mass which is suffi- ciently hard to cut glass. On account of its extreme refractoriness, magnesia is used for a variety of metallurgical purposes, such as the manufacture of crucibles, cupels, &c. Magnesium Hydroxide, Mg(HO) 2 , is found in nature as the mineral brucite. It is prepared by precipitating a magnesium salt by sodium or potassium hydroxide. At a dull red heat it loses water, and is converted into the oxide, and the magnesia so obtained has the property of rehydrating itself in contact with water, with evolution of heat. Magnesium hydroxide slowly absorbs carbon dioxide, forming the carbonate ; owing to this fact, and to the property it possesses of rehydration, magnesia that has been prepared by calcination at a low temperature can be employed as a cement. Thus, if calcined magnesite be made into a paste with water, the mixture is found to harden in about twelve hours, and ultimately to acquire a hardness equal to that of Portland cement. Magnesium Chloride, MgCl 2 — This salt is formed when mag- nesia, or magnesium carbonate, or the metal itself, is dissolved in hydrochloric acid. From this solution monosymmetric crystals of the composition MgCl 2 ,6H 2 0 are deposited. When this salt is heated it loses water, and at the same time is partially decomposed into hydrochloric acid and magnesia ; in order, therefore, to pre- pare the pure anhydrous compound, the double magnesium ammo- nium chloride is first formed, by adding ammonium chloride to a solution of magnesium chloride. On evaporation, the double salt separates out, MgCl 2 ,NH 4 Cl,6H 2 0. This salt allows itself to be dehydrated by heating, without any decomposition of the magne- sium chloride. When the dried salt is more strongly heated, ammonium chloride volatilises and leaves the anhydrous magnesium chloride as a fused mass, which congeals to a white crystalline solid. Magnesium chloride is deliquescent, and dissolves in water with evolution of heat. With alkaline chlorides it forms double salts, as the ammonium salt above mentioned. The potassium salt, MgCl 2 ,KCl,6H 2 0, occurs in large quantities as the mineral camallite ; and the calcium salt, 2MgCl 2 CaCl 2 ,12H 2 0, as tachydrite , in the Stassfurt deposits. When a strong solution of magnesium chloride is made into a thick paste with calcined magnesia, the mass quickly sets and hardens, like plaster of Paris, and is found to contain an oxychloride having the composition MgCl 2l 5MgO, Magnesium Sulphate 5 3 1 associated with varying quantities of water. The white deposit which forms in bottles containing the solution known as magnesia mixture, consists of MgCl 2 ,5MgO,13H 2 0. When magnesium oxychloride is heated to redness in a current of air, the magnesium is converted into oxide, and a mixture of chlorine and hydrochloric acid is evolved. The reaction may be represented as taking place as follows — 2MgCl 2 + H 2 0 4-0 = 2 MgO 4- 2HC1 4- Cl 2 . The Weldon- Pechiney process for manufacturing chlorine is based upon this reaction. Magnesium Sulphate, MgS0 4 {Epsom salts), is met with in many mineral springs, and in large quantities as the mineral kieserite, MgS0 4 ,H 2 0. Magnesium sulphate may be obtained by decomposing dolomite, (CaMg)C0 3 , with sulphuric acid, the nearly insoluble calcium sulphate being readily removed from the soluble magnesium salt. Magnesium sulphate is now very largely manufactured from kieserite , which in contact with water is converted from the slightly soluble monohydrated salt, into MgS0 4 ,7H 2 0, which is readily soluble, and is purified by recrystallisation. As, usually obtained, crystallised magnesium sulphate, MgS0 4 ,7H 2 0, forms colourless rhombic prisms ; but when deposited from a cold supersaturated solution, it sometimes forms prisms belonging to the monosymmetric (monoclinic) system, having the same degree of hydration. Above 50 °, monosymmetric prisms of the composition MgS0 4 ,6H 2 0 are deposited. When the ordinary salt, MgS0 4 ,7H 2 0, is placed over sulphuric acid, it loses two molecules of water : when heated to 1 50 ° it loses six molecules, and at 200 ° it becomes anhydrous. At the ordinary temperature, 100 parts of water dissolve 126 parts of crystallised magnesium sulphate ; the solution has a bitter taste, and acts as a purgative. With alkaline sulphates, magnesium sulphate forms a series of double salts, having the general formula MgS0 4 ,R 2 S0 4 , 6H 2 0. They are isomorphous with each other, crystallising in monosymmetric prisms. The potassium salt occurs in the Stassfurt deposits as schonite. When anhydrous magnesium sulphate is dissolved in hot sul- phuric acid, two acid sulphates are obtained. One, having the com- position MgS0 4 ,H 2 S0 4 , is deposited from the hot solution ; while 532 Inorganic Chemistry from the cold liquid the salt that crystallises has the composition MgS0 4 ,3H 2 S0 4 . They are at once decomposed by water. Magnesium Carbonate, MgC0 3 , occurs as the mineral magne- site , which is sometimes found as rhombohedral crystals, isomor- phous with crystals of calcite (CaC0 3 ). Magnesium exhibits a great tendency to form basic and hydrated carbonates ; the normal carbonate, MgC0 3 , is therefore not obtained by precipitating a magnesium salt with an alkaline carbonate ; the white precipitate formed under these circumstances is a basic carbonate, whose composition varies with the conditions of precipitation. If, how- ever, this precipitate be suspended in water, and the liquid saturated with carbon dioxide, the compound dissolves (more readily under increased pressure), and when the solution is heated to 300 ° under pressure, in such a manner that the evolved carbon dioxide can escape, the normal anhydrous carbonate is deposited in rhombo- hedral crystals isomorphous with calcite. If the solution be evaporated to dryness, the normal carbonate is deposited in rhombic crystals isomorphous with arragonite (CaC0 3 ). Mag- nesium and calcium carbonates are therefore isodimorphous. Basie Carbonates. — The mineral hydromagnesite is a basic carbonate of the composition 3MgCQ 3 ,Mg(H0) 2 ,3H 2 0. A number of basic carbonates are formed by the precipitation of a magnesium salt with sodium carbonate. Thus, under ordinary conditions a white bulky precipitate is obtained, known in pharmacy as magnesia alba levis. Its composition, although liable to vary through the presence of other basic carbonates, is in the main the same as that of hydromagnesite. If the precipitation be made with boiling solutions, and the pre- cipitate so obtained be dried at ioo°, a denser carbonate is ob- tained, termed magnesia alba ponder osa, 4MgC0 3 ,Mg(H0) 2 ,4H 2 0. When an excess of sodium carbonate is employed, and the mixture is subjected to prolonged boiling, a carbonate is obtained, having the composition 2MgC0 3 ,Mg(H0) 2 ,2H 2 0. CALCIUM. Symbol, Ca. Atomic weight = 39.91. Occurrence. — Calcium is only met with in nature in combina- tion. It occurs in enormous quantities as the carbonate, in a great variety of different minerals, such as marble , limestone , ealesfiar, and also as coral j and with carbonate of magnesia as dolomite, , or Calcium Oxide 533 magnesian limestone. In combination with sulphuric acid, calcium occurs as gypsum and selenite, CaS0 4 ,2H 2 0, and as anhydrite, CaS0 4 . The fluoride CaF 2 occurs as fluorspar j and the various Silicious rocks contain compound silicates of calcium and other metals. The carbonate and sulphate are present in most spring and river waters. Calcium compounds are also present in all vegetable and animal organisms. Thus, bones consist largely of calcium phosphate. Modes of Formation. — Although calcium compounds are so extremely abundant, the element itself is a rare substance. The element was first isolated in an impure state by Davy (1808). It may be obtained by the electrolysis of the fused chloride, or by fusing together calcium chloride, sodium, and zinc, when an alloy of zinc and calcium is obtained, from which the zinc is removed by distillation. Properties. — Calcium is a yellow metal, somewhat the colour of pale brass. It is sufficiently hard to be worked with a file, and may be hammered out into leaf. In moist air it is soon converted into the hydroxide, but in dry air it remains untarnished for a con- siderable time. It decomposes water at the ordinary temperature, with rapid evolution of hydrogen. When heated in the air, it takes fire and burns. Oxides of Calcium. — Two oxides are known, namely, calcium monoxide, CaO, and calcium dioxide, Ca0 2 . Calcium Oxide {lime, quicklime), CaO, is obtained by heating calcium carbonate to redness — CaC0 3 = C0 2 + CaO. On a large scale, lime is manufactured by burning limestone or chalk in kilns with coal. If much clay be present with the lime- stone, care is required to prevent the mass from fusing, when it is said to be dead burnt. Lime is a white amorphous substance, which is infusible by the oxyhydrogen flame ; but which, when so heated, emits a bright light, known as the oxyhydrogen lime-light. It absorbs moisture and carbon dioxide from the air. On account of its power of absorbing moisture, lime is frequently employed as a dehydrating agent. Thus, gases which cannot be dried by means of sulphuric acid {e.g., ammonia) may be deprived of moisture by being passed over calcium oxide. It is also used for withdrawing water from alcohol, in the preparation of absolute alcohol. When 534 Inorganic Chemistry a small quantity of water is poured upon lime, the mass rapidly becomes hot, and volumes of steam are given off ; the lime at the same time swelling up and crumbling to a soft, dry powder. This process is known as the slaking of lime, and the product is termed slaked lime , in contradistinction to quick lime. The lime enters into chemical union with water, forming calcium hydroxide, thus — CaO + H 2 0 = Ca(HO) 2 . Calcium Hydroxide, Ca(HO) 2 , is a white amorphous powder, sparingly soluble in water ; and, unlike the majority of solids, it is less soluble in hot than in cold water. One hundred parts of water at the ordinary temperature dissolve o. 14 parts of calcium hydroxide, while at ioo° the same volume of water dissolves about half that amount. This solution, known as lime water , has an alkaline reaction, and absorbs carbon dioxide, with the precipitation of calcium carbonate. Milk of Lime is the name given to a mixture of lime with less water than will dissolve it, whereby an emulsion of lime is obtained. When a thick paste of lime and water is exposed to the atmos- phere, in a few days it sets, and continues gradually to harden. On this account lime is used for mortars and cements. Mortar consists of a mixture of lime and sand with water. The sand serves the double purpose of preventing shrinkage on drying, and also of rendering the mass more permeable to atmospheric carbon dioxide. The setting of mortar is due to the combined action of evaporation and absorption of carbon dioxide. Calcium Dioxide, Ca 0 2 , is obtained by adding lime-water to hydrogen peroxide, or to sodium peroxide acidulated with dilute nitric acid : sparingly soluble crystals of Ca 0 2 , 8 H 2 0 separate out, which at 130° lose their water. When more strongly heated, the monoxide is formed, with evolution of oxygen. Calcium Chloride, CaCl^ occurs in sea and river waters, and is present in carnallile and tacky drite, of the Stassfurt deposits. It is obtained in large quantities as a bye-product in many manu- facturing processes, such as that of potassium chlorate, ammonia from ammonium chloride, &c. It may be obtained by the action of hydrochloric acid upon calcium carbonate, and is deposited on concentration, in large colourless, deliquescent, hexagonal prisms, CaCl 2 , 6 H 2 0 , which melt at 29 0 in their water of crystallisation. When heated below 200°, the crystals part with four molecules of Bleaching- Powder 535 water, and above 200° become anhydrous. As thus obtained, the anhydrous salt is a porous mass, which is extremely hygroscopic, and on this account is used as a desiccating agent, both for gases and liquids. At a red heat it fuses, and on cooling, solidifies to a crystalline, deliquescent mass. Calcium chloride combines with ammonia, forming the compound CaCl 2 ,8NH 3 . Calcium chloride, therefore, cannot be employed for drying gaseous ammonia. Crystallised calcium chloride is extremely soluble in water ; 100 parts of water at 16 0 dissolve 4°° parts of the salt, the solu- tion being attended with considerable absorption of heat. When mixed with powdered ice, or snow, liquefaction of both the solids rapidly takes place (owing to the formation of a cryohydrate), and the consequent absorption of heat, lowers the temperature of the mixture to — 40°. Bleaehing-Powder {chloride of lime), Ca(OCl)Cl. — This im- portant compound is manufactured on a large scale by the action of chlorine upon slaked lime. The hydrated lime is spread upon the floor of the bleaching-powder chambers to a depth of three or four inches, and raked into ridges or furrows with a special wooden rake. Chlorine is then led into the chambers, which are provided with glass windows to enable the operator to examine the colour of the atmosphere within. At first the absorption of the chlorine is rapid, but as the reaction proceeds it becomes slower, and the lime is from time to time raked over to expose a fresh surface. The lime is left in contact with the gas for twelve to twenty-four hours. The excess of chlorine is absorbed by projecting into the chamber a shower of fine lime-dust, by means of a mechanical fan-distributor. This, in settling, rapidly absorbs all the chlorine, and the chambers can then be opened without any unpleasant smell of chlorine being perceptible. The reaction which takes place is expressed by the equation — Ca(HO) 2 + Cl 2 = Ca(OCl)Cl + H a O. It was formerly believed that bleaching-powder was a mechani- cal mixture of calcium chloride, CaCl 2 , and calcium hypochlorite, Ca(OCl) 2 , but it has been conclusively shown that the substance does not contain any free calcium chloride. It may, however, be regarded as a compound consisting of equivalent proportions of these two salts, and its composition may be expressed by the for- mula Ca(OCl) 2 ,CaCl 2 , which corresponds to 2Ca(OCl)Cl. 53 ^ Inorganic Chemistry The relation in which bleaching-powder stands to calcium chlo- ride on the one hand, and calcium hypochlorite on the other, will be seen by the following formulae — Calcium Chloride. Cl— Ca— Cl Calcium Hypochlorite. ~ Bleaching- Powder or Calcium Chloro-hypochlorite. CIO— Ca— OC1 Cl— Ca— OC1. In practice, the absorption of chlorine by the lime is never as complete as is represented by the above equation, and the com- mercial value of the product depends upon the amount of available chlorine it contains, i.e., chlorine which is evolved on treating the compound with hydrochloric or sulphuric acid. This ranges from 30 to 38 per cent. When treated with water, bleaching-powder is converted into calcium chloride and hypochlorite, thus — 2Ca(OCl)Cl = CaCl 2 + Ca(OCl) 2 . Bleaching-powder decomposes slowly even in stoppered bottles, and more rapidly on exposure to atmospheric moisture and carbon dioxide. When acted upon by acids, chlorine is evolved, thus — Ca(OCl)Cl + 2HC1 = CaCl 2 + H 2 0 + Cl 2 Ca(OCl)Cl + H 2 S0 4 = CaS0 4 + H 2 0 + Cl 2 . When a solution of bleaching-powder is treated with very dilute acids, hypochlorous acid is first liberated, which in contact with hydrochloric acid yields chlorine — (1) Ca(OCl) 2 .Aq + 2HCl.Aq = CaCl 2 + 2HC10.Aq. (2) HCIO + HC1 = H 2 0 + Cl 2 . In the process of bleaching, the material is first steeped in a dilute solution of bleaching-powder, and then in dilute acid. The hypochlorous acid first formed is decomposed in the presence of excess of hydrochloric acid, generating chlorine within the fibres of the wet cloth. Calcium Sulphate, CaS0 4 , occurs as the mineral anhydrite , and in the hydrated condition as gypsum , CaS0 4 ,2H 2 0, of which satinspar (or fibrous gypsum ), alabaster , and selenite are different varieties. It is obtained in the hydrated condition by precipita- Calcium Carbonate 537 tion from a solution of calcium chloride, on the addition of sul- phuric acid or a soluble sulphate. When dried at no to 120 it loses a portion of its water, leaving the hydrate, (CaS 0 4 ) 2 ,H 2 0 ; at 200° it becomes anhydrous. Calcium sulphate, in the hydrated condition, is slightly soluble in water, the solubility reaching a maximum at 35°* when 1 part of the compound requires 43 ^ parts of water for its solution ; above this temperature the solubility again diminishes. Its solubility is increased by the presence of alkaline chlorides and free hydrochloric acid. When boiled in strong sulphuric acid, calcium sulphate partially dissolves, and on cooling, an acid sulphate crystallises out, having the composition CaS 0 4 ,H 2 S 0 4 . Plaster of Paris is calcium sulphate which has been partially deprived of its water of hydration by heat, and converted into the hydrate, (CaS 0 4 ) 2 H 2 0 . It is manufactured by burning gypsum in a kiln, or oven, in such a way that the carbonaceous fuel does not come in contact with the sulphate, which would result in its reduction to sulphide ; the temperature is not allowed to exceed about 130°. If heated more strongly (above 200°) the sulphate becomes anhydrous, and is said to be dead burnt j in this con- dition its property of setting when mixed with water is greatly impaired. When plaster of Paris is made into a paste with water, it rapidly sets to a hard mass : this setting is due to its rehydra- tion, whereby gypsum is reformed, thus — (CaS 0 4 ) 2 ,H 2 0 + 3 H 2 0 = 2 CaS 0 4 , 2 H 2 0 . Calcium Carbonate : CaC 0 3 .— This compound is extensively met with in nature, as limestone , chalk , marble , and innumerable varieties of calcspar . It is formed when lime is exposed to atmos- pheric carbon dioxide. It is obtained when an alkaline carbonate is added to a soluble calcium salt. Calcium carbonate is dimorphous ; it occurs as arragonite in crystals belonging to the rhombic system, and as calcspar in crystals belonging to the hexagonal system. Both these crystal- line varieties can be artificially obtained : when deposited from solutions at the ordinary temperature, the crystals are identical with calcite ; but when crystallised from hot solutions, they form rhombic crystals corresponding to arragonite. Calcium carbonate is nearly insoluble in water ; 1000 grammes of water dissolve .0018 grammes of the compound. It is more 53^ Inorganic Chemistry soluble in water charged with carbon dioxide, forming the acid carbonate of lime, CaC0 3 ,H 2 C0 3 , or H 2 Ca(C0 3 ) 2 . 1000 grammes of water saturated with carbon dioxide will dis- solve, at o°, 0.7 grammes of calcium carbonate. By increasing the pressure (thereby increasing the amount of dissolved gas) as much as 3 grammes of calcium carbonate may be dissolved. When this solution is boiled, the acid carbonate is decomposed (p. 197). Calcium Phosphate ( tricalcium orthophosphate ), Ca^PO^, is the most important of the phosphates of calcium. It is found as the mineral osteolite , Ca 3 (P0 4 ) 2 ,2H 2 0, and also as sombrerite , estramadurite , and coprolites. Apatite consists of phosphate and fluoride, 3Ca 3 (P0 4 ) 2 ,CaF 2 ; and the mineral constituents of bones consist chiefly of calcium phosphate. It is obtained in a pure state by the addition of ordinary sodium phosphate to a solution of calcium chloride, in the presence of ammonia. The precipitate is decomposed on boiling, into an insoluble basic salt, and a soluble acid salt. Although nearly insoluble in pure water, calcium phosphate dissolves in water con- taining salts in solution, such as sodium chloride or nitrate, or even dissolved carbon dioxide. On this fact depends the readi- ness with which this substance is absorbed by the roots of plants. Calcium phosphate is readily soluble in both nitric and hydro- chloric acids. It is decomposed by sulphuric acid, with the forma- tion of monocalcium orthophosphate and calcium sulphate, thus — Ca 3 (P0 4 ) 2 + 2H 2 S0 4 = 2CaS0 4 + H 4 Ca(P0 4 ) 2 . This mixture of calcium sulphate and monocalcium phosphate is known as superphosphate of lime, and is largely used as an arti- ficial manure. With a larger quantity of sulphuric acid, the phosphate is con- verted into tribasic phosphoric acid. (See Phosphorus, p. 414.) Calcium Sulphide, CaS, is formed when sulphuretted hydrogen is passed over heated lime — Ca(HO) 2 + H. 2 S = CaS + 2H 2 0 Or by heating calcium sulphate with carbon — CaSO, + 4C - CaS + 4CO. Strontium 5 39 Calcium sulphide is decomposed on boiling with water, forming calcium hydroxide and hydrosulphide, thus 2CaS + 2H a O = Ca(HO) 2 + Ca(HS) 2 . Calcium sulphide (in common with barium and strontium sul- phides), as usually obtained \ possesses the property of emitting a feeble light (or phosphorescence) in the dark, after being previously exposed to a bright light. The light emitted gradually diminishes in intensity, but on re-exposing the compound to the light, its luminosity is again restored. This property has been long known, and calcium sulphide was formerly termed Cantoris phosphorus. The material formerly known as Bononian (or Bologniari) phos- phorus is the corresponding barium compound. These various sulphides are now manufactured for the preparation of so-called luminous paint. The phosphorescence of these compounds appears to be due to the presence of small quantities of foreign substances ; thus, not only is the particular colour of the light emitted changed by the intentional introduction of minute traces of bismuth, cadmium, manganese, zinc, and many other metals, but it has been shown, in the case of calcium sulphide, that the perfectly pure substance does not exhibit phosphorescence. STRONTIUM. Formula, Sr. Atomic weight = 87. 3. Oeeurrenee. — The chief natural compounds of this element are strontianite , SrCO a , and celestine , SrS0 4 . Modes of Formation.— The metal was first obtained in small quantity by Davy, by the electrolysis of the hydroxide, or chloride, 'moistened with water. It is more advantageously obtained by electrolysing the fused chloride. An amalgam of mercury and strontium (from which the strontium may be separated by volatilising the mercury in a stream of hydrogen) has been obtained by heating a saturated solution of strontium chloride with sodium amalgam. Properties.— Strontium is a yellowish metal resembling calcium. It is readily oxidised by air, and decomposes water at ordinary temperatures ; when heated in the air it bums brilliantly. Oxides of Strontium.— Two oxides, corresponding to those of calcium, are known, namely, strontium monoxide, SrO, and dioxide, Sr0 2 . Strontium Monoxide (. strontia ), SrO, is obtained by heating the 540 Inorganic Chemistry nitrate or carbonate. It is prepared on a large scale by decompos- ing strontium carbonate by superheated steam ; carbon dioxide is evolved, and strontium hydroxide remains, which on ignition forms the monoxide. Strontia strongly resembles lime. When treated with water it slakes with evolution of heat, forming strontium hydroxide, Sr(HO) 2 . The hydroxide is more soluble in water than the lime compound, and the solution on cooling deposits quadratic crystals, Sr(H0) 2 ,8H 2 0. The solution is strongly alkaline. Strontium hydroxide combines with sugar, forming a saccharate of strontia, which is readily decomposed by carbon dioxide. On this account it is prepared on a large scale for use in the manu- facture of beet-sugar. One process by which it is obtained on a commercial scale, consists in first forming strontium sulphide, by reducing the natural sulphate with carbon, and treating the solution of the sulphide with sodium hydroxide, thus — SrS + NaHO + H 2 0 = Sr(HO) 2 4 - NaHS. Strontium Dioxide, Sr0 2 . — When hydrogen peroxide is added to a solution of strontium hydroxide, a hydrate of the peroxide separates out in the form of pearly crystals, Sr0 2 ,8H 2 0. On gently heating this compound, it is converted into the anhydrous peroxide. On heating to redness it evolves oxygen, and is converted into the monoxide. Strontium Chloride, SrCl 2 , is obtained from strontianite by the action of hydrochloric acid. The salt deposits from the solution in deliquescent hexagonal prisms, SrCl 2 ,6H 2 0, isomorphous with the corresponding calcium compound. Strontium Sulphate, SrS0 4 . — The native compound celestine occurs in amorphous fibrous masses, and also in rhombic crystals. The name of the mineral is derived from the fact that it usually has a light-blue colour. It is produced by precipitation from a strontium salt by sulphuric acid. It is only slightly soluble in cold water, and still less in hot. When boiled with solutions of alkaline carbonates, strontium sulphate is completely converted into stron- twm carbonate — SrS0 4 + Na 2 C0 3 = SrCO s -1- NagS0 4 . In this respect strontium sulphate differs from barium sulphate, which under these conditions remains unchanged. On treatment with strong sulphuric acid, strontium sulphate forms SrS0 4 ,H 2 S0 4 . Barium 54 1 which, like the corresponding calcium compound, is converted by water into sulphuric acid and the normal sulphate. Strontium Nitrate, Sr(N0 3 ) 2 , is obtained by dissolving the natural carbonate in dilute nitric acid. On concentration, the anhydrous salt separates out in regular octahedrons. From dilute solution, on cooling, it forms monosymmetric prisms, Sr(NOj) 2 , 4H 2 0, which effloresce on exposure to the air. When heated with carbon, or other readily combustible substances, the mixture in- flames, and burns with the red colour characteristic of strontium compounds ; strontium nitrate is therefore largely used in pyro- techny for the production of red fire. This property is most readily illustrated by mixing dry powdered strontium nitrate with ammonium picrate, and igniting the mixture, which burns with a brilliant red light. BARIUM. Symbol, Ba. Atomic weight = 136.86. Occurrence. — The most abundant natural compounds of barium are heavy spar, BaS0 4 , and withente , BaC0 3 . It occurs also, associated with calcium, in the mineral barytocalcite , BaC0 3 ,CaC0 3 . Modes of Formation. — The element barium is more difficult to isolate than either strontium or calcium, and it is doubtful whether pure barium has ever been obtained. Davy electrolysed various barium salts, made into a thick paste with water, using mercury as the negative electrode : in this way an amalgam of barium was formed, from which, on distilling away the mercury, a dark porous mass was obtained. Amalgams of barium and mercury have been prepared in other ways, but it has been shown that the product obtained after distilling the mercury from these, is not pure barium, but is a solid alloy or compound of barium with mercury. By the electrolysis of the fused chloride, Matthiessen obtained small globules of metal, which on exposure to the air rapidly oxidised. More recent experi- menters fail to obtain the metal by this process (Limb., compt. rend 112). Oxides of Barium. — Two oxides are known, namely, barium monoxide, BaO, and dioxide, Ba0 2 . Barium Monoxide (baryta), BaO, is usually prepared by heat- ing the nitrate. The mass fuses and evolves oxygen and oxides of nitrogen, leaving a greyish white friable residue of the oxide. It may also be obtained by heating the carbonate ; but as the tem- perature necessary to expel the carbon dioxide is very high, it is usual to mix the carbonate with lampblack, tar, or other sub- 54 2 Inorganic Chemistry stances which on heating will yield carbon, when the conversion takes place more readily, carbon monoxide being evolved, thus — BaC0 3 + C = BaO + 2CO. Small quantities may readily be obtained by heating barium iodate in a porcelain crucible, when the iodate is decomposed as follows — Ba(I0 3 ) 2 = BaO + I 2 + 50. Barium oxide is a strongly caustic and alkaline compound ; in contact with water it slakes with evolution of so much heat, that the mass may become visibly red hot if too much water be not added. When heated to a dull red heat in oxygen, or air, it takes up an additional atom of oxygen and forms the dioxide (see p. 162). Barium Hydroxide, Ba(HO) 2 , is obtained when the monoxide is slaked with water. It is manufactured by first heating the powdered native sulphate with coal, when a crude barium sulphide is formed. This is then heated in a stream of moist carbon dioxide, whereby it is converted into the carbonate, and super- heated steam is then passed over the heated carbonate — BaS + H 2 0 + C0 2 = BaC0 3 + H 2 S BaCOg + H 2 0 = Ba(HO) 2 + C0 2 . Barium hydroxide is soluble in water : the solution, known as baryta-water , absorbs carbon dioxide with the precipitation of barium carbonate. The aqueous solution deposits crystals of hydrated barium hydroxide, Ba(H0) 2 ,8H 2 0, in the form of colourless quadratic prisms, which on exposure to the air lose seven molecules of water. Barium hydroxide, when heated in a current of air, yields barium dioxide. Barium hydroxide was formerly employed in sugar-refining, but owing to its poisonous nature it has been superseded by strontium hydroxide (q.v.). Barium Dioxide ( barium peroxide ), Ba0 2 . — This oxide is obtained by heating the monoxide to a low red heat in a stream of oxygen, or of air which has been deprived of atmospheric carbon dioxide. The pure compound may be obtained by adding an excess of Barium Chloride 543 baryta-water to hydrogen peroxide, when hydrated barium per- oxide separates out in crystalline scales — Ba(HO) 2 + H 2 0 2 + 6H 2 0 = Ba0 2 ,8H 2 0. On drying in vacuo at 130° this compound loses water and is converted into the anhydrous peroxide. The commercial peroxide may be purified by treatment with dilute hydrochloric acid, whereby barium chloride and hydrogen peroxide are formed. After the removal of insoluble impurities by filtration, baryta-water is cautiously added, which causes the pre- cipitation of ferric oxide and silica. The liquid is then filtered, and to the clear liquid, consisting of a solution of barium chloride and hydrogen peroxide, an excess of strong baryta-water is added, when the hydrated barium peroxide is precipitated, as already explained. Barium peroxide is a grey powder, which on being heated to a bright red heat gives up oxygen and forms the monoxide (p. 162). Dilute acids decompose barium peroxide, with formation of hydrogen peroxide and a barium salt. Concentrated sulphuric acid forms barium sulphate and ozonised oxygen. When gently warmed in a stream of sulphur dioxide, the mass becomes incan- descent and forms barium sulphate — Ba0 2 + S0 2 = BaS0 4 . Barium Chloride, BaCl 2 , may be obtained by dissolving the natural carbonate in hydrochloric acid. It may be obtained from the natural sulphate, either by first converting it into the sulphide, and decomposing that with hydrochloric acid, or by roasting the mineral with powdered coal, limestone, and calcium chloride, when the following reactions take place — BaS0 4 + 4C = BaS + 4CO BaS + CaCl 2 = BaCl 2 + CaS. The barium chloride is dissolved in water, and an insoluble oxy- sulphide of calcium remains. Barium chloride forms colourless rhombic tables, BaCl 2 ,2H 2 0, which at 15.6° are soluble to the extent of 43.5 parts in 100 parts of water. The salt is nearly insoluble in hydrochl'oric acid, and may therefore be precipitated from an aqueous solution by the addition of this acid. 544 Inorganic Chemistry Barium chloride, in common with all the soluble salts of this element, is highly poisonous. Barium Sulphate, BaS 0 4 , is the most abundant naturally occurring barium compound. It is frequently met with as large rhombic crystals. The specific gravity of the mineral is 4.3 to 4.7 ; and on account of its high specific gravity, it received the name of barytes , or heavy spar. It is formed as a heavy white precipitate, when sulphuric acid, or a soluble sulphate, is added to a solution of a barium salt. It is insoluble in water, and only very slightly soluble in dilute acids. It is soluble in hot concentrated sulphuric acid, espexially when freshly precipitated ; and the solution deposits, on cooling, an acid sulphate, BaS 0 4 ,H 2 S 0 4 . On exposure to moisture the solution deposits crystals of BaS 0 4 ,H 2 S 0 4 , 2 H 2 0 . Both of these com- pounds, in contact with water, yield insoluble normal barium sulphate and sulphuric acid. Precipitated barium sulphate is latgely used as a pigment, known as permanent white. Barium Nitrate, Ba(N0 3 )2, is obtained by dissolving the native carbonate, or the sulphide, in dilute nitric acid. It is also formed by double decomposition, when hot saturated solutions of sodium nitrate and barium chloride are mixed. The salt crystallises in large colourless octahedra belonging to the regular system. 100 parts of water at the ordinary temperature dissolve 9 parts, and at ioo°, 32.2 parts of barium nitrate. When strongly heated, it is converted into barium oxide, with the evolution of nitrogen per- oxide, oxygen, and nitrogen. Barium nitrate is used in pyrotechny, in the preparation of mixtures for green fire. Barium Sulphide, BaS, is obtained by methods analogous to those for preparing calcium sulphide (page 538), which it closely resembles in its properties. CHAPTER VII ELEMENTS OF GROUP II. ( FAMILY B.) Zinc, Zn 65 Cadmium, Cd 111.7 Mercury, Hg 199.8 The three elements composing this family do not exhibit such a close resemblance to each other as exists between barium, strontium, and calcium ; for although zinc and cadmium are very closely related, mercury in many respects differs widely from these, and from all the other elements in the same group. Cadmium and zinc are almost invariably found associated together in nature, they are both fairly permanent in the air, and both readily take fire and burn, when strongly heated, forming the oxides. Both are acted upon by dilute hydrochloric and sulphuric acids, with evolution of hydrogen, and most of their salts are isomorphous. Mercury is peculiar in being liquid at ordinary temperatures. Zinc and cadmium melt at 420° and 320° respectively, while mercury melts at - 38°. 8 . It is quite unacted upon by oxygen at ordinary temperatures, and combines with extreme slowness when heated. Its oxide, also, is readily decomposed by heat into its elements. Dilute hydrochloric and sulphuric acids are entirely without action upon it, and it forms no hydroxide. The hydroxide of zinc, Zn(HO) 2 , differs from the corresponding cadmium compound, in being soluble in alkaline hydroxides. These three elements resemble each other, and differ from those of family A of this group, in that they can be volatilised, mercury at a temperature about 35 7 °> cadmium and zinc at temperatures approaching ieoo°. These three elements are also alike, in that their vapours con- sist of mono-atomic molecules. 545 2 M 546 Inorganic Chemistry ZINC. Symbol, Zn. Atomic weight = 65. Occurrence. — Zinc is stated to have been found in Australia in the uncombined condition ; with this exception, it is always met with in combination, chiefly as carbonate in calamine or zinc-spar , ZnC0 3 , and as sulphide in zinc-blende or black-jack , ZnS. Other ores are red zinc ore, ZnO ; and franklinite , (ZnFe)0,Fe 2 0 3 . Gahnite , or zinc-spinnelle , has the composition Zn0,Al 2 0 3 . Modes of Formation. — The ores chiefly employed for the pre- paration of zinc are the carbonate and sulphide, although in New Jersey the red oxide and franklinite are used. The process con- sists of two operations, namely, first, the conversion of the ore into oxide of zinc, by calcination ; and, second, the reduction of the oxide by means of coal at a high temperature. The calcination of the natural carbonate is readily accomplished, this compound merely giving up its carbon dioxide at the high temperature — ZnC0 3 = ZnO + C0 2 . In the case of zinc blende, the operation consists in the oxida tion of both the sulphur and the zinc by atmospheric oxygen, thus — ZnS + 30 = ZnO + SO^ Considerable care has to be exercised in order to prevent the formation of zinc sulphate, which, in the subsequent operation, would be reconverted into sulphide, and so lost. The finely- crushed calcined ore is mixed with coke or coal, and heated to bright redness in earthenware retorts, when the oxide is reduced with the formation of carbon monoxide, and the metal distils and is collected in iron receivers. Zinc ores frequently contain small quantities of cadmium, and as this metal is more readily volatilised than zinc, it passes over in the first portions of the distilled product. The two processes now almost exclusively in use for the reduc- tion of zinc, known as the Silesian and the Belgian process,* differ only in metallurgical details, &c. * The old method, known as the English process , or distillation per descensum , is entirely obsolete. For details of this and all other metallurgical processes, the student is referred to treatises on metallurgy, such as Percy. Zinc 547 Commercial zinc usually contains carbon, iron, and lead, and occasionally arsenic and cadmium. It may be obtained in a higher degree of purity by careful distillation, but pure zinc is best ob- tained by first preparing the pure carbonate by precipitation, and then calcining and finally reducing with charcoal obtained from sugar. Properties. — Zinc is a bluish- white, highly crystalline, and brittle metal. At a temperature of 300° it can be readily powdered in a mortar, while between ioo° and 1 50° it admits of being drawn into wire or rolled into thin sheet. The presence of a small quantity of lead greatly enhances this property, but is detrimental when the zinc is required for making brass. Zinc which has been either rolled or drawn, no longer becomes brittle when cold, but retains its malleability. Zinc melts at 420°, and when heated in air much beyond this point, the metal takes fire and burns with a bluish-white flame, the brilliancy of which becomes dazzling if a stream of oxygen be pro- jected upon the burning mass. The product of its combustion is zinc oxide, ZnO, which forms a soft, white, flocculent substance re- sembling wool, and formerly known as philosophers wool. Zinc is permanent in dry air at ordinary temperatures, but when exposed to moist air it tarnishes superficially ; it is also unattacked by water at the boiling temperature. It is soluble in a hot solution of sodium or potassium hydroxide, with evolution of hydrogen (p- 154). Pure zinc is scarcely acted upon by pure sulphuric or hydrochloric acid, either dilute or strong. The presence of small quantities of impurities, however, determines the solution of the metal with the rapid evolution of hydrogen, hence ordinary commercial zinc is readily attacked by these acids, and also decomposes water at the boiling-point with the evolution of hydrogen.* * The difference between the behaviour of acids towards pure and com- mercial zinc, was formerly explained on the ground that the impurities present formed with the zinc a voltaic couple, whereby local electric currents were set up, while in the case of pure zinc no such action took place. The recent observations of Pullinger (Chem. Soc., 57) and Weeren (Berichte, 24) show that this is not a complete explanation. Weeren concludes that the insolubility of pure zinc in dilute acids, is due to the formation of a film of condensed hydrogen upon the surface of the metal, which stops all further action. The addition of oxidising agents, such as hydrogen peroxide, or dilute sulphuric acid which has been electrolysed, and therefore contains persulphuric acid, tends to destroy this film by oxidising the hydrogen, and therefore promotes the solution of the 548 Inorganic Chemistry Zinc is extensively used in the process of galvanising iron, which consists in coating iron with a film of zinc, not by electrical deposi- tion, as would be implied by the name, but by dipping the iron into a bath of molten zinc. The layer of zinc preserves the iron from rusting. Galvanised iron is better able to withstand the action of air and moisture than tinned iron, hence it is extensively used for wire netting, corrugated roofing, water tanks, and other purposes where the metal is exposed to the oxidising influence of air and water. Alloys of Zinc. — Zinc forms a number of useful alloys, the most important of which are the various forms of brass (see Copper). With certain metals, such as tin, copper, and antimony, zinc will mix in all proportions ; while with others, such as lead and bismuth, it is only possible to obtain solid alloys of definite composition. When, therefore, lead and zinc are melted together, although in the molten condition the mixture is homogeneous, on cooling, the metals separate into two layers, the lighter zinc rising to the surface. The separation of the metals, however, is not perfect, for the zinc will have dissolved a certain quantity of the lead (1.2 per cent.), and the lower layer of lead is found to have dissolved a small proportion of zinc (1.6 per cent.), just as water and ether, when shaken together, separate into two layers, the uppermost being an ethereal solution of water, and the lower an aqueous solution of ether. This property is made use of in the extraction of silver from lead (see p. 517). The so-called German silver , or nickel silver , is a nearly white alloy of copper, nickel, and zinc. Bronze coinage consists of 95 parts of copper, 4 of tin, and 1 of zinc, the small proportion of zinc giving to the alloy an increased hardness and durability. Zinc Oxide, ZnO, the only oxide of zinc, occurs native as red zinc ore , the colour being due to the presence of manganese. It is formed as a soft white substance, when zinc is burnt in the air. It is manufactured under the name of zinc white by the combustion zinc. He also finds, that by mechanically removing this layer of hydrogen, either by constantly brushing the metallic surface or placing the materials under reduced pressure, the solution of the zinc by the acid is promoted. It is also found that the character of the surface of the metal, whether smooth or rough, affects the result ; zinc that is unacted upon when its surface is perfectly smooth, is more readily attacked by the dilute acid when its surface is rough. Zinc Chloride 549 of zinc, the fumes being led into condensing-chambers, where the oxide collects. Zinc oxide is a pure-white substance, which when heated becomes yellow, but again becomes white on cooling. When strongly heated in oxygen, it may be obtained in the form of hexagonal crystals ; such crystals are occasionally found in the cooler parts of zinc furnaces. The oxide does not fuse in the oxyhydrogen flame, but, like lime, under these circumstances it becomes intensely incan- descent ; for some time after being so heated it appears phos- phorescent in the dark. It is insoluble in water, and does not combine directly with water to form the hydroxide. It dissolves in acids, giving rise to the different zinc salts. Zinc oxide is largely used in the place of “white lead” as a pigment : although it does not equal white lead in covering power, or body , it possesses the advantage of not being blackened by exposure to atmospheric sulphuretted hydrogen. Zine Hydroxide, Zn(HO) 2 , is formed as a white flocculent pre- cipitate, when either sodium or potassium hydroxide, or a solution of ammonia, is added to a solution of zinc sulphate. The compound is soluble in an excess of either alkali, and is deposited from a strong solution in regular octahedra of the hydrated hydroxide, Zn(HO) 2 ,H 2 0. Both of these compounds on heating, readily lose water, and are converted into the oxide. Zinc Chloride, ZnCl 2 , is formed by the direct combination of zinc with chlorine, or by the action of hydrochloric acid upon the metal. It is also obtained in the anhydrous state by distilling a mixture of mercuric chloride and zinc, or a mixture of anhydrous zinc sulphate and calcium chloride. It is usually prepared on a large scale by dissolving zinc in hydrochloric acid, and after precipitating any manganese and iron, the liquid is boiled down in enamelled iron vessels, until on cooling it solidifies ; it is usually cast into sticks. Zinc chloride is a soft, white, easily fusible solid, which volatilises and distils without decomposition. It is extremely deliquescent, and readily soluble in water and in alcohol, its solution being powerfully caustic. From a strong aqueous solution, deliquescent crystals are deposited, having the composition ZnCl 2 ,H 2 0. When the aqueous solution is evaporated, partial decomposition takes place, hydrochloric acid being evolved, and basic compounds being precipitated, consisting of combinations of the chloride and oxide. Hence, during the concentration of the liquid in the pre- 550 Inorganic Chemistry paration of zinc chloride, hydrochloric acid is added to re-dissolve this compound. A paste made by moistening zinc oxide with zinc chloride, rapidly sets to a hard mass ; this mixture, under the name of oxychloride of zinc, is employed in dentistry as a filling or stopping for teeth. Zinc chloride unites with alkaline chlorides, forming a series of crystalline double salts having the general formula ZnCl 2 ,2RCl. Zinc Sulphate, ZnS0 4 , is formed when zinc is dissolved in sulphuric acid. It is obtained on a large scale by roasting the natural sulphide, whereby it is partially converted into the sulphate, which is then extracted with water. The salt crystallises from its aqueous solution at ordinary tem- peratures in colourless rhombic prisms, ZnS0 4 ,7H 2 0, isomorphous with MgS0 4 7H 2 0. It is extremely soluble in water; ioo parts of water at the ordinary temperature dissolve 160 parts, and at ioo 0 , 653.6 parts, of the crystalline salts. When exposed to the air, the crystals slowly effloresce, and if placed in vacuo over sulphuric acid, or if heated to ioo°, they lose six molecules of water, leaving the monohydrated salt ZnS0 4 ,H 2 0. At a temperature about 300° this is converted into the anhydrous compound, and at a white heat it gives off sulphur dioxide and oxygen, leaving the oxide. The hydrated salt, ZnS0 4 ,6H 2 0, is obtained in the form of mono- symmetric crystals, when the salt is deposited at temperatures above 40°. This compound is isomorphous with MgS0 4 ,6H 2 0. Zinc sulphate combines with alkaline sulphates, forming a series of double salts, having the general formula ZnS0 4 ,R 2 S0 4 ,6H 2 0, which are also isomorphous with the corresponding magnesium compounds (page 531). Zinc sulphate, in common with all the soluble salts of zinc, has an astringent taste, and is poisonous. Zinc Sulphide, ZnS. — The natural compound, zinc-blende , is usually dark-brown or black, and exhibits crystalline forms belong- ing to the regular system. The mineral wurtzite is a less common variety of zinc sulphide, crystallising in hexagonal prisms. Zinc sulphide is obtained as a white amorphous precipitate, when an alkaline sulphide is added to a solution of a zinc salt, or when sulphuretted hydrogen is passed through an alkaline solution of a zinc salt. Precipitated zinc sulphide is insoluble in acetic acid, but readily dissolves in dilute mineral acids, with evolution of sulphuretted hydrogen ; hence the compound is not formed when sulphuretted Cadmium 5 5 1 hydrogen is passed through a solution of a zinc salt containing a free mineral acid. Zinc Carbonate, ZnC0 3 , is obtained as a white powder, when hydrogen sodium carbonate is added to a solution of zinc sulphate. If normal sodium carbonate be employed, the precipitated zinc compound consists of a basic carbonate, whose composition varies with the conditions of temperature and concentration of the liquids. A basic carbonate, having the composition ZnC0 3 ,2Zn(H0) 2 ,H 2 0, is employed as a pharmaceutical preparation, under the name zinci carbonas. CADMIUM. Symbol, Cd. Atomic weight = 111.7. Occurrence. — Cadmium is never found in the uncombined state. The only natural compound of which cadmium is the chief con- stituent, is the extremely rare mineral greenockite , which is the sulphide, CdS. Cadmium occurs in small quantities in many zinc ores, such as the sulphide and carbonate ; and in the process of extracting zinc from these ores, the cadmium is obtained in the first portions of the product of the distillation, partly as metal, and partly as oxide. Mode of Formation. — The crude product of distillation is dis- solved in dilute sulphuric or hydrochloric acid, and the cadmium precipitated as sulphide by means of sulphuretted hydrogen. The cadmium sulphide is then dissolved in strong hydrochloric acid, and precipitated as carbonate by means of ammonium carbonate. The washed and dried carbonate is first converted into oxide by calcination, and finally mixed with charcoal and distilled. Properties. — Cadmium is a bluish-white metal resembling zinc in appearance, but much more malleable and ductile. It tarnishes superficially on exposure to the air, and, when strongly heated, bums with the formation of a brown smoke of cadmium oxidej CdO. The metal is attacked by dilute hydrochloric and sulphuric acids, with the evolution of hydrogen. It readily dissolves in nitric acid, yielding the nitrate, with the formation of oxides of nitrogen. Cadmium is less electro-positive than zinc, and is precipitated in the metallic condition from its solutions by that metal. Cadmium melts at 320 °, and when volatilised in an atmosphere of hydrogen, it forms crystals belonging to the regular system. Cadmium Oxide, CdO, is formed as a brown fume or smoke, 552 Inorganic Chemistry when cadmium burns in the air. It may be obtained by heating the carbonate or nitrate. That obtained by the ignition of the latter salt, is in the form of minute crystals, having a bluish-black appearance. Cadmium oxide is insoluble in water, but dissolves in acids yielding cadmium salts. It is infusible in the oxyhydrogen flame, but is readily reduced when heated on charcoal before the blowpipe ; and the reduced metal, as it volatilises and burns, forms a characteristic brown incrustation of oxide upon the charcoal. Cadmium Chloride, CdCl 2 , is obtained by the action of hydro- chloric acid upon the metal or the oxide. The salt is deposited from the solution in white silky crystals, having the composition, CdCl 2 ,2H 2 0. On exposure to the air the crystals effloresce, and when heated become anhydrous. Cadmium Sulphide, CdS, is obtained as a bright yellow preci- pitate, when sulphuretted hydrogen is passed through a solution of a cadmium salt. The precipitate is soluble in concentrated hydro- chloric and nitric acids, and in warm dilute sulphuric acid. Cad- mium sulphide is insoluble in ammonium sulphide ; this property readily* distinguishes it from arsenious sulphide, which in colour it closely resembles. Cadmium sulphide is used as a pigment, both in oil and water- colours. MERCURY. Symbol, Hg. Atomic weight = 199.8. Occurrence. — In the uncombined state mercury is met with in small globules, disseminated through its ores, especially the sul- phide. It is also occasionally found as an amalgam w r ith silver and gold. The principal ore is cinnabar , HgS, and the chief mines of this ore are those of Almaden (Spain), Idria (Camiola), California, and the Bavarian Palatinate. Modes of Formation. — Mercury may be obtained from the natural sulphides, by either roasting the ore, whereby the sulphur is oxidised to sulphur dioxide, and the metal liberated, or by dis- tillation in closed retorts with lime, when calcium sulphide and sulphate are formed, and the mercury set free. The first method is almost exclusively employed. At Idria the crude ore, consisting of cinnabar mixed with shale and earthy matters, is roasted in a furnace, upon perforated arches, «, «' ; p,P', Fig. 137. The action of the fire and heated air is to Mercury 553 oxidise the sulphur, and volatilise the mercury, and the gases and vapours together pass through a series of flues or chambers, C, C, where the mercury condenses. By the use of a reverberatory furnace (the Alberti furnace) the process can be made continuous. The ore is fed into the furnace through a hopper, and the calcined residue is raked out through an opening at the opposite end of the hearth. The gases are passed first through iron pipes, kept cool by water, and then Fig. 137. through a series of chambers where the remaining metal is con- densed. . , . The method adopted at Almaden is essentially the same as the Idrian process, except that the condensation takes place in a series of pear-shaped earthenware vessels, called aludels , which are con- nected together as shown in Fig. 138. Usually six rows of forty- seven such aludels, are connected with six openings in a chamber Fig. 138. immediately above the furnace. The impure mercury is freed from mechanically mixed impurities by straining or filtering through chamois leather, but from metals in solution, such as zinc, tin, lead, and others, it is purified by distillation. For laboratory purposes, pure mercury is best obtained by distillation in vacuo, by means of the apparatus shown in Fig. 139 (Clarke). In this arrangement the mercury is distilled in a Spreiigel vacuum. The mercury (previously cleaned by being thoroughly agitated with mercuric nitrate) is placed in the reser- 554 Inorganic Chemistry voir R, which is then placed upon the upper shelf S, and by means of the clamp, mercury is allowed to pass into the long wide tube T, and up into the bulb. The air in the tube and bulb escapes down the narrow inner tube, which reaches nearly to the top of the bulb, as seen in the enlarged detail, t. The mercury is allowed to rise in the bulb and fall down the long inner tube, after the manner of the Sprengel pump. The reservoir is then placed upon the lower Fig. 139. adjustable stand, and its height so arranged that the mercury in the bulb falls to the position shown in the figure. This space is a Torricellian vacuum. The mercury is then heated by a ring- burner, B, and the whole is protected from draught by the hood h. As the mercury distils, it passes down the inner tube, and by its fall continues to preserve the Sprengel vacuum within the bulb. Properties.— At ordinary temperatures mercury is a bright, silver- Mercury 555 white liquid metal (hence its old name quicksilver). When cooled to - 38.8° it solidifies to a highly crystalline solid, which is ductile and malleable, and softer than lead. When the liquid is cooled, it contracts uniformly until the solidifying point is reached, when considerable contraction takes place. Solid mercury, therefore, is denser than the liquid metal, and sinks in it. The specific gravity of liquid mercury at o° is 13.596, while that of the solid at its melting-point is 14.193. Mercury in extremely thin films appears a violet colour by transmitted light. Under a pressure of 760 mm. mercury boils at 357 - 25 °? giving a colourless vapour. The density of mercury vapour referred to hydrogen is 100.93 ; hence this element, like its associates in the family to which it belongs, consists of mono-atomic molecules when in a state of vapour. Mercury gives off vapour even at ordinary temperatures, and a gold leaf suspended over mercury in a stop- pered bottle, gradually becomes white upon the surface, owing to its amalgamation with the mercurial vapour. The vapour of mercury is poisonous, giving rise to salivation. Mercury does not tarnish on exposure to the air, and is unacted upon by a large number of gases : hence this liquid is invaluable to the chemist, affording a means of collecting and measuring gases which are soluble in water. When submitted to prolonged heating in the air, it is slowly converted into the red oxide, which at a higher temperature is again decomposed into its elements. Mercury is obtained in the form of a dull-grey powder, when it is shaken up with oil, or triturated with sugar, chalk, or lard. This operation is known as deadening , and is made use of in the pre- paration of mercurial ointment. The grey powder consists simply of very finely-divided mercury, in the form of minute globules. Mercury is not attacked by hydrochloric acid. Strong sulphuric acid is without action upon it in the cold, but when heated the metal dissolves, with evolution of sulphur dioxide. Strong nitric acid rapidly attacks it, with formation of mercuric nitrate and oxides of nitrogen. Cold dilute nitric acid slowly dissolves it, forming mercurous nitrate. Alloys of Mercury. — When mercury is one of the constituents of an alloy, the mixture is called an amalgam. Most metals will form an amalgam with mercury. In some cases, as with the alkali metals, the union is attended with great rise of temperature. In other cases, as with tin, an absorption of heat takes place. 55 6 Inorganic Chemistry Sodium and potassium amalgams are obtained by dissolving various amounts of the metals in mercury. In contact with water they are decomposed, hydrogen being evolved, and the alkaline hydroxide formed. On this account sodium amalgam is frequently used in the laboratory as a reducing agent. When heated to 440°, these amalgams leave behind crystalline compounds, K 2 Hg and Na 3 IIg, which spontaneously inflame in contact with the air. Zinc amalgams are only very slowly acted upon by dilute sul- phuric acid ; therefore, by the superficial amalgamation of the zinc plates used for galvanic batteries, the same result is ob- tained, as though the zinc were perfectly pure (see page 547), and no solution of zinc takes place until the electric circuit is closed. Tin amalgams are employed for the construction of ordinary mirrors. Amalgams of gold, and also copper and zinc, are used in dentistry as a filling or stopping for teeth. Oxides of Mercury. — Two oxides are known, namely, mercu- rous oxide, Hg 2 0 , and mercuric oxide, HgO. Mercurous Oxide, Hg 2 0 , is obtained as an unstable dark-brown or black powder, when sodium hydroxide is added to mercurous chloride. When exposed to the light, or when gently heated, it is converted into mercuric oxide and mercury. Mercuric Oxide, HgO, is produced in small quantity by the pro- longed heating of mercury in contact with air, or by igniting the nitrate. It is prepared on a large scale by heating an intimate mixture of mercuric nitrate and mercury. Obtained by these methods, it is a brick-red crystalline powder ; but when sodium hydroxide is added to a solution of a mercuric salt, the oxide is pre- cipitated as an orange-yellow amorphous powder. When heated, mercuric oxide first darkens in colour, and gradually becomes almost black, but returns to its original bright red colour on cool- ing. At a red heat it is completely decomposed into its elements. Salts of Mercury. — Two series of salts, corresponding to the two oxides, are known — (a) mercurous salts, in which two atoms of the hydrogen of the acids are replaced by the divalent radical, or double atom (Hg 2 ) ; and (/ 3 ) mercuric salts, in which the same amount of hydrogen is replaced by the single divalent atom (Hg). All the mercury salts are poisonous. Mercurous Sulphate 557 (a) MERCUROUS SALTS. Mercurous Chloride, Hg 2 Cl 2 (calomel), is met with in small quantities as the mineral horn mercury. It may be obtained by the addition of sodium chloride, or hydrochloric acid, to a solution of mercurous nitrate. On a large scale it is usually prepared by heating a mixture of mercuric chloride and mercury, when the mercurous chloride sublimes as a white or translucent fibrous cake. When a mixture of mercuric sulphate, common salt, and mercury is heated, mercurous chloride is also obtained, thus HgS0 4 + 2NaCl + Hg = Na 2 S0 4 + Hg 2 Cl 2 . Calomel is perfectly tasteless, and is insoluble in water. When heated, it vaporises without fusing. The density of the vapour that is formed by heating mercurous chloride is 1 17-59, which is half that demanded by the formula Hg 2 Cl 2 . It has been shown, however, that the compound dissociates when vaporised, into mercuric chloride and mercury* Boiling hydrochloric acid de- composes mercurous chloride into mercury, which separates out, and mercuric chloride, which dissolves. Mercurous Nitrate, Hg 2 (N0 3 ) 2 , is deposited in the form of colourless monosymmetric crystals containing 2H 2 0, from a solu- tion of mercury in cold dilute nitric acid. The salt is soluble in water acidulated with nitric acid, but an excess of water causes the precipitation of a basic nitrate having the composition — Hg 2 (N0 3 ) 2 ,Hg 2 0,H 2 0 (or 2Hg 2 (N0 3 )(H0)), which, on boiling, is converted into mercuric nitrate and mercury. If either this or the normal salt, be boiled in the presence of an excess of mercury, a basic nitrate of the composition — 3Hg 2 (N0 3 ) 2 ,2Hg 2 0,2H 2 0 (or Hg 2 (N0 3 ) 2 ,4Hg 2 (N0 3 )(H0)), is obtained. Mercurous Sulphate, Hg-jSO^ is obtained as a white crystalline precipitate, when dilute sulphuric acid is added to a solution of mercurous nitrate. It is very slightly soluble in water. * Harris and Meyer. Berichte, June 1894. 558 Inorganic Chemistry (P) MERCURIC SALTS. Mercuric Chloride, HgCl 2 ( corrosive sublimate ), is formed when chlorine is passed over heated mercury. It is prepared on a large scale, by heating a mixture of mercuric sulphate and common salt ; a small quantity of manganese dioxide being added, to prevent as far as possible, the formation of mercurous chloride. The mercuric chloride sublimes as a white translucent mass. It dis- solves in water to the extent of 6.57 parts in 100 parts of water at io°, and 54 parts in the same volume of water at ioo°, forming an acid- solution from which the salt is deposited in long white silky needles. It readily melts, and volatilises unchanged. It dissolves without decomposition in nitric acid, and in sulphuric acid, and volatilises unchanged from its solution in the latter acid on boiling. Mercuric chloride is a violent poison : the best antidote is albu- men, with which it forms an insoluble compound. It has also strong antiseptic properties, and on this account is largely used by taxidermists. With hydrochloric acid, mercuric chloride forms two crystalline double chlorides, HgCl 2 ,HCl and 2HgCl 2 ,HCl ; and with the alkaline chlorides it forms a number of similar double salts, of which the ammonium compound, HgCl 2 ,2NH 4 Cl,H 2 0, was known to the early chemists under the name sal alembroth. Mercuric Iodide, Hglg. — When mercury and iodine are rubbed together in a mortar, and moistened with a small quantity of alcohol, the red mercuric iodide is formed. It is also obtained by precipitation from a solution of mercuric chloride, upon the addition of potassium iodide. The precipitate first appears yellow, but in a few seconds becomes scarlet. Mercuric iodide is insoluble in water, but readily dissolves in either mercuric chloride or potassium iodide, and also in alcohol and in nitric acid. From its solutions it is deposited in scarlet quadratic octahedra. Mercuric iodide is dimorphous ; when heated to about 150°, the scarlet quadratic crystals are changed into bright yellow rhombic prisms. At ordinary temperatures this yellow rhombic form is unstable, and on being lightly touched it is at once retransformed into the red quadratic form. At very low temperatures, however, the yellow variety is the more stable : thus, when the red crystals Ammoniacal Mercurous Compounds 559 are exposed to the temperature of evaporating liquid oxygen, they pass into the yellow variety. Mercuric Nitrate, Hg(N0 3 ) 2 , is prepared by boiling nitric acid with mercury, until sodium chloride produces no precipitate with a sample of the liquid. If this solution be evaporated over sulphuric acid, deliquescent crystals are obtained of 2Hg(N0 3 ) 2 ,H 2 0, while the mother liquor has the composition Hg(N0 3 ) 2 ,2H 2 0. Mercuric nitrate exhibits a great tendency to form basic salts : thus, when this mother liquor is boiled, the compound Hg(N0 3 ) 2 , HgO,2H 2 0 is precipitated. When this compound, or the normal nitrate, is treated with an excess of cold water, there is formed the still more basic salt Hg(N0 3 ) 2 ,2Hg0,H 2 0. Mercuric Sulphide, HgS {cinnabar ).— When mercury ‘and sulphur are triturated together in a mortar, or when excess of sulphuretted hydrogen is passed into a solution of a mercuric salt, mercuric sulphide is obtained as a black amorphous powder. If this be sublimed, it is obtained as a red crystalline substance. Mercuric sulphide in the red condition, is also obtained by digesting the black amorphous product for some hours in alkaline sulphides. A soluble double sulphide is first formed, which when heated, is decomposed, with the deposition of red mercuric sulphide. This compound is manufactured on a large scale for use as the pigment vermilion. Mercuric sulphide is insoluble in either nitric, hydrochloric, or sulphuric acid. In the presence of an alkali it is soluble in sodium or potassium sulphide, and deposits crystals from these solutions having the composition HgS,Na 2 S,8H 2 0, and HgS,K 2 S,5H 2 0 respectively. Ammoniacal Mercury Compounds.— These may be regarded as ammonium salts, in which two atoms of hydrogen in ammonium (NH 4 ) have been replaced by either (Hg 2 ) in the mercurous , or by (Hg) in the mercuric compounds ; the two atoms so replaced, being either drawn from one and the same ammonium group, or from two. (a) MERCUROUS COMPOUNDS. Mercurous Ammonium Chloride, (NH 2 Hg 2 )Cl, is the black powder produced by the action of aqueous ammonia upon calomel, thus — Hg 2 Cl 2 + 2NH 3 aq = (NH 2 Hg 2 )Cl + NH 4 Claq. 560 Inorganic Chemistry Mercurous Ammonium Nitrate, (NH 2 Hg 2 )N0 2 , is formed, to- gether with other compounds, when aqueous ammonia is added to mercurous nitrate. Mercurous Diammonium Chloride, N ^ aC1 ! Hg 2 or (NH 3 ) 2 N H3CI ) Hg-jCljj, »s obtained when calomel absorbs dry gaseous ammonia. On exposure to the air, it gives up its ammonia, and is reconverted into mercurous chloride. (/9) MERCURIC COMPOUNDS. Mercuric Ammonium Chloride, (NH 2 Hg)Cl (1 infusible white precipitate), is formed when ammonia is added to a solution of mercuric chloride — HgCl 2 + 2NH 3 = (NH 2 Hg)Cl + NH 4 C1. Dimercurie Ammonium Chloride, (NHg 2 )Cl, is obtained by the action of water on the preceding compound. Mercuric Diammonium Chloride, \ H g, or(NH 3 ) 2 HgCl 2 fusible white precipitate ), is obtained by adding mercuric chloride to a boiling aqueous solution of ammonium chloride and ammonia, until the precipitate which first forms, no longer dissolves. On cooling, the solution deposits small crystals belonging to the regular system. Oxy-dimereuric Ammonium Iodide, (NH 2 Hg)I,HgO, is pro- duced by the action of aqueous ammonia upon mercuric iodide, thus — 4NH 3 + 2HgI 2 + H 2 0 = (NH 2 Hg)I,HgO + 3NH 4 I. It is readily produced as a brown precipitate, by adding ammonia to a solution of mercuric iodide in potassium iodide containing an excess of potassium hydroxide. The alkaline solution of potassium mercuric iodide is known as Nessler's solution , and constitutes a delicate reagent for detecting the presence of ammonia. Minute traces of free ammonia in solu- tion produce a yellow or brown coloration with this test. CHAPTER VIII THE ELEMENTS OF GROUP III Family A. Family B. Scandium, Sc Yttrium, Y . • 43*97 Boron, B . 10.9 27.04 69.86 ii3*4 203.7 89.6 Aluminium, A1 . Lanthanum, La . Ytterbium , Yb . 138.5 Gallium, Ga . 173 Indium , In Thallium, T1 With the exception of boron, aluminium, and thallium, the mem- bers of this group are amongst the rarest of the elements * Some of these occur only in minute traces in certain ores of other metals : such is the case with the elements gallium and indium, which are met with in certain specimens of zinc blende, the ore being con- sidered rich in gallium if it contains as much as 0.002 per cent, of this element. Both gallium and indium were discovered by means of the spectroscope ; the latter by Reich and Richter (1863), and named indium on account of two characteristic lines in the indigo- blue part of the spectrum ; gallium by Lecocq de Boisbaudran (1875), and named after his own country. The spectrum of this metal is characterised by two violet lines. One of the most remarkable properties of gallium is its extremely low fusing-point, the metal melting at 30. 15 0 . (For a comparison of the properties of gallium with MendelejefFs eka- aluminium, see p. 109.) Others of these elements are met with in certain rare minerals, thus, lanthanum occurs in the mineral orthite (from Greenland) ; and both yttrium and lanthanum (associated also with the rare elements cerium and erbium) are found in gadolinite, or ytterbiie (from Ytterby). Boron (the typical element of the group) is the only non-metal : all the others exhibit well-marked metallic properties. They all yield sesquioxides of the type R 2 0 3 ; in the case of boron this oxide, B 2 0 3 , is acidic. * For detailed descriptions of the rare elements, the student is referred to larger treatises, or to chemical dictionaries. 2 N 562 Inorganic Chemistry Thallium in many respects is peculiar. It forms two series of compounds ; in one class it functions as a monovalent, and in the other as a tri valent element. In some of its properties it exhibits a close analogy to the alkali metals ; thus, it forms a soluble strongly alkaline hydroxide, T1HO, corresponding to KHO. And many of its salts, such as the sulphate, T1 2 S0 4 ; perchlorate, T1C10 4 , and the phosphates, are isomorphous with the corresponding potassium compounds. Thallium also shows many properties in common with lead, which in the periodic system is the next element in the series (the fourth long series). Thus, the chloride, like lead chloride, is thrown down as a white curdy precipitate on the addition of hydrochloric acid to a soluble salt of the metal, and like lead chloride, thallous chloride is soluble in hot water. Thallous iodide also closely resembles lead iodide, being formed as a yellow crystalline precipitate when potassium iodide is added to a soluble thallous salt. Metallic thallium also bears the closest resemblance to metallic lead. In the thallic compounds this element is more closely related to the other members of this family : thus thallic oxide, T1 2 0 3 ; thallic chloride, T1C1 3 ; and thallic sulphide, T1 2 S 3 , are analogous to the corresponding boron compounds, B 2 0 3 , BC1 3 , B 2 S 3 . BORON. Symbol, B. Atomic weight = 10.9. Occurrence. — The element boron has never been found in the free state. In combination it occurs principally as boric acid in volcanic steam, and as metallic borates, of which the commonest are tincal , a crude sodium borate, or borax, Na 2 B 2 0 7 ; boracite and colemanite, or borate spar, Ca 2 B 6 O u ; and boronatrocalcite , or ulexite , Ca 2 B 6 O u , N a 2 B 4 0 7 , 16 H 2 0. Modes of Formation. — (i.) Boron may be prepared by heating boron trioxide with either sodium or potassium in a covered crucible — B 2 0 3 + 3K 2 = 3K 2 0 + 2B. The fused mass is boiled with dilute hydrochloric acid, and the Boron Trioxide 5^3 boron, which is in the form of a dark-brown powder, is separated by filtration. (2.) The element may also be obtained by heating potassium borofluoride with potassium BF 3 ,KF + 3K = 4KF + B. (3.) Boron is also formed when potassium is heated in the vapour of boron trichloride — BC1 3 + 3K = 3KC1 + B. Properties.— Boron, as obtained by these methods, is a dark greenish-brown powder. When strongly heated in air it burns, uniting both with oxygen and nitrogen, forming a mixture of boron trioxide, B 2 0 3 , and boron nitride, BN. It is unacted upon by air at ordinary temperatures. Boron has no action upon boiling water, but cold nitric acid converts it into boric acid— B + 3HN0 3 = H 3 B0 3 + 3N0 2 . When heated with sulphuric acid it is similarly oxidised 2B + 3H 2 S0 4 = B 2 0 3 + 3S0 2 + 3H 2 0. When fused with alkaline carbonates, nitrates, sulphates, and hydroxides it forms borates of the alkali metals, thus— 2B + 3Na 2 C0 3 = 2Na 3 B0 3 + 3CO. 2B + 6KHO = 2K 3 B0 3 + 3H 2 . Boron dissolves in molten aluminium, which on cooling deposits crystals of a compound of aluminium and boron* Boron Trioxide, B 2 0 3 , is formed when boron bums in the air, or in oxygen. The readiest method for its preparation consists in heating boric acid to redness, when it fuses and gives up water— 2B(HO) 3 = 3H 2 0 + B 2 0 3 . Properties.— The fused mass solidifies to a transparent, colour- * This compound was at one time mistaken for an allotropic modification of boron. 564 Inorganic Chemistry less, vitreous solid, which gradually absorbs atmospheric moisture, and becomes opaque. It is not volatile below a white heat, and on this account, although only a feeble acid, it is capable at high temperatures of displacing strong acids which are volatile, from their combinations ; thus, when boron trioxide is fused with potas- sium sulphate, potassium borate is formed, and sulphur trioxide expelled — B 2 0 3 + 3K 2 S0 4 = 2B(KO ) 3 + 3S0 3 . Boron trioxide at a high temperature is capable of dissolving many metallic oxides, some of which impart to the fused mass a characteristic colour. Boron forms three oxyacids, namely — Orthoboric acid, B(HO) 3 , or H 3 B0 3 . Metaboric acid, B 2 0 2 (HO) 2 , or H 2 B 2 0 4 , or B 2 0 3 ,H 2 0. Pyroboric acid, B 4 0 6 (HO) 2 , or H 2 B 4 O r , or 2B 2 0 3 ,H 2 0. Orthoboric Acid, or Boric Acid, B(HO) 3 , occurs naturally, both in the waters and in the jets of steam which issue from the ground in many volcanic districts, notably in Tuscany. The actual amount of boric acid in these natural jets of steam, or soffioni , is very small ; but as the steam becomes condensed in the pools of water, or lagoons , which often surround the jets, the amount of boric acid with which the water becomes charged, is suffi- cient to constitute this a profitable source of supply. To obtain the acid, large brick-work basins are built round the steam jets, in such a manner that the liquid can be caused to flow from one to another. Water is placed in the highest basin, and, after the steam from the fumaroles beneath it has blown through for twenty-four hours, the liquid is passed on to the second basin, and a fresh supply of water is run into the first. In this way the water passes on through a series of four or five such basins, receiving the steam of the soffioni for twenty-four hours in each. The muddy liquor, after passing through a settling reservoir, is concentrated by evaporation, the heat from the natural steam being utilised. The concentrated liquor, having a specific gravity about 1 . 07 , is allowed to cool in lead-lined tanks ; and the crystals, after being drained, are dried upon the floor of a chamber, also heated by the natural steam. The crude boric acid thus obtained, is purified by recrystallisa- tion. Borax 5^5 Boric acid may be prepared by the action of sulphuric acid, or hydrochloric acid, upon a strong solution of borax— Na 2 B 4 0 7 + 5H 2 0 + 2HC1 = 2NaCl + 4H 3 B0 3 . Properties. — Boric acid crystallises in lustrous white laminae, which are soft and soapy to the touch, ioo parts of water at 18 dissolve 3.9 parts of the acid. The aqueous solution turns blue litmus to a port wine red, similar to the colour produced by car- bonic acid. In contact with turmeric paper, it gives a brown stain, resembling that caused by alkalies, but readily distinguished by not being destroyed by acids, and by being turned black in contact with a solution of sodium hydroxide. Boric acid is more soluble in alcohol than in water ; and when this solution is boiled, a portion of the boric acid volatilises with the alcohol, and imparts a green colour to the flame of the burning vapour. The orthoborates are mostly unstable salts. Metaborie Acid, H 2 B 2 0 4 , is obtained when boric acid is heated to ioo° — 2H 3 B0 3 = 2H 2 0 + H 2 B 2 0 4 . The metaborates are more stable salts than the orthoborates. The acid is dibasic, and forms normal and acid salts, as well as super-acid salts, thus — Normal potassium metaborate . . K 2 B 2 0 4 . Acid potassium metaborate . ♦. HKB 2 0 4 . Super-acid potassium metaborate . HKB 2 0 4 ,H 2 B 2 0 4 . Pyroborie Acid, H 2 B 4 0 7 , is obtained by heating either meta- boric acid, or orthoboric acid, to 140° for some time — 2H 2 B 2 0 4 = H 2 0 + H 2 B 4 0 7 . 4H s B0 3 = 5H 2 0 + H 2 B 4 0 7 . Borax.— The most important salt of pyroborie acid is the sodium salt, ordinary borax, Na 2 B 4 0 7 . This compound occurs naturally as the mineral tincal. It is manufactured from boric acid by double decomposition with sodium carbonate — 4H 3 B0 3 + Na 2 C0 3 = Na 2 B 4 O r + 6H 2 0 + C0 2 . Anhydrous sodium carbonate is added to a boiling solution of 566 Inorganic Chemistry boric acid, and the liquid is then allowed to crystallise, when it forms large transparent prisms belonging to the mono-symmetric system, of the composition Na 2 B 4 O 7 ,10H 2 O. Borax is also obtained from the natural calcium borate, which has the composition Ca 2 B 6 O n . The powdered mineral is boiled with water, and soda ash is added to the mixture, when calcium carbonate is precipitated, and a mixture of borax and sodium metaborate is formed — Ca 2 B fl O n + 2Na 2 CO a = 2CaC0 3 + Na 2 B 4 0 7 + Na 2 B 2 0 4 . On crystallisation, the borax deposits, and the more soluble metaborate remains in the mother liquor. On concentrating these mother liquors, and blowing carbon dioxide through the solution, the metaborate is converted into borax, which is pre- cipitated as a fine meal, leaving sodium carbonate in solution — 2Na 2 B 2 0 4 + C0 2 = Na 2 CO s + Na^O^ When heated, borax loses its water of crystallisation, and swells up, forming a white porous mass, which finally melts to a clear glass. One hundred parts of water at io° dissolve 4.6 parts of crystal- lised borax, and at ioo°, 201.4 parts ; the solution possesses a feeble alkaline reaction. When deposited slowly from warm solutions, borax crystallises in octahedra belonging to the regular system, and having the composition Na 2 B 4 0 7 ,5H 2 0. Boron Trifluoride, BF 3 , is formed when boron is brought into fluorine : the boron takes fire spontaneously in the gas, forming the trifluoride. It is also produced when a mixture of dry powdered fluorspar and boron trioxide is heated to redness in an iron vessel, calcium borate being at the same time produced — 2B 2 0 3 + 3CaF 2 = Ca 3 B 2 0 6 + 2 BF 3 . It is more conveniently prepared by heating together fluorspar, boron trioxide, and sulphuric acid. The reaction may be regarded as taking place in two stages, thus — ( 1 .) CaF 2 + H 2 S0 4 = CaS0 4 + 2 HF. ( 2 .) B 3 0 3 + 6 HF = 3H 2 0 + 2BF 3 . Boron Trichloride 5^7 Properties.— Boron trifluoride is a colourless, pungent-smelling gas, which fumes strongly in moist air on account of its powerful affinity for water. So great is this affinity, that a strip of paper introduced into the gas is charred, by the abstraction of the elements of water. Boron fluoride neither burns, nor supports the combustion of ordinary combustibles. When potassium is heated in the gas, it burns brilliantly, forming the borofluoride. At o° one volume of water dissolves about 1000 volumes of the gas, the absorption being attended with rise of temperature. When the gas is passed into water until the solution is distinctly acid, a mixture of metaboric acid and hydrofluoboric acid is obtained ; the former separates out, while the latter remains in solution— 8BF S + 4H 2 0 = H 2 B 2 0 4 + GHBF 4 . When the gas is passed into water until the latter is saturated, a syrup-like liquid is obtained which chars organic matter and is strongly corrosive. This liquid is sometimes called fluoboric acid, and contains boron trifluoride and water in the proportions represented by the formula 2BF 3 ,4H 2 0; or it may be regarded as consisting of metaboric acid and hydrofluoric acid, as ex- pressed by the formula H 2 B 2 0 4 ,6HF.* In presence of an excess of water, this substance is decomposed into metaboric acid and hydrofluoboric acid. When mixed with its own volume of dry ammonia gas, boron fluoride forms a white crystalline compound, having the composition represented by the formula BF 3 ,NH 3 . This substance may be sublimed without change. Two other compounds with ammonia are known, namely BF 3 ,2NH 3 , and BF 3 ,3NH 3 . These are both colourless liquids, which on being heated give off ammonia, leaving the solid BF 3 ,NH 3 . , The salts of hydrofluoboric acid, HBF 4 , are known as boro fluorides t and are formed by the action of the acid upon metallic hydroxides HBF 4 + KHO = H 2 0 + KBF 4 . In many instances, their aqueous solutions redden litmus ; this is the case with ammonium borofluoride, NH 4 BF 4 , and calcium borofluoride, Ca(BF 4 ) 2 . Boron Trichloride, BC1 3 , is produced when boron is heated in a stream of dry chlorine. It is most readily prepared by passing dry chlorine over an intimate mixture of boron trioxide and charcoal, heated to redness in a porcelain tube. The volatile product is condensed in a tube immersed in a freezing mixture— B 2 0 3 + 3C1 2 + 3C = 3CO + 2BC1 3 . * It is considered very doubtful whether this substance can be regarded as a definite compound. 5 68 Inorganic Chemistry Properties.— Boron trichloride is a mobile, colourless liquid, boiling at 18 . 23 °. It fumes in moist air, being decomposed in contact with water, with formation of boric and hydrochloric acids — BC1 3 + 3H 2 0 = B(HO) s + 3HC1. Boron trichloride unites directly with dry gaseous ammonia, with evolution of considerable heat, forming a white crystalline compound, having the composition 2BC1 3 ,3NH 3 . Boron Hydride, BH 3 .— This compound has never been obtained in a state of purity. When magnesium boride (an impure substance obtained by fusing boron trioxide and magnesium in a covered crucible) is acted upon by hydrochloric acid, a gas is evolved which has a characteristic and unpleasant smell, and which produces headache and sickness when inhaled. The gas is largely hydrogen, containing, however, a certain quantity of boron hydride, which imparts to the flame a green colour, and produces boron trioxide.’ When passed through a heated tube, boron is deposited as a brown film. When burnt with a limited supply of air, or when a cold porcelain dish is depressed into the flame of the burning gas, a brown stain of boron is deposited. Boron Nitride, BN, is formed when boron is strongly heated in nitrogeh or in ammonia. It is best obtained by heating, in a covered platinum crucible, a mixture of one part of dehydrated borax, and two parts of ammonium chloride— Na 2 B 4 0 7 + 2NH 4 C1 = 2BN + B 2 0 3 + 2NaCl + 4H 2 0. Boron nitride is a white amorphous powder. It is insoluble in water, but is slowly acted upon by boiling caustic alkalies, with evolution of ammonia— BN + 3KHO = K 3 B0 3 + NH 3 . Heated in a current of steam it forms boron trioxide and ammonia — 2BN + 3H 2 0 = B 2 0 3 + 2NH 3 . Boron Sulphide, B 2 S 3 , is prepared by heating a mixture of boron trioxide and carbon (made by mixing boron trioxide and soot with oil, and heating the pellets out of contact with air) to bright redness in a stream of vapour of carbon disulphide — 2B 2 0 3 + 3C + 3CS 2 = 6CO + 2B 2 S,. Boron sulphide is a yellowish solid, consisting of small crystals. It has a strong unpleasant smell, and its vapour attacks the eyes. It is immediately decomposed by water, being converted into boric acid and sulphuretted hydrogen — BaSs + 6H 2 0 = 2B(HO) 3 + 3HaS. Aluminium 569 ALUMINIUM. Symbol, Al. Atomic weight = 27.04. Oceurrenee. — Aluminium is one of the most abundant of all the elements, although it has never been found in the uncombined state. In combination with oxygen as A1 2 0 3 , it constitutes such minerals as corundum , ruby, sapphire. As the hydrated oxide, A1 2 0 3 ,H 2 0, it occurs associated with iron oxide in the mineral bauxite , which constitutes the chief source from which the metal itself is obtained. As a double fluoride of aluminium and sodium, Al 2 F 6 ,6NaF, it occurs in the mineral cryolite , and as a hydrated phosphate in the various forms of turquoise. Aluminium is met with in enormous quantities in the form of silicate, constituting the various clays ; and as compound silicates in the felspars , and other common minerals constituting a large proportion of the solid crust of the earth. Mode of Formation. — Aluminium is prepared on a large scale from the mineral bauxite, the process being conducted in four stages:— ( 1.) and (2.) The preparation of pure aluminium oxide, free from iron. (3.) The preparation of a double chloride of aluminium and sodium. (4.) The reduction of the double chloride by means of sodium. (1.) The powdered bauxite (usually containing about 50 per cent, of alumina) is mixed with sodium carbonate and heated for five or six hours in a reverberatory furnace, when carbon dioxide is evolved and sodium aluminate is formed — A1 2 0 3 + 3Na2C0 3 = Al 2 0 3 ,3Na 2 0 + 3C0 2 . (2.) The sodium aluminate is extracted with water, leaving the iron in the form of insoluble oxide. Through the filtered liquid a stream of carbon dioxide is then passed, which decomposes the sodium aluminate, regenerating sodium carbonate, and precipitat- ing hydrated aluminium oxide — Al 2 0 3 ,3Na 2 0 + 3H 2 0 + 3C0 2 = 3Na 2 C0 3 + A1 2 0 3 ,3H 2 0. (3.) The purified alumina, after being washed and dried, is mixed with sodium chloride and powdered wood charcoal, and sufficient water added to enable the mixture to be worked up into balls. These are dried at 150° and then packed into a vertical fireclay 570 Inorganic Chemistry cylinder, where they are strongly heated in a stream of chlorine — A 1 2 0 3 + 3 C + 3 C 1 2 = 3 CO + A 1 2 C 1 6 . I he aluminium chloride combines with the sodium chloride present in the mixture, forming the double chloride, Al 2 Cl 0 , 2 NaCl, which volatilises from the retort, and is condensed in an earthenware receiver as a nearly white crystalline salt, which is almost entirely free from iron. (4-) In order to reduce the double chloride, three charges (each consisting of a mixture of 25 kilos of the salt, 11 kilos of powdered cryolite (as a flux), and 12 kilos of metallic sodium in small pieces) are thrown into a strongly heated reverberatory furnace, and are immediately followed by a fourth charge, containing the same quantity of the double chloride and of cryolite, but without sodium. A violent reaction at first takes place, and after a short time the entire mass is in a state of fusion, the metal separating out beneath the slag — Al 2 Cl c , 3 NaCl + 6Na = 2 A 1 + 8NaCl. At the present time aluminium is almost exclusively obtained by means of the electric furnace. A solution of alumina in fused cryolite is electrolysed by a powerful current, in a carbon-lined crucible ; the crucible being the cathode, and a bundle of carbon rods the anode. The alumina alone is decomposed, the pure metal collecting at the bottom of the crucible. Properties. Aluminium is a tin-white metal, possessing great tensile strength. It is very ductile and malleable, but requires frequent annealing during the process of drawing or hammering. Its specific gravity is 2.58 ; by hammering and rolling it may be raised to 2.68. Its power of conducting heat and electricity is about one-third that of silver. Aluminium is an extremely sonor- ous metal, and when struck it emits a clear and sustained note. It is not tarnished by air under ordinary circumstances, but when strongly heated it becomes oxidised ; and in the condition of thin foil it readily burns in oxygen, forming alumina, Al 2 O s . The metal melts at a temperature about 700°. Aluminium is scarcely acted upon by nitric acid of any strength, but readily dissolves in hydro- chloric acid, and in solutions of sodium or potassium hydroxide with elimination of hydrogen. When heated with strong sulphuric acid, aluminium sulphate is formed, and sulphur dioxide is evolved. Organic acids are almost without action upon aluminium, but in the presence of sodium chloride they are capable of dissolving it to a slight extent. Pure aluminium is scarcely acted upon by water or Alumina 57 * steam, but the presence of impurities such as usually occur in the commercial metal, renders it much more readily oxidised. Aluminium is a highly electro-positive element, and is capable of reducing a number of other metals from their combinations with oxygen or sulphur. Thus, when finely divided aluminium is heated with the oxides of such metals as manganese, chromium, tungsten, uranium, along with lime to form a slag, an energetic action takes place, in which the aluminium combines with the oxygen, and the metals are thrown out of combination, and are obtained as a coherent mass. Similarly, iron pyrites is reduced to the condition of metallic iron, with the formation of aluminium sulphide. Alloys of Aluminium.— The most important of these is an alloy with copper, containing io per cent, of aluminium, and known as aluminium brofize. This alloy has a yellow colour, resembling that of gold; it is scarcely tarnished by exposure to air, and is susceptible of a high polish. Its specific gravity is 7 . 69 , and it possesses a tenacity equal to that of steel, and more than twice that of the best gun-metal. The alloy is malleable, and yields good castings, and on account of its many valuable pro- perties it is employed for a variety of purposes. Aluminium Oxide {alumina), A1 2 0 3 , occurs native in a colour- less crystalline condition as corundum, and coloured by traces of various metallic oxides in such precious stones as ruby, sapphire, and amethyst. In a less pure condition, it occurs in large quantities as emery. These naturally occurring crystalline forms of alumina are extremely hard, ranking second only to diamond. Alumina is obtained in an amorphous condition, by igniting either the pre- cipitated hydroxide, or ammonia alum, thus Al 2 (HO) 6 = 3H 2 0 + A1 2 O s . A1 2 (S0 4 ) 3 ,(NH 4 ) 2 S0 4 = 2NH 3 + H 2 0 + 4S0 3 + Al 2 O s . It is also obtained by the action of carbon dioxide upon sodium aluminate (p. 569 ). In the crystalline form it is obtained by strongly heating a mixture of aluminium fluoride and boron tn- oxide — A1 2 F 6 + B 2 O s = A1 2 0 3 + 2BF 3 . The boron trifluoride volatilises, leaving alumina in the form of rhombohedral crystals. Artificial rubies have been obtained by heating barium fluoride with alumina, and adding a trace of potassium dichromate. 5 7 2 inorganic Chemistry Amorphous alumina is a soft white powder, insoluble in water, but dissolved by acids with the formation of aluminium salts ; after being strongly heated, however, alumina is attacked only with slowness by hydrochloric or sulphuric acid. Aluminium Hydroxides. — Three hydroxides, or hydrated oxides, are known. Thus, when ammonia is added to a solution of an aluminium salt, a white gelatinous precipitate is obtained, which when dried at ioo° consists of the trihydrate, A1 2 0 3 ,3H 2 0, or Al 2 (HO) 6 . If this be heated to 300° it loses 2H 2 0, and is con- verted into the mono-hydrate, A1 2 0 3 ,H 2 0, or A1 2 0 2 (H0) 2 . By the addition of ammonia to a boiling solution of an aluminium salt, and drying the precipitate at 100 6 , the dihydrate is obtained, A1 2 0 3 ,2H 2 0, or Al 2 0(HO) 4 . These compounds are soluble in acids, and all yield the same aluminium salts. Aluminium hydroxide unites with many soluble organic colour- ing-matters, and precipitates them from solution as lakes . Upon -this property depends the use of aluminium salts as mordants in dyeing and calico printing : the colouring-matter being held in the fibres of the material by the aluminium hydroxide, which is pre- viously precipitated upon the fabric. Aluminates.— Alumina is capable of acting as a feeble acidic oxide : thus, the hydroxides are dissolved by sodium or potassium hydroxide, yielding salts known as aluminates. Certain alumi- nates occur native, such as spinelle (magnesium aluminate), Al 2 0 3 ,MgO, and chrysoberyl (beryllium aluminate), Al 2 0 3 ,Be0. Sodium aluminate is now manufactured on a large scale, for the preparation of the metal (p. 569) and also of aluminium salts. It is readily decomposed even by carbonic acid (p. 569), and by aluminium chloride — Al 2 0 3 ,3Na 2 0 + A1 2 C1 6 = 6NaCl + 2A1 2 0 3 . On the manufacturing scale powdered cryolite is employed to effect this decomposition, the aluminium hydroxide being pre- cipitated, and the sodium fluoride going into solution — Al 2 0 3 ,3Na20 4 - Al 2 F 6 ,6NaF = 12NaF + 2A1 2 0 3 . Aluminium Sulphate, A1 2 (S0 4 ) 3 ,18H 2 0, is found native as the minerals hair salt and aluminite , the latter being a basic salt having the composition A1 2 0 3 S0 3 ,9H 2 0 The normal sulphate The Alums 573 is obtained by dissolving the hydrated oxide in sulphuric acid. Large quantities of commercial aluminium sulphate are made, by directly dissolving bauxite in sulphuric acid. The product, how- ever, contains iron, which is detrimental to the technical uses to which the sulphate is applied, and from which therefore it must be carefully purified. Pure aluminium sulphate is prepared from either bauxite, or cryolite, by first forming sodium aluminate : in the former case by heating the mineral with sodium carbonate (p. 569), and in the case of cryolite by boiling with milk of lime — Al 2 F 6 ,6NaF + 6Ca(HO) 2 = 6CaF 2 + 6H 2 0 + Al 2 0 3 ,3Na 2 0. The sodium aluminate, free from iron, is then decomposed by carbon dioxide, as already described, and the precipitated hydrated oxide dissolved in sulphuric acid. On concentration, the mass solidifies to a white, difficultly crystallisable solid. B The Alums. — Aluminium sulphate unites with certain other sulphates, forming double salts, which belong to a class of com- pounds known as the alums. The most important of these compounds is the double sulphate of aluminium and potassium, A1 2 (S0 4 )3,K 2 S04,24H 2 0, known as potassium alum , or simply alum. The alums have the general formula R 2 (S0 4 ) 3 ,M 2 S0 4 ,24H 2 0, in which R may be either aluminium, iron, chromium, manganese (indium or gallium), and M a monovalent element or group, such as sodium, potassium, or ammonium. 574 Inorganic Chemistry These compounds are all isomorphous, crystallising in the regular system (usually in cubes or octahedra) with twenty-four molecules of water. Fig. 140 represents a crystal of potassium alum (A) and potassium chromium alum (B). In naming the alums* it is usual, when the salt contains aluminium , only to introduce the name of the monovalent element or group : thus, ammonium alum , ox potassium alum , signifies the double sulphate of ammonium, or potassium, and aluminium. If, on the other hand, the compound contains no aluminium, the names of both metals are used, thus, potassium chromium alum , ammonium iron alum. A second class of double sulphates is known, which resemble the alums, although they are not isomorphous with them. These are termed pseudo- alums. They may be regarded as alums, in which the two atoms of the monovalent element are replaced by one atom of a divalent element, thus — Manganese aluminium, pseudo-alum . Iron aluminium, pseudo-alum Copper iron, pseudo-alum Zinc iron, pseudo-alum .... Magnesium manganese, pseudo-alum . Al 2 (S0 4 ) 3 MnS0 4 ,24H 2 0. Al 2 (S0 4 ) 3 FeS0 4 ,24H 2 0. Fe 2 (S0 4 ) 3 CuS0 4 ,24H 2 0. Fe 2 (S0 4 ) 3 ZnS0 4 ,24H 2 0. Mn 2 (S0 4 ) 3 MgS0 4 ,24H 2 0. The alums are all soluble in water, and their solutions have an acid reaction and possess an astringent taste. When heated, they gradually part with water, and at higher temperatures are broken up into oxides and alkaline sulphates ; in the case of ammonium alums, leaving only the metallic oxide. Potassium Alum, A1 2 (S0 4 ) 3 ,K 2 S0 4 ,24H 2 0, is prepared by the addition of the requisite quantity of potassium sulphate to aluminium sulphate. A considerable quantity of alum is also obtained from a naturally occurring basic potassium alum, known as alum stone , or alunite , which has the composition A1 2 (S0 4 ) 3 , K 2 S0 4 ,2A1 2 0 3 ,8H 2 0. At Tolfa this is first calcined, and after- wards lixiviated with water, which dissolves the potassium alum, leaving alumina undissolved. The alum so obtained is known as Roman alumj and although it has a reddish colour, due to the * Selenic acid (the selenium analogue of sulphuric acid) forms a similarly constituted series of double selenates, crystallising in the same form, and with the same number of molecules of water. The system of nomenclature adopted for these compounds is the same : thus, ammonium selenio-alum signifies the double selenate of ammonium and aluminium, while potassium chromium selenio-alum represents the double selenate of potassium and chromium. Alu?n 575 presence of iron, this iron is present only as the insoluble oxide, which is readily removed, and the salt is in reality extremely pure. Alunite is also converted into alum, by treating the calcined mineral with sulphuric acid, and adding the requisite quantity of potassium sulphate. A large quantity of alum is manufactured from alum shale , which is a bituminous mineral, consisting chiefly of aluminium silicate, with finely-divided iron pyrites dissemi- nated throughout the mass. The shale is usually first roasted, and is then exposed to the action of air and moisture, whereby the oxidation of the pyrites is completed. The result of this oxidation is the formation of sulphuric acid, which, acting upon the aluminium silicate, forms aluminium sulphate, while the iron is converted into ferrous and ferric sulphates, and ferric oxide. The oxidised mass is then lixiviated with water, and, after concentra- tion, the requisite quantity of potassium chloride or sulphate is added to the hot liquor. (The use of potassium chloride is pre- ferable, as by double decomposition the ferrous and ferric sulphates are converted into the very soluble chlorides, and an equivalent amount of potassium sulphate is formed.) The liquor is stirred mechanically during its cooling, whereby the alum is deposited in small crystals known as alum meal , which permit of its more ready purification by recrystallisation. Alum crystallises in fine colourless regular octahedra, which, on exposure to the air, become coated with a white efflorescene, due not to loss of water, but to absorption of atmospheric ammonia, and the formation of a basic salt. The solubility of alum in water increases rapidly with rise of temperature. Thus, ioo parts of water at o° dissolve 3.9 parts of alum; at 50°, 44.1 parts; and at ioo°, 357.5 parts. Alum is in- soluble in alcohol. When heated to 42 0 , alum loses 11 molecules of water; and when heated to 6i° in a closed vessel over sulphuric acid, it parts with 18 molecules. On the application of heat, alum first melts in its own water of crystallisation, which is gradually expelled, until at a dull red heat the salt is converted into a white porous mass, known as burnt alum. At a still higher temperature it is broken up into potassium sulphate, alumina, and sulphur trioxide. Burnt alum is only very slowly dissolved by water. The chief use of alum is as a mordant in dyeing, alum being a salt which is much more easily obtained in a state of purity than aluminium sulphate. By the addition of 576 Inorganic Chemistry sodium hydroxide or carbonate to a solution of alum, until the pre- cipitate first thrown down is just redissolved, a basic alum is pro- duced known as neutral alum — 2A1 2 (S0 4 ) 3 ,K 2 S0 4 + 6NaH O = A1 2 (S0 4 ) 3 ,A1 2 (H 0) fl ,K 2 S0 4 + 3Na 2 S0 4 + K 2 S0 4 . This solution gives up its alumina to the fabric with great ease, and on this account is used by dyers and calico printers as a mordant. When this solution is heated to 40°, ordinary alum is re-formed, and a precipitate is obtained consisting of another basic salt, hav- ing the same composition as alunite , thus — 2A1 2 (S0 4 ) 3 ,A1 2 (H0) 6 ,K 2 S0 4 = A1 2 (S0 4 ) 3 K 2 S0 4 , + A1 2 (S0 4 ) 3 ,2A1 2 0 3 ,K 2 S0 4 +6H 2 0. Aluminium Fluoride, A1 2 F 6 . — This compound may be prepared by passing gaseous hydrochloric acid over a mixture of fluorspar and alumina heated to whiteness in a graphite tube, when aluminium fluoride volatilises, leaving calcium chloride — 3CaF 2 + A1 2 0 3 + 6HC1 = 3H 2 0 + SCaCLj + A1 2 F 6 . In the form of a crystalline hydrate it may be obtained by dissolving alumina in aqueous hydrofluoric acid — A1 2 0 3 + 6HF + H a O = A1 2 F 6 ,7H 2 0. Aluminium fluoride forms colourless rhombohedral crystals, which are in- soluble in water. It combines with alkali fluorides, forming insoluble double fluorides, of which the sodium compound is the most important, Al 2 F 6 ,6NaF. This compound occurs native as the mineral cryolite. Aluminium Chloride, A1 2 C1 6 . — This compound is produced when powdered aluminium is strongly heated in chlorine, or with certain metallic chlorides, such as zinc chloride. It is best obtained by passing chlorine over a strongly-heated mixture of alumina and charcoal. An aqueous solution of aluminium chloride may be obtained by dissolving alumina in hydrochloric acid. On evaporation, the solution deposits crystals of a hydrate, A1 2 C1 6 ,12H 2 0. Aluminium chloride forms white hexagonal crystals, which fume strongly in moist air. When gently heated it vaporises, and sub- limes without fusion. When heated under pressure of its own vapour, the compound melts. It dissolves in water with the evolution of heat, and the solution, on evaporation, deposits the hydrated chloride, which, on being heated, breaks up into hydro- chloric acid, water, and alumina — Al a Cl g ,l2H 2 0 = 6HC1 + 9H a O + Al 2 O a . Thallium 577 Aluminium chloride unites with other metallic chlorides, forming double salts, of which the sodium compound Al 2 Cl 0 ,2NaCl (page 570 ) is the most important. It also combines with ammonia, forming the compounds A 1 2 C 1 6 , 6 NH 3 and A1 2 C1 6 ,2NH 3 . Aluminium Sulphide, A1 2 S 3 .— When finely divided aluminium is heated with iron pyrites, an energetic reaction takes place ; metallic iron being reduced, and aluminium sulphide being formed. The same compound is produced when sulphur is thrown upon strongly heated aluminium. As obtained by these methods, aluminium sulphide is a greyish black solid, which, when thrown into water, is converted into the oxide with evolution of sul- phuretted hydrogen — A1 2 S 3 + 3H 2 0 = A1 2 O s + 3H 2 S. The compound is decomposed in the same manner by atmos- pheric moisture, when exposed to the air. THALLIUM. Formula, TL Atomic weight = 203.7. History. — Thallium was discovered by Crookes ( 1861 ) in the seleniferous deposit from a sulphuric acid manufactory. In the spectroscopic examination of certain residues obtained in the ex- traction of selenium from this deposit, the presence of an unknown element was manifested, by the appearance of one bright green line. From its characteristic spectrum, the name thallium (signi- fying a green twig) was given to the element. Occurrence. — Thallium is found in small quantities in many varieties of iron pyrites, and when these are employed in the manufacture of sulphuric acid, oxide of thallium collects in the flue dust of the pyrites burners. Thallium also occurs associated with copper, selenium, and silver, in the rare mineral crookesite. Mode of Formation. — The metal is obtained by reducing the sulphate, by immersing strips of zinc into the solution. The thal- lium is deposited upon the zinc, as a spongy or crystalline mass, which is then pressed together, and fused beneath potassium cyanide in a crucible. Properties. — Thallium is a soft heavy metal, resembling lead. It is readily cut with a knife, and leaves a streak when drawn across paper. When preserved out of contact with air, it is a tin- 2 o 578 Inorganic Chemistry white lustrous metal ; but on exposure to the air, it tarnishes upon its surface, with the formation of black thallous oxide. Its specific gravity is u. 8 , and it melts at 290 °. When exposed to air and moisture, or when placed in water which is free to absorb atmospheric oxygen, the metal is slowly converted into thallous hydroxide, which is soluble in water, and imparts to the liquid a strong alkaline reaction. The solution absorbs carbon dioxide, with the formation of thallous carbonate. When heated in the air thallium melts, and rapidly oxidises to thallium trioxide, T1 2 0 3 ; heated in oxygen it burns, forming the same oxide. It readily burns when heated in chlorine, producing thallous chloride, T1C1. The metal is soluble in dilute acids. Oxides of Thallium.— Two oxides are known, namely, thallous oxide, T1 2 0, and thallic oxide, Tl 2 O a . Thallous Oxide, T1 2 0, forms as a dark grey film upon the surface of the metal, on exposure to the air. It may also be obtained by heating the hydroxide to ioo°. It dissolves in water forming the hydroxide. Thallous Hydroxide is obtained by the addition of barium hydroxide to a solution of thallous sulphate, the precipitated barium sulphate being removed by filtration— T1 2 S0 4 + Ba(HO ) 2 = BaS0 4 + 2T1HO. The solution, on concentration, deposits yellow needle-shaped crystals of T1H0,H 2 0. Thallous hydroxide is soluble in water, yielding an alkaline solution, which gives a brown stain upon turmeric paper. This stain soon disappears, owing to the de- struction of the colouring -matter, and is thereby distinguished from the similar stains produced by sodium and potassium hydroxides. Thallic Oxide, Tl20 3 , is obtained when thallium bums in the air, or when thallium oxyhydroxide, TIO(HO), is heated to ioo°. It forms a dark reddish powder, insoluble in water. In warm dilute sulphuric acid it dissolves, forming thallic sulphate— T1 2 0 3 + 3H 2 S0 4 = T1 2 (S0 4 ) 3 + 3H 2 0, but with hot concentrated acid oxygen is evolved, and thallous sulphate formed — T1 2 O s + HjSO, = TljSO, + O, + H,0. Thallic Oxy salts 579 At a red heat thallic oxide is converted into thallous oxide with loss of oxygen. Thallium Oxyhydroxide, TIO(HO), is formed by the action of potassium hydroxide upon thallium trichloride— T1C1 S + 3KHO = 3KC1 + H 2 0 + TIO(HO). Thallous Chloride, T1C1, is obtained as a white curdy precipi- tate, when hydrochloric acid is added to a solution of a thallous salt. It is considerably more soluble in hot than in cold water : 100 parts of water at 16 0 dissolve 0.265 parts ; and at ioo°, 1.427 parts of thallous chloride. Thallie Chloride, T1C1 3 , is formed by passing chlorine through water, in which thallous chloride is suspended. The solution so obtained, on evaporation in vacuo, deposits colourless transparent crystals of T1C1 3 ,2H 2 0. When either thallium or thallous chloride is gently heated in a stream of chlorine, a compound is obtained, having the composi- tion T1C1 3 ,T1C1, or T1 2 C1 4 . If this be further heated, it loses chlorine, and is converted into a yellow crystalline compound of the composition T1C1 3 ,3T1C1, or T1 4 C1 6 , thus — 2T1 2 C1 4 = Cl 2 + T1 4 C1 6 . Thallous Oxysalts. — The sulphate T1 2 S0 4 , and nitrate T1N0 3 , are best obtained by dissolving the metal in the respective acids. Both salts are soluble in water. Thallous Carbonate, T1 2 C0 3 , is prepared by saturating a solu- tion of thallous hydroxide with carbon dioxide. The salt forms long white prismatic (monosymmetric) crystals, which are mode- rately soluble in water, giving an alkaline solution. Thallous Phosphate, T1 3 P0 4 , is obtained by precipitation from a thallous solution, by the corresponding potassium phosphate. The monohydrogen phosphate, HT^PO^ on being heated to 200 °, is converted into pyrophosphate — 2HT1 2 P0 4 = h 2 o + ti 4 p 2 o 7 , and the dihydrogen salt, on being ignited, yields the metaphos- phate — h 2 tipo 4 = h 2 o + tipo 3 . Thallic Oxysalts. — The chief of these are thallic sulphate, 580 Inorganic Chemistry T1 2 (S0 4 ) 3 ; and thallic nitrate, T1(N0 3 ) 3 . They are obtained by the action of sulphuric acid and nitric acid respectively upon thallic oxide, T1 2 0 3 . Thallic sulphate forms colourless crystals of the composition T1 2 (S0 4 ) 3 ,7H 2 0. It is decomposed by excess of water, with precipitation of the hydrated oxide ; and when heated, yields thallous sulphate, sulphur trioxide, and oxygen — T1 2 (S0 4 ) 3 = T1 2 S0 4 + 2S0 3 + 0 2 . Thallic nitrate is deposited in colourless crystals of T1(N0 3 ) 3 ,8H 2 0 which are decomposed in the presence of much water. CHAPTER IX THE ELEMENTS OF GROUP IV Family A. Titanium, Ti Zirconium, Zr Cerium, Ce . Thorium, Th Family B. 48 Carbon, C . 11.97 90.4 Silicon, Si . 28.3 141.2 Germa?iium, Ge. . 72 232 Tin, Sn • 117.35 Lead, Pb . . 206.39 Family A consists of four rare elements.* Titanium, as the oxide Ti0 2 , occurs in the three rare minerals — rutile , brookite , and anatase. The metal is extremely difficult to isolate in a pure state, owing to the fact that it unites directly with nitrogen, form- ing a nitride. Zirconium is met with as the silicate ZrSi0 4 (or Zr0 2 ,Si0 2 ) in the mineral zircon. Like silicon, it has been obtained in two forms, crystalline and amorphous. The latter variety, when gently heated, burns in the air, while the crystalline variety requires the high temperature of the oxyhydrogen flame for its ignition. Cerium occurs associated with lanthanum, in the rare minerals cerite and orthite , and with yttrium and ytterbium in gadolinite and wohlerite. Thorium is found in the extremely rare minerals, thorite and orangeite , met with in Norway. Family B. — In this family the rare element germanium forms a link between carbon and silicon on the one hand, and tin and lead on the other. Carbon (the typical element) is essentially non-metallic, and forms an acidic oxide. Silicon approaches more nearly to the metals in its physical properties, but its oxide is still acidic, and no compounds are known in which silicon functions as a basic element. Germanium is both metallic and non-metallic ; its oxide * For descriptions of these rare elements, tht student is referred to larger treatises. 581 582 Inorganic Chemistry unites with acids ; and it also combines with alkaline hydroxides, forming germanates corresponding to silicates. Tin is a still more basic element, forming well-marked salts with acids ; but it is also acidic, and with alkalies forms stannates. Carbon and silicon exhibit a close relationship. They both form allotropes, which correspond in many respects. They both unite with hydrogen, forming the analogous compounds CH 4 and SiH 4 ; and with hydrogen and chlorine they form the similarly con- stituted compounds, chloroform, CHC 1 3 ; and silicon chloroform, SiHCl 3 . Tin and lead approach more nearly to each other, especially in their physical properties, than to the other members of the family. They both form compounds, in which the metals function both as divalent and tetravalent elements. Although in the case of lead (as often happens with the heaviest metals of a family), the element exhibits much greater readiness to act in the lower state of atomicity. Until quite recently (1893), no compound was known in which an atom of lead is united with four monovalent atoms, although lead ethide, Pb(C 2 H 6 ) 4 , had been obtained. Now, how- ever, the compound PbCl 4 has been produced, corresponding to SnCl 4 , which it resembles in many respects ; and still more recently (1894), the tetrafluoride has been obtained. Carbon, as usual with the typical elements, stands apart from the other members of the family in many of its attributes. Thus, its oxides are both gaseous : it also forms a vast number of com- pounds with hydrogen, oxygen, and nitrogen, the study of which constitutes the science of organic chemistry. This element has already been treated in Part II. (page 250). SILICON. Symbol, Si. Atomic weight = 28.3. Occurrence. — Silicon is not known to occur in the uncombined state, although in combination it is the most abundant and widely distributed of all the elements, with the exception of oxygen. In combination with oxygen, as silicon dioxide or silica , Si 0 2 , it occurs as flint, sand , quartz , rock crystal , and chalcedony j while in combination with oxygen and such metals as calcium, magnesium, and aluminium, it occurs in clay and soil, and constitutes a large number of the rocks which make up the earth’s crust. Silicon, in Silicon 583 combination with oxygen, is also met with in the vegetable kingdom, being absorbed by plants from the soil. Modes of Formation.— ( 1.) Silicon may be obtained by strongly heating a mixture of potassium silico-fluoride and potassium— K s SiF 6 + 2K 2 = Si + 6KF. The mass, after cooling, is treated with water, which dissolves the potassium fluoride, leaving the liberated silicon. (2.) This element may also be prepared by heating sodium in a stream of the vapour of silicon tetrachloride — SiCl 4 + 2Na 2 = Si + 4NaCl. As obtained by either of these methods the silicon is in the form of an amorphous, dark brown powder. (3.) Silicon is obtained in a crystalline condition, by passing a slow stream of the vapour of silicon tetrachloride over aluminium, previously melted in a current of hydrogen ; the volatile aluminium chloride passes on in the stream of gas, and the liberated silicon dissolves in the excess of aluminium — 3SiCl 4 + 4A1 = 3Si 4- 2A1 2 C1 6 . As the mass cools, silicon is deposited in the form of long, lustrous, needle-shaped crystals. (4.) The most convenient method for the preparation of crystal- lised silicon, consists in heating in a crucible a mixture of 3 parts of potassium silico-fluoride, 1 part of sodium, and 4 parts of granu- lated zinc. The regulus so obtained contains crystallised silicon. It is gently heated, and the excess of zinc drained away, the remainder being removed by treatment with acids. Properties.— Amorphous Silieon, as obtained by the reactions Nos. 1 and 2, is a dark brown amorphous powder, having a specific gravity of 2.15. When heated in the air it burns with the forma- tion of silicon dioxide, which, being non-volatile, coats the particles of the element, and protects it from complete oxidation. It burns when heated in a stream of chlorine, with formation of silicon tetrachloride. It is insoluble in water, and in all acids except hydrofluoric acid, in which it dissolves, with the formation of silico- fluoric acid and evolution of hydrogen — Si -l- 6HF = H 2 SiF 6 + 2H,. 5^4 Inorganic Chemistry On boiling with potassium hydroxide, it forms potassium silicate and hydrogen — Si + 2 KHO + H a O = K 2 Si 0 3 + 2 H* Crystallised Silicon. — As obtained by reactions Nos. 3 and 4, silicon is a brilliant, steely-grey substance, crystallised in needles derived from the rhombic octahedron. The specific gravity of the crystals is 2.34 to 2.49. Crystallised silicon does not burn in oxygen, even when strongly heated : it burns when heated in chlorine, and takes fire spontaneously when brought into fluorine. It is not soluble in any acid except a mixture of nitric and hydro- fluoric acids. Crystallised silicon is very hard, being capable of scratching glass. When silicon is exposed to a high temperature, out of contact with air, it becomes denser and harder, and has been obtained in the form of small, steel-grey nodules, showing a crystalline structure, and having a specific gravity as high as 3.0.* Silicon Hydride, SiH 4 . — This compound is evolved at the negative electrode (along with hydrogen), when dilute sulphuric acid is electrolysed, the electrodes consisting of aluminium con- taining silicon. In an impure condition, also mixed with hydrogen, this gas may be obtained by the action of hydrochloric acid upon magnesium silicide — SiMg 2 + 4 HC 1 = 2 MgCl 2 + SiH 4 . * Although silicon in combination is such an abundant element, constituting, as it does, about one-fourth of the total weight of the solid crust of the earth,' in the free state it must still be regarded as somewhat of a rarity, and con- sequently a good deal of uncertainty exists as to its properties. From differ- ences that have been observed in the substance, as obtained by different methods, and from the close analogy that exists between silicon and carbon, it was at one time believed that three allotropes of this element existed, corre- sponding to those of carbon. Amorphous silicon was considered to represent charcoal. A crystalline substance obtained by Wohler, by heating potassium silico-fluoride with aluminium, has been regarded as corresponding to graphite, and called graphitic silicon ; while the octahedral crystals of silicon prepared by reactions 3 and 4 given above (Deville), were thought to be the analogue of diamond ; and this substance has, therefore, been called diamond , or adaman- toid silicon. There is considerable doubt as to whether the silicon obtained by all these various methods was sufficiently pure to warrant this classification, and this doubt is not diminished by the recently discovered fact, that silicon unites with carbon, forming a hard crystalline substance, which has received the name carborundum. Silicon Fluoride 585 (Magnesium silicide for this reaction may be prepared by fusing together, in a covered crucible, a mixture of dry magnesium chloride 40 parts, dry sodium chloride 10 parts, sodium silico-fiuoride 35 parts, and metallic sodium 20 parts.) Pure silicon hydride is prepared by acting upon triethyl silico- formate with metallic sodium. The mode of action of the sodium is not known, the ethyl silico-formate breaks up into silicon hydride and ethyl silicate — 4SiH(OC 2 H 6 ) 3 = SiH 4 + 3Si(OC 2 H 6 ) 4 . Properties. — Silicon hydride is a colourless gas. As obtained by the first two methods it inflames spontaneously. The pure gas does not possess this property. Its ignition point, however, is very low, and if the gas be slightly warmed, or if a jet of it be caused to impinge upon an object a few degrees above the ordinary temperature, the gas at once takes fire and burns with a brightly luminous flame : it is also rendered spontaneously inflammable by reduction of pressure, or by admixture with hydrogen. When brought into chlorine the gas takes fire, with formation of silicon chloride and hydrochloric acid. Silicon Fluoride, SiF 4 . — This compound is formed when silicon is brought into fluorine, the silicon taking fire spontaneously in the gas. It is prepared by the action of sulphuric acid upon a mixture of powdered fluorspar and sand — 2CaF 2 + 2H 2 S0 4 + Si0 2 = 2CaS0 4 4 - 2H 2 0 4 - SiF 4 . Properties. — Silicon fluoride is a colourless, fuming gas. It is not inflammable, and does not support combustion. It is de- composed by water into hydrofluosilicic acid, and silicic acid, hence the gas cannot be collected over water — 3SiF 4 + 3H 2 0 = 2H 2 SiF 6 4 - H 2 Si0 3 . The silicic acid is precipitated as a gelatinous mass. Each bubble of gas as it comes in contact with the water is at once decomposed, and a little sack-like envelope of silicic acid is formed round it. On filtering the liquid, a solution of hydrofluo- silicic acid is obtained. When silicon fluoride is passed over strongly heated silicon, a white powder is obtained, having the composition Si 2 F 6 . 586 Inorganic Chemistry Silicon Chloride, SiCl 4> is formed when silicon is heated in a stream of chlorine. Under these circumstances the silicon burns in the gas. It is obtained by heating an intimate mixture of silica and carbon in a stream of chlorine, and passing the products through a cooled tube — Si0 2 + 2C + 2C1 2 = 2CO + SiCl 4 . Properties. — Silicon chloride is a colourless liquid, which fumes strongly in moist air, and boils at 58.3°. It is decomposed by water into silicic and hydrochloric acids — SiCl 4 + 4H 2 0 = Si(HO) 4 + 4HC1, and the silicic acid so formed, passes either entirely or in part into the dibasic acid, thus — Si(HO) 4 = SiO(HO) 2 + H 2 0. Disilicon hexachloride (also known as silicon trichloride), S^Clg, is formed when the vapour of silicon tetrachlorif is passed over strongly heated silicon — 3SiCl 4 + Si = 2 TIN. Symbol, So. Atomic weight = 117.35. Occurrence. — Tin does not occur in nature in the uncombined state.* It is met with chiefly as the oxide Sn 0 2 in the mineral tin-stone or cassiterite , t which is found in immense deposits, although in comparatively few localities. It is usually associated with arsenical ores, copper pyrites, wolfram (a tungstate of iron and manganese), and other minerals. Occasionally it is met with in nodules of nearly pure oxide, known as stream-tin. Mode of Formation. — Tin is obtained exclusively from tin- stone : and the process with ordinary ore consists of three opera- tions, namely— (1) calcining, (2) washing, (3) reducing or smelting. If the ore be nearly pure tin-stone it may be at once smelted. The finely crushed ore, after being washed from earthy matters, * Metallic tin has been found in Bolivia, but its origin, whether natural or artificial, is doubtful. t Cassiterides, the ancient name for the British Isles, is derived from the fact that tin-stone was found in large quantities in Devonshire and Cornwall Tin 59i is calcined in a reverberatory furnace. The sulphur and arsenic pass away as sulphur dioxide and arsenious oxide, and are led into condensing flues, where the arsenic deposits and is collected. The iron and copper are oxidised to oxide and sulphate. This calcina- tion is sometimes conducted in the revolving calciner, shown on page 446. The calcined ore is next washed, whereby copper sulphate is dissolved, and the iron oxide and other light matters are removed. The purified ore is then mixed with powdered anthracite, and smelted in a reverberatory furnace — Sn0 2 + 2C = 2CO + Sn. The metal so obtained, is purified by first heating it upon the hearth of a similar furnace, until the more readily fusible tin melts and flows away from the associated alloys ; and afterwards by stirring into the molten tin so separated, billets of green wood, which results in the separation of a scum or dross carrying with it the impurities. Properties. — Tin is a bright white metal, which retains its lustre unimpaired in the air. It is sufficiently soft to be cut with a knife, but is harder than lead, although less hard than zinc. At ordinary temperatures it is readily beaten out into leaf (known as tin-foil), and may be drawn into wire ; but at temperatures a little below its melting-point (228°) it becomes brittle and may be powdered. Tin may be obtained in the form of crystals, by melt- ing a quantity of the metal in a crucible, and when partially solidified, pouring out the remaining liquid portion. Its crystalline character is also seen by pouring over the surface of a block of cast tin, or a sheet of ordinary tinned iron, a quantity of warm dilute aqua-regia, when the surface of the metal will immediately exhibit a beautiful crystalline appearance. When a bar of tin is bent, it emits a faint crackling sound, and if quickly bent backwards and forwards two or three times, the metal becomes perceptibly hot at the point of flexure. These phenomena are due to the friction of the crystalline particles. When strongly heated, tin takes fire and burns, forming stannic oxide, Sn0 2 . It is oxidised by both sulphuric and nitric acids ; thus, when heated with strong sulphuric acid, stannous sulphate and sulphur dioxide are produced — Sn + 2H 2 S0 4 = SnS0 4 + S0 2 + 2H 2 0. 59 2 Inorganic Chemistry The strongest nitric acid (specific gravity, 1.5) is without action upon tin. Ordinary concentrated nitric acid (specific gravity, 1.24) attacks it with violence, forming metastannic acid (p. 593), while in cold dilute acid, it slowly dissolves with the production of stannous nitrate — 4Sn + 9HN0 3 = 4Sn(N0 3 ) 3 + 3H 2 0 + NH 3 . The ammonia unites with another portion of nitric acid, forming ammonium nitrate. Strong hydrochloric acid converts it into stannous chloride, with evolution of hydrogen. Tin is extensively employed in the process of tinning , which consists in coating other metals with a thin film of tin, by dipping into a bath of the molten metal. Ordinary tin-plate (or in common parlance, “tin,” the material of which articles generally called “tins” are made) is thin sheet -iron which has been thus super- ficially coated with tin. Alloys of Tin. — Tin enters into the composition of a large number of useful alloys. With lead, tin will mix in all proportions, and many alloys are in use consisting of these two metals. They are all white, and melt at temperatures lower than that of either constituent. Pewter contains 3 parts of tin to 1 part of lead. Common solder consists of 1 part tin and 1 part lead, while coarse and fine solder contain half, and twice this proportion of tin respectively. With copper, the most important alloys are the various brasses and bronzes. Britannia metal contains tin 84 parts, antimony 10 parts, copper 4 parts, and bismuth 2 parts. Tin is a constituent also of the so-called fusible alloys (see Bismuth, page 461). Oxides of Tin. — Two oxides are definitely known, namely, stannous oxide, SnO, and stannic oxide, Sn0 2 . The monoxide is a base, yielding the stannous salts j the dioxide is both a basic and an acidic oxide. Stannous Oxide, SnO, is obtained by heating stannous oxalate out of contact with air, thus — SnC 2 0 4 = SnO + C0 2 + CO. When sodium carbonate and stannous chloride are mixed, carbon dioxide is evolved, and the white hydrated oxide is precipitated, thus — 2SnCl 2 + 2Na 2 CO s + H g O = 4NaCl + 2C0 2 + 2Sn0,H 2 0. Metastannic Acid 593 When this hydrated oxide is boiled with insufficient caustic alkali to dissolve it, the undissolved portion is dehydrated and converted into the black monoxide. When heated in the air, stannous oxide becomes incandescent, burning to the dioxide. It is soluble in acids, forming stannous salts. The solution of stannous oxide in sodium hydroxide is used by the calico printer, and is known commercially as sodium stannite. Stannic Oxide, Sn0 2 {tin dioxide ), is the chief ore of tin. It is formed where the metal is burnt in the air, but is most readily pre- pared by igniting metastannic acid. It is a white amorphous powder, which changes to yellow and brown on heating, but returns to its original condition on cooling. When strongly heated in a stream of gaseous hydrochloric acid, it may be obtained in small crystals, identical with the natural com- pound. Stannic oxide is unacted upon by acids or alkalies, but in contact with fused potassium hydroxide it is converted into potassium stannate. Stannic Acid, H 2 Sn0 3 , or Sn0 2 ,H 2 0, is obtained in a hydrated condition, as a white gelatinous precipitate, when calcium car- bonate is added to stannic chloride in insufficient quantity for complete precipitation. When the precipitate is dried in vacuo, it has the composition H 2 Sn0 3 . The equation representing its formation may be expressed thus — 2CaC0 3 + SnCl 4 + H 2 0 — 2CaCl 2 + 2C0 2 + H 2 Sn0 3 . Stannic acid forms a number of salts, of which sodium and potassium stannates are the most important ; the former being extensively employed as a mordant in dyeing, under the name of preparing salt. The salts have the composition Na 2 Sn0 3 ,3H 2 0, and K 2 Sn0 3 ,3H 2 0 respectively, and are both soluble in water. Metastannic Acid, H 10 Sn 6 O 16 , is obtained as a white amorphous powder, when tin is acted upon by strong nitric acid ; the reaction may be represented thus — 5Sn + 20HNO 3 = H 10 Sn 6 O 16 + 5H 2 0 + 20NO 2 . The composition of the compound depends upon the particular temperature at which it is dried. This acid is sometimes regarded as a polymer of stannic acid, which may be expressed by the formula 5(H 2 Sn0 3 ) ; metastannic acid, however, appears to be ? P 594 Inorganic Chemistry dibasic, forming salts in which two only of the hydrogen atoms are replaced ; its composition may therefore be conveniently ex- pressed thus — H 2 Sn0 8 ,4Sn0 2 ,4H 2 0, or H 2 Sn 6 O n ,4H 2 0. Potassium and sodium metastannates are the best known salts, their formulae being — K 2 Sn0 3 ,4Sn0 2 ,4H 2 0, and Na 2 Sn0 3 ,4Sn0 2 ,4H 2 0. Stannous Chloride, SnCl 2 , is obtained by dissolving tin in hydrochloric acid, and evaporating the solution, when monosym- metric prisms separate out, having the composition SnCl 2 ,2H 2 0. When dried in vacuo they become anhydrous. The anhydrous chloride is directly obtained when tin filings and mercuric chloride are heated together — HgCl 2 + Sn = SnCl 2 + Hg. The reduced mercury volatilises and leaves the chloride, which at a higher temperature may be distilled. Stannous chloride dissolves in a small quantity of water, but with an excess of water, or on exposure to the air, an oxychloride (or basic chloride) is precipitated with simultaneous elimination of hydrochloric acid, thus — 2SnCl 2 + 2H 2 0 = SnCl 2 ,Sn0,H 2 0 + 2HC1. The composition of this oxychloride may also be expressed by either of the following formulae — Sn 2 OCl 2 ,H 2 0, or 2Sn(OH)Cl, or 2(SnO,HCl). Stannous chloride is a powerful reducing agent, as it readily combines with either oxygen or chlorine ; thus, when added to a solution of mercuric chloride, the latter is first reduced to mer- curous chloride, which, on being gently warmed, is reduced to metallic mercury — 2HgCl 2 4- SnCl 2 = Hg 2 Cl 2 + SnCl 4 . Hg 2 Cl 2 + SnCl 2 = 2Hg + SnCl 4 . By the absorption of oxygen, the above oxychloride and stannic chloride are formed, thus — 3SnCl 2 + O + H a O = SnCl2,SnO,H 2 0 + SnCl 4 . Stannic Sulphide 595 Stannous chloride boils at a temperature about 6 o 6 u . The density of the vapour only agrees with the formula SnCl 2 at tem- peratures above 900 °, at lower temperatures its vapour density approaches more nearly to that required by the formula Sn 2 Cl 4 . Stannic Chloride, SnCl 4 , is obtained by passing a stream of dry chlorine over melted tin in a glass retort ; or by heating a mixture of powdered tin with an excess of mercuric chloride, when the anhydrous chloride distils over as a colourless, mobile, fuming liquid, which boils at 113 . 9 0 . It unites with water with evolution of heat, forming hydrated compounds of the composition SnCl 4 , 3H 2 0 ; SnCl 4 ,5H 2 0, and SnCl 4 ,8H 2 0. The compound containing 5H 2 0 is employed as a mordant, and is commercially known as oxymuriate of tin. Stannic chloride combines with alkaline chlorides forming double chlorides (sometimes called chloro-stannates ), such as SnCl 4 ,2NH 4 Cl, and SnCl 4 ,2KCl. Stannous Sulphide, SnS. — When tinfoil is introduced into sulphur vapour the metal takes fire, and yields a leaden-coloured mass of stannous sulphide. In the hydrated condition, stannous sulphide is precipitated as a brown powder, when sulphuretted hydrogen is passed through stannous chloride ; on drying, this becomes black and anhydrous. Stannous sulphide dissolves in hot concentrated hydrochloric acid. It is also soluble in alkaline polysulphides, forming sulpho- stannates, thus — ( 1 ) 4SnS + K 2 S 6 = K 2 SnS 3 + 3SnS 2 . ( 2 ) SnS 2 + K 2 S = K 2 SnS 3 . On the addition of hydrochloric acid to the solution, stannic sulphide is precipitated — K 2 SnS 3 + 2HC1 = 2KC1 + H 2 S + SnS 2 . Stannic Sulphide, SnS 2 . — This compound cannot be formed by heating tin and sulphur alone, as the heat of the reaction is greater than that at which stannic sulphide is resolved into stannous sulphide and sulphur. It is obtained by heating tin amalgam, sulphur and ammonium chloride, in a retort. The action that takes place is a complicated one, various products being volatilised, and stannic sulphide remaining in the retort as a mass of golden yellow scales. Amongst the products expelled during the process are ammonium chloride, sulphur, mercuric chloride, 596 Inorganic Chemistry mercuric sulphide, and sulphuretted hydrogen. The ammonium chloride present, probably acts by the formation of ammonium stannous chloride, as an intermediate product, which is then de- composed with the production of stannic sulphide and ammonium stannic chloride, thus — 2SnCl 2 ,2NH 4 Cl + 2S = SnS 2 + NH 4 C1 + SnCl 4 ,2NH 4 Cl. Stannic sulphide is a golden yellow crystalline substance, which when heated, partially sublimes as such, but is for the most part decomposed into the monosulphide and free sulphur. It is largely used as a pigment known as mosaic gold. LEAD. Symbol, Pb. Atomic weight = 206.39. Occurrence. — Lead has been found in small quantities in the uncombined state, probably reduced from its ores by volcanic action. In combination with sulphur it occurs in enormous quantities in the mineral galena , PbS, which is the ore from which the metal is chiefly obtained. Large quantities are also met with as carbonate in the mineral cerussite , PbC0 3 . Other natural compounds are anglesite , PbS0 4 ; lanarkite , PbS0 4 ,PbO ; matlockite , PbCl 2 ,PbO ; pyromorphite , 3Pb 3 P 2 0 8 ,PbCl 2 . Modes of Formation. — Lead is very readily reduced from its compounds, and on this account was one of the earliest known metals. It was termed by the Romans plumbum nigrum. Two general processes are made use of for the reduction of lead from its ores : — In the first method (known as the reduction process ), the lead sulphide is reduced by double decomposition with lead oxide and sulphate, which are formed by roasting the ore. In the second (called the precipitation process ), the sulphide is reduced by metallic iron. ( 1 .) The galena is introduced into a reverberatory furnace, where it is partially roasted, whereby a portion of the sulphide is oxidised to sulphate and oxide — PbS + 20 2 = PbS0 4 . 2PbS + 30 2 = 2PbO + 2SO> Lead 597 The temperature is then raised, when the oxide and sulphate react upon a further portion of the sulphide, with the formation of metallic lead and the evolution of sulphur dioxide — PbS0 4 + PbS = 2Pb + 2S0 2 . 2PbO + PbS = 3Pb + S0 2 . This method of lead smelting is followed when the ore is fairly free from other metallic sulphides. The reverberatory furnace usually employed (known as the Flintshire furnace ) has a con- siderable depression, or well, in the hearth, where the metallic Fig. 143. lead collects during the process, and from which it is drawn off into a metal pot. The same process is carried out in the North of England, and in Scotland, where a very pure lead ore is employed, upon open shallow hearths (known as the ore hearth , or Scotch hearth ), built under a brickwork hood or chimney in such a manner that the fumes of lead which escape, are caused to pass into condensing chambers. Fig. 143 shows such a hearth in section. The fire of peat and coal is urged by a small blast admitted from behind, and the ore is added in small quantities at a time. The reduced metal sinking to the bottom, runs under the fire-bar, and overflows from the shallow hearth down a channel upon an inclined stone 59 ? Inorganic Chemistry surfaces (called the work-stone) into an iron pot P, which is gently heated by a small fire to enable the operator to ladle the metal out into moulds. (2.) This method of lead smelting depends upon the fact, that at a high temperature metallic iron, in contact with lead sulphide, is converted into ferrous sulphide, with separation of lead — PbS + Fe = FeS + Pb. The ores (either in the raw state, or after previous calcination) are smelted in a blast-furnace with coke and either metallic iron, or such materials as will yield iron under the furnace conditions. The sulphide of iron, along with other metallic sulphides, rises to the top of the molten lead as a matt or regulus, while above this a fusible slag collects, consisting chiefly of silicate of iron. The lead first obtained by any of these processes usually con- tains antimony, tin, copper, and other metals. These impurities are removed by heating the metal in a shallow, flat-bottomed reverberatory furnace. Most of the admixed metals oxidise before the lead, and collect in the dross which forms upon the surface. This process is known as the softening of lead. The silver, how- ever, which is always present, is not removed by this operation, but is extracted by one of the methods for desilverising lead described under silver, page 516. Properties. — Lead is a soft, bluish-white metal, which when freshly cut exhibits a bright metallic lustre. On exposure to the air its bright surface becomes quickly covered with a film of oxide. Lead is sufficiently soft to be scratched with the finger nail, and it leaves a black streak when drawn across paper. It cannot be hammered into foil, or drawn into wire, but may readily be obtained in these forms by rolling and pressing. When a quantity of melted lead is allowed partially to resolidify, and the still liquid portion poured off, the metal is obtained in the form of octahedral crystals belonging to the regular system. Its crystalline nature is also readily seen by submitting a solution of a lead salt to electro- lysis, when the metal is deposited upon the negative electrode in beautiful arborescent crystals with a brilliant metallic lustre (Fig. 144). It is deposited in a similar form, known as the lead tree , by suspending a strip of zinc in such a solution. The specific gravity of lead is 11.3 ; it melts at 330° to 335 0 , and becomes covered with a black film of the suboxide, Pb 2 0 : when more strongly heated it is oxidised to the monoxide, PbO. Lead 599 Lead is rapidly dissolved by nitric acid, but hydrochloric and sulphuric acids are almost without action upon it in the cold. Hot concentrated hydrochloric acid, however, slowly converts it into lead chloride. Lead is unacted upon by pure water, in the absence of air ; but in contact with air, lead hydroxide is formed, which is slightly soluble in water. By the action of atmospheric carbon dioxide upon this solution, a basic carbonate is precipitated, having the com- position 2PbC0 3 ,Pb(HO) 2 . The solvent action of water upon lead is greatly influenced by the presence of various dissolved sub- Fig. 144. stances in the water ; thus, water containing small quantities of ammoniacal salts, notably the nitrate, dissolves lead much more rapidly, and the same is the case with water charged with carbon dioxide under pressure. In the latter case the action is probably due to the formation of a soluble acid carbonate. Water, on the other hand, containing small quantities of phos- phates and carbonates, especially the acid calcium carbonate, are almost entirely without action upon lead. Certain drinking waters (such as the Loch Katrine water), which on account of their purity 6oo Inorganic Chemistry exert a solvent action upon the lead pipes through which they are conveyed, are rendered incapable of acting upon the lead by being first filtered through chalk or animal charcoal, which enables them to take up sufficient calcium carbonate or phosphate to prevent this action. On account of the exhaustive methods of desilverisation to which the lead is subjected, commercial lead possesses a degree of purity not found in any other metal as commonly met with ; the total amount of foreign metals present in ordinary commercial lead, ranges from o. i to 0.006 per cent. Lead is put to a large number of uses in the arts, on account of the ease with which it can be worked, and its power of resisting the action of water and many acids. In the manufacture of lead pipes advantage is taken of the extreme softness of the metal, and the readiness with which it can be pressed into shape ; the lead, in a pasty or semi-molten condition, being merely squeezed, or squirted , through a steel die, by hydraulic pressure. Lead bullets are also made by squeezing the metal into moulds ; for as lead contracts on solidification, bullets made by casting always contain a small cavity, which (unless it happens to form exactly at the point of centre of gravity) renders the flight of the bullet untrue. Oxides of Lead. — Five oxides of lead are known, having the composition Pb 2 0, PbO, Pb 2 0 3 , Pb 3 0 4 , Pb0 2 . Lead Suboxide [plumbous oxide), Pb 2 0, is the black compound which is formed when lead is heated to its melting-point. It is obtained by heating plumbic oxalate to about 300 ° in a glass tube or retort — 2PbC 2 0 4 = CO + 3C0 2 + Pb 2 0. When heated in the air it bums, forming plumbic oxide ; in the absence of air it is decomposed into the same oxide and metallic lead, the reactions being — Pb 2 0 + 0 = 2 PbO. Pb 2 0 = Pb + PbO. In contact with acids it decomposes in the same manner, lead being deposited, and the plumbic oxide dissolving in the acid to form a plumbic salt. Plumbic Oxide ( lead monoxide , litharge , massicot ), PbO, is formed when lead is strongly heated in the air, and is obtained in Red Lead 601 large quantities in the cupellation of argentiferous lead. It may be obtained by heating lead nitrate or carbonate, and it is produced when any of the other oxides are heated. Plumbic oxide is a yellowish powder, known commercially as massicot , which, when melted and resolidified, is obtained as a crystalline mass, known as litharge. Plumbic oxide is very slightly soluble in water, i part dissolving in 7000 parts of water : this solution is alkaline, and on exposure to the air absorbs carbon dioxide, forming an insoluble basic carbonate. Plumbic oxide is dissolved by acids, with formation of the salts of lead ; it also dissolves in warm potassium, or sodium hydroxide. This oxide forms two hydrated compounds, having the com- position 2Pb0,H 2 0 and 3Pb0,H 2 0. The former is obtained as a white precipitate when ammonia is added to a solution of lead acetate ; the second, by the action of ammonia on basic lead acetate at 25 0 . Lead Sesquioxide, Pb 2 0 3 , is obtained as an orange-coloured precipitate by adding sodium hypochlorite to a solution of plumbic oxide in potassium hydroxide. Heat decomposes it into oxygen and plumbic oxide. Acids convert it into the monoxide and dioxide, the former dissolving and yielding a salt of lead. This oxide may be regarded as a compound of two oxides, PbO,Pb0 2 . Triplumbie Tetroxide ( red lead , , minium ), Pb 3 0 4 , is obtained when lead carbonate, or monoxide, is subjected to prolonged heating in contact with air, at a temperature not above 450 °. At higher temperatures it again gives up oxygen. It is a scarlet crystalline powder, varying somewhat in colour, according to its mode of preparation. Dilute acids convert it into Pb0 2 and 2 PbO, the latter oxide dissolving to yield lead salts. With strong hydrochloric acid and sulphuric acid, the molecule of lead dioxide is acted upon with evolution of chlorine and oxygen respectively — Pb 3 0 4 + 8HC1 = 4H 2 0 + 3PbCl 2 + Cl 2 . Pb 3 0 4 + 3H 2 S0 4 = 3H 2 0 + 3PbS0 4 + O. Red lead* is employed as a pigment, and also in the manu- facture of flint glass. * Commercial red lead varies considerably in composition, and although it has been shown that a definite compound exists, of the composition Pb 3 0 4 (which may also be expressed by the formula 2Pb0,Pb0 2 ), it is still uncer- tain whether there are not other compounds consisting of these two oxides united in different proportions. 602 Inorganic Chemistry Plumbic Peroxide {lead dioxide ), Pb 0 2 , may be obtained by the action of dilute nitric acid upon red lead — Pb 3 0 4 (or Pb0 2 ,2Pb0) + 4HN0 3 = Pb0 2 + 2Pb(N0 3 ) 2 + 2H 2 0. Or it may be prepared by the action of oxidising agents upon the monoxide. Thus, when chlorine is passed through an alkaline solution, in which the monoxide is suspended, or when bleaching- powder is added to a solution of lead acetate, the dioxide is . produced. The dark brown deposit which forms upon the positive electrode when a solution of a lead salt is electrolysed, also consists of the dioxide. Plumbic peroxide is a brown or puce-coloured powder. It is a powerful oxidising substance, and when gently rubbed with flowers of sulphur in a warm mortar the mass suddenly inflames. When a stream of sulphur dioxide is passed over the peroxide in a tube, the two compounds unite to form lead sulphate, the mass becom- ing incandescent. Nitric acid is without action upon it, but hydrochloric and sulphuric acids act upon it in the same manner as upon red lead. When strongly heated, the peroxide gives up oxygen, and is converted into the monoxide. When plumbic peroxide is boiled with strong aqueous potassium hydroxide it dissolves, and the solution deposits crystals of potas- sium plumbate, K 2 Pb0 3 ,3H 2 0. This compound corresponds with potassium stannate, K 2 Sn0 3 ,3H 2 0, and its existence shows that lead possesses, although to a very feeble extent, the acidic properties exhibited by the other members of the same family of elements. Plumbic Chloride {lead dichloride ), PbCl^ is obtained as a white curdy precipitate, when hydrochloric acid, or a soluble chloride, is added to a solution of a lead salt. It is also produced by the action of boiling hydrochloric acid upon lead in the pre- sence of air. It is best prepared by dissolving lead oxide or carbonate in hot hydrochloric acid, when the lead chloride sepa- rates out on cooling, in long white, lustrous, needle-shaped crystals belonging to the rhombic system. Lead chloride is soluble in boiling water to the extent of about 4 parts in 100 parts of water. On cooling the solution, the greater part of the salt separates out, and at o° the liquid contains 0.8 parts in solution. The presence of hydrochloric acid and soluble chlorides diminishes the solu- bility of lead chloride. Lead Nitrate 603 When heated in contact with air, it is converted into an oxy- chloride, of the composition Pb 2 OCl 2 , or PbCl 2 ,PbO, corresponding with the natural compound inatlockite. This compound, in the hydrated condition, Pb 2 0Cl 2 ,H 2 0, is prepared on a large scale by the addition of lime-water to a solution of lead chloride, and is employed as a white pigment, known as Pattinsoris white lead. Cassel yellow is an oxychloride of lead of the composition PbCl 2 ,7PbO, obtained by heating lead oxide and ammonium chloride. Lead Tetrachloride {lead perchloride), PbCl 4 .— When plumbic peroxide is dissolved in cold concentrated hydrochloric acid, a yellow liquid is obtained, which, on warming, yields chlorine, with precipitation of lead dichloride. This liquid contains the tetrachloride of lead in solution. When lead dichloride is suspended in hydrochloric acid, and chlorine is passed through the mixture, a solution of lead tetra- chloride is obtained ; and on the addition of ammonium chloride, ammonium plumbic chloride, PbCl 4 ,2NH 4 Cl (corresponding to ammonium stannic chloride), separates out. When this compound is acted upon with strong sulphuric acid, in the cold, lead tetra- chloride separates out as a yellow oily liquid. Lead tetrachloride is a yellow, highly-refracting, fuming liquid, which decomposes in contact with moisture into lead dichloride and chloride. It may be preserved beneath concentrated sul- phuric acid. With small quantities of water, it forms a hydrated compound, but excess of water decomposes it into hydrochloric acid and lead peroxide — PbCl 4 + 2H 2 0 = Pb0 2 + 4HC1. When heated with strong sulphuric acid to about 105 °, it suddenly decomposes with explosion. Lead Nitrate, Pb(N0 3 ) 2 , is obtained by dissolving litharge in nitric acid. The salt is deposited from the solution in the form of regular octahedral crystals. It is soluble in water to the extent of 50 parts in 100 parts of water, at the ordinary temperature. When heated, it evolves nitrogen peroxide and oxygen, leaving plumbic oxide (page 217 ). On boiling an aqueous solution of lead nitrate with lead oxide, the latter dissolves, and the solution on cooling, deposits crystals of a basic nitrate, Pb(N0 3 )H0 or Pb(N0 3 ) 2 ,Pb0,H 2 0. By the addition of ammonia to a solution of lead nitrate, other basic 6°4 / norganic Chemistry nitrates are obtained, which may be regarded as consisting of compounds of Pb(N0 3 )HO with PbO, or of Pb(N0 3 ) 2 with PbO and H 2 0 in varying proportions. Lead Carbonate, PbC0 3 , is obtained as a white crystalline powder, by the addition of ammonium sesquicarbonate to a solu- tion of lead nitrate. It occurs in the form of transparent rhombic crystals in the mineral cerussite , isomorphous with arragonite. Lead carbonate is almost insoluble in water, but is appreciably dissolved in water charged with carbon dioxide. When sodium or potassium carbonate is added to a solution of lead nitrate, basic carbonates of lead are precipitated, varying in composition with the conditions of temperature. The most important of the basic carbonates is white lead , 2PbC0 3 ,Pb(HO) 2 . This compound is manufactured on a large scale by several processes, for use as a pigment. The oldest process, and that which yields the best product, is known as the Dutch method. It depends upon the action of acetic acid upon metallic lead, in the presence of moist air and carbon dioxide. The lead, cast into rough gratings in order to expose a large surface, is placed in earthen- ware pots, as shown in Fig. 145. A small quantity of dilute acetic acid (in the old Dutch process, vinegar) is placed in the pots, and the gratings of lead, which rest upon the shoulder of the pot, are piled one upon the other. These pots are then placed upon a thick bed of spent tan- bark (in the original method, dung), upon the floor of a shed, and covered with planks. Upon these another layer of tan-bark is spread, and a second row of pots similarly charged. In this manner the layers of pots are built up to the roof of the shed, and the whole allowed to remain for about three months. Such a stack will contain many tons of lead, and about 65 gallons of dilute acetic acid to the ton of metal. The acid is gradually vaporised by the heat developed by the fermenting tan-bark, which results first in the formation of a basic lead acetate— Fig. 145. 2H(C 2 H 3 0 2 ) + 2Pb + 0 2 = Pb(C 2 H 3 0 2 ) 2 ,Pb(H0)2. Lead Sulphate 605 This basic acetate is then acted upon by the carbon dioxide evolved during the fermentation, with the production of a mixture of normal lead acetate, and basic lead carbonate, thus— 3{Pb(C 2 H 3 0 2 ) 2 ,Pb(HO) 2 } +2C0 2 =3Pb(C 2 H 3 0 2 ) 2 + 2PbC0 3 ,Pb(H0) a + 2H 2 0. And the lead acetate, in the presence of air and moisture, reacts upon a further portion of the metal, regenerating the basic acetate, which is once more decomposed by carbon dioxide — Pb(C 2 H 3 0 2 )2 + Pb + O 4- H 2 0 = {Pb(C 2 H 3 0 2 ) 2 ,Pb(H0) 2 }. In this cycle of reactions, therefore, the acetic acid acts as a carrier, a comparatively small quantity being able to convert an indefinite amount of lead into white lead. White lead is also prepared by passing carbon dioxide into a solution of the basic acetate, obtained by boiling plumbic oxide (litharge) with lead acetate. The product, however, is not so opaque as that obtained by the former method, and is therefore not so valuable as a pigment. (This method is known as the Clichy , or Thenard’s process.) Milner 3 s process consists in grinding together litharge, sodium chloride, and water, whereby a mixture of an oxychloride of lead and sodium hydroxide is formed — 4PbO + 2NaCl + 5H 2 0 = PbCl 2 ,3Pb0,4H 2 0 + 2NaHO, and then passing carbon dioxide into the mixture, which converts it into white lead and sodium chloride, thus — 3[PbCl 2 ,3Pb0,4H 2 0] + 6NaHO + 8C0 2 = 6NaCl + 4[2PbC0 3 ,Pb(H0) 2 ] + 11H 2 0. White lead is a heavy, amorphous powder, whose value as a pigment, or body colour, depends upon its opacity and density. Although this compound labours under the disadvantages of being extremely poisonous, and of becoming blackened by sulphuretted hydrogen, no substitute for it has yet been found which possesses the same “ body ” or covering power. Lead Sulphate, PbS0 4 . — The mineral anglesite, PbS0 4 , occurs in the form of rhombic crystals, isomorphous with strontium and barium sulphates. Lead sulphate is obtained as a white powder, by precipitating a lead salt with sulphuric acid, or a soluble 6 o 6 Inorganic Chemistry sulphate. It is soluble in water only to an extremely slight extent, and still less in dilute sulphuric acid, but strong sulphuric acid dissolves it readily. It also dissolves in potassium hydroxide, and in many ammoniacal salts, notably the acetate, and in sodium thiosulphate. An acid sulphate, of the composition PbS 0 4 ,H 2 S 04 ,H 2 0 is obtained by boiling the normal sulphate with sulphuric acid ; and a basic sulphate, PbS0 4 ,PbO, is formed by the action of ammonia upon the normal salt. Lead Sulphide, PbS. — The natural sulphide, galena , is found in the form of cubical crystals, possessing very much the colour and the metallic lustre of freshly cut lead. It is artificially formed when lead is heated in sulphur vapour, or when sulphuretted hydrogen is passed through a solution of a lead salt. When heated in vacuo, or in a stream of an inert gas, lead sulphide melts, and sublimes in the form of small cubes.* When heated with free access of air it is converted into lead sulphate. Boiling dilute nitric acid converts lead sulphide into the nitrate, with separation of sulphur ; but strong nitric acid oxidises it into lead sulphate. It is decomposed by hot concentrated hydrochloric acid, with evolution of sulphuretted hydrogen. When sulphuretted hydrogen is passed into a solution of lead chloride, the precipitate which forms is first yellow, then reddish- brown, and finally black ; the yellow and red precipitates are com- pounds of lead chloride and lead sulphide, termed sulphochlorides. having the composition, PbS,PbCl 2 , and 3PbS,PbCl 2 . The compounds of lead are powerful poisons, and when con- tinuously taken into the system in small quantities, they act as cumulative poisons. Painters and others who constantly handle white lead, are liable to suffer from chronic lead poisoning. From the readiness with which lead sulphide volatilises when heated in a stream of sulphur dioxide, Hannay concludes ( Proc . Chem. Soc., May 1894), that the two substances unite to form a volatile compound, PbS,S0 2 , or PbS 2 O a , which when its temperature falls, again breaks up into lead sulphide and sulphur dioxide. When sulphur dioxide is generated in intimate contact with galena, as when a stream of air is passed through the molten substance, it is observed that one half of the lead is reduced, and one half volatilises with the gas, according to the equation 2PbS + 0 2 = Pb + PbS,S0 2 . No definite proof, however, has yet been adduced as to the existence of this volatile compound. CHAPTER X ELEMENTS OF GROUP V. (. FAMILY A.) Vanadium , V = 51.1 ; Niobium , Nb = 93.7 ; Tantalum , Ta = 182. The three rare metals comprising this family are closely related to each other, and also to the elements of family B of the same group, namely, the nitrogen and phosphorus series. Vanadium occurs in a few rare minerals, as vanadite, 3 Pb 3 ( V 0 4 ) 2 , PbCl 2 (the vanadium analogue of pyromorphite) ; pucherite, BiV 0 4 ; mottramite, (PbCu) 3 (V0 4 ) 2 ,2(PbCu)(H0)2. Small quantities also occur in certain iron ores, the vanadium ultimately finding its way into the Bessemer slag, in which it has been found concentrated to the extent of 1.5 per cent. Metallic vanadium was first isolated by Roscoe (1867), although its existence was previously discovered by Del Rio (1801). The metal is extremely difficult to obtain, as at a red heat it combines with oxygen with great readiness, yielding the pentoxide V 2 0 6 , and also with nitrogen, forming the nitride VN. The element is prepared by heating the dichloride in a stream of perfectly pure hydrogen — VC 1 2 + H 2 = 2 HC 1 + V. Vanadium is unacted upon by air at ordinary temperatures, but when heated burns brilliantly to the pentoxide. Niobium and tantalum are found associated together in the rare mineral tanialite or columbite. The first to be discovered was tantalum (Hatchett, 1801), and was originally named columbium ; and the name niobium (from Niobe, the daughter of Tantalus) was given to the allied element, by Rose (1846). Niobium is obtained by heating the trichloride, NbCl 3 , in a stream of hydrogen. Vanadium forms five oxides, corresponding to the oxides of nitrogen, while three oxides of niobium and two of tantalum are known V a O ; V 2 0 2 (or VO) ; V 9 0 3 ; V 3 0 4 (or V 0 2 ) ; V 2 O s . — ; NbO ; — ; Nb 0 2 ; Nb 2 0 5 . — ; — ; — ; Ta 0 2 ; Ta 2 O s . The pentoxides are obtained when the metals are burned in air or oxygen. They give rise respectively to vanadates, niobates, and tantalates, correspond ing to nitrates and metaphosphates, thus — Sodium nitrate, NaN 0 3 . Sodium metaniobate, NaNbO s . Sodium metaphosphate, NaP 0 3 . Sodium metatan talate, NaTa 0 3 . Sodium metavanadate, NaV 0 3 . 607 6o8 Inorganic Chemistry 1 he closer relation of these elements to phosphorus than to nitrogen, is seen in the formation of salts derived from ortho- and pyro-acids, corresponding to orthophosphates and pyrophosphates. The naturally occurring vanadium compounds above mentioned are vanadates derived from the hypothetical orthovanadic acid, H 3 V0 4 . Both metavanadic acid, HV0 3 , and pyrovanadic acid, H 4 V 2 0 7 , have been obtained. Unlike the phosphorus compounds, the metavanadates are the most stable of the three classes of salts, and the orthovanadates the least stable. The most important of these salts is the ammonium metavanadate, NH 4 V0 3 , which is prepared by dissolving the pentoxidc in ammonia. This salt is insoluble in ammonium chloride, and use is made of this property in the preparation of vanadium compounds from the mineral mottramite. When ammonium metavanadate is ignited vanadium pentoxide is obtained — 2NH 4 V0 3 = V 2 0 5 + 2NH 3 + H 2 0. Vanadium acts also as a feeble base. Thus, when the tetroxide, or hypo- vanadic oxide, is dissolved in sulphuric acid, hypovanadic sulphate, V 2 0 2 (S0 4 ) 2 , is formed. The solution of this salt possesses a rich blue colour. Vanadium forms three chlorides, having the composition— VC1 2 (or V 2 C1 4 ) ; VC1 3 (or V 2 C1„) ; VC1 4 . Niobium gives a trichloride, NbCl 3 , and pentachloride, NbCl 5 , while only the pentachloride of tantalum is known, TaCl 5 . Vanadium forms a number of compounds with oxygen and chlorine. Thus, when vanadium tetrachloride is acted upon by water, it yields hypovanadic chloride, V 2 0 4 C1 2 , which dissolves in the water, giving a blue solution Vanadium oxychloride, or vanadyl trichloride, VOCl 3 , corresponds to phos- phorus oxychloride, POCl 3 . From vanadyl trichloride, by treatment with zinc, vanadyl dichloride is obtained, VOCl 2 , and by the action of hydrogen at a high temperature upon this, both vanadyl monochloride, VOC1, and divanadyl, monochloride, VoOgCl, are formed. CHAPTER XI ELEMENTS OF GROUP VI. (. FAMILY A.) Chromium, Cr . . . . 52 I Tungsten , W. . . 184 Molybdenum , Mo . . 95* 9 I Uranium, U . - . 239.8 CHROMIUM. Symbol, Cr. Atomic weight = 52. Occurrence. — Chromium does not occur in nature in the un- combined state. In combination with oxygen and associated with iron, it is met with in considerable quantities in the mineral chrome iron ore , or chromite , Cr 2 0 3 ,Fe0. This ore is the chief source of chromium compounds. Other natural compounds are crocoisite , PbCr0 4 , and chrome-ochre , Cr 2 0 3 . Traces of chromium are present in various minerals, such as the emerald and green serpentine, and impart to them their green colour. Modes of Formation. — Although chromium compounds are manufactured for industrial purposes, the element itself has re- ceived no technical application. It was obtained by Wohler, by the reduction of fused chromium chloride with metallic zinc, beneath a layer of fused sodium and potassium chlorides. The regulus, or alloy of zinc and chromium, was then treated with dilute nitric acid, whereby the zinc was dissolved, and the chromium was obtained in the form of a powder. Chromium may also be obtained as bright metallic scales, by the electrolysis of a solution of chromous chloride containing chromic oxide. The metal may be prepared by the reduction of the oxide, Cr 2 0 3 , by means of carbon, at a high temperature ; or by heating the oxide with metallic aluminium. Properties. — Chromium is a hard, steel-grey metal, which is not oxidised in dry air. When heated in the oxyhydrogen flame it burns brilliantly. It dissolves in hydrochloric acid with evolu- tion of hydrogen. The metal is not magnetic. The presence of 609 2 Q 6io Inorganic Chemistry minute quantities of chromium in steel, imparts to the latter great hardness and tenacity. Oxides Of Chromium. — Two oxides of chromium are definitely known, namely — Chromium sesquioxide ( chro?nic oxide) . . Cr 2 0 3 . Chromium trioxide {chromium anhydride) . . CrQ 3 . The first is a basic, and the second an acidic oxide. Besides these two compounds, a hydrated oxide, derived from the unknown chromous oxide, also exists, having the composition CrO,H 2 0, or Cr(H 0) 2 . It is obtained as a yellowish precipitate by adding potas- sium hydroxide to a solution of chromium dichloride (chromous chloride), with the exclusion of air. It rapidly absorbs oxygen, turning dark brown. When heated out of contact with air it is converted into the sesquioxide, with evolution of hydrogen — 2Cr0,H 2 0 = Cr 2 0 3 + H 2 0 + H 2 . Other compounds of chromium and oxygen are described, whose composi- tion, however, is not definitely established ; thus, the product obtained as a brown powder, either by the partial reduction of the trioxide, or the oxidation of the sesquioxide, is regarded by some chemists as chromium dioxide, Cr0 2 , and by others as chromium chromate, Cr 2 03 ,Cr 0 3 . It is readily obtained by passing nitric oxide into a solution of potassium dichromate. Chromium Sesquioxide, Cr 2 0 3 , is obtained as a grey-green powder, when either the hydroxide, or the trioxide, or ammonium dichromate is ignited (see page 206 ). When the vapour of chromyl dichloride, CrC^CLj, is passed through a red-hot tube, chromic oxide is deposited in the form of dark-green hexagonal crystals. Chromic oxide which has been strongly ignited, is nearly soluble in acids. It is used under the name of chrome green as a pigment, and for giving a green colour to glass. Chromic Hydroxides. — Chromic oxide yields a number of hydrated compounds. When ammonia is added to a solution of chromic chloride, or other chromic salt, free from alkali, a light- blue compound is precipitated, which, when dried over sulphuric acid, has the composition Cr 2 (H0) 6 ,4H 2 0 (or Cr 2 0 3 ,7H 2 0). When this is dried in vacuo, it loses 3H 2 0, and becomes Cr 2 (H0) 6 ,H 2 0 (or Cr 2 0 3 ,4H 2 0) ; and on being heated at 200 °, it again parts with 3H 2 0, and has the composition Cr 3 0 3 ,H 2 0. Chromium Trioxide 61 When potassium dichromate and boric acid are heated to dull redness, and the mass treated with water, a rich green residue is obtained, having the composition Cr 2 0 3 ,2H 2 0. This compound, known as Guignefs green , is employed as a pigment. The first two of these compounds, which may be looked upon as consisting of the hydroxide Cr 2 (HO ) 6 in a hydrated condition, namely, Cr 2 (H0) 6 ,4H 2 0 and Cr 2 (H0) 6 ,H 2 0, are readily soluble in acids, yielding the chromic salts. Chromium Trioxide (< chromic anhydride ), Cr0 3 . — When strong sulphuric acid is added to a cold saturated solution of potassium dichromate, the trioxide separates out in long, red, needle-shaped crystals — K 2 Cr 2 0 7 + H 2 S0 4 = K 2 S0 4 + H 2 0 + 2CrO s . The liquid is decanted from the crystals, which are drained upon porous tiles, and the adhering sulphuric acid and potassium sulphate washed away by strong nitric acid. The crystals are finally heated upon a sand-bath, whereby the nitric acid is evaporated. Chromium trioxide dissolves in water to the extent of 62 parts in 100 parts of water at 26 °. It melts at a temperature about 192 0 . At 250 ° it begins to give off oxygen, and is ultimately converted into the sesquioxide — 2Cr0 3 — Cr 2 0 3 + 30. Chromium trioxide is a powerful oxidising agent, and in contact with most organic substances it is reduced. In the preparation of the compound, therefore, the liquid cannot be filtered through paper in the usual way. Warm alcohol dropped upon the trioxide at once takes fire, while in a more diluted condition it is oxidised to acetic acid ; and the reduction of the chromium trioxide is made evident by the change of colour of the liquid, from red or yellow, to olive green. Gaseous ammonia reduces the trioxide to the sesquioxide, with formation of water and nitrogen — 2NH 3 + 2CrO s = Cr 2 0 3 + N 2 + 3H 2 0, the reaction being accompanied with the evolution of so much heat that the chromic oxide produced becomes incandescent. When hydrogen peroxide is added to a dilute solution of 612 Inorganic Chemistry chromium trioxide, or to a dilute solution of potassium dichromate, acidified with sulphuric acid, a deep indigo-blue solution is ob- tained. This blue compound is believed to be perchromic acid ’ but its composition has not been definitely established. It may be regarded as a compound of chromium trioxide, Cr 0 3 , with hydrogen peroxide, H 2 C>2, in undetermined proportions. In aqueous solution, the blue colour quickly disappears, oxygen being eliminated. The compound is soluble in ether ; and, there- fore, when the aqueous solution is shaken up with that liquid, a deep blue ethereal solution rises to the top. In this solution the compound is more stable, but when evaporated it evolves oxygen, leaving chromium trioxide. It is decomposed by alkalies, forming alkaline chromates, with evolution of oxygen. The formation of this compound constitutes a delicate test for either chromium trioxide or hydrogen peroxide (see Hydrogen Peroxide, page 203). Chromous Compounds. — These correspond to chromous hydrate, Cr(HO) 2 , in which the chromium functions as a divalent element. Comparatively few of these salts are known. Chromous Chloride, CrCl 2 , is formed when the metal dissolves in hydro- chloric acid. It is prepared in the anhydrous state by gently heating chromic chloride in a current of pure hydrogen. It is a white crystalline compound, soluble in water to a blue solution, which rapidly absorbs oxygen. Chromous Sulphate, CrS0 4 ,7H 2 0, is obtained by dissolving chromous acetate in dilute sulphuric acid. It is deposited from the solution in blue crystals, isomorphous with ferrous sulphate, FeS0 4 ,7H 2 0. Chromic Compounds. — These are derived from chromic oxide, the oxide acting as a base. Chromic Chloride, CrCl 3 , or Cr 2 Cl 6 , is prepared by strongly heating a mixture of chromic oxide, Cr 2 0 3 , and carbon, in a stream of dry chlorine. The chromic chloride sublimes in the form of scales, having a reddish-pink colour. The molecular weight of chromic chloride is 158.6, showing that in the vaporous state its molecules have the formula CrCl 3 . It is nearly insoluble in water, but readily dissolves in water containing minute traces of chromous chloride, forming a green solution. The same solution is obtained by dissolving hydrated chromic hydroxide, Cr 2 (HO) 6 , 4 H 2 0 , in hydrochloric acid, and if this solution be slowly evaporated, very soluble green crystals separate out having the composition Cr 2 Cl 6 , 12 H 2 0 . If strongly heated in the air, this compound gives off water and hydrochloric acid, leaving chromic oxide, Cr 2 O a ; but when heated to 250°, in Chrome Alum 613 either gaseous hydrochloric acid or chlorine, it is converted into the pink anhydrous chromic chloride, which redissolves in water to the green solution. If heated strongly and sublimed, the com- pound obtained is nearly insoluble in water. Chromic Sulphate, Cr 2 (S 0 4 ) 3 , is obtained by dissolving chro- mium hydroxide in concentrated sulphuric acid, when a green solution is formed, which on standing changes to blue, and slowly deposits violet-blue crystals. The salt may be purified by dis- solving in cold water and precipitating with alcohol. If insufficient alcohol be added to cause immediate precipitation, the salt slowly deposits from the dilute spirit in blue octahedrons, belonging to the regular system. A cold aqueous solution, which has a violet colour, becomes green when boiled. Chromic sulphate forms double salts with the sulphates of the alkalies, which belong to the alums. Potassium Chromium Alum ( chrome alum), K 2 S0 4 ,Cr 2 (S0 4 ) 3 , 24H 2 0. — This double sulphate is formed when solutions of potas^ sium and chromium sulphates are mixed together in molecular proportions. It is most conveniently prepared, by the addition of the requisite amount of sulphuric acid to an aqueous solution of potassium dichromate, and reducing the chromic oxide by passing sulphur dioxide through the liquid — ( 1 ) K 2 Cr 2 O r + H 2 S0 4 = 2CrO s + H 2 0 + K 2 S0 4 . ( 2 ) 2CrO s + 3S0 2 = Cr 2 (S0 4 ) 3 . The resulting solution, containing the two sulphates in mole- cular proportions, deposits crystals of the double sulphate, in the form of dark plum-coloured octahedrons (Fig. 140 , B, p. 573 ), which appear red by transmitted light. Chrome alum dissolves in water, yielding a plum-coloured solu- tion, which on boiling turns green, but on long standing returns to its original colour. Sodium chromium alum is more soluble, and ammonium chro- mium alum is less soluble, than the potassium salt. Chromites.— Chromic oxide acts also as a weak base, and combines with other oxides, forming compounds resembling the aluminates. When potas- sium hydroxide is added to a solution of a chromic salt, the green hydrated oxide which is precipitated contains alkali which cannot be removed by hot water ; this is present in the form of potassium chromite. The best known 614 Inorganic Chemistry chromites are zinc chromite, Cr 2 0 3 ,Zn0; manganous chromite, Cr 3 0 3 ,Mn0, and ferrous chromite, Cr 2 0 3 ,FeO ; the latter occurs naturally as chrome iron ore. Chromates. — When chromium trioxide is dissolved in water, the solution is believed to contain chromic acid, H 2 Cr0 4 ; when the solution is evaporated, however, the trioxide alone is left. (Red crystals have been obtained, by cooling a hot saturated solution of the trioxide, which have been regarded as the acid.) Potassium Chromate, K 2 Cr0 4 , is prepared by adding potas- sium hydroxide to a solution of the dichromate — K 2 Cr 2 0 7 + 2KHO = 2K 2 Cr0 4 + H 2 0. On evaporation, the yellow chromate of potash separates out, in rhombic crystals, isomorphous with potassium sulphate. It is soluble in water at the ordinary temperature to the extent of 60 parts in ioo parts of water, forming a yellow solution having an alkaline reaction. Potassium Diehromate, K 2 Cr 2 0 7 , is manufactured from chrome iron ore by roasting the finely crushed ore with potassium car- bonate and lime in a reverberatory furnace ; the mass being frequently raked over to expose fresh portions to the oxidising action of the flames. In this way a mixture of calcium and potas* sium chromates is produced — 2Cr 2 0 3 1 F eO + 3 K 2 C0 3 -f* Ca.O + 70 = Ca.Cr0 4 + 3K 2 Cr0 4 + F ej0 3 + 3C0 2 . The yellow mass, when cold, is broken up and lixiviated with a hot solution of potassium sulphate, which by double decomposition with the calcium chromate, forms potassium chromate and precipi- tates calcium sulphate. The solution after settling, is treated with the requisite quantity of sulphuric acid to convert the chromate into the dichromate, thus — 2K 2 Cr0 4 + H 2 S0 4 = K 2 S0 4 + H 2 0 + K 2 Cr 2 O r . The dichromate being much less soluble than the normal chro- mate, a large proportion of it at once deposits as the solution cools ; and the mother liquor containing potassium sulphate is used again to lixiviate a fresh quantity of the roasted mixture. Potassium dichromate forms large red prisms or tables, belong- ing to the asymmetric (triclinic) system. It is soluble in water at Chromyl Chloride 615 the ordinary temperature to the extent of 10 parts in 100 parts of water, yielding an acid solution, which is extremely poisonous. When a film of gelatine is impregnated with potassium dichromate and exposed to light, a reduction of the chromium to chromic oxide takes place, and at the same time the gelatine is rendered insoluble. This property is utilised in photographic processes* Potassium dichromate is also known under the misnomer bichromate of potash , which would suggest that the salt was in reality hydrogen potassium chromate, corresponding to bisulphate of potash, HKS0 4 . Such a chromium compound does not exist. The dichromates correspond to the disulphates (or pyrosulphates), see page 396. Potassium Trichromate, K 2 Cr 3 O 10 (or K 2 Cr0 4 ,2Cr0 3 ), and Potassium Tetrachr ornate, K 2 Cr 4 0 13 (or K 2 Cr0 4 ,3Cr0 3 ), are also known. Lead Chromate, PbCr0 4 , is found as the mineral crocoisite. It is produced by precipitation from a lead salt, with either potassium chromate or dichromate. It forms a bright yellow powder, known as chrome-yellow , and is employed as a pigment. It melts without decomposition, and resolidifies on cooling to a brown crystalline solid. At higher temperatures it gives off oxygen, and is converted into chromic oxide and a basic lead chromate. When heated with organic compounds, the latter are completely oxidised ; lead chromate is therefore employed in organic analyses. When lead chromate is digested with sodium hydroxide, or with normal potassium chromate, a basic lead chromate is obtained as a rich red powder — 2PbCr0 4 + 2NaHO = Na 2 Cr0 4 + H 2 0 + Pb 2 Cr0 6 . This compound is known as chrome-red. Chromyl Chloride, Cr0 2 Cl 2 . — This compound is prepared by distilling a mixture of potassium dichromate and sodium chloride with strong sulphuric acid. Chromyl chloride is a deep red, mobile, strongly fuming liquid. It is decomposed by water into hydro- chloric acid and chromium trioxide, and acts as a powerful oxidising substance. When dropped upon phosphorus it explodes. When heated in sealed tubes it is converted into trichromyl chloride with loss of chlorine, (Cr0 2 ) 3 Cl 2 . Chromyl chloride may be regarded as being derived from chromic acid, Cr0 2 (HQ ) 2 (unknown), by the complete substitution * Abney, “Treatise on Photography.” 616 Inorganic Chemistry °f (HO) by Cl. The intermediate compound, chloro-chromic acid Cr0 2 (H0)Cl is also unknown, although its salts have been pre- pared ; thus, by the gentle action of hydrochloric acid upon potassium dichromate, potassium chloro-chromate is obtained as a red crystalline salt — K 2 Cr 2 0 7 + 2HC1 = 2Cr0 2 (KO)Cl + H 2 0. Molybdenum, Mo = 95.9 ; Tungsten, W= 184 ; Uranium, U = 239.8. These three somewhat rare elements are closely related to chromium. Molybdenum occurs in the mineral molybdenite , MoS 2 , (which strongly re- sembles graphite in appearance), and more rarely as molybdenum ochre , MoO», and wulfenite , PbMo0 4 . Tungsten is found chiefly in wolfram , 2FeW0 4 ,3MnW0 4 (occurring in the Cornish tin mines); more rarely as scheelinite , PbWCb, and wolfram ichre, W0 3 . Uranium occurs as an oxide, U0 2 ,2U0 3 , in pitchblende (a considerable quantity of which, associated with other uranium compounds, has recently been discovered at St. Stephens, Cornwall). Molybdenum is obtained by the action of hydrogen upon the heated oxide or chloride ; uranium, by the action of sodium upon the chloride ; while tungsten has been obtained by both methods. In their specific gravities, tungsten and uranium exhibit a marked difference from chromium and molybdenum ; thus, Cr, sp. gr. = 6 ; Mo, sp. gr. = 8.6 ; while W, sp. gr. = 19. 1 ; U, sp. gr. = 18.7. Molybdenum and uranium form a large number of oxides, some of which are regarded as definite oxides, while others are looked upon as combinations of two oxides. Only two oxides of tungsten are known. The composition of these compounds is as follows — MoO Mo 2 0 3 Mo0 2 Mo0 3 The trioxide of each metal is an acid oxide ; uranium trioxide, however, is both acidic and basic. They are insoluble in water, but by the action of alkalies they yield molybdates, tungstates, and uranates. Molybdates and tungstates, derived from the acids H 2 Mo0 4 ,2H 2 0, and H 2 W0 4 ,2H 2 0 (corre- sponding to chromic acid), are known. And all three oxides yield salts corresponding to potassium dichromate, thus — Sodium Dimolybdate. Sodium Ditungstate. Sodium Diuranate. Na 2 Mo 2 0 7 Na 2 W 2 0 7 Na 2 U 2 0 7 . Molybdic and tungstic acids also form numerous polymolybdates and poly- tungstates, by the absorption of varying quantities of the trioxide into the wo 2 W0 3 uo 2 uo 3 uo 4 . u 2 o 5 = uo 2 ,uo 3 . U 3 O s = U0 2l 2U0 3 . Molybdenum , Tungsten, Uranium 617 molecule of the normal salt (see Chromates, page 615 ). And in the case of tungsten, the compound, metatungstic acid, H 3 W 4 0 13 ,7H 2 0, is known. Uranium dioxide and trioxide are both basic oxides, the former yielding the unstable uranous salts, such as uranous sulphate, U(S0 4 ) 2 ; and the latter producing the uranyl salts, of which the sulphate (U0 2 )S0 4 and (UQ 2 )(N 0 3 ) 2 , are well known. Uranium peroxide, U0 4 , is an acid oxide, which yields per-uranates. Both molybdic and tungstic acids form complex compounds with phos- phoric acid, known as phospho-molybdic and phospho-tungstic acids : thus, when a nitric acid solution of ammonium molybdate (NH 4 ) 2 Mo0 4 , is added in excess to a solution of orthophosphoric acid, or an orthophosphate, a canary yellow crystalline precipitate of ammonium phospho-molybdate 2(NH 3 ) 3 P0 4 ,22Mo0 3 ,12H 2 0 is obtained (see page 437 ). It is soluble in alkalies and in excess of phosphoric acid, but insoluble in dilute mineral acids. When this compound is dissolved in aqua-regia , the solution deposits yellow crystals of phospho-molybdic acid, 2H 3 P0 4 ,22Mo0 3 . Compounds with chlorine having the following composition are known — MoCl. *2 WC1 2 MoC 1 3 or Mo 2 C1 ( MoC1 4 MoC1 5 WC1 4 WC1 B WCl, CHAPTER XII GROUP VII. { FAMILY A.) MANGANESE. Symbol, Mn. Atomic weight = 54.8. Occurrence.— This element is never found in nature in the free state. It is widely distributed in combination with oxygen, as pyrolusite , Mn0 2 ; braunite , Mn 2 0 3 ; and hausmannite , Mn 3 0 4 . Also as a hydrated oxide in manganite , Mn 2 0 3 ,H 2 0. It is met with also as carbonate in manganese spar> MnCO s ; and as sul- phide in manganese blende , MnS. Modes of Formation. — Manganese may be obtained by the reduction of the oxide by means of carbon, at a very high temperature. The product, however, contains carbon. In a purer state it may be prepared by the reduction of fused an- hydrous manganous chloride, by means of metallic magnesium, or by reducing the oxide with aluminium at a high temperature. Properties. — Manganese is a hard, steel -grey, brittle metal. It rapidly oxidises on exposure to moist air, and is readily dis- solved by dilute sulphuric or hydrochloric acid, with evolution of hydrogen. Oxides of Manganese. — The four most important of these are— Manganous oxide {manganese monoxide . Red manganese oxide {hausmannite) Manganic oxide {manganese sesquioxide) . Manganese dioxide {pyrolusite) . MnO. Mn 3 0 4 . Mn 2 0 3 . Mn0 2 . The monoxide and sesquioxide are basic, giving rise to man- ganous and manganic salts respectively. The oxide, Mn 3 0 4 , is also basic, but yields with acids both manganous and manganic salts. Manganese dioxide, or peroxide, MnO^ gives manganous 618 Manganese Dioxide 619 salts, with elimination of available oxygen. It also combines with certain more basic oxides, forming unstable compounds known as manganites. Manganese trioxide, MnO s , and hept-oxide, Mn 2 0 7 , have also been obtained. They are both acid oxides, giving rise respec- tively to the 7>ianganates and permanganates. Manganous Oxide, MnO, is obtained by heating any of the higher oxides in a stream of hydrogen ; or by igniting a mixture of manganous chloride, sodium carbonate, and ammonium chloride. It is a light green powder, which, if prepared at a low tempera- ture, oxidises in the air. When perfectly air-free solutions of potassium hydroxide and a manganous salt are mixed, with exclu- sion of air, hydrated manganous oxide, or manganous hydroxide, Mn(HO) 2 , is obtained as a white precipitate, which rapidly oxidises on exposure to air. Red Manganese Oxide '{jnangano-manganic oxide), Mn 3 0 4 , is the most stable of the oxides of manganese, being formed when both the higher or lower oxides are strongly heated. Thus, in the preparation of oxygen by heating the dioxide, this compound remains (page 162). With cold sulphuric acid, it yields a mix- ture of manganous and manganic sulphates ; but when heated with dilute acid, manganous sulphate and dioxide are formed — Mn 3 0 4 + 2H 2 S0 4 = 2MnS0 4 + Mn0 2 .+ 2H 2 0. Manganic Oxide ( manganese sesquioxide ), Mn 2 0 3 , occurs native as braunite , and in the hydrated condition as manganite , Mn 2 0 3 ,H 2 0. The hydrated oxide is formed by the spontaneous oxidation of manganous hydroxide, and when gently heated it yields the oxide. Both the oxide and the hydrate, on treatment with warm nitric acid, yield manganous nitrate and manganese dioxide. Manganese Dioxide, Mn0 2 , is the most important of the man- ganese ores. It may be obtained by the cautious ignition of manganous nitrate — Mn(N0 3 ) 2 = N 2 0 4 + Mn0 2 . Manganese dioxide is a hard black solid, which conducts electri- city, and is strongly electro-negative to metals. On this account it is employed in certain forms of voltaic battery. When heated it loses oxygen, and forms first the sesquioxide, and finally Mn 3 0 4 . Manganese dioxide dissolves in cold concentrated hydrochloric ° 20 Inorganic Chemistry acid, forming a dark brown solution, which is believed to contain the compound Mn 2 Cl # . On warming, it evolves chlorine, and leaves manganous chloride, MnCl* Manganites.— Manganese dioxide combines with certain me- tallic oxides, forming unstable compound oxides. Thus with lime it forms CaO,Mn0 2 ; Ca0,2Mn0 2 , and Ca0,5Mn0 2 . These com- pounds are produced in the Weldon recovery process (page 321 ). MANGANOUS SALTS. Manganous Chloride, MnCl 2 , is the only chloride of this metal that has been isolated. It is obtained by dissolving any of the oxides, or the carbonate, in hydrochloric acid ; and on evaporation, is deposited as pink crystals of MnCl 2 ,4H 2 0. The anhydrous salt is prepared by heating the crystals in a stream of hydro- chloric acid. Manganese chloride forms double salts, with chlo- rides of the alkalies, the ammonium salt MnCl 2 ,2NH 4 Cl,H 2 0 being the best known. Manganous Sulphate, MnS0 4 , is prepared by strongly heating a pasty mixture of the dioxide and strong sulphuric acid. The iron present is thereby converted into ferric oxide, and on treat- ing the calcined mass with water, manganous sulphate dissolves. The solution on evaporation, deposits, at ordinary temperatures* large pink crystals of MnS0 4 ,5H 2 0 (isomorphous with copper sulphate). Below 6 ° rhombic crystals are formed (also pink) of the composition MnS0 4 ,7H 2 0 (isomorphous with ferrous sulphate). When these salts are heated to 200 °, or when their solutions are boiled, the anhydrous salt is formed. With sulphates of the alkalies, manganous sulphate forms double salts, as potassium manganous sulphate, K 2 S0 4 ,MnS0 4 ,6H 2 0 ; and with aluminium sulphate it yields a pseudo-alum (see page 574 ), MnS0 4 ,Al 2 (S0 4 ) 3 , MANGANIC SALTS. Manganic Chloride is obtained as a dark brown solution when the dioxide is dissolved in cold hydrochloric acid. It has never been isolated, and is believed to have the composition Mn 2 Cl 6 . Manganic Sulphate, Mn 2 (S 0 4 ) 3 , is obtained as a green deli- quescent powder, by the action of sulphuric acid upon the pre- Permanganates 62 1 cipitated peroxide. On exposure to the air, the deliquesced mass becomes muddy, by the precipitation of hydrated manganic oxide, thus — Mi^(S 0 4 ) 3 + 4H 2 0 = 3H 2 S0 4 + Mn 2 0 3 ,H 2 0. On the addition of potassium sulphate to a solution of manganic sulphate in dilute sulphuric acid, potassium manganese alum is obtained, K 2 S0 4 ,Mn 2 (S0 4 ) 3 ,24H 2 0, which deposits in violet regular octahedra. In the presence of much water the salt is decomposed, and deposits the hydrated manganic oxide. MANGANATES. These salts are derived from the hypothetical manganic acid, H 2 Mn0 4 . The oxide corresponding to this acid is known, viz., Mn0 3 . It is an unstable compound, obtained as a reddish amor- phous mass, by adding a solution of potassium permanganate in sulphuric acid to dry sodium carbonate. The manganates of the alkalies are obtained by fusing manganese dioxide with potassium or sodium hydroxide. If air be excluded, the following reaction takes place — 3Mn0 2 + 2KHO = K 2 Mn0 4 + Mn 2 0 3 + H 2 0. In the presence of air or oxygen, or by the addition of potassium nitrate or chlorate, more of the manganese is converted into man- ganate. The fused mass has a dark green colour, and dissolves in a small quantity of cold water to a deep green solution, which is only stable in the presence of free alkali. When a solution of potassium manganate is largely diluted, or gently warmed, it changes from green to pink, owing to the con- version of the manganate into permanganate, thus — 3K 2 Mn0 4 + 2H 2 0 = 2KMn0 4 + Mn0 2 + 4KHO. The same change takes place when carbon dioxide is passed through the solution. PERMANGANATES. These salts are derived from permanganic acid, HMn0 4 . When potassium permanganate is cautiously added to cold strong sul- phuric acid, green oily drops of the unstable manganese heptoxide 622 Inorganic Chemistry (or permanganic anhydride) are obtained, Mn 2 0 7 . This compound dissolves in a small quantity of water to a purple solution, which contains the unstable acid Mn 2 0 7 ,H 2 0, or H 2 Mn 2 0 8 = 2HMn0 4 . 1 he solution evolves oxygen and deposits manganese dioxide. Potassium Permanganate, KMn0 4 , is the most important salt of this class. It is prepared by fusing the dioxide with potassium hydroxide and potassium chlorate, dissolving the manganate so obtained in water, and passing carbon dioxide through the solu- tion. The filtered solution, on evaporation, deposits dark purple rhombic prisms, which appear deep red by transmitted light. Potassium permanganate is isomorphous with potassium per- chlorate, KC10 4 : it dissolves in water, forming a rich purple solution. When boiled with strong caustic alkalies it loses oxygen and forms the green potassium manganate — 2KMn0 4 + 2KHO = 2K 2 Mn0 4 + H 2 0 + O. It readily gives up oxygen to oxidisable and organic compounds, and on this account is used both as a laboratory oxidising agent, and as a disinfectant. The crude sodium salt is largely employed, under the name of Condos Disinfecting Fluid , for this purpose. When solid potassium permanganate is heated to 240 ° it evolves oxygen, and forms potassium manganate and manganese dioxide— 2KMn0 4 = K 2 Mn0 4 + MnO a + O,. CHAPTER XIII THE TRANSITIONAL ELEMENTS OF THE FIRST LONG PERIOD Iron, Fe = 55.88. Cobalt, Co = 58.6. Nickel, Ni = 58.6. These three elements belonging to Group VIII. (see classifica- tion, page 102) stand in a different relation to each other than the members of the other seven groups. Iron, cobalt, and nickel belong to the same period, being the transitidnal elements falling between the first and second series of the first long period. They are related, on the one hand, through iron, to the preceding metals manganese and chromium {see such compounds as ferrates , manganates, chromates ) : while, on the other hand, through nickel, they approach the metal copper, which is the next following in the period. Iron, cobalt, and nickel are closely related elements ; in nature they are usually associated together. They are all attracted by the magnet, and are nearly white, hard, and difficultly fusible metals. In their chemical habits, however, they exhibit a gradual transition in their properties. Thus, iron forms two basic oxides, yielding two series of stable salts, viz., ferrous and ferric. Cobalt also has two basic oxides, but the basicity of the sesquioxide is very feeble, and cobaltzV salts (except double salts) are unstable, and are only known in solution. Nickel only forms one basic oxide, and yields only one series of salts corresponding to the ferrous salts : the sesquioxide of nickel behaving with acids as a peroxide. IRON. Symbol, Fe. Atomic weight = 55.88. Occurrence. — Iron is one of the most abundant and widely distributed elements. It occurs in the uncombined state in small particles disseminated through certain basalts, and also in meteoric 6a j 624 Inorganic Chemistry iron, where it is usually associated with nickel, cobalt, and copper. Masses of iron have also been found, which have been formed by the reduction of iron ores, owing to the firing of coal pits : such iron is known as natural steel. The chief ores of iron are red hamatite and specular iron ore , Fe 2 0 3 ; brown hamatite , 2Fe 2 0 3 ,3H 2 0 ; magnetic iron ore (load- stone), Fe 3 0 4 ; spathic iron ore , FeC0 3 ; clay iron stone consists of spathose iron mixed with clay ; and blackband is clay iron stone containing from 20 to 25 per cent of coal. Iron is also found in combination with sulphur, as iron pyrites , FeS 2 , and with iron and copper in copper pyrites, Cu 2 S,Fe 2 S 3 , but these compounds are not employed in the metallurgy of iron. Modes of Formation. — Iron is readily reduced from its com- pounds. Thus, if ferric oxide, or oxalate, be gently heated in a stream of hydrogen, the metal is obtained as a black powder, which spontaneously oxidises with incandescence when brought into the air. On the industrial scale the reduction is effected by means of coke and limestone. The ore is first calcined, whereby water and carbon dioxide are expelled, and any sulphides present are oxidised, with the expulsion of sulphur dioxide. By this pro- cess also, the ore is rendered more porous. The calcined ore is then smelted in a blast-furnace, with limestone and coke. Fig. 146 shows in section a modern blast-furnace. The charge is admitted at the top by means of the cup and cone arrangement, which closes the furnace, and a powerful hot-blast is forced through tuyeres, placed round the base of the furnace. The furnace gases are led off by the side pipe at the top, and are utilised for heating the blast The chemical reactions which take place in a blast-furnace are many and complex, and differ in different parts of the furnace. In the main, the following are the changes which occur. The atmospheric oxygen of the hot-blast, on coming in contact with the carbon, forms carbon monoxide (at the high temperature carbon dioxide is probably not first formed). As the charges of ore gradually work their way down the furnace, they soon arrive at a point where the ferric oxide begins to be reduced by the heated carbon monoxide, first to ferrous oxide, and then to a spongy or porous mass of metallic iron. The region where this takes place is termed the zone of reduction — Fe 2 0 3 -1- 3CO = 3C0 2 + 2Fe. In the early stages of its descent through the furnace, the lime- Iron 625 stone is converted into carbon dioxide and lime. The reduced spongy metal, as it passes down through the hotter regions of the furnace, begins to take up carbon. It is probable that carbon monoxide first combines with the reduced iron, forming iron carbonyl (see page 246), which at a higher temperature is decom- posed, with the precipitation of finely divided carbon within the pores of the mass. More and more carbon is taken up by the iron 2 R 626 Inorganic Chemistry as it descends, until it passes from a pasty condition to a state of complete fusion, when it collects upon the bottom, or hearth, of the furnace. In passing through the hottest regions, the lime combines with the siliceous materials originally present in the ore, to form a fusible slag, beneath which the molten iron collects. Other re- actions which go on in various regions of the furnace, are the reduc- tion of sulphur compounds, and of phosphates and silicates, with the absorption into the iron of a certain amount of sulphur, phos- phorus, and silicon. The precise nature of the changes suffered by the gases in the various regions of the furnace, is still obscure. The cyanogen formed by the direct union of atmospheric nitrogen with carbon, and also the hydrocarbons present, doubtless undergo a chemical change in contact with the heated iron, and probably aid in its carburisation. The molten iron is drawn off at intervals from a tap-hole into moulds, and is known as cast iron or pig iron , The slag as it accumulates, overflows in a regular stream through an opening known as the slag hole. When such a furnace is in full blast, fresh charges of materials are introduced at regular intervals, and the process continues uninterruptedly for years. The metal obtained from the blast-furnace is far from pure iron, but contains varying quantities of carbon, silicon, phosphorus, sulphur, and manganese. The carbon may be present either in combination with iron as a carbide, or distributed throughout the metal as fine particles of graphite, or in both of these forms. White cast iron contains its carbon in the combined form, while grey cast iron owes it grey colour to the presence of minute crystals of graphite disseminated throughout the metal. When grey cast iron is dissolved in hydro- chloric acid, the graphite remains behind as a black powder ; but on similarly treating iron containing combined carbon , the carbon unites with the hydrogen, forming various hydrocarbons, which impart to the escaping gas a characteristic and unpleasant smell. Average cast iron contains from 90 to 95 per cent, of iron, and 3 to 5 per cent, of carbon. Spiegel is a variety of white cast iron con- taining 3. 5 to 6 per cent, of carbon, and from 5 to 20 per cent, of manganese. With more than 20 per cent, of manganese, the metal is termed perro-manganese. Purification. — The properties of iron are greatly modified by the presence of various impurities, especially carbon, and for different purposes for which iron is used, metal of different degrees of purity is required. The purest form of ordinary commercial Iron 627 iron is known as wrought iron, while steel is intermediate between this and cast iron. The process by which cast iron is converted into wrought iron, is termed puddling j and the method is called either dry puddling or pig-boilbig, depending upon whether the cast iron is subjected to a preliminary refining or not. The chemical reactions in both cases are the same, and consist in the oxidation of the impurities ; the carbon being expelled as carbon dioxide, while the oxides of silicon, phosphorus, and manganese pass into the slag. The method of pig-boiling is almost exclusively adopted. The cast iron is melted in a reverberatory furnace, the working bottom of which, as well as the lining (or fettling), consists of a layer of ferric oxide. The decarburisation of the iron is mainly effected by means of the oxide of iron derived from the fettling ; and for some time the molten mass appears to boil, owing to the escape of carbon monoxide. As the impurities are oxidised and removed, the mass becomes pasty (owing to the fact that the melting-point of pure iron is much higher than that of cast iron), and is then worked up into lumps, or blooms, which are ultimately removed and placed under a steam hammer, whereby admixed slag is squeezed out, and the metal is welded into a solid mass. Wrought iron contains from 0.06 to 0.15 per cent, of carbon. Steel may be produced either from wrought iron, by adding carbon, or from cast iron by removing that impurity. Formerly steel was exclusively obtained by the first method, by what is known as the cementation process. This simply consists in heating the bars of iron, buried in broken charcoal, for several days to a red heat. The precise nature of the chemical change which results in the carburisation of the iron, is not definitely established. In all probability, the carbon is conveyed into the body of the metal (which is not even heated to the softening point) by the intervention of iron carbonyl ; the carbon monoxide being formed by the union of the carbon with the air retained within the layer of charcoal. At the conclusion of the operation, the iron presents a blistered appearance, and on this account is termed blister-steel. At the present time, steel is mostly produced by the Bessemer process , which consists in oxidising the impurities present in cast iron, by blowing through the molten metal a blast of air. This operation is performed in a large pear-shaped vessel, known as a converter, which is mounted on trunnions, and through the bottom of which a powerful air blast can be admitted. The converter is 628 Inorganic Chemistry tilted into a horizontal position, and a quantity of molten cast iron is run in. The air blast is then started and the converter immedi- ately swung back into a vertical position. In the course of a very short time the whole of the impurities are burnt away, and the stage at which the operation is complete is sharply marked, by the sudden disappearance of the flame from the open mouth of the converter. The converter is once more swung into a horizontal position, and the blast is stopped. The exact quantity of molten spiegel is then added, to supply the carbon required to convert the entire charge into steel, and the blast is turned on for a few moments in order to thoroughly mix the materials, after which the contents are poured out into the casting ladle. The comparative purity of the three forms of iron will be seen from the three following typical examples : — Cast Iron. Steel. Wrought Iron. Carbon . . 3.81 O.65 0.10 Silicon . 1.68 0.07 0.05 Phosphorus 0.70 O.03 0.15 Sulphur . . 0.60 0.02 0.05 Manganese 0.41 O.4O 0.07 7.20 1. 17 0.42 Iron . 92.80 98.83 99.58 100.00 100.00 100.00 Properties. — Pure iron is a white lustrous metal, capable of taking a high polish. Its specific gravity is 7.84 to 8.139. It is more difficultly fusible and more malleable than wrought iron, but at a red heat it becomes soft and can be welded. The physical properties usually associated with iron, are in reality those of iron containing varying amounts of impurities : thus, pure iron when rendered magnetic, quickly loses this property, whereas steel retains its magnetism at ordinary temperatures, losing it, however, when heated. Pure iron, when heated and suddenly cooled, does not take a temper , while steel when so treated be- comes extremely hard and brittle. Iron is unacted upon by dry air, at ordinary temperatures, but in moist air, especially in the presence of carbon dioxide, it becomes coated with rust, and the process which is slow to begin, proceeds rapidly when a film of oxide has been once formed. Iron decomposes water readily at a red heat ; in the finely divided Oxides of Iron 629 state, the metal decomposes water at ioo°. Dilute hydrochloric and sulphuric acids rapidly dissolve iron with evolution of hydrogen. Dilute nitric acid dissolves it, forming ferrous nitrate and am- monium nitrate ; with stronger nitric acid, ferric nitrate and oxides of nitrogen are formed. Concentrated nitric acid (specific gravity 1.45) is without solvent action upon iron. A strip of iron which has been immersed in such strong acid, is unacted upon when afterwards dipped into the more dilute acid, and is also incapable of precipitating metallic copper from a solution of copper sulphate. Iron in this condition is said to be passive. Other oxidising agents, as chromic acid, or hydrogen peroxide, are capable of bringing about the same result. It is believed that this condition is due to the formation of a film of the oxide Fe 3 0 4 upon the surface. Finely divided iron takes fire spontaneously in chlorine ; and when gently warmed in sulphur dioxide it combines with that gas with incandescence. It absorbs carbon monoxide with formation of iron carbonyl, Fe(CO) 6 . When heated in ammonia it forms a nitride, Fe 6 N 2 . Oxides of Iron. — Three oxides of iron are known, namely : — Ferrous oxide (j iron monoxide) . . FeO. Ferric oxide (iron sesquioxide) . . Fe 2 0 3 . Ferroso-ferric oxide (magnetic oxide) . Fe 3 0 4 , or Fe 2 O s ,FeO. The two first are basic oxides, giving rise respectively to ferrous and ferric salts : the third yields both ferrous and ferric salts. Ferric oxide combines with certain more basic oxides, form- ing compounds analogous to Fe 2 0 3 ,Fe0 ; such as Fe 2 0 3 ,CaO, Fe 2 0 3 Zn0. These are known as ferrites. Ferrous Oxide (protoxide of iron\ FeO, is formed as an inter- mediate product during the reduction of ferric oxide by hydrogen, or carbon monoxide ; but it is difficult to obtain it free from either the higher oxide, or the metal. It is also formed when ferrous oxalate is heated out of contact with air. It is a black powder, which oxidises in the air, and which dissolves in acids yielding ferrous salts. Ferrous Hydroxide, Fe(HO) 2 , or FeO,H 2 0, is obtained as a white precipitate, when potassium hydroxide is added to a solution of a ferrous salt with entire exclusion of air. In the presence of air it is green. It readily absorbs oxygen and passes into ferric oxide* 630 Inorganic Chemistry Ferric Oxide {sesquioxide of iron), Fe 2 0 3 , occurs in brilliant black crystals belonging to the hexagonal system, in specular iron ore. It is obtained as a red amorphous powder by heating hydrated ferric oxide, ferrous sulphate, or ferrous carbonate. In a crystalline condition it may be produced by carefully heating a mixture of ferrous sulphate and common salt, or by heating the amorphous oxide in gaseous hydrochloric acid. The natural com- pound, and also the artificial substance after strong ignition, is only slowly dissolved by acids. Ferric oxide is extremely hygro- scopic. When strongly heated it is partially converted into Fe 3 0 4 . The amorphous substance, obtained by distilling ferrous sulphate for the manufacture of Nordhausen sulphuric acid, is employed as a red pigment, and a polishing powder, under the name of rouge. Ferric Hydroxide, or Hydrated Ferric Oxide, Fe 2 (HO) 6 , or Fe 2 0 3 ,3H 2 0. — When an excess of ammonia is added to a solution of ferric chloride, and the voluminous brown precipitate is dried at a moderate temperature, it has the composition Fe 2 0 3 ,3H 2 0. On exposure to various temperatures, or by precipitation under various conditions, hydrated oxides of the composition Fe 2 0 3 , 2H 2 0 ; Fe 2 0 3 ,H 2 0, and others, have been obtained ; and several of these compounds occur in nature. Ordinary rust of iron has the composition 2Fe 2 0 3 ,3H 2 0, or Fe 2 O 3 ,Fe 2 (HO) 0 . The monohydrate Fe 2 0 3 ,H 2 0 has been obtained as a soluble modification, by heating an acetic acid solution of precipitated ferric hydroxide to ioo° in sealed vessels. On the addition of sulphuric acid, a brown precipitate is obtained, having the composition Fe 2 0 3 ,H 2 0, which is insoluble in acids, but soluble in water. The solution gives no reaction with potassium ferro- cyanide. Another soluble hydroxide is produced by dissolving the ordinary precipitated hydroxide in ferric chloride, and subjecting the solution to dialysis. This solution is employed in medicine under the name of dialysed iron. Ferroso-ferrie Oxide, Fe 3 0 4 , occurs native as magnetite and magnetic oxide of iron ; the magnetic variety being known also as loadstone. When iron is heated in the air, the black film which forms (the so-called iron-scale , or hammer-scale) consists of the oxide FeaO^ with more or less ferric oxide, Fe 2 0 3 , upon the outer surface. It is also produced when steam or carbon dioxide is passed over heated iron, with evolution of hydrogen and carbon monoxide respectively, these reactions being the reverse of those Ferrous Sulphate 631 by which oxides of iron are reduced by hydrogen or carbon monoxide. Ferrates. — These compounds correspond to the manganates, but neither the acid FeH 2 0 4 nor the oxide Fe0 3 are known. Potassium ferrate, K 2 Fe0 4 , is formed when chlorine is passed through a solution of potassium hydroxide in which ferric hydroxide is suspended. FERROUS SALTS. Ferrous Chloride, FeCl 2 . — The anhydrous compound is ob- tained by heating iron wire in gaseous hydrochloric acid, when the salt sublimes in the form of white deliquescent crystals. In aqueous solution, it is obtained when iron is dissolved in hydrochloric acid, and is deposited in pale blue-green crystals of FeCl 2 ,4H 2 0. When heated in the air it is converted into ferric oxide and chloride, the latter volatilising — 6 FeCl 2 4 - 30 = Fe 2 O a + 2FeCl 3 . When volatilised in an atmosphere of hydrochloric acid, its vapour density at high temperatures corresponds to the formula FeCl 2 ; at lower temperatures it lies between the values required for FeCl 2 and Fe 2 Cl 4 . When strongly heated in a current of steam it is decomposed as follows — 3FeCl 2 + 4H 2 0 = Fe 3 0 4 + H 2 + 6HC1. Ferrous Sulphate {green vitriol ), FeS0 4 ,7H 2 0, is obtained when iron is dissolved in sulphuric acid. It is prepared on a large scale by exposing heaps of iron pyrites, FeS 2 , to the action of air and moisture. The liquor which drains away contains ferrous sulphate and sulphuric acid, and the latter is converted into ferrous sulphate by the introduction of scrap iron. Ferrous sulphate forms pale green monosymmetric crystals, which effloresce on exposure to the air. They are soluble in water to the extent of 70 parts in 100 parts of water at 15 0 , and 370 parts in 100 parts at 90 °. At ioo° the crystals lose 6H 2 0, being con- verted into FeS0 4 ,H 2 0. If a crystal of zinc sulphate be thrown into a supersaturated solution of ferrous sulphate, the iron salt is deposited in rhombic prisms (isomorphous with zinc sulphate). On the other hand, if a crystal of copper sulphate be added, asymmetric (triclinic) crystals of FeS0 4 ,5H 2 0 (isomorphous with copper sulphate) are formed. 632 Inorganic Chemistry Ferrous sulphate forms double salts with the sulphates of the alkalies. Thus, when mixed with ammonium sulphate in the re- quisite proportions, ammonium ferrous sulphate, FeS0 4 ,(NH 4 ) 2 S0 4 , GH 2 0, is obtained. This salt is less readily oxidised on exposure to air than ferrous sulphate itself. Ferrous salts give, with potassium ferrocyanide (K 4 Fe(CN) fl , or 4KCN,Fe(CN) 2 ), a white precipitate of potassium ferrous ferro- cyanide (FeK 2 Fe(CN) 6 , or 2KCN,2Fe(CN) 2 ). The precipitate is quickly oxidised, and becomes blue. With potassium ferricyanide (K 3 Fe(CN) fl , or 3KCN,Fe(CN) 3 ), ferrous salts yield a blue pre- cipitate of ferrous ferricyanide ( TurnbulPs blue ) (Fe 3 {Fe(CNU 2 , or SFe(CN) 2 ,2Fe(CN) 3 ), thus— 3FeS0 4 + 2K 3 Fe(CN ) 6 = Fe 3 {Fe(CN ) fl } 2 + 3K 2 S0 4 . FERRIC SALTS. Ferric Chloride, FeCl 3 , is prepared in the anhydrous state by passing dry chlorine over heated iron wire. In solution it may be obtained by dissolving iron in aqua regia ; or ferric oxide in hydrochloric acid. The anhydrous salt forms nearly black crystals, appearing deep red by transmitted light. It readily volatilises, and at temperatures above 700 ° the density of its vapour corresponds to the formula FeCl 3 , while at lower temperatures its density agrees more nearly with the formula Fe 2 Cl 6 . Ferric chloride is extremely deliquescent, and readily dissolves in water. When the solution is slowly evaporated, yellow crystals are deposited, having the composition Fe 2 Cl 6 , 12 H 2 0 (or FeCl 3 , 6 H 2 0 ). When a dilute solution of ferric chloride is boiled, it decomposes, forming either an insoluble oxychloride, or a soluble hydroxide and free hydrochloric acid (depending upon the strength of the solution). Ferric Sulphate, Fe 2 (S0 4 ) 3 , is prepared by the addition of sul- phuric and nitric acids to a solution of ferrous sulphate — GFeS0 4 + 3H 2 S0 4 + 2HNO s = 2NO + 4H 2 0 + 3Fe 2 (S0 4 ) 3 . The brown solution, on evaporation, leaves the anhydrous salt as a white mass. When the requisite quantity of potassium sul- phate is dissolved in a strong solution of ferric sulphate at o % the double potassium iron sulphate (iron alum), K 2 S0 4 ,Fe 2 (S0 4 ) 3 , 24H 2 0, separates out in the form of violet octahedrons. Ferric salts give, with potassium ferrocyanide (K 4 Fe(CN) 6 , or Sulphides of Iron 633 4KCN,Fe(CN) 2 ), a dark blue precipitate of ferric ferrocyanidc (- Prussian blue ) (4Fe(CN) 3 ,3Fe(CN) 2 ) or Fe 4 {Fe(CN) 6 } 3 — 4FeCl 3 + 3K 4 Fe(CN) 6 = Fe 4 {Fe(CN) 6 } 3 + 12KC1. With potassium ferricyanide ferric salts give no precipitate. SULPHIDES OF IRON. Ferrous Sulphide, FeS. — When a white-hot bar of wrought iron is dipped into melted sulphur, the elements unite ; and the readily fusible monosulphide of iron falls to the bottom. It may be prepared by throwing into a red hot crucible a mixture of iron filings and sulphur. So obtained, it is a dark, yellowish -grey, metallic-looking mass. When heated out of contact with air, it does not part with sulphur, but in the presence of air is converted into ferric oxide and sulphur dioxide. Ferrous sulphide is pre- cipitated from either ferrous or ferric solutions, by alkaline sul- phides, as a black amorphous powder, which in the moist state is quickly oxidised by the air. Dilute- sulphuric acid, or hydrochloric acid, decomposes ferrous sulphide, with evolution of sulphuretted hydrogen. Iron Sesquisulphide, Fe 2 S 3 , is formed when equal weights of iron and sulphur are heated to a moderate temperature. It can- not be obtained by precipitation from a ferric salt, as the product so formed c'onsists of ferrous sulphide and sulphur — Fe 2 Cl 6 + 3(NH 4 ) 2 S = 6NH 4 C1 + 2FeS + S. It is a yellow, metallic-looking solid, which is decomposed by dilute hydrochloric acid, yielding sulphuretted hydrogen. Ferric Disulphide, FeS 2 , occurs in nature in large quantities as iron pyrites, sometimes in the massive condition, and at others in the form of brass-yellow cubical crystals. In many cases the native compound bears the impression, or assumes the shape, of various organised forms, such as wood, ammonites, &c., the mineral having been formed by the reducing action of the organic matter, upon ferrous sulphate in solution. Ferric disulphide is also found in the form of brass-like, rhombic crystals, in radiated pyrites. The compound may be prepared by heating to a low red heat a mixture of ferrous sulphide and sulphur. Ferric disulphide is unacted upon by dilute acids; hot con 634 Inorganic Chemistry centrated hydrochloric acid decomposes it, with liberation of sul* phur and sulphuretted hydrogen. When heated in hydrogen, sulphur is evolved (which partly combines with the hydrogen), and ferrous sulphide remains. When heated in the air, ferric oxide and sulphur dioxide are formed. Ferroso-ferric Sulphide ( magnetic pyrites ), Fe 3 S 4 , occurs in the form of hexagonal crystals. Like the corresponding oxide, this compound is attracted by the magnet, and is itself sometimes magnetic. COBALT. Symbol, Co. Atomic weight = 58.6. Occurrence. — With the exception of small quantities present in meteoric iron, cobalt is not found uncombined in nature. Its chief natural compounds, which are only sparsely distributed, are speiss-cobalt , or smaltine , CoAs 2 ; cobalt glance , CoAsS, in both of which the cobalt is partially replaced by nickel and iron ; and cobalt-bloom , Co 3 (As0 4 ) 2 ,8H 2 0. Modes of Formation. — Cobalt is obtained by reducing the oxide, or the chloride, in a stream of hydrogen, or by strongly heating cobalt oxalate in a closed crucible. Properties. — Cobalt is an almost white, hard metal, which, when polished, resembles nickel, but is slightly bluer. It is malleable, and when heated is very ductile. Like both iron and nickel, it is attracted by the magnet ; but unlike these, it retains this property, even at a red heat. In the massive form, cobalt is unacted upon by the air ; but the finely-powdered metal, obtained by the reduction of the oxide in hydrogen, rapidly oxidises on exposure to the air, sometimes with incandescence. When heated in the air, it forms the oxide Co 3 0 4 . Cobalt decomposes steam at a red heat, yielding cobaltous oxide, CoO. Oxides of Cobalt. — Three oxides of cobalt are recognised, namely, cobaltous oxide, CoO ; cobaltic oxide, Co 2 O s ; and cobalto- cobaltic oxide, Co 3 0 4 . Four other oxides are known, which are regarded as compounds of the two first, having the composition 2CoO,Co 2 0 3 ; 3CoO,Co 2 0 3 ; 4CoO,Co 2 0 3 ; 6CoO,Co 2 0 3 . The monoxide, CoO, is basic, and yields the cobaltous salts. The sesquioxide, Co 2 0 3 , is feebly basic, forming only unstable Cobaltous Chloride 635 salts. Stable double salts, however, corresponding to this oxide are known. Cobaltous Oxide (< cobalt monoxide ), CoO, is formed when the sesquioxide is heated to redness in a stream of carbon dioxide, or gently heated in hydrogen. It is also obtained when the carbo- nate or hydroxide is heated in the absence of air. It forms a drab- coloured powder, which is unacted upon by the air, but when heated, forms Co 3 0 4 . When heated in either hydrogen or carbon mon- oxide, it is reduced to metallic cobalt. Cobaltous Hydroxide, Co(HO) 2 .— When potassium hydroxide is added to a solution of a cobaltous salt, a blue basic hydrate is precipitated, which, on boiling, is converted into the pink hydroxide Co(HO) 2 . It turns brown on exposure to the air, by the absorp- tion of oxygen. Both the oxide and hydroxide are really soluble in acids, giving cobaltous salts. Cobaltie Oxide ( cobalt sesquioxide ), Co 2 0 3 , is obtained by care- fully heating cobaltous nitrate until red fumes cease to be evolved. It is a dark grey powder, which, when strongly heated, is con- verted into the intermediate black oxide, Co 3 0 4 . Cobaltie oxide dissolves in cold acids, forming brown solutions, which contain unstable cobaltie salts. When warmed, these are converted into cobaltous salts, with evolution of oxygen in the case of oxy-salts, and of the halogen from haloid salts. This sesquioxide, therefore, behaves as a peroxide. Cobaltie Hydroxide, Co 2 (HO) 6 , or Co 2 0 3 ,3H 2 0, is obtained as a nearly black precipitate, by the addition of an alkaline hypochlorite to a cobaltous salt. With acids it behaves as the oxide. Cobalto-Cobaltie Oxide, Co 3 0 4 , is formed as a black powder, when the sesquioxide is strongly heated in air. COBALTOUS SALTS. Cobaltous Chloride, CoCl 2 . — When the carbonate, or any of the oxides, are dissolved in hydrochloric acid, the concentrated solution deposits dark red prisms (monosymmetric), having the composition CoC1 2 ,6H 2 0. When exposed over sulphuric acid, they lose 4H 2 0, and are converted into a rose-red salt, CoC1 2 ,2H 2 0, which reabsorbs moisture from the air to form the hexa-hydrate. When the di-hydrate is heated to about ioo°, it is converted into violet-blue crystals of CoCl 2 ,H 2 0 ; and at 120 ° it becomes an- 636 Inorganic Chemistry hydrous, and is blue. The blue salts, on exposure to the air, rapidly rehydrate themselves, and become pink. Cobaltous chloride dissolves in alcohol, giving a deep blue solu- tion, which, on the addition of water, also becomes pink. This property of forming pink hydrated salts, which become blue or green when nearly or quite anhydrous, is common to most cobal- tous salts. Thus, the iodide CoI 2 ,6H 2 0 forms rose-coloured crystals. When gently heated, it changes to a moss-green salt, CoI 2 ,2H 2 0, which, when dehydrated, becomes nearly black. Cobaltous Sulphate, CoS0 4 ,7H 2 0, is obtained by dissolving the carbonate or oxides in sulphuric acid, and is deposited from the solution in dark red crystals, isomorphous with ferrous sul- phate. Cobalt sulphate, like the sulphates of iron and nickel, forms double salts with alkaline sulphates, of which cobalt potas- sium sulphate, CoS04,K 2 S0 4 ,6H 2 0, is the best known. Cobaltie Salts. — Single salts corresponding to cobalt sesqui- oxide are unstable, and exist only in solution. More stable double salts are known. Thus, when potassium nitrite is added to an acetic acid solution of cobalt chloride, a yellow crystalline precipi- tate is obtained, consisting of the double nitrite of cobalt and potassium — 2CoC1 2 + 10KNO 2 + 4HN0 2 = Co 2 (NO 2 ) 0 ,6KNO 2 + 2NO + 4KC1 + 2H 2 0. The formation of this compound is made use of for separating cobalt from nickel, the latter element yielding no corresponding double nitrite. In the presence, however, of salts of barium, strontium, or calcium, nickel forms, with potassium nitrite, triple salts, sudh as Ni(N0 2 ) 2 ,Ba(N0 2 )2,2KN0 2 , which are precipitated as yellow crystalline powders. Hence, in the presence of metals of the alkaline earths, nickel and cobalt cannot be separated by this method. SULPHIDES OF COBALT. Cobaltous Sulphide, CoS, is obtained by heating cobaltous oxide with sulphur, or by fusing a mixture of cobalt sulphate, barium sulphide, and common salt. It forms bronze- coloured crystals, which are soluble in strong hydrochloric acid. Cobalt sulphide is precipitated as a black amorphous powder, when ammonium sulphide is added to a cobalt solution. The precipi- tate slowly dissolves in dilute mineral acids, but is insoluble in Cobaltamines 637 acetic acid. When heated in a stream of sulphuretted hydrogen, it yields the sesquisulphide Co 2 S 3 ; and if mixed with sulphur, and heated in a current of hydrogen, it forms the disulphide CoS 2 . Cobaltamines (ammoniacal cobalt compounds*). Cobalt forms a large number of complex ammoniacal salts. A few of these contain the metal in the divalent condition, and are known as ammonio-cobaltous salts j but by far the larger number contain the hexavalent double atom Co 2 , and are termed ammonio-cobaltic compounds. These compounds are classified as follows t : — Ammonio-Cobaltous Salts are formed by the absorption of gaseous am- monia by anhydrous cobaltous salts, or by dissolving the salts in strong aqueous ammonia, with exclusion of air. In this way the following salts have been obtained — Ammonio-cobaltous sulphate, CoS0 4 ,6NH 3 . Ammonio-cobaltous nitrate, Co(N0 3 ) 2 .6NH 3 ,2H 2 0. Ammonio-Cobaltic Salts. — These may be arranged under the following classes and subdivisions : — I. Hexammonio Salts. — General formula, Co 2 (NH 3 ) 6 ‘R 6 , where R equals a monacid radical, or its equivalent of di or tri acid radicals. { Hexammonio-cobaltic chloride ( dichro-cobaltic chloride ), Co 2 ’(NH 3 ) 6 'C1 6 ,2H 2 0. Hexammonio-cobaltic sulphate, Co 2 *(NH 3 ) 6 '(S0 4 ) 3 ,6H 2 0. II. Octammonio Salts — (a.)‘Praseo% Salts. — General formula, Co 2 *(NH 3 ) 8 *R 6 . (j8.) Fusco Salts. — General formula, Co 2 (NH 3 ) 8 (HO) 2 ’R 4 . { Fusco-cobaltic chloride, Co 2 (NH 3 ) 8 (HO) 2 ’C1 4 ,2H 2 0. Fusco - cobaltic sulphate, Co 2 (NH 3 ) 8 (HO) 2 *(S0 4 ) 2 , 2H a O. (7.) Croceo Salts. — General formula, Co 2 (NH 3 ) 8 (N0 2 ) 4 *R 2 . F tils $ Croceo-cobaltic chloride, Co 2 (NH 3 ) 8 (N0 2 ) 4 *C1 2 . xamp s | Croceo-cobaltic sulphate, Co 2 (NH 3 ) 8 (N0 2 ) 4 ’S0 4 . * For details respecting the preparation and properties of these salts, the student is referred to larger works. f On the constitution of metaliammonium compounds generally, see Werner, Zeitschrift fur Anorganische Chemie, 1893, vol. iii. X These names denote the characteristic colours of the salts ; thus, prasimus , leek-green ; fuscus , swarthy ; crocus , yellow, &c. Ammonio-cobaltous chloride, CoC1 2 ,6NH 3 which, at 120 0 , is converted into CoC1 2 ,2NH 3 . 638 Inorganic Chemistry III. Decammonio Salts — (a.) Roseo Salts.— General formula, Co^NHjhdHjObR*. { Roseo-cobaltic chloride, Co^NHa^^HaO^Cl#. Roseo-cobaltic sulphate, Co 2 (NH 3 ) 10 (H 2 O) 2 -(SO 4 ) 3 , 3H 2 0. (/3.) Purpureo Salts — General formula, Co 3 (NH 3 ) 10 X 2 R 4 (where X and R are either the same, or different acid radicals). 'Chloro-purpureo-cobaltic chloride, Co 2 (NH 3 ) 10 Cl 2 *Cl 4 . Chloro- purpureo -cobal tic sulphate, Co^NHghoCV Examples\ (S0 4 ) 2 . Bromo - purpureo - cobaltic nitrate, Co 2 (NH 3 ) 10 Br 2 ‘ . (no 3 ) 4 . ( 7 .) Xantho Salts. — General formula, Co 2 (NH 3 ) 10 (NO 2 ) 2 'R 4 . f Xantho-cobaltic chloride, Co 2 (NH 3 ) 10 (NO 2 ) 2 -Cl 4 . Examples -{ Xantho-cobaltic bromo-nitrate, Co 2 (NH 3 ) 10 (NO 2 ) 2 * { Br 2 -(NO s ) 2 . IV. Oxy-decammonio Salts . — General formula, Co 2 (NH 3 ) 10 R 4 , X*O(HO) (where X is either (HO) or an acid radical either the same as, or different from, R). { Oxy-decammonio cobaltic chloride, Co^NH^oCV (HO)'O'(HO). Anhydro - oxy - decammonio cobalt chloride, Co 2 (NH s ) 10 C1 4 -C1 2 ’O-(HO). V. Dodecammonio Salts (luteo-cobalticsalts). —General formula, Co 2 (N H 3 ) 12 R 6 . Examples i Lute °- cobaltic chlorid e, CoatNHsbaCl,,. r t Luteo-cobaltic sulphate, Co 2 (NH 3 ) 12 (S0 4 ) 3 ,5H 2 0. When cobalt compounds are fused with borax, a clear blue vitreous mass is obtained, which contains a borate of cobalt. A similar blue colour is imparted to ordinary potash glass, when a small quantity of a cobalt salt is added to the molten material, owing to the formation of a silicate of cobalt. Under the name of smalt , this substance has been manufactured for use as a pig- ment, by fusing the roasted cobalt ore with quartz sand and pearl- ash. The fused mass of deep blue glass is then finely ground beneath water. Nickel Alloys 639 NICKEL. Symbol, Ni. Atomic weight = 58.6. Occurrence. — Nickel occurs chiefly in combination with arsenic as kupfer nickel ,* Ni 2 As 2 ; white nickel , NiAs 2 ; nickel glance , Ni 2 (AsS) 2 , also as nickel blende , NiS. Nickel ore almost invari- ably contains cobalt, and frequently antimony and bismuth. Modes of Formation. — Nickel is obtained by reducing the oxide with carbon at a high temperature. It may be obtained as a black powder by reducing nickelous oxide in a stream of hydrogen, or by heating nickelous oxalate out of contact with air. It is also obtained as a lustrous coherent deposit by the electrolysis of an ammoniacal solution of the double sulphate of nickel and ammonia. Properties. — Nickel is a lustrous white metal, with a faint yellow tinge when compared with silver. It is ductile and malle- able, and at the same time very hard and tenaceous. It is sus- ceptible of a very high polish. Nidkel is attracted by the magnet, but loses this property when moderately heated. When obtained by reduction with charcoal, the metal contains a certain amount of carbon (like cast iron), which renders it less malleable ; and when produced by reduction of the oxalate at a low temperature the powder is pyrophoric. In the massive form, nickel is unacted upon by moderately dry air, but in moist air it tarnishes, and becomes covered with a film of nickelous oxide. It decomposes steam only slowly at a red heat, and is slowly attacked by dilute hydrochloric or sulphuric acid (contrast iron). In a finely divided state, nickel absorbs carbon monoxide (see page 263). Nickel is largely used for electro-plating iron and steel articles. Nickel Alloys. — With copper, and with copper and zinc, nickel furnishes several important alloys. The small coinage in use in Belgium, Germany, and the United States, consists of 1 part of * Kupfer nickel signifies the false copper , and was applied by the Germans in the middle ages to this ore, which resembled a copper ore, because they tried in vain to extract copper from it. It is probable that this ore had been smelted along with copper ores, under the belief that it contained copper, by the early ancients. Thus, a coin, 235 b.c., has been found to contain 20 per cent, of nickel. ^ 4 ° Inorganic Chemistry nickel and 3 parts of copper; while the so-called German silver, or nickel-silver, contains in addition about 1.5 parts of zinc. Oxides of Nickel. — Three oxides of nickel have been obtained namely, nickelous oxide, NiO ; nickelic oxide, Ni 2 0 3 ; and nickelo- nickelic oxide, Ni 3 0 4 . The first alone is basic. Nickelous Oxide ( nickel monoxide ), NiO, is obtained as a greenish powder by heating nickel carbonate, or hydroxide, out of contact with air. It is dissolved by acids yielding nickel salts. When heated in hydrogen, or carbon monoxide, it is readily re- duced to the metallic state. Nickelous Hydroxide, Ni(HO) 2 , is obtained as a pale green precipitate when potassium hydroxide is added to a solution of a nickel salt : the precipitate has the composition 4Ni(H0) 2 ,H 2 0. When strongly heated it is converted into nickelous oxide and water. It is readily soluble in acids, forming the nickel salts ; and it also dissolves in ammonia and in solutions of ammonium salts. Nickel Sesquioxide, Ni 2 0 3 , is obtained as a black powder when the nitrate is decomposed by heat at the lowest temperature. With hydrochloric acid and sulphuric acid it behaves like a per- oxide ; yielding nickel salts, with the elimination of chlorine and oxygen respectively — Ni 2 O s + 6HC1 = 2NiCl 2 + 3H 2 0 + Cl* Ni 2 O s + 2H 2 S0 4 = 2NiS0 4 + 2H 2 0 + O. It is soluble in ammonia, with evolution of nitrogen — 2Ni 2 0 3 + 2NH 3 = 6 Ni(HO ) 2 + 3H 2 0 + N* Hydrated Sesquioxide of Nickel, Ni 2 (HO) 6 , or Ni 2 0 3 ,3H 2 0. When chlorine is passed through water, or sodium hydroxide, in which nickelous hydroxide, N^HO)* is suspended, a black powder is obtained having the composition Ni 2 0 3 ,3H 2 0. The same com- pound is obtained when a nickel salt is added to a solution of bleaching-powder. In contact with acids and ammonia it behaves like the oxide. Niekelo-nickelic Oxide, Ni 3 0 4 , is obtained as a grey metallic- looking mass, when nickel chloride is heated to about 400 ° in a stream of oxygen. Nickel Salts. — Nickel forms only one series of salts, corre- sponding to the monoxide. In the anhydrous state these are usually yellowish, while in the hydrated condition they are green. Nickelous Sulphide 641 Nickel Chloride, NiCl 2 , is obtained as a yellow amorphous mass, by dissolving the oxide or carbonate in hydrochloric acid, and evaporating the solution to dryness. When heated in a current of chlorine it sublimes in the form of lustrous golden yellow scales, which dissolve in water forming a green solution. From the aqueous solution, green crystals of the composition NiCl 2 ,6H 2 0 are deposited. Anhydrous nickel chloride absorbs gaseous ammonia forming the compound NiCl 2 ,6NH 3 , which when deposited from an aqueous solution, forms blue octahedrons. Nickel Sulphate, NiS0 4 ,7H 2 0, is produced when the metal, the carbonate, or the oxide is dissolved in dilute sulphuric acid, and the concentrated solution is allowed to crystallise at the ordi- nary temperature. It forms green crystals, isomorphous with magnesium sulphate. When heated to ioo° the crystals lose 6H 2 0, and above 300° the salt becomes anhydrous. The anhy- drous salt absorbs gaseous ammonia, being converted into a pale violet powder having the composition NiS0 4 ,6NH 3 . When nickel sulphate is dissolved in strong aqueous ammonia, the solution deposits dark blue quadratic crystals of NiS0 4 ,4NH 3 ,2H 2 0. With sulphates of the alkalies, nickel sulphate forms double salts, of which the ammonium salt is the most important, NiS0 4 , (NH 4 ) 2 S0 4 ,6H 2 0. It is obtained by mixing concentrated solu- tions of the two sulphates in the requisite proportions. This salt is employed in the process of nickel-plating. Nickelous Sulphide ( nickel monosulphide ), NiS, occurs as the mineral capillary pyrites. It is obtained as a bronze-like mass, insoluble in hydrochloric acid, by heating sulphur and nickel together. In the hydrated condition, nickel sulphide is precipitated as an amorphous black powder, on the addition of ammonium sulphide to a nickel salt. The precipitate is scarcely soluble in hydrochloric acid, but partially dissolves in excess of ammonium sulphide forming a brown solution. Three other sulphides have been obtained, having the composition Ni 2 S, NiS 2 , and Ni 3 S 4 . CHAPTER XIV THE TRANSITIONAL ELEMENTS OF THE SECOND AND FOURTH LONG PERIOD Ruthenium, Ru = 103.5. Rhodium, Rh = 104. 1. Palladium, 106.2. Osmium, Os = 191. Iridium, Ir = 192.5. Platinum, 194.3. These elements, although constituting two transitional groups, are very closely related to each other. In nature they all occur associated together in what is commonly known as platinum ore, and they are on this account usually spoken of as the platinum metals. Platinum Ore, or native platinum, contains all these elements in the metallic state. It is found in small grains, sometitnes in nuggets, in alluvial deposits and river sand, principally in Brazil, Borneo, California, Australia, and the Urals. Native platinum contains from 60 to 86 per cent, of platinum, the remainder consisting of the other five metals of the group, together with varying quan- tities of gold, copper, and iron. Amongst the grains of platinum ore, there are also found grains which consist essentially of an alloy of platinum and iridium (containing from 30 to 75 per cent, of iridium) known as platin- iridium : and also particles of an alloy of osmium and iridium (called osmiri- dium ), which contain from 30 to 40 per cent, of osmium, as well as small quantities of rhodium and ruthenium. They are all white lustrous metals, having high melting-points. They are unacted upon by air or oxygen at ordinary temperatures ; and, with the excep- tion of osmium (which burns when strongly heated, forming the tetroxide), they are scarcely oxidised by oxygen at any temperature. With the exception of palladium, which readily dissolves in hot nitric acid, these metals are unacted upon by ordinary acids. Aqua regia converts osmium into the tetroxide ; it dissolves platinum with formation of the tetra- chloride, and slowly acts upon ruthenium, but is without action upon rhodium and iridium. The specific gravities of the metals of the first group, although very close to one another, are widely different from those of the second group ; and it will be seen that the specific gravities fall, with increasing atomic weights, thus — Ru, sp. gr. = 12.26. Rh, sp. gr. = 12. 1. Pd, sp. gr. = 11.4. Os, ,, =22.47. Ir. >• =22.38. Pt, ,, =21.5. The element osmium is the heaviest known substance. The most easily fusible of these metals is palladium, which melts about the temperature of wrought iron. The melting-point of platinum is somewhat higher, but it may be boiled by the oxyhydrogen flame. Rhodium and 642 Platinum 643 iridium come next in order of fusibility, the latter metal being just fusible by the oxyhydrogen flame, while ruthenium has a still higher melting-point. Osmium has not been melted. When heated to the melting-point of iridium, osmium volatilises ; and if air be present, it burns. The following oxides of these metals are known— — — — — Pd 2 0 RuO OsO RhO — PdO PtO Ru 2 0 3 Os 2 O s Rh 2 0 3 I r 2^8 Ru0 2 0s0 2 RhO a Ir 0 2 PdO a Pto 2 Ru0 4 Os 0 4 — — Ruthenium, osmium, rhodium, and iridium form salts corresponding sesquioxide, such as ruthenious chloride, Ru 2 C1 6 ; rhodium sulphate, RholsOJ, • iridious chloride, Ir 2 Cl 6 . With the exception of rhodium, they all form chlorides, corresponding to the dioxides, thus— ruthenic chloride, RuC1 4 ; iridic chloride, IrCl 4 ; platinic chloride, PtCl 4 , while palladium and platinum yield palladia* and palatine compounds, corresponding to their monoxides. The tetroxides of ruthenium and osmium are remarkable in melting at an extremely low temperature (about >40°), and boiling about 100°. They yield intensely irritating vapours, which, in the case of osmium tetroxide, exerts a most injurious effect upon the eyes, and is extremely poisonous. (Osmium tetroxide is commonly known as osmic acid.) Osmium and ruthenium also exhibit a non-metallic character in forming compounds derived from the unknown ruthenic and osmic trioxides, such as potassium ruthenate, K 2 Ru 0 4 , and potassium osmate, K 2 0 s 0 4 (the corresponding ruthenic and osmic acids are unknown). Ruthenium also forms potassium per-ruthenate, KRu 0 4 (analagous to permanganate), although the corresponding acid and peroxide, r u 2 0 7 , are unknown. The most important of these elements is platinum. PLATINUM. Symbol, Pt. Atomic weight = 194.3. In order to separate platinum from the other metals with which the native platinum, (see page 642) is mixed, the ore is digested in dilute aqua regia, under slightly increased pressure. The solution so obtained contains the higher chlorides of platinum, palladium, rhodium and iridium (for although in the pure state the last two named metals are scarcely attacked by aqua regia, when alloyed with much platinum they dissolve). The solution is evaporated to dry- ness, and heated to 125 0 , whereby the palladium and rhodium are obtained in the form of their lower chlorides, PdCl 2 and Rh 2 Cl 6 (the latter of which, in the anhydrous condition, is insoluble in water). The residue is extracted with water, and to the clear solu- 644 Inorganic Chemistry tion, acidified with hydrochloric acid, ammonium chloride is added. The double chloride of platinum and ammonium (PtCl 4 , 2 NH 4 Cl), separates out as yellow crystals, while the corresponding - iridium salt, being more soluble, remains for the most part in solution, and may be obtained by concentrating the mother liquor. The ammonium platinic chloride, on being ignited, loses ammonium chloride and chlorine, leaving the metal in the form of a black spongy mass known as spongy platinum, which is then melted by means of the oxyhydrogen flame in a lime crucible. The platinum so obtained, usually contains small quantities of iridium, and traces of associated metals. Pure platinum is obtained by alloying commercial platinum with pure lead, and treating the alloy first with nitric acid, which dissolves any copper and iron, a part of the palladium and rhodium, and most of the lead ; and then with dilute aqua regia, which dis- solves the whole of the platinum and the remaining lead, with traces of rhodium. From this solution the lead is precipitated as sulphate, and the platinum is then precipitated as the double chloride, by ammonium chloride. To remove traces of rhodium which are present, the dried double chloride is ignited with hydrogen potassium sulphate, whereby the rhodium is converted into a soluble double sulphate of rhodium and potassium, while the platinum is reduced to the condition of the spongy metal. Properties.— Platinum is a lustrous, greyish-white, malleable, and ductile metal. At a red heat it may be welded with great ease. It is melted by the oxyhydrogen flame, and vessels of platinum are readily made by fusing the metal together in this way. Heated platinum absorbs large quantities of hydrogen (see page 157) ; and when the metal is melted in the oxyhydrogen flame, it exhibits the phenomenon of “spitting,” when it again solidifies (see Silver, page 518). Platinum does not combine with oxygen at any temperature, neither does the heated metal absorb this gas ; but it has the property, when cold, of condensing oxygen upon its surface. A piece of clean platinum foil or wire, when introduced into a mixture of oxygen, and a readily inflammable gas or vapour (such as hydrogen, ether, alcohol, &c.), causes their combination ; and occasionally the metal becomes red hot, and ignites the mixture. This action is more rapid in the case of platinum sponge, when a larger surface is brought into play, and a fragment of this material introduced into a detonating mixture of oxygen and hydrogen at once determines its explosion. Platinum Dichloride 645 Platinum is not acted upon by either nitric or hydrochloric acid. It is oxidised when fused with caustic alkalies, or with potassium nitrate, and is also attacked by fused alkaline cyanides. In the form of sponge, it is dissolved by boiling potassium cyanide with the evolution of hydrogen, and formation of a double cyanide. Platinum readily combines with phosphorus, silicon, and carbon. The carbide of platinum is formed when the metal is continuously heated by a smoky flame, or one in which combustion is incom- plete, hence care is necessary in the use of platinum vessels. Platinum Black is the name given to the finely-divided metal obtained by precipitating platinum from its solutions by reducing agents, or by metals. It is a soft, black powder, which is capable of absorbing, or condensing upon its surface, large quantities of oxygen. It therefore acts as a powerful oxidising agent. Platinum Alloys. — Platinum readily alloys with many metals ; hence compounds of easily reducible metals should not be heated in vessels of platinum. The most important alloys are those with iridium. The addition of 2 per cent, of iridium is found greatly to increase the hardness, and raise the melting-point of platinum. An alloy containing 10 per cent, of iridium resists the corrosive action of chemical reagents to a greater extent than pure platinum (see Fluorine, page 310). Oxides of Platinum. — Platinous oxide, PtO, and platinic oxide, Pt 0 2 , are obtained in the form of dark grey or black powders, by gently heating the corresponding hydroxides. When strongly heated they are converted into the metal. Platinous Hydroxide, Pt(HO) 2 , is obtained by the action of potassium hydroxide upon platinum dichloride. It is a black powder, which dissolves in the halogen acids, yielding platin ous compounds. Platinie Hydroxide, Pt(HO) 4 , is prepared by adding boiling potassium hydroxide to a solution of platinum tetrachloride, and treating the precipitate with acetic acid to remove the potash. When dried it forms a yellowish powder, which is soluble in acids to form platinic salts. Platinic hydroxide behaves both as a weak base, and a feeble acid. With stronger bases it forms compounds known as filatinates, which are yellow crystalline salts. The sodium salt has the composition Na 2 0 , 3 Pt 0 2 , 6 H 2 0 . Platinum Dichloride ( platinous chloride ), PtCl 2 , is produced when platinum tetrachloride is heated to about 250°. It forms a greenish powder, insoluble in water. It dissolves in hydrochloric 646 / norganic Chemistry acid, giving a reddish-brown solution which is believed to contain the double compound PtCl 2 ,2HCl, or H 2 PtCl 4 , to which the name chloro-pi at inous acid has been given : the compound has never been isolated, but a number of double salts of platinous chloride with other chlorides are known, which may be regarded as derivatives of this acid, and which are therefore termed chloro- platinites ; thus, potassium platinous chloride, 2KCl,PtCl 2 , or potassium chloro-platinite, K 2 PtCl 4 , is obtained as fine red crystals, by adding potassium chloride to a solution of platinous chloride in hydrochloric acid. This salt is used in the platinotype photo- graphic process. Platinum Tetrachloride (. platinic chloride ), PtCl 4 , is obtained by dissolving the metal in aqua regia, and removing the excess of the acids by evaporating to dryness and gently heating the residue. From its aqueous solution, the salt deposits in large red crystals having the composition PtCl 4 ,5H 2 0, which are not deliquescent. When the salt is crystallised from a hydrochloric acid solution, or when the aqua regia solution is evaporated to expel the nitric acid, with frequent addition of hydrochloric acid, the double compound of platinic chloride and hydrochloric acid is formed, PtCl 4 ,2HCl, which is deposited as reddish-brown deli- quescent crystals, with 6H 2 0. To this substance (which is commonly called platinic chloride), the name chloro-platinic acid has been given, and the double salts of platinic chloride and various chlorides are regarded as salts of this acid. The most important of these chloro-platinates are those of the alkali metals, their different solubilities being made the basis for the separation of these metals. Potassium Chloro platinate (or potassium platinic chloride ), 2KCl,PtCl 4 or K 2 PtCl 6 , is obtained as a yellow crystalline pre- cipitate, by adding potassium chloride to platinic chloride. It is soluble in 100 parts of water at the ordinary temperature to the extent of 1. 1 parts, and at ioo°, 5. 18 parts. It is insoluble in alcohol. The rubidium and caesium compounds are very similar, but are still less soluble in water, 100 parts of water at 20° dissolving 0.141 of the rubidium and 0.07 of the caesium salt. Ammonium Chloro-platinate, 2NH 4 Cl,PtCl 4 , closely resembles the potassium salt, being slightly less soluble, but more so than the rubidium compound. Sodium Chloro-platinate, 2NaCl,PtCl 4 ,6H 2 0, is a reddish- yellow salt, readily soluble in both water and alcohol. Platinamines 647 Platino-eyanides. — Just as platinous chloride combines with metallic chlorides to form chloro-platinites, so platinous cyanide, Pt(CN) 2 , unites with other cyanides, forming similarly constituted double compounds, known as platino-eyanides.* Potassium platino-cyanide, K 2 Pt(CN) 4 , or 2 KCN,Pt(CN) 2 , is formed when spongy platinum is dissolved in boiling potassium cyanide. The platino-eyanides may be regarded as the salts of platino-cyanic acid, H 2 Pt(CN) 4 . Both the acid and the salts are characterised by the wonderful play of colours they exhibit, when viewed in different lights ; and by forming different coloured crystals with varying quantities of water of crystallisation (see page 192). Sulphides of Platinum. — Platinous sulphide, PtS, and platinic sulphide, PtS 2 . are obtained as amorphous black powders by the action of sulphuretted hydrogen upon the respective chlorides. Oxysalts of Platinum. — Few well defined single salts of platinum with oxyacids are known. This element, however, exhibits a great tendency to form complex double salts. One such series of compounds is seen in the ftlatino-nitrites , which may be regarded as the salts of platino-nitrous acid, H 2 Pt(N 0 2 ) 4 . These salts are remarkable, in that the platinum they contain cannot be detected by the ordinary tests for that metal ; just as the iron present in ferro-cyanides is not detected by the ordinary reagents used in testing for that metal. Ammoniacal Platinum Bases, or Platinamines. Like cobalt, platinum forms a large number of basic compounds with ammonia, many of which are of extremely complex composition. The first of these to be discovered was a bright green salt, obtained by the action of ammonia upon platinous chloride, having the composition PtCl 2 ,2NH 3 , or Pt(NH 3 ) 2 Cl 2 , and known as the green salt of Magnus. Many of the platina- mines exhibit isomerism ; thus, a compound known as the chloride of Reiset's second base is a yellow crystalline salt having the same composition as Magnus’s green salt. Twelve distinct series of ammoniacal platinum com- pounds are known, four of which are derived from platino^j and the remainder from platim'c salts ; the former are termed platoso ammonia compounds, while the latter are distinguished as th zplatino compounds. + * The name Cyano-platinites might with advantage be applied to these compounds. f For detailed descriptions of these compounds, the student is referred to larger works on chemistry : and on the constitution of these, and metallam- monium compounds generally, the article by Werner, in the Zeitschrift fit* Anorganische Chemie , 1893, vol. iii. p. 267, may be consulted. 6 4 S Inorganic Chemistry ARGON. Sym!)ol, A. Density, 19.9. Atomic weight, undetermined. History. More than a hundred years ago Cavendish observed that when a mixture of nitrogen ( phlogisticated air) and oxygen [dephlogisticated air ) was confined in a glass tube over mercury along with a solution of caustic potash and he gases exposed to the continued action of electric sparks there was a smal residue of gas (amounting to about of the volume of the nitrogen) which was not absorbed ; and he raised the question as to whether the phlogisticated air of our atmosphere is entirely of one kind.* I his observation and speculation of Cavendish’s remained buried until 1804 when Lord Rayleigh and Professor Ramsay announced to the world the dis- covery of a new gaseous constituent of the atmosphere. t In Hiking exact determinations of the densities of gases, Lord Rayleigh found that nitrogen obtained from atmospheric sources always gave a slightly higher number than that obtained for nitrogen which was prepared from chemical compounds. On careful investigation, in conjunction with Professor Ramsay it was found that this higher density of “ atmospheric nitrogen ” was due to the presence in the air of a hitherto unknown gas, which they succeeded in isolating, and to which they gave the name Argon; (1894) Occurrence.— Argon occurs in the atmosphere, where it is present to the extent of about .80 per cent., or 1 per cent, of “atmospheric nitrogen" is argon. It is also present in the occluded gases in certain specimens of meteoric iron. Argon has not been met with in a state of combination, and no com- pounds containing it are known. Modes Of Formation. -(1.) Argon may be obtained from the atmosphere by sparking a mixture of air and oxygen. The nitrogen combines with the oxygen, and the oxidised product is absorbed by potash. When no further contraction of volume is obtained, the excess of oxygen is removed by alkaline Pyrogallate, and the residual gas is the argon. Unless a high-tension alternat- ing electric discharge is employed the process is extremely slow. (2.) Argon may also be separated from the other atmospheric gases bv first withdrawing the oxygen by means of red hot copper, and after removing the carbon dioxide and aqueous vapour, passing the remaining gas over strongly heated magnesium turnings. The magnesium combines with the nitrogen (p. 208), and leaves the argon. In order to effect the complete absorption of every trace of nitrogen, the gas is passed backwards and forwards over the heated magnesium for many hours. Properties.— Argon is remarkable for its extraordinary inertness: and it has hitherto baffled all attempts to get it to combine with any other element. No compounds of argon therefore are known. The density of the gas is ig.g and therefore its molecular weight is 39.8. Whether its atomic weight is 19 g or 39.8 depends upon whether the molecules are diatomic, like hydrogen ard oxygen, or monatomic, like mercury. This point is not yet settled (1897) Argon is about times as soluble in water as nitrogen, 100 volumes of water at 12 dissolving about 4 volumes of argon. Owing to this higher solubility the gases which are expelled from solution in rain water are richer in argon than ordinary air. Argon is more difficult to liquefy than oxygen. Its critical temperature is — 121 , at which point a pressure of 50.6 atmospheres causes it c liquefaction. Liquid argon boils at - 187°, and solidifies at - 189. 6° (Olszewski)' The most characteristic lines in the spectrum of argon are two in the red (less refrangible than those of either hydrogen or lithium), a bright yellow line (more refrangible than that of sodium), and a group of five green lines. * Experiments on Air: Phil. Trans., 75, 372, 1785. Helium 649 HELIUM. During the year 1895 Professor Ramsay discovered that the gas which is found occluded in certain minerals, and which was supposed to be nitrogen, contains a gas whose spectrum is different from that of any other known substance. Tts spectrum is characterised by a brilliant yellow line, coincident with the line D 3 of the solar spectrum, which is the characteristic line of a hitherto unknown solar element discovered many years ago by Professors Lockyer and Frankland, and by them called “ helium.” The minerals from which terrestrial helium has been obtained in largest quantity are clbveite and broggerite, although traces of the gas have been dis- covered in many other minerals containing uranium. The gas is obtained from these minerals by strongly heating them in vacuo. The density of helium, determined by Ramsay, is 2.18 ; by C16ve, 2.02. Like argon, this gas is also believed to consist of monatomic molecules, in which case its atomic weight will be the same as the molecular weight, namely, 4-3 6 - By employing liquid hydrogen as a refrigerating agent, Dewar has succeeded in liquefying helium (May 1898), and he concludes that the boiling-point of the new gas is only slightly different from that of hydrogen. No combinations of helium have as yet been discovered. OTHER GASES IN THE ATMOSPHERE. During the present year, 1898, Prof. Ramsay has obtained indications of the presence in the atmosphere, in extremely minute quantities, of other hitherto unknown gases. Considerable quantities of liquid air are allowed to evaporate in such a manner that the gas derived from the residual and least volatile portions is collected in separate fractions ; in other words, the liquid air is submitted to a process of fractional evaporation. In this way it has been found that these residual gases exhibit spectra which in some cases appear to be different from those given by any other known gases, and this fact has led to the belief that at least three new gases are present, which have been named in the meantime by their discoverer, Krypton (hidden), Neon (new), and Metargon. Investigations into the nature of these gases are now being pro- secuted, and until more facts have been brought to light nothing definite can be said as to their elemental or compound character. INDEX Absolute boiling-point, 78 ,, temperature, 69 Absorptiometer (Bunsen), 125 Absorption of gases by charcoal, 256 Acetylene, 279 Acetylide of copper, 280 Acid, antimonic, 458 „ arsenic, 448 , , arsenious, 447 ,, boric, 564 ,, bromic, 345 ,, carbamic, 275, 405 ,, carbonic, 274 chloric, 336 ,, chloro-auric, 524 ,, chlorochromic, 616 „ chloroplatinic, 646 ,, chloroplatinous, 646 , , chlorosulphuric, 400 ,, chlorosul phonic, 400 ,, chromic, 614 ,, dithionic, 398 ,, hydrazoic, 245 hydriodic, 351 ,, hydrobromic, 343 ,, hydrochloric, 325 ,, hydrofluoboric, 567 ,, hydrofluoric, 312 ,, hydrofluosilicic, 585 ,, hydrosulphurous, 386 ,, hypobromous, 345 ,, hypochlorous, 335 ^ ,, hyponitrous, 225 ,, hypophosphorous, 432 ,, hyposulphuric, 398 ,, hyposulphurous, 386 „ iodic, 353 ,, manganic, 621 ,, metaboric, 564 Acid, metantimonic, 458 ,, metaphosphor ic, 436 ,, metarsenic, 448 ,, metasilicic, 588 ,, metastannic, 593 ,, metatungstic, 617 ,, metavanadic, 608 ,, molybdic, 617 ,, muriatic, 332 ,, nitric, 210 ,, nitrosulphuric, 388 ,, nitrous, 219 ,, Nordhausen sulphuric, 395 ,, ortho-antimonic, 458 ,, ortho-arsenic, 448 ,, ortho-arsenious, 447 ,, orthoboric, 564 ,, orthophosphoric, 434 ,, orthosilicic, 588 ,, osmic, 643 ,, oxymuriatic, 314 ,, pentathionic, 399 „ perchloric, 337 ,, perchromic, 612 „ periodic, 355 ,, permanganic, 621 ,, persulphuric, 385 ,, phosphomolybdic, 617 „ phosphoric, 434 ,, phosphoric (glacial), 437 „ phosphorous, 433 ,, pyro-antimonic, 458 ,, pyro-arsenic, 448 ,, pyro-arsenious, 447 ,, pyroboric, 565 ,, pyrophosphamic, 438 ,, pyrophosphodiamic, 438 ,, pyrophosphoric, 435 ,, pyrophosphotriamic, 438 65: Index 6 52 Acid, pyrostilphuric, 395 ,, pyrovanadic, 608 ,, selenic, 408 ,, selenious, 408 ,, silicic, 588 ,, stannic, 593 ,, sulphuric, 387 ,, sulphurous, 382 ,, telluric, 410 ,, tellurous, 409 ,, tetrathionic, 399 ,, thiocarbamic, 405 ,, thiocarbonic, 404 ,, thiosulphuric, 396 ,, trithionic, 398 ,, tungstic, 617 Acid-forming oxides, T7 Acids, dibasic, 18 ,, mono-, tetra-, and tribasic, 18 Affinities, 60 Affinity, chemical, 10, 60 After-damp, 262 Alabaster, 536 Algin, 348 Alkali manufacture, 490 ,, metals, 466 Alkali-waste, 361 Alkaline earths, 527 Allotropy, 17 1 Aludels, 348, 553 Alum, 573 ,, burnt, 575 ,, meal, 575 ,, shale, 575 ,, stone, 574 Alumina, 571 Aluminates, 572 Aluminite, 572 Aluminium, 569 ,, alloys, 571 ,, bronze, 509, 571 ,, chloride, 576 , , fluoride, 576 ,, hydroxides, 572 ,, sodium chloride, 577 ,, sulphate, 572 ,, sulphide, 571, 577 Alums, 573 Alunite, 574 Amalgamation process (silver), 515 Amalgams, 555 American pot-ashes, 482 Amethyst, 571 Ammonia, 239 ,, solubility of, in water, 241 Ammonia-soda process, 495 Ammoniacal cobalt compounds, 637 ,, liquor, 281, 501 ,, mercury compounds, 559 ,, platinum compounds, 647 Ammonium, 501 ,, alum, 574 amalgam, 501 ,, boroHuoride, 567 carbamate, 275, 503 ,, carbonate, 503 ,, chloride, 501 chloroplatinate, 646 chromate, 206 cyanate, 14, 23 dissociation of, 86 , . ferrous sulphate, 632 ,, hydrazoate, 246 ,, iron alum, 574 , , magnesium arsenate, 449 .. magnesium phosphate, 435 . , manganous chloride, 620 , , meta-thio-arsenate, 451 metavanadate, 608 molybdate, 617 nitrate, 223 ,, nitrite, 207 phosphomolybdate, 437, 617 ,, plumbic chloride, 603 ,, pyro-arsenite, 447 pyro-thio-arsenite, 450 , , salts, 501 . . sesquicarbonate, 504 , , sodium phosphate, 435 stannic chloride, 596 ,, sulphate, 502 ,, thiocyanate, 504 Ammon-sulphonates, 247 Amorphous silicon, 583 Analysis, 13 Anastase, 581 Anglesite, 596 Anhydrides, 17 Anhydrite, 536 I Animal charcoal, 255 Index 653 Anions, 93 Anodes, 92 Anthracite, 258 Antimonates, 458 Antimonious oxide, 457 Antiraoniuretted hydrogen, 453 Antimony, 451 ,, amorphous, 453 ,, blende, 451 ,, bloom, 451 ,, compounds with halogens, 454 ,, chlorides, 455 „ hydride, 453 „ ochre, 451 ,, oxides and oxy acids, 456 „ oxychlorides, 456 ,, sulphides, 457 ,, sulpho-trichloride, 456 ,, tetroxide, 457 j, trioxide, 457 Apatite, 309, 538 Apollinaris water, 195 Aquafortis, 209 Aqua regia, 216 Aqueous vapour (atmospheric), 231 Argentic compounds (see Silver), 514 Argentiferous lead, 516 Argentite, 514 Argon, 648 Arragonite, 537 Arsenates, 449 Arsenic, 439 , , allotropic mo difications of, 440 „ chlorhydroxide, 444 „ chloride, 443 ,, compounds with halogens, 443 „ fluoride, 443 ,, oxides and oxyacids, 444 ,, pentoxide, 448 , , sulphides, 449 ,, trihydride, 441 Arsenical iron, 439 ,, pyrites, 439 Arsenious bromide, 444 ,, iodide, 444 ,, oxide, 445 Arsenites, 447 Arsenuretted hydrogen, 441 Arsine, 441 Asbestos, 528 Asymmetric system, 139 Atacamite, 51 1 Atmolyses, 81 Atmosphere, 227 ,, composition of, 229 ,, height of, 237 ,, suspended impurities in, 236 Atmospheric ammonia, 233 ,, aqueous vapour, 231 , , carbon dioxide, 232 ,, gases mechanically mixed, 235 , , nitric acid, 233 ,, ozone, 234 Atomic heat, 44 ,, theory, 24 ,, volumes, 43, 104 ,, weight, definitions of, 36, 42 ,, weight, determination of, by chemical methods, 35, 57 ,, weight, determination of, by means of isomorphism, 49 ,, weight, determination of, by means of specific heat, 43 ,, weight, determination of, from volumetric relations 37 Atomic weights, list of, 21 Atoms, 4 Aurates, 525 Auric chloride, 524 ,, oxide, 525 Auro-auric sulphide, 525 Aurous iodide, 524 Autogenous soldering, 392 Avogadro’s hypothesis, 39 Azote, 205 Azurite, 506 Balling furnace, 492 Barium, 541 ,, amalgam, 541 ,, bromate, 345 ,, carbonate, 541 ,, chlorate, 336 ,, chloride, 543 ,, dioxide, 162, 542 ,, dithionate, 398 654 hidex Barium hydroxide, 542 ,, hypophosphite, 432 ,, iodatc, 542 ,, monoxide, 541 , , nitrate, 544 ,, oxides, 541 , , peroxide, 542 , , sulphate, 544 , , sulphide, 544 ,, tetrathionate, 399 ,, thiosulphate, 399 Baryta, 541 ,, water, 542 Barytocalcite, 541 Basic oxides, 17 ,, salts, 19 Basicity of acids, the, 18, 433 Battery, galvanic, 91 Bauxite, 569 Beryl, 528 Berylla, 528 Beryllum, 528 ,, aluminate, 572 ,, compounds, 528 ,, specific heat of, 46 Bessemer process (steel), 627 Biaxial crystals, optically, 140 Binary compounds, 15 Bismuth, 460 ,, alloys, 461 ,, carbonate, 463 ,, compounds with halogens, 461 ,, dichloride, 462 , , dioxide, 463 ,, glance, 460 ,, nitrate, 463 ,, nitrate, basic, 463 , , ochre, 460 , , oxides, 462 , , oxychloride, 464 , , pentoxide, 464 , , tetroxide, 464 ,, tribromide, 462 ,, trichloride, 461 ,, tri-iodide, 462 ,, trioxide, 463 ,, trisulphide, 465 Bismuthic oxide, 462 Bismuthous oxide, 462 Bisulphate of soda, 383 Bittern, 488 Bituminous coal, 258 Black ash, composition of, 494 ,, furnace, 492 ,, revolving furnace, 493 Black-band, 624 Black-jack, 546 Blacklead, 253 Blast-furnace, 625 Bleaching-powder, 165, 335, 535 Blister copper, 507 ,, steel, 627 Blue vitriol, 512 Boiling-point, definition of, 112 ,, absolute, 78 ■ Boiling-points, 113 ,, of saturated saline solutions, 117 ,, effect of pressure upon, 113 ,, effect of dissolved substances upon, 118 Bolognian phosphorus, 539 Bone ash, 413 ,, black, 255 Bones, composition of, 255 Boracite, 562 Borate spar, 562 Borates, 565 Borax, 565 Borofluorides, 567 Boron, 562 ,, hydride, 568 ,, nitride, 568 ,, sulphide, 568 ,, trichloride, 567 ,, trifluoride, 566 ,, trioxide, 563 Boronatrocalcite, 562 Bort, 250 Boyle, law of, 70 Brass, 509 Braunite, 618 Brin’s process (oxygen), 162 Britannia metal, 452, 592 British Channel, composition of, 195 Bromates, 346 Bromides, 344 Bromine, 339 ,, hydrate, 342 Index 655 Bromine monochloride, 356 ,, oxyacids, 344 ,, water, 342 Bronze, 592 Brookite, 581 Brown haematite, 624 Brown iron ore, 624 Brucite, 530 Bunsen flame, the, 303 , , non-luminosity of, 304 ,, temperature of, 305 Burnt alum, 575 Cadmium, 551 , , chloride, 552 ,, oxide, 551 ,, sulphide, 552 Caesium, 466, 500 ,, spectrum of, 470 Cailletet’s apparatus, 74 Calamine, 546 Calcined magnesia, 529 Calcite, 532 Calcium, 532 ,, bicarbonate, 197, 274, 538 , , borate, 566 ,, borofluoride, 567 ,, carbide, 280 , , carbonate, 537 ,, chlorate, 478 ,, chloride, 534 ,, chloro-hypochlorite, 536 ,, dioxide, 534 ,, fluoride, 312 ,, hydroxide, 534 ,, hypochlorite, 478, 536 ,, manganite, 321, 620 ,, oxides, 533 ,, phosphate, 413, 538 ,, phosphide, 421 ,, sulphate, 536 ,, sulphide, 361, 372, 538 Calc-spar, 532 Caliche, 349 Calomel, 557 Calorie, 144, 288 Calx, 227, 283 Candle flame, 298 Canton’s phosphorus, 539 Capillary pyrites, 641 Carat, definition of, 524 Carbide of barium, 280 ,, of iron, 474 Carbon, 250, 581 ,, compounds, 259 ,, dioxide, 264 ,, ,, atmospheric, 232 ,, ,, composition of, 272 11 M solid, 271 ,, disulphide, 402 ,, hydrogen, compounds of, 276 ,, monoxide, 260 ,, oxides of, 260 , , specific heat of, 45 Carbonado, 250 Carbonates, 274 Carbonyl chloride, 263 -Carbonyls, metallic, 263 Carborundum, 584 Carboxy-hsemoglobin, 262 Carnallite, 471, 528 Carre's freezing machine, 116 Cassiterite, 590 Cast iron, 626 Catalysis, i6l Catalytic action, 12, 161 Cathions, 93 Cathodes, 92 Caustic potash, 476 ,, soda, 487 Celestine, 539 Cellulose, 215 Cementation process (steel), 627 Cerite, 581 Cerium, 581 Cerussite, 596 Chalcedony, 582 Chalk, 357 Chalybeate waters, 195 Chamber acid, 391 Chamber crystals, 388 Chance’s process, 371 Change of volume on solidification, 119 Charcoal, 254 ,, absorption of gases by, 256 , , animal, 256 , , specific heat of, 45 Charles’ law, 68 Chemical action, 11 Chemical action, modes of, 13 Index 656 Chemical action, affinity, 10 ,, combination, laws of, 24 ,, equations, 22 ,, formulae, 22 11 nomenclature, 15 ,, notation, quantitative, 52 ,, reactions, 22 ,, symbols, 20 Chili saltpetre, 497 Chlorates, 337 Chloride of lime, 535 Chlorine, 314 ,, hydrate, 324 , liquefaction of, 72 ., liquid, 324 ,, monoxide, 333 , , oxides and oxyacids, 333 ,, peroxide, 334 , , water, 323 Chloro-aurates, 535 Chloro-chromates, 616 Chloro-stannates, 595 Chromates, 614 Chrome alum, 613 ,, green, 610 ,, iron ore, 609 ,, ochre, 609 ,, red, 615 ,, yellow, 615 Chromic anhydride, 61 1 ,, chloride, 612 ,, hydroxides, 610 , , sulphate, 613 Chromite, 609 Chromites, 613 Chromium, 609 ,, anhydride, 610 ,, chromate, 610 ,, dioxide, 610 ,, oxides of, 610 ,, sesquioxide, 610 ,, trioxide, 61 1 Chromous chloride, 612 ,, hydrated oxide, 610 ,, sulphate, 612 Chromyl chloride, 615 Chrysoberyl, 528, 572 Cinnabar, 552 Clark’s process for softening water, tq8 Classification of elements, 97 Clay, 569 Clay, ironstone, 624 Coal, 257 ,, gas, 281 Coarse metal (copper), 507 Cobalt, 634 ,. bloom, 634 ,, glance, 439, 634 ,, oxides of, 634 Cobaltamines, 637 Cobaltic hydroxide, 635 ,, oxide, 635 Cobalto-cobaltic oxide, 635 Cobaltous chloride, 635 ,, hydroxide, 635 ,, oxide, 635 ,, sulphate, 636 ,, sulphide, 636 Coefficient of absorption, 124 . , solubility, 124 Coefficients of expansion of gases, 69 Coke, 254 Colemanite, 562 Colloids, 589 Columbite, 607 Combining proportions, 29 Combustibles, 284 Combustion, 283 ,, gain in weight by, 287 ,, heat of, 288 ,, supporters of, 284 Common salt, 488 Compound radicals, 22 Compounds, 7 Condy’s fluid, 622 Constant composition, law of, 30 Constitution of matter, 3 Copper, 506 ,, acetylide, 280 ,, alloys, 509 ,, arsenite, 447 ,, bromide, 512 ,, carbonates, 513 ,, chlorides, 511- — ,, ferrocyanide, 13c ,, fluoride, 512 ,, glance, 506 ,, hydroxide, 510 ,, nitrate, 512 Tndex 657 Copper nitroxyl, 219 11 oxides, 509 ,, oxychloride, 511 ,, pyrites, 506 ,, sulphate, 512 ,, sulphides, 513 Coprolites, 538 Coral, 532 Corpse light, 292 Corrosive sublimate, 558 Corundum, 569 Cream of tartar, 457, 472 Crith, 55 Critical pressure, 78. ,, temperature, 77, 117 Croceo-cobaltic salts, 637 Crocoisite, 609 Crookesite, 577 Cryohydrates, 192 Cryolite, 309, 569 Crystalline forms, 138 Crystallisation, suspended, 118 ,, water of, 192 Crystalloids, 589 Cubical nitre, 497 Cupel, 516 Cupellation process (silver), 516 Cupric carbonates, 513 „ chloride, 51 1 ,, hydroxide, 510 , nitrate, 512 ,, oxide, 510 ,, sulphate, 512 ,, sulphide, 513 Cuprous acetylide, 280. ,, chloride, 51 1 ,, oxide, 509 ,, sulphide, 513 Cyanide process, gold, 523 Dalton, atomic theory, 29 Davy lamp, 292 Deacon’s process, 316 Dead Sea, solid matter in, 195 Deep well waters, 196 Deliquescence, 193 Dephlogistigated air, 159 Dew-point, 232 Diatomic molecules, 8 Dialysed iron, 630 Dialysis, 589 Diamond, 250 ,, combustion of, 252 , , specific heat of, 45 Diffusiometcr, 79 Diffusion of gases, 78 • > ,, law of, 80 ,1 of dissolved substances, 137 Dimorphism, 141 Dissociation, 85 Disulphur dichloride, 374 Disulphuryl chloride, 401 Dithionates, 398 Divalent elements, 58 Dolomite, 528 Dry copper, 507 Dulong and Petit, law of, 44 Dutch brass, 509 ,, metal, 322 Dyad elements, 58 Earth’s crust, composition of, 160 Ebullition, 115 Efflorescence, 192 Effusion of gases, 82 Eka-aluminium, 108 Eka-boron, 108 Eka-silicon, 108 Electro-chemical equivalents, 94 Electro-gilding, 524 Electrolysis, 91 Electrolytes, 92 Electrolytic dissociation, 94 Electroplating, 92, 515 Elements and compounds, 6 ,, classification of, 102 ,, list of, 21 ,, non-metallic, 8 Elton Lake, water of, T95 Emerald, 528 Emery, 571 Empyreal air, 159 Endosmometer, 134 Endosmose, 134 Endothermic compounds, 147 English brass, 509 Epsom salts, 531 Equations, chemical, 22 Equivalence ( see Valency, 58) Equivalents, chemical, 29 1 ) electro-chemical, 93 2 T Index 658 Estramadurite, 413, 538 Ethene, 278 Ethine, 279 Ethyl silicate, 585 Ethylene, 278 ,, dibromide, 278 Euchlorine, 334 Eudiometry, 227 Evaporation, 111 ,, cold produced by, 76, 114, 243 Exothermic compounds, 147, 293 Expansion by heat of liquid carbon dioxide, 270 Expansion by heat of liquid oxygen, 170 Extincteur, 267 Faraday’s law, 93 Felspar, 569, 590 Ferrates, 631 Ferric chloride, 632 ,, ferrocyanide, 633 ,, hydroxide, 630 ,, soluble, 630 ,, oxide, 630 ,, sulphate, 632 ,, sulphide, 633 Ferrites, 629 Ferro-manganese, 626 Ferroso-ferric oxide, 630 ,, sulphide, 634 Ferrous bromide, 481 chloride, 631 ,, chromite, 614 ,, ferricyanide, 632 ,, ferrocyanide, 632 , , hydroxide, 629 ,, oxide, 629 ,, sulphate, 631 ,, sulphide, 633 Fettling, 627 Fine metal (copper), 507 Fire-damp, 277 Fire-damp caps, 294 Fixed air, 264 Fixed alkali, 467 Flame, 294 ,, candle, 297 ,, the Bunsen, 303 Flame, structure of, 296 Flames, cause of luminosity of, 301 Flint, 582 Flintshire furnace, 597 Fluorapatitc, 309 Fluorides, 315 t Fluorine, 308 Fluor-plumbates, 310 Fluor-spar, 309, 533 Forces, chemical and physical, 3 Formula weight, 53 Formulce, 22 Fraction of dissociation, the, 88 Franklinite, 546 Fulminating gold, 525 ,, silver, 520 Fusco-cobaltic salts, 637 Fusible metal, 461 Fusion, latent heat of, 120 Gadolinite, 561 Gahnite, 546 Galena, 541, 596 Gallium, 109, 561 Galvanised iron, 548 Gas carbon, 254 Gases, absorption by charcoal, 256 coefficients of expansion of 69 critical pressure, 78 , , critical temperature of, 77 ,, diffusion of, 78 ,, effusion of, 82 ,, kinetic theory of, 82 ,, liquefaction of, 71 ,, occlusion of, 257 ,, relation to heat, 68 ,, relation to pressure, 70 ,, solubility of, in liquids, 122 ,, transpiration of, 82 Gastric juice, 490 Gay-Lussac, law of, 25, 37 General properties of gases, 68 ,, liquids, no German silver, 548 Germanium, 109, 581 Gilding, 524 Glauberite, 496 Glauber’s salt, 496 Glucinum. 528 Index 659 Gold, 523 ,, alloys, 524 , , compounds of, 524 , , fineness of, 524 ,, fulminating, 525 Graduators, 488 Graham’s law, 80 Gramme-molecule, 56 Graphite, 252 ,, specific heat of, 45 Greenockite, 551 “Green salt of Magnus," 647 Green vitriol, 395, 631 Grey antimony ore, 451 ,, cast iron, 626 Guignet’s green, 61 1 Gun-cotton, 215 Gun-metal, 509 Gunpowder, 484 ,, products of combustion of, 484 Gypsum, 536 ,, fibrous, 536 Haematite, 624 Haemoglobin, 170 Hair salt, 572 Halogens, 18, 307 Haloid salts, 18 Hardness (water), 197 Hargreave’s process, 497 Hausmannite, 618 Heat, atomic, 44 , , molecular, 47 ,, of combustion, 288 ,, of formation, 146 , , specific, 44 „ specific, table of, 45 y , , units, 144 Heavy spar, 541 Helium, 649 Henry’s law, 123 Hepar sulphuris , 485 Hexagonal system, 139 Holmes’s signal, 424 Horn mercury, 557 Horn silver, 514 Hydrazine, 245 ,, hydrochloride, 245 ,, hydrate, 245 .. sulphate, 245 Hydrocarbons, 276 Hydrogen, 150 ,, chloride, 325 ,, compounds with oxygen, 179 ,, dioxide, 179 ,, disodium phosphate, 435 ,, displaceable, 18 ,, monoxide, 179 ,, nitrate, 19 ,, occlusion, of, 150, 157 , , peroxide, 199 , , persulphide, 373 ., phosphide, 420 potassium fluoride, 309 ,, potassium sulphate, 395 ,, sodium ammonium phos* phate, 435 ,, sulphate, 19 sulphide, 369 Hydrogenium, 157 Hydromagnesite, 532 Hydroxides, 17 Hydroxyl, 247 ! Hydroxylamine, 246 ! ,, disulphonate, 248 , , hydrochloride, 247 ,, mono-sulphonate, 248 Hypobismuthic oxide, 462 Hypobismuthous oxide, 462 Hypochlorites, 336 Hypochlorous anhydride, 333 Hyponitrous anhydride, 223 Hypophosphites, 433 Hypovanadic chloride, 608 ,, oxide, 608 ,, sulphate, 608 Ice, 190 ,, effect of pressure upon, 190 ,, the melting-point of, 120 Icicle, 1 16 Ignition-point, 291 Indifferent substances, 136 Indigo-copper, 513 Indium, 107, 561 Inflammable air, 150 Intestinal gases, hydrogen in, T50 Iodates, 354 I Iodic anhydride, 353 I Iodine, 346 66o Index Iodine, bromides, 357 i, chlorides, 346 „ pentoxide, 353 Ions, 93 Iridium, 642 ,, chlorides, 643 ,, oxides, 643 Irish Sea, solid impurity in, 195 Iron, 623 „ alum, 632 „ carbide, 474, 626 ,, carbonyl, 264 ,, magnetic oxide of, 630 „ monoxide, 629 „ oxides of, 629 „ passive; 629 „ pyrites, 360, 624 ,, sesquioxide, 630 ,, sesquisulphide, 633 ,, sulphides of, 633 Isodimorphism, 14 1 Isogonism, 50 Isomerism, 171 Isomorphism, 49, 141 ,, law of, 49 Isotropic crystals, 140 Jolly's apparatus, 230 Kainite, 471 Kelp, 347 Kelp substitute, 345 Kiesel-guhr, 586 Kieserite, 528, 531 Kinetic theory, 82 Kish, 253 Kupfernickel, 439, 639 Lagoons (boric acid), 564 Lakes, 572 Laminaria digitata, 347 ,, stenophylla, 347 Lamp-black, 254 Lanarkite, 596 Lanthanum, 561 Latent heat of fusion, 120 „ ,, vaporisation, 114 Laughing-gas, 223 Law of Boyle, 70 ,, Charles, 68 Law of constant heat consummation, 148 ,, constant proportion, 24 ,, Dulong and Petit, 44 ,, gaseous diffusion, 80 ,, Gay-Lussac, 25, 37 ,, multiple proportions, 26 ,, octaves, 98 ,, partial pressures, 127 ,, periodic, 100 ,, reciprocal proportions, 27 Layer crystals, 50 Lead, 596 ,, acetate, 605 ,, action of water upon, 599 ,, carbonate, 604 ,, chromate, 615 ,, composition of commercial 600 ,, desilverisation of, 516 ,, dichloride, 602 ,, dioxide, 602 Lead ethide, 582 , , nitrate, 603 ,, oxides of, 600 ,, oxychloride, 603 ,, sesquioxide, 601 ,, softening of, 598 , , squirted, 600 ,, suboxide, 600 ,, sulphate, 605 ,, sulphide, 606 ,, sulphochlorides, 606 , , tetrachloride, 603 ,, tree, 598 ,, white, 604 Leblanc process, 490 Leguminous plants, 233 Lepidolite, 499 Light red silver ore, 514 Lime, 533 » chloride of, 535 ,, dead burnt, 533 ,, milk of, 534 ,, quick, 534 ,, slaked, 534 ,, superphosphate of, 538 Limestone, 532 Liquefaction of air, 76 ,, of gases, 71 Index 66 Liquids, general properties of, no Liquor ammonia , 242 Litharge, 600 Lithium, 499 ,, carbonate, 500 ,, hydroxide, 500 „ mica, 499 „ nitride, 208 ,, oxide, 500 ,, phosphate, 500 ,, spectrum of, 469 Liver of sulphur, 485 Load-stone, 624, 630 Lothar Meyer’s curve, 105 Lucifer matches, 419 Luminous paint, 539 Lunar caustic, 522 Luteo-cobaltic salts, 638 Magistral, 515 Magnesia, 529 Magnesia alba levis, 532 ,, ponderosa , 532 ,, usta, 529 Magnesia mixture, 531 Magnesian limestone, 196, 533 Magnesite, 528 Magnesium, 528 , , aluminate, 572 ,, ammonium chloride, 530 „ ammonium phosphate, 435 ,, boride, 568 ,, bromate, 346 ,, calcium chloride, 530 carbonates, 532 ,, chloride, 530 Sl combustion of, in steam, 152 ,, hydroxide, 530 ,, nitride, 208, 529 „ oxide, 529 ,, oxychloride, 530 ,, phosphate, 435 ,, platinocyanide, 193 ,, potassium chloride, 530 ,, pyrophosphate, 436 ,, silicide, 584 „ sulphate, 531 Magnetic iron ore, 624 ,, oxide of iron, 630 pyrites, 634 Magnetite, 630 Malachite, 506 Manganates, 621 Manganese, 618 ,, blende, 618 ,, dioxide, 619 ,, monoxide, 618 ,, oxides of, 618 ,, sesquioxide, 619 ,, spar, 618 Manganic chloride, 620 ,, oxide, 619 ,, sulphate, 620 Manganite, 618 Manganites, 620 Mangano-manganic oxide, 619 Manganous chloride, 620 ,, chromite, 614 ,, hydroxide, 619 ,, sulphate, 620 Marble, 537 Marine acid air, 325 Marsh gas, 276 ,, synthesis of, 404 Marsh’s test, 442 Massicot, 600 Matches, 419 Matlockite, 596 Mechanical mixtures, 8 Mediterranean Sea, 195 Meerschaum, 528 Mendelejeff s periodic law, 100 Mephitic air, 205 Mercuric ammonium chloride, 558 , , chloride, 558 ,, iodide, 558 ,, oxide, 556 ,, potassium chloride, 65 Mercurius calcinatus per se, 159 Mercurous chloride, 557 ,, nitrate, 557 , , oxide, 556 ,, sulphate, 557 Mercury, 552 ,, alloys of (amalgams), 555 , , deadening of, 555 ,, distillation of, 554 ,, oxides of, 556 Metal slag (copper), 597 Metallic nitroxyls, 219 662 Index Metalloids, 8 Metals and non-metals, 7 Metameric compounds, 171 Metantimonates, 458 Metaphosphates, 437, 459 Metarsenates, 449, 459 Metarsenites, 448 Metastan nates, 594 Metavanadates, 607 Meteoric iron, 150 Methane, 276 Meyer, Lothar, curve of atomic volumes, 106 Microcosinic salt, 498 Milk of lime, 534 ,, sulphur, 368 Milky quartz, 588 Mineral alkali, 467 Mineral chameleon, 621 Minium, 601 Mispickel, 439 Mixed crystals, 50 Modes of chemical action, 13 Molecular combinations, 65 , , equations, 54 ,, formulae, 22 , , heats, 47 ,, lowering of vapour pressure, 118 ,, volume, 42 „ weight, 39 ,, weight, determination of, by the depression of freezing- point, 121 Molecules, 3 ,, compound, 6 ,, definition of, 4 , , elementary, 6 ,, mean free path of, 83 „ size of, 3 Molybdates, 616 Molybdenite, 616 Molybdenum, 616 ,, chlorides, 617 ,, ochre, 616 ,, oxides, 616 Monad elements, 58 Mono-atomic molecules, 8 Monoclinic system, 139 Monosymmetric system, 139 Monovalent elements, 58 Mordants, 572 Mortar, 534 ,, the setting of, 534 Mosaic gold, 596 Mottramite, 607 Mundic, 439 Muntz metal, 509 Multiple proportions, law of, <6 Natukal waters, 194 Natural steel, 624 Nessler's solution, 241, 560 Neutral alum, 576 Nickel, 639 ,, alloys of, 639 ,, blende, 639 ,, carbonyl, 263 ,, chloride, 641 ,, glance, 439, 639 ,, monosulphide, 641 , , monoxide, 640 ,, oxides of, 640 ,, sesquioxide, 640 ,, silver, 548 ,, sulphate, 641 Nickelo-nickelic oxide, 640 Nickelous oxide, 640 ,, sulphide, 641 Niobates, 607 Niobium, 607 ,, oxides of, 607 Nitrates, 215 ,, detection of, 215 Nitre, 483 ,, plantations, 483 Nitric anhydride, 216 ,, oxide, 220 Nitrification, 483 Nitrites, 219 Nitro-cellulose, 215 Nitrogen, 205 ,, iodide, 249 ,, oxides and acids of, 209 ,, pentoxide, 216 , , peroxide, 217 , , tribromide, 249 ,, trichloride, 248 Nitro-metals, 219 Nitro-sulphuric acid, 368 Index 663 Nitrosyl chloride, 225 ,, hydrogen sulphate, 225 ,, sulphate, 388 Nitrous anhydride, 209 Nomenclature, 15 Non-metals, 7 “Nordhausen” acid, 395 Notation, chemical, 21, 52 Occluded hydrogen, 151 Occlusion of gases, 157 Olefiant gas, 278 Opal, 586 Ore hearth, 597 Orangeite, 581 Organic chemistry, definition of, 259 Orpiment, 439 Orthite, 561 Orthoclase, 590 Osmiridium, 642 Osmium, 642 ,, oxides of, 643 ,, tetroxide, 643 Osmotic pressure, 134 Osteolite, 538 Oxides, 16 Oxygen, 159 ,, allotropic, 172 ,, Brin’s process, 162 ,, Tessin du Motay process, 167 Oxyhcemoglobin, 170 Oxyhydrogen flame, 29 t O zone, 172 , , atmospheric, 234 ,, constitution of, 176 tube, Siemens’, 173 ,, ,, Andrews’, 17? Palladium, 642 ,, absorption of hydrogen by, 157 ,, chlorides, 643 ,, hydride, 157 ,, oxides, 643 Parkes’s process, 516 Partial pressures, law of, 127 Passive iron, 629 Pattinson’s process, 5x6 ,, white lead, 603 Pearl-ash, 4S2 Perchlorates, 338 Percy- Patera process, 5x8 Periclase, 529 Peridote, 590 Periodates, 356 Periodic classification, 102 Permanent white, 544 ,, hardness, 197 Permanganates, 621 Permanganic anhydride, 622 Persulphates, 385 Fersulphuric anhydride, 384 Petalite, 499 Petzite, 523 Pewter, 592 Phenacite, 528 Phlogiston, 283 Phosgene gas, 263 Phospham, 438 Phosphates, 435 Phosphine, 420 Phosphites, 434 Phosphonium bromide, 422 ,, chloride, 422 ,, iodide, 423 Phosphoretted hydrogen, gaseous, 42c ,, liquid, 423 ,, solid, 424 Phosphorous oxide, 430 Phosphorus, 412 ,, allotropic, 418 ,, compounds with sulphur, 438 ,, oxides and oxyacids, 429 ,, oxychloride, 428 ,, oxyfluoride, 428 ,, pentabromide, 427 , , pentachloride, 426 ,, pentafluoride, 425 ,, pentasulphide, 438 ,, pentoxide, 431 ,, red, 418 ,, tetriodide, 427 ,, tribromide, 427 ,, trichloride, 425 , , trifluoride, 424 ,, triodide, 427 Phosphoryl chloride, 428 ,, fluoride, 428 , , nitride, 438 ,, triamide, 438 664 Index Photo-salts, 522 Pig-boiling, 627 Pig iron, 626 Pitchblende, 6x6 Planes of symmetry, 139 Plaster of Paris, 537 Plastic sulphur, 367 Plate sulphate, 348 Platinamines, 647 Platinates, 645 Platinic hydroxide, 645 ,, chloride, 646 Platiniridium, 642 Plati no-chlorides, 646 ,, cyanides, 647 ,, nitrites, 647 Platinotype process, 646 Platinous chloride, 645 ,, hydroxide, 645 Platinum, 643 , , alloys, 645 ,, blade, 645 ,, oxides of, 645 ., oxysalts, 647 ., sodium chloride, 65 ,, spongy, 644 ,, sulphides of, 647 ,, tetrachloride, 646 Platoso-ammonia compounds, 6 Plumbago, 253 Plumbic chloride, 602 ,, oxalate, 600 ,, oxide, 600 ,, peroxide, 602 Plumbous oxide, 600 Plumbum nigrum , 596 Pollux, 470 Polybasite, 514 Polyhalite, 481 Polymerism, 17 1 Pot-ashes, 481 Potash, caustic, 476 Potassium, 471 ,, alum, 574 ,, aluminate, 572 ,, aluminium chloride, 576 ,, antimonate, 458 ,, borofluoride, 563, 567 „ bromate, 346, 480 „ bromide, 480 Potassium carbonate, 481 ,, chlorate, 477 i, chloride, 477 ,, chlorochromate, 616 ,, chloroplatinate, 646 ,, chloroplatinite, 646 ,, chromate, 614 1 1 chromium alum, 574, 613 ,, dichromate, 614 ,, ferrate, 631 ,, ferricyanide, 632 ,, ferrocyanide, 261, 63a ,, fluoride, 476 ,, fluor-plumbate, 310 ,, hydride, 475 ,, hydroxide, 476 I ,, hyponitrite, 225 „ iodate, 355 ,, iodide, 480 I ,, manganate, 621 | ,, metaborate, 565 ,, metantimonate, 458 ,, metarsenite, 448 ,, metastannate, 594 ,, meta-thio-arsenite, 450 ,, nitrate, 483 ,, nitrite, 220 , , osmate, 643 ,, oxides of, 475 ,, ortho-thio-antimonate, 460 ,, ortho-thio-antimonite, 460 ,, ortho-thio-arsenate, 451 ,, ortho-thio-arsenite, 450 ,, pentasulphide, 485 ,, pentathionate, 400 ,, perchlorate, 480 ,, periodate, 356 ,, permanganate, 622 ,, peroxide, 475 ,, platinic chloride, 646 ,, platino-cyanide, 647 ,, platinous chloride, 646 ,, plumbate, 602 , , pyro-antimonate, 458 ,, ruthenate, 643 , , silico-fluoride, 583 ,, silver thiosulphate, 398 ,, stannate, 593 ,, sulphate, 481 ,, sulphite, 362 Index 66 5 Potassium, sulphides of. 485 ,, tetrachromate, 615 ,, trichromate, 615 ,, zinc oxide, 154 Powder of Algaroth, 455 Praseo-cobaltic salts, 637 Preparing salt, 59.3 -Producer gas, 164 Proustite, 514 Prussian blue, 633 Pseudo-alms, 574 Pucherite, 607 Puddling, 627 Purple copper ore, 506 Purpureo-cobaltic salts, 630 Pyrargyrite, 514 Pyrites burners, 390 Pyrolusite, 618 Pyromorphite, 596 Pyrophosphates, 436 Pyrosulphuric chloride, 401 Quadratic system, 139 Quantitative notation, 52 Quartz, 586 Quicklime, 533 Radiated pyrites, 633 Radicals, compound, 22 Rain water, solid impurity in, 196 Realgar, 439 Red antimony, 451 ,, copper ore, 509 ,, haematite, 624 ,, lead, 601 ,, manganese oxide, 618 ,, phosphorus, 418 ,, zinc ore, 546 Refinery slag (copper), 507 Regular system, 139 Reiset’s second base, chloride of, 647 Relation of gases to heat, 68 , , ,, „ pressure, 70 Reversible reactions, 85 Rhodium, 642 Rhombic system, 139 Rochelle salt, 519 Rock crystal, 586 Rock salt, 486 Rodonda phosphates, 413 Roll sulphur, 364 Roman alum, 574 Roseo-cobaltic salts, 638 Rouge, 630 Rubidium, 500 Rubies, artificial, 571 Ruby, 569 Ruby ore, 506 ,, silver ore, 514 ,, sulphur, 439 Rust, 288 Ruthenium, 642 ,, chlorides of, 643 ,, oxides, 643 Rutile, 581 Sal alembroth, 558 ,, ammonia, 501 Salt-cake process, 491 Salt-forming oxides, 17 Salterns, 488 Saltpetre, 483 Salts, acid, 19 ,, basic, 19 „ haloid, 18 ,, normal, 18 ,, oxy-, 18 ,, thio-, 18 Sand, 582 Sapphire, 569 Satinspar, 536 Saturated solutions, 130 ,, vapours, 111 Scandium, 561 Scheele’s green, 447 Scheelinite, 616 Schlippe’s salt, 460 Schonite, 531 Schweinfurt green, 448 Scotch hearth, 597 Seaweed, iodine in, 347 Selenite, 536 Selenium, 405 ,, alums, 574 ,, di chloride, 407 ,, dioxide, 407 Selenuretted hydrogen, 406 Seltzer water, 195 Semipermeable membranes, 134 Serpentine, 528, 590. 666 Index Siemens’ ozone tube, 173 Silica, 586 Silicates, 589 Siliciuretted hydrogen, 584 Silicon, 582 ,, chloride, 586 ,, chloroform, 582 ,, dioxide, 586 ,, fluoride, 585 ,, hexachloride, 586 ,, hexafluoride, 585 ,, hydride, 584 Silver, 514 ,, allotropic, 519 ,, alloys, 519 alum, 522 bromide, 521 , , chloride, 520 ,, flashing of, 516 ,, fluoride, 521 ,, fulminating, 520 ,, glance, 514 ,, iodide, 521 , , nitrate, 522 ,, oxides, 519 ,, oxybromide, 522 ,, oxychloride, 522 ,, periodate, 355 ,, phosphates, 435, 437 ,, plating, 519 ,, spitting of, 518 ,, standards, 59 ,, suboxide, 520 ,, sulphate, 522 ,, sulphide, 514 Singly refracting crystals, 140 Slaked lime, 534 Smalt, 638 Smaltine, 634 Smoky quartz, 588 Soda, 496 Soda-ash, 495 ,, caustic, 487 ,, crystals, 495 Sodium, 486 ,, acetate, 277 ,, alloy with potassium, 486 ' ,, aluminate, 569 ,, aluminium chloride, 570 „ amalgam, 555 Sodium antimonate, 458 ,, antimonite, 457 arsenate, 449 ,, benzoate, 245 ,, bicarbonate, 496 ,, bromide, 490 ,, carbonate, 49c ,, chloride, 488 ,, chloro-platinate, 646 ,, hydrazoate, 245 ,, hydride, 486 ,, hydroxide, 487 ,, hypophosphite, 420 ,, hyposulphite, 386 ,, iodide, 490 ,, metaniobate, 607 ,, metaphosphate, 437 ,, metastannate, 594 ,, metatantalate, 607 ,, metavanadate, 607 ,, nitrate, 497 ,, oxalate, 155 ,, oxides, 486 ,, permanganate, 622 ,. phosphates, 498 ,, pyro-arsenate, 449 ,, pyrophosphate, 436 , , sesquicarbonate, 496 , , silicate, 588 ,, silver thiosulphate, 518 ,, stannite, 593 , , sulphate, 496 >> >> solubility curve, 132 ,, sulphide, 491 , , thio-antimonate, 460 ,, thiosulphate, 397 ,, tungstate, 616 , , , uranate, 616 zinc chloride, 65 Soffioni, 564 Solar prominences, 150 Solder, 592 Solfatara, 360 Solidification, suspended, t 18, 365, 4 16 Solidifying points of liquids, 119 , , points of liquids, effect of dis- solved substances upon, 121 ,, points of liquids, effect of pressure on, 119 Solubilities, diagram of, 131 Index 667 Solubility of gases in liquids, 122 ,, of liquids in liquids, 128 ,, of mixed gases, 126 ,, of solids in liquids, 129 Solution, 122 Solutions, saturated, 130 ,, supersaturated, 130 Sombrerite, 413, 538 Spathic iron ore, 624 Specific gravity of gases, 39 ,, liquids and solids, 104 ,, heat, 44 ,, heats, tables, 45 Spectra of alkali metals, 469 Spectroscope, 468 Specular iron ore, 624 Speiss-cobalt, 634 Spiegel, 626 Spinelle, 572 Spirits of hartshorn, 239 Spitting of silver, the, 518 Spodumene, 499 Spring water, 196- 1 99 Stalactites, 197 Stalagmites, 198 Standard temperature and\ pressure, 69, 70 Stannates, 594 Stannic chloride, 595 ,, sulphide, 595 Stannous chloride, 594 ,, hydrated oxide, 592 ,, nitrate, 592 A ,, oxide, 592 ,, oxychloride, 594 ,, sulphate, 591 ,, sulphide, 595 Stassfurt deposits, 471, 481, 530 Steam, 189 ,, volume composition of, 184 Steel, 627 Steel mill, 291 Stephanite, 514 Stereotype metal, 452 Stibnite, 451 Still-liquor, composition of, 319 Stream-tin, 590 Stromeyerite, 514 Strontia, 539 Strontianite, 539 Strontium, 539 , , chloride, 540 ,, dioxide, 540 ,, hydroxide, 540 ,, nitrate, 541 ,, oxides, 539 ,, sulphate, 540 Substitution, 344 Suint, 471 Sulphates, 395 Sulphides, 371 Sulphites, 382 Sulpho-acids, 17 Sulpho-thionyl chloride, 374 Sulphur, 359 , , allotropic modifications, 365 , , chlorides of, 374 ,, dioxide, 376 ,, flowers of, 363 , , milk of, 368 , , oxides and oxyacids of, 375 , , oxychlorides of, 400 ,, plastic, 367 , , prismatic, 365 , , recovery of, from alkali waste, 361 ,, recovery of (Chance’s process), 37 i , , rhombic, 365 , , sesquioxide, 384 , , tetrachloride, 375 , , trioxide, 382 Sulphuretted hydrogen, 369 Sulphuric anhydride, 375 ,, chlor hydrate, 401 Sulphurous anhydride, 375 Sulphuryl chloride, 400 j Supercooling of water, 118 Superphosphate of lime, 538 Supersaturated solutions, 130 Suspended solidification, 119, 365, 416 Sylvanite, 523 Sylvine, 471 Symbols, 21 Sympathetic inks, 193 Synthesis, 13 Tachydrite, 530 j Talc, 528 668 Tank liquor, 494 Tantalite, 607 Tantalum, 607 11 oxides of, 607 Tartar emetic, 457 Tellurates, 410 Telluretted hydrogen, 409 Tellurites, 410 Tellurium, 408 Temporary hardness, 197 Tenorite, 510 Tessid du Motay process, 167 Tetradymite, 409 Tetratomic molecules, 8 Tetravalent elements, 58 Thallic chloride, 579 ,, nitrate, 580 11 oxide, 562, 578 sulphate, 580 , , sulphide, 562 Thallium, 577 ,, oxides of, 578 ,, oxyhydroxide, 579 ,, perchlorate, 562 ,, sulphate, 562 Thallous carbonate, 579 ,, chloride, 579 ,, hydroxide, 578 ,, iodide, 562 ,, oxide, 578 n phosphate, 579 Thdnardite, 496 Thermochemistry, 142 Thio-acids, 17 Thio-antimonates, 460 Thio-antimonites, 460 Thio-arsenates, 451 Thio-arsenites, 450 Thiocarbonates, 404 Thionyl chloride, 400 Thiophosphoryl chloride, 429 Thiophosphoryl fluoride, 428 Thorite, 581 Thorium, 581 Tin, 590 ,, alloys of, 592 ,, dioxide, 593 ,, oxides of, 592 ,, oxymuriate, 595 Tin-plate, 592 Index Tin-stone, 590 Tin-white cobalt, 439 Tincal, 562 Tinning, 592 Titanium, 581 Tombac, 509 Transitional elements, 100, 623 Transpiration of gases, 82 Triad elements, 58 Triclinic system, 139 Tridymite, 586 Triethylamine, 129 Triethyl silico-formate, 583 Trivalent elements, 58 Trona, 496 Tungstates, 616 Tungsten, 616 1, chlorides, 617 1, oxides, 616 Turnbull’s blue, 632 Turpeth mineral, 395 Turquoise, 569 Type metal, 452 Typical elements, 100 Twin crystals, 588 Ulexite, 562 Uniaxial crystals, optically, 140 Unit of heat, 144, 289 1 1 volume, 42 Unsaturated compounds, 61 Uranates, 616 Uranium, 6x6 1, chlorides, 6x7 ,, oxides, 616 Uranous salts, 6x7 ,, sulphate, 617 Uranyl salts, 617 Urea, 13, 23, 257 Valency, 58 Vanadates, 608 Vanadite, 607 Vanadium, 607 ,, chlorides of, 608 ,, oxides of, 607 ,, oxychlorides of, 608 Vaporisation, latent heat of, 114 Vapour densities of elements, 40 „ pressures of solutions, 1 17 Index 66g Vapour tension, in Verdigris, 513 Vermilion, 559 Vinasse. 482 ,, cinder, 482 Vital force, 259 Vitriol chambers, 393 Volatile alkali, 467 Water, 180 Clark’s process for softening, 198 colour of, 188 compressibility of, 189 electrolysis of, 183 freezing of, 115 gas, 260 gravimetric composition of, 185 hardness of, 197 maximum density of, 191 of constitution, 194 of crystallisation, 192 rain, 196 solubility of gases in, 127 ,, salts in, 131 solvent power of, 191 supercooling of,> 118 volumetric composition of, 181 Waters, chalybeate, 195 dangerous, 199 deep well, 196 fresh, 196 hard, 197 mineral, 194 natural, 194 potable, 198 river, 196 safe, 199 sea, 196 spring, 196 suspicious, 199 Wavellite, 413 Weldon’s process, 319 Welsbach burner, 303 White arsenic, 446 ,, cast iron, 626 ,, lead, 604 ,, metal (copper), 507 ,, nickel, 639 ,, vitriol, 194 Witherite, 541 WoMerite, 581 Wolfram, 616 ,, ochre, 616 Wood’s fusible metal, 461 Wrought iron, 627 Wulfenite, 616 Wurtzite, 550 Xantho-cobaltic salts, 638 Ytterbite, 561 Ytterbium, 561 Yttrium, 561 Zier vogel process, 517 Zinc, 546 ,, alloys of, 548 ,, aluminate, 546 ,, amalgam, 556 ,, blende, 546 ,, carbonate, 551 ,, chloride, 549 ,, chromite, 614 ,, granulated, 153 ,, hydroxide, 549 ,, methyl, 277 V ,, nitrate, 215 „ oxide, 548 „ spar, 546 ,, spinnelle, 546 ,, sulphate, 550 ,, sulphide, 550 ,, white, 548 Zinci carbonas, 551 Zinc-copper couple, 152, 277 Zircon, 581 Zirconium, 581 Printed by BALLANTYNE, HANSON &• Co. 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