GIFT OF Bertha Stut . OO"^ VTL^Crt MODERN CHEMISTRY WITH ITS PRACTICAL APPLICATIONS BY FREDUS N. PETERS, A.M. INSTRUCTOR IN CHEMISTRY IN CENTRAL HIGH SCHOOL, KANSAS CITY, MISSOURI AUTHOR OF " EXPERIMENTAL CHEMISTRY," ETC. NEW YORK MAYNARD, MERRILL, & CO. 1905 PY3 COPYRIGHT, 1001, BY MAYNARD, MERRILL, & CO. 6/ff PREFACE IN preparing this book for use in secondary schools, I have endeavored to look at the science from the viewpoint of the students themselves. The fault with many texts upon the same subject is that the position of the learner has been disre- garded; the books have been encyclopaedic; they have pre- sented a great number of facts as a skeleton or framework, but this skeleton has not been clothed with muscle and ani- mated with life. No more fascinating subject finds a place in the curricula of our secondary schools, yet to the average student chemistry is too often but an irksome task. In the present work I have omitted much that is often given in an elementary text, while at the same time entering more into that detail which gives lively interest to the subject. It has been my aim to show, whenever possible, the practical application of the science to the everyday affairs of life; in other words, to emphasize industrial and commercial chem- istry. At the same time the fundamental principles of the science have not been forgotten ; on the contrary, they have been emphasized even more than is usual in an elementary chemistry. This has been rendered possible by the omission of much that can never be either of interest or of value to the beginner. Recognizing the fact that science must be taught inductively by experiment, some authors have assumed that the student must gain all his knowledge of chemistry in this way. No greater mistake could be made. The science has been hun- dreds of years in reaching its present development, and much must be accepted by the student without any effort to work it out for himself. In this text a large amount of experi- M 1832 4 PREFACE mental work is given, sufficient to meet the requirements of all our best colleges. The experiment is always supplemented by notes and suggestions which enable the student to draw proper conclusions, and give him such information as he cannot hope in his limited time to gain for himself. It will be noticed also that the laboratory directions are largely in the form of questions, so as to compel even the least ener- getic students to secure the benefits of personal investigation. This plan is always followed except in cases where the student would be in danger of going astray. To the pedagogical treatment of the difficult parts of the sci- ence I desire to call attention. The subject has been presented much in the same way as in my own classes, where the method has met with success. Beginning with the study of that most familiar of substances, Water, the text enters into a discussion of its composition and then proceeds to a detailed statement of its constituents. This work is prefaced merely by a short chapter connecting the science of Chemistry with that of Physics and by one chapter upon Valence. This some- what intricate subject of Valence was introduced early by request of a number of teachers, to avoid many difficult questions that must arise when it is deferred. Logically, it would come later, and may be deferred if the teacher so desires. However, its simple, graphic presentation is such that the student can hardly fail to grasp the meaning. I recognize the demands, coming from all higher institutions of learning, for more quantitative work, and believe I have fully met all such requirements. The student of the subject, as taught hitherto, has been in danger of coming to the con- clusion that very little in chemistry is exact ; whereas nothing could be further from the truth. The pupil is shown this by the series of quantitative experiments which have been carefully worked out in the laboratory. For this work I have sought to make use of such apparatus only as may be or should be found in any secondary school. PREFACE 5 To the regular text is appended detailed instruction for various laboratory manipulations, preparing solutions, etc. A chapter has also been added for the benefit of any who may desire to continue along qualitative lines the work introduced at various places in the text. I desire to acknowledge, with gratitude, the valuable sug- gestions offered by Professor Irving P. Bishop, of the State Normal School, Buffalo, and Professor M. D. Sohon, of the Boys' High School, New York City, both of whom have read the book critically in manuscript ; furthermore, I wish to say that to many of my students of the past I am indebted for descriptions, original with them, and more appropriate than any I have found in any text. To Dr. Paul Schweitzer, for many years Professor of Chemistry in the Missouri State Uni- versity, the true friend of the student, to whom I owe much for his great sympathy and encouragement, and his words of fatherly advice, I desire to express especial gratitude. Finally, I acknowledge with pleasure the help and inspiration I have gained in my private study and research from the writings of those who have been permitted to drink long and deep from this fountain of science. TO THE TEACHER IT is not expected that everything given in the text will be demanded of the pupil, unless possibly in reviews. Some of the manufacturing processes, for example, I have deemed of sufficient importance and interest to be given; yet it may seem best in the judgment of the instructor to omit these. The experiments, as a rule, may be performed by the stu- dents, and apparatus is suggested which most schools will be able to provide. The number of pupils will determine to some extent what experiments should be performed by the teacher and what by the students themselves. If the classes are small, so that the teacher can give very close personal attention, the pupils may attempt almost all; on the other hand, if the classes are large, it may be necessary for the teacher to perform some of the more difficult and dangerous ex- periments himself. On pages 355 to 380 whi be found many useful suggestions to the student for the care and manipula- tion of apparatus, making up of solutions, etc. These should be read before the student begins his work in the laboratory, and frequent reference should afterward be made to them. It is presumed that a school year of nine or 'ten months will be given to the work in this text, but by omitting some of the less important elements, much of the theory and many of the practical applications of chemical science may be obtained in five or six months. Besides the various chapters devoted to the fundamental laws of chemistry, special study should be given to the following elements and a few of their important compounds : hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, iodine, carbon, sulphur, sodium, calcium, zinc, lead, and iron. 6 MODERN CHEMISTRY CHAPTER I INTRODUCTION V. 1. With what is Chemistry concerned ? Nature pre- sents herself in ever-changing forms, and to one who is not familiar with these variations she is a mystery. The untaught inhabitant of the tropics, who has never been beyond the confines of his native state, taken to a colder climate would see no relation between the snowflake or the icy covering of our northern rivers and the rain-drop as it falls upon his native hills. To him they are entirely different substances. 2. So the diamond, the filling of the ordinary " lead " pencil, and the coal that we burn in our furnaces seem altogether dissimilar, and yet they are practically the same thing. Likewise the emery with which the seamstress sharpens her needle and the mechanic his tools, and such valuable stones as the oriental emerald and the ruby, though seemingly so different, have really the same composition. The purpose of the science of chemistry is the investiga- tion of the objects that lie all about us in nature, the study of their composition and of their relations to one another, the explanation of the various phenomena in connection with them, and the ability to apply this knowledge to practical uses. 3. Importance of the Subject. A knowledge of chemis- try adds a charm to many of the common things of life, 7 8 MODERN CHEMIST UY clothing them with new beauty. Later it will be noticed that the science of chemistry enters into all or nearly all of the great manufacturing industries of the world, and that without the application of its laws and principles all such enterprises would result in failure. Whether studied, tiiei'ef ore, ' 1 j atoms of another element taken as a standard. This standard, primarily, as shown above, * If in the judgment of the teacher the subject of valence can be more easily grasped by the student later in the course, this chapter may be deferred until after the study of carbon and its compounds. 21 22 MODERN CHEMISTRY is hydrogen, and by it the valence of other elements is measured or determined. It may be illustrated in this way: suppose the first line represents the combining power of hydrogen, which is our standard. Then with this " yard stick " we will measure the combining power of the other elements. In water, H 2 O, the valence of the oxygen atom is determined by applying the "yard stick," and is seen to be two ; in NH 3 the standard is used three times, and the valence of the nitrogen atom is three. In the same manner the valence of the carbon atom is determined as four. 3. Suppose, however, hydrogen did not combine with carbon, could we still determine its valence ? We are familiar with the compound, carbon dioxide. In this molecule, CO 2 , the oxygen atom is used twice with the carbon atom, hence the latter must have twice the combin- ing power of the oxygen. This has already been shown to be two, hence carbon would be four. To illustrate roughly, we sometimes speak of the atoms as having a certain number of "bonds" or "poles of attraction," represented as below : @p Atoms with one " bond." Atoms with two " bonds." Atoms with three "bonds." Atoms with four "bonds." From this illustration it will be seen that an atom in the second group in combining would have two bonds by VALENCE 23 which to hold the two bonds of two individual atoms of the first. Mercuric Bromide. Water. N and Sb both have valence (&) (ri^ of three. Ammonia. Antimony Chloride. C and Si, valence of four. Silicon Dioxide. 4. Such atoms as combine with one of hydrogen or its equivalent are said to be univalent, or are sometimes called monads; those which combine with two of hydrogen are bivalent, or dyads; with three, trivalent, or triads ; with four, quadrivalent, or tetrads; with five, quinquivalent, or pentads. 5. Variation in Valence. In studying a number of the compounds of any element it will be noticed that while the valence of the element in most of them is the same, there will be some compounds which show it to be dif- ferent. Many of these are believed to be merely apparent exceptions and may be readily explained ; while others, as yet not thoroughly understood, may be real variations. For example, the oxygen atom is always regarded as biva- lent, yet we shall meet with the compound, hydrogen peroxide, H 2 O 3 , in which oxygen is apparently univalent. 24 MODERN CHEMISTRY It is believed, however, that the atoms have an arrange- ment in the molecule which may be represented thus : or This simply means that one bond of each atom of oxygen is held by a bond of the other. In a similar way the apparent double valence of a great many other ele- ments is explained. Thus copper and mercury, ordinarily bivalent, also form the compounds Cu 2 Cl 2 , and Hg 2 Cl 2 . But these are exactly parallel to the case above. Bivalent. Apparently Bivalent. Apparently Univalent. Univalent. 6. Again we shall study the two compounds of carbon, CO and CO 2 , the first of which would indicate a valence of two for the atom, while in the second it would be four. The second is believed to show the true valence, and carbon monoxide is regarded as an unsaturated compound, that is, one in which the valence of the atom is not satisfied, or one in which a part of the bonds is not held by any other element. We may represent it thus : Saturated Unsaturated Compound. Compound. This theory is accepted because carbon monoxide very readily takes up one more atom of oxygen and forms the dioxide. VALENCE 25 7. Double Valence. In the examples of double valence, noticed above, the irregularity is only apparent. There are many cases, however, in which all the indications thus far would show that the valence of the atom is variable. Thus we have said the nitrogen atom is trivalent, and this is so in ammonia and nitrogen trioxide ; but we shall also meet with nitrogen pentoxide, N 2 O 5 , in which the valence is five ; the monoxide, N 2 O, wherein it is appar- ently one, etc. There are many such variations that will trouble the student, but for our present work we shall need, as a rule, to aid us in writing formulae and reactions, only a knowledge of the ordinary valence of the atoms. 8. Valence of Groups or Radicals. We shall find also that many groups of atoms react in the game way as individual atoms ; such groups are called radicals. They have combining power or valence just as individual atoms have. Thus when sulphuric acid reacts with zinc, we shall find that the group (SO 4 ) is not broken up ; the same is true in hundreds of other instances. As it is combined with two atoms of hydrogen in H 2 SO 4 , and always does combine in the same way, we say its valence is two ; this may be shown graphically in this way. (SO 4 ) Group, showing Sulphuric Acid, showing all bonds two bonds unused. of the (SO 4 ) group saturated. While we cannot prove that such is the arrangement of the atoms in the group, still it is believed to be true ; at any rate it serves to illustrate how the valence of the group is two. In the same way we would show the valence of any other radical. In sal ammoniac, NH^Cl, we 26 MODERN CHEMISTRY find the group (NH 4 ) in combination with one atom of chlorine, hence its valence is one. Then if the radical (NH 4 ) combines with (SO 4 ), it must be used in the proportion of two of the former to one of the latter, thus (NH 4 ) 2 S0 4 . Cl] Sulphuric Acid. Ammonium Ammonium Sulphate. Or we may show the same facts in this way : Ammonium, showing Ammonium Sulphate, showing how (SO 4 ) one bond unused. can unite with two of (NH 4 ). Exercise in Valence.* Applying the principles set forth in the pre- ceding paragraphs, let the student write the formulae for the fol- lowing : when Ba unites with I, Cl, Br, O, NO 3 , Cr0 4 . when Na unites with O, S, Cl, CIO,, SiO 4 , SO 4 . when Cu unites with Cl, S, SO 4 , IIO, O, I. when NH 4 unites with T, PO 4 , SO 4 , S, HO, Br. when Bi unites with Cl, O, S, NO 3 , SO 4 . * Let the teacher give further exercises until the student can write with assurance the compound resulting from the union of any two of these elements or radicals. VALENCE 27 The valence of each of the above and certain others is shown below : MONADS MONADS DYADS DYADS TRIADS TKTEADS I Li Ba Cu Sb C Br (NH 4 ) Zn Fe Bi Si Cl (NO S ) Ca S As (SiOJ F (CIO,) Sr (S0 4 ) (PO^ Na (HO) O (Cr0 4 ) K Hg lonization. Closely related to the idea of valence is that of ionization. Attempts to pass a current of elec- tricity through a vessel of pure water meet with very indifferent success. Again, if the electrodes of a battery be placed in a cup of dry salt, no current is transmitted; however, if a little of the salt be dissolved in the pure water, the solution becomes a good conductor. 10. Explanation. It is well known to students of physics that water containing a solid in solution boils at a higher temperature than pure water. For example, water saturated with common salt boils at 108 C. ; with potassium nitrate, 116; with calcium chloride, 179. Numerous experiments have shown that this rise of boil- ing point is proportional to the amount of the substance dissolved ; furthermore, to secure a like change in the lowering of the vapor tension when different substances are dissolved, it is found that amounts proportional to the molecular weights of the substances must be used. (See page 68.) To illustrate, suppose the molecular weight of the compound A is 342 and of B 46, then to secure like results in lowering of vapor tension, we should be required to dissolve portions of the salts in the ratio of 342 to 46. 11. Exceptions. When, however, we dissolve such substances as common salt, NaCl, we find a lowering of the vapor tension about double what we should expect, 28 MODERN CHEMISTRY and with calcium chloride, CaCl 2 , about three times. Furthermore, such solutions are good conductors of elec- tricity, while the others are not. 12. Conclusion. Putting all the facts together, we are led to believe that those substances which lower the vapor tension abnormally, when dissolved in water, are broken up into parts or ions. Thus, common salt becomes largely ionized into sodium and chlorine ions, and calcium chlo- ride into calcium and chlorine ions, etc. 13. Negative and Positive Ions. If a V-shaped tube is filled with a dilute solution of common salt and a cur- rent of electricity passed through it, the sodium ions will tend to collect at one electrode and the chlorine at the other. This may be proved by adding a few drops of a solution of phenolphthalein. It will be found that the sodium ions collect at the negative electrode and the chlorine at the positive. From the well-known law of electrical attraction, we know, therefore, that the sodium ions are positive and the chlorine are negative. In all such compounds we find the two kinds of ions, the nega- tive or anions neutralizing the positive or kathions. 14. Relation to Valence. From this we see that atoms having a valence of two, as calcium, for example, would carry double the electricity that an atom like chlorine would with a valence of one. In other words, in calcium chloride, CaCl 2 , the calcium ion has a positive charge equal to the negative charge carried by the two ions of chlorine. SUMMARY OF CHAPTER Valence Meaning of term. Illustrations. Classification of elements as to valence. Synonymous terms for univalent, etc. Variation in valence. WATER 29 Unsaturated compounds. Real Illustrations of. Valence of radicals. Illustrations. Application of a knowledge of valence. Writing formulae of compounds. Relation to ionic theory. CHAPTER III WATER : H 2 1. Its Abundance. Water is one of the most abundant substances known ; it covers about three-fourths of the surface of the earth, besides existing in vast quantities in other forms. In the arctic regions in the form of an ocean of compressed snow it covers the entire surface of the land to a depth of many feet; in a similar form it caps all the loftiest mountain peaks from which great rivers of ice flow down the valleys until they are melted at the snow- line. In the form of vapor it exists in the atmosphere, invisible except when condensed in fogs, clouds, etc. In any given locality this moisture in the air varies largely at different times, but not often is there more than sixty- five per cent of what the air is able to hold. Even with this amount, however, it has been estimated that were the vapor in the air all condensed, it would form over the surface of the entire earth a layer of water five inches deep. 2. The human body is about sixty per cent water, and daily throws off through the pores and from the lungs over three pounds of moisture. Many vegetable articles of food contain eighty to ninety per cent of water, and some even more. 30 MODERN CHEMISTRY \\ 3. Water of Crystallization. Water also exists in an- other form not so familiar as those already mentioned; that is, water of crystallization. A great many compounds in solidifying from their aqueous solutions take up a con- siderable amount of water. This does not exist in a free state like water in the pores of a sponge, or in a piece of soft wood that has been submerged for some time, but is in combination crystallized in with the molecules them- selves. Such substances in crystallizing cause the disap- pearance of a considerable amount of water, which may, however, usually be obtained again by subjecting the body to a greater or less degree of heat. Some astronomers even believe that the absence of water upon the moon may be accounted for by the fact that such bodies as those mentioned have taken it all up in crystallizing from their aqueous solutions. An idea of the amount contained by such substances may be gained from the following experi- ments. EXPERIMENT 12. Put into a test-tube a crystal of native gypsum and heat in the Bunsen flame. What do you see depositing upon the cooler portions of the tube ? How is the crystal of gypsum affected ? Repeat the experiment, using borax or alum instead of gypsum, and state results. EXPERIMENT 13. Expose to the air for several hours a crystal of ferrous sulphate. Notice its appearance before the exposure; how has it changed in the air ? 4. Efflorescent Substances. Many such substances as x ferrous sulphate and copper sulphate, upon being exposed to an atmosphere more or less dry, give up all or part of their water of crystallization ; at the same time they usually change in color and crumble to a powder. The process is the same as when the substances are heated, but not so rapid. By adding water to them the color is WATER 31 usually restored, and they crystallize as before. Such substances as these that give up their water of crystalliza- tion to the air are said to be efflorescent. *VJ. Deliquescent Substances. There is another class of substances which have the power of abstracting moisture from the air or surrounding bodies, and of dissolving them- selves either in whole or in part in this moisture. Such are called deliquescent bodies. A familiar example of these is a substance commonly sold by grocers under the name of " lye," which on being exposed to the air rapidly takes up moisture. Another noted example is phospho- rus pentoxide, a white solid formed when phosphorus is burned in the air or oxygen ; also calcium chloride and caustic potash. EXPERIMENT 14. Put into a dry evaporating dish or beaker a small lump of fused calcium chloride and allow it to stand several hours or over night. Notice how it has changed. In the same way expose a small piece of caustic potash. Notice how rapidly it changes. Only a few minutes will be necessary in this case. Common salt, stick candy, and some forms of taffy are very familiar examples of deliquescent bodies. 6. Distinguishing Characteristics of Water. In the pure state, water is practically colorless, but when of great depth it is seen to be of a blue color. It is odorless and tasteless, but we are so accustomed to drinking impure water that when we use that which is distilled, or perfectly pure, it tastes- "flat," just as unseasoned food does to those who are habituated to the use of salt, pepper, and other condiments. Pure water, on being evaporated to dryness, leaves no residue whatever, and this, in connection with the fact that it affects vege- table coloring matter in no way, is one method of test- ing it. 32 MODERN CHEMISTRY , Solvent Powers of Water. Pure water is seldom found, owing to its great solvent powers. To a greater or less extent it may be said to be almost a universal solvent. Even glass and similar substances immersed in water show appreciable loss after a considerable length of time. From this property result the various kinds of " hard " or mineral waters, medicinal, saline, etc. It is owing to the solvent powers of water, and the fact that evaporation leaves all mineral matter behind, that the ocean contains such vast quantities of different kinds of salt. 8. For example, in a hundred pounds of sea water, there are over three pounds of solid matter ; of this the greater portion is common salt, but compounds of magnesium and calcium in the form of what are usually known as epsom salts and gypsum also occur. It has been estimated that if the ocean were of an average depth of one thousand feet, the common salt in solution would occupy a space of about three and a half million cubic miles, or a volume more than five times as great as that of the Alps. On this basis, if the depth of the ocean averages what is now claimed for it, the amount of salt surpasses in bulk our greatest mountain ranges. Qj>" Composition of Water. By many of the ancients, water, along with fire and air, was regarded as an element; but about 1800 A.D. it was proved to be a compound body. There are two methods of proof, which taken together are quite conclusive. The first proof is by Electrolysis. EXPERIMENT 15. Fill the tubes of the electrolytic apparatus shown in Fig. 3 with water slightly acidulated with sulphuric acid, the latter being added simply to increase the conductivity. Then connect the platinum electrodes with a strong battery. As the WATER 33 current passes through the water, bubbles of gas will be seen rising from the two strips of platinum, P, and from one of them considerably faster than from the other. It will be found that twice as much gas collects in one tube as in the other. These two gases, we shall learn before long, are hydro- gen and oxygen. Open the stop-cock of the tube containing the greater amount of gas, and hold a lighted match to it ; notice that it burns with a very pale flame. Test the gas in tho other tube, using a pine splinter with a spark upon the end; notice that it bursts into a flame. FIG. 3. xtt). The second proof is by Synthesis. EXPERIMENT 16. Put into the eudiometer, Fig. 4, 8 cc. of oxy- gen, and twice as much or more hydrogen, the instrument being already partly filled with mercury, M. Hold the thumb over the open end and pass a spark from a galvanic battery. * The two gases will combine with ex- plosive force, producing water in the form of vapor. If the proportions were exactly two of hydrogen to one of oxy- gen, when the apparatus has become cool, there will be no gaseous residue, showing that the two unite in this pro- portion to form water. A in the figure is a cushion of air left to break the force of the explosion. FIG. 4. 11. Conclusions. From the above experiments we learn that water is the result of the union of two invisible gases, one of which burns with a pale flame, the other of which causes various substances to burn vigorously. We see also that from the union of the two gases, which to- gether form a very explosive mixture, there results an 34 MODERN CHEMISTRY exceedingly stable compound, which not only does not burn, but which has the power of quenching thirst and of overcoming the greatest fires. These two gases were given the names, hydrogen and oxygen. 12. We noticed also in the proof by analysis, that the hydrogen was given off in volume double that of the "oxygen, and further, that in mixing the two gases for the synthetic proof we caused them to unite in the same ratio. From theso experiments we may conclude that the com- position of water by volume is two parts of hydrogen to one of oxygen, a fact which we represent by the expression H 2 0. 13. Analysis by Other Methods. The analysis of water may be effected by means other than electricity. For example, if a current of steam is made to pass through a tube containing charcoal or coke heated red hot, the steam is decomposed ; the oxygen combines with the carbon of the charcoal, forming an oxide of carbon, and at the same time the hydrogen is set free. 14. Synthesis by Other Methods. In a similar way the synthesis of water may be effected. If a current of hydrogen is passed through a tube containing some me- tallic oxide, heated to redness, for example, copper oxide, the oxygen is removed from the compound by means of the hydrogen, and water is formed and may be collected. EXPERIMENT 17. Into a of the small bulb-tube put a little black oxide of copper, and weigh both tube and oxide care- fully. Next fill a U-shaped tube with lumps of calcium chloride, weigh and quickly connect with the other tube. Now pass a current of hy- FIG. 5. drogen, generated as on WATER 35 page 39, over the copper oxide, heated to redness. The hydrogen should first be dried by passing through sulphuric acid or over calcium chloride. After some time, disconnect the apparatus, and weigh the U-tube ; the gain in weight will represent the amount of water pro- duced. When the bulb-tube is cool, weigh it : the loss will represent the amount of oxygen removed. Subtracting the weight of the oxygen from the weight of water found will give the amount of hydrogen. Allowing for errors, this should give eight parts of oxygen to one of hydrogen, by weight. From this experiment we are able to conclude as to the quantita- tive composition of water, just as by the others we learned of the volumetric. The action of hydrogen in thus removing oxygen from an oxide is called reduction. Water I By volume: H J dro g en 2 5 Oxygen, 1. ( By weight : Hydrogen, 1 ; Oxygen, 8. SUMMARY OF CHAPTER Water Various forms in which it occurs. Water of crystallization. Meaning of term. Examples of substances containing it. Proof of its presence by experiment. Efflorescent substances. Deliquescence Meaning of term. Illustrations. Some special characteristics of water. Composition of water Proof of. a. By analysis Details of work. Apparatus used. b. By synthesis Explanation of process. Drawing of apparatus. Composition by weight. Proof by experiments. y CHAPTER IV ^ r HYDROGEN : H = 1 1. History. The term hydrogen is from two Greek words, which mean water producer, and the gas is so named because this element enters so largely into the composition of water. It was first isolated in quantities sufficient for experiment by Cavendish in 1766, and on account of its combustibility was called by him inflam- mable air. 2. Where found. Hydrogen is seldom found uncom- bined, though its chemical affinity for most substances is not very marked. It exists abundantly in composition water being the most important example ; it enters into nearly all organic compounds ; it is given off, together with other gases, by some volcanoes ; and by the spectro- scope we know that it exists in the atmospheres of the sun and of some of the stars. 3. Methods of obtaining Hydrogen. We have seen already in Experiment 15 that hydrogen may be ob- tained by the electrolysis of water. This gives a very pure gas, but does not produce it rapidly enough for ordinary experimental purposes. Just as electricity has the power of decomposing water, so do certain metals. When iron is exposed to moisture, we say it rusts; in reality, it takes up a certain amount of oxygen from the water and sets free a corresponding amount of hydrogen.* * Rust is an oxide of iron ; that is, a compound of iron and oxy- gen, represented by the formula Fe 2 3 . The chemical change which 36 HYDROGEN 37 4. Decomposition of Water. Again, there are some metals, like calcium, a constituent of common limestone, which have the power of decomposing water at the boiling point, setting free a part of the hydrogen and forming at the same time a compound, such as lime water. The chemical action may be expressed by the following equation : Ca + 2 H 2 = Ca (OH) 2 + H 2 . 5. There are two common metals, sodium and potas- sium, which decompose water rapidly at ordinary tempera- tures. Of these, the second acts much more violently, generating almost instantly sufficient heat to ignite the hydrogen given off from the water and volatilizing a por- tion of the metal itself. This is seen in the violet color which is imparted to the flame. That sodium is setting free a combustible gas in the same way may be shown by bringing a lighted match close to the metal, when the hydrogen will be ignited as it was spontaneously with potassium. EXPERIMENT 18. Fill a test-tube with water and invert it over a trough or basin of water, as shown in Fig. 6. Put into a wire gauze spoon, or wrap in a piece of flexible wire FIG. 6. probably takes place when iron is thus exposed may be expressed as 2 Fe + 6 H 2 = 6 H + Fe 2 O 3 , 3 H 2 O, in which Fe 2 O 3 is rust. Likewise, if iron filings be heated red hot, and a current of steam slowly passed over them, the filings take up the oxygen from the steam and are converted into an oxide of iron, Fe 3 04, differing somewhat from rust, while hydrogen is set free. The following equation expresses the chemical changes that take place : 3 Fe + 4 H 2 O = Fe 3 O 4 + 4 H 2 . 38 MODERN CHEMISTRY cloth, a small piece of sodium and hold under the mouth of the tube. Bubbles of gas will rapidly form, will rise into the tube and displace the water. Test the gas obtained to see whether it acts as did the hydrogen obtained by electrolysis in Experiment 15. Does it seem to be the same kind of gas? Sometimes, before putting the sodium into water, it is treated with a small quantity of mercury, whereby the rapidity of the action is greatly decreased. 6. Caustic Soda. The chemical change which has taken place in the above experiment may be expressed as follows : or, as it is usually and most simply written, H 2 + Na = NaOH + H. The graphic equation above shows that in each molecule of water one atom of the hydrogen is replaced by one of sodium, represented by Na, and that thus a new compound, NaOH, called caustic soda or sodium hydroxide, is formed. In other words, the water molecule becomes one of caustic soda, thus : becomes 7. Proof of the Above. The equation written above is not a matter of theory, but is determined by experiment. Add a drop of phenolphthalein to the water in which the sodium was placed. In the same way, test a solution of HYDROGEN 39 caustic soda ; also a little pure water in another tube. What are the results? 8. Other Methods of obtaining Hydrogen. The above methods of obtaining hydrogen, while of interest, are too expensive where considerable quantities are needed for experimental work. All acids contain hydrogen, and just as some metals decompose water, so certain others act with acids. 9. Laboratory Method. In obtaining hydrogen for laboratory purposes this is the plan usually pursued. The metal used is generally iron or zinc, and the acid, sulphuric or hydrochloric. EXPERIMENT 19. Fit to a flask a cork doubly perforated. Through one of the holes insert a delivery tube, and through the other a this- tle tube which extends nearly to the bottom of the flask. Put into the flask several pieces of granulated zinc made by pouring the metal in a molten condition into cold water. Add water until the zinc is nearly covered, and then pour in slowly a small quan- tity of strong sulphuric acid. After allowing the first gas which comes over to escape, because it is mixed with air, collect several bottles over water as described on page 362 and preserve for experiments a little later. The action may be hastened by adding a little copper sulphate to the flask a few minutes before the acid is introduced. 10. The Chemical Reaction. It will be noticed that the zinc in the above experiment gradually disappears. We shall find by testing the gas which is evolved that FIG. 7. Hydrogen Apparatus. p = pan. g = generating flask. t = thistle tube. r = receiving flask. d = deli very tube. 40 MODERN CHEMISTRY it is hydrogen. But what remains in the flask ? When the action has ceased, decant or filter off the clear solu- tion from any pieces of zinc or sediment, evaporate over a cup or beaker of boiling water nearly to dryness, and allow to cool. Drain off from the crystals any liquid remaining, rinse with a little cold water, and dry them. To prove that they are zinc sulphate, see pages 181 and 338. We may then express the changes thus : or Zn + H 2 S0 4 = ZnSO 4 + H 2 . Zinc sulphate is a white compound which collects upon the zinc, and would soon cover it so completely as to stop the chemical action ; the water, however, being added dis- solves it as fast as formed, and leaves a clean surface exposed to the acid. 11. Method of obtaining Large Quantities. When hydro- gen is desired in very large quantities, as in filling balloons, iron, being cheaper, is used instead of zinc. The gas thus obtained is somewhat less pure, but it is not on this account specially objectionable. Large vessels or retorts are used, which are lined with lead, a metal which is not affected by dilute sulphuric acid. The hydrogen obtained is passed through water and lime to purify it, after which it is transferred to the balloon. 12. Mond's Method. This method, although it has thus far been used only to a limited extent, promises to give satisfaction. We have seen that when steam is HYDEOGEN 41 passed over red-hot charcoal the former is decomposed just as when passed over red-hot iron (page 37) and two similar products are formed, both gases ; thus, H 2 O + C = H 2 + CO. The apparatus may be represented conventionally as follows : FIG. 8. Mond's Method. F is a furnace, B the boiler in which the steam is generated, C a tube containing lumps of coke which are heated red hot by a gas furnace beneath, Ni a tube containing powdered metallic nickel, L an apartment containing lime water. The steam passing through C is decomposed as stated above ; the mixture of hydrogen and carbon monoxide formed here passes over the nickel, also heated red hot. In this tube a part of the carbon unites with the nickel, and carbon dioxide is formed. This mixed with the hydrogen passes on through the " washer " L, containing lime water, which absorbs the carbon dioxide, leaving the hydrogen comparatively pure. The reactions hi the different parts of the process may be shown as follows : H 2 O + C = H 2 + CO (in the coke tube). H 2 + 2 CO -f Ni = NiC + H 2 + CO 2 (in the nickel tube). H 2 + CO 2 + Ca(OH) 2 = H 2 + CaCO 3 + H 2 O (in the " washer "). The nickel carbide formed is readily converted back again into metallic nickel by heating in the air, so that it may be used over and over, 42 MODERN CHEMISTRY 13. Characteristics of Hydrogen. The following experi- ments will illustrate well the most striking peculiarities of hydrogen : EXPERIMENT 20. Remove one of the bottles of hydrogen from the water, keeping it inverted, and thrust up into it a burning candle. Notice whether the candle continues to burn in the gas ; notice also what happens as you remove it again. Can you see anything burning at the mouth of the bottle? EXPERIMENT 21. To show the lightness of hydrogen. Bring a bottle of the gas, a, mouth downward, up close to another inverted bottle, t, of about the same size. Then gradually tip the hydrogen bottle, a, as shown in Fig. 9, just as you would pour water from one vessel into another, only in a reverse order. Now test both bottles to learn which con- FIG. 9. - Upward Decantation. hydrogen> gtate the results< EXPERIMENT 22. Start the generator again, and replace the de- livery tube with one which has been drawn to a fine jet. Let the gas flow a few minutes until the air is all expelled, and then ignite it. 14. CAUTION. A mixture of air and hydrogen is very explosive, and before lighting the jet a towel should be wrapped about the generating flask. It will do equally well to inclose the flask in a pasteboard box as shown in the figure. When first lighted, how does the hydrogen burn ? How does it soon change? This is due to the sodium in the glass, which colors the flame. A burning jet of hydrogen is some- times called the " philosopher's lamp." Does the gas burn with much heat ? Hold a clean dry bottle or test-tube over the flame. Do you see any deposit forming upon the upper part of the tube? What is it? Now try several tubes and bot- tles of different sizes in the same way ; notice the different pitch of the "singing tones" produced. FIG 10 Hvdroaen When the tube is thus sounding, notice the flame Jet m Bo*. carefully. Can you explain the tones produced ? HYDROGEN 43 EXPERIMENT 23. Allow a jet of hydrogen from a generator work- ing rapidly to strike against a platinum sponge. State the results. EXPERIMENT 24. The hydrogen pistol shows the explosiveness of a mixture of air and hydrogen.* Load the pistol by pouring into it a small bottle of hydrogen, as shown in Experiment 21, and fire by bringing a flame to the touch-hole. A loud explosion should follow. EXPERIMENT 25. Hydrogen soap-bubbles to show lightness of hydrogen. For success in this experiment a good soap solution is necessary. To a little soft water add a few shavings of castile or other good soap, and when dissolved add about one-third as much glycerine as soap solution. Shake well. Now attach to a delivery tube, from which is flowing a current of hydrogen, a clay pipe, or even an ordinary spool ; dip into the soap solution, and let the bubble form in the usual way. Detach from the pipe by a gentle jerk and notice whether the bubble rises or falls. Touch a light to one of the bubbles. What happens? This experiment is sometimes made more striking by filling the bubble with a mixture of hydrogen and oxygen, which, when touched with a flame, explodes violently. 15. CAUTION. The greatest care must be taken to avoid bringing the flame near the delivery tube, lest the whole mixture be exploded with serious results. EXPERIMENT 26, To show the presence of hydrogen in oils, alco- hol, etc. We have already seen that when hydrogen burns, water is produced. This is true whether we have hydrogen free, or in the form of compounds. Light a small candle and hold over it a cold beaker. Notice the water condensing upon the cooler portions of the beaker. In the same way try a small spirit lamp. State the results. 16. Conclusions from our Work with Hydrogen. By the above work with hydrogen we have learned that it is a colorless gas ; is without odor if pure, and very light. * The pistol may easily be made from a small tin can. With an awl punch a hole in the side of the can near the bottom, and for a bullet use a cork snugly fitted to the mouth of the can. When a light is brought to the touch-hole, several seconds may elapse before the explosion follows, but the experiment almost invariably succeeds. 44 MODERN CHEMISTRY Its density is but little more than one-fifteenth that of air. It is this which causes it to diffuse so rapidly, and renders it valuable for filling balloons. A liter of the gas weighs .0896 g. It is very inflammable, and burns with a pale, almost non-luminous flame. As noticed above, the hydrogen flame from a glass jet has a yellow color, but this is due to a compound of sodium in the glass, just as the hydrogen arising from the sodium on the water burned with a yellow flame. The heat of this flame is intense, as is seen by the rapidity with which the glass jet becomes red hot. When hydrogen burns in the air or in oxygen, water is the only product, the union being expressed by the following equation : or, 2H O 2 = 2H 2 0. 17- Combination of Hydrogen with Other Substances. The explosiveness of hydrogen when mixed with oxygen has already been noticed. One of its most remarkable properties is that of being absorbed or occluded by certain metals. By finely divided platinum the absorption is so rapid that the metal becomes red hot, and the jet of hydrogen is quickly ignited. Likewise, if a piece of spongy platinum be lowered into a mixture of hydrogen and oxygen, the rapid absorption in a short time causes sufficient heat to explode the mixture. At the usual temperature, hydrogen has very little HYDROGEN 45 affinity for most substances. As will be seen in Experi- ment 65, it explodes violently when mixed with chlorine, either on the approach of a strong light or by means of a spark. It unites vigorously with oxygen also" on the application of a flame, but a light has no effect. 18. Liquid Hydrogen. Hydrogen is one of the most difficult gases to reduce to the liquid condition. This has been accomplished, however, by reducing the temperature to 205 C., and allowing it to escape rapidly from a pres- sure of 180 atmospheres into a vacuum. At the same time this space is surrounded by a temperature of 200 C. Considerable quantities have been obtained in this way. In April of the year 1900, Dewar even succeeded in solidi- fying the gas. He surrounded liquid hydrogen with lique- fied air, and then by a pump caused so rapid an evaporation of the hydrogen that he soon obtained the remainder in a white, opaque solid. 19. Uses of Hydrogen. As a gas it has but few prac- tical uses. Its suitability for filling balloons has been mentioned, but in such cases it is generally used in a very impure form, mixed with various hydro-carbons given off with the hydrogen in the later distillation of coal. In the nascent state, that is, at the instant it is set free from some compound, it has great chemical activity and has the power of reducing many metals from their compounds. This use has already been seen in the passage of hydrogen over copper oxide, and will be further illustrated in our work with silver, iron, and other metals.* * The use of hydrogen in an automatic cigar-lighter is occasionally seen. As shown in the figure, a small glass cylinder has a cubical block, a, of porcelain in the bottom : upon this rests an inverted glass cylinder, c, with a tubular neck and stop-cock, s ; above this jet is supported a plat- inum sponge, p\ under the small cylinder upon the porcelain block is 46 MODERN CHEMISTRY SUMMARY OF CHAPTER Hydrogen Origin of term and meaning. Occurrence of hydrogen. Methods of making hydrogen. By decomposing water. With sodium or potassium. Describe method and apparatus. Chemical action Proof of, by experiment. By decomposing acids. With zinc. Draw apparatus and explain method. Commercial methods. For filling balloons. Characteristics of hydrogen. Experiments to illustrate. Density. Inflammability. Explosiveness, etc. Liquid hydrogen. Uses of hydrogen. Special points. Explain the hydrogen pistol. Philosopher's lamp. Singing flame. FIG. 11. placed some zinc, z, and in the outer cylinder diluted sulphuric acid. As the acid and zinc react upon each other, hydrogen fills the in- verted cylinder, forces out the acid, and the action ceases. If now the stop-cock is opened, the hydrogen flows out of the jet, the acid reenters, and the generation of gas continues. As already seen, the hydrogen jet is quickly ignited by the platinum. When the customer has lighted his cigar, the stop-cock is again turned, and the action soon ceases. It ought to be said, perhaps, that such apparatus is of more interest as a novelty than as of real lasting utility. ; CHAPTER V 4 / OXYGEN, COMBUSTION, OZONE OXYGEN : = 16 1. Its Discovery. The term oxygen is derived from two Greek words, meaning acid-former, and was given to this element because it was believed to be essential to the production of all acids. Oxygen was discovered by Scheele in 1773, but he did not publish his discovery until 1775 ; and as in the meantime Priestley had isolated the same gas and had published an account of his experiments, the latter is generally given the credit. 2. Abundance of Oxygen. Oxygen is found in the atmosphere in large quantities, uncombined, constituting about one-fifth of the whole. It has been estimated that there is in the atmosphere alone over two and a half million billions of pounds. A liberal estimate of the amount used annually in respiration, and all forms of combustion, is about two and a quarter billion pounds. At this rate, in a century the entire world would use only one ten -thousandth part of the whole. At the same time it must be remembered that plant life is pouring the oxy- gen back again into the air, so that there is no danger of the equilibrium being destroyed. Oxygen also forms by weight eight-ninths of water, and being absorbed by the same, exists therein in considerable quantities in a free state. It is this free oxygen which is breathed by fishes. On account of its great affinity for other substances, it is 47 48 MODERN CHEMISTRY found in combination with nearly all known elements, and forms in this way about 45 to 50 per cent of the earth's crust. 3. How to produce Oxygen. As a matter of historical interest, the method employed by Priestley is still some- times used. It is as follows : EXPERIMENT 27. Place in a hard-glass test-tube about a half gram of mercuric oxide, HgO, and heat strongly. Notice the change in the appearance of the oxide. Insert into the tube a pine splinter with a spark upon the end. What happens? What do you notice collecting upon the sides of the tube? What substances have there- fore been obtained by heating this oxide ? 4. Explanation. The heat used has served to decom- pose the molecules of mercuric oxide into their constituent parts, thus : - -f HEAT or, 2 HgO + heat = 2 Hg + O 2 . The two molecules of red oxide of mercury have yielded two molecules of mercury, one atom in each, and one molecule of oxygen, having two atoms. By continuing the operation the entire amount of the red oxide would disappear, while the deposit of mercury upon the sides of the tube would gradually increase. 5. Other Methods of obtaining Oxygen. The above method, though of interest, furnishes too limited a quan- tity of oxygen for practical purposes. A better and more OXYGEN, COMBUSTION, OZONE 49 common way is to heat potassium chlorate, KC1O 3 , with manganese dioxide, MnO . AST l_ FIG. 12. EXPERIMENT 28. Mix together in a good-sized test-tube, or small flask, 1 or 2 g. of potassium chlorate and half as much manganese dioxide. Support upon a ring-stand with a wire screen protection, as shown in the accompanying figure, and attach the cork and delivery tube. Heat gen- tly at first, and then more strongly, but moderately, so as to regulate the flow of gas and not let it become too rapid. Allow the first that comes over to escape, then collect several bottles of the gas over water as you did the hydro- gen, and use for the following experi- ments. Save the contents of the flask for further use. EXPERIMENT 29. Slip a sheet of glass or paper under a small bottle of oxygen, and place it in an upright position upon the table. Now plunge into the oxygen a taper, or pine splinter, with a spark upon it. Do you obtain the same results as before in the case of the oxide of mercury ? EXPERIMENT 30. Into another bottle of oxygen lower a deflagrat- ing spoon containing some burning sulphur; does it burn any differ- ently than in the air? If no deflagrating spoon is at hand, the student can prepare one by hollowing put the end of a short stick of gas carbon, or of crayon, and attaching a wire handle of suitable length. EXPERIMENT 31. Fasten a piece of soft or bark charcoal to a stout iron wire, hold it in the burner flame until it begins to glow, then plunge into a jar of oxygen. If the charcoal is soft, the results will be very striking. Describe them. EXPERIMENT 32. Twist together three or four fine iron wires, fasten to the end a small pine splinter, or warm and dip into sulphur and ignite. Plunge quickly into a large jar of oxygen which contains about an inch of water in the bottom. Describe the results. Do you see anything falling to the bottom of the bottle? A knife-blade or watch-spring may be thus burned, by first drawing the temper and using a larger amount of kindling material. 50 MODERN CHEMISTRY EXPERIMENT 33. Put into a deflagrating spoon a little red phos- phorus, ignite it, and thrust it into a large jar of oxygen. Describe the combustion and the fumes that fill the jar. These are phosphorus pentoxide, a substance mentioned under deliquescent bodies. 6. The Chemical Action. In preparing oxygen as above, the manganese dioxide remains unchanged. This and other facts shown in the equation below will be proved in Exp. 34. What has really taken place is the same as in the use of the mercuric oxide. The heat has simply decomposed the molecules of potassium chlorate, setting free the oxygen and leaving behind a new compound con- taining only potassium and chlorine, called potassium chloride. The change may be shown thus : + HEAT = The two molecules of potassium chlorate shown here have each given up three atoms of oxygen, which have combined to form three molecules of oxygen, while two molecules of potassium chloride remain behind. These facts are more usually written thus : 2 KC1O 3 + heat = 2 KC1 + 3 O 2 . 7. Effect of the Manganese Dioxide. If potassium chlorate be used alone, instead of mixing with manganese dioxide, as we did above, the same results are obtained. OXYGEN, COMBUSTION, OZONE 51 but considerably more heat is required. Apparently, therefore, manganese dioxide has simply acted by its presence, or, as it is called, by catalysis. It is believed, however, that the dioxide is first converted into another compound, which at the temperature present is unstable, and that this in breaking up yields oxygen and the dioxide again. EXPERIMENT 34. To prove that the manganese dioxide remains unchanged and that potassium chloride is formed. To the residue in the flask in Experiment 28, add about 50 cc. of water, let it stand a few minutes, shaking occasionally, warm gently and then filter. Boil this clear filtrate to dryness in an evaporating dish. While this is proceeding, add a little water to the black residue on the filter paper once or twice to wash it, throw the water away, and let the black residue dry. When the solution in the evaporating dish is perfectly dry, scrape it out, mix with a little fresh manganese dioxide, transfer to a test-tube, and heat. Do you observe any indication of oxygen being given off? If not, we may conclude under the present circum- stances that the oxygen was all removed in the previous heating, and that the white solid residue is KC1, potassium chloride, and not potas- sium chlorate, KC1O 3 . When the black residue on the filter paper is dry, mix with it a little potassium chlorate, transfer to a test-tube and heat. Is oxygen given off readily ? What proof ? Is there any reason for believing that the black residue is still manganese dioxide? 8. The Proof by Weighing. It will be well to try an experiment by which the facts discovered in preceding experiments may be proved. Such is the following experi- ment. EXPERIMENT 35. Put into a test-tube or flask about a gram of potassium chlorate, put it upon the scales and balance it with shot or sand in a small box. Now add about a half gram of manganese dioxide, and then weigh carefully. As the box and shot counter- balance the flask and potassium chlorate, the weights added show at once the amount of dioxide used. Now connect a delivery tube and .heat to drive off the oxygen. When the operation is complete, known 52 MODERN CHEMISTRY by the fact that the gas no longer bubbles up through the water, re- move the delivery tube from the water, and let the flask cool. When cold, add a few cubic centimeters of water to dissolve the potassium chloride, then filter and wash the black residue as before. When thoroughly dry, weigh the residue and filter paper and subtract the weight of the paper. The latter may be obtained by weighing ten of them, or, if the balance is not very delicate, a hundred, and then tak- ing the fractional part. Does the weight of the black compound now agree with its weight before heating? 9. Commercial Methods of making Oxygen. Most of these methods consist in abstracting oxygen from the air by using a substance which when heated or when under pressure will absorb oxygen, and then when cooled or when the pressure is removed will again give it up. One of the best known of these methods is Erin's, which con- sists in using barium oxide, BaO, as the chemical agent. When gently heated in the air, it takes up an additional amount of oxygen, forming barium dioxide, BaO 2 , thus, BaO + O = BaO 2 . If, now, the heat is increased, the barium dioxide is unable to retain the additional atom of oxygen taken from the air and gives it up again, thus, BaO 2 + heat = BaO + O. Or, if the pressure under which the barium dioxide was formed is decreased, the same results follow at considera- bly less expense. 10. Motay's Method. The principle of this is about the same as that of Brin's, but different substances are used. Manganese dioxide and caustic soda, when heated moderately in a current of air, form a compound which at a higher temperature is again decomposed, yielding up the oxygen previously taken from the air. OXYGEN, COMBUSTION, OZONE 53 11. Other Methods. These are not important as a means of producing oxygen for commercial or experimental purposes, but the principle underlying them is involved in a number of the processes of chemistry with which we shall deal later, and should consequently be understood. It will be noticed that in all the methods of preparing oxygen used above, we have employed substances con- taining a large per cent of that element. There are sev- eral other substances of similar composition which may be made to furnish oxygen. Thus, manganese dioxide, MnO 2 , potassium dichromate, K 2 Cr 2 O 7 , and potassium permanganate, KMnO 4 , when heated with sulphuric acid, H 2 SO 4 , will yield oxygen. EXPERIMENT 36. Put a half gram of manganese dioxide into a test-tube and add about a cubic centimeter of sulphuric acid. Warm gently, collect a small bottle of the gas, and make the usual test for oxygen. What are the results ? It will be found that the dichromate and the permanga- nate act in a similar way, except that the quantity of gas obtained is considerably greater. The chemical action may be shown thus : - Mn0 2 + H 2 S0 4 = MnS0 4 + H 2 O + O. 54 MODERN CHEMISTRY In a similar way the reaction of potassium dicliroraate and sulphuric acid upon each other may be shown ; K 2 Cr 2 7 + 4 H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4 H 2 O + 3 O, and of potassium permanganate and sulphuric acid, 2 KMnO 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 3 H 2 O + 5 O. See pages 322, 325, for application of this property of the above substances. 12. Characteristics of Oxygen. Oxygen is an odorless, colorless gas, a little heavier than air, the weight of a liter being 1.43 g. As already noted it is slightly soluble in water, and upon this fact depends the life of aquatic ani- mals, which abstract this free oxygen from the water. It may be liquefied by extreme cold and pressure. This was first accomplished about a quarter of a century ago by Cailletet and Pictet, who succeeded in preparing a small quantity at great cost. At present it is made in almost any amount by first liquefying air and then allowing the nitrogen to boil out. (See page 100.) 13. Peculiarities of Liquid Oxygen. In the liquid con- dition oxygen is of a pale blue color and boils at about 180 C., a few degrees higher than the boiling point of nitrogen. It presents many striking peculiarities ; a rod of carbon heated red hot and plunged into the liquid oxygen at a temperature 180 below zero burns vigorously, while a stout iron wire similarly heated is rapidly con- sumed with a brilliant display of sparks. Cotton rags saturated with it and confined in a cylinder, when ignited, OXYGEN, COMBUSTION, OZONE 55 explode so violently as to burst tubes made of iron or brass. 14. Chemical Affinity of Oxygen. The strongest chem- ical property of oxygen is its affinity for other substances. This was seen in the rapidity of combustion of the various ignited substances when placed in .an atmosphere of oxy- gen. From these experiments it is not difficult to see what would be the results were the air undiluted oxygen. The smallest spark would be sufficient to start the fiercest conflagration, while our stoves, furnaces, etc., would be rapidly consumed, accompanied by a most brilliant display of sparks. 15. Uses of Oxygen. As is well known, oxygen is absolutely necessary for life. It is absorbed by the blood through the walls of the air-cells of the lungs and carried by the red corpuscles to all parts of the body. Here it unites with the waste material, burning it to carbon dioxide and other compounds, and at the same time warm- ing the body. The carbon dioxide is carried back to the lungs, from which it is thrown off into the air. In cases of asphyxiation pure oxygen is sometimes used as a restora- tive, but ordinarily, if breathed for any length of time, the temperature of the body rises owing to the increased de- struction and consumption of tissue, and general feverish symptoms follow. A limited number of experiments by the author show that small animals, such as mice, when placed in an atmosphere of pure oxygen soon exhibit great activity, followed by apparent relaxation and complete exhaustion. Experiments have also shown that animals in oxygen under pressure would very quickly die, as if the gas in this condition were an active poison. EXPERIMENT 37. Into a flask, 0, the weight of which is known, supported in a ring-stand as shown in Fig. 13, put 2.5 g. of potassium 56 MODERN CHEMISTEY chlorate. Let the tube, d, just reach through the cork of A, nearly filled with water. Make e in two parts, joined at a by several inches of rubber tubing. By suction fill e with water, and put pinch clamp upon the rubber. Make the corks air-tight. With the tube in posi- tion, fill B to the same height as water in .4, open clamp an instant, Fio. 13. Apparatus to use in finding the Weight of Oxygen. then empty B. Remove the clarnp and heat, carefully at first, until water is no longer forced into B. When is cool, raise or lower B till water stands at same height in both bottles, fasten clamp and measure water in B. This equals volume of oxygen at pressure and temperature of the room. Reduce to standard conditions. (See p. 96.) Weigh 0; loss equals weight of oxygen expelled. Knowing the weight of a certain number of cubic centimeters, find weight of one liter. By similar methods, the weight of a liter of various other gases, insoluble in water, may also be determined. 16. Oxidation and Combustion. When any substance combines with oxygen to form a new compound, it is said to be oxidized, and the process is known as oxidation. This may be slow or rapid. When it takes place so rapidly as to be accompanied by heat and light, the pro- cess is called combustion. To illustrate : when a piece of iron is exposed to the air in the presence of moisture, it soon becomes covered with rust, which is really an oxide of OXYGEN, COMBUSTION, OZONE 57 iron; in other words, the iron has been oxidized. Again, when we tipped the iron wires with sulphur and ignited it, they were rapidly consumed in the jar of oxygen with much heat and considerable light. This was combustion. A pile of brush will gradually decay, or oxidize, without any perceptible heat, but by setting it on fire we quickly destroy it by the process of combustion. ^t. Combustible Substances and Supporters of Combus- tion. Substances which thus burn in oxygen or its diluted form, the air, are said to be combustible, while the substance in which they burn is called a supporter of com- bustion. Thus, when a jet of hydrogen burns in a jar of oxygen, the former would be spoken of as the combustible substance, and the latter as the supporter of combustion. It is true, however, that if we thrust a delivery tube from which a current of oxygen is issuing, up into a jar of hydrogen which is burning at the mouth, as seen in Fig. 14, the jet of oxygen will be seen to burn in the atmosphere of hydrogen, just as before the hydrogen did in the oxy- gen. Yet in view of all the facts it seems better to adhere to the statement previously made, that it is really the hydrogen which burns, and the oxygen which supports tl^e combustion. 18. Kindling Temperature. It is well known that some substances ignite much more readily than others. This, chemically speaking, simply means that some combine with oxygen at a lower temperature, or much more readily, than do others. Thus, substances like alcohol and many FIG. 14. 58 MODERN CHEMISTRY oils need but little heat to ignite them; phosphorus, like- wise. Pine wood needs a higher temperature, and hard wood still higher. The point at which any substance takes fire is said to be its kindling temperature. 19. What is a Flame ? A flame is simply burning gas. Whenever a substance will not burn with a flame, it is because there is either no gas present or there is nothing which may be converted into a gas. For example, when a lamp burns, the oil drawn up through the wick by capil- lary attraction is volatilized by the heat, and it is the burning of this gas that makes the flame. On the other hand, charcoal and the hardest natural coals do not burn with a flame, because previous heating has driven out all the gaseous products. However, they may be heated suf- ficiently to be partially converted into carbon monoxide, a,gas which burns with a pale blue flame. ^0. The Oxy-hydrogen Blowpipe. This is a lamp arranged for burning hydrogen thoroughly mixed with oxygen, and affords one of the hottest flames known. Its construction will be understood from the figure, which gives a sectional view. The inner tube, m, is connected with the oxygen tank, and the outer, n, with the hydrogen. In this way, as the inner tube is somewhat shorter, the gases become thoroughly mixed before leaving the tube at JEJ, hence the combustion is perfect. The FIG. is. The Oxy-hydro- pressure should be so regulated gen Blowpipe. . . 111 as to furnish twice as much hydro- gen by volume as oxygen. This blowpipe is used for melting very refractory substances. It is also used espe- OXYGEN, COMBUSTION, OZONE 59 cially in furnishing light for stage effects, stereopticon views, and for illuminating moving floats in street parades given after dark. When used for these purposes the almost non-luminous blue flame is allowed to strike upon a stick of prepared lime supported in a socket just in front of the blowpipe. This is often called the calcium or Drummond light, and is of dazzling whiteness, rivaling the electric arc light. OZONE 21. Its Discovery. This substance, on account of its peculiar odor, was named from a Greek word which means to smell. It was first observed in passing electrical sparks through a tube of oxygen, and is always noticeable when an electric discharge takes place in the air. 22. What is Ozone ? For some time it was regarded as a compound body, but is now known to be simply a con- densed form of oxygen. Quite a number of substances, such as sulphur and phosphorus, appear in a form other than the usual one : this is known as the allotropic, a word which means simply another form. 23. Methods of obtaining Ozone. It is impossible to prepare ozone in large quantities in a pure condition, be- cause by the best methods only a small per cent of the oxygen used is converted into its allotropic form. Usu- ally not over one or two per cent is obtained, and under the most favorable circumstances only about twenty per cent. In the ordinary methods of making oxygen, ozone is almost always obtained in appreciable quantities. One of the easiest methods of preparing it is given in the fol- lowing : EXPERIMENT 38. Scrape a stick of phosphorus perfectly clean, put it into a bottle, and add water sufficient to cover about half of it. In a few minutes the presence of ozone may be detected by suspending 60 MODERN CHEMISTRY in the bottle a strip of white paper moistened with a solution of potassium iodide and starch. The paper will turn decidedly blue. A solution of potassium permanganate treated with strong sulphuric acid also gives the test for ozone, along with the oxygen thus evolved. 24. Ozone in the Air. Ozone is believed to exist in appreciable quantities in the atmosphere, being produced mainly by electrical discharges. Its presence in dwell- ings is never perceptible, and scarcely ever in large cities, but in the country a strip of starch paper exposed to the breeze for some time shows the characteristic blue color. 25. How Ozone differs from Oxygen. As already stated, it is a condensed form of oxygen. Experiment has shown that if a given amount of ozone is decomposed so as to form ordinary oxygen, the volume increases one-half ; that is, 100 cc. of ozone would become 150 cc. of oxygen. On the other hand, if a closed volume of oxygen be subjected to a silent discharge of electricity so as to convert a por- tion of it into ozone, a corresponding decrease of volume takes place. 26. For example, suppose 150 cc. of oxygen be thus treated, and that 30 cc. are converted into ozone. It is found that by absorbing this ozone so as to separate it from the remain- ing oxygen, and again setting it free, there are only 20 cc. of ozone, Molecule of Molecule of wn il e but 120 cc. of oxygen remain. Oxygen. Ozone. * If, however, this 20 cc. of ozone be heated strongly, so as to convert it into oxygen again, we shall find the volume increases to 30 cc. The mole- cule of ozone therefore would differ from that of oxygen, in that it contains three atoms of oxygen, while the other has only two. This is shown in the accompanying figure. OXYGEN, COMBUSTION, OZONE 61 From this it naturally follows that ozone is 50% more dense than oxygen. 27. Properties of Ozone. The properties are also dif- ferent from those of oxygen. Its odor has already been mentioned. If placed in a long glass tube, so as to give considerable depth, it is seen to have a blue tinge. It readily destroys the color of such vegetable solutions as indigo and litmus and quickly attacks such metals as mer- cury and silver, which remain unchanged in the air and which are little affected by oxygen when heated. We have seen its effect upon potassium iodide above. Free iodine always turns starch blue. In the test made, the ozone united with the potassium in the potassium iodide to form an oxide with the metal, and the iodine was thus set free. In the same way if a drop of ammonium hydroxide be let fall into a jar of ozone, a dense white cloud forms, owing to the fact that a white solid compound of ammonia is formed, thus : 2 NH 4 OH + O 3 == NH 4 NO 2 + 3 H 2 O. 28. Liquid Ozone. Ozone may be liquefied at ordinary atmospheric pressure by reducing the temperature to 106 C., a point considerably higher than that at which oxygen liquefies. Ozone is also a very unstable body, changing back readily into oxygen ; an illustration of this is seen in the fact that if a quantity of ozone be suddenly compressed and heated, it explodes with violence. It is because of this instability that ozone is so strong an oxi- dizing agent. The nascent oxygen liberated, if inhaled, attacks the mucous linings, causing an irritation some- what like that of dilute chlorine. More than this, head- ache soon follows, if much ozone is inhaled, even though diluted with considerable quantities of oxygen. 62 MODERN CHEMISTRY TABULAR VIEW OF DIFFERENCES OXYGEN Colorless. Odorless. Density; slightly heavier than air. Two atoms in molecule. Strong oxidizer. Liquefies at - 180C. Stable. OZONE Blue. Peculiar odor. Density; considerably heavier than air. Three atoms in molecule. Very strong oxidizer. Liquefies at - 106 C. Unstable. 29. Value of Ozone. It is believed to have a beneficial effect in destroying disease germs and in oxidizing decay- ing organic matter. ^^80. Isomeric and Polymeric Bodies. Just as ozone is another form of oxygen, so we shall find that phosphorus and certain other elements present allotropic forms as unlike the usual forms as oxygen and ozone. When we come to the study of compound bodies we often find two substances not at all alike in properties, which, upon analysis, are found to contain exactly the same elements united in exactly the same ratios. Thus aldehyde and oxide of ethylene both have the same -composition, repre- sented by the formula C 2 H 4 O, but their properties are very different. Such substances are said to be isomeric. 31. Sometimes while they have the same percentage composition, the vapor density of one will be several times that of the other. Thus acetylene is C 2 H 2 , and benzine C 6 H 6 . In each case the carbon is ^|, or 92.3 per cent, of the molecule ; but the molecular weight of one OXYGEN, COMBUSTION, OZONE 63 is three times that of the other. Such substances are said to be polymeric. HYDROGEN DIOXIDE: H 2 2 32. Composition. This is a compound which in some of its characteristics resembles ozone. In composition it is most like water, having one additional atom of oxygen. It is believed to exist in the air in minute quantities, and some of the effects attributed to ozone may be due to hydrogen dioxide. 33. How to. obtain It. For experimental purposes it is usually prepared by treating barium dioxide with dilute sulphuric or hydrochloric acid. EXPERIMENT 39. Add to about a gram of barium dioxide a little water, and then dilute sulphuric or hydrochloric acid. Stir for a moment or two with a glass rod. To prove the presence of hydrogen dioxide, add a few drops of potassium dichromate and about a half cubfc centimeter of ether, and shake well. The hydrogen dioxide forms a blue solution with the dichromate, which is taken up by the ether and thus concentrated within little space. 34. Some of its Peculiarities. Like water, it is a color- less liquid, but is thicker or sirupy, and has a bitter taste. It is very unstable, decomposing at all tempera- tures into water and oxygen ; it is therefore a good bleaching agent, the bleaching being done by the nascent oxygen. Like ozone, it readily tarnishes silver and decom- poses potassium iodide, giving in the same way a test with starch paper. It is soluble in water, and thus diluted it will bleach the skin, but when concentrated it burns or blisters it. 35. Uses. This compound, more usually sold under the name hydrogen peroxide, is now manufactured very 64 MODEEN CHEMISTRY cheaply, and is used to a considerable extent as a bleach- ing agent, especially for hair and feathers. It is used largely by dentists and in surgery as an antiseptic, and to some extent in cleansing oil paintings and engravings. SUMMARY OF CHAPTER Discovery of oxygen. Meaning of term. Abundance of oxygen Various forms in which it occurs. Methods of preparing oxygen. Priestley's method. Ordinary method. Chemicals and apparatus used . Chemical changes involved. Proof of these changes Experimental. Proof by weight. Other methods. Character of substances used. Characteristics of oxygen Compare with hydrogen in Color, odor, density, combustibility. Power of supporting combustion. How would you distinguish the two ? Some peculiarities of liquid oxygen. Uses of oxygen. Special. Meaning of terms combustion, oxidation, flame, kindling point. Illustration of the terms. Description and drawing of oxyhydrogen blowpipe. Uses for it. Ozone Meaning of the word. What is its relation to oxygen ? Methods of obtaining. Compare with oxygen, showing differences. Value of ozone. Hydrogen dioxide Formula. Compare with water in properties. Uses for the compound. CHAPTER VI CHEMICAL NOTATION, SYMBOLS, FORMULA, EQUATIONS, PROBLEMS 1. Symbols. The student will have noticed that in chemistry we frequently employ a short-hand method of expressing the different elements and their compounds. Thus we have seen that hydrogen is represented by H, oxygen by O, and so on. These are called symbols. 2. Their Form. Frequently, the symbol of an element is its initial letter ; often, however, this is the same for a number of elements, as for example, carbon, calcium, cad- mium, copper, etc. In such cases, the most common usually is designated by the initial letter; another by the first and second letter, as Ca, calcium; another by the first and some other distinctive letter, as Cd for cadmium. Frequently, the first or the first and second letters of the Latin term for the same substance are used, as Cu for copper, from cuprum. In lixe manner, sulphur, silicon, selenium, silver, are designated by the symbols S, Si, Se, and Ag (from the Latin argentuni). The Latin has fur- nished a number of the symbols of the common elements : thus, sodium, Na (natrium), potassium, K (kalium), iron, Fe (ferrum). The symbol, Hg (hydrargyrum), for mer- cury is from the Greek. 3. Strict Meaning. Strictly speaking, the symbol of an element not only represents that element, but a defi- nite amount of it ; that is, one atom. Hence, to speak of an element by using its symbol when we mean an indefinite amount is unscientific and should not be practiced. 65 66 MODERN CHEMISTRY 4. Formulae. As the elements are represented by symbols, so compound bodies are by formulae ; that is, i>y an aggregation of symbols. Compounds are usually named by simply combining the terms representing the elements entering into the composition, the more electro- positive being placed first ; thus, potassium iodide consists of two elements, potassium and iodine. The formulae are always arranged in the same way: thus, KI, potassium iodide ; KC1, potassium chloride. It will be seen, there- fore, that as a symbol represents an atom of an element, so a formula represents the smallest amount of a compound body, a molecule. 5. Sub-figures. When the elements enter into com- position, in other than a single atom of each, that fact is indicated by putting a small figure below and at the right of the symbol ; thus, H 2 O, the formula for water, indi- cates that there are two atoms of hydrogen in the mole- cule, and, H 2 O 2 , for hydrogen peroxide, indicates that there are two atoms of each. These sub-figures are some- times appended to a group of elements, in which case the group is inclosed in parentheses and the figure placed outside ; for example, lime-water, or calcium hydroxide is Ca(OH) 2 . This might also be written CaO 2 H 2 , but the former method is preferable, as will be seen later. If we desire to indicate more than one molecule of a substance, this is done by prefixing a coefficient to the formula. Thus, 2 HC1 indicates two molecules of hydrochloric acid ; 5 H 2 O, five molecules of water. 6. Radicals. By a radical we mean a group of ele- ments which in most chemical reactions seem to hold together, but which do not by themselves form a distinc- tive compound. For example, (HO) seen in the formula for lime-water is a radical known as hydroxyl, which enters REACTIONS 67 into a great many compounds. Again, (NH 4 ) is a group called ammonium, which is very common, and ordinary ammonia water is NH 4 OH, composed of two radicals (NH 4 ) and (OH), not written NH 5 O, because of this fact. 7. Reactions. Equations in chemistry which show the chemical changes that take place when two or more substances react with each other are called reactions. We have already seen a number of these ; thus : 2 Na + 2 H 2 O = 2 NaOH + H 2 ; Zn + H 2 SO 4 = ZnSO 4 + H 2 . 8. The first indicates that two atoms of sodium uniting with two molecules of water will produce two molecules of caustic soda and two atoms or one molecule of hydro- gen. Another thing must be noticed, and that is that every atom appearing in one member of the equation must also be found in the other. Thus, the two atoms of sodium are seen in the second member of the equation in the two molecules of caustic soda, the four atoms of hydro- gen in the water appear partly in the caustic soda and partly as free hydrogen; likewise the two atoms of oxy- gen in the water are found in the caustic soda. It must be borne in mind that a coefficient before a formula mul- tiplies every symbol in that formula. Thus, 2 KC1O 3 means that there are two atoms of potassium, K; two of chlorine, Cl; and six of oxygen. This will be seen from the fol- lowing illustration : a represents a molecule of water containing two atoms of hydrogen and one of oxygen ; b represents a second molecule having the same composition. Tak- ing both together, or two molecules of water, 2 H 2 O, we see there are four atoms of hydrogen and two of oxygen. 68 MODERN CHEMISTRY 9. Atomic Weights. We cannot think of matter with- out assigning to it some weight. So the atoms of the elements, though the smallest conceivable portions of matter, are assumed to have definite weights. Hydrogen, being the lightest of substances, is taken as the standard,* and its atomic weight is assumed to be one, or by some, as one micro-crith. Of course, a weight as small as this has never been determined, and is therefore merely an abstract idea ; but something is necessary for comparison. When we speak of the atomic weight of an element, there- fore, we simply mean its density compared with hydro- gen. Thus, we say the atomic weight of oxygen is 16, of carbon, 12 ; we mean that these elements are, respectively, sixteen and twelve times as heavy as hydrogen, or, if one cubic foot of hydrogen weighs a gram, one cubic foot of oxygen will weigh sixteen grams. 10. Molecular Weight. By molecular weight we mean the sum of the weights of the atoms entering into the composition of the molecule. For example, H 2 O repre- sents a molecule of water ; the two atoms of hydrogen weigh 2, the one of oxygen, 16, or all together, 18. The molecular weight of water is therefore 18. Now if we examine any chemical equation, we will find that the sum of atomic weights in one member must equal the sum of those in the other member. Take the following : Na + H 2 O = NaOH + H, and substituting the atomic weights as given in the table on page 9, we have 23 + (2 + 16) = (23 + 16 + 1) + 1, or 41 = 41 ; and this must always be so in any true reaction. * There has long been a controversy whether Hydrogen, H = 1, or Oxygeu, = 16, should be the standard. The latter is increasing in favor. EQUATIONS 69 11. Writing Equations. A chemical equation is valu- able in that it shows at once in concise form not only the substances which enter into the reaction, but also the products formed, and the exact amount of each. At first the student will experience some difficulty in completing even the simpler reactions, but he must remember that they only show what has been proven by experiment. Thus on page 82 we decomposed water by electricity and ob- tained two gases, one double the other in quantity. These were shown to be hydrogen and oxygen. We represent these facts by the reaction, H 2 = H 2 + O. When we treated water with sodium, we obtained not only a gas, which was hydrogen, but a solution of caustic soda. Representing our experiments in brief form, we wrote H 2 O + Na = NaOH + H. So all reactions are determined experimentally, and the student at first will be called upon to write but few which he has not worked out himself. 12. Practical Value of the Equation. Having deter- mined by experiment the products that are formed in any chemical reaction, and having therefrom written the equa- tion, we can readily ascertain the amount of each product that will be formed from a certain amount of another ; or if required to produce a definite quantity of any body, we can calculate what it will be necessary to use in obtaining it. To illustrate, suppose we are required to determine how much zinc will be necessary for the preparation of 50g. of hydrogen. We would first write the equation, showing the preparation of hydrogen : Zn + H 2 S0 4 = H 2 + ZnS0 4 . 70 MODERN CHEMISTRY In the table we find the atomic weight of zinc is 65 ; looking at the reaction, then, we would see that 65 parts of zhic by weight produce, when reacting with the acid, 2 parts of hydrogen. The gas obtained is therefore fa by weight of the metal used. Then, 50g. = - 6 V g L = i of 50 g. = 25 g. || = 65 x 25 g. = 1625 g. of Zn. PROBLEM 1 . How much sulphuric acid is it necessary to put with 260 g. of zinc in preparing hydrogen? Using the same reaction as above, we find first the molecular weight of H 2 SO 4 , which is 98. We see then that 65 parts of zinc unite with 98 of acid, or the acid used is | of the rnetal. Then, || of 260 g. Zn = 98 x 26 = 392 g. H 2 SO 4 . 65 Some prefer to solve such problems by proportion, thus : The wt. of the Zn : wt. of acid : : wt. of Zn in g. : wt. of acid in g. ; or, 65 : 98 : : 260 : x. , = Mx2eo = m 65 PROBLEM 2. In using 260 g. of zinc in preparing hydrogen, how much zinc sulphate, ZnSO 4 , will be obtained? PROBLEM 3. How many grams of oxygen may be obtained from 450 grams of potassium chlorate? PROBLEM 4. How much caustic soda will be produced in prepar- ing 10 g. of hydrogen by using metallic sodium and water? SUMMARY OF CHAPTER Symbols and formulae Difference between them. Composed of what. Exact meaning of each. NITROGEN AND ITS COMPOUNDS 71 Radicals. . Atomic and molecular weights. Meaning of the term. Meaning of the terms. Illustrations. Illustrations. Chemical equations. Value of. Problems. CHAPTER VII NITROGEN AND ITS COMPOUNDS NITROGEN : N = 14 1. History. Nitrogen, meaning niter producer, was given this name because of its being an important con- stituent of saltpeter, often called niter. It had previously been called azote, a term which meant that it would not support life. 2. Where found. As already stated, nitrogen con- stitutes about four-fifths of the air, and is uncombined. It also exists in various compounds, such as saltpeter, potassium nitrate, KNO 3 , and Chile saltpeter, sodium nitrate, NaNO 3 . It also enters into the composition of many vegetable and animal products, and in their decom- position is given off into the air in the form of ammonia, NH 3 . 3. How to prepare Nitrogen. As nitrogen exists so abundantly in a free state in the air, this is the best source from which to obtain it. Any method by which we can remove the oxygen and leave the nitrogen will do. For this purpose phosphorus is generally used. EXPERIMENT 40. Cover a large flat cork with a coating of plaster of paris and float it upon a pan or basin of water. A small iron saucer serves well instead of the cork. Put upon it a small piece of 72 MODERN CHEMISTRY FIG. 16. phosphorus and ignite. Quickly place over the burning phosphorus a large wide-mouthed jar. Notice that the water gradually rises in the jar to take the place of the consumed oxygen, and that in a few minutes the white fumes are absorbed by the water. Owing to the expansion caused by the heat some bubbles of air almost always escape in the early part of the experiment. Jf the phosphorus is not ignited, but the combination with the oxygen is allowed to take place slowly, this loss may be avoided, but several hours are required. Sometimes it is more con- venient to use a deflagrating spoon instead of a cork ; if so, the handle must be bent V-shaped, so as to bring the phosphorus above the water even after it has risen in the jar. Notice about how much the water rises. When the fumes have all disappeared, lift the jar and put a burning candle up into the gas. What happens? Compare it with similar tests with oxygen and with hydrogen. 4. Other Ways of preparing Nitrogen. The method already given, while the easiest and most commonly used, does not give as pure nitrogen as may be obtained in some other ways. If a current of air be made to stream slowly over a tube con- taining copper turnings Mr heated to redness, the oxy- gen will combine with the copper, forming copper ox- ide, and the nitrogen will remain. Then if this is allowed to bubble through FIG. 17. Nitrogen, prepared bypass- a bottle of lime-water, the ing Air over Copper ' carbon dioxide will be absorbed, and we shall obtain a fairly pure nitrogen. The illustration will show the method. NITROGEN AND ITS COMPOUNDS 73 Nitrogen may also be obtained by heating certain com- pounds containing it. EXPERIMENT 41. Into a small flask put 1 or 2 g. of sal am- moniac, NH 4 C1, and the same amount of sodium nitrite, and add about 30 cc.-bi_ water. Heat gently and cautiously, and collect the gas over water as you did oxygen and hydrogen. Test the gas for nitrogen. What are your conclusions? 5. Peculiarities of Nitrogen. From the experiments made the student will notice that the gas has no color ; it is odorless, lighter than air, will neither burn as does hydrogen, nor support combustion as does oxygen. It has no affinity for other substances at ordinary tempera- tures. It will combine with red-hot magnesium in the absence of oxygen, and with oxygen when a discharge of electricity takes place, both of which methods have been used in preparing argon from its mixture with atmos- pheric nitrogen. It will not support respiration any more than it will combustion, and is one of the most inac- tive substances known. This inactivity, or feeble chemi- cal affinity of nitrogen, is the reason for the instability of many of its compounds, as seen in the explosiveness of gunpowder and nitroglycerine. 6. Value of Nitrogen. The use of nitrogen, except in the form of many valuable compounds, seems to be simply to dilute the oxygen of the air as already stated. COMPOUNDS OF NITROGEN 7. In an indirect way nitrogen forms a large number of compounds, many of which are very valuable. Among these we shall first consider ammonia. 8. Ammonia, NH 3 . As already stated, ammonia is one of the products formed in the decomposition of nitroge- 74 MODERN CHEMISTRY nous organic matter ; that is, organic matter which con- tains nitrogen in addition to the usual carbon, hydrogen, and oxygen. It finds its way into the air from these sources, and being absorbed by the moisture of the air is brought down in the rain, and usually exists in very small quantities in cistern and river water. With these excep- tions, ammonia does not occur free to any extent, but is found abundantly in certain compounds, especially sal ammoniac or ammonium chloride, NH 4 C1. 9. The commercial supply of ammonia is obtained from the distillation of coal in the manufacture of common \lluminatinggas. (See page 153.) The decay of organic matter, attended by the formation of ammonia, occurs as follows : when the nitrogenous matter decomposes, the former arrangement existing among the atoms of carbon, oxygen, nitrogen, and hydrogen is broken up, and in the rearrangement the nitrogen and hydrogen unite to form ammonia. 10. Ammonia prepared from Coal. In the distillation of coal the process is really the same, but more rapid, and ammonia is one of the impurities given off with the hydro- carbon gases. These are all passed through a tank filled with water, which absorbs the ammonia and forms an aqueous solution, known as aqua ammonia or ammonium hydroxide. This, more or less impure, is drawn off at inter- vals and treated with hydrochloric acid, which converts it into a salt of ammonia, ammonium chloride, NH 4 C1, as shown by the following reaction : NH 4 OH + HC1 = NH 4 C1 + H 2 O. 11. Then by treating this chloride with some strong alkali like caustic potash or soda, and heating, ammonia is again liberated, and being passed into water, produces NITROGEN AND ITS COMPOUNDS 75 the aqua ammonia of commerce. The following shows the reaction which takes place : NH 4 C1 + KOH = NH 3 + KC1 + H 2 O. 12. On account of its cheapness, slaked lime, Ca(OH)2, a compound very similar in properties to caustic potash or soda, is ordinarily used with the sal ammoniac to liberate the ammonia. The reaction is seen below : 2 NH 4 C1 + Ca(OH) 2 = 2 NH 3 + CaCl 2 H 2 O. FIG. 18. Preparation of Ammonia. m, , o, cylinders containing solutions of impure ammonium chlo- ride, as obtained from coal-gas factories, mixed with lime ; S, S, S, stirrers to keep the lime from settling ; F, furnace to heat the mixture and expel the ammonia: P. B, condensers for cooling the ammonia gas ; C, cylinder of pure water to absorb the ammonia and thus form aqua ammonia ; 7), trough of acid to combine with any fumes escap- ing from C- In this trough, if hydrochloric acid is used, there would form ammonium chloride. EXPERIMENT 42. To illustrate the preparation of ammonia. Put about a half gram of sal ammoniac, NH 4 C1, into a test-tube and add to it about 1 cc. of water, then a little caustic soda or potash solution, 7G MODERN CHEMISTRY and heat gently. Is there any gas given off having an odor? Hold in the mouth of the test-tube a piece of moistened red litmus paper and note the effects. Try also a piece of turmeric paper in the same way. How is it affected ? EXPERIMENT 43. To about 2 g. of ammonium chloride in a tube or flask add 1 or 2 cc. of slaked lime, made by adding a little water to some lime ; adjust upon a ring-stand and attach a delivery tube. Warm the flask gently and collect a jar of the gas by upward displacement, as described in appendix. To tell when the flask is filled, hold near the mouth a piece of red litmus paper, as in the pre- ceding experiment. Keeping the bottle inverted, insert a burning taper up into the bottle. Does ammonia burn? Does it support combustion ? 13. Peculiarities of Ammonia. Ammonia is a colorless gas having a strong pungent odor, and if inhaled in con- siderable quantities produces strangulation and fills the eyes with tears. It is lighter than air, having a density of 0.59 ; it will not support combustion, nor burn in the air; but in oxygen a jet if ignited will continue to burn for some time with a yellow flame. It has remarkable affinity for chlorine, as will be seen when we come to study that gas. It also combines readily with hydrochloric acid, forming dense white fumes. This will be noticed if two bottles, one of each, be opened close together. It is well shown also in the following experiment. EXPERIMENT 44. Put into a bottle two or three drops of strong hydrochloric acid and cover with a glass or paper. Now fill another bottle with ammonia gas and invert over the bottle containing the acid. Remove the cover separating the two and notice the results. 14. Solubility in Water. Ammonia is very soluble in water, as the following experiments will show. EXPERIMENT 45. Fill the bottle again with ammonia gas as before and place it mouth downward into a basin of water. Let it stand two or three minutes and notice whether the water rises in the bottle. NITROGEN AND ITS COMPOUNDS 15. Ammonia Fountain. The most striking illustration of the solubility of ammonia in water is the "ammonia fountain." EXPERIMENT 46. Fit to a round-bottomed flask or strong bottle, of a gallon capacity or more, a rubber cork through which passes a long glass tube that will reach half way to the bottom of this flask and nearly to the bottom of another similar one. Draw out the upper end to a jet. Fasten in position or hold over the lower bottle or jar as shown in the figure. To the water in the lower flask add a few drops of some acid and a little litmus solution, or a few drops of phenol-phthalein solution. Now fill the upper flask with ammonia as in Experiment 44, or by warming gently a solution of strong aqua ammonia the latter will be much quicker and collecting by up- ward displacement as before. When well filled, quickly insert the cork and long jet- tube and support upon the other flask of water, as shown in Fig. 19. In a few sec- onds the water will begin to rise in the tube, owing to the gradual absorption of the ammonia, and will soon flow into the upper flask. The absorption then will be very rapid, and the water will be forced up, forming a beautiful fountain. As it enters the upper flask it will change in color, owing to the effect of the ammonia upon the litmus or the phenol put into the solution. 16. Effect of Heat on the Solubility of Ammonia. At C. one liter of water will absorb about 1150 liters of ammonia. As the temperature rises, the amount absorbed rapidly decreases. This is seen in the fact that if a few cubic centimeters of aqua ammonia in a flask be warmed gently, the gas bubbles out so rapidly that the liquid seems to be boiling vigorously when it scarcely feels more than warm to the hand. FIG. 19. Ammonia Fountain. 78 MODERN CHEMISTRY 17. Effects of Platinum and Charcoal on Ammonia. If a small flask containing strong ammonium hydroxide be warmed gently, and a spiral of platinum wire previously heated to redness be held in the neck of the flask, the wire will continue to glow for a considerable time. Ammonia is also absorbed rapidly by charcoal. EXPERIMENT 47. To show absorption of ammonia by charcoal. Fill a large test-tube with ammonia and place it inverted over an evaporating dish containing a quantity of mercury, as shown in the figure. Slip under the tube a piece of charcoal. In two or three minutes the mercury will begin to rise in the tube to fill the space formerly occupied by the FIG. 20. gaSt 18. Uses of Ammonia. Immense quantities of ammonia are manufactured and used annually. For cleansing pur- poses and for softening or " breaking " water it is found in almost every household. In a medicinal way it is used as a restorative in cases of fainting, and overdoses of chlo- roform and other anaesthetics. Considerably diluted it is employed in neutralizing the effects of acids upon the clothing or upon the hands and face ; in a similar way, by inhaling it cautiously, it will counteract the effects of chlorine, bromine, sulphur dioxide, and similar irritating gases. 19. As a Refrigerant. Perhaps the most extensive use of ammonia is as a refrigerant in the manufacture of ice. The principle underlying this process is as follows : Am- monia may be readily liquefied by moderate pressure ; if this pressure is suddenly removed, very rapid evaporation takes place, producing a low degree of cold. NITROGEN AND ITS COMPOUNDS 79 EXPERIMENT 48. To show the freezing of water by rapid evapora- tion. Put upon a block of wood a few drops of water, and upon this a watch crystal. Into the crystal pour 1 or 2 cc. of carbon disulphide, and blow a current of air by means of a blowpipe across the liquid. By the time the disulphide is all evaporated the crystal will be frozen tightly to the block. Lift the block by taking hold of the crystal. FIG. 21. < 20. Manufacture of Ice. It is upon the same principle that ice is manufactured. The first apparatus for this purpose was devised by Carre, and is shown in the i ii FIG. 22. Carre's Apparatus. accompanying figure. (I) a is a tank containing strong ammonia water, underneath which a fire is placed. This causes the ammonia to bubble out of the water very rapidly, whereupon it flows over into 6, and there liquefies by its own pressure, the water surrounding it keeping it cool, c is a cylinder of pure water fitting into 6. 21. After a half hour or so, when the ammonia has about all been driven out of the solution in a, the posi- tion of the two, a and 6, is reversed (II). A partial 80 MODERN CHEMISTRY vacuum forms in a as it cools, the ammonia in b begins to evaporate to fill the vacuum, and as fast as it flows over into a is absorbed by the water there. The rapid evaporation is thus kept up for a consider- able time, and the cylinder c, containing pure water sur- rounded by the ammonia chamber 6, has its contents frozen. 22. Manufacture of Ice for Commerce. The first ice machines used for the manfacture of ice upon a large scale were made upon this principle. At present, however, in- stead of creating a vacuum by cooling A, pumps are used to remove the vapor from B as fast as it forms. This causes, as in the other class of ice machines, a rapid evap- oration and a consequent cooling of the adjacent water. FIG. 23. Cross-sectional View of Carre's Apparatus. NH B FIG. 24. Modern Ice Plant. 23. Figure 24 will show the essentials of the improved methods of ice manufacture. A is a strong cylindrical tank containing liquid ammonia. is a large rectangular vat filled with strong salt water, through which are coiled a series of pipes, xx, which connect with A. Through the NITROGEN AND ITS COMPOUNDS 81 top of this vat are let down oblong galvanized iron boxes containing the water to be frozen. They are thus sur- rounded by the salt water through which the ammonia pipes, 32?, pass. P is a pump worked by steam, which is continually exhausting the pipes and keeping up a rapid evaporation in A. The pump, at the same time that it exhausts xx, is also condensing the ammonia again in the tank M, from which, at intervals, it is allowed to flow back again by the pipe y into A. In this way the ammo- nia is used over and over without appreciable loss. The rapid evaporation lowers the temperature of the salt water in below the freezing point of pure water, and in from 36 to 60 hours the ice is ready to be drawn from the boxes. 24. Oxides of Nitrogen. There are five of these com- pounds, though not all are of much importance. They Nitrous Oxide, Laughing Gas, or Nitrogen Monoxide, N 2 O Nitric Oxide, Nitrogen Dioxide NO Nitrous Anhydride, Nitrogen Trioxide .... N 2 O 3 Nitrogen Peroxide, Nitrogen Tetroxide .... NO 2 Nitric Anhydride, Nitrogen Pentoxide .... N 2 O 5 The formulae for the second and fourth, for certain reasons, are sometimes written N 2 O 2 and N 2 O 4 . 25. Nitrous Oxide. This is ordinarily called " laugh- ing gas." It has been stated already that many of the compounds of nitrogen are unstable. So, if we heat am- monium nitrate, NH 4 NO 3 , it first melts, then begins to boil, and is decomposed to form nitrous oxide and water, thus : NH 4 N0 3 + heat = N 2 O + 2 H 2 O. 82 MODERN CHEMISTRY EXPERIMENT 49. Put into a test-tube 1 or 2 g. of ammonium nitrate, attach a delivery tube, and suspend upon an iron ring-stand. Heat moderately and collect two or three small bottles of the gas over warm water. Be careful not to heat so strongly as to cause a vigorous ebullition, lest some of the impurities always present in the nitrate may be carried over and thus vitiate the nitrous oxide. When two or three bottles of the gas have been collected, remove the cork and notice the odor. Has the gas any color? Test a bottle of it with a glowing pine splinter as you did the oxygen. What are the results V Try also a small piece of phosphorus ignited ; how does it burn V 26. Peculiarities of Nitrous Oxide. Laughing gas is colorless, somewhat heavier than air, having the odor of sugar when being heated or slightly burned. It is solu- ble to a considerable extent in cold water, will not burn, but supports the combustion of most bodies nearly as well as oxygen. Upon the human system it acts as an intoxi- cant, producing first a sense of hilarity, and afterward unconsciousness. Because of this fact it is frequently used in a purified form as an anaesthetic in dentistry. It is easily liquefied by cold and pressure, and is generally used in this form. 27. Nitric Oxide, NO. This gas is almost always one of the products formed when a metal is treated with nitric acid. EXPERIMENT 50. Into a flask put 2 or 3 g. of copper turnings, and make connections as for collecting oxygen over water. Add a few cubic centimeters of nitric acid, somewhat diluted. What kind of fumes first fill the flask? Notice that they disappear, being carried over and dissolved in the water. Collect three or four bottles of the gas. What can you say of its color and density? Test it to learn whether it will burn. Try a blazing pine splinter, also a burning candle in the gas ; do they continue to burn ? Try also in a deflagrat- ing spoon a well-ignited piece of phosphorus ; what results ? Can you explain ? NITROGEN AND ITS COMPOUNDS 83 28. Peculiarities of Nitric Oxide. As seen above, it is a colorless gas, heavier than air, is non-combustible and a non-supporter of ordinary combustion. It is noticed, however, that substances which burn with great heat, such as phosphorus, sodium, and the like, continue to burn in nitric oxide with great brilliancy. The reason is apparent. Ordinary air is about 20 per cent oxygen ; nitric oxide is about 50 per cent oxygen ; such bodies therefore as have sufficient heat in burning to decom- pose the gas continue to burn more brilliantly, while those which kindle at a low temperature have not the power to use the large proportion of oxygen present. 29. Affinity for Oxygen. The strongest chemical prop- erty of the gas is its great affinity for oxygen. This is seen whenever it is allowed to escape into the air, brown fumes of nitrogen tetroxide being formed. EXPERIMENT 51. Into a bottle of nitric oxide inverted over a basin of water, pass slowly a current of oxygen. This may be gener- ated in a test-tube by using a small amount of potassium chlorate and manganese dioxide, or by treating the latter with sulphuric acid. Notice how the colorless gas changes; what else happens? 30. Two molecules of nitric oxide unite with one of oxygen, O 2 , as follows : 2 NO + O 2 = 2 N0 2 . If two or three drops of carbon disulphide be put into a bottle of nitric oxide and allowed to stand a few minutes, or until the disulphide vapor has filled the bottle, on the approach of a flame the mixture of gases will burn with a brilliant flash, pale violet in color. 31. Nitrous Anhydride, N 2 3 . The term, anhydride, means without water, and is applied to certain oxides, which, when water is added to them, form acids; that 84 MODERN CHEMISTRY is, the anhydride is the acid without the water. Such oxides were formerly called acids, and carbon dioxide is sometimes even now spoken of as carbonic acid ; but all true acids contain hydrogen, and theoretically at least are formed by adding water to the oxide or anhydride. Nitrogen trioxide is thus the anhydride of nitrous acid, and is of interest to us only because of this fact. Thus : N 2 3 + H 2 = 2 HN0 2 . In this case if the oxide is passed into water it is readily absorbed, forming nitrous acid, as shown by the reaction. EXPERIMENT 52. Into an evaporating dish put a little starch, and with a little water rub it to a thick paste. Transfer this to a test-tube into which you have put about 2 cc. of nitric acid. Attach a delivery tube and let the end dip into 15 or 20 cc. of water in a bottle or flask. Heat the starch in the test-tube for some time, or until the fumes are given off readily. What color are they? When the gas ceases to come over, test the solution in the bottle with blue litmus paper to learn whether it is acid in character. You should thus obtain nitrous acid from the brown fumes of nitrogen trioxide which were driven off. Let us test it to determine. Put one or two cubic centimeters of the solution into a test-tube and add a few drops of a solution of ferrous sulphate, made by dissolving a crystal of the salt in water. Does it turn brown in color ? If so, nitrous acid is indicated. 32. Instability of Nitrous Acid. Although it is char- acteristic of nitrogen compounds to decompose readily, nitrous acid is more unstable than most of the others. In fact, it breaks up of its own accord at ordinary temperatures. EXPERIMENT 53. Put a part of the nitrous acid prepared above into a test-tube, and when it has been standing a few minutes, or when gently warmed, hold a sheet of white paper behind the tube and notice carefully whether brown fumes are being given off. Continue to heat gently for a few minutes, or until these do not seem to appear, and test with blue litmus paper. Is the solution still acid ? If so, test NITROGEN AND ITS COMPOUNDS 85 FIG. 25. a part of it to determine whether it is still nitrous acid. If it is, warm it a little longer and test again. If not, test it for nitric acid. This is usually done thus : To about 1 cc. of the solution to be tested add about as much strong sulphuric acid; shake the two together and cool well by holding the tube in a stream of cold water. Next prepare a fresh solution of ferrous sulphate, and pour it very cautiously upon the solution to be tested so as not to mix them. To do this it will be necessary to hold the two tubes almost in a horizontal position, as shown in the figure, and let the ferrous solution run slowly upon the other. Set aside the tube in a vertical position and let it stand for a few minutes, when a dark brown ring should have formed at the junction of the two liquids. The test requires consid- erable care, but is very satisfactory when well done. The test is not distinctive, however, if nitrous acid is present, as this also will form a ring ; but the latter ring is usually much broader and forms much more quickly. A simpler plan is to drop a crystal of ferrous sulphate into the solution to be tested, and then pour down the side of the tube upon it a little sulphuric acid. A brown ring forms about the ferrous sulphate. 33. Nitrogen Tetroxide, N0 2 . This gas is of little importance, yet it is one frequently seen in the action of nitric acid upon metals in the presence of air. We noticed that when nitric oxide was exposed to the air it quickly turned brown. So when a metal is treated with ordinary nitric acid, the nitrogen dioxide at first formed quickly unites with oxygen from the air and forms the brown fumes of nitrogen tetroxide. For experimental purposes it may be prepared directly by heating almost any nitrate. 34. Characteristics of Nitrogen Tetroxide. It is at ordi- nary temperatures a brownish red gas, heavier than air, having a very offensive suffocating odor; is non-com- 86 MODERN CHEMISTRY bustible and a non-supporter of ordinary combustion. It is soluble in water, as may be seen by inverting a bottle of it over water. The brown fumes will disappear and the water will rise in the bottle. 35. Nitrogen Pentoxide, N 2 5 . The only fact of inter- est in connection with this compound is its relation to nitric acid, of which it is the anhydride. Hence it is often called nitric anhydride. The relation is exhibited in the following reaction : N 2 5 + H 2 = 2 HN0 3 . Hi 36. Nitric Acid, HN0 3 . When a strong electrical dis- ' charge takes place in the air, as in the case of violent thunder storms, small quantities of the nitrogen and oxy- gen are caused to unite, forming nitrogen oxides, which dissolved in the falling rain form nitric acid. This is sometimes in appreciable quantities. Compounds of nitric acid, such as sodium and potassium nitrate, are found in abundance, especially the former. These salts are now known to be produced by the action of certain bacteria upon nitrogenous matter, and in some countries the sodium nitrate needed is prepared by introducing these bacteria. 37. Formation of Nitric Acid. Nitric acid may be ob- tained by decomposing any nitrate with sulphuric acid. EXPERIMENT 54. Put into a retort 40 or 50 g. of sodium nitrate, NaNO 3 , cover with concentrated sulphuric acid, and insert the long neck of the retort into a flask surrounded with ice and salt, as in Fig. 26. Instead, the retort may be connected with a short Liebig condenser, kept cool by a stream of water. Apply a moderate heat to the retort ; nitric acid will distil over and condense in the receiver. The reaction may be represented in two ways, according to the amount of Chile saltpeter used : 2 NaNO 3 + H 2 SO 4 = Na 2 SO 4 + 2 HNO 3 . NaNO 3 + H 2 SO 4 = NaHSO 4 + HNO 3 . NITROGEN AND ITS COMPOUNDS 87 FIG. 26. Preparation of Nitric Acid. 38. It will be remembered that we prepared ammonia, a volatile compound, by treating a salt of ammonia with caustic lime, a compound of similar properties which is not volatile. In exactly the same way nitric acid may be easily expelled from a liquid by heating, while sulphuric acid cannot be. The latter, therefore, simply takes the place of the former in combination with the metal, and the nitric acid boils out and condenses in the receiver. This is shown in the above reactions. 39. Characteristics of Nitric Acid. Aqua fortis, as this acid is frequently called, is colorless when pure, though, owing to impurities present,, it is generally slightly yellow- ish in color. It is a volatile acid and gives off fumes which are very irritating. It colors the skin and finger nails yellow, and the color is intensified rather than removed by the application of ammonia. Like other strong acids, it attacks all organic matter, rapidly destroys the fibres of clothing, and the discoloration of the cloth cannot be removed by the application of any alkali, as is the case with other acids. Though a comparatively stable compound, a flask of it exposed to bright sunlight, or 88 MODERN CHEMISTRY heated, soon becomes filled with a brownish gas, nitrogen peroxide, NO 2 , and oxygen. This is seen in the follow- ing reaction : 2 HN0 3 + heat = O + H 2 O + N 2 O 4 . On account of this property of giving up a part of its oxygen with moderate ease it is frequently used as an oxidizing agent. This is seen in the following experi- ments : EXPERIMENT 55. Warm slightly a little turpentine in an evap- orating dish and pour upon it some strong or fuming nitric acid. Usually only a copious evolution of fumes is the result, but sometimes the oxidation is so rapid that the whole mass bursts into a flame. EXPERIMENT 56. Heat in an iron spoon a few small pieces of charcoal ; when red hot drop quickly into a beaker containing some strong nitric acid. Notice that the charcoal continues to glow for some little time, owing to the oxygen obtained from the acid ; notice also that brown fumes fill the beaker. Upon a small quantity of warm strong nitric acid in an evaporating dish, drop a very small piece of phosphorus. Notice that it is instantly set on fire, and small particles are thrown out in all directions. EXPERIMENT 57. To a little tin-foil in a test-tube add some strong nitric acid and heat. Notice that the metal is not dissolved, but con- verted into a white solid, which is really an oxide of tin in combina- tion with water, SnO 2 , H 2 O. By heating, the water is evaporated and the oxide remains. 40. Uses of Nitric Acid. Nitric acid finds a great many uses in the laboratory, frequently as an oxidizing agent, as will be seen from time to time. It is used con- siderably in the manufacture of sulphuric acid, which will be described later, and in making nitro-glycerine and other explosives. 41. Aqua Regia. This is a mixture of nitric and hydro- chloric acids in the proportion of one of the former to three of the latter, and is so named because it will dis- NITROGEN AND ITS COMPOUNDS 89 solve gold, the "king of the metals." It is the strongest solvent known, and attacks several metals which are unaffected by single acids. 42. Nitro-glycerine and Dynamite. Nitro-glycerine is prepared by treating glycerine with a mixture of fuming nitric and sulphuric acids. It is in the form of a liquid, and hence not convenient for uses under all circumstances. Dynamite differs from nitro-glycerine in that it contains about 25 per cent of siliceous or infusorial earth. It is, therefore, more convenient and less liable to explode by accident. Guncotton is a similar compound which is pre- pared by treating cotton wool with nitric and sulphuric acids. It is, therefore, not very different from nitro-glycer- ine in composition. It has the advantage of being perfectly safe when wet, and is, therefore, kept damp when carried on board men-of-war. In this condition it is exploded by igniting with a small charge of fulminating mercury. Its combustion is five hundred times as rapid as that of the best gunpowder. The heavy charges now used for tor- pedoes give an impact that no man-of-war can withstand. All of these explosives, as well as gunpowder, are valuable because of the great instability of the nitrates present or formed in the preparation of them. ARGON : A = 40 ? 43. Its Discovery. For some time previous to the dis- covery of argon, in 1894, it had been observed that nitro- gen obtained from the atmosphere was heavier than that from its compounds. In that year Lord Rayleigh and Professor Ramsay observed that, by passing atmospheric nitrogen over red-hot magnesium, a small residue was obtained which could not be made to enter into combina- tion. This residue was the new gas now called Argon. 90 MODERN CHEMISTRY Its name comes from the Greek word argon, which means idle or inactive. 44. Characteristics of Argon. This element is an odor- less, colorless gas, somewhat heavier than air, constituting about eight-tenths of one per cent of the atmosphere. As far as is known it is a perfectly inert substance, hitherto resisting all attempts to make it enter into combination. No compounds of the gas being known, it is impos- sible to assign it a positive atomic weight, but it is believed to be about forty. SUMMARY OF CHAPTER Origin of the term nitrogen. Abundance of the element and of certain compounds. Easiest method of preparing nitrogen. What is the purpose of the phosphorus ? The source of the nitrogen ? Would a candle do as well as phosphorus ? Why ? Two other ways of preparing nitrogen. Chemical action in each case. Characteristics of nitrogen. Compare with oxygen and hydrogen How similar How different. How test each ? Compounds of nitrogen. Ammonia How formed in nature. Old method of preparing " hartshorn." Present source of ammonia. Wherein are these three methods similar ? Characteristics of ammonia. Experiments to illustrate these. Uses. Experiments to illustrate the most impor- tant. Carre's ice machine. Present ice machines. THE ATMOSPHERE 91 Oxides of nitrogen. Names and formulae. Most important. Why? Method of preparing this one. Use Physiological effects. Acids of nitrogen. Names and formulae. Anhydride of each Meaning of. How distinguish each by test. Preparation of nitric acid. Characteristics and uses of. Aqua regia. Explosives Explanation of their explosive character. CHAPTER VIII THE ATMOSPHERE 1. What it is. We are living at the bottom of an ocean as wonderful as the watery one that washes the shores of our continent. The atmosphere covers the entire earth to a depth variously estimated at from fifty to two hundred miles. Some of the recent investigators believe that, in an extremely attenuated form, the air extends through space, even reaching and commingling with the atmospheres of other planets. Centuries ago the air was regarded as one of the elements, just as water was, and the other gases, as discovered, were all called air ; for example, hydrogen was known as inflammable air, carbon dioxide as fixed air, etc. So the perfumes that were exhaled from various flowers were regarded as air, slightly changed in some unknown way. 2. Constituents of the Air. We know now that the air is not an element, but a mixture of several substances. Three of these, nitrogen, oxygen, and argon, are con- 92 MODERN CHEMISTRY stant, but the watery vapor and carbon dioxide vary from time to time. Many efforts have been made to learn whether the air is a compound of oxygen and nitrogen, mixed with the other constituents named. Analyses have been made in all parts of the world, thousands of feet above the earth, in the crowded cities, on the North and South American prairies ; but though the proportion of the gases, 79 of nitrogen to 21 of oxygen, by volume, is found in all cases approximately the same, yet the varia- tion is too great to permit one to believe that they are united to form a compound. The argon constitutes about 1 per cent of what has usually been taken as nitrogen, or about 0.8 per cent of the air. The carbon dioxide varies somewhat, but seldom amounts to more than three or four parts in 10,000, except in poorly ventilated rooms. The aqueous vapor varies greatly. When the air contains all it is able to hold, it is said to be saturated, or to contain 100 per cent. Ordinarily, however, the humidity is not above 60 to 70 per cent. The amount may be estimated by passing a certain volume of air through a tube filled with calcium chloride and noting the increase of weight. 3. Diffusion of Gases. We find that the air contains five gases, of densities ranging all the way from nine to forty times that of hydrogen. Were it not for the law of diffusion, we should find the argon, perhaps, nearest the ground. The next above this, forming a layer twelve feet or more deep, would be the carbon dioxide; then the oxygen, nitrogen, and water vapor, in the order named. Such conditions would be fatal to all animal life. As it is, however, owing to the constant circula- tion of the atmosphere and the rapid diffusion of gases, no more carbon dioxide is found close to the surface of the earth than hundreds of feet above. Two or three THE ATMOSPHEEE 93 exceptions to this ought to be noted, among them the deadly Upas Valley, where the carbon dioxide is exhaled from volcanic sources more rapidly than diffusion can carry it away.* 4. Boyle's Law. Many years ago Boyle discovered and formulated the law, which now bears his name, that the volume of a gas, the temperature remaining constant, varies inversely as the pressure. In other words, if we double the pressure, the volume decreases by half ; or if we lessen the pressure by half, the volume becomes twice as great. In the accompanying figure we have 10 cc. of the gas, a, under the pressure of the atmos- phere, simply confined by the mercury in the bottom of the bent tube ; if now we pour in more mercury at the open end, the volume of a will constantly decrease, and when we have added as much as corre- sponds to the pressure of an additional atmosphere, the volume will have decreased to 5 cc. 5. Standard Pressure. Atmospheric pressure is meas- ured by the barometer, which at sea level stands about 30 in. high. In chemical calculations, however, we use the metric system, and the equivalent of 30 in. is 760 mm. Hence when we say that a gas is under standard pressure, we mean 760 mm. * This valley is located in the island of Java, is about a half mile in circumference and thirty-five feet deep, surrounded at no great distance by hills. The bottom is comparatively smooth and is devoid of vegeta- tion. Loudon, in describing his visit there, says that " skeletons of human beings, tigers, pigs, deer, peacocks, and all sorts of birds" are to be seen everywhere, bleached by the exposure till they are as white as ivory. A fowl thrown in died in one and a half minutes. 94 MODERN CHEMISTRY PROBLEM. 500 cc. of oxygen under standard pressure would be how much under 750 mm.? As the pressure has decreased, the volume would have increased. We would solve then by the fol- lowing proportion : F: F'::P':P; or 500 cc. : a::: 750: 760. If desired, this problem may be solved without using a proportion. As the pressure has decreased, we know that the volume will be corre- spondingly increased ; that is, V will equal |f of F; or V = m x 500. V 1 = ? 2. What volume would 300 cc. of hydrogen, at 750 mm. pressure, occupy at 780 mm. pressure ? 3. 25 liters of air at 380 mm. pressure, would be how many at 5 atmospheres' pressure V 6. Law of Charles. Just as heat causes solids and liquids to expand, so it affects gases. In the case of the latter the rate of expansion is practically constant, and is in the Law of Charles stated thus : The pressure re- maining constant, all gases expand or contract uniformly under the same increase or decrease of temperature. This has been studied carefully, and it has been proven that for an increase or decrease of 1 C., a volume of gas ex- pands or contracts 2T_ of the volume it occupies at C. To illustrate, suppose we have in a vessel 273 cc. of oxygen at C. If by any means the temperature is raised to 10 C., the volume would increase -^fa of 273 cc. or 10 cc., and would occupy 283 cc. It obviously follows from this that were the law to hold true, and were the gas reduced to a temperature of 273 below zero, it would disappear entirely. However, all gases thus far tried THE ATMOSPHERE 95 become liquids before reaching this low temperature, so that the law no longer applies. 7. Absolute Zero. From the fact that a gas would disappear entirely at 273 below zero, according to the Law of Charles, 273 has been called absolute zero, the point at which the molecules of a body would have no vis viva, or absolutely no heat energy. This point has never yet been reached, though recent investigators have approached within a few degrees of it. It is necessary to have a clear understanding of what is meant by the abso- lute zero, as it is used in making all calculations for correction of the volume of gases for temperature. 8. The Absolute Thermometer. In Fig. 28 we have the Fahrenheit, Centigrade, and Absolute thermometers represented in F, 0, and A, respectively. It must be re- membered that the last is not a thermometer really in existence, but serves merely for illustra- tion. The boiling points on the three are marked 212, 100, and 373; the freezing points 32, 0, and 273. The absolute zero therefore would }>e the same as 273 on the Centigrade, as the degrees on these two are of the same size. Let us apply this in a problem. o 10$ 273- A 373- 273 o- -Freezing Pt. FIG. 28. PROBLEM. 500 cc. of oxygen at C. would occupy what volume at25C.? Expressing Charles's Law in the form of a proportion, we would have F:F::*:f, 96 MODERN CHEMISTRY in which V and t represent the volume and temperature of the gas at the beginning of the experiment, and V and t' at the end ; and it must be remembered that t and t' always mean absolute temperatures. Applying this to the problem, we see that t = C., or 273 A., and t' = 273 + 25, or 298 A. Substituting, we have : 500 : V : : 273 : 298. v , _ 298 x 500 273 PROBLEM 2. What volume would 250 cc. of gas at 20 C. occupy at -10C.? Here, t = 20 C. = 273 + 20 = 293 A. t = - 10 C. = 273 - 10 = 263 A. F=250cc. Substituting, 250 : V : : 293 : 263. T/ , = 250 x 263 293 PROBLEM 3. If 400 cc. of hydrogen is heated from 15 C. to 30 C., what volume would the gas then occupy? PROBLEM 4. What would be the result in problem 2, if at the same time the barometer fell from 760 mm. to 740 ? This may be solved by first finding the value of F', as shown above at the temperature t', and substituting this in the proportion for determining V under P' pressure. Suppose, in problem 2 above, F' = 225+, then solving for pressure, we would have V" = |fS of 225 + ; or F: F"::P":P'; or 225 : V" : : 740 : 760. T// , = 225 x 760 740 Or the problem may be solved by using a compound proportion : 250 : x : : j ^ + ^ : 273 L ; 293 x 740 x = 263 x 760 x 250. PROBLEM 5. 500 cc. of gas under 4 atmospheres and at - 25 C. would have what volume at 760 mrn. and at 20 C. ? Let the teacher furnish a number of similar problems for practice. THE ATMOSPHERE 97 9. Weight of Air. The weight of a liter of air may easily be found by the following experiment : EXPERIMENT 58. M in the figure is a flask of about 500 cc. capacity. Fit to it a cork with a glass tube somewhat drawn out, as shown. Put into the flask about 50 cc. of water and boil for several minutes, so as to expel all the air. Immediately remove the cork and insert another, not perforated. When the flask has cooled to the temperature of the room, weigh the whole. Suppose this to be m. Remove the cork, thus allowing the air to enter, and again weigh flask and cork. Suppose this to be n. The gain in weight, FIG. 29. n m = wt. of air in flask. To determine the volume of the air contained, take a graduated flask, or cylinder, and fill the flask M with water. Suppose this to be r cc. Then r cc. of air weighs n m grams, from which the weight of 1000 cc. = 1 liter may be determined. 10. Liquefaction of Air. The air is so well known that it is not necessary to say anything regarding its proper- ties. At the present time, however, considerable atten- tion is being given to it in the liquid form. A large number of experiments with it have been made by Dewar, Pictet, Linde, Tripler, and others, with a view to ascer- taining its properties and practical value. It is said that the first ounce of liquid air ever produced cost about 13000 and the next pint about $80; with improved methods, however, it may now be prepared for a few cents per gallon. 11. Dewar's Bulbs. Dewar has invented a double- walled glass globe in which liquid air may be kept for a number of hours with little loss; here in this country 98 MODERN CHEMISTRY it is often shipped several hundred miles in large double- walled tin cans, heavily lined with felt, but at the expense of 20 to 40 per cent of the liquid. The Dewar bulbs vary somewhat in construction, but the general plan is the same in all. Into the space between the inner and outer walls of the globe, a drop or two of mercury is introduced; the air-pump is then attached, and a vacuum of very high degree obtained. F,a. 30-Dewar's Buibs. As the air is P Um P ed OUt ' the mercury vaporizes and fills the space. When liquid air is introduced into the inner globe, the mercurial vapor is condensed upon the outer surface of the inside flask, and forms a perfect mirror. Thus we have not only a vacuous chamber, but also a mirror to prevent the access of heat rays to the liquid air, and the insulation is well nigh perfect. A modified form of this Dewar bulb, holding about two gallons, is now used for shipping liquid air. The insulation is so perfect that the liquid may be kept two weeks with little loss. It is obvious that the ordinary closed tank is unsuit- able on account of the high pressure which would soon obtain. 12. Linde's Apparatus. The plan used for liquefying air may be understood from the accompanying figure, which represents the apparatus used by Linde. P is a pump which, when the piston is raised, opens a valve at G- and allows the air from D to enter; as the piston descends, the valve Gr closes and H opens. The air is thus forced up through the coils in the tank J", kept cold by running water, and passes on through B. At C the THE ATMOSPHERE 99 pipe B enters within a larger one, and continues thus until at the point E it again emerges. The ingoing cur- rent of air flows" through the inner pipe under pressure and issues from a small aperture at R into a chamber, T, under low pressure. As expansion is a cooling process, the air is thus reduced in temperature ; at the next stroke Used by Courtesy of the Scientijic American. FIG. 31. Linde's Apparatus for liquefying Air. of the piston the vacuum formed in the pump again opens the throttle valve at 6r, and the cooled air in T flows back through the outer pipe, back through D. As this opera- tion is constantly repeated, the outgoing current being cooled by its expansion into T, continually lowers the temperature of the ingoing current, until finally liquid air will trickle down into the chamber T, and may be drawn off at Fmuch the same as water from a reservoir. 100 MODERN CHEMISTRY 13. Effects of Liquid Air upon Certain Substances. It is found that such articles as rubber, beefsteak, eggs, etc., immersed in liquid air, become exceedingly brittle ; while an ordinary tin cup dipped into the liquid and dropped upon the floor breaks into fragments like glass. All these effects are due to the intense degree of cold of the liquid air, and not to any chemical action. 14. As the boiling point of nitrogen is lower than that of oxygen, the former boils out the more rapidly, and in a short time a vessel of liquid air, freely exposed, will contain almost pure liquid oxygen. If into this a red- hot iron rod be thrust, it will burn vigorously, notwith- standing the fact that the temperature of the surrounding liquid is nearly 1700 C. below the melting point of iron. It should be said, however, that the two are probably not in contact, but that a layer of gaseous oxygen next to the iron rod supports the combustion. Felt, saturated with liquid air, burns explosively, and if confined in metal tubes, bursts them with violence. 15. Practical Uses of Liquid Air. Numerous applica- tions for liquid air have been suggested, but as yet these are in the experimental stage. Among them may be named the following : (1) as a substitute for compressed air ; (2) as a refrigerant ; (3) in blasting ; (4) in surgery for removing diseased tissues without the use of the knife ; (5) as a smoke consumer, and for burning garbage in cities. SUMMARY OF CHAPTER Composition of the atmosphere. Old ideas of the air. Present ideas. Explanation of uniformity of composition. Boyle's Law Statement of. Meaning of term standard pressure. THE HALOGENS 101 Charles' Law statement o. - Meaning of term absolute ^^\ *' Problems. Density of air. \ , //' i j *" '- -" Methods of finding weight of one liter. Liquefaction of air. Present method. Dewar bulbs. Effects of liquid air. Suggested uses. CHAPTER IX THE HALOGENS 1. Members of the Group. The term halogen is from two Greek words, meaning salt producer, and is given to this group of elements because with the metals they form a large number of salts. The group includes fluorine, chlorine, bromine, and iodine. The first two are gases, the third a liquid, and the fourth a solid. They possess prop- erties very similar to each other, differing in degree rather than otherwise. It will be found that as the atomic weights increase, the chemical activity decreases. FLUORINE : F = 19 2. Characteristics. Fluorine is an element which had not been prepared until a few years ago. It is a greenish- colored gas, of a very irritating odor, and readily attacks almost all substances. By extreme cold and pressure it has been liquefied, and when in that condition loses much of its chemical activity. It is of little practical value, and is considered only because of one or two compounds which it forms. 102 MODERN CHEMISTRY 3. . Compomid of Fluorine. There is only one com- pound cf this elem-8'ai m which we are specially interested, a.ud. ;that is hy4roflu9ric acid, HF. It is prepared by treating iludr spar, calcium fluoride, with strong sul- phuric acid, the reaction being CaF 2 + H 2 S0 4 = CaS0 4 + 2 HF. Hydrofluoric acid is a very irritating, colorless gas, which readily dissolves or corrodes glass, and hence is sometimes used in glass etching. EXPERIMENT 59. Warm a sheet of glass 3 or 4 in. square by holding it at some height above the burner flame, and drop upon it a few shavings of paraffine. Move the glass about so as to distribute the melted wax evenly, and allow it to cool. Now with a sharp pen- cil or stylus draw any desired figure in the wax, being sure to cut through to expose the glass. Lay this face down over a lead saucer,* into which you have put about 2 g. of calcium fluoride and as much strong sulphuric acid. Support upon a ring-stand and warm for a minute very gently, so as not to melt the wax. In a few minutes the etching should be completed. This can be determined by testing with the point of a knife blade, when the glass will feel rough where the figure was drawn in the wax. When the experiment is finished, the paraffine may be removed with a dull knife or by immersing in warm water. CHLORINE: 01 = 35.5 4. History. Chlorine, the most important element of the halogen group, was first prepared by Scheele in 1774, in treating black oxide of manganese the same com- pound we have used in preparing oxygen with hydro- chloric acid. He did not know, however, that he had discovered a new element, but supposed it to be a com- * Instead of the lead saucer a small evaporating dish may be used. If so, notice whether it also is attacked on the inside by the hydrofluoric acid. THE HALOGENS 103 pound of oxygen and hydrochloric acid, and called it dephlogisticated marine acid air. Hydrochloric acid was then called marine acid. Later, when chlorine was found to be an element, it was given its present name from the Greek word chloros, meaning green. 5. How found. Because of its great chemical affinity, chlorine, like fluorine, is never found uncombined. Its most widely distributed compound is common salt, NaCl, which is found in extensive deposits in nearly all parts of the United States, and constitutes a large per cent of the solids held in solution in the ocean. 6. How to prepare Chlorine. For laboratory purposes the simplest way of preparing chlorine is that used by its discoverer, by treating manganese dioxide with hydro- chloric acid and warming gently. EXPERIMENT 60. Into a good-sized test-tube or generating flask put 1 or 2 g. of manganese dioxide and about 2 cc. of hydrochloric acid. Attach a delivery tube and warm gently. Collect two or three bottles of chlorine by downward displacement, as described on page 362, and preserve for future experiments in studying its properties. 7. The reaction that takes place in preparing chlorine as above may be indicated thus : Mn0 2 + 4 HC1 = C1 2 + MnCl a 4- 2 H 2 O. From this we see that only half the chlorine, in the hydro- chloric acid used, is obtained free, the other half having united with the manganese to form manganese chloride, a compound which has no application in the arts. An immense quantity of chlorine is used every year in the manufacture of bleaching powder, and cost of production is a very important consideration. The method described above is, therefore, not strictly followed commercially, but is so modified that the manganese chloride is converted 104 MODERN CHEMISTRY into the dioxide again. This is much cheaper, and is known as the Weldon process. 8. The Weldon Process. In the preparation of chlorine $ ] for manufacturing processes, pyrolusite, a natural ore of manganese and an impure form of the dioxide, is treated with hydrochloric acid in large stone tanks. When the chlorine is no longer given off, any excess of acid in the residual liquor is neutralized with common limestone, finely powdered. The reaction may be represented thus : MnCl 2 + H 2 + 2 HC1 + CaCO 3 = MnCl 2 + CaCl 2 + CO 2 + 2 H 2 O. /Residual liquor and excess\ , ni mes t one ) = /mixture manganese and cal-\ \ of acid / \ cium chloride in water. / Therefore, we now have a mixture of manganese chloride and calcium chloride in solution. Next, lime water, pre- pared by treating ordinary lime with water, is added. CaO (lime)+ H 2 O = Ca(OH) 2 (lime water). This precipitates the manganese in the form of the hydrox- ide, Mn(OH) 2 , thus : - 1 + Ca(OH) 2 = Mn(OH) 2 + 2 CaCl 2 . ^vg J Now by heating this and at the same time passing a cur- rent of air through the solution, the manganese hydroxide, Mn(OH) 2 , is converted into the dioxide, thus : Mn(OH) 2 + O (air)= MnO 2 + H 2 O. The calcium chloride, being very soluble, remains in solu- tion. The mixture is now allowed to flow into settling basins, where the dioxide is slowly deposited as a dark- THE HALOGENS 105 colored ooze, known as Weldoris mud. This is now ready to be passed again into the stills for a second treatment with hydrochloric acid.* 9. The Chemical Changes in the Above Method. By studying the reaction MnO 2 + 4 HC1 = MnCl 2 + 2 H 2 O + C1 2 , we see that the oxygen in the manganese dioxide has been set free from the manganese and has united with the hy- drogen in the acid ; or, as we sometimes say, the chlorine has been set free by the oxidation of the hydrogen with which it was combined. In like manner other substances, besides manganese dioxide, may be used with hydrochloric acid in preparing chlorine. In every instance the princi- ple is the same : the oxygen is first set free, and, combining with the hydrogen in the acid, liberates the chlorine. Let us prove this. EXPERIMENT 61. Treat a few crystals of potassium chlorate, KC1O 3 , a substance from which we obtained oxygen, with a little hydrochloric acid. Warm very gently if necessary to start the action, and then remove the test-tube from the flame. Notice the rapid evo- lution of gas. With the chlorine thus obtained we have also an oxide of chlorine, C1O 2 , as seen in the reaction 4 KC1O 3 + 12 HC1 = 4 KC1 + 9 Cl + 3 C1O 2 + 6 H 2 O. Add to this a few cubic centimeters of water, which will give a yellowish solution known as euchlorine or chlorine water. Preserve it in a dark-colored, tightly stoppered bottle. * It perhaps ought to be stated that a small excess of lime water usu- ally remains mixed with the precipitated manganese hydroxide. When the current of air is passed through the solution, this lime water, Ca(OH)2, forms with a portion of the manganese hydroxide, calcium manganite, CaMnO 3 , which may be regarded as CaO . MnO 2 . This, with hydrochloric acid, yields chlorine, as does manganese dioxide. 106 MODERN CHEMISTRY 10. It will be remembered that we prepared oxygen also by using potassium dichromate. If now we treat this compound with hydrochloric acid, chlorine is obtained as in the other instances. 11. Practical Application of this Principle. In all the above methods the chlorine is set free by bringing into contact with hydrochloric acid some highly oxygenized substance which will give up a part of its oxygen to unite with the hydrogen of the acid. Hence was conceived the idea of using atmospheric oxygen as the most economical source of supply. 12. Deacon's Process. This idea is applied in Deacon's process. Theoretically the reaction that takes place ac- cording to this method is as follows : 2 HC1 + O = H 2 + C1 2 . In reality, however, the process is not so simple. In the preparation of oxygen from potassium chlorate and man- ganese dioxide, we have seen that the latter compound remains unchanged. It acts, as was said, by catalysis, in a manner not thoroughly understood, causing the potas- sium chlorate to yield up its oxygen at a temperature much lower than would otherwise affect it. 13. The Catalytic Agent. In Deacon's process for the preparation of chlorine some catalytic agent is necessary, because a mixture of oxygen and gaseous hydrochloric acid, when heated, is only slightly decomposed. As a catalytic agent some such compound of copper as the sulphate or the chloride is used. Clay balls or bits of brick are saturated with the copper solution and placed in an iron pipe called the decomposer. Through this the mixed gases, air and hydrochloric acid, previously heated to about 500 C., are made to pass. The acid is oxidized THE HALOGENS 107 and the chlorine set free. The chemical action of the cuprous chloride, Cu 2 Cl 2 , is not thoroughly understood ; but it is believed that two or three reactions take place, in the course of which cupric chloride, CuCl 2 , is formed, which at the temperature present is unstable and gives up a part of its chlorine, leaving cuprous chloride again. 14. Another Method of preparing Chlorine. Another method is frequently employed in the laboratory instead of the first one given. EXPERIMENT 62. Into a test-tube put a small quantity of common salt, NaCl, mixed with a little manganese dioxide, and about a cubic centimeter of sulphuric acid. Warm gently. Is there any evidence that chlorine is being generated ? 15. Comparison of the Two Methods. We shall find that when common salt is heated with sulphuric acid they react with each other, forming hydrochloric acid. That is what we have done in this case. We see by comparing the two reactions, and Mn0 2 + 4 HC1 = MnCl 2 + 2 H 2 O + C1 2 2H a SO 4 =MnSO 4 +Na a SO 4 +2H a O+Cl a , that in the first case we treated the dioxide directly with hydrochloric acid, but in the second indirectly by the use of two substances, which in reacting prepare the hydro- chloric acid needed. It will be seen, however, in the sec- ond instance that all the chlorine is set free, while in the first only one-half. 16. It is probable that in the second case the reaction is a little more complicated, perhaps as follows : First, a part of the sulphuric acid reacts with the com- mon salt, forming hydrochloric acid, thus : 2 NaCl + H 2 S0 4 = 2 HC1 + Na 2 SO 4 . 108 MODERN CHEMISTRY Then another part reacts with the manganese dioxide also present, setting free oxygen, as we have seen before, thus : MnO 2 + H 2 SO 4 = MnSO 4 + H 2 O + O. Then this nascent oxygen immediately attacks the hydro- chloric acid present, oxidizing it and liberating the chlorine, thus : 2 HC1 + O = H 2 O + C1 2 . Putting these three reactions together, we would have 17. Experiments with Chlorine. With the chlorine prepared make the following experiments in study of its properties : EXPERIMENT 63. Note the color of the ga s ; the odor. Put a burning match into a bottle of chlorine ; try also a burning candle. State the results. Does the gas burn? Does it support combustion? EXPERIMENT 64. To show its chemism for certain metals. Sift into a bottle of chlorine, by means of a fine wire-gauze spoon, some powdered metallic antimony; try in the same way metallic arsenic. Describe the results. EXPERIMENT 65. To show the chemism of chlorine for hydrogen. In a room partially darkened, fill a strong bottle with chlorine and hydrogen, mixed. Wrap a towel about it, ignite a piece of magnesium ribbon, and bring it toward the mouth of the bottle. A violent explo- sion is the result. Bright sunlight has the same effect. Try also the following experiment to show the same fact. EXPERIMENT 66. Attach a jet to a hydrogen generator, H, and when it has been in action long enough to expel all the air, ignite it, and insert into a jar of chlorine, C, as shown in Figure 32. Does it continue to burn? How does the flame change in appearance ? What becomes of the green gas ? After a few moments add about 1 or 2 cc. of water to the gas, and shake well. Drop into the solution a piece of blue litmus paper ; what is indicated ? It is best to dry the hydrogen by passing through a drying tube, D, filled with calcium chloride. THE HALOGENS 109 FIG. 32. EXPERIMENT 67. To show affinity of chlorine for hydrogen in compounds of the latter. Into a jar of chlorine thrust a narrow slip of blotting paper which has been moistened in moderately warm tur- pentine. State the results. Turpen- tine consists of carbon and hydrogen, C 10 H 16 . What has the chlorine really done ? EXPERIMENT 68. Practical appli- cation of the preceding experiment. Into a jar of chlorine pour a few cubic centimeters of any solution containing organic colors, as litmus, logwood, or carmine. Shake it up and notice the effects. EXPERIMENT 69. With the same purpose as in Experiment 68. In another jar of chlorine, suspend a piece of blue or pink calico mois- tened with water. Try another simi- lar piece without moistening it. Are the results different? EXPERIMENT 70. To show affinity of chlorine for ammonia. At- tach to a small flask, into which you have put 25 or 30 cc. of strong aqua ammonia, a delivery tube with jet attached. Warm the flask gently as in preparing ammonia for the "fountain," Experiment 46, and insert the tube into a bottle well filled with chlorine. What happens? What becomes of the chlorine? 18. Characteristics of Chlorine. Chlorine is a greenish yellow gas, with a very irritating odor, producing tem- porarily a catarrhal affection of the nasal passages. It is somewhat soluble in water, forming a solution yellowish in color, with the characteristic odor of chlorine. This solution is, however, unstable, as the chlorine gradually combines with the hydrogen of the water to form hydro- chloric acid, while the oxygen is set free. Chlorine is about two and a half times as heavy as air and does not support ordinary combustion. It will be found, however, 110 MODERN CHEMISTRY that sodium and phosphorus, when well ignited, burn vig- orously in an atmosphere of chlorine. EXPERIMENT 71. Put a small piece of sodium, heated in a defla- grating spoon until it takes fire, into a bottle of chlorine. State results. Notice the white deposit of common salt that forms. In the same way try a piece of phosphorus, without first igniting it. State results. 19. Chemical Affinity of Chlorine. From our experi- ments in oxygen we learned that considerable heat was necessary to effect its rapid union with any other element. The iron wire, the sulphur, and the phosphorus, all had to be raised to the kindling point. In the case of chlorine we find that union often takes place at ordinary tem- peratures, showing its chemism to be far greater. Thus arsenic and antimony sprinkled into the gas took fire spontaneously, as did also the phosphorus and the tur- pentine. In the latter case the chemical action is due to the affinity between the hydrogen in the turpentine and the chlorine ; the same remarkable affinity of these gases for each other was also seen in exploding the mix- ture of the two by means of light, and in the hydrogen jet which continued to burn in the chlorine. 20. Chlorine Hydrate. If a saturated solution of chlo- rine water be surrounded by a mixture of ice and salt, in a few minutes yellowish crystals of chlorine hydrate, represented by the formula Cl, 5 H 2 O, are formed through- out the liquid. Chlorine may be liquefied at 34 C. under ordinary atmospheric pressure, or at with a pressure of six atmospheres. In this condition it is of a bright yellow color. It has also been solidified by reducing to 102 below zero, and in this form closely resembles the liquid in color. THE HALOGENS 111 21. Uses of Chlorine. Chlorine is used to a consider- able extent in the extraction of gold from its ores, because it is a good solvent of that metal. A large amount of that now used for this purpose is put up at the factories in the liquid form in steel cylinders lined with lead, and then shipped wherever desired. 22. As a Bleaching Agent. Chlorine is a powerful bleaching agent, but acts indirectly. We noticed that dry calico was but little affected by chlorine. The reason for this is that chlorine in its great chemism for hydrogen abstracts it from the water, and the nascent oxygen unites with the coloring matter of the cloth, converting it into colorless compounds ; whereas in the dry cloth there was" comparatively little moisture to furnish the necessary oxygen. 23. Its most extensive use in manufactures is in bleach- ing cotton and linen goods and paper pulp. Here, how- ever, it is used in the form of bleaching powder. This is a compound, which when treated with dilute acid readily gives up its chlorine. The following diagram will illus- trate the method employed in bleaching cloth. FIG. 33. Cloth-bleaching Apparatus. The cloth is seen in a roll at A ; from here it passes down under rollers at the bottom of the vat B, which con- tains bleaching powder in water, next up over rollers and down into a second vat containing dilute hydrochloric acid, into a third vat with bleaching powder, and so on until the cloth is sufficiently bleached. The excess of chlorine 112 MODERN CHEMISTRY must now be removed, or it will attack the fibers of the cloth and make them weak. To prevent this the cloth is drawn through another vat, jD, containing an antichlor ; that is, a solution which combines with the chlorine still present and forms such compounds as will not attack the fibres. For this, sodium thiosulphate is frequently used. Then, after passing through a vat of water for washing, the cloth comes out pure and white= Chlorine is also used to some extent as a disinfectant, but generally in the form of bleaching powder for this purpose also. HYDROCHLORIC ACID, HC1 24. History. This acid, sold usually under the name muriatic acid, has been known for four centuries, and was formerly called spirit of salt. Later it received the name of marine acid. 25. Where found. It is found uncombined only in very small quantities. It is said to exist in the stomach and to aid digestion, and is sometimes emitted from vol- canoes in eruption. 26. How to prepare Hydrochloric Acid. The method of preparation has already been suggested in one of the experiments for making chlorine. EXPERIMENT 72. Into a generating flask put about 25 g. of sodium chloride, NaCI, and cover with strong hydrochloric acid. Then add sulphuric acid, drop by drop, by means of a separatory funnel, warm- ing gently. Collect two or three bottles of the gas by downward dis- placement and preserve for a study of the properties. Keep them covered to prevent diffusion. The bottle is full when a moistened piece of blue litmus paper held near the mouth is quickly turned red. 27. Manufacture on a Large Scale. The method illus- trated by this experiment is really the one used in prepar- ing hydrochloric acid on a large scale. It is nearly all THE HALOGENS 113 obtained as a by-product in the manufacture of soda crys- tals preparatory to the making of soap. Like many other valuable articles of commerce, it was at one time allowed to go to waste as of no value. 28. In the manufacture of sodium carbonate, common salt was treated with sulphuric acid as above, and the gas obtained was allowed to escape from the flues. But being heavier than the air it settled to the ground, destroying vegetation and rendering all life in the neighborhood almost unendurable. In some places it was produced so abundantly as to corrode even the tools of workmen. It thus became so great an evil that laws were passed pro- hibiting any manufacturer from allowing the escape of such gas, just as the consumption of coal smoke is de- manded in most large cities to-day. An attempt was also made to conduct the gases into streams of water, but this resulted in the death of animals living in the streams. 29. Finally uses were found for the acid, and then plans were thought of and efforts made to save and use it. The gas is conducted into towers filled loosely with coke, down which water is allowed to trickle slowly. In this way the gas is practi- cally all absorbed, and there results a moder- ately strong aqueous solution of hydrochlo- ric acid. Sometimes FIG. <*i. Hydrochloric Acid Factory. the gases are conducted through large Woulff bottles partly filled with water, where solution is effected in the same way. 114 MODERN CHEMISTRY The reaction that takes place may be represented as follows : NaCl + H 2 SO 4 = NaHSO 4 + HC1. If, however, the heat is increased, a larger amount of hydrochloric acid is obtained by using the same amount of sulphuric acid with more salt. Thus : 2 NaCl + H 2 S0 4 = Na 2 SO 4 + 2 HCL 30. Experiments with Hydrochloric Acid. Many char- acteristics of hydrochloric acid may be learned by the following experiments : EXPERIMENT 73. Into a bottle of the gas collected above put moistened pieces of blue and red litmus paper. How are they affected ? Lower a candle into the bottle. What happens? Does the gas burn? EXPERIMENT 74. To show the solubility of the gas in water. Add to a bottle of hydrochloric acid gas a little water and shake for a moment. Hold a piece of moistened blue litmus paper within the bottle. Is it affected? Drop it into the solution. What happens? What has the water done ? Has the solution any taste ? EXPERIMENT 75. Purpose same as the preceding. This is a repetition of the "ammonia fountain" experiment. In pre- paring for it one or two additional points should be noticed. It is better to use appa- ratus somewhat smaller than before, and the gas must be collected by downward displace- ment. The lower flask in this case had bet- ter be fitted with a two-hole rubber coik, through one of which the long tube passes. Through the other should be passed a short tube bent at right angles, as shown in the figure accompanying. When the flask is well filled with gas, make the connections all tight, then blow through the bent tube b to start the flow= Otherwise it will be necessary to wait several Illlllllllllllllllllllllllllllllllllllllllllllllllllllllil FIG. 35. Hydrochloric Acid Fountain. THE HALOGENS 115 minutes before the water will enter the upper flask. The experiment works well, but will be more attractive if the water is colored by lit- mus or some vegetable solution which will change color upon absorbing the acid in the upper flask. A drop or two of ammonia and a few of phenol phthalein in the water serve excellently. The deep purplish red solution becomes perfectly colorless as it enters the upper flask. 31 . Characteristics of Hydrochloric Acid. Hydrochloric acid is a colorless gas, somewhat heavier than air, and has a very irritating odor. It neither burns nor supports com- bustion ; it turns blue litmus paper red, and is very soluble in water. At C. 1 liter of water will dissolve about 500 liters of hydrochloric acid gas. So great is its affinity for moisture that whenever it escapes into damp air, heavy, white clouds appear. 32. The commercial acid, which is simply an aqueous solution of the gas, contains about 32 per cent of acid. Very dilute solutions of hydrochloric acid may be concen- trated by heating until the solution contains 20 per cent of acid, but the process can be carried no further. On the other hand, very strong acid, if exposed to the air, or if heated, loses strength. 33. Hydrochloric acid has great affinity for ammonia ; if a bottle of hydrochloric acid and a bottle of ammonia remain undisturbed side by side for some time, they become thickly coated about the top with ammonium chloride, a white salt formed by the union of the two gases. 34. Uses of Hydrochloric Acid. The chief use of this acid is in the preparation of chlorine for the manufacture of bleaching powder. It is also used very largely in all chemical laboratories as a reagent, in gas works to neu- tralize the ammonia solutions drawn off from the "washer," and in the preparation of various chlorides. 116 MODERN CHEMISTRY BROMINE : Br = 80 35. Where found. Because of its great chemical activ- ity, bromine, like chlorine, does not occur free, but is found in sea water and in salt wells combined with sodium and magnesium. Its discovery dates from the year 1826, when Balard found it in sea water. 36. Commercial Supply. The greater part of the com- mercial supply of bromine is obtained from Germany and the United States. The greater amount used in this country comes from Pomeroy Bend, Ohio, where there are a large number of salt wells. Bromine appears there in the form of magnesium and sodium bromide. The salt water from these wells is boiled down to a certain extent, the common salt (NaCl) crystallizing out, while the other compounds remain in solution. This residue is known as the "mother liquor." The next step in the process is to put the solution into stills hewn out of solid rock, adding to it manganese dioxide and sulphuric acid. The whole is then heated by steam introduced into the liquid through pipes. Bromine distills over and is condensed under water. 37. Formerly bromine was expensive, but, owing to cheaper methods of production, the price has been so re- duced that many of the salt works no longer prepare it. The method of preparation described above is illustrated in the following experiment : EXPERIMENT 76. Into a test-tube put a few small crystals of sodium or potassium bromide, add a little manganese dioxide, and cover with sulphuric acid. Warm slightly and notice the dark red gas given off. What other gas have we prepared that resembles this somewhat? Describe the odor. How does it affect the eyes? Try its bleaching effects upon a moistened piece of calico or litmus paper. How does it compare with chlorine in this respect? Does anything condense upon the cooler portion of the tube? What is its physical condition ? Its color ? THE HALOGENS 11? 38. Laboratory Method of obtaining Bromine. If some bromine is desired for class experiments, it may be pre- pared as above. Attach a delivery tube and conduct the gas into cold water. As soon as the water is saturated, the bromine will condense in the bottom of the jar. It may then be obtained from the water by a separating funnel, or by pouring into a burette and drawing off the heavier liquid as needed for experiment. Preserve both the bromine and the water. 39. The reaction that takes place is the same as in the preparation of chlorine by a similar method, thus : 2 KBr + 2 H 2 SO 4 + MnO 2 = K 2 SO 4 + MnSO 4 + 2 H 2 O + Br 2 . 40. Another Method. Sometimes another method is used when the purpose is merely to determine whether bromine is present in a solution In this process the bromine is set free from its compound by the use of chlorine. EXPERIMENT 77. To the solution supposed to contain bromine add a little chlorine water as prepared in Experiment 61. If bromine is present, the solution should turn darker in color, due to the libera- tion of the bromine by the chlorine. Which does this experiment show to have the greater chemism ? To prove that this color is due to the presence of free bromine add about a half cubic centimeter of carbon disulphide, shake well, and allow it to settle. If free bromine is present, the disulphide will be turned brown from the fact that it has taken up all the free bromine in the solution. 41. Characteristics of Bromine. Bromine is a dark reddish brown liquid. It is the only non-metallic ele- ment that is a liquid. It is very volatile, giving off at all temperatures heavy brown fumes. At seven degrees below zero it solidifies. It has a very disagreeable odor, 118 MODERN CHEMISTRY and attacks not only the throat and nostrils, but also the eyes. It differs from chlorine in that the odor is more sickening, and it was this fact that gave to the element the Greek name bromos, meaning offensive odor. 42. The vapors are non-combustible, yet, like chlorine, they allow of the continued combustion of a jet of hydro- gen. As the hydrogen burns, the red vapors gradually disappear, and colorless hydrobromic acid gas takes their place. Powdered arsenic, sifted into the vapors, burns, and a small bit of antimony dropped upon liquid bromine burns brightly, and the heat generated by the chemical action melts the metal, which spins around upon the sur- face like sodium upon water. Bromine is soluble to a considerable extent in water, and if the temperature of such a solution is reduced by surrounding it with a freez- ing mixture, light brown crystals of bromine hydrate separate, as did the crystals of chlorine hydrate under similar circumstances. 43. Experiments with Bromine. Let the teacher prove the above facts by experiments with bromine before the class. EXPERIMENT 78. Place a small piece of phosphorus in a defla- grating spoon and put it into a jar of bromine vapor. Allow it to remain a few minutes. Does it burn? Compare bromine with chlo- rine in this regard. EXPERIMENT 79. To test the bleaching effects of bromine upon colored solutions. Pour into a bottle a little bromine vapor, and add a few cubic centimeters of logwood, litmus, or carmine solution. Shake it up. Notice the effect upon the color. 44. Uses of Bromine. The principal use of bromine is as a disinfectant. It is also used in organic work in chem- istry and in the preparation of some dyes. For organic colors it is a strong bleaching agent, though not as active THE HALOGENS 119 as chlorine. There are also several compounds which have application in medicine. Of these magnesium bro- mide, MgBr 2 , and potassium bromide, KBr, are the most important; the former is found in the water of many mineral springs and is regarded as of medicinal value ; the latter is used as a sedative in the case of nervous headache. A third compound, silver bromide, AgBr, is used in photography for sensitizing various printing papers, j/ IODINE : 1 = 127 *45. The Source of Supply. Until within recent years, the iodine of commerce was obtained from certain varieties of sea-weeds. These weeds were collected in large quan- tities and burned, and the ashes treated with water to dissolve out the sodium carbonate which was wanted for making soap. If such sea- weeds are burned at a low tem- perature, the iodine will remain in the ashes in the form of sodium and potassium iodide. From these it can be obtained as shown below. 46. The greater part of our present supply of iodine comes from Chile. There is a desert in that country many square miles in area, where are found vast deposits of sodium nitrate mixed with considerable quantities of soil and small amounts of iodine compounds. This mix- ture is treated with water, which dissolves out the sodium nitrate and the iodates; the solution is then evaporated till the sodium nitrate crystallizes out, as in the manufac- ture of bromine in Ohio, leaving the iodine compounds still in solution. The residual solution, called the " mother liquor," is treated with manganese dioxide and sulphuric acid, and gently heated. 47. Preparation for Commerce. When treated as above, from the mother liquor, violet fumes of vaporous iodine 120 MODERN CHEMISTRY are given off abundantly ; they are passed over into cool chambers, where they condense. To further purify the iodine, it is resublimed at a low temperature and con- densed in a series of conical-shaped flasks (see Fig. 36). FIG. 36. Iodine Apparatus. At the left is a small brick furnace, in the upper part of which is an oven. The iodine to be purified is placed in the oven, and gently heated. The final reaction in the separation of the iodine is the same as in the case of the bromine and chlorine. 2 NaI + 2 H 2 SO 4 +MnO 2 =Na 2 SO 4 + MnSO 4 +2 H 2 O + I 2 . The essential features of this method of preparing iodine may be shown by the following experiment : EXPERIMENT 80. Into a small test-tube put a crystal or two of potassium iodide, add a little manganese dioxide, and cover with sul- phuric acid. Warm gently ; notice the fumes that are given off and what condenses upon the cooler portion of the tube. 48. Another Method of preparing Iodine. The follow- ing method is employed to some extent in France in obtaining iodine from the ashes of sea-weeds. It is also the usual method pursued in the laboratory in testing for iodine. The plan consists simply in adding free chlorine THE HALOGENS 121 to the iodine solution, whereby the latter is liberated from its compounds. As a commercial method it is open to the objection that if too little chlorine is added, not all of the iodine is liberated, and if too much, a portion com- bines with the chlorine. EXPERIMENT 81. To a solution containing iodine in combina- tion add a few drops of chlorine water. What change in color takes place ? This indicates free iodine, as may be proven by adding starch paste solution. The starch will turn blue, as it did with ozone. 49. Experiments with Iodine. Many characteristics of iodine may be learned from the following experiment : EXPERIMENT 82. Put a small crystal of iodine into a test-tube and warm gently. What happens? Describe the color and odor of the vapors. Hold a piece of moistened starch paper near the mouth of the test-tube ; how is the starch affected ? Close the mouth of the tube with your finger and notice the stain that is formed. See whether you can remove it by moistening with caustic potash or ammonia. 50. Characteristics of Iodine. Iodine is a solid of a dark bluish black color, with a metallic luster. At or- dinary temperatures it is somewhat volatile, and when gently heated it is readily converted into vapors of a beautiful violet color. It was from this fact that the element received its name, iodine being derived from a Greek word which means violet. The odor of the vapors resembles somewhat that of dilute chlorine, but is less irritating. It has the power of turning the skin yellow, but the stain may be removed by treatment with some alkali. It has feeble bleaching properties, and turns starch paste blue. This is so delicate a test that one part of iodine in several hundred thousand of water will be clearly shown. Its affinity for phosphorus is so strong that if a crystal of iodine be dropped upon a small piece of phosphorus, the latter will be ignited almost instantly. 122 MODERN CHEMISTRY 51. Solvents for Iodine. Among the better solvents for iodine are chloroform, ether, alcohol, carbon disul- phide, and a solution of potassium iodide. EXPERIMENT 83. Put a crystal of iodine into a test-tube with a little cold water. Shake for a moment or two, and then pour off a part of it into another tube and test with starch paste to determine whether any has dissolved. What are your conclusions? Warm the remainder; what indications are there that the iodine is dissolving? Test the solution again with starch paste, or with carbon disulphide, thus : add about a half cubic centimeter of the disulphide to the iodine solution, and shake well. Notice the beautiful violet color imparted to the disulphide. Try alcohol also as a solvent. Before testing the solution with starch or carbon disulphide, dilute until pale yellow in color. What are the results? Try in the same way a solution of potassium iodide upon an iodine crystal, and state results. 52. Uses for Iodine. In the form of a tincture, or alcoholic solution, iodine is used to a considerable extent in medicine to prevent the spread of eruptive diseases, like erysipelas, in skin affections, sore throat, and the like. In the compound iodoform it is used by physicians as a deo- dorant and disinfectant. As potassium or sodium iodide it is frequently used as a reagent in the laboratory, and to a limited extent in making aniline dyes. In these vari- ous ways 300 tons or more are used annually. 53. Some Comparisons. It has probably been observed that the same method is used in preparing chlorine, bro- mine, and iodine. Notice the following reactions : Cl Br MnO 2 + 2 H 2 SO 4 + 2NaCl = MnSO 4 + Na 2 SO 4 + 2 H 2 + C1 2 . MnO 2 + 2H 2 SO 4 + ZNaBr = MnSO 4 + Na 2 SO 4 2 H 2 = MnSO 4 2H0+/. THE HALOGENS 123 SUMMARY F CHAPTER Meaning of term halogen. Names, symbols, and atomic weights of the halogens. Comparison of the halogens. a. Method of preparing. b. Physical condition at ordinary temperatures; at lower tem- peratures. c. Color. d. Odor. e. Density. /. Chemical activity. g. Bleaching powers. h. Affinity for certain substances, as hydrogen, phosphorus, etc. {. Uses. j. Hydrogen compounds. Compare hydrofluoric and hydrochloric acids as to 1. Method of preparation. 2. Characteristics. 3. Uses. Special points for study. Method of etching glass. What kind of substances may be used instead of manganese dioxide in preparing chlorine? Why? Proof of this by experiments. Practical application of this. Compare these two methods of making chlorine : 1. MnO 2 + HCL 2. MuO 2 + XaCl + H 2 SO 4 . How similar ? How different ? Explain how chlorine bleaches. Write the equation. Source of commercial supply of hydrochloric acid. Tests for bromine and iodine with carbon disulphide compare results. Method of obtaining and purifying iodine. Describe experiments which illustrate chief properties of chlorine, bromine, and iodine. Solvents for chlorine, bromine, and iodine. CHAPTER X ACIDS, ALKALIES, AND SALTS 1. Neutralization. There are a great many substances which, if put together, have the power of destroying the characteristic properties of each other. EXPERIMENT 84. To show this fact, put into an evaporating dish about 10 cc. of dilute hydrochloric acid and dip into it a small piece of blue litmus paper. Notice that it is changed to red. Now add slowly, stirring all the time with a glass rod, a solution of caustic soda, until the litmus paper just turns blue again; then add one drop of hydrochloric acid. You ought now to have a solution that will not affect either red or blue litmus paper. Boil this solution to dryness. What is the appearance of the solid thus obtained? Taste it. Does it seem familiar ? Dissolve it in a little water and test with both red and blue litmus paper. Is the paper affected? Now boil a little hydrochloric acid to dryness. Does it leave a residue? Examine a specimen of solid caustic soda and compare with the white solid obtained above. Are the two solids the same? Do they both affect litmus in the same way ? 2. From this experiment we see that the acid and the caustic soda, on being put together, have both lost their characteristic properties and have reacted to form a new substance having the properties of neither. In other words, they have neutralized each other. 3. Bases. Such substances as have the power of neu- tralizing the properties of acids are called bases. This was shown in Experiment 84 above. We have already seen that the compound of any element with oxygen is called an oxide; many of the oxides combine with water to form 124 ACIDS, ALKALIES, AND SALTS 125 zvater oxides, or, to use the ordinary term, which is from the Greek, hydroxides or hydrates. We have seen also that some oxides or anhydrides, when united with water, form acids, as, for example, nitrogen trioxide. Strictly speaking, such compounds are hydroxides, but we never apply that term to them ; it is restricted to the compounds of metallic oxides with water. In brief, bases are metallic hydroxides. ~Ty%" Alkalies. Bases, soluble in water, which have ex- - ceedingly strong basic properties, are called alkalies. The four most common alkalies are the three hydroxides of sodium, potassium and calcium, and ammonia. If we study the 'formulae of the hydroxides, we shall see that water may be taken as the type upon which all the others are built. Thus : Water HOH Caustic Potash . . , . . . KOH Caustic Soda NaOH Lime Water -. Ca(OH) 2 Ammonium Hydroxide . . . NH 4 OH The only difference is that one atom of hydrogen in the water has been replaced by some metal or group of elements. 5. As most of the metals themselves have certain characteristic properties of bases, they are, by some, spoken of as bases. Possibly there is no serious objec- tion to this, but it should be remembered that all bases are compounds ; thus, sodium may have many of the properties of a base, but it is not a base any more than bromine is an acid. . Acids. It is a difficult matter to define acids. They are substances which have certain properties the opposite of bases. They possess the power not only of turning 126 MODERN CHEMISTRY blue litmus red, but of similarly affecting various other vegetable colors, all of which are restored again by the use of an alkali. They also have a sour taste, though this is not a distinctive feature, as many bodies not acids also have the same property. 7. Their Composition. If we recall the formulae of the few acids we have already used, nitric, hydrochloric, and sulphuric, we see that they all contain hydrogen ; it is true also that most contain oxygen, together with some third element which seems to give the distinctive proper- ties to the acid. It was at one time supposed that all acids contained oxygen, and in accordance with this idea oxygen received its name. Later, however, were discov- ered hydrochloric and other acids, which contained no oxygen whatever. A distinctive property of acids is that they all have the power of giving up the whole or a part of their hydrogen, and of combining instead with some metal or base. This we have seen. They are compounds of water with non-metallic oxides, and sometimes their formulae are written as if they were hydroxides; thus, H 2 S0 4 , S0 2 (OH) 2 ; HN0 3 , NO 2 (OH), etc. 8. Salts. A salt is a compound formed by the union of an acid with a base or metal, possessing properties different from those of either of its constituents. We have been accustomed to think of salt as a particular sub- stance used in seasoning food, but we must now remember that it is a term applied to a large number of compounds, called salts because they resemble common salt in appear- ance or properties. They are all produced in the same way. We saw above that when we neutralized hydro- chloric acid with caustic soda and boiled to dry ness, we obtained a white solid, resembling and tasting like com- mon salt, which it really was. ACIDS, ALKALIES, AND SALTS 127 EXPERIMENT 85. In the same way as in Experiment 84, neu- tralize about 10 cc. of hydrochloric acid with caustic potash and boil to dryness as before. Compare the salt produced, in taste and appear- ance, with that obtained before. 9. Normal or Neutral Salts. There are two general classes of salts, neutral or normal, and acid. A normal salt is one in which all the displaceable hydrogen of the acid used in making the salt has been replaced by some base. For example, when caustic potash and sulphuric acid neutralize one another, the following reaction takes place : H 2 SO 4 + 2 KOH = K 2 SO 4 + 2 H 2 O. We see here that all the hydrogen in the sulphuric acid, two atoms, has been replaced by an equivalent amount of the metal, potassium, and the salt produced, potassium sulphate, K 2 SO 4 , is a normal salt. 10. Again, if lead is treated with vinegar (acetic acid), which is represented by the formula HC 2 H 3 O 2 , we have this reaction : Pb + 2 HC 2 H 3 2 = Pb(C 2 H 3 2 ) 2 + H 2 . It will be noticed that in the salt resulting, Pb(C 2 H 3 O 2 ) 2 , a quantity of hydrogen remains. Lead acetate is, notwith- standing, a neutral salt, because only one atom of hydro- gen, the first, in each molecule of acid can be displaced. 11. Acid Salts. If, however, we use only half the amount of caustic potash shown by the first reaction above with the sulphuric acid, we shall replace only half of the hydrogen in the acid, and the salt resulting will be an acid salt, thus : H 2 SO 4 + KOH = KHSO 4 + H 2 O. 128 MODERN CHEMISTRY 12. Reading the Formulae of Salts. The compound, K 2 SO 4 , is read, normal potassium sulphate, or usually, simply potassium sulphate. The acid salt, KHSO 4 , is read, acid potassium sulphate, or potassium hydrogen sulphate. Sometimes the prefix mono is applied, but usually only in the case of salts of acids having three or more replacealle hydrogen atoms, as phosphoric, H 3 PO 4 , or silicic, H 4 SiO 4 With these acids we may form the following salts : Phosphoric Acid, H 3 PO 4 Mono-sodium Phosphate NaH 2 PO 4 Di-sodium Phosphate Na 2 HPO 4 Normal sodium Phosphate Na 3 PO 4 Silicic Acid, H 4 SiO 4 Mono-sodium Silicate NaH 3 SiO 4 Di-sodium Silicate .... ... Na 2 H 2 SiO 4 Tri-sodium Silicate Na 3 HSiO 4 Normal sodium Silicate Na 4 SiO 4 EXERCISES. In the following formulae, state which represent acids, which bases, and which salts, giving reasons therefor. Give also the name of the substance 'represented. If salt, state whether acid or normal : Na 2 S0 4 , KOH, H 2 SO 4 , P(OH) 3 , ZnSO 4 , KNO 3 , Ca(OH) 2 , BaSO 4 , K 3 P0 4 , K 2 HP0 4 , KH 2 P0 4 , HC1, NaOH, NaHSO 4 , NaNO 3 . 13. Nomenclature of Acids. It will be noticed that with one exception the acids we have met with so far all have names ending in ic ; thus : Sulphuric .... H 2 SO 4 Nitric HNO 3 Phosphoric .... H 3 PO 4 Silicic ..... H 4 SiO 4 , etc. Sulphur : acid, H 2 SO 4 . Sulphuric Nitrogen : HNO 3 . . Nitric Phosphorus : H 3 P0 4 . . Phosphoric Silicon : H 4 SiO 4 . Silicic ACIDS, ALKALIES, AND SALTS 129 The greater number of acids with which we shall have to deal, as already stated, contain three elements, the first of which is hydrogen, the third oxygen, and a second which gives the name to the acid. Thus the middle sym- bols in the above formulas are : S . N . P . Si . 14. Sometimes, however, this second element forms more than one acid with hydrogen and oxygen. In such cases the most common, and hence the earliest known, received the name with the termination ic. Then the acid having a smaller amount of oxygen is given, the termination ous. This we have seen in the two nitro- gen acids : Nitric . . . HNO 3 . . Oxygen, 3 atoms. Nitrous . . HNO 2 . . 2 " 15. Sometimes even three or four acids are formed from the same three elements, the amount of oxygen only vary- ing. In such cases, the one with the least quantity of oxygen is given the prefix hypo, meaning under or lesser, and the one with the most oxygen has the prefix per, beyond or above. These may be illustrated by the follow- ing series : Sulphured . * r 5 , H 2 Sulphurous. . . . H 2 . H O 4 . . Oxygen, 4 atoms O 3 . . " 3 " O "2 " 130 MODERN CHEMISTRY CHLORINE ACIDS Perchloric . H Chloric . . . H Chlorous H Cl Cl 01 O 4 . . Oxygen, 4 atoms. O 3 . . " 3 " Oo " 2 Hypo-chlorous . . H Cl O . . " 1 " 16. Nomenclature of Salts. All of the acids named above have the power of combining with various metals, or their hydroxides, to form salts. All such as result from the union of a base with an ic acid are given names with the "termination ate. Thus, all salts of sulphuric acid are sulphates; of nitric acid, nitrates; phosphoric acid, phosphates, etc. To illustrate : H 2 SO 4 , sulphuric acid, gives with zinc, ZnSO 4 , zinc sulphate ; with potassium, K 2 SO 4 , potassium sulphate ; . with calcium, CaSO 4 , calcium sulphate. HNO 3 , nitric acid, gives with potassium, KNO 3 , potassium nitrate ; with sodium, NaNO 3 , sodium nitrate ; with ammonia, NH 4 NO 3 , ammonium nitrate. 17. Salts formed from the ous acids receive names end- ing in ite. (It may aid the memory in associating the pronouns singular, J, plural objective us, with the ous acids and ite salts.) Thus, from H 2 SO 3 , sulphurous acid, we have K 2 SO 3 , potassium sulphite ; Na 2 SO 3 , sodium sulphite, etc. ACIDS, ALKALIES, AND SALTS 131 In the case of salts formed from the hypo and per acids, the corresponding prefix is simply given to the salt. Thus : - NaCIO, sodium hypochlorite, from HC1O, acid, %p0chlorous, and NaClO 4 , sodium perddorate, from HC1O 4 , acid, perchloric. 18. Binary Compounds. All of the above salts are formed from what are sometimes called ternary acids; that is, those consisting of three (or more) terms. In like manner, a binary compound would be one which consists of only two elements. The following are ex- amples : Common Salt .... NaCl Calcium Chloride . . . Water ....... H 2 O Turpentine It will be noticed from these formulae that though in a binary compound there are but two elements, the number of atoms of each of these elements is quite variable. 19. As already stated, there are a few acids which contain no oxygen. Salts obtained from them would, therefore, all be binaries. Thus : from Hydrochloric acid, HC1, we obtain the chlorides ; Hydrobromic " HBr, " bromides; Hydriodic " HI, " iodides; Hydrofluoric " HF, " fluorides; Hydrosulphuric " H 2 S, * ; sulphides. 132 MODERN CHEMISTRY 20. How Binary Compounds are Named. It will be noticed that all binary salts are given names ending in ide. Furthermore, it is seen that it is the negative ele- ment in every case which gives the name to the substance ; thus : NaCl ^| KOI -,,,-, \- are all chlorides, while the positive MnCl 2 I CaCl 2 J element indicated in the formula is simply descriptive in character, or the adjective that tells what kind of a chloride. Thus the above are Sodium Potassium , . . , TV* \ Chloride ; Manganese Calcium just as we might say Stone Brick [ House. Frame Log 21. It frequently happens that two elements, just as in the case of the ternary acids, may unite in different pro- portions to form two or more compounds. We have already seen this in studying the oxides of nitrogen, p. 81. When two such exist, as for example the two oxides of mercury, HgO and Hg 2 O, the one having the smaller proportion of the negative element, as indicated by the formula, is the ous compound, just as in the case of the acids already studied, while the one having the greater proportion of the same element is the ic compound. ACIDS, ALKALIES, AND SALTS 133 22. Again, we noticed in studying the oxides of n> trogen : N 2 O, ratio of oxygen to nitrogen, 1 : 2 N 2 2 , " " 2:2 = 1:1. Jn the first we found one-half as many atoms of oxygen as of nitrogen ; in the second the same number ; they were therefore called nitrons and nitric oxides. In a few instances, instead of using the English name with the terminations ous and ic, for the sake of euphony, the Latin forms are taken. Thus : Cu 2 O, Cuprous Oxide CuO, Cupric " FeCl 2 , Ferrous Chloride Fe 2 Cl 6 , Ferric 23. Returning to the series of nitrogen compounds, it will be noticed that they were given two names. This is very often done, one of them using a prefix to indicate the exact number of atoms of the last element in the formula of the compound. Thus we have N 2 O, Nitrogen Monoxide. N 2 O 4 , " Tetroxide P 2 O 5 , Phosphorus Pentoxide, etc. 24. Old Forms. Occasionally we use the old terms, pro, per, and sesqui. The first is a prefix, meaning before, and is given to some uncommon or unstable compounds which in the case of a series would be the first or lowest. Thus, FeO is sometimes spoken of as iron protoxide. Nitrogen tetroxide, N 2 O 4 , is also called peroxide, as is 134 MODERN CHEMISTRY hydrogen dioxide as well, it being the compound com- ing in the series beyond the others. Sesqui is applied to binary compounds in which the two elements unite in the ratio of 2 to 3, as in Fe 2 O 3 , iron sesquioxide. SUMMARY OF CHAPTER Neutralization Meaning of the term. Experiments to illustrate. Three classes of compounds. Compare bases and acids. a. In composition. b. In properties. Alkalies What are they ? Examples. Salts What are they ? Two classes. How formed. How distinguished by name. Examples to illustrate. Nomenclature. a. Of acids. b. Of salts. Examples to illustrate both. Binary compounds. Meaning of term Illustrations. Six important classes Examples. Nomenclature Compare with acids. CHAPTER XI CARBON AND A FEW COMPOUNDS. C = 12 1. Abundance. With the exception of oxygen, carbon is the most widely distributed element, and is also very abundant. In the form of compounds it is found in the air as carbon dioxide, resulting from combustion and respiration, and in limestone, CaCO 3 , which constitutes a large portion of the rocky crust of the earth. It also occurs in almost all food products, such as sugar, flour, starch, vegetables, and fruits, and forms a large part of the woody structure of plants and trees. 2. Forms. In the free state carbon may be considered under two divisions : a. Crystallized, including 1. The Diamond. 2. Graphite or Plumbago. b. Amorphous (without crystalline form), 1. Coal. 2. Lampblack. 3. Gas Carbon, etc. 3. Diamonds. The diamond occurs in octahedral crys- tals. It is found in South America, Africa, Australia, and India. By some the stones are thought to be of meteoric origin and not native to the earth, but the theory seems not well founded. Moissan, the French chemist, has suc- ceeded in making a few diamonds in the electrical furnace, 136 MODERN CHEMISTRY but they have all been exceedingly small, and black in color, so as to have no value except in a scientific way. In nature they occur rough and covered with a layer of partially decomposed rock. The most highly prized are perfectly transparent, but many of various colors have been found. The diamond has strong refractive power, is the hardest of all minerals, and can be cut and polished only by its own dnst. 4. Their Practical Uses. Diamonds are used, not only as ornaments, but also in cutting glass ; and the cheaper, imperfect varieties are employed as tips on drills for cut- ting through hard rocks. That the diamond consists of carbon may be proved by burning it between electric ter- minals in an atmosphere of oxygen ; the diamond and oxygen disappear, and carbon dioxide, CO 2 , remains. 5. Graphite. Next to the diamond, graphite or plum- bago is the purest form of carbon. It is sometimes called black lead, but it contains no lead whatever. It is often found in hexagonal prisms, is steel-gray in color, has a greasy feeling, and as a mass is comparatively soft, though the particles themselves are very hard. 6. That it consists of carbon may be proved by testing it in the electric furnace, as in the case of the diamond, similar results being obtained. 7. Uses. The most common use of graphite is in making what are known as lead pencils, so named because plumbago was at first supposed to be a compound of lead. In making pencils the graphite is thoroughly pulverized and mixed, according to the grade of pencil, with different proportions of fine clay, also well ground. The whole is then made up with water into a dough and pressed into moulds and dried, or while still soft is forced through plates with apertures the size of the lead in the pencil. CARBON AND A FEW COMPOUNDS 137 Graphite is also used as a lubricant, as a stove polish, and in making crucibles. 8. Amorphous Carbon. The most important uncrys- tallized forms of carbon are the various coals, anthra- cite, semi-anthracite, bituminous, lignite, peat, jet, cannel, and the artificial form, charcoal. Of great importance also are gas carbon, lampblack, and coke. 9. Coal. Coal is supposed to be the result of pressure and heat applied to a luxuriant vegetable growth in the presence of moisture. Peat is the newest of the coals, being in process of formation in swamp-lands to-day. It consists almost entirely of a mass of roots. Next in age is lignite, in which the woody structure is still apparent. 10. Anthracite and Bituminous Coals. Anthracite dif- fers from bituminous coal in that the former, being sub- jected to greater heat and pressure, has been deprived of its volatile products. These furnish in part, at least, the petroleum and natural gas of the present time. Petroleum is really a mixture of a number of different oils, with boil- ing points differing greatly. These, in the process of re- fining the crude oil, distill over at different temperatures. Such light oils as naphtha and benzine are obtained at a low temperature, a somewhat higher temperature produc- ing kerosene, and higher still parafnne. This method of separating substances through differences in their boiling points is called fractional distillation, while that "in which the substance heated is decomposed is called destructive distillation. 11. Charcoal. Charcoal, because of the abundance of timber, has usually been prepared in a simple, but very wasteful, manner. Large piles of wood are covered with earth and set on fire. Most of the air is excluded in this way, and only enough heat is produced to expel the vola- 138 MODERN CHEMISTRY tile products from the wood. At present, however, in some sections the wood is heated in iron retorts, and the volatile products are condensed and refined, much in the same way as with petroleum. 12. Coke. Coke bears the same relation to soft or bituminous coal that charcoal does to wood. It is an artificial product obtained by expelling all the volatile products from the coal. Part of the supply comes from the gas factories as a by-product, but where the local supply is insufficient, it is prepared specially for smelters in large brick ovens. See the figure below. FIG. 37. Coke Oven. aa, openings for slight draught at first. DD, doors for removing coke. The coal, in car loads, is shoveled in from above ; it is then ignited, and the openings on the side almost entirely closed. In the course of several hours the combustion of the lower layer of coal has converted the remainder into coke, the doors are opened, and the coke drawn out. 13. Gas Carbon. Gas carbon is another by-product of coal-gas manufacture. Just as soot collects in stove-pipes and flues, so on the inside of the retorts there is gradually deposited a very hard, black substance, known as gas carbon. This is occasionally removed, ground up fine, CARBON AND A FEW COMPOUNDS 139 and moulded into the familiar carbon rods in our electric arc lights, and into plates for electric batteries. 14. Lampblack. Lampblack is the result of the imper- fect combustion of any substance rich in carbon. It is usually prepared by burning some hydrocarbon, such as turpentine, C 10 H 16 , in a limited supply of air. The dense black smoke resulting is allowed to deposit upon canvas in a cool room, from which it is shaken, and is then ready for commerce. It is used in making black paint, printers' ink, etc. SOME USES OF CARBON 15. As a Reducing Agent. In the form of charcoal or coke, at a high temperature, carbon is a great reducing or deoxidizing agent. By this we mean that when it is heated with the oxides of various metals, it has the power of combining with the oxygen and reducing the oxide to the metallic condition. This will be made clear by the fol- lowing experiment. EXPERIMENT 86. Make a small cavity near one end of a stick of charcoal, and put therein a little litharge, PbO, or red lead, Pb 3 O 4 , and heat strongly with the reducing flame. Notice that in a few minutes a bead of lead appears instead of the oxide that we had. The carbon has combined with the oxygen in the lead oxide to form carbon djioxide, and the lead has been reduced to the metallic form. 16. As an Absorbent. Carbon in the form of charcoal is an excellent absorbent, not only of gases, but of certain other substances as well. EXPERIMENT 87. Thrust a piece of charcoal under water and hold it there a minute or so. What is seen escaping from the char- coal? Heat another piece red-hot and plunge under water. Are the results different? Why?* * In this connection refer to Exp. 47, under ammonia. 140 MODERN CHEMISTRY EXPERIMENT 88. Soak some vegetable matter in a vessel of water until it has become very offensive, on account of decomposi- tion. Put a little of this water into a flask and add some bone-black or powdered charcoal, and shake well. Notice that the disagreeable odor disappears. 17. As a Purifier. Application is made of this fact in purifying cisterns which have become foul with decom- posing organic matter. The charcoal should be removed after a time and heated to redness to destroy thoroughly the organic matter which may have been absorbed. It is believed that partial oxidation takes place within the pores, but unless the charcoal is heated they eventually become clogged. EXPERIMENT 89. Fit a filter paper smoothly to a funnel as described in Appendix C, page 365, and partly fill it with bone-black. Now pour upon it, slowly at first, a few cubic centimeters of logwood or some other colored vegetable solution. How is it affected? Try also in the same way a solution of copper nitrate. Are the results the same? 18. .In Refining. An application of the power of char- coal to absorb vegetable colors is made in refining sugars. At first they are brown, not very different from maple sugar in appearance. This raw sugar, as it is called, is dissolved in water an/i passed through filters of bone- black which absorb the coloring matter and leave the solution clear. This may be shown by filtering a solu- tion of molasses in water. EXPERIMENT 90. In like manner, charcoal has the power of ab- sorbing various organic flavors. Pass through a powdered char- coal filter an infusion of tea or coffee, and taste it after it has gone through. How is it changed? 19. It would be impossible to enumerate the various uses of carbon in its different forms. Many of these are CARBON AND A FEW COMPOUNDS 141 familiar to the student, and others will be learned from time to time. Many of them have already been named in the sections immediately preceding this. COMPOUNDS OF CARBOX. THE OXIDES 20. Carbon Monoxide, CO. This is a gas obtained when carbon is burned in a limited supply of air. It may be prepared by passing steam over red-hot coke or charcoal, whereby the steam is decomposed, thus: H 2 + C = CO + H 2 . It is also produced in grates and base-burners. At the lower portions of the fire where the heat is most intense the carbon is completely burned, producing carbon dioxide; as this passes up through the red-hot coal, it unites with another portion of carbon and forms the monoxide. Again, on reaching the upper surface, the monoxide unites with the oxygen of the air and is burned into carbon dioxide. 21. Carbon monoxide may be prepared in an impure form by heating oxalic acid or potassium ferrocyanide with sulphuric acid, or by passing a current of carbon dioxide slowly through a tube containing red-hot charcoal or coke. 22. Characteristics of Carbon Monoxide. Carbon mo- noxide is a colorless gas, having a faint, peculiar, but some- what unpleasant and stifling odor; it is a little lighter than air and burns with a pale blue flame. It is not soluble in water, is only slightly explosive when mixed with air or oxygen, and is poisonous when inhaled. It has the power of decomposing the blood, and thus of ren- dering it incapable of carrying oxygen and removing the waste of the body. On this account serious results some- times follow its escape into rooms from coal stoves when 142 MODERN CHEMISTRY the drafts are closed at night. Open charcoal fires also produce the same gas, and have sometimes been the means of causing death. 23. As a Reducing Agent. It has been seen that carbon is a strong reducing agent. Carbon monoxide has the same properties, owing to the fact that it has strong affinity for more oxygen, to form carbon dioxide. The reduction of metallic ores in blast furnaces is, to a con- siderable extent, due to this property of carbon monoxide. It may be seen by passing a current of carbon monoxide over lead oxide, PbO, heated red hot in a tube. The monoxide abstracts the oxygen from the lead oxide, form- ing carbon dioxide and metallic lead. The reaction is as follows: - pbQ + CQ = pb 24. Carbon Dioxide, C0 2 . Where found. This gas is always found in the air, being produced by the combustion of organic bodies and by respiration. The proportion varies somewhat, but seldom exceeds four parts in 10,000 parts of air. Another source of this gas as found in the atmosphere is fermentation and decay. 25. Produced in Decomposition. As already mentioned in considering ammonia, organic substances are very un- stable and readily break up to form simpler compounds. The molecules of most so-called organic compounds consist of carbon, hydrogen, oxygen, and often nitrogen, and are usually very complicated. In the processes of decay the atoms rearrange themselves, and carbon dioxide is one of the new products. The process is the same when fer- mentation is induced by bacteria or germs, such as those of ordinary yeast. If into a flask containing some water sweetened with sugar or molasses a little yeast be intro- duced, fermentation very soon begins, and the bubbles of CARBON AND A FEW COMPOUNDS 143 ga3 which pass off may be collected and proved to be carbon dioxide. 26. How prepared. For laboratory purposes carbon dioxide is usually prepared by treating some carbonate, as marble, CaCO 3 , with dilute acid. EXPERIMENT 91. Put into a small flask some marble, coarsely powdered, and add some dilute hydrochloric or nitric acid. Notice the rapid effervescence and evolution of colorless gas. EXPERIMENT 92. Collect by downward displacement a bottle of the gas, generated as above. Lower into it a burning match or candle ; what are the results ? Ignite a piece of magnesium ribbon and hold it in a bottle of carbon dioxide; what are the results? What two products are formed? Why does the ribbon continue to burn? Mention some other gas that supports the combustion of phosphorus, but not that of ordinary substances. EXPERIMENT 93. To show the density of the gas. Put into a good- sized bottle or beaker a small candle and pour in upon it another bottle of carbon dioxide. You cannot see anything being turned out, but the results are apparent. This is sometimes made more effective by fastening at short intervals upon the bottom of a trough several candles. Lift one end of the trough and pour down it a large bottle of carbon dioxide. The candles will be extinguished, one after another, as the gas reaches them. EXPERIMENT 94. Purpose same as preceding. Put into an evaporating dish a little gasoline and ignite it. Take a large bottle of carbon dioxide and pour suddenly upon the burning oil. The flame will be instantly extinguished. EXPERIMENT 95. To show effect of carbon dioxide on limestone. Pass a current of carbon dioxide through a few cubic centimeters of lime water. Notice the formation of a white precipitate, which is calcium carbonate, CaCO 3 , of the same composition as limestone. Continue passing the gas through the milky solution ; what change takes place ? Can you explain ? 144 MODERN CHEMISTRY 27. Characteristics of Carbon Dioxide. From the above experiments we learn that carbon dioxide is a colorless^ odorless gas, considerably heavier than air. It is non- combustible and a non-supporter of ordinary combustion^ though 'such substances as magnesium, which burns with great intensity, are able to decompose the gas and make use of the oxygen. It is slightly soluble in water and gives to the latter a faint acid taste and reaction. The presence of carbon dioxide may always be determined by its effect upon lime-water. It forms in the water a white precipitate which dissolves slowly again in excess of the dioxide. Limestone caves are a manifestation on a large scale of the principle shown in the simple experiment above. Water under pressure absorbs considerable quan- tities of carbon dioxide, which gradually dissolves the limestone and forms caverns. 28. Liquid Carbon Dioxide. Carbon dioxide may be liquefied in strong cylinders by pressure ; if the pressure is suddenly withdrawn, a portion of the liquid is rapidly vaporized, producing such cold as to convert the remainder into a white crystalline solid like snow. The temperature of this carbonic acid snow is sufficiently low to freeze mercury. The solid carbon dioxide vaporizes without first melting. 29. Choke Damp. Because of its density, carbon di- oxide frequently collects in deserted mines and deep wells, and is called by miners "choke damp." Its presence in such places, however, may always be detected by lowering to the bottom a burning candle or lantern. Carbonic anhydride is another name for the same gas, it being the anhydride of the unstable acid, H 2 CO 3 . It is still popu- larly called carbonic acid gas. 30. Uses of Carbon Dioxide. Carbon dioxide is used extensively in making "soda water." It is confined in CARSON AND A FEW COMPOUNDS 145 strong cylinders under great pressure, and allowed to flow into cold water in strong tanks also under pressure. The water is thus thoroughly charged. When the stop- cock is turned and the water flows into the glass, the pressure being removed, the carbon dioxide rapidly bubbles out. It is this gas which gives the sharp biting taste to soda water and also to the water of many mineral springs. It is the same gas that causes the effervescence in beer and the sparkling appearance of some wines. 31. In some of our cities carbon dioxide is now being put upon the market in small oval-shaped steel vessels into which the gas is forced under great pressure. When ready for use, a valve is opened, the gas rushes into a glass of water flavored and sweetened, and the soda water is ready. These sparklets, as the steel vessels are called, are very small, and a large number may be carried without great inconvenience. In Germany the same article is sold under the name of Sodors. Certain fire extinguishers owe their value to the large quantities of carbon dioxide contained ; and instances are on record in which fires in coal mines which have defied all other means have been extinguished by passing in carbon dioxide. 32. Though this gas cannot be inhaled in any consider- able quantities, it is not poisonous, but like water causes death by shutting out the oxygen. Hence a person might drown in a well or vat of carbon dioxide just as readily as in one of water. To plant life, however, the gas is in- dispensable ; it is inhaled by plants as oxygen is by animals, and in the presence of light the life forces of the plant are sufficient to decompose the compound into its constituents. The carbon is stored up in the woody structure of the tree or plant, and the oxygen is given off again to the air. Thus a considerable portion of the 146 MODERN CHEMISTRY carbon in all our forests and coal beds was once in the form of gaseous carbon dioxide in the atmosphere. THE HYDROCARBONS 33. Definition. By this term we mean those com- pounds consisting of carbon and hydrogen, of which there are many. The most important are the three follow- ing: Marsh Gas .... CH 4 Olefiant Gas .... C 2 H 4 Acetylene . . . . C 2 H 2 34. Marsh Gas. This is also known as methane, and by miners asfire damp. It is often found in coal mines as the result of the decomposition of organic matter, and in swamps from the same source. By stirring the leaves and similar matter that collect upon the bottoms of ponds, bubbles of gas, consisting largely of marsh gas, are seen to rise. It is always produced in the destructive dis- tillation of any organic matter, such as the preparation of charcoal in retorts or that of illuminating gas from coal. 35. Characteristics of Marsh Gas. Marsh gas is a color- less, odorless gas, the lightest of all except hydrogen, having a specific gravity compared with air of less than 0.6. It is highly inflammable, burning with a pale blue flame, and with air or oxygen forms a dangerous explosive mixture. It is by this gas that most explosions in coal mines are caused, and on this account it is called fire damp, the word damp with miners being a generic term meaning gas. Marsh gas is somewhat soluble in water, and is neutral to test paper, affecting neither red nor blue. It is an important constituent of ordinary coal gas, and when burned produces much heat. CARBON AND A FEW COMPOUNDS 147 36. Protection against Fire Damp. If you hold a wire screen over the flame of a Bunsen burner, you will see that the flame does not pass through it, although if you bring a lighted match above the screen, you will find there a combustible gas. This is because the wire cloth, being a good conductor of heat, withdraws it from the burning gas and so lowers the temperature that what has passed through no longer burns. Now hold the screen in the flame until it becomes red hot ; the gas above will be ignited and continue to burn. An observation of these facts led Sir Humphry Davy to design the " safety lamp " which now bears his name. It is little more than an ordinary miner's lamp sur- rounded by a wire screen. If the miner enters a chamber filled with fire damp, though the gas may burn on the inside of the screen, there is no danger unless he remains until the wire becomes hot enough to ignite the gas outside. 37. Olefiant Gas, C 2 H 4 . This also is a constituent of common illuminating gas, and is formed in the destructive distillation of wood and coal. It may be prepared by heating ethyl alcohol with sulphuric acid. Really, two reactions take place ; first an ethyl ester of sulphuric acid is formed, C 2 H 5 HSO 4 , which very soon breaks down into ethylene and sulphuric acid. The final result is C 2 H 5 OH + H 2 S0 4 = C 2 H 4 + H 2 S0 4 , H 2 O. 38. Characteristics of Olefiant Gas. Ethylene, as this gas is also called, is of about the same density as air, is colorless, has a faint odor, and burns with a yellowish white light, such as is seen in the ordinary gas jet. It is somewhat explosive when mixed with air or oxygen ; at 40 atmospheres' pressure it is reduced to a liquid. 148 MODERN CHEMISTRY ACETYLENE, C 2 H 2 39. How prepared. This gas is formed in small quao* tities together with other hydrocarbons in the distillation of wood and coal. It is prepared now in large quantities by treating calcium carbide, CaC 2 , with water, as follows : CaC 2 + H 2 = CaO + C 2 H 2 . ' The lime, CaO, thus formed immediately reacts with an- other molecule of water, forming slaked lime, or calcium hydroxide, Ca(OH) 2 , thus : CaO + H 2 O = Ca(OH) 2 . The final reaction then would be indicated by CaC 2 + 2 H 2 O = C 2 H 2 + Ca(OH) 2 . 40. Calcium Carbide. Calcium carbide is a dark gray solid, more or less crystalline in appearance, always giving off the odor of acetylene, owing to its decomposition by the moisture in the air. In America the greater portion of the commercial supply comes from Niagara Falls, where it is prepared by fusing at intense heat in electrical fur- naces pure lime intimately mixed with charcoal finely pulverized. When taken from the furnace, it is packed in metallic drums, sealed air-tight, and is then ready for shipment. EXPERIMENT 96, Into a test-tube put a small lump of calcium carbide, cover with water, and quickly insert a cork with delivery tube and jet attached. Notice the violent chemical action and the odor of the gas. Light the jet and notice with what kind of a flame it burns. 41. Another Method. Sometimes this method is varied slightly by using a flask fitted with a two-hole cork. Through one hole passes the delivery tube, through the CARBON AND A FEW COMPOUNDS 149 other a funnel with a stop-cock. In this way the flow of water can be regulated and the rapid evolution of gas prevented. Precaution must be taken in this case not to light the jet too soon, as acetylene mixed with air is dan- gerously explosive. 42. Acetylene Generators. This illustrates one class of acetylene generators now offered upon the market, in which the water is allowed to drip on the carbide. The objection to this is that with the small supply of water the carbide becomes so warm as to bring about a partial decomposition of the acetylene into other undesirable hydrocarbons. EXPERIMENT 97. For class-room work an excellent generator may be prepared thus : procure a tin can, holding a quart or two, and having a screw top. (A can in which some varieties of coffee are sold will do.) To the inside of the screw top solder a hook. Upon this suspend a small bucket or basket made from a tin can, and having a wire-cloth or perforated bottom. Cut out the bottom of the larger can as shown in the cross-sectional view 6 of the accom- panying figure ; then solder two strong bent wires, W, W, upon opposite sides of it. Xow obtain another can just large enough to allow the first to move up and down easily within it. Melt or cut out the top ; then cut down two flaps about three-fourths of an inch deep, and bend them to a horizontal position, as at F. Through each flap punch a hole large enough to receive the bent guide-wires soldered on to the other can. Near the bottom cut a round hole and insert a rubber cork, through which passes a bent delivery tube extending up nearly to the top of the can. When ready for work fill this can nearly full of water, put some carbide into the basket, sus- pend it upon the hook and then lower the first FIG. 39. 150 MODERN CHEMI8TEY can into position, the guide-wires passing through the openings in the flaps. The screw top enables one to refill the basket without remov- ing the entire cylinder. As soon as the carbide touches the water, acetylene will begin to form, and, mixed with air, will flow from the delivery tube T. This generator, which illustrates another class now upon the mar- ket, is automatic. In case the delivery tube becomes clogged, the increasing pressure of the gas will lift the inner cylinder, and with it the basket of carbide, from the water. Or, if the guide-wires become caught and prevent this, the pressure on the water will cause it to flow out over the flaps. In either case the rapid evolution of gas will soon cease. CAUTION. Before beginning the generation of acetylene be sure no lights are in close proximity, and allow the first gas generated to escape. It contains too much air for good results and is too dangerous. With these precau- tions the gas may be used direct from the generator, or first passed into an ordinary gasometer, which any tinner can make cheaply. 43. Acetylene Burners. But to secure steadi- ness of flow and safety, it is always better to pass the gas through an acetylene burner or tip, which differs from the tip of an ordinary gas jet only in that instead of a slit there are two very small openings drilled, oblique to each other. See Fig. 40 ; a is a cross-sectional and b the top view. These tips are. very cheap, and safe because the openings for the exit of FIG. 40. gas are so small that the flame cannot pass back into the generator ; c shows another form of tip frequently used. The two openings compel the issuing jets of gas to strike each other obliquely, as in a. 44. Characteristics of Acetylene. Acetylene is a color- less gas, of an ethereal odor when perfectly pure, but as CARBON AND A FEW COMPOUNDS 151 ordinarly obtained it is very offensive to the smell. It is soluble, volume for volume, in water and very explo- sive when mixed with oxygen or air. An ordinary jet of acetylene burns with a yellowish flame, and owing to the large proportion of carbon, over 92 per cent, it gives off considerable soot. With a burner like the one described above it furnishes an intensely white light, rival- ing the calcium or Drummond light in brilliancy ; so that it is now frequently used for projecting lantern slides upon screens and for bicycle lamps. 45. Intense Heat. Fine iron wires held in the flame are quickly consumed, throwing off sparks as if burning in oxygen. When used in a blast lamp instead of common gas, acety- lene burns with a bluish-white flame. The intensity of this is sufficient to melt copper wires readily, and ordinary platinum wires in two or three minutes ; furthermore, it will even soften porcelain. Iron wires a sixteenth of an inch in diameter are quickly fused and burn with a most brilliant shower of sparks, especially when a molten globule of iron upon the end of the wire is suddenly oxi- dized, and being thrown out into the air breaks into a shower of stars. Watch-springs and knife-blades may be as easily burned away. In a darkened room the display is very beautiful. 46. Blowpipe for Experiments. The blowpipe best suited to this work may be made by almost any student. See Fig. 41. The outer part, B, is the ordinary black japanned blowpipe, costing only a few cents. A hole is cut through at the point E, for the insertion of an inner 152 MODEEN CUEMISTRY tube, which may be made by carefully straightening an ordinary eight-inch brass blowpipe. Solder this firmly in place, plug the mouth end of the outer pipe with a piece of brass through which a small hole has been drilled, and the acetylene blowpipe is complete. Connect the inner tube with the foot-bellows furnishing the air, and the outer tube with the acetylene tank, through the acetylene tip. Regulate by means of a stop-cock the flow of gas, so that when in operation the acetylene is completely burned, with the flame almost entirely blue. 47. From these experiments it will be seen that the heat of this flame is intense, reaching probably 2000 C. EXPERIMENT 98. To show the explosiveness of acetylene. In the center of the bottom of a pound baking-powder can punch a small hole. Place the can, bottom upward, for a minute or so over a tube delivering acetylene, then set upon the table in the same position. Bring a flame to the touch-hole, when, if the proportions are suitable, a violent explosion will ensue, and the can will be thrown several feet into the air. If too much acetylene has been introduced, it may burn quietly a moment at the opening, until, as more air enters at the bottom to take the place of the gas burned, an explosive mixture is formed and a report follows. ILLUMINATING GASES 48. One of the most important of these has just been considered. It is new as an illuminant, and some problems in connection with it have not been entirely solved, but it is already being extensively applied in many of the smaller towns where no gas plant exists, for railway lighting, bi- cycle lamps, etc. The fact that thus far no appliance has been invented for using it in cooking, for the reason that the excess of carbon covers the utensils with a deposit of soot, has prevented a much more extensive use. CAEBON AND A FEW COMPOUNDS 153 OTHER ILLUMINANTS 49. Besides acetylene, ordinary or coal gas, "water" gas, and Pintsch gas deserve notice. 50. Coal Gas. This is obtained by the destructive distillation of coal in iron retorts. The following diagram illustrates the essential features of a gas plant. FIG. 42. A Gas Plant. 51. Preparation. Soft coal is shoveled into the retort, beneath which is the furnace, F. When the retort is filled, the door is luted on air-tight. The heat from the furnace drives out the gaseous products from the coal in the retorts, and they are carried up to the hy- draulic main, H. From here the gas is forced by means of pumps, not shown in the diagram, through the con- densers, a series of pipes several hundred feet in length, where it is cooled and the tar condensed. This by-product is drawn off by pipes, P, to the tar- well, T, from which it is pumped into barrels. 52. From the condensers the gas goes through the scrubber, a large cylindrical tank filled with coke or lattice work, over which water slowly trickles. The 154 MODERN CHEMISTRY partition through the center causes the gas to flow down one side and up the other; the coke 'breaks up the gas into bubbles, so as to secure a thorough washing. Here the ammonia is mostly removed, and the impure aqua ammonia thus obtained is drawn off at intervals, neutral- ized with acids, and treated with lime for the preparation of the ammonia of commerce, as already described. 53. The gas next passes through the lime purifiers, a number of low cylindrical tanks, containing lime spread upon horizontal shelves. The lime dries the gas and at the same time removes the sulphureted hydrogen and the carbon dioxide. In some works, ferric oxide, Fe 2 O 3 , is used for the same purpose. From the purifier the gas passes to the gas-holder, a very large tank, where it is stored for use. On the inside of the retorts, as previously stated, there gradually collects a fine, hard deposit, known as gas car- bon, which is now a very useful by-product. 54. Water Gas. This gas receives its name from the fact that steam is used in one part of the process of manu- facture. From the boilers steam is passed into chambers, or pipes, containing charcoal or coke, heated red hot. Here the vapor and coke react upon each other, the former being decomposed, thus : C + H 2 = CO + H 2 . Two gases, carbon monoxide and hydrogen, mixed to- gether, are thus obtained. Both are combustible, and in burning produce great heat, but neither gives any light. This mixture, therefore, is next allowed to pass into re- torts, kept at a high temperature, into which kerosene, or some similar oil, is sprayed. The heat vaporizes and de- composes the oil into hydrocarbons that do not liquefy CARBON AND A FEW COMPOUNDS 155 again upon cooling, and which burn with a luminous flame. This last step is called "carbureting," and by it a gas is obtained not very different in composition from coal gas. 55. Pintsch Gas. This is the gas so frequently used for lighting street cars and railway coaches. It received its name from its inventor, who sought to improve the old and -very unsatisfactory method of lighting coaches in England by means of candles. The essential features of manufacture are similar to those of the coal gas plant. Naphtha is sprayed into retorts heated suffi- ciently to decompose the vaporized oil into other hydro- carbons. These are then passed through an improved form of condenser, a washer, and lime purifiers into the gasometer. 56. Next the gas is drawn through a cylinder known as the " freezer," or " dryer." Here, owing to the action of the pumps, it expands, and being cooled thereby, loses all its moisture. The same pumps force the gas into large tanks called "accumulators," from which it is drawn off into smaller tanks for shipment from place to place, or directly into the storage cylinders, so frequently seen under railway coaches. This light possesses not only the advantages of intensity and whiteness, which coal gas, as ordinarily burned, lacks, but unlike ordinary gas, its illuminating power is only slightly decreased by strong pressure such as is necessary for transportation in storage cylinders. 57. Natural Gas. Natural gas is formed by the de- composition of organic matter, and the main constituents are about the same as those of the other mixed gases used for illumination. 58. Composition of Illuminating Gases. With the ex- ception of acetylene, the illuminating gases noticed are all 156 MODERN CHEMISTRY mixtures. The most important constituents of coal and " water " gas are given below : COAL GAS Hydrogen . . . . H, about 46 per cent. Marsh Gas .... CH 4 " 38 Olefiant Gas .... C 2 H 4 , " 2 Carbon Monoxide . . CO, " 11 Small amounts of higher hydrocarbons, and such impuri- ties as hydrogen sulphide, ammonia, and carbon dioxide. WATER GAS* Marsh Gas CH 4 Carbon Monoxide . . .CO Hydrogen H Small amounts of higher hydrocarbons. SUMMARY OF CHAPTER Classification of free forms of carbon. Description, preparation, and uses of. a. Diamonds. b. Graphite. c. Coals. Origin of natural coal. Varieties of and differences. Petroleum and products from it. d. Charcoal. e. Coke. /. Gas carbon. g. Lampblack. * Water gas contains a larger proportion of carbon monoxide than ordinary coal gas; otherwise the two are not very different. CARBON AND A FEW COMPOUNDS 157 Reducing power of carbon. Meaning of term. Experiment. Absorbing power. For various substances. Practical applications of this power. Compounds of carbon. The oxides Names and formulae. Preparation of CO. Characteristics. Sources of CO 2 in the air. Laboratory method of preparing. Characteristics of CO 2 . Experiments to illustrate same. Practical uses of CO 2 . Soda water, sparklets, sodors, etc. Hydrocarbons Meaning of term. Three important hydrocarbons. Marsh gas Sources of, in the air. Characteristics. Protection against explosions. Olefiant gas Where found. Characteristics of. Compare witli marsh gas. Value of each in coal gas. Acetylene How prepared for use. Manufacture of carbide. Description of acetylene generators. Description of acetylene tips. Characteristics of acetylene. Experiments to show its lighting, heating, and explosive properties. Other illuminating gases. Coal gas. Method of preparing Apparatus. Plans for purifying. Different forms of gas-burners. Valuable ^-products How secured Use. 158 MODEEN CHEMISTRY Water gas. How prepared. Characteristics of. Comparison with coal gas in composition. Pintsch gas. Method of preparing and purifying. Used where. Comparison with coal gas. . CHAPTER XII FUNDAMENTAL LAWS OF CHEMISTRY 1. Quantitative Work. It may have *seemed to the student that the quantity of a reagent used in any experi- ment makes little difference. While definite amounts are usually specified, care is not often taken to use exactly that quantity. Generally the result will be the same, but if more than the necessary amount of a substance is used, the excess remains and is simply wasted. This fact is usually stated in what is known as V 2. The Law of Definite Proportions. Briefly, it is this: Two or more elements, in uniting to form a compound, always do so in the same proportion by weight. This has been illustrated somewhat in the earlier part of the book in discussing compound bodies. It is a very important law, and upon it much of the science of chemistry depends., To illustrate it more fully the student should make Jjie following experiments, using the utmost care to insure accuracy. Let him not draw his conclusions beforehand and then endeavor to make his results conform to these. EXPERIMENT 99. Fill two burettes, one with a solution of caustic soda and the other with dilute hydrochloric acid, and support them upon a stand. Carefully take the reading of each, using the lowest HP FUNDAMENTAL LAWS OF CHEMISTRY 159 part of the meniscus in doing this, as shown in the figure. Here the lowest part of the curve coincides with 38.4, and this would be the reading. Now find the weight, as accurately as possible, of a small evaporating dish. Much time can be saved here if each student will provide himself with a small pasteboard box and cover, such as blank labels are packed in. Put the box and cover upon the scale pan opposite to the dish, and pour in fine shot or sand until it is exactly counterpoised. This Represents the -weight of the dish. Put the box with its contents away where it will be safe from accident. Now, from the caustic soda burette allow 10 cc. to flow into the evaporating dish, and add one drop of phenolphthal- ein solution, or,, if more convenient, enough litmus solution to give a decidecf blue color. From the other burette, with constant stirring, let the acid flow in slowly until the color given by the phthalein barely disappears, or until the blue litmus just shows pink. Take the reading of the acid burette, and by subtracting the previous reading determine how much hydrochloric has been used. The change in the color noted above indicates that sufficient acid has been added to neutralize the alkali and form therewith .a salt." Now place the evaporating dish upon a ring-stand, or better upon a sand-bath, and evaporate slowly to dryness. Do not let the liquid boil, as some will be lost by spurting out, and be careful toward the close to withdraw the heat before the solution is entirely dry, lest the dish become so warm as to decompose some of the salt. If the heat * of the dish does not complete the evaporation, warm it very gently for another moment. When perfectly dry let the dish cool, and weigh it. In doing this put the small box and shot upon the opposite pan as be- * fore, then whatever weights are necessary to add will represent the vwight of the salt obtained. If the shot are not used, subtract the first weight from the second. Tabulate results as below: Wt. of dish + salt . . 17.103 Caustic soda used . . 10:0 cc. Wt. of dish .... 15.217 HC1 used 6.4 cc. Wt. of salt 1.886 EXPERIMENT 100. Purpos^ a continuation of the preceding. Repeat the preceding experiment, using the same amount of caustic 160 MODERN CHEMISTRY soda, but twice as much acid. The litmus or phthalein need not be added. Use the same precautions as before. Tabulate results. Wt. of dish + salt . . Caustic soda used . . 10.0 cc. Wt. of dish .... 15.217 HC1 used 12.8 cc. Salt EXPERIMENT 101. Purpose, same as above. Repeat, using 5 cc. of caustic soda solution, a few drops of litmus or one of phthalein, and then enough hydrochloric to neutralize, as in Experiment 98. Cool and weigh as before. 3. Comparison of Results. Comparing the results ob- tained, we may formulate them as below : Exp. 99. NaOH used . . 10.0 cc. Salt (NaCl) obtained . . HC1 . . 6.4 cc. Exp. 100. NaOH used . . 10.0 cc. NaCl obtained HC1 . . 12.8 cc. Exp. 101. NaOH used . . 5.0 cc. NaCl obtained HC1 . . 4. What evidence in the above experiments do you find in proof of the law of definite proportions ? Is there any agreement between the first and second of the above? Between the second and third ? Why ? EXPERIMENT 102. Further proof of the law. Carefully weigh an evaporating dish or find its equivalent in shot as before, then add a half- gram weight to the pan on which the shot is, and put into the evaporat- . ing dish sodium carbonate crystals to balance. Add a few centimeters of pure water to the carbonate, and then add dilute hydrochloric acid, a little at a time. Keep the dish covered with a sheet of glass or watch crystal so as not to lose any by its spattering out. In this way cap- tiously add the acid until the carbonate is all dissolved, or until it no longer effervesces. Now rinse off the cover-glass into the evaporating dish, and evaporate to dryness with the same precautions used before. Cool, weigh, and determine the amount of salt obtained. Sod. Carb. : Na 2 CO 3 + dish . . NaCl + dish . . Wt. of dish . . . Dish ~*^ NaCl FUNDAMENTAL LAWS OF CHEMISTRY 161 EXPERIMENT 103. Same as preceding. Pursue the same method as above, using 1 g. of the carbonate instead of a half gram. Used : Na 2 CO 3 + dish . . Obtained : NaCl + dish . . Wt. of dish . Dish . Na 2 C0 3 NaCl EXPERIMENT 104. Purpose, same as before. Repeat the preced- ing, using this time 1 g. of sodium carbonate crystals. Results : Used: Na 2 CO 3 + dish . . Obtained : NaCl + dish . . Dish . Dish . 1.500 NaCl 5. Summary. In each of the last three experiments find the ratio existing between the carbonate used and the salt obtained. 1. Na 2 CO 3 : NaCl : : 1 : x = 2. Na 2 C0 3 : NaCl : : 1 : y = 3. Na 2 C0 3 : NaCi : : 1 : z = Is there any uniformity in the value of these ratios? Do your results afford further evidence of the law of ; definite proportions ? If so, in what way ? 6. The Law of Multiple Proportions. We have learned that when two or more elements unite to form a compound they do so in a constant ratio. We have seen, however, that the same two or three elements may unite to form several compounds, and at first this may seem contrary to the statement of the preceding law. It is a modification, but not a contradiction. If a new and different compound is formed when other proportions are used, in this the quantity of the elements that enter into combination is always some multiple of the lowest. An illustration will make this plain. Thus, we are familiar with the series of nitrogen oxides : 162 MODERN CHEMISTRY Nitrogen Monoxide " Dioxide . " Trioxide " Tetroxide " Pentoxide N 2 N 2 2 N 2 3 N 2 4 N 2 6 F 2 :O L:0 fl '2 * 28:16 28:32 28 : 48 28:64 28:80 It is seen that while the weight of the nitrogen entering into combination remains constant, the oxygen is in the ratio of 2, 3, 4, and 5 times what it is in the lowest of the series. 7. This law may be proved experimentally by estimat- ing the amount of oxygen that a given weight of potas- sium chlorate, KC1O 3 , will yield, by the method previously suggested. Then, determine the amount in potassium perchlorate, KC1O 4 . In these two compounds we should find the ratio to agree with that expressed in the formulae, that is, 3 and 4 times what would be contained in a mole- cule like mercuric oxide, HgO. FIG. 44. EXPERIMENT 105. To prove the law, the work may be conven- iently done as shown in the accompanying figure. Instead of the flask, 0, a hard-glass test-tube may be used. Put into the flask about 1 g. of manganese dioxide, MnO 2 , and weigh carefully the flask and FUNDAMENTAL LAWS OF CHEMISTRY 163 contents. Then add to it about 1J g. of potassium chlorate, and weigh accurately. The difference will be the chlorate. A is a liter bottle fitted with a two-hole rubber cork. The delivery tube, d, just reaches through the corks of O and A. The tube, e, should be made in two parts, joined by rubber tubing several inches long, and should reach nearly to the bottom of both A and B. A pinch clamp will be needed at e. Xearly fill A with water, and by suction the tube connecting A and B, completely. Fasten the clamp at e. See that the corks fit air-tight, and fill B to the same height as A. Open the clamp an instant, then empty B. Replace 5, remove the clamp, and heat the chlorate carefully until water is no longer driven out of A. Let the tube cool, equalize the pressure in the two bottles as in Experiment 37, and again fasten the clamp. Determine the volume of water in B ; this will give the volume of the oxygen at the tempera- ture and pressure of the room. According to methods already given on page 96, reduce this volume to what it would be under standard conditions. Knowing the weight of a liter of oxygen, 1.43 g., find the weight of the determined volume. Weigh also the cooled flask, O, and determine its loss ; this also represents the oxygen. Next, arrange the apparatus as at the beginning. Into the hard- glass tube put about 1..4 g. of potassium perchlorate, KC1O 4 , and, after making connections, heat strongly as before until no more gas is pro- duced. Cool and weigh the flask or tube ; the loss will represent the oxygen, which may be checked up by determining the volume of the gas given off as before and reducing to standard conditions. Let the student now compare results. The two reactions are as follows : KC10 3 -f heat = KC1 + 3 O, and KC1O 4 + heat = KC1 + 4 O. From other experiments we know that the oxygen is entirely removed and that potassium chloride, KC1, remains. Then we should have the proportion Mol. wt. KC1O 3 : wt. of O in 1 mol. KC1O 3 : : 1.25 g. : m g., in which m = no. grams found by experiment above. Substituting, 122.5 : x : : 1.25 g. : m g., _ m x 122.5 1.25 164 MODERN CHEMISTRY Then, as 16 is the weight of one atom of oxygen, there would be as many atoms of oxygen in the molecule as 16 is contained times in x ; the result should agree very closely with the assumed number. In the same way, Mol. wt. KC1O 4 : wt. of O in 1 mol. KC1O 4 : : 1.40 g. : n g., in which n = no. grams found in second instance above. Or, 138.5 : y : : 1.40 g. : n g., _ n x 138.5 1.40 How does the value of y agree with the known value? 8. Combining Weights. We have previously learned that when elements or compounds react with each other in the formation of new substances, they always do so in a fixed or definite proportion. We have seen also that when several compounds are formed from the same two elements, there is one smallest quantity of which all the others are multiples. This smallest amount in the case of the nitrogen oxides was found to be 16 for the oxygen, and all the others were multiples of this. Therefore, 16 is regarded as the atomic weight of oxygen, and in all chemical reactions into which it enters, this, or some multiple of it, is its combining weight. EXPERIMENT 106. To find the combining weight of copper. Put into a beaker 2| g. of clean, bright copper, accurately weighed, and dissolve slowly in nitric acid somewhat diluted. Use every precaution to prevent loss by spurting, just as in other similar work, and when the copper is all dissolved, transfer the solution to a weighed evapo- rating dish, as small as will conveniently hold the solution ; carefully rinse off the cover-glass and the beaker into the evaporating dish, and evaporate to dryness. We now have a blue salt, copper nitrate. Be sure it is perfectly dry, and then remove the sand-bath, or any other protection used for the dish, and gradually increase the heat until all particles of the blue salt have been changed to a black compound. A dull red heat is generally necessary for this. We now have copper FUNDAMENTAL LAWS OF C&EMISTRY 165 oxide, CuO. Cool and weigh carefully. Determine how much oxygen has combined with the copper by subtracting the amount of copper used from the weight of oxide obtained. Dish + CuO . m + x ; m = wt. of dish ; x, of CuO. Dish + Cu . m + n; . n = wt. of Cu = 2 g. O m + x - (m + n) = y. Numerous experiments have shown that the combining weight of oxygen is 16 ; using this as a basis, we can determine what it is for copper : Wt. of O found : wt. of Cu used : : comb. wt. of O : comb. wt. of Cu. or, Wt. of O : 2J g. Cu : : 16 : z. From this proportion the combining weight of copper should be found to be approximately what is given in the table on page 9. The sources of error are liable to make the difference comparatively great, but the result should not vary too much. EXPERIMENT 107. Purpose, same as above. Use 3 or 4 g. of finely powdered copper nitrate which has not been exposed to the air any length of time. Be sure the exact weight is known, then heat in a small evaporating dish, or better, in a porcelain crucible, cautiously at first ; when the nitrate is converted into the black oxide as before, cool and find the weight. Experiment has shown that 1 g. of crystal- lized copper nitrate contains 0.2619 g. of metallic copper. From this determine the amount of copper represented by the 3 g. (or 4 g.) of nitrate used. Wt. of dish + CuO ...:... Wt. of dish CuO Cu O Wt. of O : wt. of Cu : : 16 : x. Does this give practically the same combining weight for copper that the preceding did? If the results do not correspond fairly well with each other and with the table, the experiments should be repeated. EXPERIMENT 108. To find the combining weight of tin. For this use granulated tin. If not at hand, procure a quantity of pure tin 166 MODERN CHEMISTRY foil, melt it in an iron ladle, and pour into cool water. Remove from the water and dry it, when it will be ready for use. Weigh out care- fully 2 g. of the granulated tin, and treat with nitric acid in an evapo- rating dish. Take care always to avoid loss by spurting. Evaporate slowly to dry ness, and then gradually heat the white residue to dull redness. Wt. of dish + SnO 2 . Wt. of dish SnO 2 Wt. of. Sn 2 It has been found by analysis that the amount of oxygen in this compound indicates two atoms to the molecule, hence in making our calculations that amount must be used. Then we have wt. of O found : wt. of Sn used : : 32 (2 x 16) : x. EXPERIMENT 109. Repeat the above experiment, using 2| or 3 g. of tin, and make calculations as before. How do the results in the two experiments agree ? If they do not correspond fairly well with the atomic weight given in the table, allowing for errors in weighing, the experiment should be repeated. 9. Such experiments as the above might be endlessly multiplied. We have found in these, as has been the case in an indefinite number of instances in which chem- ists have done the work with the utmost care, that every element combines with' others in some exact proportion by weight, and whether we use much or little of the element, in the same compound the ratio never changes. This fact is of the utmost importance, for upon it depends much of the science of chemistry. It is this that enters into the application of chemistry to the arts and manufac- tures, and renders its results so sure and unchanging. 10. Some Application of the Laws of Combination. Knowing that 'the laws of combination are true, we may make use of the principles in determining the strength FUNDAMENTAL LAWS OF CHEMISTRY 167 of acid or alkaline solutions. The following work will illustrate this. EXPERIMENT 110. To determine the strength of any hydrochloric acid solution in the laboratory. Put the acid to be tested into a burette and take the reading. From this allow 10 cc. to flow into an evaporating dish, add a drop or two of litmus or phthalein, and then, from another burette, after taking the reading, run in a solution of caustic soda until the solution in the evaporating dish is neutralized, as in previous work. Evaporate slowly to dryness, cool, and weigh. Subtract the weight of the dish to determine the salt obtained. Sup- pose this is 0.585 g. Now we know that caustic soda and hydrochloric acid react with each other according to the following equation : NaOH + HC1 = NaCl + H 2 O. From this we see that one molecule of pure hydrochloric acid yields one of sodium chloride, or 36.5 parts by weight of acid give 58.5 of salt. The 0.585 g. of salt would thus correspond to 0.365 g. of acid, the amount in 10 cc. of the solution used. Then in a liter, 1000 cc., there would be 100 times this amount, or 36.5 g. of pure acid. The liter of acid then ought to weigh 1000 g. -f 36.5 g. = 1036.5 g. The question simply is this : 36.5 g., the amount of pure acid, is what per cent of 1036.5 g., the weight of the acid solution ? This is found to be about 3 per cent. EXPERIMENT 111. Repeat the preceding experiment, neutralizing 10 cc. of the hydrochloric acid with caustic potash. Make your cal- culations from the following equation : KOH + HC1 = KC1 + H 2 O. Do your results agree with the preceding as to the per cent strength of acid? 11. To determine the Amount of Caustic Soda or Potash in the Solutions used above. We know that when an acid and an alkali are put together, they neutralize each other to form a salt. If then we know how much acid is con- tained in a solution, and measure the amount of the latter used, having some means of knowing when the alkali is 168 MODERN CHEMISTRY exactly neutralized, we can easily calculate the amount of alkali contained in a given volume of solution. EXPERIMENT 112. Suppose we are required to determine the number of grams of sodium hydroxide in 1 liter of the solution. We know the reaction is NaOH + HC1 = NaCI + H 2 O, or by weight, 40 + 36.5 = 58.5 + 18. That is, 40 g. of caustic soda are necessary to neutralize 36.5 g. of hydrochloric acid. Suppose now we have a solution of acid that con- tains 3.65 g. of pure hydrochloric acid to the liter, then 1000 cc. HC1 would neutralize 4.0 g. of NaOH. Then, if with 100 cc. of caustic soda solution we used 20 cc. of the acid solution, we should have this proportion : 1000 cc. HC1 : 4.0 g. NaOH : : 20 cc. HC1 : x g. NaOH ; x = .08. That is, in 100 cc. of the solution of caustic soda there are .08 g. of the solid alkali dissolved ; then in 1 liter there would be 10 times as much, or .8 g. For such work as this we very often use oxalic acid instead of hydrochloric, because it is easy to weigh out, and forms a good working solution. Its formula is H 2 C 2 O 4 , 2 H 2 O. With caustic soda the reaction is 2 H 2 O, H 2 C 2 O 4 + 2 NaOH = Na 2 C 2 O 4 + 4 H 2 O, or by weight, 126 + 80 =134 +72. That is, 126 g. of oxalic acid will neutralize 80 g. of caustic soda. Suppose for work we weigh out 6.3 g. of oxalic acid and dissolve in 1000 cc. of pure water. This will be our standard solution of acid. To find how much caustic soda in 1 liter of solution. Measure out accurately into a beaker 50 cc. of the alkali solution, and add one drop of phenolphthalein, or about 1 cc. of litmus solution. Next take the reading of a burette containing the standard oxalic acid solution, and with constant stirring let the acid drop in slowly until, finally, by the addition of a single drop the red color of the phenol disappears, or the FUNDAMENTAL LAWS OF CHEMISTRY 169 blue of the litmus changes to red. Again read the burette and deter- mine how much acid has been used. Suppose it has been 10 cc. Then to calculate, 1000 cc. acid 4.0 g. NaOH : : 10 cc. : x NaOH ; x = .04 g. NaOH, the amount in 50 cc. of NaOH solution used. In 1000 cc. there would be 20 times as much, or .8 g. PROBLEM 1. Let the teacher make up a solution of caustic potash with distilled water, and have the student determine the number of grams used to the liter. PROBLEM 2. In the same way let the student determine the amount of common salt in a solution by using in the burette a solution of silver nitrate containing 17 g. per liter. To determine when sufficient silver nitrate is used, add to the common salt solution sufficient potassium chromate solution to give a yellow color. With constant stirring run in the silver nitrate until the precipitate that forms shows the faintest red tinge. The reaction is NaCl + AgNO, = AgCl + NaNO 3 , or by weight, 58.5 + 170 = 143.5 + 85. That is, 170 g. of silver nitrate will precipitate the chlorine in 58.5 g. of salt, or if 17 g. of silver nitrate were used to make a liter of the solution, then 1000 cc. of silver nitrate would precipitate the chlorine in 5.85 g. of salt. Or, 1000 cc. AgNO 3 : 5.85 g. NaCl : : m cc. AgNO 3 x g. NaCl, in which m is the number of cubic centimeters of silver nitrate solu- tion used with the amount of common salt solution taken. If this latter is 20 cc., or ^ of a liter, then 50 m = number cubic centimeters AgNO 3 necessary to precipitate the chlorine in 1 liter. 12. Displacing Power of Metals. We have seen in pre- paring hydrogen that various metals have the power of reacting with certain acids to displace the hydrogen contained. Of course this displacing power is in accord- ance with the valence of the element (see chapter on 170 MODERN CHEMISTRY Valence), and the following plan may be used to deter- mine it : EXPERIMENT 113. Let a part of the students perform this experi- ment, another portion No. 114, and another, 115. Arrange apparatus as for Experiment 37 or 105. Use a wide-mouth, 4 oz. bottle instead of a flask or hard glass tube. Put into this 1 g. of finely granulated zinc and a short test-tube containing 15 or 20 cc. strong hydrochloric acid. Make connections air-tight. Fill the delivery tube with water and equalize pressure in the two bottles as in other experi- ments. When all is ready, tip the generating bottle so that the acid shall be spilled upon the FIG. 45. zinc. When the zinc has disappeared and the generator has cooled, equalize the pressure in the two bottles and attach the clamp. Measure the water expelled. Assuming this to be the volume of the hydrogen displaced from the acid by 1 g. of zinc, allow for vapor tension and reduce to standard conditions. Suppose, for example, it is found that m cc. of water have been forced over by the hydrogen, and that by reducing this volume to standard conditions, we obtain n cc. as a result, then as 1 liter of hydrogen weighs .0896 g., n cc. would weigh n x .0896 , , , 10QQ = w g. of hydrogen. Then we should have the proportion w g. of H : 1 g. of Zn : : x : 65 ; that is, the weight of the hydrogen obtained is to the weight of the zinc used in displacing it as x is to 65, the atomic weight of zinc. This should give for the value of x, approximately, 2. Then as the hydrogen atom is the standard, or 1, in this case x represents the weight of two atoms of hydrogen. In other words, the zinc atom has the power of displacing two atoms of hydrogen. EXPERIMENT 114. AVith apparatus arranged as in the preceding experiment, let the student use one gram of aluminum wire cut into small pieces. No heat will be necessary if strong hydrochloric acid be used, and the chemical action, slow at first, will soon become very rapid. Determine as before the volume and weight of the hydrogen set free. Then we have 8ULPHV& AND ITS COMPOUNDS 171 wt. of H obtained : wt. of Al us d : : x : 27, atomic wt. of Al, and x - ? From this what can you say is the displacing power of the aluminum atom? EXPERIMENT 115. In exactly the same way try 1 gram of magne- sium ribbon, cut into small pieces. Hydrochloric acid somewhat diluted had better be used, as the action is very rapid. Make your corrections for temperature and pressure, and calculate as before. What do you find for the displacing power of the magnesium atom ? SUMMARY OF CHAPTER Statement of Law of Definite Proportions. Experiments illustrating it. Law of Multiple Proportions. How illustrated. Combining weights. Method of determining by experiment. For copper. For tin. Practical application. Method of determining amount of acid or alkali in a solution. Method of determining valence or displacing power of metals. .CHAPTER XIII SULPHUR AND ITS COMPOUNDS 1. Where found. Sulphur is an element that has been known from very early times. By some of the alchemists it, together with mercury, was regarded as forming all of the metals. It is a native of volcanic regions, and is found in abun- dance in Sicily and to some extent in Iceland. There are said to be some deep deposits in the Southern States, but 172 MODERN CHEMISTRY. they have not been developed. In 'the form of compounds with the metals, sulphur is found abundantly and very widely distributed. Some of the more common compounds are gypsum, CaSO 4 ,2H 2 O, iron pyrite, FeS 2 , galena, PbS, and zinc blende, ZnS. In the form of hydrogen sul- phide, it is found in many mineral springs and is often emitted from volcanoes. 2. For many years Sicily had a monopoly of the sul- phur trade. It occurs there in almost unlimited quan- tities, mixed with earthy matter. This mixture may be partially purified by a method similar to that employed in the preparation of charcoal. Large piles of the crude sulphur are heaped up and covered with earth and sod. FIG. 46. The mass is then ignited and a part of the sulphur in burning melts the remainder, which runs out into trenches or vats, leaving the earthy matter behind. 8ULPHUE AND ITS COMPOUNDS 173 3. For many purposes the sulphur thus obtained needs further purification. It is heated and vaporized in retorts, the vapors passing over into cool chambers and condensing upon the walls in the form known as flowers of sulphur. If the operation continues for a length of time, however, the walls become heated enough to melt the sulphur that forms upon them. It is then allowed to run out into molds, in which form it is known as brimstone or roll sulphur. (See Fig. 46.) S is a cylinder in which the sulphur is melted, V, a retort where it is vaporized, and E, the condensing chamber. 4. New Source of Supply. From the fact that Sicily controlled the sulphur trade, prices rose so high at one time that the English manufacturers were obliged to resort to some other source of supply. Sulphur was used extensively in making sulphuric acid for the manu- facture of soda crystals. It was found that by roasting iron pyrite, FeS 2 , a compound that had been hitherto alto- gether worthless, the sulphur dioxide could be obtained ; or if the ore was heated in retorts sealed up to prevent access of air, the sulphur was not oxidized, and could be condensed. As the pyrite is very abundant, and the method of obtaining the sulphur cheap, this at the present time furnishes not only about all the sulphur needed in making sulphuric acid, but even more, so that the demand for Sicilian sulphur has greatly decreased. The reaction that takes place when iron pyrite is heated in sealed retorts is g ^ + heat = 2 S + Fe 3 S 4 . 5. Characteristics of Sulphur. Sulphur is a yellow, brittle solid, twice as heavy as water. It is seen in a number of forms, of which the flowers and roll sulphur have been mentioned. It also occurs crystallized. 174 MODERN CHEMISTRY EXPERIMENT 116. Into a test-tube put about a cubic centimetei of carbon disulphide, and add a little sulphur. When the latter has dissolved, pour off the clear solution upon a watch crystal, and allow it to evaporate slowly to dryness. The sulphur will form in crystals, the shape of which may be recognized if the evaporation is slow. If necessary, however, examine with a magnifying glass. What form have they ? EXPERIMENT 117. Fill a small crucible nearly full of sulphur, and heat till it is melted. Allow it to cool, and when a crust has formed over the surface, break an opening in the top and pour out what remains molten. Let it cool a little more and break open the mass. What kind of crystals have formed? EXPERIMENT 118. Put 4 or 5 g. of sulphur into a test-tube and warm. Note how it changes, first melting to form a light yellow- colored liquid, then becoming quite thick again and very dark, then thin again. At this last stage, pour out the sulphur into cold water and note its condition. This is called amorphous sulphur, or some- times plastic sulphur. 6. Sulphur is found in a large variety of crystallized forms. The octahedral and the long, needle-like crystals have been seen. Upon standing for some time these gradually change into other forms, modifications of the two. Amorphous sulphur is dark-colored and elastic, somewhat like rubber. It is regarded as an allotropic form. Sulphur is insoluble in water, hence has no taste ; it is also without odor. As we have seen, it is soluble in carbon disulphide. It is combustible, burning with a pale blue flame, and forming the well-known irritating gas, sulphur dioxide. At high temperatures sulphur combines readily with most of the metals, forming sulphides. This has been shown already in preparing ferrous sulphide by heating iron filings mixed with sulphur. Copper turnings serve equally well. 7. Comparison of Ozone with Allotropic Sulphur. In the case of ozone, we have seen that its molecule is differ- SULPHUR AND ITS COMPOUNDS 175 ent from that of the oxygen molecule. The same is be- lieved to be true of sulphur arid its allotropic form, as well as of all other elements which show the same variation. We cannot prove this for sulphur, but there are some facts which make this theory strongly plausible. Thus, if the vapor is weighed at 1000 temperature, it is only one-third as dense as when weighed at 500. 8. Uses of Sulphur. Sulphur is used largely in the manufacture of gunpowder, the other two constituents being charcoal and saltpeter. ' These are united in about the following proportions : Sulphur 12 per cent Charcoal 13 " Saltpeter 75 " " Greek fire," which played so important a part in the early centuries, and the composition of which was kept a secret for several hundred years, differed very little from the gunpowder of the present time. Sulphur is employed to some extent in the manufacture of rubber goods, espe- cially vulcanite, and considerably in fumigating buildings; it is used largely in making sulphuric acid. Because of its low kindling-point sulphur has been used very exten- sively in the manufacture of matches, but the irritating gas produced, and the slowness with which such matches burn, have led to the substitution of other substances. COMPOUNDS OF SULPHUR 9. Hydrogen Sulphide, H 2 S. This gas, known also as sulphureted hydrogen, occurs in many mineral springs, which give it off abundantly; it is sometimes emitted from volcanoes, and is noticed in the decay of eggs and other similar substances. 176 MODERN CHEMISTRY 10. How prepared. For laboratory purposes hydrogen sulphide is always prepared by treating ferrous sulphide, FeS, with sulphuric acid. EXPERIMENT 119. Owing to the offensive odor of the gas, it should be prepared in very small quantities, and kept from access to the room as much as possible. Put into a test-tube a small bit of ferrous sul- phide, cover it with water, and add a few drops of strong sulphuric acid. Action will begin at once. Notice the odor of the gas. Has it any color ? Attach a jet and ignite it. With what kind of a flame does it burn ? Notice the odor given off by the burning gas. Hold a cold beaker over the flame. What do you- see depositing upon it? What are the two products formed when hydrogen sulphide burns? Write the reaction. The reaction that takes place when hydrogen sulphide is prepared is seen below : FeS + H 2 SO 4 = FeSO 4 + II 2 S, 9 FeS + 2 HC1 = FeCl 2 + H 2 S. Hydrochloric acid may be used instead of sulphuric acid with good results. 11. Characteristics of Sulphureted Hydrogen. It is a colorless gas, having a very disagreeable, nauseating odor ; is somewhat poisonous, and should not be inhaled. It is inflammable, burning with a bluish flame, is a little heavier than air, and somewhat soluble in water. It has the power of forming precipitates with solutions of many metallic salts. EXPERIMENT 120. Into a test-tube put a little of a mercuric chloride solution, into another a solution of antimony tartrate, into a third arsenic trioxide dissolved in hydrochloric acid and water. Attach a delivery tube to a hydrogen sulphide generator, and pass the gas through each of the solutions. Notice the color of the precipitates obtained. Lead salts are very sensitive to the action of hydrogen sulphide, and are used in testing for its presence. 12. Use of Hydrogen Sulphide. Mineral waters con- taining this gas in solution are supposed to be beneficial SULPHUR AND ITS COMPOUNDS 177 co health. With this exception, about the only use for hydrogen sulphide is in the laboratory, as a reagent, espe- cially in making analyses of unknown solutions. Many of the metals in the form of salts are converted by hydro- gen sulphide into insoluble sulphides. Such metals, there- fore, when treated with the gas, may be separated from others which are not so precipitated. EXPERIMENT 121. The above statements will be made plain by this experiment. Into a beaker put a few cubic centimeters of a solu- tion of mercuric nitrate and as much of zinc sulphate ; add a few drops of hydrochloric acid, and pass through it a current of sulphureted hydrogen until the odor of the gas is still perceptible after shaking the solution. Then filter out the black precipitate and test the clear filtrate for zinc with ammonia, as you have done in Chapter VII, Sec- tion 12. Have you succeeded in separating the two metals ? 13. Oxides of Sulphur. There are two of these com- pounds, the dioxide and the trioxide, SO 2 and SO 3 . It is only the first that is of special importance or interest to us. 14. Sulphur Dioxide, S0 2 . This is also known as sul- phurous anhydride, because by passing it into water sul- phurous acid is formed. It is the familiar, irritating gas always produced when sulphur is burned in the air. 15. How prepared. For laboratory purposes sulphur dioxide is prepared by treating copper turnings with strong sulphuric acid. The reaction is usually indicated as follows : Cu + 2 H 2 S0 4 = CuSO 4 + 2 H 2 O + SO 2 . If this is compared with the reaction of zinc and sulphuric acid upon each other, it will be seen to be very different. Zinc is acted upon by cold, dilute acid, while copper re- 178 MODERN CHEMISTRY quires the acid hot and concentrated. It is probable that, as with zinc, hydrogen is first formed, thus : Cu + H 2 SO 4 = H 2 + CuSO 4 , and that this nascent hydrogen immediately attacks an- other molecule of sulphuric acid, decomposing it, thus : H 2 S0 4 +H 2 =2H 2 + S0 2 . Putting these two reactions together, we have the one given above. EXPERIMENT 122. Put into a test-tube a few copper turnings and nearly cover with strong sulphuric acid. Heat moderately until the fumes begin to come off, and collect two or three bottles of the gas as you have carbon dioxide, by downward displacement. What is the odor of the gas ? Test it to learn whether it will support combustion or will burn. What can you say of its density ? Try its effect upon moistened red and blue litmus paper ; state results. Pour into one bottle of the gas a few cubic centimeters of litmus, cochineal, or some other vegetable solution, and shake it. What happens? Suspend in another bottle some colored paper, or silk or straw goods, moistened, and allow to remain some time. State results. Invert another bottle or test-tube filled with sulphur dioxide, over a small evaporating dish of water. Does the water rise in the tube? Why ? Test the water with blue litmus paper ; what effects ? What has been formed with the water? 16. Characteristics of Sulphur Dioxide. It is a very irritating, colorless gas, considerably heavier than air. It is soluble in water, forming an acid solution, which, how- ever, is very unstable. It will neither burn nor support combustion, though magnesium ribbon will burn in it with difficulty as it does in carbon dioxide. It is readily lique- fied by passing the gas through a spiral tube, surrounded by ice and salt. In the liquid condition it is limpid, trans- parent, and very slightly yellow in color. ^ / ' SULPHUR AND ITS COMPOUNDS 179 17. Sulphur dioxide is a great reducing agent, like carbon, but more active. That is, it has the power of abstracting oxygen from other substances. If sulphur dioxide is passed into a bottle containing nitrogen te- troxide, NO 2 , the red fumes will soon disappear because the tetroxide has been deprived of a portion of its oxygen and converted into the dioxide, thus : S0 2 + N0 2 = S0 3 + NO. Likewise a current of sulphur dioxide passed into a solution of potassium dichromate, or permanganate, will deprive them of a portion of their oxygen, changing the first to a compound, green in color, and rendering the second colorless. It will be important to remember this property on account of its relation to the manufacture of sulphuric acid, to be shown later. 18. Uses of Sulphur Dioxide. These have already been mentioned. It is used frequently as a disinfectant or fumigant, and for bleaching silk and straw goods. Evap- orated fruits, especially apples and peaches, owe their white, almost natural, color to the bleaching effects of sulphur dioxide, which is allowed to flow over the fruit as it is put into the evaporator. Its most extensive use is for making sulphuric acid. 19. Sulphur Acids. Sulphur forms several acids with hydrogen and oxygen, not all of which are important. The best known is sulphuric, H 2 SO 4 , also called oil of vitriol. EXPERIMENT 123.* Arrange three flasks as shown in Fig. 47, one for the generation of sulphur dioxide by the treatment of copper with * If it is found necessary to use simpler apparatus, fill a flask with sul- phur dioxide, and introduce into it a swab of cloth upon the end of a glass rod, moistened with nitric acid. Soon, brown fumes will begin to appear, 180 MODERN CHEMISTRY sulphuric acid, another containing nitric acid and copper turnings f OT the preparation of nitrogen dioxide, and a third containing water to be converted into steam. Connect with a large flask, Z>, which has a fourth tube to allow the entrance of air. t( FIG. 47. When the nitrogen dioxide enters the flask, Z>, containing air, it combines with the oxygen, forming the tetroxide, thus : 2 NO + O 2 = 2 N0 2 . Immediately the sulphur dioxide attacks this compound of nitrogen, taking away two atoms, reducing it to the dioxide again, thus : NO 2 + SO 2 = NO + SO 3 . The dioxide thus formed, with the oxygen of the air, again combines to form the tetroxide, and so serves as a carrier of oxygen from the air to the sulphur dioxide. Next, the sulphur trioxide combines with the water introduced in the form of steam, producing sulphuric acid, thus : SO 3 + H 2 O = H 2 SO 4 . showing that the nitric acid is being decomposed and the sulphur dioxide converted into the trioxide. Now add a few cubic centimeters of water, and shake. The flask will contain a dilute solution of sulphuric acid. SULPHUR AND ITS COMPOUNDS 181 20. To test the Acid prepared. Put a little of the acid into a test-tube and add 1 or 2 cc. of a solution of barium chloride. If a white precipitate forms, which is not soluble in nitric or hydrochloric acid, or both to- gether, sulphuric acid is indicated. 21. The Manufacture of Sulphuric Acid. This acid is now prepared in immense quantities. The United States and Great Britain each produce annually about one million tons, and Germany is not far behind. Formerly, sulphur was used to prepare the sulphur dioxide for the manufac- ture of this acid, but, as stated above, the attempt to con- trol the entire output of the Sicilian mines raised the price to such an extent that sulphuric acid manufacturers sought other sources, and finally discovered the present method. The pyrite is roasted in the presence of plenty of air, and the following reaction takes place : 2 FeS 2 + 11 O = Fe 2 O 3 + 4 SO 2 . 22. These fumes are conducted into large chambers lined with sheet lead, into which jets of steam are con- stantly sprayed, together with nitric acid vapors, obtained by treating sodium nitrate with sulphuric acid. The reactions that take place in these lead chambers are the same as already described. The amount of nitric acid necessary is very small, and theoretically might be used indefinitely, but practically it is gradually carried by the draughts of air into the flues and must be replaced. The sulphuric acid thus prepared collects upon the floors of the rooms, which are so large that a dancing party of a hundred couples could easily be held in them, and is called chamber acid. It is only moderately strong, and is next evaporated in leaden vessels until a specific gravity of a little over 1.7 is reached, when it begins to attack the 182 MODERN CHEMISTRY lead. It is next concentrated in glass or platinum retorts. (See Fig. 48.) The following simpler plan has recently FIG. 48. Apparatus for condensing Sulphuric Acid. been introduced. Sulphur dioxide and air are passed over finely divided platinum or ferric oxide, whereby sulphur trioxide is formed. This is then treated with water. 23. Characteristics of Sulphuric Acid. It is a colorless, sirupy liquid ; it received the name oil of vitriol on this account, and because it was made horn green vitriol, ferrous sulphate. It is not a volatile acid, and, unlike nitric or hydrochloric, gives off no odor. It is very heavy and very corrosive. Organic matter exposed to it is charred black, as already noticed. It has great affinity for water, so much so that a beaker two-thirds filled with strong acid will in a few weeks, if left exposed to the air, absorb enough moisture to cause the beaker to overflow. Like- wise when strong acid is added to water, or vice versa, the mixture becomes very hot, reaching nearly 100 C., owing to the strong affinity of the two for each other. 24. It is upon this principle that sugar, C 12 H 22 O n , is charred. The hydrogen and oxygen, being sufficient to form eleven molecules of water, are abstracted, and the carbon remains behind as a black mass. Upon the same SULPHUR AND ITS COMPOUNDS 183 principle depends its use as a drying agent for various gases. They are made to bubble up through a bottle of strong sulphuric acid, and by this means lose their moisture. 25. Uses for Sulphuric Acid. It will be concluded from the vast quantities manufactured that sulphuric acid is a very important article of commerce. It is the most useful of acids, and almost all the others are dependent upon it for their preparation. In the manufacture of soda crystals, Na 2 CO 3 , by the Leblanc process (see page 211), sulphuric acid is indispensable. This salt, Na 2 CO 3 , is the basis for all soap manufacture as well as for glass, baking powders, etc. We can thus see the commercial importance of sulphuric acid. 26. Another very extensive use is in the manufacture of artificial fertilizers from bones. When they have had the bone oil and gelatine removed, and as bone-black are no longer valuable for clarifying sugar, the bones are treated with sulphuric acid. This converts the phosphates present into a soluble form that may be used by plants. Sulphuric acid is also used in the manufacture of such explosives as nitroglycerine and gun-cotton, for making glucose, and in some of the processes of electroplating and electrotyping. 27. Other Acids of Sulphur. Sulphurous Acid, H 2 S0 3 . This acid has already been mentioned, as well as its method of formation and its instability. We also have Hyposulphurous, H 2 SO 2 . Fuming Sulphuric, H 2 S 2 O 7 , which is really ordinary sulphuric acid, charged with sulphur trioxide, SO g . 1$4 MODERN CHEMISTRY 28. Thiosulphuric Acid, H 2 S 2 3 . This last is of some interest because it is the basis of the salts known as thiosulphates, the best known of which is sodium thiosul- phate, Na 2 S 2 O 3 . From a mistaken idea of its composition sodium thiosulphate was first named hyposulphite, and is still commonly sold under that name. This is the photog- rapher's "hypo." SUMMARY OF CHAPTER Sulphur Where found. Forms in which it occurs. Sources of commercial supply. Methods of purification. Characteristics of sulphur. Various forms How prepared. How different. Uses of sulphur. Compounds. With hydrogen Two names for the gas. Occurrence. Method of preparing. Characteristics of, and proof by experiment. Use of. With oxygen Names and formulae. Preparation of the more important. Comparison of method with that of making hydrogen. Characteristics of SO 2 . Uses. With hydrogen and oxygen. Most important Commercial name. How manufactured Explanation of the chemical changes involved. Characteristics of H 2 SO 4 . Uses. CHAPTER XIV SILICON AND ITS COMPOUNDS GLASS SILICON : Si = 28 1. Abundance. Silicon is never found free, but in the form of compounds is one of the most widely dis- tributed as well as one of the more abundant of the non-metallic elements. Sand, an oxide of silicon, SiO 2 , is familiar to all ; quartz, crystallized or massive, including the agate, amethyst, opal, and other stones, is another variety of the same substance. All soils contain it to a greater or less extent, and it is taken up by plants and enters into their structure. Combined with sodium, cal- cium, magnesium, aluminum, and other metals it forms silicates which are very abundant. In this class may be placed granite, mica, and many other substances. 2. Character of Silicon. Silicon has been prepared in such limited quantities that not a great deal is known about it. It occurs in three forms, the amorphous and the crystallized or lustrous being the two most impor- tant. At high temperatures it combines readily with oxygen or with carbon dioxide, forming the dioxide. COMPOUNDS OF SILICON 3. As already stated, silica or silicon dioxide, SiO 2 , is the most abundant compound. In the crystallized form it is often called rock crystal, and is found in hexagonal prisms, often more or less modified. Owing 185 186 MODERN CHEMISTRY to the presence of foreign substances, silica often assumes a variety of colors, and is known as rose quartz, smoky quartz, etc. It is very hard, being seven in the scale, is brilliant, highly refractive when cut, and is often used for ornaments in imitation of diamonds. It melts at about 2000 C., and is soluble in alkalies as well as in hydro- fluoric acid. 4. It is from the fact above mentioned that siliceous incrustations occur about many geysers. These springs are alkaline in character, and at the high temperature present beneath the surface dissolve considerable quanti- ties of silica; when the water becomes cold and exposed to the action of the air, it is not able to hold the silica, and this is deposited upon any bodies on which the water may fall. The power of alkalies to dissolve silica may often be observed in the laboratory, where bottles contain- ing ammonia, caustic soda and potash, sodium carbomtte solutions, etc., become etched or rough on the inside, and the glass stoppers so tight as to render their removal an impossibility. 5. The Silicates. Theoretically, silica is the anhy- dride of silicic acid ; that is, Si0 2 + 2 H 2 = H 4 Si0 4 . But water added to the oxide in this case produces no reaction. The silicates, however, are based upon this acid. They are abundant, and many of them are very complicated in composition. As silicic acid is tetrabasic, the hydrogen may be replaced by a variety of elements, even in the same molecule ; thus, we might have NaAlSiO 4 : Sodium Aluminum Silicate. CaMgSiO 4 : Calcium Magnesium Silicate. SILICON AND ITS COMPOUNDS GLASS 187 NaKCaSiO 4 : Sodium, Potassium Calcium Silicate. H 8 Mg 5 Fe 7 Al 2 Si 3 O 18 : Mica, etc. 6. Preparation of Silicic Acid. Silicic acid may be pre- pared from " water glass," that is, silica dissolved in boiling caustic soda, or potash, by adding a little strong hydro- chloric acid till the solution is no longer alkaline. Then a jellylike mass will be precipitated, which is silicic acid. By filtering this out and igniting when dry we again obtain the oxide. EXPERIMENT 124. Let the student thus prepare some soluble " water glass " and the silicic acid from it. 7. Though silica is insoluble in water and has such a high melting point that only such temperatures as that secured by the oxy hydrogen blowpipe or the electric fur- nace will fuse it, still, if mixed with sodium carbonate and strongly heated in a blast lamp for a few minutes, it is converted into a soluble form, sodium silicate, and may then be readily taken up by water. 8. Glass. This is an artificial silicate that has been manufactured in some form or other for probably 4000 years. Several of the nations of antiquity were famous for their wonderful glasswork ; in beauty of coloring, their achievements have probably not been surpassed in modern times. But the applications of glass are now so varied and so adapted to the necessities of life, as well as to the luxuries, that it would seem impossible to do with- out it. Every year sees the manufacture of hundreds of millions of bottles, and tons of other kinds of glassware ; and the art of glass blowing and working has reached such a high state of perfection that glass objects, from their nature almost inconceivable, are now of frequent manu- facture. 188 MODERN CHEMISTRY We have seen above that if silica is fused with sodium carbonate, a new compound is formed which is quite soluble. If, however, we mix calcium carbonate, or chalk, with the silica, together with the sodium carbonate, and fuse the mixture, we then obtain a double silicate of sodium and calcium that is quite insoluble in water and in all acids, except hydrofluoric. 9. Varieties of Glass. There are many varieties of glass. As potassium salts are so closely related to those of sodium, it is obvious that potash could be used instead of soda. In fact, glass was first made entirely in this way. Nearly all the best chemical and physical apparatus is still made from potash salts, and this variety is known as Bohemian glass. It is much harder to melt than glass made from sodium carbonate. 10. If an oxide of lead is used with the silica and potash, we obtain a glass that is very soft and easily worked, known as flint glass. It has a very high refractive power, and on this account is used for telescopes and all kinds of optical instruments. In the purest form it is known as strass or paste, and from this are made the so-called paste diamonds. These are so lustrous and highly refractive that, except in hardness, it is difficult for any but experts to distinguish them from the genuine article. 11. Ordinary glass, known as crown glass, from which windows and the great majority of glass utensils are made, is a silicate of lime and sodium, as already described. Ordinary sand contains a considerable amount of iron in the form of an oxide. This gives to the glass used for ordinary bottles and for all the cheaper grades the well-known greenish color, which, however, may be re- moved by the addition of a small quantity of manganese dioxide. SILICON AND ITS COMPOUNDS GLASS 189 12. Much of the plain glassware used at present is molded just as any casting would be in an iron foundry. Window glass is first blown into a long cylinder ; this is cut open and flattened while still hot by means of heavy rollers. Plate glass for large windows and heavy mirrors is cast. The molten glass is poured upon a table of the desired size, allowed to cool, and the surface afterwards ground and polished. 13. All glass articles must be carefully annealed, other- wise they would be so brittle as to have little value. The glass, as soon as shaped, is placed in an oven, and during several days is cooled so slowly that the molecules have time to adjust themselves to stable positions. Indeed, so well is this annealing done that glass vessels are made for use in chemical work that may be heated strongly and plunged into cold water immediately without danger of breaking. SUMMARY OF CHAPTER Silicon. Abundance of it in nature. Some familiar forms. Compounds of silicon. The oxide Some common forms. Characteristics of. Glass Importance of. What glass is. Kinds of glass. How different in properties. Making of window glass and otker forms. Annealing of glass articles. CHAPTER XV PHOSPHORUS AND ITS COMPOUNDS PHOSPHORUS : P = 31 ^ 1. Occurrence. Phosphorus has been known for about two and a quarter centuries, but it is only since 1833 that it has had any real practical value. Owing to its strong affinity for oxygen it is never found free, but in the form of compounds it is very widely distributed. It is a con- stituent of many rocks, and, from their decomposition, also of soils. From this source plants take it up and store it away in the seeds and fruits ; plants, being used as foods, transfer it to animals, where it is found in the nerve centers and bones. Q^. Manufacture of Phospho- rus. It is obtained almost altogether from bones. These are put into retorts and heated, much as coal is for the prepa- ration of illuminating gas. The volatile products are thus driven off and their valuable portions condensed. The bones are re- duced to what is known as bone-black, or, if not desired for clarifying sugar, to bone-ash. To this sulphuric acid is added, which converts the cal- cium phosphate in the bones into a salt that is soluble in 190 FIG. 49. Manufacture of Phos- phorus. PHOSPHORUS AND ITS COMPOUNDS 191 water. This is dissolved out and the solution evaporated to dryness, then mixed with carbon and strongly heated. The phosphorus is thus set free ; it distills out and is condensed under water and molded into sticks. (See Fig. 49.) R, -R, are the retorts into' which the mixture of charcoal and phosphorus compounds are put ; F is the furnace, and TP, IF, the water tanks where the phosphorus is condensed. The process is very deleterious to health, the fumes from the retorts often producing dangerous ulcerations of the jawbones, a disease which is practically incurable. (3p Characteristics of Phosphorus. Phosphorus is a very pale, amber-colored, translucent solid, somewhat waxy in appearance. When exposed to the air it almost imme- diately begins to give off luminous fumes having a faint garlic odor, and in the course of a short time takes fire. A little friction will readily ignite it, hence it should be cut under water. Burns from it are very serious and require weeks to heal. If heated to 240 out of contact with the air, it changes to an allotropic form, known as red or amorphous phosphorus. Unlike the ordinary phosphorus, this is not poisonous, does not readily take fire, is not solu- ble in carbon disulphide, and does not glow in the dark. EXPERIMENT 125. To show the ready combustibility of phos- phorus when finely divided. Dissolve a small piece of phosphorus, half as large as a pea, in a little carbon disulphide. Pour the solution upon a piece of filter or blotting paper, and let it dry. Notice how quickly it ignites. EXPERIMENT 126. To show that phosphorus will burn under water. Put into a small bottle about 1 g. of potassium chlorate, add a few small pieces of phosphorus, and cover with water. By means of a pipette or funnel tube introduce beneath the water into contact with the potassium chlorate a little sulphuric acid. Notice that the phos- phorus begins to ^urn. Explain. 192 MODERN CHEMISTRY EXPERIMENT 127. To show the affinity of phosphorus for chlo- rine, bromine, and iodine. Put a small piece of phosphorus into a deflagrating spoon and introduce it into a jar of chlorine. What happens in a few moments? Cut a thin slice of phosphorus, and upon it place a crystal of iodine. Notice that the phosphoras is soon ignited. Try a drop of bromine in the same way. Uses for Phosphorus. About 3000 tons of phos- phorus are manufactured annually, most of which is used in preparing matches. Small quantities also are employed in making poisons. Matches were first made in Austria by tipping small pine sticks with sulphur to which a little phosphorus had been added. This method was employed for a good many years, but the sulphur has now been largely replaced by other substances rich in oxygen, such as potassium chlorate, saltpeter, etc., together with paraffine. 5. Matches are now made entirely by machinery, and with wonderful rapidity. The wood, being sawed into convenient lengths, is pressed against knives, which split it up into the proper size for matches. These are dipped into paraffine, then tipped with a paste made of a little glue containing phosphorus and the other ingredients already mentioned, toge^rkr with some coloring matter. After drying they are packed in boxes. In this way a single machine will make and pack several million matches in a day. In the case of safety matches the phosphorus is placed in the prepared surface upon the box, and the matches can be ignited only by friction on this surface. 6. Compounds of Phosphorus. One of the most inter- esting of these is hydrogen phosphide, PH 3 . It is also called phosphine and phosphoreted hydrogen. It is readily evolved when phosphorus is heated in a solution of any strong alkali, such as caustic soda or potash. PHOSPHORUS AND ITS COMPOUNDS 193 FIG. 50. EXPERIMENT 128. Suitable for class-room. Into a small flask put about 50 cc. of strong caustic soda or potash solution, and add several small pieces of phos- phorus. Pour in about a cubic centimeter of ether, and close the flask quickly with a cork and long delivery tube. Support the flask upon a ring-stand, as- shown in the figure, and heat moderately. Presently smoky-look- ing fumes will fill the flask, and then the bubbles issuing from the mouth of the de- livery tube will take fire spontaneously. If the room is free from draughts of air, beautiful rings of smoke, grow- ing gradually larger, will float upward. Notice the vortex motion of the rings. The ether is introduced to expel the air before any phos- phine is generated ; the heat should be regulated so as not to allow too rapid an evolution of gas, otherwise the rings will follow in such rapid succession as to break one another. What is the odor of the gas? Color? 7. Oxides of Phosphorus. Pentoxide, P 2 5 , and Tri- oxide, P 2 3 . The first of these has been seen on various occasions : when phosphorus was burned in oxygen, in preparing nitrogen, etc. The dense white fumes noticed consisted mainly of phosphorus pent oxide. This com- pound is always obtained when phosphorus is burned in a plentiful supply of oxygen. When the amount is limited, or when the combustion is slow, phosphorus trioxide is obtained. The peiitoxide is a white solid which has great affinity for moisture, and if dropped into water combines with it with a hissing sound as of a hot iron in cold water. 194 MODERN CHEMISTRY 8. Acids of Phosphorus. The two oxides named above, like the corresponding oxides of nitrogen, are the anhy- drides of certain acids, thus : P 2 O 3 + 3 H 2 O = 2 H 3 PO 3 . . . Phosphorous Acid P 2 O 5 + 3 H 2 O == 2 H 3 P0 4 . . . Phosphoric " The latter is the more important. It will be noticed th tit- its molecule contains three atoms of hydrogen, all of which may be replaced by a metal. Such acids are called tribasic. Phosphoric acid is a white crystalline substance, which may be prepared by treating bone-ash with sulphuric acid. At high temperatures it will give up a part of the water that was taken in its formation and yield metaphosphoric acid, HPO 3 , which is monobasic. The reaction may be shown thus : H 8 PO 4 + heat = HPO 3 + H 2 O. This is frequently sold under the name glacial phosphoric acid. 9. Compounds with Phosphoric Acid. Phosphates. The most common of these is calcium phosphate, Ca 3 (PO 4 ) 2 , found in the bones. Immense deposits of this are found in Florida, where it is mined and used as a fertilizer in various parts of the world. From the fact that all grain plants absorb the soluble phosphates from the soils, unless these salts are replaced in some way the land rapidly loses its productive power. A considerable portion of the phos- phates in the grain fed to animals is thrown off in the ex- crement and is returned to the soil in this way. 10. Immense quantities of bones are reduced to animal charcoal, and then, by treatment with sulphuric acid, con- verted into soluble phosphates and returned to the soil PHOSPHORUS AND ITS COMPOUNDS 195 in this way as artificial fertilizers. Another source of considerable supply is from the reduction of phosphorus- bearing iron ores by the Thomas- Gilchrist process ; and 'a matter of interest in this connection is that the calcium phosphate thus obtained is already in the soluble form and needs no further treatment. SUMMARY OF CHAPTER Phosphorus Occurrence. Source of supply. Method of preparing phosphorus. By-products and their uses. Characteristics. Two forms of phosphorus. Compare with forms of sulphur. Experiments to illustrate characteristics. Uses. Method of making matches. Chemicals used. Compounds. With hydrogen. How prepared. With oxygen. Names and formulae. How prepared. W T ith hydrogen and oxygen. How related to the oxides. Salts formed from these acida. Uses. CHAPTER XVI AVOGADRO'S LAW ATOMIC WEIGHTS PROBLEMS 1. Avogadro's Law. This law, or hypothesis, was for- mulated by the Italian physicist and chemist, Avogadro, and afterward, independently, by the Frenchman, Ampere. It may be stated thus : - r^> Equal volumes of all substances in the gaseous condition f _fj under the same pressure and temperature contain the same number of molecules. To illustrate, suppose a liter of hydrochloric acid gas contains a billion molecules, then a liter of nitrous oxide, or of any other gas, would also contain a billion molecules. 2. Proof of this Theory. No absolute proof of this law has ever been given, but many facts seem to favor such a theory. For example, we have seen that all gases expand and contract in the same ratio under the influence of heat and pressure. As expansion and contraction mean simply a change in the distance which separates the mole- cules from each other, this being greater when the body is heated, and less when cooled,. it would seem that bodies could expand alike only if composed of the same number of molecules, or if containing what means the same thing, the same number of intermolecular spaces. 3. Ratio of Molecular Weight to Specific Gravity. It \ has been found also that there is a constant ratio existing between the molecular weight of a gaseous body and its specific gravity ; that is, if we divide the molecular weight of any gas by its specific gravity, we always obtain prac- 196 ATOMIC WEIGHTS 197 tically the same quotient. This ratio is about 28.88. Thus, the molecular weight of carbon dioxide is 44, its specific gravity is 1.524, the ratio of 44 to 1.524 is 28.87 ; carbon monoxide has a molecular weight of 28, specific gravity of 0.967, the ratio is 28.94. EXERCISE. To apply this fact, suppose a given volume of nitrous oxide, X 2 O, weighs 1.52 grams, and the same volume of hydrochloric acid gas weighs 1.27 grams. It is evident that the weight of any volume of gas divided by the weight of one molecule would give the number of molecules in that volume. Thus : wt. of 1 1. x.,o = no mol N o in l uter wt. of 1 mol. nd - wt. of 1 1. HC1 wt. of 1 mol. HC1 We can readily find the weight of a liter of each of these gases, and also the molecular weight of each, but the first is in grams and the second in microcriths, that is, so many times as heavy as a hydrogen atom ; but unfortunately we have no means of knowing how many microcriths in a gram, hence we cannot perform the division indicated above nor assign to the quotient any concrete name. If we make the division, however, we find that the quotient is always practically the same ; that is, wt. of 1 vol. N 2 O wt. of 1 mol. X 2 O ** f * V |- *TC1 = 28.83 = no< moh HC1 in ! voL wt. of 1 mol. HC1 Hence, as the ratio in each case is the same, in accordance with the axioms of geometry, no. mol. N 2 O in 1 vol. = no. mol. HC1 in 1 vol. Putting this law into the form of a proportion, it would read : mw m w i , , = , or mw : m'w' '.'.sis'. s s' in which mw and m'w 1 represent the molecular weights of any two gases, and s and s' their specific gravities. 198 MODERN CHEMISTRY 4. Application of this Law. The truth of Avogadro's Law having been accepted long ago, it is now made use of in determining the molecular weights of new compounds. Having found lay actual work the weight of 1 liter of the gas, and knowing the weight of 1 liter of air, the specific gravity is found. Then, substituting in the formula, ?= 28.88, S we can easily find the value of mw. 5. Finding Atomic Weights. This law is further used in determining the atomic weight of a newly discovered element. Let m represent this element, and suppose we are attempting to find the atomic weight by studying some compound of it with oxygen. We should find the weight of a molecule of the oxide as shown above. Suppose this is found to be 28. Next, by chemical analysis we should determine what per cent of the compound is the new element, m. Suppose the analysis shows this to be 42.86 per cent, then we should have this proportion: mol. wt. : wt. of m in the mol. : : 100 per cent : per cent of m; or, 28 : x : : 100 : 42.86. 100 x = 42.86 x 28. x =12. In the same way we should determine the weight of the element, m, in a molecule of a number of other compounds containing it ; then, the one having the least amount would be taken as a compound con- taining but one atom of the element, and the value of x in that com- pound would represent the atomic weight. To illustrate, suppose in this way we find in our analyses, and subsequent determinations, that moi - x is equal to 24, 12, 36, 120. The second, 12, being the smallest amount found in any compound, would be accepted as the atomic weight. This would not be absolute proof, however, as later another compound might be discovered which contained a smaller amount of m, in which case that smaller amount would be taken as the atomic weight. ATOMIC WEIGHTS 199 6. Constitution of the Molecules of Elements. How many atoms are there in a molecule of an elementary sub- stance, like oxygen, hydrogen, etc. ? In writing some of the reactions in the earlier part of this book the molecules were shown as having two atoms. With some exceptions, this is true ; that is, a molecule of hydrogen, oxygen, chlorine, and of many other elements contains two atoms. How do we know this? A proof in the case of one or two elements will illustrate for the others. We have seen that when hydrogen and chlorine are caused to unite, they form hydrochloric acid. It is found also by further experiment that in uniting thus the volume is not decreased ; that is, if we put together a liter of chlorine and one of hydrogen, after causing them to combine, we have 2 liters of hydrochloric acid. Perhaps it will be clearer, stated in the form of an equation, thus : 1000 cc. of Cl + 1000 cc. of H = 2000 cc. of HC1. Now, according to Avogadro's Law, there would be the same number of molecules in a liter of chlorine as of hydrogen or of hydrochloric acid. Dividing the entire equation through by this common factor, the number of molecules of chlorine in 1 liter, or 1000 cc., we should have 1 mol. of Cl. + 1 mol. of H = 2 mol. of HC1. Two Molecules. Chemical analysis shows that in hydrochloric acid the hydrogen and chlorine are united in the ratio of 1 to 35.5, or one atom of each, as represented by the formula HC1, or by the figure. 200 MODERN CHEMISTRY It is evident, therefore, that two molecules of hydrochloric acid contain two atoms of hydrogen and two of chlorine, and as we only had one molecule of each of these elementary gases, each of those molecules must have contained two atoms. In a similar way we would prove for bromine, fluorine, oxygen, and others. 7. Most Molecules Diatomic. Such molecules as these are called diatomic. There are a few, sodium, potassium, cadmium, mercury, and zinc, whose molecules contain only one atom, and such are called monatomic. Their molecule is, therefore, identical with the atom. Only one triatomic elementary molecule is known, and that is the allotropic form, ozone. A few, like phosphorus and arsenic, are tetratomic ; that is, the molecule is made up of four atoms. 8. Application of this Fact. It often becomes necos- sary in chemical problems to know the weight of a liter of a gas. This may very easily be found, but we must first know its vapor density; that is, its density compared to hydrogen. With the elementary substances this is, as a rule, the same as the atomic weight ; for example, the atomic weight of hydrogen is 1, the molecular weight is 2 ; the atomic weight of nitrogen is 14 ; the molecular weight 28. Hence, whether we take the atomic weight of nitrogen, or its molecular weight and divide by the molec- ular weight of hydrogen, we obtain the same results. Then, as the hydrogen molecule weighs two, we find the vapor density of any other substance by dividing its molecular weight by 2. Thus : 1 mol. N 2 O weighs 2 x 14 + 16 = 44 1"H " 2x1 = 2 1 " N 2 O " 44-7-2 times as much as 1 mol. H ATOMIC WEIGHTS 201 Again, 1 mol. CO weighs 12 4- 16 = 28 1"H 2x1=2 1 " CO " 28 -T- 2 times as much as 1 mol. H Therefore, vapor density ofCOis28-f-2= 14. Thus find the vapor density of CO 2 , N 2 O 3 , O, HC1, SO 2 , C1,N. ^ jt $ !*T ^ \V* 9^ To find Weight of One Liter of Any Gas. Having found the weight of a gas compared to hydrogen (its vapor density), it is only necessary to multiply the weight of 1 liter of hydrogen by this figure. A liter of hydro- -3 gen has been found to weigh .0896 g., a number which '- should be remembered. Suppose now we desire to find the ' , weight of a liter of carbon monoxide, CO. Above we found its vapor density to be 14. Then, as a liter of <* hydrogen weighs .0896 g., one of carbon monoxide will /, weigh 14 x .0896, or 1.2544. J, Thus find the weight of 1 liter of the gases whose densi- ties were found above. Also of N 2 O, NH 3 , H 2 S. rw 10. The Formulae of Compound Bodies. We have learned that the formula of a compound is a short method of expressing its composition. It may be of interest to know how to determine the formula of a compound. The substance is first carefully analyzed, and the percentage composition determined. Suppose \ve have in mind a compound which analysis shows con- sists of carbon and oxygen, 27.27 per cent of the former, and 72.73 per cent of the latter. We should next weigh a liter of it ; suppose we find this to be 1.9712 g. As a liter of hydrogen weighs .0896 g., the unknown gas is 1.9712 H- .0896, or 22 times as heavy. We have seen that the molecular weight is twice the vapor density, then the weight of the molecule would be 2 x 22, or 44. Now, as the 202 MODERN CHEMISTRY carbon is 27.27 per cent of this, it equals .2727 of 44 = 11.9988, and the oxygen, 72.73 per cent, or its weight in the molecule is 72.73 per cent of 44 = 32 + . Previous experiments have shown that the atomic weight of carbon is 12, hence the weight found above, 11.9988, practically corresponds to one atom, and that would be the amount of carbon in the compound. In the same way as the atomic weight of oxygen is known to be 16, the amount found in the compound, 32, would indi- cate two atoms. The substance in question, therefore, would contain carbon, 1 atom, oxygen, 2 atoms, and would be carbon dioxide, for- mula CO 2 . PROBLEMS. 1. A liter of a certain gas weighs 0.8064. It consists of hydrogen and oxygen f . Find its vapor density, molecular weight, r< and the formula. 2. A gas consisting of carbon and oxygen has 42.86 per cent of \5 the former, and 57.14 + per cent of the latter. If 1 liter of it weighs Qj 1.2544, what is its formula? 3. What per cent of turpentine, C 10 H 16 , is carbon? Hydrogen? 4. The vapor density of a body is found to be 50.5. If analysis shows that 2.359 g. of it contain 1.12 g. of oxygen, how many atoms p "' of oxygen are there in the formula representing the substance? 5. What is the molecular weight of a certain substance if 50 g. of it contain 32.65 g. of oxygen, knowing that there are four atoms of oxygen in the molecule of the substance? 6. Find the percentage composition of nitric acid. SUMMARY OF CHAPTER Avogadro's Law Statement of the law. Illustration. Proof of the law. a. As seen in effects of heat. b. Ratio of molecular weight to specific gravity. Value of the law. a. Finding atomic weights Illustration. b. Constitution of molecules Proof. Meaning of terms monatomic, etc. Problems. Method of finding weight of a liter of any gas. How to determine the formula of a compound. CHAPTER XVII THE METALS PERIODIC LAW 1. Metals and Non-metals. It has been customary to divide- the elements into two great classes, the metals and non-metals, of which the former includes by far the greater . number. This classification, however, is based largely upon the external characteristics or appearance rather than upon the chemical deportment. In appearance the metals have a peculiar luster, known as the metallic luster, considerable density, with few exceptions have high melt- ing points, and are electro-positive in character. As A rule, their oxides are not anhydrides, and yet there ^re many exceptions to this statement, for we find various compounds of tin, arsenic, antimony, chromium, aluminum, etc., in which these metals seem to serve as the acid-form- ing element. And some even possess more chemical char- SON-METALS METALS Their oxides, with water, form acids, as for example : S0 2 ... H 2 S0 3 , P 2 O 5 . . . HPO 3 , etc. Many are gaseous. Many are transparent. Poor conductors of heat and electricity. Their oxides, with water, form bases, as: CaO . . . Ca(OH)2, Na 2 . . . NaOIL K 2 O . . . KOH, etc. Most are solids. All are opaque. Good conductors of heat and electricity. 203 204 MODERN CHEMISTRY acteristics in common with the non-metals than with the metals. It must be concluded, therefore, that there is no clearly dividing line between the two classes. Neverthe- less, some distinctions in addition to those mentioned above may be noted. 2. Tabular Classification. It will be seen that the above division is almost purely an arbitrary one. At the present time it is customary to classify the elements into a number of groups in accord with what is known as the periodic law. TABLE OF ELEMENTS I II III IV V VI VII VIII Period I H = l Li = 7 Gl=9 B = ll Coll N = 14 O=16 F=19 " II Na- -tf Mg'^ 1 ! AL' tf Si P 8 Cl Fe, Co, * 5 b yO Ni f i. " III < K fe c Scf Ga 7^ ^ e V As Cri Se Mn Br " {t Rb f /Ag Sr'"" Y : - In %n Cb Mo '- I Ru, Rh,' ] Pd " V-f^ Cs*/ Ba La''' ^Ce * } " VI (B Au Hg Yb Tl Pb Bi W Os, Ir, Pt " {* Th U " V\ 3. Recurrent Characteristics in the Table. If the above table is studied in connection with the atomic weights of the elements, it will be seen that, reading from left to right, THE METALS PERIODIC LAW 205 they are arranged with reference to their weights. Thus, in the first period, we have Li = 7, Gl = 9, B = 11, C = 12, N = 14, O = 16, F = 19 ; in the second, Na = 23, Mg=24, Al=27, Si=28, P = 31, S=32, 01=35.5. 4. In thus arranging them it was noticed that, starting with lithium, not until we reach the eighth element beyond, do we come to another, sodium, similar to lithium in char- acteristics ; and from sodium there are seven more before another is reached similar to this. From these observa- tions the above table was arranged, and though it is far from complete, wonderful results have come from it. We notice that in group VII, we have fluorine, chlorine, bromine, and iodine, four elements that we have found to have very similar properties. We shall hereafter find the same to be true of lithium, sodium, and potassium of the first group ; magnesium, calcium, strontium, and barium of the second, and so on. If we take these vertical columns or groups and compare their atomic weights, we notice some interesting facts. * 2 o | The atomic weight of sodium is exactly halfway between the other two. Ca= 40 Sr = 87 Ba = 136 P = 31 As= 75 Sb = 120 The weight of strontium is practically the mean of the other two. The same is true of the middle element. 206 MODERN CHEMISTRY S = 32' Se = 79 The same is true of the middle element. Te = 125 . 5. If we study the compounds that the elements form, we shall find that those falling in the same group are strikingly similar in their chemical behavior. Thus, lithium, sodium, potassium, rubidium, and caesium in the first group are all univalent and form oxides with the general formula, M 2 O, in which M represents any metal of the group. Furthermore, they form no compounds with hydrogen. If we take the second group, they are all bivalent, forming oxides with the general formula, MO, as MgO, CaO, etc. They form no hydrogen com- pounds. The members of the third group are trivalent, as seen in their oxides, A1 2 O 3 , general formula M 2 O 3 . And so we might go on through the table. 6. Vacancies in the Table. It will be noticed that there are many vacant places, but it is an interesting fact that when the table was first worked out there were many others that have since been filled. And strange to say, from this table the author of the plan not only predicted that these very elements would be found, but even gave in a general way their characteristics, and in accordance therewith suggested names for them. In the same way, it is possible that many of the places now vacant will sometime be filled by elements as yet undiscovered. NOTE. Some teachers may prefer to defer a close study of the Periodic Law until after the completion of the work with metals. SUMMARY OF CHAPTER Classes of the elements. Characteristics of the two classes, Wherein different. Wherein alike. THE ALKALI METALS 207 The Periodic Law. Recurrence of certain characteristics. Relation of atomic weights. Similarity of chemical behavior. Value of the law and table. CHAPTER XVIII THE ALKALI METALS SODIUM: NA = 23 1. Its Discovery. Up to the year 1807 caustic soda and caustic potash had been regarded as elementary sub- stances ; by electrolysis, however, Sir Humphry Davy in 1807 proved both of these substances to be compounds, and hydroxides of the metals sodium and potassium. 2. Where found. Sodium is very widely distributed, traces of it being found everywhere. On account of the strong affinity existing between it and water, it never occurs in the metallic state. Its most abundant com- pound is common salt, NaCl, which constitutes a large per cent of the solid matter found in sea water, salt lakes, and springs ; vast deposits of it, more or less pure, occur in many parts of the West as well as in other portions of the world. Sodium nitrate, NaNO 3 , is found in immense quantities in Chile and elsewhere. Other com- pounds occur in smaller proportions, but in some form or other sodium can be detected in the particles of dust that may be seen floating in the sunbeams. 3. Reduction of Sodium from its Compounds. Since the isolation of the metal by Davy, various other plans have been tried, but they are all modifications of the 208 MODERN CHEMISTRY original. What is known as the Castner process is the one generally used at present. See Figure 51. A in the figure is a large iron vessel, B another, similar but smaller, inverted over A and dipping into the fused caustic soda in the lower vessel. It is held in position by insulated supports not shown. Through the bottom at D is inserted the negative electrode, and B serves as the positive. When the cur- rent from the dynamo is passed through, the caustic soda is electrolyzed, B gradually fills with hydrogen which bubbles out underneath, while the metallic sodium col- lects upon the surface of the fused mass. In this way it is prepared for about two dollars a pound. 4. Characteristics of Sodium. It is a silvery white metal, so soft at ordinary temperatures that it may be molded with the fingers, about like stiff putty. At 20 C., however, it becomes hard. It tarnishes so rapidly in the air that only for an instant after being cut can its true color be seen. It takes up moisture and carbon di- oxide from the air, forming first caustic soda, and after- ward sodium carbonate. In course of time a piece of sodium left more or less exposed is entirely converted into amorphous sodium carbonate. It is usually preserved in naphtha or some similar light oil containing no oxygen. 5. Sodium is soluble in liquid ammonia and forms with it a blue solution. Its properties are strongly alkaline. If heated and plunged into a jar of chlorine it burns vigor- ously, forming common salt. Thrown upon water it is immediately melted, owing to the heat generated by the . THE ALKALI METALS 209 strong chemical action, and the water is decomposed, as already shown in our study of hydrogen. If a burning match is touched to. the sodium as it spins about on the water, the hydrogen will burn with a yellow flame, due to the vaporization of a small portion of the sodium. Upon moderately warm water the gas will take fire spontane- ously. If a small piece of sodium is dropped upon a moistened blotting paper, it is quickly ignited. If, when it begins to burn, the molten sodium is allowed to roll off and drop upon the floor, it will burst into many particles which will spin about, burning with the characteristic yel- low flame. EXPERIMENT 129. Moisten a piece of blotting paper with water, to which a little phenolphthalein has been added. Drop a small piece of sodium upon the blotter. Notice the red track it leaves as it slowly moves about from place to place. You have seen similar results in previous work. Let the molten globule of sodium roll off upon the floor and notice what happens. Compounds of Sodium 6. Caustic Soda, Sodium Hydroxide, NaOH. This com- pound is prepared by treating sodium carbonate, Na 2 CO 3 , in solution with lime-water. The reaction is Na 2 C0 3 + Ca(OH) 2 = 2 NaOH + CaCO 3 . The calcium carbonate, thus formed, is insoluble in water, hence is precipitated. The sodium hydroxide is drawn off, evaporated to dry ness, purified, then fused and molded into sticks ; in this form it is put upon the market. It is a white solid, deliquescent, with strongly alkaline properties. 7. Sodium Chloride, NaCl. As already stated, this compound occurs very abundantly. In some places it is mined much as rock or metallic ores are mined. In other 210 MODERN CHEMIST RT places, where the deposits are upon the surface, mingled with considerable quantities of sand and earthy matters, it is dissolved out and the strong solution evaporated. In some of our states wells are sunk into the deposits, and water pumped in to dissolve the salt. This is again drawn out and evaporated. In some places along the Mediter- ranean the sea water is pumped up and allowed to trickle down over brush or lattice work, whereby it is much con- centrated in strength, then this solution is evaporated to dry- ness in large shallow pans. It crystallizes in cubes, as may be seen if a strong solution is allowed to evaporate slowly. Sodium chloride, if chemically pure, is not deliquescent, but owing to impurities present that which is generally put upon the market soon becomes damp when exposed to the air. It is used very extensively in the manufacture of other important compounds of sodium ; also largely in our' food. A part of it is said to be decomposed by the diges- tive fluids of the stomach and to form hydrochloric acid. 8. Sodium Carbonate, Na 2 C0 3 . This is a very impor- tant compound used in the manufacture of soap, glass, and for a variety of other purposes. In the early part of the last century soda crystals, as this compound is often known in commerce, sold for over $300 a ton, while now the same quantity is worth scarcely 20. There are two general processes of manufacture. The simplest and the one most in favor at the present time is the Solvay Process. This consists of passing a current of ammonia into a strong solution of sodium chloride until it is saturated ; carbon dioxide is next forced in and with the ammonia forms ammonium bicarbonate. This reacts with the common salt, forming sodium bicarbonate. These processes may be shown thus : NH 4 OH + C0 2 = NH 4 HC0 3 , NaCl + NH 4 HC0 3 = NaHCO 3 + NH 4 C1. THE ALKALI METALS 211 The sodium bicarbonate crystallizes out much more quickly than the ammonium chloride, and in this way the two compounds are separated. The bicarbonate of soda is then heated to expel a portion of the carbon dioxide, and sodium carbonate results, thus: 2 NaHCOg + heat = Na 2 CO 3 + CO 2 + H 2 O. This process is very cheap because the salt can be had for a few cents per hundred pounds, the ammonia is obtained abundantly from all gas factories, and the carbon dioxide can be had by calcining limestone in making lime. Or, as seen by the last reaction above, the carbon dioxide driven off from the bicarbonate of soda may be utilized for this purpose, and from the ammonium chloride obtained in the second step ammonia may be evolved by treating it with lime. It will be seen, therefore, that the result of one part of the process may serve in another part and thus reduce the final cost of manufacture. The Leblanc Process. This is more complicated than Solvay's, and more expensive ; hence, were it not for the value of some by-products which are obtained, it would no longer be used. It really consists of three steps. First, common salt is treated with sulphuric acid and heated, at first moderately and then more strongly. In the beginning the salt is converted into acid sodium sulphate, thus : NaCl + H 2 SO 4 = XaHSO 4 + HC1. Xext, this acid salt reacts with another part of sodium chloride, form- ing normal sodium sulphate, thus : NaCl + NaHS0 4 = Na 2 SO 4 + HC1; or, putting the two together, we have 2 NaCl + H 2 SO 4 = NajSO 4 + 2 HC1. The sodium sulphate thus obtained is called salt cake. The hydro- chloric acid vapors are passed into flues, down which water constantly trickles and absorbs the acid. This is a valuable by-product, and serves in some places to keep alive the Leblanc industry. Second, this salt cake is mixed with powdered coal and limestone, and heated, when sodium carbonate, mixed with several other sub- stances, is obtained. The mixture is black in color and is known as black ash. The reaction shows the chemical changes that take place : 4 -f CaCO 3 + 2 C = Na 2 CO 3 + CaS + 2 CO 2 . 212 MODERN CHEMISTRY This black ash is treated with water to dissolve out the sodium car- bonate, and the solution is concentrated and purified by calcining, dissolving, and recrystallizing. In connection with almost every Leblanc factory is also one for the manufacture of bleaching powder on a large scale, by using the hydro- chloric acid obtained as a by-product, with native manganese dioxide. 9. Sodium Nitrate, NaN0 3 . This is known as Chile saltpeter on account of the locality from which it is ob- tained and its close resemblance to potassium nitrate. The crude salt found native 'in Chile is dissolved in water and concentrated, whereupon the pure crystals separate. From the fact that it absorbs moisture from the air, it cannot be used in making gunpowder to any great extent. It is used largely, however, in the manu- facture of nitric acid and also for artificial fertilizers. 10. Sodium Sulphate, Na 2 S0 4 . This is frequently called Glauber's salt. It is a white crystalline salt obtained in the preparation of sodium carbonate as described above. 11. Sodium Bicarbonate, NaHC0 3 . This is common cooking soda, and is usually prepared by the Solvay process of making soda crystals, hence is very inexpen- sive. In making bread the " soda " is put with some such acid as sour milk. The acetic or lactic acid, or whatever it may be, reacts with the soda, setting free carbon dioxide, which raises the dough by struggling to escape through it. At the same time the acid disappears in the formation of a neutral salt. This may be seen by the following reaction of " soda " with acetic acid : NaHC0 3 + HC 2 H 3 2 = NaC 2 H 3 O 2 + H 2 O + CO 2 . "N, 12. Soap. This is a substance which has been made in greater or less quantities for probably two thousand years. At first, however, it was used simply as an ointment in a THE ALKALI METALS 213 medicinal way, and not till about 200 A.D. was it applied as it is to-day, and even then only to a limited extent. Soap is made by combining some alkali, as caustic soda or potash, with some fatty substance or oil. The fat con- tains an acid which combines with the alkali, hence we see that soap is really a salt. It retains some alkaline proper- ties, however, just as many other salts do, simply because when dissolved in water it is hydrolyzed, that is, decom- posed by the water, forming caustic soda or potash. Herein lies the chemical value of the soap. It has been said that soda crystals, Na 2 CO 3 , are used in making soap. They must first be converted into caustic soda, however, and this is done by treating the solution with milk of lime, Ca(OH) 2 , as described. 13. Hard and Soft Soap. We have two kinds of soap, hard and soft; the former is made from sodium com- pounds, the latter from potassium. Wood ashes contain considerable quantities of potassium carbonate ; formerly, these were saved by farmers, placed in large " hoppers," lime added, and then leached. A dark-colored, strongly alkaline solution filtered out, containing a considerable per- centage of caustic potash. This was treated with waste fat, and boiled, when in the course of a few hours a strongly alkaline soft soap was obtained, which always remained pasty. By adding common salt to this, it could be converted into a dark-colored solid mass ; for many years this was the only hard soap known. Sodium com- pounds yield hard soap directly on combination with fats, hence they are most used at the present time in the manufacture of ordinary hard soap. 14. The practical value of soap lies in the fact that on account of its slightly alkaline properties it has the power of uniting with the oil secreted by the glands of the skin, 214 MODERN CHEMISTRY and which holds the particles of foreign matter ; this "dirt," therefore, may be removed by the mechanical action of the water. This also explains why frequent bathing with the application of strong soap will tend to cause the skin to chap, by the removal of the oil which keeps it soft and pliable. 15. Test for Sodium. Sodium may always be detected by what is known as the flame test. EXPERIMENT 130. Heat a platinum wire in the Bunsen flame until it no longer imparts any color to the flame. Then dip it into the sodium solution and again hold in the flame. The bright yellow color is distinctive. POTASSIUM : K = 39 16. Where found. Because of its great affinity for other substances, potassium never occurs free. It is very widely distributed, however, in the form of compounds ; it is a constituent of many rocks, and by their decomposi- tion becomes a part of various soils. Being stored up by plants it enters into the animal economy, and by some ani- mals, especially sheep, it is exuded from the skin and col- lects in considerable quantities upon the wool in an oily substance called suint. 17. How obtained in Metallic Form. Potassium, like sodium, may be obtained by electrolysis, but is usually reduced by treating caustic potash with charcoal. The reaction shows the chemical changes : 6 KOH + 2 C = 2 K 2 CO 3 + 3 H 2 + K a . The potassium distills out and is collected under oil. 18. Characteristics of Potassium. It is a metal some- what softer than sodium ; has a bright luster and white color, but it tarnishes instantly when cut in the air, so great is its affinity for oxygen and moisture. At zero it THE ALKALI METALS 215 becomes crystalline in structure, hard, and brittle. When thrown upon water it immediately begins to decompose the water, and with such energy that it is melted and the hydrogen given off is ignited, burning with a violet color. This is due to the vaporization of a small portion of the potassium. As hydrogen is set free from the water caustic potash is formed, according to a reaction previously seen : - H,O + K = KOH + H. With the halogens, chlorine and bromine, potassium ignites spontaneously, and with liquid ammonia it forms a blue solution. It possesses all the strong alkaline charac- teristics of sodium in a degree even more marked. 19. In the metallic form potassium has no uses in the arts. Its compounds, however, are very valuable. Any potassium salt may be tested in the same way as are the sodium compounds, with a platinum wire. The violet flame is characteristic. If both sodium and potassium are present it will be necessary to observe the flame through a blue glass. This transmits the potassium rays, but absorbs those of the sodium. Compounds of Potassium 20. Potassium Hydroxide or Hydrate, KOH. In earlier days the most common source of potassium compounds was wood ashes, which were boiled with water in iron pots. The potassium salts were dissolved out in this man- ner, and from them was prepared caustic potash, KOH. This is now obtained by a method similar to that used in the preparation of caustic soda, viz. by treating potassium carbonate with milk of lime, Ca(OH) 2 , when this reaction takes place : K 2 CO 3 + Ca(OH) 2 = 2 KOH + CaCO 3 . 216 MODERN CHEMISTRY The latter compound is insoluble and is precipitated ; the former is drawn off, concentrated, purified by redissolving in alcohol, again dried, fused and molded in the familiar round sticks. It is very deliquescent, and quickly dis- solves in the moisture it obtains from the air. It is used largely as a reagent in the laboratory. 21. Potassium Carbonate, K 2 C0 3 . As already indicated, this was at one time obtained almost exclusivjely from the ashes of wood. These were treated with water, by which the potassium carbonate was dissolved out ; the solution was boiled dry, forming a white salt known as pearl ash. Now large quantities are obtained by washing sheep's wool in hot water, then drawing off the greasy products obtained and heating them very strongly to expel the oil. The potash salts remain and are dissolved out by water. Another source of considerable quantities is the beet- sugar industry. The beet sap is boiled down to a sirup, and from this sirup is extracted the sugar, leaving a sort of molasses, in which still remain the potassium compounds that the beets had obtained from the soil. This is gen- erally first fermented and distilled; the residue is boiled to dry ness and calcined. Then from the ashes the potash salts are obtained by lixiviation. Potassium carbonate may be prepared from the chloride by the Leblanc process. 22. Potassium Chlorate, KC10 3 . This is a white, crys- talline salt, often sold under the misleading name potash. It has a not unpleasant, cooling taste, and is used some- what for throat affections. In the laboratory it has num- berless applications, many of which are familiar to the student. In the arts it is used in making matches, for fireworks, etc. It is prepared by passing a current of chlorine into a solution of caustic potash, by which both THE ALKALI METALS 217 potassium chloride and potassium chlorate are formed. The former is much more soluble, hence in concentrating the solution the potassium chlorate will crystallize out first, leaving the chloride still in solution. 23. Potassium Nitrate, KN0 3 . This is commonly known as saltpeter. It is a white, crystalline salt, found native in various parts of the world. As we have seen, it is pro- duced by the decomposition of organic matter, especially the refuse from stables. This decomposition is supposed to be brought about by the presence of certain bacteria, and in some countries the process is now carried on artifi- cially to a considerable extent. 24. The refuse is mixed with ashes and lime, and fre- quently stirred to increase the rapidity of decomposition. After a time the whole is leached with water to dissolve out the nitrate. The solution thus obtained is concentrated and the salt allowed to crystallize. Considerable quantities are now made by treating sodium nitrate, which occurs in almost inexhaustible quantities in Chile, with potassium chloride, whereby this double reac- tion takes place : KC1 + NaN0 3 = KN0 3 + NaCl. Potassium nitrate is used extensively in making gun- powder. 25. Potassium Iodide, KI. This is a white crystalline salt. It is used frequently in the laboratory as a reagent, and to some extent in medicine. 26. Potassium Bromide, KBr. This is a white salt, very similar in- general appearance to the iodide. It is used frequently in medicine as a sedative. EXPERIMENT 131. Take any potassium solution and make, the flame test just as you did for sodium in Experiment 130. Notice 218 MODERN CHEMISTRY color of the flame. Now mix with it a solution of some sodium com- pound, and again test. Can you see the potassium flame? Next observe the flame through a sheet of blue glass. State results. COMPARATIVE REVIEW OF THE ALKALI METALS Sodium and Potassium. As found in nature Two important native compounds of each Where found. Which the more important. Comparison of the two metals. Color. Tendency to oxidize. Hardness. Affinity for water. Melting point. Affinity for the halogens. Experiments that illustrate most of these properties. Proof that hydrogen is set free from water by these metals. Proof of the hydroxide formed. Compounds. The Hydroxides Method of preparing Reactions. Usual form Appearance Properties Uses. The Carbonates Source of supply. Former method of obtaining K 2 CO 8 . Present sources. Two methods of preparing Na 2 CO 3 . Uses of the carbonates. Review work in glass. Kinds of glass Differences. Soap making Chemistry of. Kinds of soap. Common salt. Preparation for market. Cooking soda Chemical name and formula. Chemistry of in bread making. Saltpeter Chemical name and formula,. Preparation Appearance Uses. Potassium chlorate Formula. Appearance Uses, CHAPTER XIX THE ALKALINE EARTHS MAGNESIUM : Mg = 24 1. Occurrence. Magnesium in the form of certain com- pounds is widely distributed. Among the most important of its compounds may be named the familiar minerals, asbestos and meerschaum. The first is a silicate of magne- sium and aluminum, and the second a silicate of magnesium. Magnesium limestone, or dolomite, CaMg(CO 3 ) 2 , occurs in considerable quantities. 2. Peculiarities of the Metal. Magnesium is silvery white in color, and melts at a red heat. In dry air it does not tarnish, but moisture quickly affects it. While at ordinary temperatures it is slightly brittle, as it nears the melting point it becomes malleable and may be drawn into wires. These, flattened into ribbons, are the usual com- mercial form, though the powder is also frequently seen. The metal is easily ignited and burns with a dazzling white light, rich in actinic properties. This combustion is so vigorous that it will decompose even carbon dioxide and certain other similar oxides. (See carbon dioxide, page 143.) 3. Uses. On account of the light furnished by burn- ing magnesium, it is frequently used in taking flash-light pictures of caverns, and other interior views. It is like- wise used to a limited extent in making fireworks. In the form of a powder it is often used like zinc in the 219 220 MODERN CHEMISTRY reduction of ferric to ferrous salts (see page 307), on account of the rapidity with which, in the presence of sul- phuric or hydrochloric acid, it yields hydrogen. It is also used by chemists in cases of supposed arsenic poison- ing, in making Marsh's test. Zinc nearly always contains traces of arsenic, whereas magnesium is obtained prac- tically pure ; for this reason it is substituted for the zinc. Compounds of Magnesium 4. Magnesium Sulphate, MgS0 4 . One of the most common compounds is epsom salts, magnesium sulphate, MgSO 4 , 7 H 2 O. This is a salt found in the water of many mineral springs. It has a very bitter taste and is used largely in medicine, also extensively in finishing cotton goods. 5. Magnesia, Magnesium Oxide, MgO. This is a white solid obtained when magnesium is burned in the air or in oxygen. It is often prepared by heating magnesium car- bonate to expel the carbon dioxide, just as lime is pre- pared from limestone (see lime, page 221). The reaction is seen below : MgCO 3 + heat = MgO + CO 2 . It is used as a face powder, and, because of its high melt- ing point, sometimes for making or lining crucibles. CALCIUM: Ca = 40 6. Occurrence. In the form of compounds, calcium is one of the most abundant and most widely distributed elements known. Because of its strong affinity for water, however, it never occurs free. The carbonate of calcium, CaCOg, is the most abundant form and includes many well- THE ALKALINE EARTHS 221 known substances, such as marble, limestone, and chalk. Some of the more highly crystallized forms are Iceland spar, calcite, and dog-tooth spar, while stalactites, corals, and shells have the same composition. The next most abundant natural compound of calcium is gypsum, calcium sulphate, CaSO 4 , 2 H 2 O. 7. Production of the Metal. Calcium has seldom been prepared, and then only for the purpose of studying its properties. Sir Humphry Davy, who first isolated potas- sium and sodium from their hydroxides by means of an electric current, in the same way decomposed calcium chloride and obtained calcium in the metallic form. 8. Characteristics. Calcium is of a brassy yellow color, and somewhat malleable and ductile. It has a density of about 1.6, and like sodium readily decomposes water, forming the hydroxide, Ca(OH) 2 . It is readily soluble in dilute acids, and at a temperature a little above its melt- ing point it burns with a reddish yellow light. The cost of its production is too great to admit of any practical use. Compounds of Calcium 9. Although as a metal calcium is of so little value, it would be difficult to estimate the worth of the compounds. 10. Lime, Calcium Oxide, CaO. This is one of the most important compounds known. It is easily prepared from limestone by heating it to a red heat, at which tem- perature carbon dioxide is expelled, thus: - CaCO 3 + heat = CaO + CO 2 . Lime is prepared in kilns, which are simply square rooms or ovens 15 to 20 feet high, and 10 to 15 feet each way. See Fig. 52. The limestone is thrown in from above and 222 MODERN CHEMISTRY strongly heated with dry cordwood or coke in alter- nate layers. In a few days the limestone is converted into lime, then the fire is removed, the mass is allowed to cool and the lime with- drawn, and if intended for shipment packed in barrels. Some kilns are arranged below so as to enable the workmen to remove the lime without putting out the fire. Such are contin- te uous ly ^d from above, and ' the operation goes on with- FlG - 52 ' out ceasing. 11. Properties of Lime. Prepared as above it is in the form of white lumps, but if left exposed to the air it begins at once to take up moisture and in a short time crumbles to a fine powder. It is then said to be "air- slaked," although it is really the water in the air that has caused the change. The reaction is as follows : CaO + H 2 O = Ca(OH) 2 . 12. If a lump of freshly prepared lime be treated with water, the change indicated above takes place rapidly, accompanied by the evolution of considerable heat. The hydroxide, Ca(OH) 2 , thus obtained is soluble in water, though very much less so than ordinary caustic soda or potash. The solution of caustic lime is known as lime- water. 13. Uses of Lime. Lime is indispensable in the erec- tion of almost all structures. Mixed with sand it forms the mortar for nearly all stone and brick work except THE ALKALINE EARTHS 223 such as is laid under water and much of the plaster for indoor work. Unmixed with sand it is frequently used to give the white or finishing coat in plastering, though various plasters are now beginning to take the place of ordinary lime in this respect. 14. It is also used extensively in the lime purifiers of illuminating gas works, in the manufacture of bleaching powder, of ammonia, in removing the hair from hides in the process of tanning, and for numerous other purposes where a cheap and easily prepared alkali is demanded. 15. Calcium hydroxide, exposed to the air, absorbs car- bon dioxide and forms calcium carbonate, thus : Ca(OH) 2 + C0 2 = CaC0 3 + H 2 O. The same reaction takes place in mortar, hence that which has been properly prepared will grow gradually harder, in time being converted back again into a siliceous lime- stone. If a beaker containing lime-water be left exposed to the air, in a little while a white film will be seen to cover the surface of the liquid. This is really a pre- cipitate of calcium carbonate, resulting from the absorp- tion of the carbon dioxide of the air by the lime-water. If the breath from the lungs be blown through a clear solution of lime-water, it quickly becomes clouded from the same cause. 16. Calcium Carbonate, CaC0 3 . In the natural form this is known in the several varieties mentioned above. Artificially, it may be prepared as a white precipitate by adding some alkaline carbonate, as sodium or ammonium carbonate, to a calcium chloride solution. The following reaction takes place : CaCl 2 + (NH 4 ) 2 CO 3 = 2 NH 4 C1 + CaCO 3 . 224 MODERN CHEMISTRY It is insoluble in pure water, but when an excess of carbon dioxide is present, it slowly dissolves. EXPERIMENT 132. Through a few cubic centimeters of lime-water in a flask or beaker, pass a current of carbon dioxide, or blow the breath for some time. What finally becomes of the white precipitate which forms at first? Preserve the water. In this way water charged with carbon dioxide percolating through limestone rocks gradually dissolves them, and has formed many of the great caves known in this country. This same water, dripping from the roof of caverns, being no longer under pressure, gives up its carbon dioxide, and the calcium carbonate, no longer held in solution by the gas, is deposited in the form of stalactites and stalagmites. 17. Calcium Chloride, CaCl 2 . This is a white salt which may be prepared from any form of the carbonate by treat- ing with hydrochloric acid. It is a by-product formed in the preparation of carbon dioxide from limestone : CaC0 3 + 2 HC1 = CaCl 2 + CO 2 + H 2 O. It is strongly deliquescent, and is often used in drying gases, damp cellars, etc. 18. Calcium Sulphate, CaS0 4 , 2 H 2 0. In the natural form this is the gypsum already mentioned. It occurs in vast quantities in many of our states, notably Kansas, New York, Illinois, etc., both in the form of rich, heavy deposits, and mixed with various impurities upon the surface. It is used extensively in making plaster of Paris. This is manufactured simply by strongly calcining the powdered gypsum till half the water of crystallization is expelled. During this time, as the water escapes from the powdered mass, the whole seems to boil vigorously. After two or three hours the process is complete, and the plaster is ready to be mixed with the " retarder," if neces- sary. This plaster has the property of " setting " or hard- THE ALKALINE EARTHS 225 ening quickly when water is added to it. This is due to tlie fact that the anhydrous salt again takes up the water of crystallization expelled in the previous calcination. If the -plaster which has been used once be again calcined, it acquires again its property of "setting." 19. Uses of Plaster of Paris. It is employed extensively in making molds for many of the finer castings, in dental work and surgery, for statuettes, as a finishing coat in plastering, and for stucco and other ornamental work on the interior of buildings. For most purposes, a plaster that does not harden so rapidly is desirable, hence it is customary to mix with it some kind of clay, or other substance, which causes it to "set" more slowly. This clay has already been spoken of as the "retarder." 20. Cements. Cements are a species of lime which have the power of hardening or setting rapidly, like plaster of Paris. They are prepared by calcining lime- stone, which contains a large percentage of silica and alumina, SiO 2 and A1 2 O 3 . Dolomitic or magnesium lime- stones, containing also the silica and alumina, when cal- cined, produce a cement that will harden under water, known as hydraulic cement. It has been stated that ordinary plaster hardens by the absorption of carbon diox- ide from the air, forming again calcium carbonate. This is, necessarily, a slow process. Cements, as already stated, are produced by driving out the water of crystallization ; hence, when they are mixed with water for use, they very rapidly take this up again, forming practically the original rock. Hydraulic cement is used in laying the piers of bridges, building jetties, and other work that is to be under water. Ordinary cements are used extensively for laying pavements, building roadbeds, for the concrete foundation for various kinds of masonry, etc. The fol- 226 MODERN CHEMISTRY lowing shows the composition of some cement rocks from various localities : LOCALITY CaC0 3 MgC0 3 Si0 2 Fe 2 3 A1,0, H 2 UNDETER- MINED Rosendale, N.Y. 45.91 26.14 15.37 11.38 1.20 Utica, 111. 42.25 31.98 21.12 1.12 1.07 2.46 Milwaukee, Wis. 45.54 32.46 17.56 3.03 1.41 Cement, Ga. 43.50 22.00 22.10 1.80 5.45 4.95 Siegfried, Pa. 78.90 2.66 11.62 6.25 0.55 Ft. Scott, Kan. 73.95 2.26 18.75 2.32 2.15 0.37 0.20 Ft. Scott, Kan., No. 2 65.21 10.65 15.21 4.56 4.37 21. Hard Water. Hardness in water is due to the presence of certain salts in solution, very commonly some compounds of calcium. This hardness may be either tem- porary or permanent, according as it may be removed by boiling or by adding ammonia, or not at all. EXPERIMENT 133. Prepare a soap solution by dissolving a shaving of soap in warm distilled water. Allow it to stand a few minutes. It should be perfectly clear. To a few cubic centimeters of the lime- water, through which the breath was blown till clear again, add a little of the soap solution. What happens? Why? Take another portion of the clear lirne-water and boil it for a few minutes. Has any sediment formed in the flask ? The heat has expelled the carbon dioxide ; why does the precipitate form ? Decant a portion of it and test with the soap solution : is the water still " hard " ? What effect has the boiling had ? To another portion of the same hard water (which has not been boiled) add a few drops of ammonia and again test to see whether the water is still hard. What are the results? Add a little powdered calcium sulphate, CaSO 4 , to some water, and after some time test a portion of it to learn whether it is hard. Now try to remove the hardness by the methods previously used. State results. THE ALKALINE EAETHS 227 22. Water the hardness of which may be easily re- moved is said to be temporarily hard, while that which cannot be so changed is permanently hard. When the hands are washed with soap in hard water, the soap pre- cipitates the salts in the water, of which a portion settles upon the skin, giving it an unpleasant feeling. Another part of the precipitate is usually seen as a scum upon the surface of the water. 23. Bleaching Powder, Ca(C10) 2 + CaCl 2 . This is also called hypochlorite of lime. It is a white powder which is prepared by passing chlorine into chambers containing common lime spread loosely upon shelves. The reaction may be represented thus : 2 CaO + 2 C1 2 = Ca(ClO) 2 + CaCl 2 . When treated with any dilute acid, chlorine is again set free ; for this reason the compound is used extensively as a source of chlorine in bleaching muslin and other cotton goods.* 24. From the fact that chlorine does not bleach dry cloth, it is believed to be not the direct bleaching agent, but simply that which sets free another. It will be seen later, in studying the compounds of manganese, that log- wood, litmus, and other colored vegetable solutions are rapidly bleached by the use of potassium permanganate, in the presence of some acid. Experiment shows that this is due to the oxygen that is set free from the per- manganate. Similarly the chlorine, which has most won- derful affinity for hydrogen (see page 108), sets free the oxygen from the water with which the cloth is moistened, and this in the nascent state oxidizes the coloring matter and converts it into colorless compounds. * See wo A under Chlorine, page 111. 228 MODERN CHEMISTRY 25. When a current of carbon dioxide is passed through a solution of bleaching powder, chlorine is liberated, and can be detected by the odor, just as when treated as above with a dilute acid. Exposed to the air, bleaching powder yields up its chlorine, owing to the action of the carbon dioxide always present ; but naturally the process is very slow. On account of this fact, and because chlorine is an excellent germicide and disinfectant, bleaching powder is used frequently in sick rooms and hospital wards. The generation of the chlorine is so slow as to be scarcely noticeable, and yet sufficient to keep the atmosphere in a wholesome condition. STRONTIUM : Sr = 87 26. Its Name. Strontium is a rare metal, which re- ceived its name from Strontian, a place in Scotland, where it was discovered. One of its chief sources is the mineral strontianite, SrCO 3 . Compounds of Strontium 27. Strontium Nitrate, Sr(N0 3 ) 2 . This is a white crys- talline salt, soluble in water. It is used considerably in fireworks and in making " red fire." EXPERIMENT 134. Mix thoroughly about a gram each of stron- tium nitrate and potassium chlorate, finely pulverized, and about as much in bulk of powdered shellac. Place the mixture in an iron saucer and ignite with a match. State the results. 28. Strontium Carbonate, SrC0 3 . This is a white pre- cipitate, like calcium carbonate, obtained when ammonium or sodium carbonate is added to a neutral or alkaline solution of a strontium salt. Sr(N0 3 ) 2 + (NH 4 ) 2 C0 3 = SrCO 3 + 2 NH 4 NO 3 . THE ALKALINE EARTHS 229 29. Strontium Hydroxide, Sr(OH) 2 . When water is added to strontium oxide, SrO, like lime, it is slaked, evolves much heat, and is converted into the hydroxide, Sr(OH) 2 . In this form it is used considerably in the manufacture and refining of beet sugar. BARIUM : Ba = 137 30. Its Name. This metal, also rare, received its name from a Greek word, meaning heavy, and was so called be- cause its chief natural ore, barite, BaSO 4 , has great den- sity. It is also found as a carbonate, BaCO 3 , known as witherite. Compounds of Barium 31. Barium Chloride, BaCl 2 . This is a white crystal- line salt, readily soluble in water. It is used in the laboratory in testing for sulphuric acid. 32. Barium Sulphate, BaS0 4 . This is a heavy white precipitate, insoluble in water and acids. It is easily pre- pared by adding sulphuric acid or any soluble sulphate to a solution of barium chloride. It is used considerably as an adulterant for white lead (see page 280), and to some extent in weighting paper. 33. Barium Carbonate, BaC0 3 . This is a white precipi- tate formed when ammonium or sodium carbonate is added to a neutral or alkaline solution of a barium salt. It is insoluble in water, but soluble in weak acids. EXPERIMENT 135. Let the student prepare both of the above compounds, using barium chloride for the barium solution. Note the differences between the two and test their solubility in the common acids. State results. 34. Barium Nitrate, Ba(N0 3 ) 2 . This is a white crystal- line salt. It is used to a considerable extent in the making of green fire for fireworks. 230 MODERN CHEMISTRY EXPERIMENT 136. Repeat Experiment 134, substituting barium nitrate for the strontium nitrate, and state results. Sulphur or pow- dered charcoal may be used instead of the shellac, but the sulphur yields very irritating fumes of the dioxide, and the charcoal does not burn so readily. 35. Barium Hydroxide, Ba(OH) 2 . This is a compound obtained from barium oxide, BaO, by the addition of water, just as slaked lime is prepared. Like calcium hydroxide, it forms a precipitate of the carbonate upon the addition of carbon dioxide. It was formerly used extensively in clarifying beet sugar, but as it is very poisonous, and traces of it sometimes remain in the sugar, its use has been supplanted by that of strontium hydroxide. 36. Flame Tests. The metals of this group, calcium, strontium, and barium, may be detected by the flame test. EXPERIMENT 137. Just as you tried sodium and potassium, now take some solutions of these three metals and make the flame test in the same way. State results as to color and duration of flame. REVIEW OF WORK IN ALKALINE EARTH METALS Magnesium, Calcium, Strontium, Barium. 1. Occurrence Compare native compounds. Crystallized forms of calcium compounds. Uncrystallized forms. Special forms. 2. Artificial compounds. a. The Oxides of Mg, Ca, Sr, Ba. Wherein is their preparation similar? Why are such compounds used ? Importance of CaO and MgO. b. The Hydroxides Similarity of preparation. Uses of Ca(OH) 2 and Sr(OH) 2 . Preparation of mortar; chemical change it under- goes as it hardens. Hydraulic cement ; other cements ; uses ; explanation. THE ALKALINE EARTHS 231 c. The Nitrates Use in the arts of Sr(N0 3 ) 2 ; Ba(N0 3 )2. How used. Chemical action of each constituent. d. The Sulphates Two important ones. Preparation of Plaster of Paris Compare with prepa- ration of CaO. Uses of CaSO 4 and BaSO 4 . Chemical change which takes place when Plaster of Paris hardens. Compare with hardening of mortar. e. Hard Waters Due to what compounds. Two classes, how different. Methods of softening water. Chemistry of these methods. / Some special calcium compounds. CaF 2 Use, and method of using. Bleaching powder Uses Compare Cl and SO 2 as bleaching agents Chemical action of each. Use of bleaching powder as a disinfectant How is chlorine set free ? 3. Flame tests Method of making test. Comparison of colors imparted. 4. Comparative value of the metals in metallic form. 5. Exercise Given some marble, HC1, H 2 SO 4 , H 2 O, and Na 2 CO s . Tell how to prepare CaO, Ca(OH) 2 , CaSO 4 , CaCl 2 , CaCO 3 (amorphous). Write all reactions concerned. CHAPTER XX THE COPPER-SILVER GROUP COPPER, SILVER, GOLD COPPER: Cu = 63 1. History. Copper has been known from earliest antiquity, its use being mentioned by Jewish, Assyrian, and other ancient historians. By the Greeks it was ob- tained from the island of Cyprus, and from this fact probably received the name kuprum, and its present symbol, Cu. In England copper-mining was begun before the close of the twelfth century. It met with little success, however, till about five hundred years later. In the United States, the oldest mines are those of the Lake Superior region. The remains of prehistoric tribes about the mines indicate clearly that these deposits were known and used in very early times. The metal was obtained by stripping the rock and earth from the outcropping strata. When the rock had been broken or cracked off, the thin sheets of copper were removed and hammered into vessels of various shapes. 2. Sources of Supply. Besides the mines of northern Michigan, which yield almost pure copper, large quantities are obtained from the silver ores of Montana and Colorado. Many of the mines of Michigan are exceedingly productive, some of them yielding annually about 25,000 tons, but in recent years the mines of Montana have furnished about 40 per cent of the world's supply. Among the ores found in the Western mines may be mentioned malachite, CuCCX, Cu(OH) 3 , green in color; azurite, 2 CuCCX, Cn(OH) 2 , 232 THE COPPER-SILVER GROUP 233 a beautiful blue, usually associated with the malachite; chalcopyrite, or copper pyrite, CuFeS 2 , a brass-colored ore, resembling fool's gold, but often having a purplish cast ; and bornite, a sulphide of iron and copper of varying proportions, usually Cu 3 FeS 3 . 3. Reduction of the Ore. In the case of the copper from the Lake Superior mines, scarcely any refining is necessary. It is passed through crushers to break up the rock associated with the metal, then by washing and other mechanical processes the separation is effected. 4. Methods in the West. When the ore is a carbonate, like malachite or azurite, or the oxide, it is simply mixed with coke and reduced according to the general plan of reducing metallic ores. Thus, CuO + C = Cu + CO. Usually, however, there is a high per cent of sulphur present, and the process is much more complicated. There are, in reality, four stages necessary before blister copper, that is copper about 98.8 per cent pure, is obtained. These four are concentration, calcination, reverberation or blast re- duction, and converting. The first consists in the separation of the silica or rock from the copper ore. This is done by mechanical washing with "jiggers." By calcination the sulphur is partially removed. After th~ ore has been roasted, either one of two plans may bo followed. Accord- ing to one method, the red-hot ore is placed in reverbera- tory furnaces and melted. The sulphide of copper, mixed with the sulphide of iron, always present, and the silver and gold, being heavy, settle to the bottom. This molten mixture is drawn off and is known as matte. 5. Sometimes the ore, even without concentration or calcining, is put directly into blast furnaces. In this case 234 MODERN CHEMISTRY limestone rock is mixed with the ore ; when the mass is heated the silica and limestone unite to form a glassy slag which takes up about 75 per cent of the iron. The slag, being relatively light, is drawn off above the metal. The sulphur in excess is removed by the strong draughts of air which are forced through the blast furnace. A matte is thus obtained similar in composition to that pro- duced by reverberation. 6. The fourth stage consists in converting this matte into blister copper. This is done in a converter, which in its essentials is not unlike the Bessemer converter described in detail in the chapter on iron. The molten matte has fine streams of air driven through it, and in a few minutes is converted into copper about 98.5 per cent pure. This still contains small quantities of iron, arsenic, gold, and silver, which are finally separated at the refineries. EXPERIMENT 138. Put upon charcoal a little copper oxide, CuO, mixed with sodium carbonate, and heat strongly in the reducing flame. Note the color of the granular mass remaining. Test its malleability with a hammer. What have you obtained ? 7. Characteristics of Copper. Copper is a very tena- cious, malleable, ductile metal, of a reddish color. It does not tarnish in dry air, but in the presence of moisture and carbon dioxide is slowly converted into a green carbonate of copper. With the exception of silver it is the best conductor of electricity known. Its melting point is high, being nearly 1100 C. In the oxidizing flame it is con- verted into the black oxide of copper, CuO. It is soluble in nitric acid and in hot concentrated sulphuric acid. From its solutions it is easily precipitated by iron, zinc, and certain other metals. THE COPPER-SILVER GROUP 235 8. Applications in the Arts. With the exception of iron, copper, probably, has more varied uses than any other metal. It is employed very extensively in alloys, among them being the following : Brass : consisting of copper and zinc in varying pro- portions. Bronze : copper, zinc, and tin. Bell-metal : copper and tin. Coinage : gold and silver with copper. Aluminum bronze : aluminum and copper. A peculiarity of the last is that, with about 1 to 3 per cent of copper, it is of a beautiful silver-white color, much whiter even than aluminum; with 10 per cent of copper it somewhat resembles gold. In the latter propor- tions it is used largely for making various fancy articles and novelties. 9. Unalloyed, copper is used for roofing, for the sheath- ing of vessels, for making various utensils, and for wire for trolley, telegraph, and telephone systems, and for electric lighting. Compounds of Copper 10. Two Classes of Salts. Copper, like several other metals, forms two classes of salts, cuprous and cupric, though as a rule only the latter are of importance. 11. Cupric Sulphate, CuS0 4 , 5 H 2 0. This is commonly known as blue vitriol. It forms in beautiful blue crystals, and is obtained when metallic copper is dissolved in boil- ing sulphuric acid. The commercial supply is obtained mostly as a by-product from the great gold and silver refineries, such as those of Kansas City and Omaha. The smelters at the former place produce monthly about eight- een hundred tons, worth between $100,000 and 1200,000. The silver ores contain more or less copper in the form of 236 MODERN CHEMISTRY cupric sulphide, which in the roasting of the ore is con* verted into cupric sulphate. CuS + 2 O 2 = CuSO 4 . This, being soluble in water, is washed out and concentrated, whereupon the crystals separate out from the solution. 12. Characteristics and Uses. The salt is somewhat efflorescent, and when exposed to the air gradually gives up a portion of its water of crystallization. At the same time it breaks up and becomes almost white in color. By heating, the water of crystallization may be entirely re- moved and the blue color destroyed. This may be restored, however, by digesting for some time in water. Blue vitriol is very poisonous, and is used extensively in making Paris green and Bordeaux mixture for spraying fruit trees to destroy moths and other insects. It is employed largely in electroplating and electrotyping, also in galvanic batteries, though the dynamo is now taking the place of these batteries. 13. Cupric Nitrate, Cu(N0 3 ) 2 , 3 H 2 0. This is a deep blue solid, soluble in water, obtained when copper is treated with nitric acid. 14. Cupric Chloride, CuCl 2 . This is a beautiful tur- quoise-blue, finely crystallized salt. 15. Cupric Sulphide, CuS. This is a black precipitate obtained when a current of hydrogen sulphide is passed through a solution of a copper salt. It is soluble in hot nitric acid, and partially so in warm yellow ammonium sulphide. 16. Cupric Acetylide, CuC 2 , H 2 0, or CuC 2 . Cupric acetylide is a reddish brown precipitate formed when acetylene is passed through a copper solution. In drying it gives up its molecule of water and becomes very explo- THE COPPER-SILVER GROUP 237 sive, a slight jar being sufficient to touch it off. Metallic copper which has for some time been in contact with moist calcium carbide is partially converted into the acetylide, and shows the same explosive tendencies. 17. Cupric Oxide, CuO. This is a black powder, ob- tained when copper is heated to redness in the air, or when cupric nitrate is treated in a similar manner. In the hydrated form, CuO, H 2 O, it may be obtained by treating a copper solution with caustic soda or potash and boiling for a few minutes. EXPERIMENT 139. To prepare some of the above compounds. The nitrate and sulphate have already been prepared. Review the work with nitrogen dioxide and sulphur dioxide. Add to a few cubic centimeters of copper nitrate solution a few drops of ammonium sulphide, (XH 4 ) 2 S, or pass through it a current of hydrogen sulphide. Xote the color of the precipitate formed. What is it? Put into a crucible or small evaporating dish a half gram of pow- dered copper nitrate, and heat gradually to dull redness. How is the nitrate changed ? What gas did you see expelled? What have you obtained ? Save the powder. Put into a test-tube a few cubic centimeters of a solution of copper nitrate or sulphate, and add a little caustic soda or potash. A blue precipitate of cupric hydrate is obtained, Cu(OH) 2 . Boil it for a few minutes. Notice the change in color. What have you obtained? Make a borax bead upon a platinum wire and fuse into it a little of the cupric oxide prepared above. What colored bead do you ob- tain ? The oxide is thus used sometimes in preparing emerald glass. EXPERIMENT 140. Practical Work. To determine the composi- tion of brass. Dissolve a few brass filings in warm nitric acid. Notice the color of the solution obtained. What metal is indicated by the color ? Evaporate nearly to dryness, and take up with 40 to 50 cc. of water. W T arm gently and pass a current of hydrogen sulphide for several minutes, or until no further precipitate will form. This may be determined by filtering out a little and passing the gas through it. If no precipitate forms, the whole may be filtered. Punch a hole in the bottom of the filter as it rests in the funnel, and wash the black 238 MODERN CHEMISTRY precipitate through into a beaker with a little nitric acid diluted. Heat until it dissolves. What is indicated by the color of the solu- tion ? To prove, add ammonia until alkaline. Do you obtain a deep blue solution? If so, copper is indicated. The nitrate obtained above from the black precipitate will con- tain the other metal or metals found in the brass. Add to it a few drops of ammonium hydroxide and then a little ammonium sulphide. Do you obtain a starchy white precipitate ? If so, zinc is indicated. EXERCISE. Write reactions showing the preparation of cupric sulphate, nitrate, sulphide, acetylide, hydrate, oxide, and the reactions in the analysis of brass as far as possible. SILVER : Ag = 108 18. Ores of Silver. This metal has been known from remote antiquity, because of the fact that it frequently occurs free in small particles disseminated through quartz and other rock. Occasionally large masses have been found, and in the museum at Copenhagen there is to be seen one weighing about five hundred pounds. Usually, however, silver is in combination with other elements. One of the most important ores is horn silver, AgCl, named from its general resemblance in color and texture to the horns of cattle. Another important ore is argentite, Ag 2 S. As the greater part of the lead ore smelted contains more or less silver, lead furnaces yield the largest portion of the world's output of silver. 19. Reduction of the Ores. The following experiment will illustrate roughly one of the methods by which silver ores are reduced. EXPERIMENT 141. To about 10 cc. of a solution of silver nitrate, add a little hydrochloric acid. The precipitate is silver chloride, AgCl; shake the contents, warm slightly, and when the precipitate has settled, decant the moderately clear solution. Transfer the curdy white precipitate to a piece of charcoal, cover with sodium carbonate, THE COPPER-SILVER GROUP 2S9 and heat strongly in the reducing flame. Presently a bright globule of silver will appear. This may be preserved for tests upon the metal if desired, or dissolved in dilute nitric acid and converted again into silver nitrate. 20. Other Methods. Various processes are used in reducing silver ores, depending upon the character of the ore. But so large a proportion of the silver output results from lead reduction, that we shall confine ourselves here to only one or two of the methods employed. When these argentiferous lead ores are reduced (see Lead, page 275), the two metals, silver and lead, are formed together as an alloy, and they must then be separated. There are two methods for doing this. When the alloy is rich in silver, Pattison's method is employed. 21. Pattison's Method. It has been found that when such an alloy is allowed to cool slowly the lead will crys- tallize before the silver. Hence, as the lead crystals begin to form they are skimmed out with perforated ladles, thus dividing the alloy into two portions, one containing the silver with a very little lead remaining in it, and the other the lead, with very small quantities of silver. The first of these is then submitted to cupellation. The alloy is gradually run into a large cupel, or basin, constructed upon a hearth within the furnace. A blast of air and flame is directed upon the surface of the alloy, and the lead is oxidized to litharge, PbO. The current of air constantly drives this film of oxide off into another vessel so placed as to receive it. In this way the lead is entirely removed, and the completion of the process is known by the brilliant appearance of the molten silver. 22. Parke's Process. Zinc will readily alloy with sil- ver but not with lead, and this principle is made use of in Parke's process of separating lead and silver. Zinc is 240 MODERN CHEMISTRY added to the alloy, and the whole is melted. The alloy of zinc and silver, being lighter than the lead, rises to the surface, and as it begins to solidify is skimmed off in the form of crystals. "Thus there is obtained an alloy of zinc and silver with very little lead adhering. This alloy is now very carefully heated in a -furnace, the bottom of which is inclined ; the lead melts and runs off before the fusing point of the alloy is reached. The zinc still remaining is next removed by heating strongly in retorts, when it is vaporized and passes off. EXPERIMENT 142. Making use of the bead of silver obtained above in Experiment 141, test its hardness and malleability. Try to oxidize it in the oxidizing flame. Does any coating form upon the charcoal? For just a moment put a silver coin into a solution of hydrogen sulphide or sodium sulphide. What are the results? Next immerse it in a moderately strong solution of potassium cyanide, and allow it to remain some time, if necessary. State the results. This last suggests a method of cleaning tarnished silverware, but it should be used with caution, as the cyanide is deadly poison. EXPERIMENT 143. Add to 2 or 3 cc. of silver nitrate a little hydro- chloric acid, spread the white precipitate smoothly upon a sheet of paper, place upon it any figure cut from thick paper, and expose it to the light. In a few minutes, notice what has happened. This illustrates the method of printing from photographic negatives upon sensitized paper. Before Fie. 53. After X,, . . , , Exposure. Exposure. The experiment may be varied, and with care and patience most beauti- ful prints may be obtained. Immerse in a solution of silver nitrate a sheet of drawing paper, and allow it to dry in the dark. Next immerse in a solution of common salt, and again let it dry in the dark. When ready to print, place upon this paper, thus sensitized, an old negative, or even a fern leaf or any similar object, and expose to bright sun- light, under a sheet of glass to hold in place. Notice when a deep pur- ple is obtained, then immerse in a solution of sodium thiosulphate, the photographer's " hypo," and rinse thoroughly in water several times. THE COPPER-SILVER GROUP 24i 23. Characteristics of Silver. Silver is a white, lustrous metal, malleable and ductile, an excellent conductor of electricity and heat, of medium hardness and density. It is quickly attacked by many sulphur compounds and by the members 'of the halogen group, although it does not tarnish in the air at any temperature. In living rooms silverware is tarnished by the action of the sulphur gases thrown off in the combustion of coal or of ordinary illuminating gas. Eggs and various other articles of food tarnish silverware for a similar reason. What is known as " oxidized " silver is really that which has been treated with some compound of sulphur, producing silver sulphide upon the surface. 24. Uses for Silver. Owing to its brilliancy and dura- bility, silver has long been used for jewelry and various other articles of ornament. Alloyed with some other metal to make it harder, it is employed extensively in coinage ; is also used in amalgams for dentistry and for the backs of high grade mirrors, and for plating innu- merable articles of use and ornament. Compounds of Silver 25. There are only a few compounds that are of interest, and but one or two that are of any considerable value. 26. Silver Nitrate, AgN0 3 . This is important because most of the other silver compounds are prepared from it, and because it has numerous applications in the arts. It occurs in slab-like, almost transparent, white crystals, which are soluble in water. It is prepared by dissolving silver in nitric acid. When exposed to the light, espe- cially if in contact with any organic matter, it turns dark. It is used for sensitizing paper for photographic work, as the principal ingredient of indelible ink, and in hair dyes. 242 MODERN CHEMISTRY In the form of lunar caustic, which is simply crystallized silver nitrate fused and molded into sticks, it is used in cauterizing wounds, such as dog bites, for ulcerated sore throat, in removing warts and other similar excrescences of the skin. 27. Silver Chloride, AgCl. This is prepared from a solution of silver nitrate by adding to it hydrochloric acid or any soluble chloride, like common salt. It is a white precipitate, curd-like in appearance, especially when shaken for a moment. It is Soluble in ammonium hydroxide, and in sodium thiosulphate, "hypo." It is much more sensi- tive to light than the silver nitrate, and hence for photo- graphic work the latter salt is generally converted into the chloride, or bromide^ which is even more sensitive. It is believed that the light gradually converts this compound back into metallic silver, which is insoluble in the " hypo," while the unchanged portions of silver chloride are dis- solved out and the paper thus de-sensitized. EXPERIMENT 144. To about 1 cc. of a solution of silver nitrate add a few drops of hydrochloric acid. Notice the appearance of the precipitate that forms. What is it? Write the reaction. To a por- tion of it add a little ammonium hydroxide and shake it. What results? To another portion add a solution of "hypo" and state the results. 28. Silver Chromate, Ag 2 Cr0 4 . This is a blood-red powder obtained as a precipitate when potassium chro- mate is added to a solution of silver nitrate. EXPERIMENT 145. Prepare the chromate as indicated, and note its appearance. 29. The formation of silver chloride and the chromate, with their characteristic appearance and the ready solu- bility of the former in ammonia, serve to distinguish a solution of silver, and may be used as tests. THE COPPER-SILVER GROUP 243 EXERCISE. Write the reactions, showing the preparation of sil- ver nitrate, silver chloride, and the chromate; also silver bromide and iodide, from silver nitrate with potassium bromide, and with sodium iodide. 30. Photography. At the present time almost all young people take more or less interest in this wonderful art. The first experiments along this line were made as early as 1727, but they were nothing more than what the student has done in the first part of Experiment 143, and the print soon disappeared. From that time to this many different plans have been tried, but we can only notice briefly that used at present. 31. Preparation of the Plates. The plates upon which the negatives are made are prepared as follows : for the most sensitive plates, potassium or ammonium bromide with gelatin and silver nitrate added is dissolved in water and heated to boiling. Thus the silver is converted into silver bromide : KBr + AgNO 3 = AgBr + KNO 3 . An excess of water is added, and the potassium nitrate formed is readily washed away. This gelatin emulsion, as it is known, is poured upon glass plates and allowed to harden ; they are then ready for use. 32. Exposure and Developing. As previously stated, when such plates are exposed to light, the silver salts are decomposed. In the camera the exposure is so brief that the decomposition is only partial; when, however, the plate is put into the developer , this solution continues the action begun by the light. Hence those portions of the plate which have received the most light have the larger amount of the silver salts decomposed, and are dark in color. If allowed to remain in the developer long 244 MODERN CHEMISTRY enough, all the silver would be reduced, and the plate would be uniformly dark. 33. Fixing. When it is seen by examination that the development has proceeded long enough, the plate is rinsed in water and placed in the fixing bath. This is a solution containing sodium thiosulphate, which is an excellent solvent for many silver compounds. The fixing bath soon removes from the gelatin film the silver bromide or chlo- ride that remains unaffected by the light or by the devel- oper. The plate is thus cleared or fixed, and is no longer sensitive to light. As the lights and shadows are all reversed, it is called a negative. After thorough washing it is allowed to dry, when it is ready to be used in making prints. 34. Printing. Various kinds of paper are now used for making prints, among them being the solio, velox, platinotype, carbon, and Hue print. The first and last of these require the least skill. Solio has a sensitized film of silver chloride ; in printing, this is placed against the film side of the negative, which causes the objects to appear in the picture in their natural position. As the dark portions of the negative transmit the fewer light rays, the picture appears as a positive, or like the original as to high lights and shadows. The advantage of solio is in the fact that it is only moderately sensitive, and that it readily shows when it has been exposed long enough. More sensitive papers, such as the velox, are like the gelatin plates in that they show no image until treated with a developer. Solio prints require toning, and all varieties need fixing by some method or other. In the platinotype papers, a compound of platinum is used which yields the dark appearance now so much admired. THE COPPER-SILVER GROUP 245 35. Solio papers cost so little that it would be easy for a class to make some experiments along this line. Let any of the pupils who may have them bring in some of their negatives and printing frames, and do some work of this kind. 36. Blue prints* are the simplest of all, are cheap, and yet for landscapes often give most excellent effects. They possess the advantage of requiring no toning or fixing except such as is secured by thorough washing. Place the paper under the negative in direct sunlight, and allow it to remain until the high lights begin to look somewhat muddy in appearance ; then put into a basin of water with the printed side down. Allow the print to remain there until the light portions are quite clear, then wash for ten minutes in running water. The beauty of these prints will be enhanced by leaving a pure white border around the picture ; this may be secured by using a black mat between the negative and print so as to cover the portion which it is desired to have white. * If he desires, the instructor may prepare his own blue print paper. Make Solution A Citrate of iron and ammonia . . 1 g. Water . . . . . . . . 10 cc. Solution B Potassium ferricyanide . . . 1 g. Water . . . . . . . 10 cc. When ready for use mix A and B in a dark room, and apply to the paper with a brush ; or, the paper may be floated in the solution. This must be used within a day or two after it is prepared, as it does not keep well. A few drops of a 10 per cent solution of potassium bromide added to A and B above will render the keeping qualities of the paper much better. 246 MODERN CHEMISTRY GOLD: Au = 197 37. Occurrence. From the fact that gold occurs free, it has been known from the earliest antiquity. It is widely distributed over various portions of the earth and usually occurs in fine grains and nuggets disseminated through the rocks. These are gradually disintegrated and brought down by rains and streams in the form of sand and gravel, with which the gold is associated. The best- producing gold regions are those of the western part of the United States, Australia, Southern Africa, and the Klondike. Gold also occurs in quartz veins deeply buried in the earth's strata. 38. Methods of Mining. The original method consisted simply in cradling or panning the sand and gravel ; thus the nuggets and larger grains find their way to the bot- tom, while the lighter stone and earthy matter is washed out. By this method only the larger particles are saved. Placer mining consists in washing the gold-bearing sand down through sluices, along the bottom of which are arranged pockets of mercury, or over plates of copper amalgamated with mercury. This readily amalgamates with the gold, and the other portions are carried aw T ay by the current. The gold amalgam thus obtained is heated in retorts, by which the mercury is vaporized, leaving the gold behind. The vapors of mercury are conducted into cold chambers where they are condensed, so that very little loss occurs. Hydraulic mining differs from the above only in that streams of water are directed with great force against the loose rock and cliffs over- hanging, washing them down into the sluice-ways. 39. Vein Mining. Vein mining differs from placer mining in that the latter is surface mining, while in the THE COPPER-SILVER GROUP 247 former the ore is taken from greater or less depths, usually from quartz veins ; hence it is sometimes called quartz mining. Gold sometimes occurs in combination with iron in pyrites, and it is then obtained by the wet or Morination process. The ore is roasted, then moistened and treated with chlorine, obtained usually from bleach- ing powder. The chlorine dissolves the gold, forming gold chloride, AuCl 3 . This is now dissolved out and ferrous sulphate added, which precipitates gold in the metallic condition, as seen in the following reaction : - 6 FeSO 4 + 2 AuCl 3 = 2 Au + 2 Fe 2 (SO 4 ) 3 + Fe 2 Cl 6 . 40. Cyanide Process. Potassium cyanide is an excel- lent solvent for gold, and at the present time is used extensively in separating it from its ores. The process is valuable where the gold occurs in a finely divided form; another advantage is that the ore does not need the roast- ing that is necessary in the chlorination process. After the gold-bearing quartz has been finely crushed, it is treated with a solution of potassium cyanide in water. The gold is dissolved out, thus : 4 Au + 8 K(^T+ 2 + 2 H 2 = 4 KAuCy 2 + 4 KOH. 41. The oxygen shown in the reaction is derived from the air, and it has been found that, unless the surface of the ore is left well exposed, the process is not satisfactory. The double cyanide of gold and potassium thus obtained is treated with zinc, which precipitates the gold, as shown in the reaction : 2 KAuCy 2 + Zn = K 2 ZnCy 4 + 2 Au. 42. There is always some zinc left in a more or less finely divided form which cannot be separated mechani- cally from the gold ; hence, when melted down the metal 248 MODERN CHEMISTRY is seldom over 80 per cent pure. For this reason some companies prefer to deposit the gold by electrolysis upon lead terminals. By this method, after oxidizing the lead in cupels, the gold remains in a very pure form. 43. Characteristics. Gold is a bright yellow metal, which, seen in light reflected several times, looks red. It is so soft that for ordinary purposes it must be alloyed with some other metal ; it is heavy, is not affected by the oxygen of the air at any temperature, is very ductile and malleable. Advantage is taken of this property in ham- mering out the metal into gold leaf, the thickness of which is not over 3-5- oW^ P art ^ an i ncn 1500 of which sheets are necessary to make one as thick as ordinary note paper. Pure gold is not affected by single acids, but is readily attacked by aqua regia, forming gold chloride, AuCl 3 . However, if richly alloyed with several other metals, it becomes soluble in single acids. 44. Uses. These are too well known to need specifi- cation. In the arts gold leaf has numerous uses, such as in making display signs, covering high grade moldings, for filling teeth, etc. SUMMARY OF CHAPTER COMPARATIVE STUDY Copper, Silver, Gold. Histories Wherein are they similar? Why ? Occurrence In what forms ? Most productive regions. Some important ores. Various forms of gold mining Description. Reduction of the ores. Special plans for copper. Special processes for gold reduction. Chlorination and cyanide. Special plans for separation of silver. Pattison's and Parke's. THE COPPER-SILVER GROUP 249 Comparison of the three metals as to a. Color. b. Density. c. Melting point. d. Permanency in the air. e. Malleability. f. Conductivity. g. Solubility in acids. Uses of the metals. a. Important alloys. b. Other uses Why so used. Compounds Most important. Of Copper The Sulphate Commercial name and formula. How obtained. Characteristics and uses. Of Silver The Nitrate Commercial name and formula. How prepared. Appearance and uses Why so used ? What other compounds prepared from this one ? How ? Special points. Meaning of the terms blister copper, matte, concentration, . calcination, converting. Describe method of determining the composition of brass. Meaning of terms cupel, cupellation. Describe experiment illustrating principles of photography. Method of sensitizing photographic plates. Chemistry of the developing and fixing of negatives. Reactions showing the preparation of CuSO 4 , CuO, AgCl, AgBr, Agl. CHAPTER XXI ZINC, CADMIUM, MERCURY ZINC : Zn = Go 1. History. Brass, an alloy of copper and zinc, has been known for centuries, but it was formerly made by fusing together copper and a mineral called calamine, which we now know is an ore of zinc. It was not until about the close of the seventeenth century that zinc was recognized as a distinct metal and its characteristics care- fully determined. 2. Ores of Zinc. Zinc occurs abundantly in many parts of the United States and Europe. In Missouri the mines of Joplin and Webb City are the best known. There thousands of tons are produced annually. Kansas also yields a considerable quantity. The ore most generally found in these states is the sulphide, ZnS, known as zinc blende. By the miners it is called "jack," or in its purer forms " rosin jack," because of the general resemblance of a broken specimen of the ore to rosin. In New Jersey the ore franklinite is the most abundant. It is a mixture of zinc oxide and ferric oxide, (ZnFe)Fe 2 O 4 . Other sections yield the carbonate, ZnCO 3 , known as smithsonite. It is said that the metal is sometimes found pure in Australia. 3. Reduction of the Ores. The general method em- ployed in the reduction of the greater number of metallic ores is used in the case of zinc. They are first ground fine aad roasted. This not only drives out certain volatile 250 ZINC, CADMIUM, MEEGUEY 251 impurities, such as arsenic, but converts the ore into the oxide, ZnO, the most convenient form for the next step. The reaction that takes place when the ore is roasted may be seen from the following : ZnS + 3 O = ZnO + SO 2 , ZnCO 3 + heat = ZnO + CO 2 . 4. The oxide thus obtained is mixed with powdered coke and heated red hot in earthen cylinders about 4J feet long, placed horizontally over one another. The zinc is thus reduced to the metallic form, and at the tempera- ture obtained is vaporized. The vapors pass out into conical-shaped earthen condensers attached to the outer end of the retorts, where they liquefy. Twice in twenty- four hours these condensers are " tapped " and the molten zinc drawn off and run into molds. The chemical change taking place is a familiar one : ZnO + C = Zn + CO. The retort is shown by R in the figure and the con- denser by C. The condensers are readily detached, and when the retorts have been charged or filled with the mixed ore and coke they are again attached and luted on nearly air-tight with clay. ^ru 4.1 FIG. 54. When in operation there is usually some escape of vaporized zinc with other gases, and these, in burning at the mouth of the condensers, give a beautiful display of colors, yellow and blue and white, which, especially at night, is exceedingly striking. 5. A single charge requires about twenty-four hours for complete reduction, and as the workmen are usually paid by the amount of metal they " draw off " they gen- 252 MODERN CHEMISTRY _ erally work twenty-four hours successively, and then are off during the next twenty-four. The zinc obtained in this way is more or less impure ; it almost always contains some cadmium, and usually some arsenic, and is known as " spelter." 6. Characteristics of Zinc. Zinc is a bluish white metal of moderately low melting point, about 420 C. ; it tarnishes but slightly in the air, and then only upon the surface. At a temperature slightly above the melting point it burns with a brilliant, bluish white flame, and if a jet of oxygen be directed upon it the light is almost dazzling. EXPERIMENT 146. Examine a piece of zinc and note its color, malleability, hardness, and tendency to oxidize. Test also its melting point by heating a small piece on charcoal with the blowpipe. Try it also with, the oxidizing flame and note the deposit upon the char- coal, both when hot and when cold. State the results. EXPERIMENT 147. To learn the solvents for zinc. Try a small piece of the metal in a test-tube with hydrochloric acid. How is it affected? What gas is obtained? What proof can you offer? Write the reaction. In the same way try nitric acid, and compare results with the above. Into each of two test-tubes put a small piece of zinc. To one add about a cubic centimeter of copper sulphate solution, and cover the other with water. After a few moments, to each add a little sulphuric acid. Is there any difference in the rapidity of the chemical action in the two cases? Why? EXPERIMENT 148. Sift some zinc dust through a wire sieve of fine mesh upon a Bunsen burner flame and note the results. 7. Further Characteristics of Zinc. Ordinary com- mercial zinc as it comes from the smelter is brittle, but if it is heated to something over 120, and then rolled into sheets or drawn into wires, it is found to be malleable, and will remain so. As it approaches the melting point, how- ZINC, CADMIUM, MEECUET 253 ever, it again becomes brittle, and may be ground into a powder known as zinc dust. It is of medium density, being a little lighter than iron, is not magnetic, and when chemically pure is but slightly soluble in dilute acids. When impure, or if in contact with some other metal, as copper or platinum, the solution is rapid. 8. Uses for Zinc. In the metallic form zinc is used extensively in many varieties of galvanic batteries, also as linings for refrigerators, bathtubs, and for various other domestic purposes. One of its most important applica- tions is in coating or " galvanizing " iron wire and other forms of iron as a protection from moisture. Galvanized iron is prepared by thoroughly cleansing the iron to be coated, heating it, and plunging it into a bath of molten zinc until a thin covering of the latter metal adheres. There are also three important alloys : Brass : consisting of zinc and copper in varying propor- tions ; Bronze : zinc, copper, and tin ; German silver : zinc, copper, and nickel. In the chemical laboratory zinc is frequently used : in making hydrogen ; in reducing ferric compounds to the fer- rous condition ; and for precipitating various metals from their solutions. Compounds of Zinc 9. Zinc Sulphate, ZnS0 4 , 7 H 2 0. White Vitriol. This is a white crystalline salt which has been obtained in the preparation of hydrogen by treating zinc with sulphuric acid. It is very soluble in water, and is used mainly for calico printing. It has a bitter, astringent taste. 10. Zinc Chloride, ZnCLj. This is a white solid, ob- tained when zinc is dissolved in hydrochloric acid. It has great affinity for water, and is, therefore, often used in 254 MODERN CHEMISTRY ' chemistry as a drying agent. It is also frequently used as a soldering solution, but as it is poisonous, serious results have sometimes followed its use in soldering tin cans con- taining fruits and other food products. 11. Zinc Hydroxide, Zn(OH) 2 . This compound of zinc may be studied in the following experiment : EXPERIMENT 149. To a few cubic centimeters of a solution of any zinc salt, as the chloride or sulphate, add a few drops of ammo- nium hydroxide. What are the results? Add more ammonia; does the precipitate dissolve? Describe the precipitate, Zn(OH) 2 , that formed, and write reaction. In the same way prepare a little zinc hydroxide by using a solution of caustic soda or potash instead of am- monia as above. Test a portion of the precipitate with hydrochloric acid ; does it dissolve ? Write the reaction. 12. Zinc Sulphide, ZnS. Many characteristics of zinc sulphide may be discovered from the following experi- ment : EXPERIMENT 150. To a few cubic centimeters of a solution of some zinc salt add two or three drops of ammonium sulphide. De- scribe the precipitate that forms. It is zinc sulphide. Test its solu- bility in dilute hydrochloric or nitric acid. 13. Zinc Oxide, ZnO. Zinc White. This was the white deposit formed on charcoal when the zinc was heated by the oxidizing flame. It is now used extensively as a sub- stitute for white lead in painting, and is preferable in localities where much coal is used as fuel, because of the discoloration of lead compounds by the considerable quan- tities of Hydrogen sulphide found in coal smoke. * CADMIUM, Cd = 112 14. Supply. Cadmium is a rare element, discovered about 1817. It received its name from a Greek word,, * This is an unimportant element, and its study may be omitted, if desired. ZINC, CADMIUM, MERCtTRT 255 kadmeia, an ore of zinc, now known as calamine, with which cadmium is usually associated. Our present supply is obtained mostly from zinc ores, with which it is found, in the form of a sulphide, CdS, called greenockite. 15. Reduction of the Ore. In smelting cadmium-bear- ing zinc ores, they are first roasted in retorts, where both sulphides are converted into oxides, thus : These oxides are then mixed with coke or charcoal and again heated, when the usual reduction takes place : = ZnOJ IZn To separate the two metals thus obtained they are dis- solved in hydrochloric acid, and the solution treated with rods of zinc, by which the cadmium is reduced to the metallic form, thus : and + Zn = Cd + 2 ZnCl 2 . 16. Appearance and Characteristics. Cadmium is usu- ally marketed in the form of small rods, 8 or 10 inches in length. It is a white metal, closely resembling tin, and is of about the same hardness, but it has a melting point not very different from lead, 315, and boils at 860. Cadmium tarnishes slowly in the air, becoming coated with a very thin covering of yellow oxide. It is malleable and ductile, and when bent, like tin, gives a creaking sound. With mercury it forms a silvery white amalgam which 256 MODERN CHEMISTRY soon becomes hard and brittle. It is easily soluble in nitric acid, less so in hydrochloric and sulphuric acids. It is but little used in the arts, though it has been applied somewhat as a filling for teeth ; but as a cadmium amal- gam gradually turns dark, it has not found favor with dentists. Compounds of Cadmium 17. Cadmium Nitrate, Cd(N0 3 ) 2 . This is a white salt obtained when the metal is dissolved in nitric acid. 18. Cadmium Sulphide, CdS. This is a yellow powder used in oil and water colors. Artificially, it is obtained when a current of hydrogen sulphide is passed through a solution of any cadmium salt. It resembles the sulphides of arsenic and tin, As 2 S 8 and SnS 2 . Unlike the arsenic, however, cadmium sulphide is not soluble in ammonium carbonate, and unlike the tin, is insoluble in yellow ammo- nium sulphide. EXERCISE. Write reactions showing the formation of cadmium nitrate, chloride, sulphate, and sulphide. I ) MERCURY: Hg = 200 \* \ / 19. Historical Facts. Mercury was one of the seven elements known to ancient chemists, and by them was dedicated to the god from which it received its name. Its symbol is taken from the Greek word, hydrargyrum, by which name it was also known. This term means water (or liquid) silver. Similarly, at the present time it is spoken of as quicksilver. By Geber, the famous alchemist of the eighth century, mercury and sulphur were regarded as the two elements from which all metals could be made. He claimed that any one knowing the proper proportions could prepare any of the noble metals from these two. ZINC, CADMIUM, MERCURY 257 20. The Source of Supply. The commercial supply of mercury comes from its chief ore, cinnabar, or vermilion, HgS. This is an exceedingly heavy, brick-red mineral, found in Spain, India, Bavaria, California, Mexico, etc. EXPERIMENT 151. Near one end of a piece of hard glass tubing place a little vermilion, HgS, as much as will remain on the point of a knife-blade. Now, with this end down, hold in a slanting position in the Bunsen burner and heat strongly. Notice the formation on the upper, cooler portion of the tube. What gas, detected by its odor, is given off from the upper end of the tube ? Name the two products resulting from the heating of mercuric sulphide. Compare with the preparation of oxygen from mercuric oxide. 21. Reduction of the Ore. This experiment illustrates the reduction of cinnabar in the preparation of mercury for commerce. The ore is placed upon shelves in an oven over a furnace (see Fig. 55). Hot blasts of air flow up through the shelves, oxidizing the sulphur to sulphur dioxide, and at the same time vaporizing the mercury. These gases pass out together into cool chambers, where the mer- cury condenses, while the sulphur dioxide escapes. As thus obtained the mercury is more or less impure. It is purified first by being strained through porous leather or chamois skin, and then distilled at moderate temperatures. 22. Characteristics. Mercury is the only metal that is liquid at ordinary temperatures. At 39 below zero it becomes a solid, and in that condition possesses some of the properties of lead. It has about the same color, is malleable, and soft enough to be cut easily. Mercury is a silver- white metal, which does not tarnish in the air, but which slowly vaporizes at all temperatures. 258 MODERN CHEMISTRY 23. Amalgams. The most remarkable property of mercury is its power of dissolving many of the metals and forming with them what are known as amalgams. If the mercury be largely in excess, the other metal dis- appears as a lump of sugar does in a cup of tea ; if a smaller proportion be used, the mercury simply combines with the outer portions of the other metal, changing more or less its appearance and general properties. There are two methods of forming amalgams. 24. a. By bringing metallic mercury into contact with a metal perfectly clean. If this is broken up into small pieces, or in the form of dust or filings, and is then heated with the mercury, the amalgamation takes place quickly. EXPERIMENT 152. Into a few drops of mercury in an evaporating dish, put a perfectly clean strip of zinc. After a few moments, ex- amine it; has it changed in appearance? Bend it; has it changed in properties? Try in the same way a five-cent piece, a penny, a nail, or any other convenient metals. Be careful, however, "of any gold rings, as mercury amalgamates very readily with gold. 25. b. The second general method is by immersing the metal to be amalgamated in a solution of some salt of mercury. Try in this way the following : EXPERIMENT 153. Put into a beaker, or evaporating dish, a few cubic centimeters of a solution of mercurous nitrate, Hg 2 (NO 3 ).>. Immerse in it a brass pin, or a thimble, a copper penny, a key ring, etc. After remaining a few minutes, they may be removed and rubbed a little, if dull in appearance. State which have been amalgamated. 26. This second method is employed frequently by street fakirs as a means of " silver plating." They pre- pare the solution by dissolving mercury in nitric acid and then adding some coloring matter. This very rapidly ZINC, CADMIUM, MEECUET 259 -X- - " plates " certain metals, but the amalgamated Articles retain their brilliancy but a short time. 27. Solvents for Mercury. The best solvent for mer- cury is nitric acid, which attacks the metal even at or- dinary temperatures. When heated, sulphuric acid also dissolves it, with the formation of sulphur dioxide gas. Compare this with the preparation of sulphur dioxide as given on page 177, section 15. 28. Uses for Mercury. Mercury is employed exten- sively in the manufacture of thermometers and barom- eters ; in the laboratory it is often used instead of water in the pneumatic trough for collecting such gases as are soluble in water, especially when they are desired perfectly free of air. Large quantities are also used in placer mining of gold and silver (see page 246). In the form of amalgams it is used with various other metals for filling teeth ; with tin or silver for the backs of mirrors, for rendering zinc plates to be used in batteries less solu- ble in acids, and sometimes for amalgamating surfaces which are to be silver plated. This is done because silver seems to adhere better to a surface which has been thus treated. Compounds of Mercury 29. Like several other metals, mercury forms two series of compounds, the mercurous and mercuric. 30. The Nitrates, Mercurous, Hg 2 (N0 3 ) 2 ; Mercuric, Hg(N0 3 ) 2 . These may be prepared by treating mercury with nitric acid ; for the former, using dilute acid with the mercury in excess ; for the latter, concentrated, with the acid in excess. Mercuric nitrate is a white salt of fine silky crystals, soluble in water. In dissolving it yields at the same time a yellowish powder, known as basic nitrate, having the formula HgNO 3 , Hg(OH) 3 . Mercurous ni- 260 MODERN CHEMISTRY trate is of a pale yellow color, almost white. It usually occurs in crystals larger than those of the mercuric nitrate and is soluble in water. Both are used in the laboratory, and occasionally for the preparation of other compounds of mercury. 31. The Chlorides, Mercurous, Hg 2 Cl 2 ; Mercuric, HgCl 2 . The former is known as calomel, the latter as corrosive sublimate. Mercurous chloride may be prepared by add- ing to mercurous nitrate, hydrochloric acid, whereupon it falls as a heavy white precipitate. On a large scale it is manufactured by thoroughly mixing in the proper propor- tions mercuric chloride and mercury, heating them strongly to vaporize, whereupon they combine and are condensed in cold chambers. Calomel is a white, flour-like substance, insoluble in water. It is used largely in medicine. 32. Mercuric chloride is prepared by subliming, as de- scribed above in making calomel, a mixture of mercuric sulphate and common salt. It is a white, crystalline salt, somewhat soluble in water, and very poisonous. It is used in the laboratory as a reagent, is a constituent of some vermin exterminators, and has frequent use in surgery as an antiseptic. 33. Mercuric Oxide, HgO. This orange-red salt, com- monly known as red precipitate, is prepared by heating mercuric nitrate for a considerable length of time. It is used sometimes for preparing small quantities of oxygen, and in some quantitative determinations in the laboratory. 34. Mercuric Sulphide, HgS. As an ore it is known as cinnabar, but the artificial product is sold under the name vermilion. It is of a bright scarlet color, and is used in making tube paints and in coloring sealing-wax. As ordinarily prepared in the laboratory, it is black, but under certain conditions is obtained in varying shades of red. ZINC, CADMIUM, MEECURY 261 EXPERIMENT 154. To prepare certain compounds of mercury. Put a drop of mercury into a test-tube and add about a cubic centimeter of dilute nitric acid, warm gently, and after a few minutes, or when the action has ceased, decant the solution ^and boil it nearly dry in an evaporating dish. Now add a few cubic centimeters of water and pour into three test-tubes. To the first add a little potassium iodide ; to the second, hydrochloric acid ; to the third, ammonia. Notice the color of the precipitate in each case. Write the reactions in the first two, and state what compound is formed. Tabulate results as follows : Hg 2 (N0 3 ) 2 Hg(N0 3 ) 2 KI HC1 NH 4 OH q^L You should have prepared mercurous nitrate by the above treat- ment of mercury with nitric acid. To another drop of mercury in a test-tube add some strong nitric acid, and warm until the mercury is all dissolved. Transfer to an evaporating dish, boil nearly dry, and add a few cubic centimeters of water. You should now have a solution of mercuric nitrate. Divide into three parts and treat with the same three reagents that you used with the mercurous nitrate, and tabulate the results. * 35. From the above experiments, it will be seen that the two series of mercury salts may be easily distinguished by the precipitates which they form with different reagents. EXPERIMENT 155. Let the student be given some mercurous~ahd mercuric solutions, and have him determine what each is. EXERCISE. Write out the reactions that take place in preparing the two nitrates, mercurous chloride, mercurous and mercuric iodide, 262 MODERN CHEMISTRY and mercuric sulphate. Before attempting the last, unless you know the results, put into a test-tube a small drop of mercury, add a little strong sulphuric acid, and heat until some familiar gas is produced. COMPARATIVE STUDY Zinc and Mercury Early history. Ores of these metals Most important of each Localities where found. Plan of reduction. Wherein are they alike? How different? Why is carbon not necessary for the mercury? Reactions for each. Description of furnaces. Comparison of the two metals in a. Color. b. Melting point. c. Density. d. Ease of oxidation. e. Malleability. /. Conductivity. g. Solubility. Special properties of each. Brittleness of zinc at certain temperatures. Condition of -mercury at low temperatures. Power of forming amalgams. Names of metals which will amalgamate and of those which will not. Two methods of making amalgams. Important uses of each metal. Alloys of zinc. Amalgams of mercury. Other uses. Compounds. The oxides Appearance and use of each. The chlorides One of zinc, two of mercury. Preparation of each Use Commercial name. Two classes of mercury compounds Methods of distin- guishing them. CHAPTER XXII ALUMINUM AND ITS COMPOUNDS ALUMINUM : Al = 27 1. Abundance. This metal was first isolated about 1827, being reduced by metallic sodium. For some years all that was used in the arts was prepared by strongly heating aluminum chloride, and passing the vapors into which it was converted over sodium. The reaction may be represented^ us : AlClg 4 3 Na = Al + 3 NaCl. By this method about three pounds of sodium were re- quired for the preparation of a single pound of aluminum, and the cost was about one dollar an ounce. 2. No metal occurs more abundantly than aluminum, and but one or two non-metallic elements are more widely distributed. It forms a large per cent of feldspar and of various other ro the jet as before. Is any deposit formed? What that you have already seen does it closely resemble ? The gas being generated is arsine. Now, hold a beaker or test-tube over the burning jet, and notice whether there are not two different deposits formed. Can you decide what they are? Write the reaction that takes place when arsine burns. 5. Arsine. Arsine, AsH 3 , is also known as arseniu- reted hydrogen, or hydrogen arsenide. It is a compound of considerable interest, because it is always prepared in testing for arsenic in cases of suspected poisoning. The method used is the one described in the experiment above. This is known as Marsh's test, and is so exceed- ingly delicate that mere traces of arsenic, even so low as one part in several hundred thousand, can be detected. Care should be taken, however, to see that tke zjjic is 288 MODERN CHEMISTRY perfectly free from arsenic. Antimony gives a spot con- siderably like that of arsenic seen above, but the latter may be detected by treating with a solution of bleaching powder, in which the arsenic spots are soluble, while the others are not. 6. Let us study the reactions that take place. First, by the reaction of sulphuric acid and zinc upon each other hydrogen is produced, thus : Zn + H 2 SO 4 = H 2 + ZnSO 4 . The hydrogen atoms in the nascent condition, instead of uniting with one another to form molecules of hydrogen, unite with the arsenic present, forming hydrogen arsenide, AsH 3 . This may be represented thus : AsCl 3 + 6 H = AsH 3 + 3 HC1. 7. Characteristics of Arsine. This is a colorless, ex- ceedingly poisonous gas, which burns with a pale violet flame, giving off white fumes of the trioxide As 2 O 3 . 2 AsH 3 + 3 O 2 = As 2 O 3 + 3 H 2 O. Both of these products may be seen if a cold beaker or test-tube be held over the burning jet of arsine. If a cold dish is held against the flame, the temperature is lowered below that required for the combustion of arsenic, and it is therefore deposited in the metallic form, while the hydrogen continues to burn. What does the experiment teach regarding the kindling point of hydrogen ? 8. The Oxides of Arsenic. Corresponding to the two oxides of phosphorus we have two of arsenic, the trioxide, As 2 O 3 , and pentoxide, As 2 O 5 . Only the former is of im- portance. It occurs in two^or three forms, the white being the most common. It is usually sold under ABSENIC, ANTIMONY, BISMUTH 289 the name " arsenic " or white arsenic, but is also called arsenious acid. It has a sweetish taste, is slightly soluble in cold water, more so in hot, in hydrochloric acid, and in caustic soda. It is very poisonous, but acts somewhat slowly. An antidote for it is ferric hydroxide, prepared by treating a ferric salt in solution with ammonia; the precipitate must be filtered out and washed. Magnesia, MgO, is also suggested, and is used more often because it is to be had already prepared. 9. Arsenic trioxide is used by taxidermists in curing the skins of animals ; it is an ingredient of many poisons, but is also often prescribed by physicians as a blood purifier, especially for removing facial eruptions. It is thought to beautify the complexion, and has a tendency to produce fat. Because of the latter property it is sometimes fed to old horses to prepare them for the market. It stimulates the action of the heart and renders breathing easier ; on this account it is said to be used by some mountain climb- ers. These apparent benefits are but temporary, however, and a discontinuance of its use is attended by all the symptoms of serious arsenic poisoning. EXPERIMENT 173. Examine a sample of arsenic trioxide and note its general appearances. Test its solubility in diluted hydrochloric acid, also in caustic soda. Which is the better solvent? Use only small quantities of the trioxide. Save the solution. 10. Paris Green; Scheele's Green. This is a very poisonous, bright green powder, used often for coloring and tinting and as an insect exterminator. EXPERIMENT 174. Let the student prepare this compound, thus: To a few cubic centimeters of a solution of copper sulphate in a test-tube add ammonia, drop by drop, until the precipitate which forms at first just dissolves. Now, add gradually a solution of arsenic ; a bright green precipitate will form. If too blue, not enough arsenic ODERN CHEMISTRY 290 has been added. This is one of the easiest methods of detecting arsenic if present in considerable quantities. It is known as Scheete's test. 11. Arsenic Trisulphide, As 2 S 3 . This is a bright yel- low powder obtained by passing a current of hydrogen sulphide through a solution of arsenic. It is soluble in ammonium carbonate, which distinguishes it from a similar compound of tin, SnS 2 , also yellow. It is also soluble in yellow ammonium sulphide, but not in hydro- chloric acid. EXPERIMENT 175. Let the student prepare this compound by passing hydrogen sulphide through a solution of arsenic trioxide in water acidulated with hydrochloric acid. Divide the yellow precipi- tate into two or three parts and test its solubility in hydrochloric acid and in ammonium sulphide and carbonate. ANTIMONY: Sb = 120 \ 12. Source of Antimony. This element is found free in very small quantities only, but frequently occurs with the ores of other metals, such as lead, copper, and iron. Its principal ore is stibnite, Sb 2 S 3 , and from this the commer- cial supply is obtained. 13. Reduction of the Ore. There are two methods used for reducing antimony ores. The first consists in heating the sulphide in a reverberatory furnace, whereby the ore is reduced to an oxide, thus: Sb 2 S 3 + 5 O a = Sb a O 4 +-3 SO 2 . Then the tetroxide, tlius formed, is mixed with char- coal, and again heated in a furnace, when metallic antimony is obtained, thus : ARSENIC, ANTIMONY, BISMUTH 291 EXPERIMENT 176. In a cavity in a piece of charcoal place a little antimony tartrate, mixed with sodium carbonate, and moisten with a few drops of water. Now heat strongly with the reducing flame. What do you obtain? Preserve for the next experiment. 14. This illustrates the method of reduction described above, and, it will be noticed, is in accord with the general plan of reducing metallic ores, first reducing them to the form of an oxide by roasting them, and then deoxi- dizing them by heating with carbon. 15. Another Method. This consists in mixing the ore, antimony sulphide, with iron, and melting the whole in a furnace. The iron combines with the sulphur, and pre- cipitates the antimony, thus : Sb 2 S 3 + 3 Fe = 3 FeS + 2 Sb. 16. Characteristics of Antimony. Owing to the fact that, like phosphorus, nitrogen, and other non-metallic ele- ments, antimony forms oxides which are the anhydrides of acids, it is sometimes regarded as a non-metallic element. It is, however, of a highly lustrous metallic appearance, steel-gray in color, notably crystalline in structure, heavy, and so very brittle that it is easily reduced to a powder. 17. Antimony combines energetically with chlorine, bromine, and iodine, in contact with all of which, when finely powdered, it quickly takes fire. Upon bromine, sufficient heat is generated to melt the antimony, and it spins around as does sodium upon water, burning all the time. At ordinary temperatures, the metal does not readily tarnish in the air, but by means of the oxidizing blowpipe flame it is converted into a white oxide, Sb 2 O 3 . It is only slightly acted upon by hydrochloric acid, but nitric acid converts it into a white powder, and it is readily soluble in aqua regia, forming antimony chloride, 292 MODERN CHEMISTRY SbCl 3 . One of its most valuable properties is that of expanding somewhat upon solidifying. EXPERIMENT 177. To illustrate some of the above-mentioned properties. Take the metallic bead obtained in the preceding experi- ment, and learn whether it is magnetic. Test it with a hammer on an anvil to learn whether it is malleable. Notice its color and appear- ance. Put a portion of it on charcoal and try the oxidizing flame. What are the results? How does it differ from arsenic treated thus? Test the solubility of the metal in nitric acid; in aqua regia. State results in each case. Boil nearly to dry ness the latter solution, and add water. What happens ? Treat this with tartaric acid, and state results. 18. Uses. Because of its property of expanding when it solidifies, antimony is used very extensively in mak- ing type metal, britannia ware, and other similar alloys. Antimony may be obtained in a powdered or amorphous condition by immersing a strip of zinc in a solution of some antimony salt, as the chloride or tartrate. The principle underlying is the same as that in the second method of reducing the ore, described already. This antimony black, as it is called, is a dark-colored, finely divided powder, and is sometimes used in giving plaster figures a metallic appearance. Compounds of Antimony 19. There was a time when the compounds of antimony were extensively employed in medicine, but owing to their exceedingly poisonous character, their use was prohibited by law, and their applications now are considerably limited. 20. Stibine, Antimoniureted Hydrogen, SbH 3 . This gas, known also as hydrogen antimonide, corresponding to similar compounds of arsenic and phosphorus, is usually ARSENIC, ANTIMONY, BISMUTH 293 prepared from nascent hydrogen and some antimony com- pound, just as arsine was prepared in Experiment 172. It is a combustible gas, which burns with a green flame, and deposits upon a cold dish held against this flame a black spot resembling that of arsenic, but not so lustrous. It is also less volatile if heated, and is insoluble in a solu- tion of calcium or sodium hypochlorite. EXPERIMENT 178. Prepare stibine exactly as you did the arsine, using the same precautions. Test the spots with a solution of bleach- ing powder or sodium hypochlorite, and verify the statements made above. 21. Oxides of Antimony. None of the three oxides of antimony is of any importance. The trioxide, Sb 2 O 3 , and pentoxide, Sb 2 O 5 , are the anhydrides of the acids, antimo- nous and antimonic, corresponding to those of nitrogen from the similar oxides. Sb 2 3 + 3 H 2 = 2 H 3 SbO 3 . Sb 2 5 + 3 H 2 = 2 H 3 Sb0 4 . 22. The Chlorides of Antimony. When antimony is dissolved in aqua regia, as in Experiment 177, above, and the solution evaporated, antimony trichloride, SbCl 3 , a white crystalline salt, is obtained. It was formerly known as "butter of antimony," from the thick oily appearance which it assumes before solidifying. Upon adding water to this compound, a white precipitate is formed, which is known as basic antimony chloride, or antimony oxychlo- ride, SbOCl. The reaction may be expressed thus : SbCl 3 + H 2 = SbOCl + 2 HC1. The trichloride has given up two atoms of its chlorine, and has taken in their place one atom of bivalent oxygen. 294 MODERN CHEMISTRY This oxychloride is soluble in tartaric acid, but not in water. 23. Antimony Trisulphide, Sb 2 S 3 . This is obtained arti- ficially by passing a current of hydrogen sulphide through an antimony solution. It is of a beautiful orange color, soluble in yellow ammonium sulphide, and also in strong hydrochloric acid. BISMUTH: Bi = 208 24. Source of Supply. Most of the commercial supply of bismuth is obtained from Saxon}'. It is usually found free, but alloyed with small quantities of several other metals. It also occurs in two ores : the sulphide, Bi 2 S 3 , known as bistnuthite, and the oxide, Bi 2 O 3 . 25. Reduction. When obtained from native bismuth, as it usually is, the process consists of little more than simply heating to melt the bismuth ; the other metals found with it have a higher melting point, and remain unchanged. In the case of the ores, if bismuthite is used, it is treated as the sulphides of other metals are, first con- verted into an oxide, and then heated with charcoal. Let the student write the reactions representing the two steps. 26. Characteristics. Like antimony, bismuth is a hard, brittle, distinctly crystalline metal. It is steel-gray in color, having somewhat of a golden reflection, or upon some surfaces a purplish hue. It has a low melting point, being just above tin in this respect, expands upon solidi- ARSENIC, ANTIMONY, BISMUTH 295 fying, and is permanent in -the air at ordinary tempera- tures ; at a red heat it oxidizes to a light yellow powder. It unites readily with bromine and chlorine, and if sifted into them takes tire at once. EXPERIMENT 179. If no bismuth is to be had in the laboratory, prepare a little by heating bismuth nitrate, mixed with sodium carbo- nate ami moistened, on charcoal with the reducing flame. Note the color of the metallic bead; test its hardness and mallea- bility, and learn whether it is magnetic. Dissolve a portion of the bead obtained in nitric acid, boil nearly dry, and add water. What forms? Treat with tartaric acid in solution. Compare results with similar tests with antimony. How do they differ? 27. Uses. In the metallic form bismuth has but little use, except in alloys. To these it imparts the properties of low fusing points and of expansibility. For these reasons it is used in stereotyping, and for similar purposes -where clearly defined copies are demanded. Bismuth is also used for making safety plugs in boilers, and for very fusible alloys, sucli as Wood's alloy, which melts at about 60 C. EXPERIMENT 180. Put into an iron spoon about 2 g. of bismuth, 1 g. of lead, and 1 g. of tin, and melt them. When cold put into a beaker of boiling water. What happens ? 28. Most of the bismuth produced at the smelters is converted into its compounds and used in a medicinal way. Compounds of Bismuth 29. Two Classes of Compounds. Like antimony, bis- muth forms two classes of compounds. These may be rep- resented by the nitrates ; the ternitrate, Bi(N O 3 ) 3 . in which it is readily seen that the bismuth atom is tnvalent, and the basic nitrate, BiONO 3 , in which one atom of oxygen 296 MODERN CHEMISTRY has replaced two of the groups of N0 3 . It may be graphically shown as follows : Ternitrate. Basic (Bisniuthyl) . 30. The first of these is a white crystalline salt, which is prepared by dissolving metallic bismuth in nitric acid,, It has little use, except in the preparation of other com- pounds of bismuth. The basic nitrate, sold at drug-stores as the subnitrate, or simply as "bismuth," is a white powder, obtained from the ordinary nitrate by the addi- tion of water, whereupon a fine white precipitate falls, thus : Bi(NO 3 ) 3 + H 2 = BiONO 3 + 2 HNO 3 , or more properly, considering the water of crystallization, Bi(N0 3 ) 3 , 2 H 2 + H 2 = BiON0 3 , H 2 O + 2 HNO 8 +H 2 O. This is used largely as a cosmetic, and for relieving the irritation of chafed or chapped skin ; also in cholera and kindred diseases, and in acute dyspepsia. 31. Bismuth Trioxide, Bi 2 3 . This is also called bis- muth ocher, the chief ore of bismuth, but may be obtained artificially by heating the metal in the oxidizing flame. It is of a deep yellow color when hot, but yellowish white when cold. Its principal use is as a paint. 32. Bismuth Trichloride, BiCl 3 . This may be prepared by heating bismuth in chlorine gas. If water is added to it, the basic bismuth chloride, or oxychloride, BiOCl, is formed, as is the case with antimony. The latter, how- ARSENIC, ANTIMONY, BISMUTH 297 ever, is soluble in sodium tartrate or tartaric acid, but the former is not. Basic bismuth chloride is a fine white powder, and is used as a paint, known as "pearl white." 33. The Nitrogen Group. From the similarity of their compounds, and their chemical affinity, nitrogen, phos- phorus, arsenic, antimony, and bismuth are often classed together and called the nitrogen group. The following table will give a comparative view of their more impor- tant compounds : N=*14 P = 31 As = 75 Sb = 120 Bi = 208 NlTROGEX PHOSPHORUS ARSENIC ANTIMONY BISMUTH NH 3 PH 3 AsH 3 SbH 3 N,0, NA P 2 o 3 PA As 2 3 As 2 5 AsCl 3 Sb 2 3 Sb 2 5 SbCl 3 SbOCl Bi 2 3 Bi 2 5 BiCl 3 BiOCl SUMMARY OF CHAPTER Comparative Study of Arsenic, Antimony, and Bismuth. Sources of the metals. Wherein alike. Wherein different. Reduction of the ores. Wherein similar How similar to reduction of other metallic ores. In what respects different. Description of experiments illustrating methods. Characteristics of the group. Compare two of them with the non-metals. Wherein are they all metallic in character. Compare in Color. Melting point. Density. Tendency to oxidize. Hardness. Solubility in acids. Malleability. 298 MODERN CHEMISTRY State any special characteristics, not common. Compare bismuth and antimony as to certain classes of salts formed by each. Compare arsenic and antimony in the same way. Uses of each. Special use for metallic arsenic Reason. Same for antimony, and reason. Antimony black How made? Use? Same for bismuth, and reason. Compounds. Compare the hydrogen compounds of arsenic and antimony as to Method of preparing and chemical action. Characteristics of each. How distinguish one from the other? Products formed when Asll 3 burns. Experimental proof. Oxygen compounds. Names and formulae. Important one of arsenic Why ? Appearance and uses. Physiological action Compare with antimony. Antidotes. Solvents. Appearance and use of bismuth oxide. Sulphides. Names and formulae. Method of preparing. Appearance of each. How distinguish As 2 S 3 from SnS 2 ? How distinguish As 2 S, from Sb^? How distinguish an antimony salt from one of bismuth? Special compounds. Paris green Experimental preparation. Appearar.ce Uses. Butter of antimony Chemical name and formula. Means of identifying. For arsenic Marsh's test ; Scheele's test. For bismuth and antimony. Comparison of the nitrogen group. Compounds with hydrogen, oxygen, chlorine, etc. CHAPTER XXV IRON, NICKEL, COBALT IRON: Fe = 56 1. Distribution. Iron, the most useful of all metals, is also the most abundant and most widely distributed. It is found in nearly all clays and soils, and from these is taken up by plants, and through them makes its way into the animal economy. The color of many soils, rocks, and minerals is due to the presence of iron in some form. Pure iron does not occur in any considerable quantities, except in meteorites, of which some weigh many tons. The largest meteorites ever found were discovered by Lieutenant Peary in his Arctic explorations. One of these, weighing nearly one hundred tons, was brought back and placed in the Brooklyn Navy Yard. Meteorites are found to consist of iron, about 93 per cent, and nickel, 7 per cent. 2. Iron Ores. A large number of iron ores are known, among which are the following: 3. Magnetite, Fe 3 4 . This is also known as lode-stone, on account of its magnetic properties. 4. Hematite, Fe. 2 3 . This ore received its name from the Greek word for Hood, because of the red streak it gives on porcelain. It is a very abundant ore : two knobs, Iron Mountain and Pilot Knob, of the Ozark Range in Missouri, consist almost entirely of hematite, in masses ranging all the way from "the size of a pigeon's egg to that of a medium-sized church." 299 300 MODERN CHEMISTRY 5. Limonite, 2Fe 2 3 , 3H 2 0. An ore resembling hema- tite, which gives a yellow streak on porcelain. 6. Siderite, FeC0 3 ; Spathic Iron Ore. This is common in some localities', has a gray to brownish red color, and often contains manganese. 7. Iron Pyrites, FeS 2 , is a very abundant ore, but on account of the difficulty experienced in reducing it, it is not used, except for the manufacture of sulphuric acid. It is commonly known as "fool's gold." 8. Reduction of Ores. This is accomplished in a blast furnace, the essential features of which are shown in Figure 61. The furnace is from 50 to 75 feet in height, supported by masonry, and strengthened with boiler plate. It is lined inside with fire-brick. Near the bottom some pipes, PP, enter the furnace. These are called tuyeres, or blast-pipes, and furnish a powerful blast of air. Just below these is an opening, S, where the slag is drawn off, and below this another opening, /, for drawing off the iron. 9. The furnace is charged from the top ; first, wood for kindling being placed in the bottom, then alternate layers of coke and iron ore mixed with limestone. These are all dumped upon the cone-shaped top, (7, which fits air tight and works automatically. When the ore and other mate- rials fall upon the top, it lowers mechanically and allows the charge to roll into the furnace. Many of the gases formed in the interior of the mass are combustible, and are conducted off through the pipe M, and are burned in other furnaces. FIG. 61. IRCXZV, NICKEL, COBALT 301 Figure 62 shows a perspective view of the blast furnace. The various materials are lifted to the top by elevators ; FIG. 62. Perspective View of Blast Furnace. the molten iron is drawn off and molded in trenches in the ground under the shed. 802 MODERN CHEMISTRY 10. In accordance with the usual method of reducing metallic ores, the oxides of iron are mixed with coke and limestone and strongly heated. If an ore, not an oxide, is used, it is first calcined to convert it into an oxide. The coke serves as a deoxidizing agent, and the lime- stone, used as a flux, is melted, and renders the iron ore more fusible. The limestone then combines with the silica always present in the ore, and forms a molten glass, or slag, which floats upon the iron and prevents its oxida- tion by the strong currents of air. The iron thus obtained, nevertheless, absorbs in the intense heat of the furnace considerable quantities of sulphur, phosphorus, carbon, and silica, and in this impure form is not suited to many of the numerous demands for iron. From the blast fur- nace it is run off through trenches into molds, 2 to 4 feet long, which are called "pigs," and the iron is known as cast or pig iron. It is very brittle, coarse grained, and contains from 5 to 10 per cent of impurities. 11. Wrought Iron. This variety is prepared in what are known as puddling furnaces. In these the low arch- ing roof deflects the flames down- ward upon the broken cast iron. The furnace is lined with ferric oxide, Fe 2 O 3 , and as the cast iron melts, the carbon which it contains combines with the oxy- gen in the lining. By stirring the molten mass, or puddling, as it is called, the whole is gradually purified, until finally, as it is much more difficult to melt pure iron, the whole mass becomes pasty. This pasty mass, bloom, as it is called, is then removed and hammered with trip-hammers, FIG. 63. IRON, NICKEL, COBALT 303 a process which drives out any remaining slag, and renders the iron malleable. 12. Steel. Formerly steel was made from wrought iron by embedding bars of the latter in finely powdered charcoal and keeping at a red heat for about ten days. During this time the bars of iron slowly absorbed more or less carbon, and were converted into steel. Besides the expense and the length of time required in this process, there were two other serious objections to it : first, that there was no possible way of controlling accu- rately the amount of carbon taken up by tha iron, and second, a steel bar was obtained which was not at all uniform in quality and texture. 13. Present Method of Manufacture. At present, steel is made directly from cast iron, by the Bessemer process. An egg-shaped vessel, called a converter, is used. It is securely bound with boiler iron, and is lined with ganister, a siliceous earth, fusible only at a very high temperature. The converter, which will hold ten or more tons of iron, is mounted on axes, or trunnions. One of these, A, in Fig. 64, is hollow, so that a blast of air may be forced through it when the converter is in a vertical posi- tion. This trunnion opens into a pipe, P, which extends down the outside of the converter and opens into the tuyere box, B, beneath the body of the converter. Through the tuyere box, numerous small openings admit the air to the converter and its con- tents. FIG. 64. A Converter. 304 MODERN CHEMISTRY 14. Bessemer Process. About ten tons of pig iron are placed in a cupola furnace, that is, one resembling a* blast furnace in most of its details, but considerably smaller. When the iron is melted, it is run into the converter. Immediately the blast of air is turned on, and, bubbling up through the molten iron, the oxygen unites with the carbon and other impurities, burning them out. No heat is used in the operation except what is evolved by the combustion of the impurities themselves. About twenty- five or thirty minutes are required for the completion of the operation, during most of which time a brilliant shower of sparks is thrown from the mouth of the con- verter. This is represented in colors by the frontispiece. The converter on the right is shown in action ; the one on the left, at the close of the process, discharging the molten steel into a pot, from which it will be poured into molds. 15. When the mass of flame and sparks no longer issues from the converter, the workmen know that the cast iron has had its impurities entirely removed, and is now wrought iron. Next, a Aveighed quantity of spiegeleisen or manganese iron, containing a known amount of carbon, is thrown into the converter, and in a moment or two the process is complete. In this way, in thirty minutes or less, ten tons of steel are obtained at a cost only a fraction of what it would be by former methods. 16. Basic-lining Process. If the iron ore contains much phosphorus, the converter is lined with lime- stone, which, during the process of oxidation, takes up the phosphorus from the iron, and is converted into calcium phosphate. This is known as the basic-lining process and was put into practical use by the inventors, Thomas and Gilchrist. IRON, NICKEL, COBALT 305 17. Tempering Steel. Tempering consists in harden- ing steel, by heating and then suddenly plunging into cold water or oil. Tempered in this way, it becomes much less malleable, but can take and hold a sharp edge. Different instruments require steel that has been heated to different temperatures ; thus, surgical instruments after being hardened are again heated to about 225 or till a yellow film of oxide appears upon the surface. For ordinary cutlery, a temperature of about 250 is used, indicated by the appearance of a brown film, while clock and watch springs and such forms as require great elas- ticity are made of steel heated till blue, or about 290. By heating any form of steel strongly and then cooling very slowly, the temper is " drawn," or removed, and the metal becomes like ordinary wrought iron. 18. Comparison of the Three Forms. CAST IKON. STEEL. WROUGHT IRON. Impurities 5 to 10%. 1 to 2 %. 0.36 to 0.5 %. Brittle. Somewhat malleable. Very malleable. Coarse grained. Fine gf%ined. Very fine grained. Cannot be tempered. May be tempered. Cannot be tempered. Lowest melting point. Medium melting point. Highest melting point. 19. Uses of Iron. This is preeminently the " Steel Age." Day by day the uses of iron are increasing. The continual cheapening of both steel and wrought iron by improved methods has caused their use in thousands of ways where wood was formerly demanded. These applica- tions are too well known, however, to need mentioning. Compounds of Iron 20. Ferrous and Ferric Compounds. Like several other metals, iron forms two general classes of compounds, the ferrous and ferric. The former are very unstable, and 306 MODERN CHEMISTRY when exposed to the air gradually change to the ferric. The reaction in the presence of free acid may be indicated thus : 2 FeSO 4 -h H 2 SO 4 + O(air) = Fe 2 (SO 4 ) 3 H 2 O. If there is no free acid present, a part of the ferrous salt is converted into the ferric and another part into an insoluble basic salt. EXPERIMENT 181. To distinguish between ferrous ami ferric units. Pulverize a crystal of ferrous sulphate and dissolve in a feu cubic centimeters of water; divide into three portions. To one portion add promptly a lew drops of ammonium hydroxide, to the second a few drops of potassium sulphocyanide solution, to the third a few drops of potassium ferrocvani.le solution. Notice the results in each case. Now dissolve a little ferric chloride or nitrate in water, divide into three parts, and repeat the same tests. Compare results and tabulate as below. NII 4 OII KSCy K 4 FeCy G FeSO, Fe.C! G EXPKRIMKNT 182. To show the instability of ferrous salts. Quickly dissolve in cold water a little powdered ferrous sulphate, and divide into two parts. To one add a few drops of ammonium hydroxide, and note the color of the precipitate. Allow both portions to stand for some time. How does the greenish precipitate change in color? Into what is it apparently converted? Has the other portion changed any in appearance? How? Test it with potassium sulphocyanide or ammonia to learn what kind of a salt it is now. What are your conclusions? 21. This experiment will show the tendency of ferrous salts. What is thus accomplished slowly by the action of IRON, NICKEL, COBALT 807 atmospheric oxygen at ordinary temperatures is effected rapidly by nitric acid at the boiling point. As already seen, this acid is a strong oxidizing agent, readily giving up a part of its oxygen when heated, thus : 2 HN0 3 + (heat) = H 2 O + 2 NO 3 + O. This nascent oxygen rapidly attacks any oxidizable substance that may be present. With ferrous chloride in the presence of hydrochloric acid, the following reaction takes place : * **, f (s^f * #? t^* EXPERIMENT 183. To show the effects of nitric acid upon a ferrous sail. Dissolve a little ferrous sulphate in water, add a few drops of sulphuric acid and then some nitric acid, and heat to the boiling point for two or three minutes. Does the solution change any in color? Now test a part of it in two or three ways to learn whether it has been converted into a ferric salt. What are your conclusions? 22. Ferric salts, on the other hand, may be reduced to the ferrous by treatment with hydrogen sulphide or nascent hydrogen. The reaction may be shown thus : Fe 2 Cl 6 + H 2 = 2 FeCl + 2 HC1. ^ EXPERIMENT 184. To prove the above statement. Put into a test- tube about 5 ec. of a solution of ferric chloride or nitrate and drop into it a good-sized granule of zinc. Now add a little strong sulphuric or hydrochloric acid to cause a rapid evolution of hydrogen. In from 5 to 7 minutes the yellow color of the ferric solution should have disap- peared. Test with ammonia or potassium sulphocyanide. What are your conclusions in the matter ? 23. How to distinguish Ferrous from Ferric Salts. From the preceding work it will be learned that ferrous salts in solution are usually colorless or very pale green, while ferric salts are light brown. Potassium sulpho- 308 MODERN CHEMISTRY cyanide serves as the most delicate method of detecting a ferric salt, because even exceedingly small quantities will show the characteristic wine-red color ; with ferrous salts, however, it shows no reaction, hence will not indicate their presence. Ammonia gives precipitates with both classes of salts, deep reddish brown with the ferric, and greenish with ferrous. Potassium ferrocyanide and ferricyanide may also be used to distinguish between the two. 24. Sulphates of Iron. Ferrous sulphate, FeSO, 7 N 2 0, the only common ferrous salt, is formed when iron is dis- solved in sulphuric acid. It is commonly known as cop- peras or green vitriol, and occurs in light green crystals. It is somewhat efflorescent, and gradually gives up its water of crystallization, turning white and breaking up into a powder, anhydrous ferrous sulphate. It is used con- siderably in making black ink and dyes, also as a deodorizer and disinfectant. EXPERIMENT 185. To show one method of making ink. Prepare a strong solution of ferrous sulphate, and add to it a little of another solution made by soaking some powdered nutgalls in water. Notice the bluish black color obtained. Allow it to stand a few minutes, and notice whether the color deepens. Now add to the solution a few drops of a solution of oxalic acid. What happens? This suggests a method for removing ink stains without injuring the fiber of the paper or cloth. 25. Ferric Chloride, Fe 2 Cl 6 . This is a brownish yel- low salt, which rapidly absorbs moisture when exposed to the air. It is obtained when iron is treated with aqua regia or dissolved in hydrochloric acid with the addition of a crystal of potassium chlorate. It has little use except in the laboratory. 26. The Sulphides. Ferrous, FeS ; Ferric, .Fe 2 S 3 . The former is a dark gray substance somewhat resembling cast IRON, NICKEL, COBALT 309 iron. It is made by fusing together, in the proportion of their atomic weights, iron and sulphur, and is used exten- sively in the laboratory for making hydrogen sulphide. Ferric disulphide is the native ore, pyrite, or " fool's gold." It is of a brassy yellow color, and frequently occurs in beautiful cubes or modified forms of the cube. It is very abundant, but has little use except in the preparation of sulphur dioxide for the manufacture of sulphuric acid. 27. The Oxides. Ferric, Fe 2 3 . This is met with in the ore, hematite, already mentioned. It is also formed when iron is exposed to moisture, and is known as rust. In the hydrated form, Fe 2 (OH) 6 , ferric hydroxide, it is formed when any ferric solution is treated with ammonia. As a reddish brown precipitate it has already been seen in several of the preceding experiments. It is sometimes used as an antidote for arsenic poisoning. Magnetite, Fe 3 O 4 , is regarded as ferrous ferrite, Fe(FeO 2 ) 2 , a salt of ferrous acid. Compare Pb 3 O 4 . The greenish precipitate obtained in some of the preceding experiments by adding ammonia to a solution of ferrous sulphate is ferrous hydroxide, Fe(OH) 2 , or FeO, H 2 O; that is, the hydrated form of the protoxide, FeO. NICKEL: Ni = 58.7 28. Distribution. Like iron, nickel is never found pure except in meteorites, of which, as already stated, it often constitutes from 5 to 7 per cent. Its ores are fairly well distributed, but are nowhere in great abundance, and with them are always associated cobalt and iron. 29. Characteristics of Nickel. Nickel is a silvery white metal with the faintest yellow tinge ; it is susceptible of a very high polish and does not tarnish in the air. It is 310 MODERN CHEMISTRY very hard, melts at about white heat, may be welded like iron, is magnetic, and becomes brittle like cast iron upon the addition of such impurities as cast iron always con- tains carbon and silicon. Its density is but little greater than that of iron. It is soluble in nitric acid. In most respects, therefore, it is very similar to iron, and strikingly different in one respect only. 30. Uses. Nickel is used very extensively in alloys, among them being certain coins; in gernian silver, con- sisting of nickel, zinc, and copper; and with steel for armor plating in making what is known as Harveyized steel, noted for its hardness and toughness. Nickel is also used largely in plating various articles of ornament and utility. 31. Compounds. The general color of the more com- mon i ickel salts is green. Among these may be named the nitrate, Ni(NO 3 ) 2 , chloride, NiCl 2 , sulphate, NiSO 4 ; also Ni(OH) 2 , nickel hydroxide. This last may be pre- pared from a solution of any of the preceding salts by adding a few drops of ammonia or caustic potash. EXPERIMENT 186. To 3 or 4 cc. of a solution of any of the above salts, add a little caustic soda. Describe the precipitate that forms. Te>t its solubility in hydrochloric acid. Write the two reactions taking place. 32. A fifth compound which may be mentioned is the sulphide, NiS. It is prepared, as is the sulphide of other kindred metals, by adding ammonium sulphide to a neutral or alkaline solution of any nickel salt. EXPERIMENT 187. Add a little ammonium sulphide to 4 or 5cc. of a solution of any nickel salt. Describe the nickel sulphide that forms. Test its solubility in hydrochloric acid. Also in aqua regia. Write the reactions. IRON, NICKEL, COBALT 311 33. Nickel salts fused in a borax bead impart to it a smoky yellow or brown color according to the amount of the nickel present. EXPKRIMEXT 188. Make a small loop in the end of a platinum wire, heat it in the burner flame and dip into some powdered borax. Kow hold again in the flame until the borax which swells up at first has formed a clear transparent glassy bead. Dip into a solution of some nickel salt, or touch it to a tiny particle of nickel salt and fuse again. If you use the solution, it may be necessary to dip the bead several times. Note the color imparted. EXPERIMENT 183. To fml ihe composition of a coin. Put a "nickel" into an evaporating dish and treat with warm nitric acid for a few minutes. Remove the coin and add a few cubic centimeters of water. Pass a current of hydrogen sulphide through the solution for several minutes, and filter out the black precipitate. After \\ashing it, punch a hole in the bottom of the filter and wash the precipitate through into a beaker with a little nitric acid. Heat until it dissolves. What colored solution is obtained? What metal is indicated by this color ? Add ammonia till alkaline ; is a deeper blue solution obtained V What metal is it? Boil nearly to dryness the filtrate from the black precipitate above ; note the color. Does this indicate any salts with which you are familiar? Make a borax bead as directed in the preceding section and test the solution ; what are your conclusions? Of what two metals is the coin composed? If you can obtain one of the lighter-colored pennies seen occasionally, test it in the same way. COBALT: Co = 59 34. Characteristics. This is a somewhat rare metal that is usually found associated with nickel. It is very similar to iron and nickel in its characteristics, being steel-gray in color, very hard, magnetic, and of about the same melting point. It is permanent in the air. The metal itself has no application in any of the arts. 35. Compounds. Cobalt forms salts with the three common acids ; the nitrate, Co(NO 3 ) 3 ; cJiloridt, CoCl 3 ; 312 MODERN CHEMISTRY and sulphate, CoSO 4 . These are all some shade of red in color, but when heated so as to lose their water of crystallization they become blue. 36. The hydroxide and sulphide are prepared just as the similar compounds of nickel are. EXPERIMENT 190. Prepare the last two as you did the correspond- ing compounds of nickel in Experiments 186 and 187, and test their solubility in the same way. 37. None of the above has much use except occasionally in the laboratory. There are one or two others, however, which have extensive application in the arts. Among these may be named 38. Smalt, a silicate of cobalt. When fused with glass or pottery ware this imparts a beautiful blue color, and is largely used for that purpose. It may be illustrated in the following experiment : EXPERIMENT 191. Prepare a borax bead as you did for nickel, and fuse with it some salt of cobalt. Note the color imparted. If it is too dark to recognize, it is because too much cobalt has been intro- duced. Break out the bead, and repeat the experiment, using less of the compound. 39. Sympathetic Inks. They are inks which under ordinary circumstances are invisible, or nearly so, on paper ; when heated or treated by some other method they become legible. Many of these have some compound of cobalt as their basis. EXPERIMENT 192. Mix a solution of some compound of cobalt with one of ferrous sulphate. Using this as an ink, write with it upon paper, and when the inscription is dry heat it. Do you obtain a dis- tinct green color, though before it was nearly invisible ? In the same way try potassium iodide mixed with the cobalt solution. Results? IKON, NICKEL, COBALT 313 SUMMARY OF CHAPTER Iron, Nickel, Cobalt. Occurrence Wherein are iron and nickel similar. History of some large meteorites. Some important ores of iron. Names and formulae. Localities where found. Reduction of iron ores Description of blast furnace. Drawing of essential features. Method of charging the furnace. Chemical action that takes place. Plan of molding pig iron. Varieties of iron Three. How different in composition and properties ? Description of the puddling furnace. Chemical action. Meaning of term bloom. Description of the converter. Explanation of the chemical changes. Plans used for phosphorus-bearing iron ores. Tempering steel. Meaning of the term. Process used. Characteristics Compare iron and nickel as to Color. Susceptibility of polish. Hardness. Permanency in the air. Melting point. Several other similarities. Magnetic properties. Uses Compare nickel and iron. ' Compounds Classes of iron compounds. Compare them as to stability. Plans for distinguishing the two. Method of converting each into the other. Explain the chemical action in each case. Names of three or four compounds of iron and their uses. Compare the compounds of nickel and cobalt in color and method of preparation. CHAPTER XXVI THE PLATINUM GROUP PLATINUM : Pt = 195 1. Where obtained. Platinum is a rare metal, usually found uncombined, but almost always associated with iridium, and smaller quantities of palladium and osmium. The greater portion of the commercial supply comes from the Ural Mountains in Russia, though small quantities have been obtained in California, Arizona, and some parts of South America. 2. Characteristics. Platinum is a hard, silvery white metal, unaffected by the air at any temperature. It is somewhat malleable, but becomes less so if alloyed with a small per cent of iridium, though by this means its hard- ness is increased. It is a very dense metal, with a specific gravity of 21.5, -osmium, the heaviest metal known, having a density of only 22.5. The melting point of platinum is about 1900 C., and it can be fused only by such intense heat as that of the oxyhydrogen blowpipe, or acetylene blast lamp. Like gold, it is soluble only in nitre-hydro- chloric acid, forming therewith platinic chloride, PtCl 4 . 3. Property of occluding Gases. The most remarkable property of platinum is that of occluding or absorbing various gases within its pores. It is estimated that at ordinary temperatures it will absorb 200 times its own volume of oxygen. In the spongy form, that is, when finely divided, as in the case of a metallic precipitate, the power of 314 THE PLATINUM GROUP 315 absorption is especially striking. If a current of hydrogen be directed against the platinum sponge, so rapid will be the absorption that almost instantly the metal will become red hot, and in two or three seconds the jet will be ignited. EXPERIMENT 193. Repeat Experiment 23 with hydrogen. 4. If into ajar of hydrogen and oxygen, mixed in the proportion of two to one, a platinum sponge be introduced, the gases will be made to unite with explosive violence. This power of occlusion may be seen in the case of certain other gases, as ammonia and common coal gas. EXPERIMENT 19i. Support upon a ring-stand a small flask con- taining some strong aqua ammonia; warm it gently so as to secure a constant and rapid evolution of gas from the liquid. Now heat to bright redness in the Buusen flame a spiral of platinum wire, made by coiling it about a small glass rod, and hold it in the neck of the flask. The wire will continue to glow, and the intensity of the heat will often be increased. Take this same platinum coil and flatten it a little so as to bring the parts of the spiral closer together; hold it in the Bunsen flame until red hot, then turn off the gas. When the redness has just disappeared from the wire, again turn on the gas. The wire will quickly grow red again, and in two or three seconds will re-ignite the escaping gas. This may be repeated over and over again. The same may be tried with a spirit lamp. 5. Platinum Alloys. Platinum readily alloys with lead, silver, antimony, and other metals which are easily reduced from their compounds ; hence it should never be strongly heated in contact with them. It is likewise injured by heating in a smoky flame, or by placing upon red-hot charcoal, which blisters the surface. Platinum vessels are usually cleaned by fusing in them for a few minutes some acid potassium sulphate, KHSO 4 , and are polished by rubbing gently with a little fine sea-sand. 316 MODERN CHEMISTRY 6. Uses. The rarity of the metal and the long, com- plicated processes involved in preparing it in the pure form, make it almost as expensive as gold. It is worth from 50 cents to 75 cents per gram, or about $300 per pound. It is made into wire, foil, and various articles for use in the chemical laboratory, such as crucibles, dishes, tips of forceps, etc. To the chemist it is simply indispensable in analytical work. SUMMARY OF CHAPTER Names of the elements in this group. Source of the supply of platinum. Characteristics of platinum. Experiments that illustrate these. Alloys of platinum. Uses and value of the metal. Compare with metals studied previously as to Color. Melting point. Density. Tendency to oxidize. Power of occluding gases. Solubility in acids. CHAPTER XXVII CHROMIUM AND ITS COMPOUNDS CHROMIUM : Cr = 52 1. Where found. Chromium is a rare metal which received its name from the Greek word, chromos, meaning color, and is so named because of the striking colors of most chromium compounds. It occurs chiefly in the Shetland Islands in the form of chromite, or chrome iron, Cr 2 O 3 , FeO, also written FeCr 2 O 4 . It is also found as crocosite, PbCrO 4 , in Siberia, Brazil, and the Philippine Islands. Compounds of Chromium 2. Classes. In the metallic form chromium has but little use. Its compounds, however, have various applica- tions. They may be divided into two important classes : 3. Chromium as a Basic Element. Those in which chromium acts as a basic element, with the power of replacing the hydrogen in acids to form salts. Of these, as in the case of iron, mercury, and others, there are two divisions, the chromous and chromic, but only the latter are important. As examples, we have chromic chloride, CrCl 3 , chromic nitrate, Cr(NO 3 ) 3 , and chromic sulphate, Cr 2 (SO 4 ) 3 . These as a rule are green in color, but the double sulphate of potassium and chromium, K 2 Cr 2 (SO 4 ) 4 , is violet. Solutions of the chromic salts are precipitated by caustic potash or ammonia, giving the hydroxide, Cr(OH) 3 . 317 818 MODERN CHEMISTRY 4. Chromium as an Acid Producer. Those in which chromium serves as a noil-metallic element, forming acids. Of these there are three classes, but only two merit notice, the chromates and the die hro mates. The chronmtes are based on the theoretical chromic acid, H 2 CrO 4 , wherein the chromium atom is that which distinguishes the acid, as does the sulphur in sulphuric acid, H 2 SO 4 . The general color of the chromates is yellow, though there are some exceptions. The best-known example is potassium chro- mate, K 2 CrO 4 . EXPERIMENT 195. To prepare some other chromates. To 3 or 4 cc. of a solution of potassium chromate in a test-tube add a few drops of lead nitrate or acetate in solution. Notice the color of the lead chro- mate formed. In the same way prepare some barium chromate by us- ing barium chloride with the potassium chromate. Compare its color with the preceding. Now prepare some silver ehromate by using silver nitrate solution with the potassium chromate. Note its appearance. 5. Potassium Bichromate, K 2 Cr 2 O r orange-red in color, is the best-known example of the dichromates. Tabular view of the compounds : I. Chromium as a true metal : 1. Chromous. 2. Chromic a. Chloride, CrCl 3 . b. Nitrate, Cr(NO 3 ) 3 . c. Sulphate, Cr 2 (SO 4 ) 3 . II. Chromium as an acid former : 1. Chromates a. Potassium, K 2 Cr0 4 . b. Lead, PbCrO 4 . c. Barium, BaCrO 4 . 2. Dichromates a. Potassium, K 2 Cr 2 O 7 . CHROMIUM AND ITS COMPOUNDS 319 6. Conversion of Ons Class of Compounds into Another. Though the chromium salts are stable, they may easily be converted from one into another. By adding a little acid and passing a current of hydrogen sulphide through a solution of potassium chromate, the latter is changed into a salt of the first class (chromic). The change of color- to green indicates that the reduction has taken place; at the same time free sulphur is precipitated. Thus : 2 K a Cr0 4 + 3 H 2 S + 10 HC1 = 4 KC1 + 2 CrCl, + 8 H 2 O + 3 S. EXPERIMENT 196. To prove the above. Put into a test-tube a few cubic centimeters of a solution of potassium eliminate, and add a little hydrochloric acid. Now pass through the solution a current of hydrogen sulphide. What change in color is noticed? Is the sulphur precipitated? 7. Sulphur dioxide has a like reducing effect upon a chromate solution. EXPERIMENT 197. Put into a test-tube 4 or 5 cc. of a solution of sodium sulphite, Na 2 S0 3 , and a little hydrochloric or sulphuric acid. Notice that sulphur dioxide gas is being liberated. Now add a little potassium chromate or dichroinate. How does the chromium solution change in color? If sodium sulphite is not to be had, fill a bottle with sulphur dioxide gas, pour in the dichromate, and shake. EXPERIMENT 193. To show the reduction of the Hlchromnles to the chromic sail*. Put into an evaporating dish 10 or 15 cc. of a solution of potassium dichromate, add some hydrochloric acid, and boil a few minutes. The addition of a little alcohol will hasten the action. Notice the change in color. What compound of chromium is probably formed? 8. The above experiments prove that either the chro- mates or dichromates may be reduced to salts of the first class. The reaction that takes place in the latter is as follows : K a Cr s 7 + 14 HC1 = 2 KC1 + 2 CrCl 3 + 7 H 2 O + 6 CL 320 MODERN CHEMISTRY EXPERIMENT 199. To 2 or 3 cc. of potassium chromate solution in a test-tube add a few drops of hydrochloric or nitric acid. How does its color change ? What other salt of chromium in solution does it resemble ? In like manner treat 2 or 3 cc. of potassium dichromate solution with a few drops of caustic potash or any alkali. Notice the change in color ; what chemical change has taken place ? 9. It will be seen by the above experiments that the chromates and dichromates may readily be converted, the one into the other. The reactions taking place are shown thus : 2 K 2 Cr0 4 + 2 HC1 = K 2 Cr 2 O 7 + 2 KC1 + H 2 O, and K 2 Cr 2 O 7 + 2 KOH = 2 K 2 CrO 4 + H 2 O. 10. The Oxides of Chromium. Or* 0* and Or 0, chro- & o o~ mium sesquioxide and trioxide. The former is basic in properties, the latter acid. The former is green, and is used in imparting a green color to glass and enamel ; the latter is a dark red crystalline solid. EXPERIMENT 200. Make a borax bead and dip it into a solution of some chromium salt, then fuse in the burner flame. If a good color is not secured the first time, repeat the operation. 11. Chromium trioxide may be prepared by adding strong sulphuric acid to a saturated solution of potas- sium dichromate. After standing for some time, beautiful red needle-like crystals separate from the liquid, thus: K 2 Cr 2 O 7 + H 2 SO 4 = 2 CrO 3 + K a SO 4 + H 2 O. These cannot be filtered out by ordinary methods, as the trioxide is a strong oxidizing agent and readily gives up a part of its oxygen to any organic compound, itself being changed into the sesquioxide, thus : 2CrO 3 =Cr 2 O 3 + 3O. CHROMIUM AND ITS COMPOUNDS 321 12. Chromium trioxide is theoretically the anhydride of chromic acid, H 2 CrO 4 , and seemingly ought to produce it when added to water, thus : CrO 3 + H 2 = H 2 Cr0 4 . But the action is merely one of solution, and the acid is not formed. 13. Chromic Hydroxide, Cr(OH) 3 . This is a green precipitate formed when ammonia or caustic potash is added to any chromic salt, as the chloride or sulphate. CrCl 3 + 3 KOH = Cr(OH) 3 + 3 KC1. 14. Uses of the Compounds. Some of the uses of chro- mium compounds, among others those of the sesquioxide and of lead chromate, have been mentioned. Both the chromate and dichromate of potassium are used as reagents in the laboratory, and in the arts for dyeing and calico printing. If the reaction, K 2 Cr 2 O 7 + 8 HC1 = 2 KC1 + 2 CrCl 3 + 4 H 2 O + 3 O, is studied, it will be seen that potassium dichromate, treated with hydrochloric acid, is a strong oxidizing agent. Each molecule gives up three atoms of oxygen. If no other salt is present, this nascent oxygen unites with the hydrogen in six additional molecules of hydrochloric acid, thus : 6 HC1 + 30 = 3 H 2 + 6 Cl. Combining the last two reactions, it will be seen that we have the one given on page 319, showing the reduction of potassium dichromate to chromic chloride. However, if any oxidizable salt be present, as, for example, a ferrous compound, the nascent oxygen readily converts it from 322 MODERN CHEMISTRY \ the ferrous to the ferric condition. This is shown in the following reaction : 2 FeCl 2 + 2 HC1 + O = Fe 2 Cl 6 + H 2 O. On account of this property, potassium dichromate is fre- quently used by chemists in estimating the amount of iron present in a solution. EXPERIMENT 201. To illustrate this use and the oxidizing power of potassium dichromate. Dissolve a little ferrous sulphate in a few cubic centimeters of water and add some hydrochloric acid. Now add gradually drop by drop a solution of potassium dichromate. Notice how the solution changes to green. Test a portion of it with potassium sulphocyanide and learn whether the solution has been oxidized to the ferric condition. What are your conclusions? Study some of the foregoing reactions and see whether you can determine why the solution became green. SUMMARY OF CHAPTER Origin of the term chromium. Why applied to this metal. Classification of the chromium compounds. Names and formulae of the most important. Relation of the classes of compounds. Method of converting those of second class to first. Indication of the change. Method of converting chromates into dichromates, and vice versa. Compare the two oxides in Appearance. Properties. Commercial uses of certain compounds. Chromium sesquioxide. Chrome yellow. Laboratory uses. What uses as a reagent. How used as an oxidizing agent. Experiment to illustrate. CHAPTER XXVIII MANGANESE AND ITS COMPOUNDS MANGANESE : Mn = 55 1. Where found. This is a somewhat rare metal, often associated with iron ores. The most abundant natural compound is the dioxide, MnO 2 , known as pyrolusite. In the metallic form, manganese has little use, but some of its compounds are valuable. Compounds of Manganese 2. Classes. These may be classified as follows : 3. As a Metal. Those in which manganese acts as a metal, that is, having the power of replacing hydrogen in acids. These may be divided into a. Manganous, b. Manganic, of which only the former are important. The most com- mon of these are manganous chloride, MnCl 2 , and manganous sulphate, MnSO 4 , both crystalline salts, pink in color. From these may be prepared the hydroxide, Mn(OH) 2 , by adding ammonia to a solution of either salt ; also the sulphide, MnS, by adding ammonium sulphide. EXPERIMENT 202. Using a solution of either manganous chloride or sulphate, prepare the hydroxide and sulphide as indicated above and describe their appearance. Test their solubility in hydrochloric acid. State results. 323 324 MODERN CHEMISTRY 4. Manganese Dioxide. In this connection we shall notice the most important of the oxides, MnO 2 , man- ganese dioxide. It is a black compound, and is used in pre- paring oxygen, bromine, chlorine, and iodine. Notice the similarity in method of the last three. MnO 2 + 2 NaCl + 2 H 2 SO 4 = C1 2 + MnSO 4 + Na 2 SO 4 + 2 H 2 O " =Br 2 + " + " + " " =L + " + " + " Cl Br I 5. As an Acid Former. Compounds in which man- ganese serves as an acid-forming element. Of these, there are two classes, a. Manganates, b. Permanganates. The first of these is based upon a theoretical acid, man* ganic, H 2 MnO 4 ; they are not of special interest to us. The best-known example of the second is potassium perman- ganate, KMnO 4 . 6. Potassium Permanganate. This is a dark purple crystalline salt, soluble in water. It is used frequently in the laboratory as a reagent, in a technical way for the estimation of iron in iron ores, and for the testing and purification of cistern water. Like nitric acid and potas- sium dichromate (see pages 88, 321), it is a strong oxidizing agent. When treated with hydrochloric or sulphuric acid, it gives up oxygen, thus : 2 KMnO 4 +3 H 2 SO 4 =K 2 SO 4 + 2 MnSO 4 +3 H 2 O + 5 O. The nascent oxygen thus obtained may be used in oxidiz- ing ferrous salts to the ferric condition, or in destroying MANGANESE AND ITS COMPOUNDS 325 (oxidizing) the organic matter contained in a solution. In the case of the iron the reaction may be shown thus : 10 FeSO 4 + 8 H 2 S0 4 + 2 KMnO 4 = K 2 SO 4 + 2 MnSO 4 + 5 Fe 2 (SO 4 ) 8 + 8 H 2 O ; or, the five atoms of oxygen set free as shown above de- compose five additional molecules of sulphuric acid, thus : 5 O + 5 H 2 S0 4 = 5 H 2 + (SO 4 ) 6 . Then the five (SO 4 ) groups or ions unite with the 10FeSO 4 , forming the ferric salt, 5 Fe 2 (SO 4 ) 3 . Sometimes it is written thus : 10 FeO + 5 O = 5 Fe 2 O 3 , which expresses in a simple form the same change from a ferrous to a ferric condition. EXPERIMENT 203. To show the oxidizing power of potassium per- manganate. To a fresh solution of ferrous sulphate add one or two cubic centimeters of sulphuric acid, and then slowly, drop by drop, potassium permanganate until the solution just begins to turn pink. Now test it with potassium sulphocyauide or ammonium hydroxide. Have you obtained a ferric salt? In this connection study the preced- ing reactions. In the same way test some cistern water that has an offensive odor. Before adding the permanganate heat the water nearly to boiling. Does it lose its odor by this treatment ? In the same way try some cistern water discolored with cedar shingles; is the color removed? Try also a strong solution of logwood; can you remove the dark color? What instances can you give in which nitric acid has served as an oxidizing agent ? Potassium dichromate ? 7. The sulphuric acid is added simply to dissolve a dark- colored precipitate that would otherwise form and obscure the results. In purifying cisterns, of course the acid can- not be used, but the brown solid in a short time settles to the bottom and remains there. The amount of organic 326 MODERN CHEMISTRY matter in cistern water may be learned by measuring the amount of potassium permanganate added before the water begins to turn pink. Sometimes a manganese solution or salt is proved by the color it imparts to the borax bead. EXPERIMENT 204. Prepare a bead as in the case of nickel or cobalt, and fuse with some salt of manganese. Notice the beautiful color imparted. SUMMARY OF COMPOUNDS Class I A true metal in its chemism. Class II An acid-forming -j 2. element. a. Chloride, MnCl 2 . b. Sulphate, MnSO 4 . 1. Manganous c. Hydroxide, Mn(OH) 2 . d. Sulphide, MnS. 2. Manganic, Dioxide, MnO 2 . 1. Manganates, not important. Permanganates, Potassium, KMnO 4 . Compare the above compounds with those of chromium and note the few differences. SUMMARY OF CHAPTER Occurrence of manganese. How associated. Chief ore. Classification of its compounds. Compare with the compounds of chromium, showing wherein similar and wherein different. Uses of certain compounds. Manganese dioxide. Appearance. What laboratory uses. What commercial uses. Potassium permanganate. Appearance. Laboratory uses. Experiments to illustrate. Practical uses. Experiment to illustrate. APPENDIX A QUALITATIVE ANALYSIS IT is not intended in the following pages to give any- thing like a complete system of qualitative analysis. Such would be impossible, keeping within the necessary bounds of a high-school text. As a matter of reference, however, and to meet the demand of any who may care to pursue to some extent this line of work, the following brief outline is offered. The student has noticed already that a reagent which will precipitate some metals from their solutions may have no effect upon various other metals. Taking advantage of this fact, we are able to divide the metals into groups, and then to separate the members of these groups one from another. Accordingly, depending upon the. reagents used for precipitating the metals, five divisions are usually made as follows : Group I Group II 1. Lead 2. Mercury (ous salts) 3. Silver Antimony Tin Arsenic Mercury (ic salts) Copper Bismuth Cadmium Precipitated as chlorides, PbCl 2 , Hg 2 Cl 2 , AgCl, by using hydrochloric acid. Precipitated as sulphides Sb 2 S 3 , SnS or SnS 2 , etc., with sulphureted hydrogen. The first three are soluble in yellow ammonium sulphide or sodium sulphide; the others, not. 327 328 MODERN CHEMISTRY Group III Iron Aluminum Chromium Cobalt Nickel Manganese Zinc Group IV Group V The first three are precipi- tated as hydroxides with am- monia, and constitute division one of this group. The last four are precipitated by am- monium sulphide as sulphides. Precipitated as carbonates, Calcium Strontium Barium j with ammonium carbonate Magnesium j from an alkaline solution. CaCO 3 , SrCO 3 , etc., Lithium Ammonium Sodium Potassium Not precipitated by any com- mon reagents. Most of them usually tested by color impart- ed to flame, or the spectrum. . The General Plan. Suppose now we have a solution which may contain salts of any or all of the above metals. By adding hydrochloric acid, those of the first group would be precipitated and their chlorides separated by filtering. The filtrate would contain the remaining four groups. This would now be treated with hydrogen sulphide, whereby the second group metals may be precipitated and filtered out. In a similar way the separation of the third, fourth, and fifth groups would be effected. All that re- mains is to separate the metals of each individual group and prove their presence by means of some distinctive test. Ionic Theory. A clear understanding of the processes underlying any qualitative analysis is rendered much easier by what is known as the Ionic theory. It has long been observed that certain elements or groups of elements APPENDIX A 329 always give the same distinctive tests with certain re- agents. For example, a silver solution gives the same characteristic precipitate with any soluble chloride, whether it be hydrochloric acid, sodium chloride, or any other. Suppose in analyzing an unknown solution we have found four bases and four acid radicals : each base might have been combined with each of the acid groups, making in all sixteen possible cases. Were we compelled to test for each one of these possible compounds, analysis would be very tedious ; but, as already stated, each base affords the same test as if it existed alone. It seems, therefore, that when substances are dissolved, they become more or less dissociated. For example, hydrochloric acid becomes largely broken up into hydro- gen and chlorine atoms; potassium chlorate into potas- sium, K and C1O 3 , groups. As the solution becomes more dilute, this dissociation as a rule increases. Ions. These dissociated atoms or groups of atoms are called ions, and the process itself, ionization. They are regarded as being charged with electricity, and are of two kinds, anions or negative ions, and cathions or positive ions. The metals, ammonium, and hydrogen are cathions ; the acid radicals and elements, like NO 3 and Cl, and the group HO, hydro xyl, are anions. This is often called the theory of electrolytic dissociation, and concisely stated is that when compound substances are dissolved in water, they are to a greater or less extent broken up into their constituent anions and cathions. Application of the Theory. In the brief space of this text it is impossible to make application of the theory to any extent. For this the student is referred to Ostwald's Analytical Chemistry, translated by McGowan. An illus- 330 MODERN CHEMISTRY tration may, however, make the theory somewhat clearer. Suppose we have a solution containing lead nitrate, Pb (NO 3 ) 2 , silver nitrate, AgNO 3 , and mercurous nitrate, HgNO 3 . According to the ionic theory, the solution contains, not molecules of the three compounds mentioned, but largely individual ions of Pb, Ag, Hg, and (NO 8 ); hence, tests need be made only for these four. Now, when we add dilute hydrochloric acid, we introduce two other ions, H and Cl. When those of Pb, Ag, and Hg meet with the Cl ions, compounds form, which in the main are insoluble in water, hence are not dissociated or broken up into ions, and therefore fall as precipitates. The same is true in any other chemical reactions. Details of the Work. Group I. Take about two-thirds of the unknown solution, " Solution A," and add to it a little hydrochloric acid ; if any of the first group metals are present, they will come down as a white precipitate. Filter out and save the clear filtrate for work with the remaining groups. We will label this "Solution B." To be sure that enough hydrochloric acid has been used, add a drop or two to this filtrate. If it becomes turbid more must be added, and the whole solution again passed through the filter paper. Now wash the precipitate on the paper two or three times with cold water, and throw out the wash water. Next punch a hole in the bottom of the paper, and by directing a stream of water from the wash bottle upon the precipitate wash it through into a beaker. Do not use too much water, however ; usually 50 to 75 cc. will be sufficient. If the precipitate is not easily loosened by the stream of water, remove it with a spatula or stirring rod, and add it to what has already been washed into the beaker. Next, heat this to. the boiling point and after a minute or two filter quickly. APPENDIX A 331 If any precipitate remains upon the filter, wash once or twice with hot water. Tests for Lead and Mercury. Lead chloride is very soluble in hot water, and if it was present it will now be found in the filtrate. Test a portion of it with potassium dichromate, K 2 Cr 2 O 7 ; another portion, with potassium iodide, KI, or sulphuric acid. The first two give dis- tinctive yellow precipitates, the third, a heavy white one, somewhat soluble in water, but almost entirely insoluble in alcohol. Any precipitate left on the filter paper above will contain the mercurous and silver chlorides, if any were present. The latter of these is very soluble in am- monia ; so pour upon the filter paper a few cubic centi- meters of ammonium hydroxide. If mercury is present, the precipitate will turn black, and further proof is un- necessary. TABLE I SEPARATION OF LEAD, MERCURY, AND SILVER To the unknown solution, add HC1, filter out the chlo- rides, and wash the precipi- tates. Save the filtrate for de- termining metals of Group II and those following. Mark it "Solution B." Transfer the precipitates to a beaker; add H 2 O, and boil. Filter, and wash with hot water, if any precipitate remains. Test filtrate for Pb as in 1. De- termine Hg and Ag in the precipitate as in 2 and 3. 1. Test the hot water filtrate for Pb with K 2 Cr 2 O 7 , KI, and H 2 SO 4 . For results, see preceding work. 2. To the precipitate left undis- solved by the hot water, add NH 4 OH. If it turns black, mercurous salts are indicated. Test filtrate that runs through, for Ag by 3, below. 3. To the filtrate from 2, above, add HNO 3 till odor of NH 3 is no longer perceptible. A white precipi- tate indicates silver. 332 MODERN CHEMISTRY Test for Silver. To determine whether silver is present put the ammonia solution that has just filtered through into a test-tube and add nitric acid until no longer alkaline. This will be known by the absence of the odor of ammonia. If there is any silver present, a white precipitate will form, which may again be dissolved by adding ammonia. Group II. Through "Solution B," the filtrate from the chlorides of the first group, pass a current of hydro- gen sulphide, until, after shaking the solution, the odor of the gas is very perceptible. Any metals of this group will now be in the form of sulphides. Warm somewhat to collect the precipitates, and filter quickly. Preserve the filtrate, "Solution C," for determining metals of the third and succeeding groups. Now wash the precipitates left on the filter and reject the wash water. Transfer the precipitates to an evapo- rating dish and add a few cubic centimeters of yellow ammonium sulphide or sodium sulphide in solution, and warm gently for several minutes. This will dissolve the sulphides of division 1 of this group, that is, those of arsenic, tin, and antimony ; while those in the second division, mercuric salts, copper, bismuth, cadmium, and, as lead chloride is somewhat soluble in water, sometimes lead, will remain as precipitates. It should be stated, however, that copper sulphide is partially soluble in strong yellow ammonium sulphide; hence, when its pres- ence is suspected from the color of the original solution, it is better to use sodium sulphide to separate division one from two. When the sulphides have been digested as stated, filter and wash the remaining precipitate with water to which a drop or two of ammonium sulphide has been APPENDIX A 333 added. Save the filtrate to test for arsenic, tin, and antimony. Test for Mercury. Transfer the precipitates of mer- cury, copper, etc., to a beaker, add a few cubic centi- meters of dilute nitric acid, and boil. All the sulphides will dissolve except that of mercury, which will remain as a heavy black residue. Disregard any dark-colored par- ticles that remain floating upon the liquid, for they consist merely of sulphur colored with small portions of the black sulphides not yet dissolved. The student can prove this by collecting them upon a small loop in a platinum wire and igniting in the bunsen flame. The mass will burn with characteristic flame and odor. The indications of mercury shown by the black residue may be verified by filtering out, washing, and dissolving in aqua regia. Boil to dry ness, take up with water, and test one portion with stannous chloride. A white precipitate, turning gray when heated, or when more of the stannous solution is added, is distinctive. Test another portion with potassium iodide, adding a drop at a time. A bright red precipitate, soluble in excess of the reagent, should form. The filtrate from the mercuric sulphide, containing copper, bismuth, etc., should be boiled nearly to dryness, and water added to dissolve the salts. Before proceeding farther, it is always better, if lead has been .found in the first group, to test a small portion of this solution in water in a test-tube with sulphuric acid and a little alcohol added. If a precipitate of lead sulphate forms, it must be removed in the same way from the whole solution, using very little sulphuric acid. Test for Copper. Now add ammonia to the solution, from which you have removed the lead, until alkaline. 834 MODERN CHEMISTRY If the solution turns darker blue, copper is indicated ; at the same time bismuth will come down as a fine white precipitate. As the quantity of bismuth in solution is usually small, the student must be careful not to overlook it ; at the same time he must not mistake for bismuth a fine sediment sometimes carelessly allowed to collect in the reagent bottle used for ammonia. Test for Cadmium. To determine whether cadmium is present, after filtering out the bismuth, add to the blue solution potassium cyanide in solution, drop by drop, until the blue color has entirely disappeared ; then pass a current of hydrogen sulphide, by which the cadmium, if present, will be precipitated as a bright yellow sulphide. Tests for Arsenic, Tin, and Antimony. For separating and determining the presence of arsenic, tin, and antimony, various plans have been suggested, but nearly all are more or less tedious and require considerable care. The following plan, perhaps, is as satisfactory as any. To the ammonium sulphide solution of these metals, saved above, add dilute hydrochloric acid till the solution is no longer alkaline. The three metals will again be precipitated as sulphides. If the precipitate is pale yellow, or nearly white, and small in quantity, it probably consists mainly of sulphur, and none of the metals need be sought. If it is dark colored, gold or platinum may be present, or if copper has been found in the other division of this group, and ammonium sulphide was used instead of sodium sul- phide, the precipitate may be only copper. Filter, and throw ou& the filtrate, as it contains no metals. Wash the precipitate, as usual, and transfer it to a beaker. "/-Now add a little strong hydrochloric acid and warm gently ; the sulphides of antimony and tin will dissolve, but the arsenic will be unaffected. Filter, and test the I APPENDTX A 335 filtrate as follows : put into it a bright iron wire or nail, and after warming gently let it stand about fifteen minutes. The antimony is reduced to the metallic form, and the stannic chloride to the stannous. Filter or de- cant and test the solution for tin with mercuric chloride. (3 The results are those given in testing for mercury with stannous chloride in the other division of this same group. Wash thoroughly the precipitated antimony, and add to it a little strong hydrochloric acid and a few drops of nitric acid. The antimony will dissolve. Boil the solution \l nearly dry and add water. A white precipitate indicates antimony, which may be verified by passing a current of hydrogen sulphide through the solution. An orange- colored precipitate will result. The arsenic left undissolved by the hydrochloric acid above may be tested in several ways. Transfer the ar- senic sulphide to a beaker, add to it some strong nitric acid, and heat. The arsenic will dissolve. Now fill a test-tube about half full of a solution of ammonium molybdate, add to it a few drops of the arsenic solution prepared above, and boil. A yellow crystalline precipi- tate indicates arsenic. Sometimes the following method works satisfactorily. After adding concentrated hydrochloric acid to dissolve the precipitates of antimony and tin sulphide obtained from the ammonium sulphide solution, decant the clear solution into a test-tube. Now 'slowly pour hydrogen sulphide water down the inside of the tube. Presently the antimony will begin to precipitate, forming an orange- colored ring of the sulphide. Continue adding the hy- drogen sulphide solution, when above the antimony a ring of yellow stannic sulphide will form. 336 MODERN CHEMISTRY TABLE FOR GROUP II < cc *3 G 3 *T o -2? fc J* O CO ^"^ _^> fl} 3 -j-d ^-i ^^ '"' fA ~.22 s 3 -V 2 *> J H 3 ' .a o o *+"* ,M O cS 3 .& c3 O ^ H cS fl a. Put into the filtrate a bright iron wire or nail and let stand about 15 minutes. A black scaly precipitate of antimony forms. Filter out and test by b. To the filtrate add HgCl 2 , drop by drop, as a test for the tin. b. Dissolve the precipitated antimony in aqua regia, boil nearly dry, and add water. A white precipitate indicates antimony, verified by H 2 S, which gives orange-colored precipitate. c. Heat the precipitate of arsenic sul- phide with a little nitric acid and add some ammonium molybdate solution. A yellow crystalline precipitate will indicate arsenic. a PH a "o ^ 5 o a . o .2 2 " g a. Deep blue color indicates copper. b. White precipitate indicates bismuth. To verify, filter out, dissolve in HC1, boil nearly dry, and add H 2 O. While precipitate forms, filter. c. After filtering out the bis- muth, add KCy solution till the blue color has disappeared. Pass a current of H 9 S. A yellow pre- cipitate indicates cadmium. Group III. Like Group II, this is also usually separated into two divisions for convenience in analysis. The first includes iron, aluminum, and chromium, precipitated by ammonia ; the second, manganese, zinc, nickel, and cobalt, with ammonium sulphide as the precipitant. APPENDIX A 337 To " Solution C," the filtrate saved from Group II, after filtering out the sulphides, add a few drops of nitric acid and boil a short time. Now add ammonium chloride, NH 4 C1, and ammonium hydroxide till alkaline. The latter reagent precipitates the metals of the first division as hydroxides. Warm the solution, filter and wash as usual. Save the filtrate for the second division of this group and the succeeding groups. Transfer the hydrox- ides of iron, chromium, and aluminum to a beaker, add 20 or 25 cc. of strong potassium hydroxide solution, and boil several minutes. This will dissolve the aluminum and leave the others unchanged. Filter and wash. To a por- tion of the filtrate, after acidulating with hydrochloric acid, add ammonia till alkaline. A white, flaky, some- times starchy precipitate indicates aluminum. Test for Iron. Take a portion of the iron and chro- mium precipitate left undissolved and add hydrochloric acid. Test the solution obtained for iron, by using either potassium sulphocyanide, KSCy, or potassium ferro- cyanide. Test for Chromium. Next, take a rectangular piece of platinum foil and bend up the sides so as to form a small boat or pan. A piece of broken porcelain dish may serve the same purpose, but more heat will be needed. Put into the boat the remaining iron and chromium precipitate, add an equal amount of potassium nitrate, KNO 3 , and as much sodium carbonate, Na 2 CO 3 , and heat red hot until the whole mass has fused well together. Upon cooling, if chromium is present, it will assume a yellowish appearance. Put the boat and contents into a beaker containing a little water and dissolve the mass. Acidulate the solution with acetic acid and test a portion with silver nitrate. A brick or blood red precipitate of silver chromate, Ag 2 CrO 4 , indi- 838 MODERN CHEMISTRY cates the presence of chromium. Test another portion with lead acetate, Pb(C 2 H 3 O 2 ) 2 . Tests f Oi Nickel and Cobalt. To the filtrate saved for the second division of this group, add some ammonium sulphide. If precipitates of light color are obtained, nickel and cobalt are not present, as their sulphides are black. If nickel is present, the filtrate will often be of a dark- brown color, which is apt to lead the student to think the solution is not filtering well. Disregard this, mark it " Solution D," and save for work with the fourth group. After washing the precipitates, transfer them to a beaker and treat with dilute hydrochloric acid ; the sulphides of zinc and manganese will dissolve, while those of nickel and cobalt will remain as a black residue. Filter and wash. Test the black residue with the borax bead ; cobalt gives the well-known blue in the oxidizing flame, and nickel, yellow to brown or black, according to the amount intro- duced into the ^ead. If both metals are present, the cobalt blue will obscure the brown, and further tests are neces- sary ; for these the student is referred to any manual on qualitative analysis. Tests for Zinc and Manganese. To the solution sup- posed to contain zinc and manganese, after boiling for two or three minutes, add caustic potash till strongly alkaline. Allow it to stand for some time, for manganese precipitates slowly. If present, it may be filtered out and the precipi- tate tested with the borax bead. It imparts a beautiful amethyst color. Acidulate the filtrate with acetic acid and add ammonium sulphide till alkaline. A white precipitate indicates zinc. This is usually verified by heating on charcoal, moistened with a solution of cobaltous nitrate. A green mass is obtained ; aluminum compounds treated in the same way give a blue mass. APPENDIX A TABLE FOR GROUP HI 339 c 5 l-H ^ :i le - 1 i a. Test for Co with borax bead ft !i ^ ^ z ~ blue. z > Z ^ c -^ I ~ < ~ > 2 | b. If Ni is present with no cobalt, borax i | ^ ~ * c 1 bead will become yellow to brown in oxi- g s aT 1^. '^ ^ ? -IT * dizing flame. pC tj i rzr ^ O " 3 CD ^ - 1 " 5 -r J c. Add considerable excess of KOH ri 1 s C OD p s and let stand for some time. A slow- c" X --^ s 3 A; "^ - forming precipitate indicates Mn. Filter - 0) > ^0 Q ^> " i ^ and test filtrate by d. Verify Mn with g-* 1 X -^- x ' fl ^ * * borax bead. Amethyst color. tl _= gt J 'E, | B "^ | -^^ d. Acidulate filtrate from c with acetic 42 eB - 'o 2 ., ^ T 'f 2 "S acid, add (NH 4 ) 2 S till just alkaline. tJ -r Cu ^ J &.5S Zinc forms white precipitate of ZnS. it ^5 QQ b 1 ES B c ,i E. a. Filtrate may contain aluminum. Acidulate g ^ . ^ r^ b. Dissolve a small portion of the precipitate b - bO a | ~ in HC1 and test for iron with KSCy, I^FeCy,., | '-Zi ~ ~ a. O 1 'i -3 &0 or NH 4 OH. | ""-r ~ ^ 1 .2 EB = < I O ; f. $ 'S c g c. Fuse the remainder in platinum dish with KNO,and Na 2 CO 3 . Dissolve in H 2 O, acidulate o r- 73 ^s _ 13 i i^ 5 with HCgHgOj. Test for Cr with AgNO 3 , also 9 o 5 ^ c3 45 with Pb(C 2 H 3 6 2 )2. Group IV. For this use " Solution D " saved from Group III. It is better to make a preliminary test be- fore proceeding with the whole. To do this, add to a small portion of the solution to be analyzed a little disodium phosphate. If a white precipitate forms, some of the metals at least are present, and all must be tested 340 MODERN CHEMISTRY for. If so, add to the whole ammonium chloride, am- monium hydroxide, and ammonium carbonate. A white precipitate may contain calcium, strontium, and barium in the form of carbonates ; filter and wash. Save the filtrate to test for magnesium and fifth-group metals. Test for Barium, Strontium, etc. Transfer the precipi- tates to a beaker and dissolve in acetic acid. Test a small portion of this with potassium dichromate ; a light yel- low precipitate indicates barium, which may be verified by the flame test. If present remove it from the entire solu- tion by adding the dichromate and filtering. To the filtrate add caustic potash till alkaline and a little more potassium dichromate, when the strontium, if present, will be precipitated, as strontium chromate is insoluble in alkaline solutions. Remove this by filtering, and test the filtrate for calcium by adding ammonium oxalate. This gives a fine white precipitate of calcium oxalate. It is customary to verify the strontium by the flame test, as its salts impart a crimson color which is very persistent. There are other plans for effecting a separation of the metals of this group, of which the following is frequently used. After removing the barium, to a small portion of the filtrate add a little strong solution of calcium sul- phate. If a white precipitate forms, strontium is present and must be removed. Add to the remainder of the solu- tion a very little sulphuric acid ; the strontium will slowly precipitate. After a few minutes filter and test filtrate for calcium. To do this add sufficient ammonia to neu- tralize any excess of sulphuric acid present, and then add ammonium oxalate solution as in other methods. To a small portion of the filtrate saved above, after pre- cipitating the barium, strontium, and calcium with am- APPENDIX A 341 monium carbonate, add a little disodium phosphate ; a white precipitate which may form slowly will indicate magnesium. TABLE FOR GROUP IV S o g CS - hjT 2 ts 5 8 - 1. To a small portion of the solution, add a little K.,Cr 2 O 7 solution. If Ba is present, indicated by the forming of a light yellow precipitate, treat the whole of the solution in the same way, and filter. The Ba pre- cipitate may be verified by flame test. Test the filtrate by 2. 2. Render the filtrate alkaline by adding KOH, and then add a little more K 2 O 2 O r If strontium is present, it will be precipitated and may be filtered out. Or, from the filtrate from 1 above, the strontium may be removed by adding a little sulphuric acid. Let it stand a few minutes and then filter, and test filtrate for Ca by 3. The precipitate may be verified by flame test. 3. When the barium and strontium have been re- moved, if the filtrate is not already alkaline, render it so by adding NH 4 OH. Then add ammonium oxalate; white precipitate is distinctive of calcium. May be verified by flame test, orange-yellow. 4. To a small portion of the filtrate saved from "Solution D," add a little disodium phosphate. A ~ ~? white precipitate indicates magnesium, distinctive in fc. 3 j the absence of other metals of this group. Filter, and *** rt save filtrate for Group V, " Solution E." Group V. Sodium, Potassium, Lithium. The salts of the metals of this group are all soluble in water, hence none of the reagents used in the previous steps of analysis pre- cipitate them. The flame, especially with the spectroscope, is usually all that is necessary for their identification. 342 MODERN CHEMISTRY As sodium is so widely distributed, a slight test for it may nearly always be obtained, but the student must learn to disregard any except a decidedly strong indication. As already seen, if sodium is present, the potassium flame can be perceived only by making the observation through a sheet of cobalt glass. Before making these flame tests, boil the solution, saved from Group IV, to dry ness, and heat gently until ammonia fumes are no longer driven off. Dissolve the residue in water, and acidulate with hydro- chloric acid. Sodium gives bright yellow flame, Potassium gives violet flame, Lithium gives bright red flame, lasting but a moment. Potassium may also be tested in another way. To the solution used in making the flame tests, add some platinic chloride in solution and a little alcohol. Allow it to stand for some time, stirring occasionally with a glass rod. A small quantity of a yellow precipitate of potassic-platinic chloride, K 2 PtCl 6 , is slowly deposited. A large watch crystal serves well for making this test. Test for Ammonia. Ammonia must be looked for in the original solution, as so many ammonium compounds are used as reagents in making the analysis. Put a few cubic centimeters of the original solution into a beaker and add caustic soda or potash until strongly alkaline. Moisten the under side of a watch crystal with a drop of water and upon it place a short strip of red litmus paper. Put the. crystal over the beaker, and warm the solution gently. If ammonia is present, it will be liberated by the non-volatile alkali added, and will turn the litmus paper blue. APPENDIX A TABLE FOR GROUP V 343 1 Boil filtrate from Table IV to dry- ness, ignite to expel NH 4 compounds, dissolve in H 2 O and acidulate with HCI. Test for Na, K, Li, by flame as in 1 ; K by wet method, as in 2 ; NH 4 , as in 3. 1. Make test with platinum wire; Na gives yellow ; K, violet when alone ; Li, red. Use cobalt glass, if Na is present, to distinguish the violet rays. 2. Sometimes K must be detected otherwise than by the flame test. To the acidulated solution, add PtCl 4 ; a yellow precipitate, slowly forming, indicates K. 3. To a portion of the original solution, add caustic soda or potash till alkaline, and warm gently. Suspend a strip of red litmus in the fumes arising. If it turns blue, NH 4 is indicated. The five tables given above simply show in condensed form the methods already described ; and when the student has once seen the details and understands them, he will find the tables very convenient for rapid work. For a successful analysis, neatness is absolutely essential, and great care must be used in washing the precipitates so as to remove all of the metals contained in the filtrate. Detection of Acids. As a rule, the beginner will only meet with a few of the more common acids, and these only will be noticed here. They may be placed in groups, somewhat as the metals are, according as they are affected by certain reagents. Group I. This includes those which form a precipitate with barium chloride. The only one with which the student will meet often is sulphuric acid. As already seen, this gives with barium chloride, barium sulphate, BaSO 4 , insoluble in all acids. If the solution be neutral, phosphoric acid or the phos- 344 MODERN CHEMISTRY phates also give a white precipitate with barium chloride ; but this is soluble in hydrochloric acid. After being thus dissolved, if the solution be made alkaline with ammonia, the precipitate will again fall. Sulphurous and thiosulphuric acids are usually put in this group. They may be easily distinguished, how- ever. To the solution add a little strong hydrochloric acid ; both sulphurous and thiosulphuric acid and their salts will give off fumes of sulphur dioxide which may be readily detected. The latter, however, at the same time, throws down a milky or pale yellow precipitate of sulphur, while the former remains clear. The reaction is shown below : Na 2 SO 3 + 2 HC1 = 2 NaCl + H 2 O + SO 2 (sulphurous), Na 2 S 2 3 + 2 HC1 = 2 NaCl + H 2 O + SO 2 + S (thiosulphuric). Group II. This includes such as form no precipitate with barium chloride, but do with silver nitrate. The most common are: Hydrochloric, HC1, curdy white precipitate, very solu- ble in ammonia. Hydrobromic, HBr, pale yellowish white precipitate, slowly soluble in ammonia. Hydriodic, HI, pale yellow precipitate, very slightly soluble in ammonia. Methods for testing each of these and its compounds have been given in the text, and the student is referred to them. Hydrogen sulphide, H 2 S, if free, is known by the odor. In the form of compounds, it may usually be detected by adding some acid and heating, whereby the gas is liberated and its characteristic odor becomes perceptible. APPENDIX A 345 Group III. Here belong those acids which form no precipitate with either barium chloride or silver nitrate. The only common one is nitric, but the salts of nitrous and chloric acids, HNO 2 and HC1O 3 , especially those of the latter, have occasional use in the laboratory. A plan for testing and distinguishing between nitrous and nitric acids was given in the text. The following plan, however, usually works satisfactorily, and by some is preferred to the other. Into a test-tube, containing the solution to be tested, drop a crystal of ferrous sulphate, and then pour down the sides of the tube a few drops of strong sulphuric acid. A brown ring will form about the crystal of ferrous sulphate. The chlorates, for example, potassium chlorate, KC1O 3 , heated with sulphuric acid, yield chlorine, and chlorine peroxide, a very explosive gas. Usually, if sulphuric acid is added to a crystal of the chlorate, a sharp explosion occurs, throwing the materials out of the tube. Group IV. We might place here certain organic acids, which require special tests for identification. The only common one is acetic, HC 2 H 3 O 2 , though the student oc- casionally meets with one or two others. Acetic acid and its salts are tested by adding a solution of ferric chloride and boiling. The solution becomes a deep red color which may be destroyed by using hydrochloric acid or mercuric chloride. Oxalic acid, H 2 C 2 O 4 , might be placed here, though more properly in Group I, as its salts form a white precipitate with barium chloride in neutral or alkaline solutions ; this precipitate is soluble in hydrochloric acid, but not in acetic. Preliminary Work. Before testing any solution for acids, the metals present should be determined, other- 346 MODERN CHEMISTRY wise the student may be greatly misled. If any are present which would interfere with necessary tests, that is, if there are any which would form precipitates with the reagents necessarily used in making the acid tests, they must be removed before proceeding with the deter- mination. Again, if the unknown substance is in solution, it would be useless to look for the acids whose salts are insoluble in water. For example, if we have found lead or barium present in a given solution, obviously it would be unnec- essary to look for sulphuric acid. Hence a knowledge of the solubility of salts is very important, and the fol- lowing incomplete table is given, showing the solubility of a few of the more common salts: Acetates, soluble in water. Bromides, nearly all soluble ; exceptions, those of first group metals and mercuric. Carbonates, only those of Group V, the alkali metals. Chlorides, nearly all, Group I excepted. Iodides, nearly all, Group I excepted, also certain iodides of bismuth and copper. Nitrates, all soluble. Nitrites, nearly all soluble. Phosphates, only those of Group V. Sulphates, many insoluble, such as those of barium, mercury, lead, and silver ; and calcium and stron- tium, nearly so. Sulphites, only those of Group V. Sulphides, only those of Groups IV and V. If the substance, of which the acid radical is to be determined, is not in solution, it is often of great advan- tage to make certain preliminary tests. Put a small por- APPENDIX A 347 tion of it into a test-tube and add a little strong sulphuric acid. Warm gently, and notice the color and odor of the gas obtained. The more common are shown below : Acetates, odor of vinegar, no color. Bromides, sickening odor, brown color, resembling that of nitrogen tetroxide. Odor is more offensive and peculiarly irritating to the eyes. Carbonates, strong effervescence, no special odor, colorless gas. Chlorides, very irritating gas (HC1), colorless. Iodides, peculiar odor, resembling weak chlorine, violet color. Nitrates, very irritating gas, no color. Nitrites, irritating gas, brown in color ; not so offen- sive as bromine. Phosphates, no special action. Sulphates, no special action. Sulphites, saffocating gas (SO 2 ), colorless. Sulphides, offensive odor (H 2 S), colorless. Thiosulphates, suffocating gas (SO 2 ), colorless. The student must remember that these are merely pre- liminary steps and must be verified by distinctive tests already described. APPENDIX B SOME ADDITIONAL QUANTITATIVE WORK IT is believed that all the quantitative work that the ordinary class can do has been introduced into the text. There may be occasions, however, when it will seem desir- able to vary the work or even to furnish more to certain students ; to meet such a demand, the following experi- ments are offered. 1. To estimate Amount of Carbon Dioxide in any car- bonate soluble in acids. (Adapted from Fresenius.) Fit two small bottles with rubber stoppers and glass tubing, as shown in the figure. E is a short piece of rubber tubing which may be closed air-tight by means of a screw clamp. The carbonate to be used, calcite for example, CaCO 3 , is accurately weighed, placed in M, and covered with water. N is filled over half full of pure concentrated sulphuric acid. Find the weight of the whole, which should not exceed 60 to 70 g., tighten the clamp at E, and test the appa- ratus to see that it is air-tight. Now by suction at (7, partially exhaust the air in JV; this will have a like effect upon M^ and upon readmitting the air to N the acid will be forced over into M. The 348 FIG. 65. APPENDIX B 349 carbonate will thus be decomposed, and the carbon dioxide will escape into JV, being dried as it bubbles through the acid. When the carbonate has all been decomposed, and the evolution of gas has ceased, open the clamp at E, and by means of an aspirator or by suction remove the carbon dioxide from Msmd N, and when the apparatus has become cool, weigh again. The loss represents the amount of carbon dioxide expelled by the acid. Carbonate used (for example) . . 1.0 g. Apparatus and contents, say . . 68.0 g. After decomposition by acid : Total weight . . . . x g. Loss 68.0 -a;. CO 2 = 68 - x g. 2. To determine the Water of Crystallization in a Com- pound. This is usually done by heating a known weight of the compound, and noting the loss. To illustrate, put into a small porcelain crucible, the weight of which is known, about a gram of magnesium sulphate, and weigh carefully. Support the crucible in a clay triangle upon an iron ring-stand. With the Bunsen burner heat cau- tiously at first, increasing to red heat, cool slowly and weigh. Heat a second time four or five minutes and weigh again. Do this until two successive weighings show the same results, then calculate the per -cent of water of crystallization. Tabulate results thus : Weight of crucible -f MgSO 4 . . . a Weight of crucible alone .... b Weight of MgSO 4 . . . . a - b 350 MODERN CHEMISTEY After the second and third heating, when weight was the same : Crucible + MgSO 4 .....~ ample, suppose we pass into the eudio- _ air meter 20 cc. of air and then 10 cc. of hydrogen. We now have a total amount -botkgases o j- ^Q cc . ^^ an e i ectr j c S p ar k to ex- plode the hydrogen and oxygen. As two parts of hydrogen unite with one of oxygen, one-third of the loss would represent the oxygen, and the other two-thirds the hydro- gen, which has combined to form water. The residue will contain the nitrogen of the air and any excess of hydrogen. Take the quantities used above : Air 20 cc. Air + H 30 cc. Residue after exploding . . 18 cc. Loss 12 cc. \ of loss = 4 cc., the oxygen of air used. 20 cc. air = 4 cc. oxygen -h 16 cc., nitrogen of air. APPENDIX B 351 Let the student arrange his own apparatus for the above experiment, making all corrections necessary for accurate results, and prove the usual statement that air is one-fifth o-cv^en and four-fifths nitrogen. 1 The Volumetric Composition of Ammonia. The .{position of ammonia may be determined, but the ex- periment requires time and is somewhat tedious. The plan is as follows: into a eudiometer, over mercury, introduce a few cubic centimeters of dry ammonia gas, and pass sparks from an induction coil for twenty or thirty minutes or until the volume of the ammonia seems no longer to increase. This, in accordance with the law of Gay-Lussac, should now be double what was intro- duced into the eudiometer. Next add sufficient oxygen to explode with the hydrogen obtained from the ammonia, and pass a spark. It is obvious, from the proportions in which hydrogen and oxygen combine, that two-thirds of the loss represents the hydrogen, which, subtracted from the volume of the gases after electrolysis, gives the amount of nitrogen contained in the ammonia. Take the following example : Ammonia gas introduced . . . 8 cc. Vol. of mixed gases after passing sparks, 16 cc. Oxygen added 8 cc. Total amount 24 cc. Residue after exploding . . . 6 cc. Loss . ... . . .18 cc. Two-thirds of loss = hydrogen, which was obtained from the ammonia gas . lt;c. Volume of mixed gases . . . . 16 cc. Subtract volume of H . . . 12 cc. Volume of N .4 cc. 352 MODERN CHEMISTRY This proves that hydrogen and nitrogen unite in the proportion of three to one to form ammonia ; furthermore we have seen that the four volumes of the mixed gases upon combining are condensed to two. Let the student arrange his own apparatus, taking every precaution to insure accuracy, and, using different quantities from those mentioned above, prove the truth of the preceding statements. 5. Composition by Volume of Hydrochloric Acid. This may be learned by the interaction of sodium and hydro- chloric acid, by which is formed common salt and free hydrogen. In order to lessen the rapidity of the reaction, an amalgam of sodium should be used. This may be prepared by putting a small quantity of mercury into a mortar, and then, by means of forceps, thrusting small pieces of sodium, one at a time, into the mercury. Do this until a pasty mass is obtained, which upon cooling becomes solid. In preparing the amalgam do not hold the face too close to the mortar, as the combination some- times takes place with considerable violence. The hydrochloric acid gas for this experiment must be dried, either by bubbling through strong sulphuric acid or by passing through a drying tube containing bits of porcelain or pumice stone moistened with sulphuric acid. The gas may be generated by the reaction of common salt and hydrochloric acid with sulphuric acid as de- scribed on page 112, Sec. 26. If dried by passing through sulphuric acid, the rapidity of evolution of gas may be observed and regulated by increasing or decreasing the aniouiff of heat applied. It is better, if possible, to collect the gas over mercury rather than by downward displacement, for in this way it may be obtained free from air. APPENDIX B 353 For this experiment, some straight graduated tube should be used, such as the eudiometer shown in some of the illustrations for the synthesis of gases. If this is not to be had, you may use a burette, the capacity of which, both above and below the graduations, is accurately known. (See Fig. 67 for the general arrangement of the FIG. 67. apparatus.) When the graduated tube is completely filled with gas, put around it, as near the mouth as possible, a paper test-tube holder. This is made by folding a sheet of paper into a strip about an inch wide ; for use it is simply placed around the tube as shown in the accom- panying figure at N, and grasped tightly between the thumb and fingers. The paper, being a poor conductor of heat, serves to prevent the transmission of the warmth of the hand to the glass so as to expand the hydrochloric acid. By means of this holder seize the tube, hold the thumb firmly over its mouth, and place in an upright position. Next, FIG. 68. 354 MODERN CHEMISTRY quickly drop into the tube a few grams of the sodium amalgam already prepared, and instantly replace the thumb, holding it as tightly as possible. Tip the tube back and forth a few times to hasten the action, and when this seems complete, place the mouth of the tube beneath the surface of the mercury and remove the thumb. The mercury instantly rises in the tube to fill the space formerly occupied by the chlorine, but now existing in the form of solid sodium chloride. Measure accurately the amount of gas remaining, and compare with the capacity of the tube ; what are your conclusions ? Test the residual gas with a light ; what is it ? What evidence have you that common salt is formed ? If the student finds he cannot hold his thumb tightly enough over the mouth of the tube to prevent leakage, he may use a short rubber stopper instead, and after the reaction of the sodium with the gas the tube may be opened over water. 6. Analytic Proof of the Composition of Hydrochloric Acid. This may be furnished by the electrolysis of hydrochloric acid and the measurement of the gases obtained. Let the student arrange his own apparatus, and, taking such precautions as are necessary to avoid possible errors (mentioned in describing certain forms of electrolytic apparatus), make the experiment, and note results. APPENDIX C SOLUTION 1. Meaning of Solution. Ordinarily we think of solu- tion as the disappearance of a solid in a liquid, but in re- ality the term is much broader. Not only may gases, solids and liquids disappear in a liquid, but one solid may dissolve another, or a solid may dissolve a liquid. To il- lustrate, Sir Roberts-Austen of London, placed some lead cylinders upon some sheets of gold and allowed them to remain for four years. At the end of that time an analy- sis showed that the gold had penetrated the lead to the depth of 8 mm. So other illustrations might be given. By a solution, then, we mean the homogeneous mixture of two or more substances. The one in excess we call the solvent; the other, the solute. NOTE. The student will frequently find in works on volumetric analysis the term normal solution. By this we mean that in a liter of the solution there is of the solute the equivalent of one gram of hydrogen ; that is, of hy- drochloric acid, HC1. there would be 36.5 g. (its molecu- lar weight in grams); of sodium hydroxide, 40 g., and so on; of sulphuric acid, H 2 SO 4 , one half its molecular weight, or 49, because sulphuric acid contains two atoms of hydrogen ; of oxalic acid, H 2 C 2 O 4 , 2 H 2 O, one half of its molecular weight in grams, or 63. A solution containing 1/10 as much of the solute as given above is called deci- normal and marked N/10. The solutions suggested for 355 356 MODERN CHEMISTRY use in the experiments on pages 168 and 169 are deci- normal. 2. Characteristics of a Solution. As solutions of solids in liquids are by far the most common, unless otherwise stated we shall understand that such are meant. One of the most notable facts is that a solvent has its boiling point raised whenever any substance is dissolved in it. This matter will be taken up at more length later in this chap- ter. Another characteristic is that the freezing (solidify- ing) point is lowered. Thus it is well known that pure water, which freezes at zero, if saturated with common salt, requires a temperature many degrees lower. Another interesting example is that of the fusible metals ; for ex- ample, "Woods' alloy." It consists of lead, bismuth, tin, cadmium, with melting points ranging from 335 to 235 C. and fuses at 65, much below the boiling point of water. Here lead, which constitutes the greater part of the alloy, may be regarded as the solvent, with a melting point of 335, which by the presence of the others is reduced to 56. As there is always a contraction of volume when one substance is dissolved in another, there is necessarily an increase in the specific gravity of the solution. 3. Theory of Solution. In every substance there exist two opposing molecular forces, the one attractive, the other repellent. When a solid is brought into contact with a liquid, the adhesion of the liquid and solid mole- cules acts in conjunction with the repellent force among the molecules of the solid, weakening it to such an extent that the particles of the solid pass into the liquid. This continues until as many molecules of the solid return to it in a given time as are given off, when equilibrium obtains, and the solution is said to be saturated. In dilute solutions we have conditions of the solute sim- . APPENDIX C 357 ilar to that of a gas. That is, the molecules of the dissolved substance are separated to such an extent that, as Van't Hoff, the great Dutch chemist, has shown, they practically obey all the laws of gases. It is thus that we are able to explain the osmotic pressure of liquids. 4. Influence of Heat upon Solution. As a rule the solu- bility of solids is increased when the temperature is raised, In some cases the increase is very marked. The table below shows the solubility of a number of substances in 100 cc. of water at different temperatures : C. 100 C. Potassium alum 3.9 357.48 Silver nitrate 120.0 1000.00 Common salt 35.60 39.80 5. Solution of Gases. Gases when brought into con- tact with a solvent may disappear in it, forming a true solution, or they may unite with the liquid, forming what we call a chemical solution. In both cases a rise of tem- perature decreases the solubility of the gas. Like solu- tions of solids, this effect of temperature is very different. Notice the solubility of the several gases in 1 cc. of water at the different degrees: C. 50 C. 100 C. CO 2 . . . 1.78 cc. 0.5 cc. Occ. H 2 S . . . 4.37 2.00 NH 3 . . . 1148.00 306.00 HC1 . . . 503.00 364.00 12 6. Effect of Pressure. According to the kinetic theory of gases their pressure is due to the bombardment, so to speak, of the molecules upon the containing surface. It follows, therefore, that if the pressure upon a gas in con- tact with a liquid be doubled, twice as many impacts of the gaseous molecules with the liquid must result. Hence, other conditions being the same, twice as much of the gas 358 MODERN CHEMISTRY must be dissolved in the liquid. This has been expressed in what is known as Henry's Law, which may be stated thus : The amount of gas dissolved by a certain volume of any liquid is proportional to the pressure. Ordinary soda water is a familiar instance of the solution of a large quantity of gas, owing to high pressure. In this case, as soon as the water is allowed to flow into the open glass, where only the pressure of the air is upon it, the gas begins to escape rapidly. To illustrate further, 100 cc. of water at zero and one atmosphere pressure will dissolve 180 cc. of carbon dioxide, while at 4 atmospheres, 720 cc. When we think of the great pressure upon sub- terranean streams of water and remember that limestone is soluble in water charged with carbon dioxide, the for- mation of vast caves seems very reasonable. 7. Conductivity of Substances. In studying electroly- sis of water, it will be remembered that a little sulphuric acid was added to the water in order, as was said, to increase the conductivity. If an attempt be made to pass a current of electricity through pure water, it will be found that water is a non-conductor. Further, it will be found that pure sulphuric acid or liquid hydrochloric acid will not conduct the current. Again, if the terminals of a battery be brought into contact with a lump of salt or any similar compound, no current is transmitted. Never- theless, if either of the acids mentioned, or the salt, be added to the pure water, which it will be remembered is a non-conductor, the solution becomes a good conductor. These phenomena indicate that some important change has taken place in one or both of the substances. 8. Other Phenomena. It has already been stated that any solvent has its boiling point raised when a substance is dissolved in it. For example, water saturated with APPENDIX C 359 common salt boils at 108 C.; saturated with potassium nitrate, 116; with calcium chloride, 179. Numerous exper- iments have shown that this elevation of the boiling point is proportional to the quantity of the substance dissolved; furthermore, it has been found that to secure a like change in the lowering of the vapor tension, with different sub- stances, amounts proportional to their molecular weights must be used. (See page 68.) To illustrate, suppose the molecular weight of the compound A is 342 and of B 46; then to secure like results in lowering of vapor tension, we should be required to dissolve portions of the compounds in the ratio of 342 to 46. Such substances are sometimes said to give a normal rise in boiling point. 9. Exceptions. A large number of substances, like common salt, cause a lowering of the vapor tension about double what we should expect from the above statement. Many others, like calcium chloride, cause a lowering three times, etc. 10. Freezing Points. Such substances as show what we have spoken of as a normal lowering of vapor tension likewise lower the freezing points of their solvents in the same way. For example, if a gram-molecule (grams equal to the molecular weight of the compound) of cane sugar, 342 g., be dissolved in a liter of water, the solution freezes at 1.8 C. Likewise a gram-molecule of grape sugar, 180 g., dissolved in a liter of water, lowers the freezing point to 1.8. And so all those substances which raise the boiling point normally, lower the freezing point approximately as given. But if we dissolve a gram -molecule of common salt in a liter of water, the freezing point is lowered about 3.5, which is practically twice that of the others given. What is true of this is also true of all other similar compounds. 360 MODERN CHEMISTRY 11. Osmotic Pressure. The same variations are found in osmotic pressure. Such substances as lower the freezing point normally exert an osmotic pressure proportional to their molecular weights, but substances belonging to what we have noted as exceptions, show an increase in pressure as in other respects named. 12. Some Related Facts. A further study of the sub- stances shows that those which cause a normal lowering of vapor pressure, etc., in solution are non-conductors, while the others render the solutions good conductors. Thus, cane sugar, belonging to the normal class, shows a con- ductivity exceedingly small. Acids, bases, and salts (see Chapter X) in solution are excellent conductors. 13. Dissociation by Heat. At a temperature of about 2500 C. steam is about half decomposed into oxygen and hydrogen molecules. It is impossible, however, to decom- pose all of the molecules, no matter how long the heat is applied, for there are as many molecules of water formed by the recombination of its components as are being broken up by the heat. Such decomposition as this is spoken of as dissociation. Iodine above 600 C. has a density only half what it has below 600; this indicates that the molecules of iodine have become dissociated by the heat. (See Chapter XVI.) The same has been found to be true of sulphur. (See section 7, page 175.) It must be borne in mind that dissociation is not decomposition such as we have when mercuric oxide is heated strongly. In that case the mercuric oxide molecules are not reformed, while in dissociation, when the producing cause is removed, the molecules go back to their former conditions. Thus, H 2 ; H 2 + O, NH 4 C1 ^ NH 3 + HC1, I 4 ^ 2 I 2 . 14. Dissociation by Solution. As various substances are dissociated by heat, so others are when taken up by APPENDIX C 361 water or other solvents. This fact has been learned by determining their molecular weights in solution. Without showing how the formula was derived, the following has been found to be true : in which Mis the molecular weight of the dissolved sub- stance ; K, a constant factor ; j9, the concentration of the solution, that is, the amount of the solute per 1-00 g. ; d, the lowering of the freezing point (or vapor tension). By using dilute solutions and determining the lowering of the freezing point, or of the vapor tension, and substitut- ing in this formula, we are able to determine the molecular weight of the substance dissolved. As long as we are dealing with what we have spoken of as normal substances, the results are very nearly in accord with results de- termined by other methods. But when we come to deal with the abnormal substances, we obtain results one half, one third, etc., what they should be theoretically. It must be concluded, therefore, that the molecules have been dissociated by the solvent as they were in other cases by heat. Such dissociation by a liquid we call ionization, and the dissociated parts are known as ions. 15. Relation to Conductivity. From the fact that only solutions which show the irregularities above mentioned are conductors, as well as for some other reasons, it is concluded that electric conductivity is by means of ions. Such substances as are thus dissociated are called electro- lytes, and the decomposition is spoken of as electrolytic dissociation. 16. Positive and Negative Ions. Whenever any com- pound is dissociated, two kinds of ions are formed, positive and negative. For example, common salt gives positive 362 MODERN CHEMISTRY + sodium ions, Na, and negative chlorine ions, Cl. This may be readily shown by a simple experiment. If a V- shaped tube be partly filled with a dilute solution of com- mon salt and a current of electricity be passed through it, the sodium ions will move toward the negative electrode and the chlorine toward the other. This may be seen by using a few drops of phenol-phthalein in the salt solution or sufficient blue litmus solution to color it. From the law of electrical attraction and repulsion, therefore, we know that the sodium ions are positive and the chlorine negative. 17. Ions not Atoms. It is important to know that ions are not atoms, but electrically charged atoms or groups of atoms. Thus we may have either simple or complex ions: common salt dissociates necessarily into sodium and chlo- rine, simple ions positively and negatively charged. Sal ammoniac, ammonium chloride, breaks up into ammonium, + NH 4 , a complex positive ion, and chlorine, a simple nega- tive ion. Again, potassium chlorate dissociates into potas- sium, simple positive, and chlorate, C1O 3 , complex negative, ions. 18. Application to Electrolysis. According to this theory, electrolysis is not due to the breaking up of the water molecules by the electric current; but the ions pres- ent simply serve as carriers for the current and give up their charges to the electrodes. For example, when com- mon salt is put into water, it is broken up into the ions sodium and chlorine. When the current is turned on, the sodium cathions are repelled from the anode, and move across to the cathode, where they become discharged from contact with the cathode. But sodium atoms in the pres- ence of water, as the student knows, at once decompose APPENDIX C 363 the water molecules, forming sodium hydroxide and free hydrogen. This gas soon begins to collect in bubbles upon the cathode and shortly after to rise through the water. The chlorine ions are attracted to the anode, and in their turn give up their negative charge to the positive electrode, becoming free chlorine atoms. But chlorine in water at once begins to decompose it (see section 18, page 109), forming hydrochloric acid and setting free oxygen. Thus bubbles of oxygen soon begin to rise from the anode. The hydrochloric acid formed at the anode is then again dissociated and the process continues indefinitely. 19. Ions present in a Solution. When one tenth of a gram-molecule of potassium chloride is dissolved in a liter of water, about 75 per cent of the molecules are dissoci- ated. We might represent it thus: 100 KC1 75 K + 75 Cl + 25 KC1. This means that some of the ions are constantly uniting to form molecular potassium chloride, while other like ions are being formed. But water is very slightly ionized and + + gives H and HO ions. Now whenever a H ion meets a Cl ion, a molecule of hydrochloric acid forms, until a cer- + tain equilibrium is reached. In the same way the K and HO ions form molecular potassium hydroxide. Hence in + + such a solution we should have not only the ions, K, H, Cl, HO, but also the molecules KC1, H 2 O, NaHO, and HC1. 20. Chemical Action Ionic. From a large number of experiments which cannot be given here, and for other reasons, it is believed that chemical action always takes place between ions. For example, granulated zinc in dry 3G4 MODERN CHEMISTRY liquid hydrochloric acid shows no chemical action what- ever; so silver nitrate added to hydrochloric acid dissolved in benzene in which the acid is not ionized, gives no trace of the familiar white precipitate. When we neutralize a solution of caustic potash with hydrochloric acid (see Chapter X), at first we have the four kinds of ions and the four of molecules already named. As all strong acids and bases as well as salts ionize to about the same extent, we soon have an equilibrium of all these eight substances, which holds until we begin to boil the solution down. As evap- oration proceeds, the solution becomes saturated, and crys- tals of the chloride begin to separate from the solution. As these deposit, the equilibrium is destroyed and more of the molecular acid and alkali ionize and more molecular potassium chloride forms. This in turn crystallizes, and so the process continues until all the molecules of hydro- chloric acid and of potassium hydroxide have ionized, so that when the water has been entirely evaporated, only molecules of potassium chloride remain. APPENDIX D SOME CARBON COMPOUNDS \\ T E have already studied a few carbon compounds. The three hydrocarbons, methane, ethylene, and acetylene, may serve as the starting point for a very large number of similar compounds. This will be seen from the following : CH 4 . Methane C 2 H 4 . Ethylene C 2 H 6 . Ethane C 3 H 6 . Propylene C 3 H 8 . Propane C 4 H 8 . Butylene C 4 H 10 . Butane C 5 H 10 Ainylene C n H 2n + 2 General formula C n H 2w . General formula C 2 H 2 . . Acetylene C 3 H 4 . . Allylene C n H 2n _ 2 General formula for the series In all three series it will be noticed that the difference in the formulae of any two consecutive compounds is always CH 2 . Derivative Compounds. If methane be treated with chlorine, the hydrogen may be partly or wholly removed, giving a series of chlorine compounds with hydrogen. Thus: CH 4 + C1 2 -^ CH 3 C1 + HC1. By further treatment we may remove the other hydrogen atoms, forming successively CH 2 C1 2 , CHC1 3 , CC1 4 . We call such compounds derivatives or substitution products. It will be readily seen that by using a variety of substances, ^365 366 MODERN CHEMISTRY a vast number of substitution products might be obtained. For example, we might substitute bromine or iodine, or a large number of other substances, for the hydrogen, as we have chlorine. Chloroform and lodoform. Two of the simpler deriva- tives from methane are chloroform, CHC1 3 , and iodoform, CHI 3 , both of which are extensively used in medicine, the first as an anaesthetic, the second as a deodorant or germicide. Preparation. Chloroform is prepared by treating alcohol with bleaching powder and distilling. lodoform is obtained when a solution of sodium carbonate in water and alcohol is heated to 60 or 75 C., and iodine care- fully added. Yellow crystals of iodoform are slowly deposited. Characteristics. Chloroform is a colorless liquid, with a specific gravity of 1.526, boiling point 62 C., pleasant ethereal odor, and produces anaesthesia when inhaled for some time. lodoform is a light-yellow solid, melting point 119 C., peculiar disagreeable odor, and strong germicide. Other Substitution Products. If, instead of substituting an element, such as chlorine or iodine for hydrogen, in such compounds as methane, we introduce a group of elements, as hydroxyl, we obtain another well-known series. Thus : CH 4 gives CH 3 HO, Methyl alcohol C 2 H 6 " C 2 H 5 HO, Ethyl alcohol C 3 H 8 " C 3 H 7 HO, Propyl alcohol C 4 H 10 " C 4 H 9 HO, Butyl alcohol C s H i2 " C 5 H n HO, Amyl alcohol 0,11*+, " C n H 2ra + 1 HO, General formula for alcohols It will be seen that the alcohols of this series differ in APPENDIX D 367 their formula by CH 2 , as was the case with the hydro- carbons named above. Ethyl Alcohol. This is the best known of the alcohol series, and is often called grain alcohol or spirit of wine. Fermentation. The commercial supply of alcohol comes from the fermentation of that variety of sugar known as glucose or grape sugar. The chemical change may be represented thus : Ordinary or cane sugar will not ferment unless first in- verted into grape sugar. The process of fermentation is caused by germs found in the air in abundance. As these germs differ, so we have different kinds of fermentation. One of the most familiar is that of lactic acid fer- mentation, in which milk sugar is changed into lactic acid. Another species of ferments, as these germs are called, has the power of breaking up very dilute solutions of alcohol and changing it into vinegar or acetic acid. It is thus that the farmer makes his cider vinegar. A third variety is that found in yeast, and is known as vinous ferment. It has the power of converting sugar into alcohol, with carbon dioxide set free. This has been shown above in the equation. Distillation. The alcohol obtained by the inversion of the starch in grain to grape sugar and subsequently fermenting it, is not only impure, but contains a large amount of water. By repeated fractional distillation and filtering through charcoal a product about 96 per cent alcohol is obtained. To remove this remaining 4 per cent of water and obtain absolute alcohol is somewhat difficult, and such dehydrating agents as lime and anhydrous copper sulphate are used. 368 MODERN CHEMISTRY Characteristics of Alcohol. Ethyl alcohol is a colorless liquid of pleasant odor. It has a boiling point of 78.3 C., and is readily solidified by liquid air. Like paraffine and the various resins, its melting point is rather indefinite. Upon being removed from a liquid air bath it is first brittle, but gradually softens, becomes pasty, and finally melts. It is an excellent solvent, and is thus used in preparing various tinctures for medicines, perfumes, ex- tracts, varnishes, etc. It burns with a slightly luminous flame, and is thus used often for heating purposes. Methyl Alcohol. This is also known as u wood spirit." It is obtained in impure form when wood is distilled in the preparation of charcoal. Characteristics. Wood alcohol is a colorless liquid with a -rather unpleasant odor. Its boiling point is 66.7 C. It is even more poisonous than ethyl alcohol, is an excellent solvent for resins and similar substances, and being cheaper than ordinary alcohol is commonly used for such purposes. Organic Acids. A series of acids corresponding to the alcohols is known, the two most common of which are formic and acetic acid. The relation is shown below. ALCOHOLS ACIDS CH 3 OH . . Methyl CH 2 O 2 . . Formic C 2 H 5 OH . . Ethyl C 2 H 4 O 2 . . Acetic C 3 H 7 OH . . Propyl C 3 H 6 O 2 . . Propionic C 4 H 9 OH . . Butyl C 4 H 8 O 2 . . Butyric By examining the formulae of the acid series, it will be seen that each one has two less atoms of hydrogen than the corresponding alcohol, and one more of oxygen. Theoretically, the acid in each case may be formed by the oxidation of the alcohol, thus : C 2 H 5 HO + 2 -> H 2 + C 2 H 4 2 . APPENDIX D 369 Formic Acid. In nature formic acid is found in the bodies of red ants, from which it receives its name, and in some nettles. It is a colorless liquid, which blisters the skin, when pure, causing intense pain. Acetic Acid. Dilute acetic acid, more or less impure, is known to every one in the form of vinegar. To a limited extent it is often prepared by allowing the juice of apples or other fruit to ferment. The germ which brings about this decomposition is popularly known as "mother of vinegar," and technically, as Mycoderma aceti. Poor wine is often made into vinegar by allowing it to remain exposed to the air for a considerable time. If, however, the wine is allowed to pass slowly through tanks filled with shavings which have been treated with mother of vinegar, the process of oxidation is very rapid. Commercial acetic acid is often obtained as one of the products of distilling wood. The impure acid thus formed is converted into sodium acetate by treating the distillate from the wood with sodium car- bonate; then sulphuric acid is added, and the mixture distilled, when acetic acid comes over. Characteristics. Pure acetic acid is a colorless liquid with a boiling point of 119 C., melting point 16.7 C. Such acid is known as glacial acetic. It has a sharp, pene- trating odor and readily takes up moisture from the air. Aldehydes. We have seen above the relation between the alcohols and acids. If the oxidation is not allowed to go so far, we may obtain products between the alcohols and acids. Thus : - ALCOHOLS ALDEHYDES ACIDS CHoOH CH 2 O CH 2 O 2 C,H 5 OH C 9 H 4 C 2 H 4 6 2 C,H 7 OH C 3 H 6 C,H 6 2 CH 3 OH + O -^ CH 2 O + H 2 O CH 3 OH + O 2 -> CH 2 O 2 + H 2 O 370 MODERN CHEMISTRY It will be seen that the aldehydes are alcohols with two atoms of hydrogen removed, while the acids are aldehydes with an additional atom of oxygen. Formic Aldehyde, Methyl Aldehyde, CH 2 0. Formic aldehyde is prepared by passing the vapor of methyl alcohol mixed with air over a red-hot copper gauze or platinum spiral. The hot metal seems to serve as a catalytic agent. Formic aldehyde is a gas, with a boiling point of 21 C. The commercial solution contains about 40 per cent of the gas. It has a very irritating odor, and is used largely as a preservative. Ethers. Their relation to the marsh gas series may be seen from the following: CH 4 . . Methane (CH 3 ) 2 O . . Methyl ether C 2 H 6 . . Ethane (C 2 H 5 ) 2 O . .' Ethyl ether The second of these is the one familiar to students of chemistry and is often called "sulphuric ether." It was given this name from the fact that sulphuric acid is used in its manufacture. The process consists in the careful distillation of a mixture of alcohol and sulphuric acid. The ether distills over and is condensed. There are two reactions : - C 2 H 5 OH + H 2 S0 4 -> ]_j5 S0 4 + H 2 0. This corresponds to the formation of an acid salt when sulphuric acid is treated with caustic soda or potash, thus : NaHO + H 2 SO 4 -> NaHSO 4 + H 2 O. The compound C 2 H 5 HSO 4 is sometimes called the ethyl ester of sulphuric acid. Next, c 2 H 5 OH + c ff 5 y so 4 -> ^HS \ o + H 2 so 4 . APPENDIX D 371 It will be seen, therefore, that the sulphuric acid is regenerated in the second reaction, and only additional amounts of alcohol must be used. Characteristics. Ordinary ether is a very volatile liquid with a peculiar pleasant odor. It boils at 34.9 C. It is an excellent solvent for fats, oils, and various other organic bodies. It is used often as an anaesthetic. Petroleum. Petroleum is a very complicated mixture of hydrocarbons, that from different sources varying greatly. When pumped to the surface, the lower mem- bers of the paraffine series, being gases, escape. In the process of refining, the light oils, rhigoline, gasolene, naphtha, benzine, come over before the kerosene. Later we have the solid paraffines and, in some cases, a residue of asphaltum. To render the kerosene safe it is necessary that the lighter, more volatile oils be carefully removed by distilling. How well this has been done is learned by determining the " flash point " of the oil. CARBOHYDRATES Thus far we have been studying hydrocarbons and their derivatives. We shall now consider a few carbohydrates; that is, compounds containing hydrogen and oxygen in such proportions as to form a certain number of molecules of water. Among these may be named the simple sugars, as glucose ; the complex sugars, as cane sugar ; and starch and cellulose. Glucose C 6 H 12 O 6 Cane sugar .... ^12^22^11 Starch (C 6 H 10 6 ) Glucose. Glucose is ordinarily prepared by treating corn starch with dilute sulphuric acid and removing the excess of acid with calcium carbonate. The solution is 372 MODERN CHEMISTRY then boiled down to a sirup or to the point of crystalliza- tion. The sulphuric acid serves as a catalytic agent and causes the starch to take up a molecule of water, as shown by the following equation : - C.H 10 8 + H 2 -* C 6 H 13 6 Starch Glucose By the action of ferments, as already stated, cane sugar may be changed to glucose and a closely allied sugar, fructose, thus: C 12 H 22 O n + H 2 0-C 6 H 12 6 + C 6 H 12 6 Cane sugar Glucose Fructose It will be noticed that glucose and fructose have the same formula; but glucose turns the plane of polarized light to the right, while fructose turns it to the left. Fructose is about as sweet as cane sugar, while glucose is only about three fifths as sweet. Cellulose, (C 6 H 10 5 ) W . Cellulose belongs to the same class of compounds as starch, each having the same empir- ical formula. It is the basis of all vegetable fiber. Flax, cotton, hemp, rags, paper, etc., consist largely of cellulose. From its formula and what has been said about the prep- aration of glucose, it will be readily seen that any of the above forms of cellulose, or sawdust, for example, might be converted into glucose. The chemical action induced by sulphuric acid is the same as already shown in the inver- sion of the starch. Ethereal Salts. If we examine the formula of an alco- hol, we shall see that it resembles somewhat that of a metallic hydroxide. Thus : C 2 H 5 HO . . . . Ethyl alcohol NaHO Sodium hydroxide The one is the hydroxide of an organic radical, the other APPENDIX D 373 of a metal. In some respects other than in the one named the two classes of compounds are similar. As has been shown, acids may react with the alcohols, forming salts. For example, we have seen C 2 H 5 HO + H 2 S0 4 -> C 2 H 6 HSO 4 + H 2 O. Such compounds we often call ethereal salts. Ethyl Acetate, C 2 H 5 C 2 H 3 2 . This compound, also called acetic ether, may be prepared by treating alcohol with acetic acid, as shown by the following reaction : C 2 H 5 HO + CHgCOOH -*- C 2 H 5 CH 3 COO + H 2 O. It will be seen that the reaction is similar to that of an alkaline hydroxide with an acid. Ethyl acetate is a color- less, neutral liquid, of a rather pleasant odor, with a boiling point of 73 C. It is used somewhat as a medicine. Glycerine, C 3 H 5 (HO) 3 . Glycerine is a triacid alcohol; that is, an alcohol with three replaceable hydroxyl groups. It is a by-product of the manufacture of soap. It is a familiar, colorless, sirupy, sweet-tasting liquid, readily soluble in water. Glyceryl Salts. Such fats as those found in beef and mutton suet and others are mixtures mainly of three ethe- real salts, in which glycerine as the base is combined with certain fatty acids. For example, beef fat consists mainly of stearin, palmitin, and olein. Stearin . . . C 8 H 6 (C 17 H M COO) 8 Palmitin . . . C 3 H 5 (C 15 H 31 COO) 3 Olein .... C 3 H 5 (C 17 H 33 COO) 3 Saponification. When the fats are treated with an alkali, as caustic soda, the ethereal salt is decomposed, free glycerine is formed, and a metallic salt of the organic acid, known as soap. Thus : 374 MODE11X CHEMISTRY C 3 H 6 (C 17 H 35 COO) 8 + 3 NaHO -> C S II S (HO), Glycerine + 3 NaC 17 H 35 COO Soap (sodium stearate) Chemical Action of Soap. When soap is dissolved in water, being a salt, it is very largely ionized, as all salts are. At the same time some molecular sodium hydroxide and stearic acid will form from union with the ions from the water. But sodium hydroxide, being a strong base, is almost completely ionized, while stearic acid, being exceed- ingly weak, is ionized very slightly. Hence we have pres- ent a very large number of sodium ions, which are the active principle in cleansing by such soap. When soap is put into hard water, the salts in the water react with the soap, forming a calcium salt, for example, with the organic acid. This is a compound insoluble in water and it ap- pears as a precipitate upon the surface of the water, and we speak of it as scum. The chemical action may be rep- resented thus : - CaSO 4 + 2 NaC 17 H 35 COO -> Na 2 SO 4 + Ca(C 17 H 35 COO) 2 APPENDIX E PROBLEMS CHAPTER VI 1. Find the molecular weight of crystallized zinc sul- phate, which contains seven molecules of water of crystalli- zation. Also, of crystallized magnesium sulphate, having the same amount of water. 2. In 5 g. magnesium sulphate crystals how much water of crystallization ? 3. How many grams of oxygen can be obtained from 2 kg. of mercuric oxide ? From 28 g. of water ? 4. How many grams of sodium hydroxide can be pre- pared by treating 20 g. of sodium with water ? 5. To prepare 50 1. of oxygen at standard conditions how much potassium chlorate is needed ? 6. What is the per cent of oxygen in potassium chlorate ? 7. How much crystallized zinc sulphate could you pre- pare by dissolving 10 g. of zinc in dilute sulphuric acid ? What weight of hydrogen would be set free at the same time ? What would be its volume ? (Standard conditions.) CHAPTER VII 1. By the method of preparing nitrogen as given in Experiment 41, page 73, what weight of sal ammoniac will be needed to prepare 10 g. of nitrogen ? How many liters of nitrogen would this be, if nitrogen is 14 times as heavy as hydrogen ? 375 376 MODERN CHEMISTRY 2. How much common salt would be obtained in the above experiment ? 3. How many grams of ammonia could be obtained from 1000 g. of ammonium chloride ? How much slaked lime would be needed ? 4. What weight of nitrous oxide could be prepared from 10 g. of ammonium nitrate ? 5. How much nitric acid could be prepared from 50 g. of sodium nitrate ? How much sulphuric acid would be needed ? 6. How much water in one ton of ferrous sulphate (FeS0 4 . 7 H 2 0)? CHAPTER VIII In the proportion, v : v : : t : , where v is the volume of a gas at zero temperature, if we substitute, we have 273 v = v (273 + 0; or v = v (1 + 1/273^) v = v (1 + 0.0036650. If the student prefers, he may use this expression instead of the method he has followed on page 96. If we raise the temperature of a gas without allowing it to expand, we increase the pressure in the same proportion. Hence we may substitute p and p for v and v in the above formula. We then have p=p (1 + 0.0036650- 1. If the gauge of a steam boiler showed a pressure of 4 atmospheres, what would be the temperature of the steam contained in the boiler ? 2. If a balloon rises to a height where the barometer registers half what it does at sea level, what must be the temperature of the atmosphere if the gas in the balloon has not changed in volume ? APPENDIX E 377 3. What must be the pressure in a boiler if the temper- ature has been sufficient to melt the fusible plug, point of fusion being 225 C. ? 4. If a boy pumps air into a bicycle tire, already filled under a pressure of 760 mm., until it is under a pressure of 5 atmospheres, assuming that the tube does not expand, theoretically how much would the air rise in temperature ? 5. A certain amount of potassium chlorate yielded 27.3 1. of oxygen in a room where the barometer read 740 mm. and the thermometer 23. What would be the volume under standard conditions ? (1 1. O = 1.43 g.) 6. I collected over water 950 cc. of hydrogen in a room at a temperature of 25 and a pressure 750; find the true volume of gas under standard conditions, allowing for aqueous tension. (See page 404, section 22.) 7. How much oxygen should I obtain from 5 g. of potassium chlorate, collecting over water in a room at a temperature of 200 C. and barometer reading 740 mm. ? 8. If air is 14.44 times as heavy as hydrogen at stand- ard conditions, what must be the pressure that the density may be the same as that of hydrogen ? 9. I have a liter flask which is able to withstand an internal pressure of 4 atmospheres. It is filled with oxygen at 740 pressure and 11 C. To what temperature must the gas be raised to cause the pressure to break the flask ? 10. A liter flask is filled with oxygen under normal temperature and pressure. To what temperature must the gas be raised to lower its density from 16 to 14 ? What will be the volume of the gas then ? 11. To what temperature must a volume of hydrogen be raised that its density shall be one half what it is at standard conditions ? 378 MODERN CHEMISTRY CHAPTER IX 1. What weight of chlorine could be prepared by using 50 g. of pyrolusite (80 per cent MnO 2 ) with strong hydro- chloric acid ? 2. What weight of chlorine is needed to combine with 10 g. of hydrogen? What weight of hydrochloric acid would be obtained ? 3. What is the greatest weight of hydrochloric acid that could be prepared from 50 kg. of common salt ? 4. How much bromine would be set free from a solution of potassium bromide by means of 5 g. of chlorine ? How much iodine from potassium iodide ? 5. What weight of iodine could be obtained from 1 kg. of sodium iodide ? How much manganese dioxide would be needed in setting the iodide free by the usual method ? 6. With chemicals at the following prices, manganese dioxide, 40 c.; hydrochloric acid, 36 per cent, 8 c.; com- mon salt, 1 c. ; and sulphuric acid, 10 c. per pound, which is the cheaper method of making chlorine : by using manganese dioxide and hydrochloric acid, or from salt, sulphuric acid, and manganese dioxide, provided the by-products are not used in either case ? CHAPTER XI 1. What weight of carbon dioxide can be obtained from 1 ton of limestone, 80 per cent pure ? 2. What weight of calcium chloride would be obtained at the same time ? 3. If 60 kilos of marsh gas are burned, what weight of carbon dioxide and of watel* would be obtained ? 4. If 60 kilos of ethylene are burned, what weight of air would be necessary for perfect combustion ? APPENDIX E 379 5. How much acetylene can be obtained from 5 Ib. of calcium carbide ? What weight of calcium hydroxide ? 6. With carbide worth 10 c. per kilogram, knowing acetylene to be thirteen times as heavy as hydrogen, what would be the cost of acetylene per liter ? 7. If the equation, 3 C + SiO 2 -> 2 CO + SiC, represents the preparation of carborundum in the electric furnace, what weight of sand is necessary to prepare 500 kg. of carborundum ? 8. One hundred cubic centimeters of water at C. and one atmosphere pressure will dissolve 180 cc. of car- bon dioxide, and under 4 atmospheres, 720 cc. If you drink a glass of soda water, 500 cc., with the temperature of the water zero, how much gas would you drink provided none escaped on being drawn from the faucet, assuming the pressure to be 4 atmospheres ? What would be the weight of the gas if it is twenty-two times as heavy as hydrogen ? 9. If a spherical balloon, 8 m. in diameter is filled with a mixture of hydrogen and marsh gas, half of each, marsh gas being eight times as heavy as hydrogen, what will be the buoyancy of the balloon if the weight of the car and the balloon itself is 10 kg. ? 10. When carbon dioxide is passed into lirnewater, this reaction takes place : Ca(HO) 2 + CO 2 >CaCO 3 + H 2 O. If 20 1. of the gas are passed through lime water what weight of dry calcium carbonate would be obtained? 11. An analysis of a sample of coal shows that it contains: carbon, 83 per cent ; hydrogen, 5 per cent ; sulphur, 1.5 per cent; oxygen 2.5 per cent; nitrogen 1 per cent ; ash, 6 per cent. If air by weight is 23 per cent oxygen, how much air by weight and volume will it require to burn 1000 Ib. of the coal ? Assume that the 380 MODERN CHEMISTRY hydrogen burns to form water, the carbon to carbon dioxide, and the sulphur to sulphur dioxide. The ash and nitrogen do not burn. CHAPTER XII 1. In Experiment 102, page 160, what weight of salt should be obtained from the one gram of sodium car- bonate ? What weight if dry sodium carbonate were used ? 2. In Experiment 106, page 164, what weight of copper nitrate will be obtained from 2| g. of copper ? What weight of nitric acid theoretically would be needed ? What should be the weight of the copper oxide obtained ? 3. To neutralize 500 cc. of hydrochloric acid made to contain 3.65 g. of pure acid per liter, what weight of pure sodium hydroxide would be needed ? 4. I used 50 cc. of sodium hydroxide solution, made to contain 4 g. of pure hydroxide in 100 cc., to neutralize 250 cc. of commercial acetic acid. If its specific gravity is 1.014, what is the strength of the acid (grams in 100 g. of acid) ? 5. How much silver nitrate will be needed to precipi- tate the chlorine in 5.85 g. of common salt ? 6. If I dissolve 1.7 g. of silver nitrate in 1000 cc. pure water and then use 100 cc. of this solution to precip- itate the chlorine in 500 cc. of spring water, how many g. of salt, assuming the chlorine to be present in the form of salt, would there be in a liter ? 7. What weight of zinc will be necessary to displace 1 g. of hydrogen from sulphuric acid ? From hydro- chloric acid ? 8. What weight of magnesium would be needed for the same purpose ? Of aluminum ? APPENDIX E 381 CHAPTER XIH 1. What weight of sulphur could be obtained from 1000 kg. of iron pyrites if heated in sealed retorts ? 2. What weight of sulphur dioxide would be obtained from 1000 kg. of iron pyrites if heated with plenty of air ? 3. How much sulphuric acid could be prepared from the sulphur dioxide obtained from 1000 kg. of pyrites ? 4. How much air in grams is necessary to roast the pyrites for the sulphur dioxide in problem 3, if air is 14.5 times as heavy as hydrogen ? It may be assumed that the air is 23 % oxygen. 5. How much sodium nitrate is necessary to prepare 100 g. of nitric acid ? What weight of sulphur dioxide would this amount of nitric acid oxidize ? 6. For every kilogram of sulphur used in making sulphuric acid, what weight of sulphur dioxide would be obtained ? What weight of nitric acid would be neces- sary to oxidize it ? What weight of sulphur trioxide would be obtained ? What weight of steam would be required to convert the trioxide to sulphuric acid ? CHAPTER XVI 1. This equation has been found to be true : 2 mol. hyd. -f- 1 mol. oxy. > 2 mol. water, or 2H 2 + 2 -^2H 2 0. If 100 cc. of hydrogen are exploded with the required amount of oxygen, what volume of steam would be pro- duced ? 2. It has been found that H 2 +Cl 2 -2 HC1. If one liter of hydrogen is exploded with the necessary amount of chlorine, what volume of hydrogen chloride will be obtained ? 3. This equation is true : 3 H 2 + N 2 -> 2 NH 3 . What 382 MODERN CHEMISTRY volume of hydrogen is required to combine with 100 cc. of nitrogen ? What volume of ammonia would be obtained ? The above problems illustrate the Law of Simple Volumes; that is, that the volumes of gases combining with one another always bear some simple ratio to each other and to the resulting product, if that product is gaseous. 4. To produce as above 2000 cc. of ammonia what volume of nitrogen is necessary ? What volume of hydrogen ? 5. To produce 1000 cc. of hydrogen chloride, what volume of chlorine is needed ? Of hydrogen ? 6. To produce 2000 1. of steam, what volume of hydro- gen is required ? What of oxygen ? 7. What volume of hydrogen would be set free from sulphuric acid by 12 g. of magnesium in a room at C. ? How much at a temperature of 20 C. ? (Pressure normal in both cases.) 8. What would be the volume of acetylene obtained from 1000 g. of carbide if the barometer reads 740 and the thermometer zero ? 9. What would be the volume of the carbon dioxide obtained by burning 1000 g. of coal, 80 per cent carbon, if volume is estimated at 20 C. and the barometer reads 750? NOTE. It has been found that 22.32 1. of hydrogen at standard conditions weigh 2 g., or, as it is spoken of in the chapter on Solution, a gram-molecule. In the same way, according to Avogadro's Law, the same volume of any gas, 22.32 1., should weigh a gram-molecule of that gas. This is found to be true. Thus, 22.32 1. of oxygen weigh 32 g. ; 22.32 1. of carbon dioxide weigh 44 g. ; etc., all being estimated under normal conditions. 10. In problem 1 above, if 22.32 1. of hydrogen are used, what volume of oxygen would be needed ? APPENDIX E 383 11. In problem 3, 33.48 1. of hydrogen would require what volume of nitrogen ? 12. If 3.5 1. of oxygen are mixed with 1.5 1. of marsh gas and exploded, what volume of carbon dioxide will be formed ? What of steam, and how much oxygen will remain ? 13. What volume of air, 21 per cent oxygen, will be needed to completely burn 50 1. of ethylene, and what will be the volume of the carbon dioxide which will be formed ? 14. Determine the same in the case of acetylene. GENERAL PROBLEMS 1. In the process of analyzing a sample of flint glass the lead contained was converted into lead sulphate ; 3 g. of the glass gave 1.25 g. of lead sulphate. Find the per cent of lead oxide in the glass. 2. A specimen of dolomitic limestone, calcium and magnesium carbonate, was found to contain 5 per cent of silica. If a hundred pounds of the sample yielded 43.07 Ib. of carbon dioxide, what were the per cents of calcium and magnesium carbonate in the sample ? 3. A mixture of silver chloride and bromide weighing 2 g. was heated and a stream of chlorine passed over it ; when cooled and weighed again, it showed a loss of .3 g. What per cent of the mixture was silver chloride ? 4. According to Dulong and Petit's Law u all atoms have the same capacity for heat." This means that the atomic weight of any element multiplied by its specific heat is a constant quantity. For example, the specific heat of iron is 0.114; its atomic weight is 56. The prod- uct of these two factors is 6.384. This product for all the elements is approximately 6.4 and is called the atomic 384 MODERN CHEMISTRY heat of the element. The specific heat of lead is 0.0315. Determine its atomic weight. 5. The specific heat of arsenic is 0.083 ; of zinc, 0.0955 ; of platinum, 0.0325. Find their atomic weights. 6. The atomic heat of bromine is 6.7; of iodine, 6.87; of calcium, 6.8. Find their specific heats. 7. The specific heat of magnesium is 0.25. How does this show that the atomic weight is 24 and not 12 ? APPENDIX F LABORATORY SUGGESTIONS 1. Neatness. To the best success in any chemical experiment neatness is absolutely essential; indeed, the merest traces of substances foreign to those with which we are working may cause a complete failure of the experi- ment. A student hardly knows what neatness is until he has had a thorough training in chemical analysis. The apparatus should always be clean when put away, and then before using should be rinsed with pure water. Never lay a cork or stopper down upon the table, as it will gather dust and thus pollute the reagent. If you desire to use some solution contained in a bottle, take the stopper between the first and second fingers with the palm of the hand upward and remove it from the bottle ; then without laying it down seize the bottle with the thumb on one side and the fingers on the other. In this way the stopper will not come in contact with the side of the bottle and soil it, neither will dust and dirt be gathered from the table. The reagent bottles should be frequently wiped, as they soon become more or less covered with deposits which form from the gases gener- ated in the laboratory. The table also should be kept lean, and water and other liquids should not be allowed to remain if accidentally spilled. 385 386 MODERN CHEMISTRY 2. Order. Great advantage will also be secured by having everything in its allotted place. Especially is this true of the reagent bottles, and the more there are of these the more important it is that they should be kept in order. For the larger schools probably about twenty reagent bottles will be furnished each student, and these will be arranged upon two shelves, one above the other. In such case, the following order is suggested as being as good as any : LOWER SHELF Beginning at left hand : Sulphuric Acid . . Hydric Sulphate . . H 2 SO 4 Hydrochloric Acid . Hydric Chloride . . HC1 Nitric Acid . . . Hydric Nitrate . . HNO 3 Ammonium Hydroxide or Hydrate .... NH 4 OH Ammonium Chloride NH 4 C1 Ammonium Sulphide ....... (NH 4 ) S Ammonium Carbonate ...... (NH 4 ) 2 CO 3 Barium Chloride BaCl 2 Potassium Dichromate . Potass. Acid Chromate . K 2 Cr 2 O 7 Potassium Ferrocyanide K 4 FeCy 6 UPPER SHELF Calcium Hydroxide or Hydrate ..... Ca(OH) 2 Mercuric Chloride ....... HgCl 2 Silver Nitrate . . Argentic Nitrate . . AgNO 3 Ferric Chloride ........ Fe 2 Cl 6 Acetic Acid . . . Hydric Acetate . . HC 2 H 3 O 2 Lead Acetate . . Plumbic Acetate . . Pb(C 2 H 3 O. 2 ) 2 Potassium Iodide ........ KI Sodium Carbonate . Crystals or powder . . Na 2 CO 3 Borax, powdered Na 2 B 4 O 7 Some of the above reagents are known by different names, and in such cases two of them, the most common, have been given above. APPENDIX F 387 3. Apparatus needed. Each student should be assigned a locker where he may safely keep the apparatus supplied to him, and for the care of this he should be held respon' sible. The following apparatus is suggested : 3 Test-tubes, 5 x f. 1 Test-tube Brush. 3 Test-tubes, 6 x $. 1 Pair Forceps. 3 Test-tubes, 6 x f. 1 Glass Stirring Rod, 1 Evaporating Dish, small. 1 Blowpipe. 1 Evaporating Dish, medium. 1 Platinum Wire. 1 Beaker, 2 oz. 1 Rubber Cork, one hole. . 1 Small Flask, 2 oz. 1 Rubber Cork, two holes. 1 Delivery Tube. 1 Small Mortar. Directions will be given later for preparing the delivery tube, stirring rod, and some other desirable apparatus. The student should also have the following, and will furnish them himself : An apron, reaching to the ankles. This may be made of denim, oil cloth, or rubber cloth. The last is the most serviceable in many ways, but is the most expensive. A Towel. An Iron Spoon. A Bar of Soap. A Clay Pipe. A Small Magnet. A Candle. A Small Triangular File. The candle will be needed very frequently during the first half of the work in studying the properties of gases. Common Property. In addition to the individual prop- erty assigned above, certain articles on account of their size or for other reasons are used in common. There should be enough of them so that each member of the class may be supplied. Among these may be named ; An Iron Pan, 8 x 14 and about 2J inches (Jeep, to, b used for a pneumatic trough. 388 MODERN CHEMISTRY Test-tube Rack. Iron Ring-stand. Bunsen Burner, with Con- Funnel. nections. Wash-bottle (?). Wire Gauze. Sand Bath. MANIPULATIONS 4. Cutting Glass. To cut tubing, with a sharp-cornered file scratch the glass somewhat deeply where you desire to cut it. Now grasp the tube with both hands, the fingers above, and the thumbs below nearly meeting at the line scratched by the file. Now bend the tube downward and pull strongly apart at the same time. With a little prac- tice good square cuts may be made. The rough ends thus secured will cut any rubber connections used. To prevent this hold them in the Bunsen flame until the glass by becoming softened loses its sharp edges. Sometimes it becomes necessary to cut a bottle or large tube in two ; this may be done in two ways, but both de- pend upon the unequal heating of the glass. Tie around a bottle where you desire to cut it an ordinary twine string ; saturate it with kerosene and ignite it. Some- times it will be found necessary to apply the oil the second time, as soon as the first has ceased to burn, and again ignite it. In this way, if the oil has been applied carefully, a nar- row line extending around the bottle is heated strongly, and if the glass be cooled suddenly by pouring over it cold water, the bottle will be neatly severed. 5. To prepare a Delivery Tube. This may be made of rubber and glass tubing, or of glass alone. The former is often preferable because it allows of more freedom in manipulation. If made entirely of glass, two bends are necessary, and one should be within an inch of the end. Hold the tubing in the Bunsen burner, moving it back APPENDIX F 389 FIG. 70. and forth and rolling it around so as to warm all por- tions equally. When the glass begins to soften, allow its own weight to bend it, and take care that you do not form a right- angled tube, but one of a gentle curve like the elbow of a stove pipe. When the bend has cooled just a little, close the openings at the bot- tom of the burner and hold the glass in the luminous flame until it is well covered with soot. This will cause the glass to cool slowly and hence make it less liable to fracture. Com- plete by making the second bend in the same way, forming an obtuse angle as shown in the figure. If rubber connections are used, a second bend is unnecessary. 6. To make a Jet. Frequently a tube drawn to a fine point is desirable. Take a piece of glass tubing 5 or 6 inches in length and heat as in making a delivery tube. When it begins to soften, draw it slowly apart until a tube of small diameter is obtained at the center, as shown in a in the adjoining figure. When some- what cooled, cut in two at a ; then make a bend in one of the shorter FlG - TL tubes, as shown ir> b. Round off the sharp edges and anneal as previously de- scribed. You will now have two jets, one straight and one bent, for both of which you will find uses. 7. To make a Wash-bottle. Any good-sized bottle or flask will do for this. The tube, a, should be drawn to a. 390 MODERN CHEMISTRY jet as shown in the figure, and after being bent should reach nearly to the bottom of the flask. The other tube after being bent should just reach through the cork. By blowing through >, a jet of water may be directed wherever desired; or if a larger stream is desired, it may be poured out at b. The bottle is more convenient if the jet, a, is attached to the rest of the tube by a rubber tube 2 or 3 inches long ; the stream of water may then be turned readily in any direction. The wash-bottle is indispensable for qual- itative work in washing precipitates. A rubber band should be slipped over the lower end of the tube, a, so that if it strikes the side of the flask in removing the cork and tubing it will not be broken. If the lockers are too small to receive the wash-bottle, one may be used in common by the students working at each laboratory table or section. In such case it is better for each student to have a short tube with rub- ber connections to attach to 6, whenever he desires to use the bottle. 8. To repair a Test-tube. Test-tubes are frequently broken by the beginner, but they may be easily mended, and will then be almost as useful as at first. Hold the broken end in a hot Bunsen burner flame, roll the tube about to heat all sides evenly. When the glass becomes soft, by means of a glass rod, which will cohere to the softened tube, draw off the viscous portion, and thus seal the tube. Usually a small mass of softened glass will re- main upon the end of the tube. This must be drawn off in the same way, until the bottom is very thin, like the rest of the tube. Then by alternately heating and blow- ing into the tube, it may be rounded out and made almost APPENDIX F 391 as perfect as a new tube. After a little practice students may become skillful at this work. 9. Blowpipe Work. In metallurgy, the blowpipe must be used frequently, and two kinds of flames are employed, the reducing and the oxidizing. In preparing for either one, turn down the jet to about a quarter its usual force, or until you have a flame not much larger than that of a good-sized candle, and close the openings at the bottom so as to render it luminous. In the figure, a shows the small luminous flame ready for the use of the blowpipe, 6 shows the reducing flame. The tip of the blowpipe is placed FIG. 73. in the outer edge of the flame, and a gentle but steady stream of- air forced into the flame. In this way a small luminous cone, ?, will remain in about the center of the flame, and in this the metallic oxide should be held. This lumi- nous portion contains red-hot particles of carbon, and they have the power of reducing oxides of metals to the metallic condition. If this luminous cone is not apparent, too much air is being forced into the gas. Either blow more gently, or turn the gas on a little stronger. With a little practice the student will learn to breathe and blow at the same time, and will not find the work especially tiresome. For the oxidizing flame, (?, above, the tip of the blowpipe is placed in the very center of the jet. In this way the air introduced and the gas become thoroughly mixed, and complete combustion ensues. The cone should be perfectly 392 MODERN CHEMISTRY FIG. 74. Collecting over water. non-luminous, and the metal to be oxidized should be held about where n is in the cut. The flame is exceedingly hot, and having an excess of oxygen readily converts into oxides such metals as are oxidizable. 10. Collecting Gases. There are several methods for collecting gases, varying according to the characters of the gases. Those which are insoluble in water are usu- ally collected over water. Students will find an or- dinary baking pan, 2 inches deep and about 6 inches broad by 12 long, sufficiently large. The bottle to receive the gas is first filled with water and inverted over the pan, Fig. 74. This is done by holding tightly a sheet of paper or glass over the mouth of the bottle until inverted and placed under the water in the pan. The delivery tube, T, dips under the bottle and conducts the gas from the generating flask, 6r, into the bottle. 11. Collecting by Downward Displacement. Gases soluble in water obviously cannot be col- lected by the method already described. If it is necessary to have them absolutely pure, mercury is frequently substi- tuted for the water. Ordina- rily, however, if heavier than air they are collected by downward displacement. By this method the bottle is simply left standing upon the table, FIG. 75. APPENDIX F 393 and the delivery tube reaches down into the bottle. Thus the heavier gas is introduced below the air, and gradually displaces it. Such gases as chlorine or carbon dioxide are collected in this way. If the gas is lighter than air and soluble in water, it is usually collected by upward displacement. The receiving bottle is held in an inverted position, and the delivery tube runs up to the bottom of the bottle, gradually displacing the air in the bottle. In Fig. 75, a shows the arrangement for collect- ing by downward displacement, and 6, that for upward displacement. 12. Measurements. Frequent reference is made through- out this work to the cubic centimeter and gram, and the student should have fairly definite ideas of these terms. This can come only by practice. For the volumetric, a test-tube and beaker may be graduated. From a burette run into a test-tube 1 cc. of water ; indicate its height by fastening upon the tube just above the lowest part of the meniscus a narrow strip of mucilage paper. Add another cubic centimeter and mark the height in the same way. Thus grad- uate the tube up to 5 cc. ; mark it also for the 10 cc. Now that the graduation l cc may be permanent, with a file scratch carefully the marks, after which the paper may be removed ; a shows the meniscus for each cubic centimeter, and b the small strip of paper. In the same way graduate a beaker for 5, 10, 15, 20, and 25 cc. As different compounds vary so greatly in density, it is more difficult to obtain an accurate idea of a gram, but the FIG 394 MODERN CHEMISTRY student should be able to approximate it. Put upon one scale pan of a balance a small evaporating dish, and coun- terbalance it with shot or sand upon the other. Then add a gram weight to the shot. Into the evaporating dish now slowly add common salt until the gram weight is balanced. Thus try some other amount, as 2 g. or 5 g. If the classes are large, one portion may be graduating the test-tubes and beakers, while another is doing the gravimetric work. This will greatly expedite matters. 13. Precipitates. A precipitate is any solid matter thrown down in a solution by adding to it some reagent. It may be very dense, so as to be quite jelly-like, or it may form merely a cloud in the solution. To illustrate, put one drop of sulphuric acid into a beaker half or two-thirds full of water and add 2 or 3 cc. of barium chloride solution. The dilute solution should thus give a slight precipitate only. Now powder about a gram of ferrous sulphate and dissolve in as little water as possible, 2 or 3 cc., then add a few drops of ammonia. A thick gelatinous precipitate should form. 14. Decanting and Filtering. These are processes for separating a precipitate from the solution in which it is formed. When the precipitate is one that has considerable density and settles quickly, leaving a clear solution, this supernatant liquid may be decanted or poured off. There is no objection to this method unless the presence of small particles of the precipitate in the decanted portion, or of the solution in the precipitate, will interfere with subse- quent tests. To illustrate, a solution of lead acetate may be precipitated with hydrochloric acid, and after warming slightly and allowing the precipitate to settle, the solution may be decanted. But in cases where the separation must be complete, APPENDIX F 395 Folded Twice Opened FIG. 78. filtration is necessary, that is, passing the solution through a filter paper. There are two ways of folding filters : the simplest, and one used when the precipitate is to be removed from the pa- per, is as follows : fold the paper to form a semicircle, &, then this to form a quadrant, making one fold slightly smaller than the other. This is done because funnels are seldom perfectly made, and one "quarter" will fit them better than another. Usually this is the larger. Now open out one of the quarters, and press down neatly into the funnel. If the quarter tried does not seem to fit, the other one may do so better. Now moisten with a little water, and with the fingers press the paper against the sides of the funnel to remove any air bubbles that may exist there. In filtering, pour in slowly at first, especially if the precipitate is very finely divided. If the solu- tion does not come through clear, it may be necessary to filter again through the same filter paper. The pores will soon become partially filled, and the fil- Dliiillillll trate will be perfectly clear. FIG. 79. 306 MODERtf CHEMISTS? In filtering, the stem of the funnel should always be made to touch the side of the beaker or vessel into which the liquid is being passed, so that no drops may spatter out. Furthermore, in pouring a liquid from any vessel, it should always be allowed to run down a moistened stir- ring rod into the funnel. By observing these precautions, neatness in transferring liquids from one vessel to another will be secured, 15. Opening Bottles. The common acids and aqua ammonia, as well as some other reagents, are frequently put up in bottles with glass stoppers. They are sealed by dipping the stopper into melted paraffin before inserting into the bottle. To remove the stopper the paraffin must be melted. This may be done by turning down the gas- jet moderately low, taking the bottle in both hands, hold- ing the neck over the flame, not too close, and rolling it rapidly around so as to heat all sides alike. Be careful to heat the glass only gently. In a moment or two the wax will be melted and the stopper may be very easily removed. With a little practice bottles may be opened in this way without ever breaking or cracking. Be careful, however, in removing the stopper, never to have the face directly over the bottle. 16. Platinum Wires. These are used in making flame and borax-bead tests for various metals. For the sake of convenience in handling, they are generally fused into a short piece of glass tub- ing. Take a few inches of small-size tubing and FlG> 80 * draw out, as in making a jet such as has already been described for use in testing the combustibility of gases. Cut the glass in two, as before, and insert the platinum wire into the tubing to APPENDIX F 397 a distance of 3 or 4 cm. ; again hold in the flame until the glass is softened. Upon cooling, the wire will be securely fastened in the tubing. (See Fig. 80.) 17. Electrolytic Apparatus. If necessary, the student may prepare his own apparatus for experiments in electrol- y S * S ut ^ ot ^ er a PP aratu s that he will find at hand. Take two pieces of heavy platinum wire, each about a foot long, and make into spiral coils by wrapping around a pencil. Leave two or three inches straight at one end, as shown at a, Fig. 81. Fit to a short-necked bell jar with an open top a rubber cork with two small holes, and support the bell jar upon an iron ring- stand, fastening it securely in position. Next, take two pieces of small glass tubing, each long enough to reach through the cork 0, and extend just into the body of the jar. Insert the straight ends of the two platinum spirals, already made, through these tubes, and fuse the glass at the ends so as to fasten the wires firmly in the glass ; make a small loop in the wire at the lower end. See b in the figure. Insert the two electrodes thus prepared through the holes in the cork, and see that everything is water tight. Next take two burettes with glass stop-cocks and deter- mine accurately the capacity of each below the point of graduation, that is, from m to n in Fig. 82. This must be done if we desire to measure accurately the amount of gas collected. Now by means of clamps support these FIG. 81. 398 MODERN CHEMISTRY FIG. 82. two burettes inverted over the two spiral electrodes, and the apparatus is complete. For use, fill the bell jar with the liquid to be electrolyzed to some distance above the mouth of the burettes. Attach a rubber tube to the tip of the burettes, open the stop-cock, and by suction fill each with the liquid and close the stop-cock. Turn on the current, and the capacity of each burette above the point of graduation having been determined, the amount of gas which collects in each tube is quickly read. Instead of the burettes, test-tubes 8 inches by one-half in diameter may be used with good results, except that the gases cannot be accurately measured. 18. A Simple Electrolytic Apparatus. Occasionally it may be desired to electrolyze a substance without sepa- rating the gaseous products. For such purposes a very simple form of apparatus may be employed, as shown in the figure. Prepare the two electrodes as described for the more complicated form, and fit them to a 3-hole stopper as shown in Fig. 83. Through the other opening pass a bent delivery tube, T, for conducting off the mixed gases which will collect in the top of the bottle when the current is passed. Such apparatus as this may be used to show the explosive character of the mix- ture of hydrogen and oxygen obtained by the electrolysis of water, or of hydrogen and chlorine resulting from the decomposition of hydrochloric acid. To prevent the contents of the bottle becoming too warm, it should be placed in a vessel of cold water. Use hydro- FIG. 83. APPENDIX F 399 chloric acid of specific gravity about 1.1, and allow the current to pass for some time before collecting the gases, in order that the liquid may become saturated with the chlorine. If it is desired to collect bottles of the mixed gases over water, let the water be first saturated with common salt. 19. Eudiometers. The eudiometer is an instrument used to test the composition of mixed gases. The most con- venient form for all purposes is the U-shaped one, in which mercury is used to confine the gases. The air left in one limb of the tube serves as an air cush- ion to receive the shock of the explo- sion. The straight eudiometer, how- ever, is cheaper, and with a few addi- tional attachments may be used satis- factorily. A in the figure is an open-top bell jar, such as has been used in other experiments. The neck of A is closed with a tight-fitting, 1-hole rubber stopper, through which passes a glass tube having an en- largement blown upon the lower end, at B. Another rubber cork, which must fit the eudiometer, E, very tightly, is put upon the glass tube as shown in the figure. This must also fit very tightly. T is simply a piece of glass tubing about one inch in diameter, which should have a capacity somewhat greater than E, It 15 closed at the FIG. H4. 400 MODERN CHEMISTRY lower end with a cork, through which passes a short glass tube. A rubber tube connects the two portions of the apparatus, and just above B is fastened by some fine insu- lated copper wire wrapped about it. For use the eudiometer is filled with water and sup- ported in position over A. The gases to be exploded are introduced separately, and each measured carefully, the eudiometer being held by a paper test-tube holder at such height that the water stands at the same level inside and outside. Now press JB firmly down upon its cork, and lower T as much as possible in order that the confined gases may have the pressure upon them reduced ; grasp the rubber tubing near B firmly with the thumb and finger, and pass the spark. After the explosion, adjust the level inside and outside of E as when the gases were introduced, and measure the residue. If this adjustment cannot be secured by lowering E^ it may remain connected as when the spark was passed, and the level secured by changing the height of T. 20. Aspirators and Aspirating Bottles. As an aid in filtering certain classes of precipitates, an aspirator is fre- quently used. This acts upon the principle of the Sprengel air-pump. The aspirator con- sists merely of two tubes, A and B, secured at right angles to each other. A is attached to a water faucet, and B, by means of heavy- walled rubber tubing, to a filter flask. As FIG. 85. ^ ne wa ter flows through A, the air is gradu- ally withdrawn from the flask ; the pressure being thus removed from beneath the filter containing the precipitate, the liquid is forced through much more rapidly. The filter flask is usually shaped like an Erlenmeyer flask (see Fig. 86), and has a side tube for connecting APPENDIX F FIG. 87. FIG. 86. with the aspirator at B. It is made of heavy glass so as to withstand any ordinary atmospheric pressure. For use it is fitted with a rubber stopper having one hole, through which the stem of a funnel is inserted. In the apex of the funnel is placed a small platinum cone, perforated with minute open- ings. This cone is used to prevent the breaking of the filter paper by the atmospheric pres- sure ; at the same time the numerous small holes permit the outflow of the filtrate with comparative freedom. For certain experiments an aspirat- ing bottle is almost indispensable. For example, suppose the experimenter desires to cause a regular flow of air or of some other gas through some vessel, suitable apparatus is necessary and may be very easily made. Large bottles, holding 3 or 4 liters, will serve best. To each fit a cork with two holes, and insert glass tubing as shown in the accompanying figure. The bent tube, 6r, has attached a short piece of flex- ible rubber tubing, upon which is placed a screw clamp, at H. By means of this the FIG.S& flow of gas issuing 402 MODERN CHEMISTRY from N is regulated. The bottle, M, is placed upon a box so as to elevate it considerably above N. A rubber tube, E, connects the two bottles, and, being flexible, allows of the elevation of either bottle above the other. If you desire to fill N with any gas not soluble in water, place both down upon the table, and fill N completely with water. Open the clamp at H, and insert the cork with the tubing into N. The water will be forced out into (7, and expel the air therefrom ; this done, connect at H with the generating flask (not shown in the figure), after having waited until all air has been expelled from it. By the gas pressure, the water will be forced from N over into MI continue until N is nearly filled, close the clamp at H tightly, and remove the generator. Elevate M to its position upon the box, and the aspirator is ready for use. By simply opening the screw clamp, the siphon connect- ing the two bottles transfers the water from M to N as rapidly as the exit of gas at H will allow. If the gas has been permitted to fill completely the bottle JV, and has forced the water out of the siphon tube, it is only neces- sary to apply a little pressure at D. If a dry gas is de- sired, it must be obtained by passage from N through some suitable drying tube attached at H. If the gas to be used is ordinary air, the action of this apparatus may be made continuous, except for a momentary delay in changing the connections, by placing first M, and then jV, upon the box, and connecting the receiver with the tubes, G- and D, respectively. The apparatus may be used in this way for showing the presence of carbon dioxide in air, by forcing it through lime-water. In other cases, where the amount of gas needed is not in excess of the capacity of the bottle JV", this apparatus will work with entire satisfaction. APPENDIX F 403 21. Gas Generators. It is often desirable to have a generator, automatic in action, which will furnish a steady flow of gas and be ready for use at a moment's notice. Kipp's apparatus meets such a demand ; but at much less expense one which works equally well may be prepared for any laboratory. In the fig- ure, A is a bottle of about 500 cc. capacity, fitted with a cork and tube at P, to keep out dust. Through the bottom at K, with a glass drill, make a hole and insert a rubber cork with one perforation. Through B near the bottom drill a hole and insert a rubber cork with a glass tube and short rubber connection clamped with a Hoffman screw. This is for the purpose of drawing off the spent acid. In the top of B fit a stopper with two holes ; through one of these pass a long tube reaching to the bottom of B and extending up into A. To the other hole fit the bent tube, Z>, which has rubber connections for joining with any other apparatus. When not in use, this is kept *tightly closed with a screw clamp. If you desire to use this apparatus as a hydrogen gener- ator, place a half pound or more of zinc in B, close tightly the screw clamp at D, and pour diluted sulphuric or hydro- chloric acid into A until about two- thirds full. Open the screw clamp ; the acid will run down into the lower bottle and will continue to react with the zinc as long as the ga FIG. 89. 404 MODERN CHEMISTRY has free exit at D. If, however, the clamp is closed, the pressure in B soon becomes sufficient to force the acid up the longer tube into the upper bottle, and the evolution of gas ceases. Ci The bottle, A, is held in position by a clamp at the neck, and rests upon a ring of the support. The holes at K and E may be drilled by using a large file broken off, together with emery dust. To use the generator for hydrogen sulphide or carbon dioxide, the zinc would be replaced with ferrous sulphide or marble. 22. Correction of Barometric Reading. In the various problems given in the text in connection with the Law of Charles, it was assumed without being stated that we were dealing with dry gases. Further than this, in the quanti- tative work with gases, certain corrections have been neg- lected. For exact work, however, in the measurement of gases, not only must the temperature be known, and the barometric pressure as well, but also certain other facts. If the gas has been collected over water, the exact volume will not be obtained .by methods already used, for the reason that the presence of water vapor increases the tension of the gas, and hence the volume. In reducing the volume of gases, therefore, to standard conditions, allowance must be made for this tension. This has been carefully estimated, and for the ordinary range of tempera- ture is shown below : 19 C. . . 16.35mm. 22.0 C. . . 19.66mm. 19.5 C. . . 16.86 " 22.5 C. . . 20.27 " 20.0 C. . . 17.39 " 23.0 C. . . 20.89 " 20.5 C. . . 17.94 " 23.5 C. . . 21.53 " 21.0 C. . . 18.50 " 24.0 C. . . 22.18 " 21,5 C. , 19.07 " 24.5 C. . 22.86 " APPENDIX F 405 To illustrate, suppose we have 40 cc. of gas, the tern- perature of the room being 21 C., the barometric pressure 740. According to the law, stated previously, V. V :.P' :P, V X P' or in which V represents volume under standard pressure P, which is 760, V 1 the given volume of gas under the pres- sure P'. Substituting, rr = 40 x 740 760 But making correction for aqueous tension, we have v _ F'xQP' -p) P in which p is the tension of the aqueous vapor. From the table given above, we find that at 21 C. this is 18.5 mm. Substituting in the formula, we have, F= 40 x (740 -18.5) 760 which will give the true volume of the gas under standard conditions. 23. Drying Tubes. Drying may usually be accom- plished by forcing a strong current of air through the tube by means of a foot-bellows ; if the tube has been previously moistened with alcohol, the process will be materially hastened. In like manner flasks may be dried. By means of rubber tubing connect a glass tube, long enough to reach to the bottom of the flask, to a foot-bel- lows, and direct a strong current of air into the fla,ski 406 MODERN CHEMISTRY 24. Recording Results of Experiments. In the first place, the student should understand exactly what he is expected to learn from the experiment ; then he must know what steps are necessary in order to secure the correct results. Do not make the mistake of drawing conclusions before the experiment is complete, and then endeavoring to make the results conform to your precon- ceived ideas. Learn to see everything that occurs, and draw your conclusions in accordance with what really happens. These results should be recorded in suitable note-books, and, were it possible, always completed in the laboratory. Note the results neatly and concisely in good rhetorical sentences. When they admit of being tabulated, such a form is always desirable. If the notes are not written up in the laboratory, a brief record should be made th re, and at home put into permanent form in the note-book without delay. These records should be examined fre- quently by the teacher, at least after the completion of each distinctive portion of the work ; for instance, in studying the halogen group, when the work in chlorine has been done, the notes should be examined ; after that in bromine is completed, a similar examination should take place. PREPARING SOLUTIONS 25. For ordinary work, reagents which are " commer- cially pure" will do, and are much cheaper. It is better to use distilled water in making up all solutions, but for some, such as caustic potash, soda, and such as form pre- cipitates with water that is more or less "hard," pure water is essential. 26. Acids Hydrochloric, Nitric, and Sulphuric. For ordinary work these acids should be diluted with twice/ APPENDIX F 407 their own volume of water. In the case of the last acid the water must be added very cautiously, as great heat is generated. It is better to take what water is to be used in diluting the acid, and very gradually add the sulphuric acid to it. Acetic acid may also be diluted. When an acid stronger than the one prepared in this way is de- manded, it is so stated in the text. 27. Ammonia. Ordinary aqua ammonia should be diluted with about three parts of water. 28. Ammonium Chloride. This should be made up with about 100 g. of the salt to a liter of water. 29. Ammonium Carbonate. About 200 g. to liter. 30. Ammonium Oxalate. About 40 g. to liter. 31. Ammonium Sulphide. This may be prepared by the teacher if preferred. It is done by taking ammonium hydroxide as diluted above and passing into it a current of hydrogen sulphide until saturated. If yellow ammo- nium sulphide, (NH 4 ) 2 S X , is desired, add to the ammonia at the beginning a little sulphur in the form of flowers. When the solution is saturated, it is customary to add to it about two-thirds as much more of the ammonium hydroxide. 32. Barium Chloride. About 100 g. to the liter. 33. Potassium Dichromate. About 50 g. to the liter. 34. Potassium Ferrocyanide. About 75 g. to the liter. 35. Calcium Hydroxide. Saturated solution. 36. Mercuric Chloride. Saturated solution. 37. Mercurous Nitrate. About 50 g. to the liter, witft about one-twentieth part of nitric acid added. Otherwise a basic salt forms in the solution. It is a very good plan to put a few drops of mercury into the bottle containing the solution. 408 MODERN CHEMISTRY 38. Silver Nitrate. About 50 g. to the liter. Keep the solution in an amber-colored bottle and away from contact with organic substances. 39. Ferric Chloride. About 50 g. to the liter. 40. Ferrous Sulphate. This must be made up as desired. About 100 g. to the liter. 41. Lead Acetate. About 100 g. to the liter. 42. Potassium Iodide. About 50 g. to the liter. OTHER SOLUTIONS USED OCCASIONALLY 43. Arsenic Chloride. Dissolve arsenious oxide, As 2 O 3 , in caustic soda, and then add hydrochloric acid until the solution gives an acid reaction. 44. Antimony Chloride. Add hydrochloric acid to ^water until well acidulated, and then a small quantity of antimony trichloride ; a solution of antimony may be obtained from the antimony tartrate in the same way. 45. Bismuth Nitrate. This must be prepared in the same manner as the antimony chloride. Dissolve a few crystals of the salt in water to which considerable nitric acid has been added. 46. Calcium Chloride. About 100 g. to the liter. 47. Calcium Sulphate. Saturated solution. 48. Cobalt Nitrate. About 50 g. to liter. 49. Chromium Chloride. Prepare as indicated in the text. To a solution of potassium dichromate add about one-twentieth as much hydrochloric acid and a little alcohol, and boil. The green solution obtained will be chromium chloride. 50. Copper Sulphate. About 50 g. to liter. 51. Di-sodium Phosphate. About 100 g. to liter. 52. Potassium Cyanide. About 100 g. to liter. 53. Potassium Chromate. About 50 g. to liter. APPENDIX F 409 54 Potassium Hydroxide. About 100 g. to liter. 55. Sodium Hydroxide. About 100 g. to liter. 56. Magnesium Sulphate. About 100 g. to liter. 57. Sodium Carbonate. About 100 g. to liter. 58. Lead Nitrate. About 100 g. to liter. 59. Stannous Chloride. First add about one-twentieth part of hydrochloric acid to the water, and then about 75 g. of the solid to a liter. It is better to put a piece of granulated tin into the solution. 60. Cochineal Solution. Grind up the solid in a mortar and dissolve in water or in a.10 per cent solution of alcohol. 61. Indigo Solution. Treat about 1 g. of indigo with about 10 g. of sulphuric acid. After standing several days, dissolve the whole in water. 62 . Litmus Solution. Dissolve the blue solid, powdered, in water. 63. Phenol -phthalein. Dissolve about 1 g. in 100 cc. of 50 per cent alcohol. 64. Ammonium Molybdate. Dissolve 15 g. of am- monium molybdate crystals in 100 cc. of aqua ammonia as prepared above. To this add an equal volume of distilled water, and finally 125 cc. of nitric acid, specific gravity about 1.4. SUPPLIES NEEDED. 65. Chemicals. For ten students. Acid, Acetic 1 Ib. " Hydrochloric .... 10 " Nitric 6 " Oxalic k Sulphuric 10 " Tartaric i Alcohol 1 qt. Alum 1 Ib. Aluminum 2 oz. Ammonium Carbonate ... 1 Ib. " Chloride .... 1 Ib. Ammonium Ferric Citrate " Hydroxide . " Nitrate . . " Sulphate " Sulphide Antimony, Metallic . . " Potassium Tartrate Trichloride . . . Arsenic, Metallic Arsenic, Trioxide Barium Chloride loz. 81b. * loz. Jib. i Ib. 410 MODERN CSEMISTRT *' Nitrate ' Mercurous Nitrate V ^" i " Bark Charcoal (( Nickel " i " Bismuth Metallic . : . . . (( 1 OZ. " Nitrate Paraffin 4 " Bleaching Powder Phenolphthalein .... 1 OZ Bromine 2 OZ. Platinum Wire 3ft. Calcium Carbide 3 lb. Phosphorus, Ordinary " Chloride . . 4 " Plaster of Paris 3 " " Fluoride 4 " Potassium Hydroxide sticks ' . 1 " " Sulphate 4 " Iodide 2 " Carbon Bisulphide .... Charcoal, Powdered, animal . " Stick 4 " 4 " 1 doz. Bromide .... Carbonate . . . Chlorate .... l (( I " 1 " " Wood, powdered . Cobalt Nitrate Jib. i " Chromate . . . Cyanide . 4 " 1 " 1 " Dichromate 5 " Copper, Metallic, turnings . . " Nitrate 2 " i Ferrocyanide . . Ferricyanide 1 (C " Oxide I " Metallic .... A " " Sulphate 4 " Nitrate .... 3 " Ether J ' Nitrite .... k " Permanganate . i ( 1 < Sulphocyanide . . i " IS ' Shellac 1 OZ. IS ' Silver Nitrate i lb. 3 < " Filings l < Tetraborate (Borax) . 4 " " Sulphate Carbonate .... 1 " " Sulphide 2 ' Chloride 1 " " Wire i ( Hydroxide, sticks 1 " ' Nitrate . . . . 1 " " Metallic Nitrite 5 " " Nitrate i < Phosphate, Di . . . 5 " i < Sulphide 5 " ' Sulphite 1 " 1 ' Thiosulphate . . . 3 " Starch k " 10 Strontium Nitrate 1 n Sucrar 1 " 1 OZ. 1 " " roll . .... 1 " " Powdered 4 " Tin Metallic 3 " u Sulphate 1 <' " Chloride i " Manganese Chloride .... " Dioxide 4 " 1 " " Tetrachloride .... Turpentine 1 OZ. 1 lb. Marble 2 " Zinc, Granulated 2 " " " Dust . . ... i " " Sheet 1 " . *^ t 4 " 4 " Oxide . 4 " APPENDIX G REFERENCE LIBRARY No text on chemistry can hope to give more than a glimpse at the subject. Naturally, therefore, it should be the aim of every teacher to build up a reference library for the use of himself and students. Among the many good books to be obtained, the following are suggested : Newth's Inorganic Chemistry Longmans. Newth's Chemical Lecture Experiments Longmans. Mendeleeff's Principles of Chemistry Longmans. OstwalcTs Outlines of General Chemistry Macmillan. Ostwalds Foundations of Analytical Chemistry Mac- millan. Walker- Dobbin 1 s Chemical Theory for Beginners Mac- millan. Roscoe and Schorlemmer' s Treatise on Chemistry, Vols. I and II Apple ton. Remsen's Chemistry, Advanced Course Holt. Remssn's Theoretical Chemistry Lee. Ramsay's Experimental Proofs of Chemical Theory Macmillan. Cornishes Practical Proofs of Chemical Laws Longmans. Johnston's Chemistry of Common Life Appleton. Lassar-Cohn's Chemistry of Every-day Life Lippiucott. Ramsay's Gases of the Atmosphere Macmillan. Meyer's History of Chemistry Macmillan. Thorpe's Essays in Historical Chemistry Macmillan. ,411 412 MODERN CHEMISTBT Sutton's Volumetric Analysis Blakiston. Addymaris Agricultural Analysis Longmans. Alembic Club Reprints Chemical Pub. Co., Easton, Pa. Foundations of the Atomic Theory. Experiments on Air. Foundations of the Molecular Theory. Discovery of Oxygen. Elementary Nature of Chlorine. Liquefaction of Gases. Early History of Chlorine. Muirs Heroes of Science Young & Co. Shenstone's Glass Blowing Longmans. Thorpe 's Chemical Preparations Ginn. APPENDIX H BIOGRAPHICAL THE following are among those who have contributed to chemical literature or to the advancement of the science. AGE OF ALCHEMY G-eber. Arabian alchemist of eighth century ; author of several chemical works, and discoverer of aqua regia. Albertus Magnus. Died 1280. Advanced the theory that the metals were composed of water, arsenic, and sulphur. Bacon, Roger. Thirteenth century. English alche- mist. Advocated experimental proof of chemical theory. Inventor of gunpowder. Valentine, Basil. Fifteenth century. Wrote several works on chemistry. Probably a fictitious name of Johann Tholde. APPENDIX H 413 MEDICAL ERA OF CHEMISTRY Paracelsus, a name coined for himself by Theophrastus Bombastus von Hohenheim. Early part of the sixteenth century. By his study and preparation of a large number of medicines, he earned for himself the title, " Father of Medicine." Libavius. Died in 1616. Proceeded with the work begun by Paracelsus. Wrote a Handbook of Chemistry. Van Helmont, Jean Baptiste. 15771644. Discov- ered several gases. Boyle, Robert. 1627-1691. Real founder of the sciences of physics and chemistry. Formulated Boyle's Law, and advanced the true theory as to the composi- tion of matter. Becker, Johann Joachim. 1635-1682. German chem- ist. Author of theory that when a metal burns terra pinguis escapes from it. AGE OF PHLOGISTON Stahl, G-eorg Ernst. 1660-1734. Founder of the phlo- gistic theory of combustion, that all combustible sub- stances contained an unknown something called phlogiston which escaped when the substance burned. It was an outgrowth of Becher's theory. Hoffmann, Christoph Ludwig. 1721-1807. Physicist and chemist. His theory of the reduction of a metal was about the same as that held to-day. He believed that the calces of the metals contained the metals themselves and some other substance, which he called sal acidum. Black, Joseph. 1728-1799. Professor of chemistry in Edinburgh. Discovered carbon dioxide arid proved that 414 MODERN CHEMISTRY the carbonates of the alkalies and alkaline earths are not elements. Cavendish, Henry. 1731-1810. Discovered hydrogen ; studied the composition of water and the air, and made a large number of experiments with the latter. Prepared nitric acid by synthesis. Priestley, Joseph. 1733-1804. Discoverer of oxygen, and strong advocate of phlogistic theory. Scheele, Carl Wilhelm. -- 1742-1786. Discoverer of chlorine; made some investigations in organic chemistry; prepared glycerine and prussic acid. MODERN ERA OF CHEMISTRY This coincides roughly with the nineteenth century. Lavoisier, Antoine Laurent. 17431794. Founder of modern chemistry. Made a beginning in quantitative work, and overthrew the theory of phlogiston. Advanced the idea of the conservation of matter. Gray-Lussac, Joseph Louis. 1778-1850. Author of the law of combination of gases by volume. Made an extensive study of the general properties of gases ; deter- mined the relation between the volume of a gas and its temperature, thus supplementing Boyle's work. Berzelius, Johann Jacob, Baron. 1779-1848. Studied the atomic weights of the elements ; improved the usual methods of chemical analysis, and investigated the law of combining proportions. Proust, Louis Joseph. 1760-1826. Advocated the theory that the elements combine always in definite proportions, now known as the "Law of Definite Proportions." Dalton, John. 1766-1844. Advanced the atomic theory of matter, and formulated the " Law of Multiple Proportions." APPENDIX H 415 BertMlet, Claude Louis. 1748-1822. Made a long series of experiments, studying the behavior of ammonia, hydrogen sulphide, chlorine, and other gases. Davy, Sir Humphry. 1778-1829. Studied the prop- erties of various gases ; proved that the alkalies, caustic soda and potash, are not elements. Dulong and Petit. Early part of nineteenth century. Made a study of the metals. Formulated the law that the specific heats of the metals are inversely proportional to their atomic weights. Dumas, Jean Baptiste Andre. 1800-1884. Made an extensive study of vapor densities. Faraday, Michael. 1791-1867. Succeeded in liquefy- ing many of the gases; studied physical chemistry, and determined the effects of an electric current upon electro- lytes. He formulated the " Law of Definite Electrolytic Action," that an electric current decomposes electrolytes so that equivalent amounts of the substance are liberated at the kathode and anode. LieUg, Justus, Freiherr von. 1803-1873. Studied organic chemistry ; investigated the phenomenon of isomerism. Mendeleeff, Dmitri Ivanovich. Born 1834. Russian chemist. Formulated the " Periodic Law of the Ele- ments." Author of general chemistry. Pictet and Cailletet. Physico-chemists of the present time. They have done much work in producing low temperatures, and in liquefying air, hydrogen, and oxygen. Ramsay, William. Born 1852. Discoverer of argon in 1894. English scientist of to-day. Dewar, James. Born 1842. English scientist of the present time. Has studied carefully low temperatures, 416 MODERN CHEMISTRY Moissan, Henry. French chemist of the present time. Has succeeded in preparing artificial diamonds ; has also studied carefully the properties of liquid fluorine. MEANING OF ALCHEMISTIC TERMS The student in attempting to read the reports of the chemists of the eighteenth century will find much difficulty in understanding the alchernistic terms so universally employed. The following are among those most commonly met with, and are given to encourage the student to read these accounts himself. The Alembic Club Reprints, men- tioned among the books suitable for reference, furnish the most desirable portions of the writings of such investigators as Scheele, Dalton, Priestley, and others. It will be noticed that often several terms are used for the same substance. This was in accordance with the plans of alchemy to keep secret the discoveries and mystify any who might attempt to decipher the records. OLD TERMS PRESENT MEANING Acid ..... Anhydride (oxide). Acid of chalk . . . Carbon dioxide. Acidum salis . . . Hydrochloric acid Aer fixus .... Carbon dioxide. Air Gas. Alkali of tartar . . . Potassium carbonate. Aqua fortis . . . Nitric acid. Aqua regis .... Aqua regia. Azotic gas .... Nitrogen. Blanc d'Espagne . . Bismuth Subnitrate. Calx Oxide. Calx of silver . . . Silver oxide. Colcothar .... Ferric oxide. Dephlogisticated air . . Oxygen. Draco mitigatus . . Mercurous chloride. Fire air . . . Oxygen. Fixed air . . . . Carbon dioxide. Fixed alkali . . . Sodium carbonate. Gas fuliginosurn , . Combustible gas, APPENDIX H 417 OLD TERMS Gas pingue Gas siccum Gas sylvestre Grey calx of lead Hartshorn . Liver of Sulphur Magnesia alba . Marcasite . Marine acid . . . Mephitic air Mercurius calcinatus . Mercurius dulcis Mercurius Niter Mercurius precipitatus per se Mercurius precipitatus ruber .... Mercurius sublimatus Mercurius vitae . Mors metallorum Niter Nitrous air Nitrous gas Phlogiston .... Phlogistic air Pulvis angelicus . Spirit of niter . Spirit of sulphur Spiritus igneo aerius Spiritus salis Terra pinguis Usifur Vital air . Vitriol Vitriolated tartar Volatile alkali . PRESENT MEANING Combustible gas. Combustible gas. Carbon dioxide. Lead sesquioxide. Ammonia. Potassium persulphide. Magnesium carbonate. Ferric sulphide. Hydrochloric acid. Nitrogen. Mercuric oxide. Mercurous chloride. Mercuric nitrate. Mercuric oxide. Mercuric oxide. Mercuric chloride. Antimony oxychloride. Mercuric chloride. Potassium nitrate. Nitrogen dioxide. Nitrogen dioxide. A hypothetical substance, be lieved to exist in all com bustible bodies. Nitrogen. Antimony oxychloride. Nitric acid. Sulphuric acid. Oxygen. Hydrochloric acid. Same meaning as phlogiston Artificial mercuric sulphide. Oxygen. Sulphate. Potassium sulphate. Ammonium Carbonate. 418 MODERN CHEMISTRY TABLE OF THE ELEMENTS AND THEIR ATOMIC WEIGHTS NAME Aluminum Al Antimony Sb Argon A Arsenic As Barium Ba Bismuth Bi Boron B Bromine Br Cadmium Cd Caesium Cs Calcium . Ca Carbon . . C Cerium Ce Chlorine Cl Chromium Cr Cobalt ... o .... Co Columbium Cb Copper Cu Erbium E Fluorine F Gadolinium Gd Gallium . Ga Germanium Ge Glucinum Gl Gold Au Helium He Hydrogen H Indium In Iodine I Iridium Ir Iron Fe Krypton Kr SYMBOL ATOMIC WEIGHTS O = 16 27.1 120. 39.9 75. 137.4 208.5 11. 79.96 112.4 133. 40. 12. 140. 35.45 52.1 59. 94. 63.6 166. 19. 156. 70. 72. 9.1 197.2 4. 1.01 114. 126.85 193. 56. 81.8 APPENDIX H 419 TABLE OF THE ELEMENTS AND THEIR ATOMIC WEIGHTS Continued NAME SYMBOL ATOMIC WEIGHTS H = Lanthanum La Lead Pb Lithium Li Magnesium Mg Manganese Mn Mercury Hg Molybdenum Mo Neodymium Nd Neon Ne Nickel Ni Nitrogen ... = ... N Osmium Os Oxygen O Palladium Pd Phosphorus P Platinum Pt Potassium K Praseodymium Pr Rhodium Rh Rubidium Rb Ruthenium Ru Samarium Sm Scandium Sc Selenium Se Silicon Si Silver Ag Sodium Na Strontium Sr Sulphur S Tantalum Ta Tellurium Te Terbium Tr 138. 206.9 7. 24.36 55. 200.3 96. 143.6 20. 58.7 14.04 191. 16. 106. 31. 194.8 39.15 140.5 103. 85.4 101.7 150. 44.1 79.1 28.4 107.93 23.05 87.6 32.06 183. 127. 160. 137.6 205.36 6.97 24.1 54.6 198.50 95.3 142.5 9 58.25 13.93 189.6 15.88 106.2 30.75 193.4 38.82 139.4 102.2 84.75 100.9 149.2 43.8 78.6 28.2 107.11 22.88 86.95 31.83 181.5 126.5 158.8 420 MODERN CHEMISTRY TABLE OF THE ELEMENTS AND THEIR ATOMIC WEIGHTS Continued NAME SYMBOL ATOMIC WEIGHTS O = 16 Thallium Tl Thorium Th Thulium Tm Tin Sn Titanium Ti Tungsten W Uranium U Vanadium V Xenon X Ytterbium ....... Yb Yttrium Y Zinc Zn Zirconium Zr 204.1 232.5 171. 118.5 48.1 184. 239.5 51.2 128. 173. 89. 65.4 90.7 202.61 230.8 169:4 118.1 47.8 182.6 237.8 51.0 ? 171.9 88.3 64.9 89.7 The above table shows two columns of atomic weights; the first as- sumes O = 16 as the standard, the second, H = 1. GLOSSARY OF CHEMICALS AND MINERALS Agate. A variety of quartz, occurring often in variegated colors, arranged concentrically. alabaster. A fine-grained, white variety of gypsum, alum. A double sulphate, of general formula, M 2 R 2 (SO 4 ) 4 24 H a O. alumina. Aluminum oxide, A1 2 O 3 . amethyst. A variety of quartz, anthracite. Natural coal, possessing little or no oil or other volatile : products. Hard coal. antichlor. A reagent used to neutralize chlorine when in excess, aragonite. A variety of calcite, CaCO 3 . argentite. Native silver sulphide, arsenic. The popular name for arsenic trioxide. arsenious acid. Another name for arsenic trioxide. arsine. Hydrogen arsenide, AsH 3 . azurite. An ore of copper, blue in color, composition Cu(OH) 2 , 2 CuCO 3 . Baryta. Barium oxide, baryta water. Barium hydroxide, bauxite. A hydrated oxide of aluminum, A1 2 O 3 , H 2 O, used as a source for aluminum. benzine. A light oil obtained from petroleum, bicarbonate of soda. Cooking soda, NaHCO 3 . bismuth ocher. Bismuth oxide, Bi 2 O 3 . bismuthite. Native bismuth sulphide. bituminous. Containing bitumen or oil. Applied to soft coals, blanc de fard. Bismuth subuitrate, BiONO 3 . blende. Native zinc sulphide, blue vitriol. Copper sulphate, borax. Sodium tetraborate, Na 2 B 4 O 7 . braunite. Native Mn 2 O 3 . butter of antimony. An old name for antimony trichloride. Calamine. An ore of zinc, Zn 2 SiO 4 , H 2 O. 421 422 MODERN CHEMISTRY calchopyrite. A sulphide of iron and copper, Cu a S, Fe 2 S 3 . calcite. Crystallized calcium carbonate. calomel. Mercurous chloride, Hg 2 Cl 2 . carbonado. A variety of diamond occurring in black pebbles 01 masses. carborundum. A hard substance, made by combining, at high tempera- tures, silica and carbon. cassiterite. Native stannic oxide, SnO 2 , the chief ore of tin. caustic potash. Potassium hydroxide. caustic soda. Sodium hydroxide. celestite. Strontium sulphate. cement. A variety of lime prepared from limestone containing from 40 to 50 per cent of slate. chalcedony. A variety of quartz. chalk. A soft variety of limestone, composed of the shells of diatoms. chloride of lime. A common name for bleaching powder. chrome alum. A sulphate of potassium and chromium. chrome red. Basic lead chromate, Pb 2 CrO 5 . chrome yellow. Lead chromate, PbCrO 4 . cinnabar. The chief ore of mercury, HgS. clay. A hydrated silicate of aluminum, containing various impurities. colcothar. Ferric oxide, Fe 2 O 3 . copperas. Ferrous sulphate. corrosive sublimate. Mercuric chloride, HgCl 2 . corundum. Anhydrous alumina, uncrystallized. cryolite. A fluoride of sodium and aluminum, NaAlF 4 . Dolomite. A native carbonate of magnesium and calcium. Emerald. (Oriental.) Crystallized alumina, green in color. emery. Massive, opaque alumina. epsom salts. Magnesium sulphate. euchlorine. A solution of chlorine in water. Fat lime. Lime made from pure limestone. feldspar. A silicate of potassium and aluminum, which, decomposed, forms clay. fool's gold. Ferric disulphide, FeS 2 . fuller's earth. A variety of clay. fuming liquor of Libavius. Anhydrous stannic chloride. Galena. The chief ore of lead, PbS. green vitriol. Ferrous sulphate. GLOSSARY 423 gypsum. Native calcium sulphate. Hartshorn. An old term for ammonia. heavy spar. Native barium sulphate. hematite. An important ore of iron, of the composition FegOg. horn silver. Native silver chloride. hydraulic cement. Lime containing from 10 to 30 per cent of silica, having the property of hardening under water. hypo. The photographer's name for sodium thiosulphate. Iceland spar. A transparent, crystalline variety of calcium carbonate. infusorial earth. A grayish white earth, composed largely of silica, resulting from the secretion of diatoms. Jeweler's rouge. An oxide of iron, red in color, used in polishing and as a pigment. Kaolin. A pure variety of clay, formed by the decomposition of feld- spar. kelp. The ashes of seaweeds, used as a source of certain potash salts and of iodine. kerosene. Popularly called coal oil. An oil obtained by the distilla- tion of petroleum. kieserite. Native magnesium sulphate. kupfer nickel. Nickel arsenide, NiAs. Labarraque's solution. Sodium hypochlorite. lac sulphuris. Sulphur precipitated from a solution of it in lime- water. laughing gas. Nitrous oxide, N 2 O. lean lime. Lime made from impure limestone. lime. Calcium oxide, CaO. limestone. Calcium carbonate, uncrystallized. lime-water. Calcium hydroxide. litharge. Impure lead oxide, PbO. lunar caustic. A commercial term for silver nitrate. Magnesia. Magnesium oxide. magnesite. Native magnesium carbonate. magnetic pyrites. A mixture of FeS and Fe 2 S 3 . This mixture is given its name because of magnetic properties. malachite. An ore of copper, CuCO 3 , Cu(OH) 2 . marble. Crystallized limestone. marcasite. A variety of ferric sulphide, FeS 2 . massicot. Lead oxide, PbO. 424 MODERN CHEMISTRY milk of lime. Calcium hydroxide, containing more or less lime in suspension. milk of sulphur. Same as lac sulphuris. minium. Red lead, Pb 3 O 4 . mispickel. An important ore of arsenic, FeSAs. Naphtha. A light oil, obtained from petroleum. Nessler's solution. A solution used in testing for ammonia, niter. Another name for potassium nitrate. Nordhausen's acid. The same as fuming sulphuric acid, H 2 S 2 7 . Oil of vitriol. Sulphuric acid. opal. A variety of silica, SiO 2 . oriental. A term applied to the true emerald and certain other gems, to distinguish them from less valuable stones similar in appearance. orpiment. A sulphide of arsenic, yellow in color, having composition Paraffin. A wax obtained in the later distillation of petroleum. Paris green. A popular name for Scheele's and Schweinfurth's green, compounds of arsenic, pearl ash. Pure potassium carbonate. pearl white. Bismuth oxychloride, BiOCl. petroleum. Rock oil, found native in various parts of the world, plaster of Paris. Calcined calcium sulphate. plastic sulphur. A dark-colored, allotropic form of sulphur, somewhat resembling rubber. potash. Another name for commercial potassium carbonate. Also a loose name for potassium chlorate. powder of Algaroth. A variable compound of antimony, approximately SbOCl. purple of Cassius. A purplish-colored precipitate obtained in testing a solution of gold with stannous chloride. pyrites. A common name for ferric sulphide, FeS 2 . pyrolusite. Native manganese dioxide. Quartz. Silicon dioxide. quicklime. The same as lime. Realgar. Red sulphide of arsenic, As 2 S 2 . red lead. The same as minium. red precipitate. Mercuric oxide. rose quartz. A variety of quartz, somewhat pink in color. Sal ammoniac. Ammonium chloride. GLOSSARY 425 sal soda. Commercial sodium carbonate. salt. A compound formed by the union of an acid and a base. salt cake. Sodium sulphate. saltpeter. Potassium nitrate. sapphire. Crystallized alumina. Scheele's green. Copper arsenite, CuHAsO 3 . silica. Silicon dioxide. slaked lime. Lime treated with water. smalt. A silicate of cobalt and potassium. smoky quartz. A variety of silica, brown or smoky in color. soda. Same as sal soda. soda, cooking. Same as sodium bicarbonate, NaHC0 3 . spathic iron. Native iron carbonate, FeCO 3 . specular iron. A variety of hematite. spiegeleisen. A variety of iron containing manganese and carbon. stibine. Same as antimoniureted hydrogen, SbH 3 . strontianite. Native strontium carbonate. subnitrate of bismuth. Basic bismuth nitrate, BiONO 3 . sugar of lead. Lead acetate. Topaz. Crystallized alumina with small quantity of coloring matter Vermilion. Artificial mercuric sulphide. White arsenic. Arsenic trioxide. white lead. Basic lead carbonate, used as a paint. white vitriol. Zinc sulphate. witherite. Native barium carbonate. Zinc white. Zinc oxide, ZnO, used as a paint. GLOSSARY OF TECHNICAL TERMS IN CHEMISTRY Acidify. To make acid. acidulate. To add acid to, until no longer alkaline or neutral. actinic. Referring to light rays, having the power to effect chemical changes. air-bath. A small oven used for drying substances, alkali. A compound of hydrogen, oxygen, and some metallic element, soluble in water, having the power to neutralize acids ; as caustic soda, NaOH. allotropic. Literally, another form; a term applied to the unusual form of an element. 426 MODERN CHEMISTRY allotropism. The phenomenon of existing in two or more forms, alloy. The product resulting from fusing together two or more metals. amalgam. An alloy, one constituent of which is mercury, amorphous. Without any special form, uncrystallized, massive, anaesthetic. An agent used to produce insensibility. anhydride. An oxide, usually non-metallic, which forms some acid upon the addition of water, anhydrous. Without water. An anhydrous salt is one from which the water of crystallization has been removed, anion. A negative ion. See ion. antiseptic. A substance used to prevent decay, or to destroy noxious germs. argentiferous. Silver-bearing. aspirator. Apparatus used to secure the passage of air or any other gas through certain vessels. assay. Determination of the quantity of the various constituents of a metallic ore. Basic. Having the properties of an alkali or base, binary. A compound consisting of two elements. brightening. The sudden brilliant appearance of the silver assay when the lead has all been removed by cupellation. bumping. A term applied to the violent boiling of the liquid in a vessel, causing it to jump, burette. A graduated tube, with stop-cock, used in volumetric work for measuring accurately a liquid. Calcine. To heat strongly. carbureting. Adding hydrocarbon compounds to an illuminating gas, as in making water gas. cathion. An electropositive ion. cementation. An old process of making steel by imbedding wrought iron in powdered charcoal and heating several days, chemism. The so-called affinity that one element or substance has for another. commercial. A term applied to chemicals as usually furnished to the trade; not absolutely pure; in distinction from chemically pure reagents. concentrated. Strong; undiluted, converter. A large, egg-shaped furnace, used in making steel from cast iron and in purifying copper. GLOSSARY 427 c. p. Chemically pure. crucible. A small vessel, made to withstand great heat. Named from the Latin word crux, because the old alchemists thus marked their crucibles. crystalline. Composed of crystals. cupel. A small cup, made of bone ashes ; used by assayers in deter- mining the gold and silver in an ore. cupellation. The process of separating lead and silver by the oxida- tion of the former. Decant. To pour off the liquid from a precipitate, after the latter has settled. decrepitate. To burst in pieces with a crackling sound, as many salts do when heated with the blowpipe. deflagrate. To burn vigorously. deflagrating spoon. A small metallic cup or spoon with a long wire handle attached. Used for holding combustible substances when burning in oxygen or other gases. deliquesce. To take up moisture from the air. deoxidizing agent. See reducing agent. desiccate. To dry. desiccator. A vessel used in drying or keeping dry a substauce which is to be weighed accurately. destructive distillation. The process of heating in closed retorts a substance to such a temperature as to effect its decomposition. digest. To warm gently. disinfectant. A substance used to cleanse and purify unwholesome places, as well as to destroy disease germs. displacement. A method of collecting a gas in a vessel filled with air, or some other gas, depending upon the difference in density of the two. dissociate. To break up a compound body into parts. distill. To evaporate a liquid and condense again in another vessel. distillate. The liquid obtained in the process of distillation. dyad. An element having a valence of two. Ebullition. Rapid boiling. effervescence. The act of bubbling, as seen upon the application of an acid to a carbonate. effloresce. To give up at ordinary temperatures the water of crystal- lization. 428 MODERN CHEMISTRY electrode. The terminal of a battery. electro-positive. A term applied to elements attracted to the negative electrode. equivalence. A term sometimes used instead of valence, escharotic. An agent which corrodes or destroys ; a caustic, evolve. To set free. excess. A quantity more than sufficient to secure certain chemical action. Filtrate. The liquid obtained after passing through the filter paper, in removing the precipitate. fixed. The opposite of volatile. flocculent. Flaky. flux. Any substance used to lower the melting point of another ; as limestone with iron ore in the blast furnace. formula. A combination of symbols used to represent a molecule of a compound body. fractional distillation. The process of separating by distillation the several constituents of a mixture of liquids, by means of their different boiling points. Gangue. The impurities contained in an ore or mineral. gelatinous. Like starch paste in appearance. generate. To produce or set free, as a gas. germicide. A substance used to destroy bacteria or germs. granulated. In irregularly shaped small particles, secured by pouring the fused metal into cold water. graphitoidal. Resembling graphite. gravimetric. Measurement or estimation by weight. Halogen. Literally, salt producer; applied to the members of the chlorine group. hydrated. Containing water, hydroxyl. A term applied to the radical OH. hygroscopic. Applied to substances which readily absorb moisture from the air. Ignite. To set fire to. indicator. A substance used to show the completion of a chemical reaction. inflammable. Combustible, ion. An atom or group of atoms in a solution, which serves as a carrier of electricity. GLOSSAET 429 ionization. The separation of a substance into ions. isomeric. Applied to substances having the same percentage composi- tion, though differing in characteristics. isomorphous. Of the same crystalline form. Leach. To treat with water; to remove the soluble salts from a mixture of substances by means of water. liquation. The process of separating one metal from another by cautiously fusing, so that one will flow out before the melting point of the other is reached. lixiviate. Synonymous with leach. lute. To seal air-tight. Manipulation. Setting up or arranging apparatus for experiment. matte. A mixture of metallic sulphides obtained in the early stages of the reduction of copper ores, containing lead, silver, etc. meniscus. The upper curved surface of a liquid contained in a small tube. monad. An element the valence of which is one. mono-basic. A term applied to an acid having only one replaceable atom of hydrogen. mordant. A substance used to set the color in dyeing. mother liquor. The liquid remaining after the principal salt con- tained in solution has been removed by crystallization. Nascent. Applied to a gas when first liberated from its compound. It is believed to exist then in the atomic condition. native. Not in combination, free. neutral. Neither acid nor alkaline. neutralization. The combination of an acid with an alkali so as to destroy the properties of each, and produce a salt. nitrogenous. Containing nitrogen. Organic matter containing nitro- gen is thus characterized. Occlude. To condense upon the surface or within the pores. Especially seen in the action of platinum upon hydrogen. oxidation. The union of a substance with oxygen. oxidizing agent. A substance which readily gives up a portion of its oxygen to combine with some other substance. oxygenized. Containing considerable oxygen. Paste. A special variety of glass, used sometimes for making imita- tion diamonds. 430 MODERN CHEMISTRY pigs. The term applied to cast iron as molded when first drawn from the blast furnace. Applied also to the molds themselves. pipette. A small graduated glass tube used in measuring small quantities of a liquid. pneumatic. Pertaining to gases ; applied to the trough or pan used in collecting gases. polymerism. A term referring to the cases of compounds which have the same percentage composition, but different molecular weights. precipitate. A solid thrown down in a liquid by some reagent. Qualitative analysis. The determination of the kind of matter which enters into a substance. quantitative analysis. The determination of the amount of a sub- stance contained in a compound. Radical. A group of atoms which seems to act as a single atom, reaction. The action of two or more substances upon each other, reagent. A substance used to bring about some chemical change, reducing agent. A substance used to convert a compound from a higher to a lower order, as from an ic to an ous compound ; or, to remove the oxygen from an oxide, residual. That which remains, reverberatory. A variety of furnace, usually of low, arching ceiling. See Fig. 57 in text. roast. To heat strongly; to oxidize metallic ores, expelling the sulphur as SO 2 . Sand-bath. A small iron saucer containing sand, used the same as a wire screen in protecting glassware when being heated, saturated. Fully satisfied; containing all it can hold, scintillate. To burn with sparks. siliceous earth. Material consisting largely of silica, slag. The dark-colored glass formed in the reduction of metallic ores from the flux used and the gangue present. solvent. A liquid which dissolves some particular substance, spit. Silver on being melted absorbs considerable oxygen. Upon cooling it again expels this, sometimes with considerable energy, throwing out fine particles of the molten silver. This is termed spitting. stable. Not easily decomposed, sublimate. The substance obtained by sublimation. GLOSSARY 431 sublimation. The vaporizing of a solid and recondensing. The same in reference to solids that distillation is with liquids. supernatant. Said of a liquid overlying a precipitate after the latter has subsided. suspension. Said of a solid in the form of fine particles floating throughout the liquid. symbol. A letter or letters representing an atom of an element. Thio. From a Greek word, meaning sulphur. treat. To apply or add to. triad. An element having a valence of three. tubulated. Applied to a flask having a small tube-like opening in the side, fitted with a stop-cock. tubulure. A small, tube-like opening. tuydre. A blast or air pipe for conducting the strong currents of air into the blast furnace. Valence. The power which an atom or group of elements has of com- bining with some other element taken as a standard. volatile. Easy to vaporize. volatilize. To drive off in the form of vapor. volumetric. Estimation of the quantity of a substance ti& Jhfeasuring. INDEX Absolute thermometer, 95. Absolute zero, 95. Acetic acid, test for, 345. Acetylene, burners for, 150. characteristics of, 150. experiments with, 151, 152. generators, 149. preparation of, 148. Acids, 125. classes of, 343. composition of, 126. detection of, 343. nomenclature of, 128. preliminaries to testing, 345, 347. properties of, 125. Air, estimation of its constituents, 350. estimation of weight, 97. liquefaction of, 97. Air-slaked lime, 222. Alchemistic terms, 386. Alkali earths, 219. Alkalies, 125. Alkali metals, 207. Allotropism, 59. Aluminum, 263. bronze, 235. characteristics of, 264. hydroxide, 269. source of supply, 263. test for, 337. uses, 264. Alums, 266. kinds of, 267. preparation of, 266. uses of, 267. uses of, for clarifying water, 268. Amalgams, 258. methods of making, 258. Ammonia, 73. absorption of, by charcoal, 78. as a refrigerant, 78. characteristics of, 76. commercial supply, 74. decomposition of, by platinum, 78. estimation of the constituents, 351. fountain, 77. preparation for commerce, 74. test for, 342. uses, 78. Ammonium, 67. Anhydride, 83. Anions, 329. Antichlor, 112. Antimoniureted hydrogen, 292. Antimony, 290. amorphous, 292. black, 292. characteristics of, 291. chloride, 293. oxides, 293. oxy chloride, 293. reduction of ore, 290. sulphide, 294. test for, 334. uses, 292. Apparatus for pupils, 357. Aqua regia, 88. Argentite, 238. Argon, characteristics of, 90. discovery, 89. Arrangement of bottles, 356. Arsenic, characteristics of, 286. Marsh's test for, 287. 434 INDEX oxides, 288. reduction of ores, 285. source of supply, 285. sulphide, 290. uses of, 286. Arsenical pyrite, 285. Arsine, 287. Asbestos, 219. Aspirators, 370. Atmosphere, 91. Atom, definition of, 11. Atomic weights, 68. determination of, 198. Avogadro's Law, 190. application of, 198. proof of, 196. Azurite, 233. Banca tin, 270. Barium, 229. carbonate, 229. chloride, 229. hydroxide, 230. nitrate, 229. separation from calcium, 340. sulphate, 229. tests for, 340. Barometric reading, correction of, 374. Baryta, 229. Base, 124. Bauxite, 264. Bell metal, 235. Bessemer steel, 303. Binary compounds, 131. Biographical appendix, 382. Bismuth, 294. characteristics of, 294. compounds, classes of, 295. nitrate, 295. ocher, 296. oxychloride, 296. trichloride, 296. trioxide, 296. uses, 295. Bismuthite, 294. Bismuthyl compounds, 295. Black ash, 211. Black lead, 136. Blast furnace, 300. Bleaching powder, 227 Bloom, 302. Blowpipe work, 361. Blue prints, 245. Blue vitriol, 235. Bohemian glass, 188. Bordeaux mixture, 236, Bornite, 233. Bottles, opening of, 366. Boyle's Law, 93. Brass, 235. Bromides, test for, 344. Bromine, characteristics of, 117. commercial supply, 116. experiments with, 118. occurrence of, 116. preparation of, 117. test for, 117. uses, 118. Bronze, 235. Burnt alum, 267. Cadmium, 254. characteristics of, 255. nitrate, 256. reduction of, 255. sulphide, 256. test for, 334. Calchopyrite, 233. Calcite, 221. Calcium, 220. carbide, 148. carbonate, 223. characteristics of, 221. chloride, 224. history of, 221. hydroxide, 223. oxide, 221. sulphate, 224. Calomel, 260. INDEX 435 Carbon, abundance of, 135. as an absorbent, 139. as a reducing agent, 139. forms of, 135. uses of, 140. Carbon dioxide, 142. characteristics of, 144. estimation of, 348. liquid, 144. preparation of, 143. source of, 142. uses of, 144. Carbon monoxide, 141. Carre's ice machine, 79. Cassiterite, 270. Cast iron, 302. Castner's process for sodium, 208. Catalysis, 51. Cathions, 329. Caustic soda, 209. Cements, 225. composition of, 226. Chamber acid, 181. Charcoal, 137. Charles's Law, 94. problems with, 96. Chemical changes, 15. experiments to illustrate, 15, 1C, 17, 18. Chloric acid, 345. Chlorine, as a bleaching agent, 111. characteristics of, 109. chemistry of its preparation, 106. Deacon's process, 106. experiments with, 108. history of, 102. liquid, 110. occurrence, 103. preparation, 103. uses of, 111. water, 105. Weldon's process, 104. Choke damp, 144. Chromic acid, 321. Chromium, 317. compounds, 317. conversion of compounds, 319. hydroxide, 321. oxides, 320. test for, 337. uses of, 321. Chromite, 317. Cinnabar, 257, 260. Clay, 265. Coal, 137. Coal gas, 153. Cobalt, 311. compounds, 311. test for, 338. Coke, 138. Combination, laws of, 166. Combining weights, 164. Combustible substances, 57 C Combustion, 56. Compounds, 10. saturated, 24. Converter, 303. Copper, 232. alloys of, 235. blister, 233. characteristics of, 234. pyrite, 233. reduction of, 233. salts, 235. supply of, 232. tests for, 233. Copperas, 308. Corals, 221. Corrosive sublimate, 260. Corundum, 265. Crocosite, 317. Crown glass, 188. Cryolite, 264. Cupel, 239. Cupellation, 239. Cupola furnace, 304. Cupric acetylide, 236. chloride, 236. nitrate, 236. oxide, 237. 430 INDEX sulphate, 235. sulphide, 286. Cyanide process for gold, 247. Decanting, 364. Definite Proportions, Law of, 158. Deliquescent bodies, 31. Delivery tubes, preparation of, 358. Dewar bulbs, 97. Diamonds, 135. practical uses, 136. Diatomic molecules, 200. Diffusion of gases, 92. Dissociation, 329. Distillation, destructive, 137. fractional, 137. Dolomite, 219. Downward displacement, 362. Drying of tubes, 375. Dyads, 23. Dynamite, 89. Efflorescent bodies, 30. Electrolysis of water, 32. Electrolytic apparatus, 367. Elements, classes of, 204. definition of, 8. table of, 8, 204, 388. vacancies in table, 206. valence of, 8. Emerald, 265. Emery, 265. Epsom salts, 220. Equations, 27. exercise in, 28. writing, 69. value of, 69. Etching glass, 102. Ethylene, 147 Euchlorine, 105. Eudiometer, 33, 369. Experiments, recording, 376. Feldspar, 265. Ferric chloride, 308. oxide, 309. salts, how changed to ferrous, 307< salts, how distinguished, 306. sulphate, 308. sulphide, 308. Ferrous salts, how changed to ferric, 307. how tested, 306. Fertilizers, 194. Filter flask, 370. Filtering, 364. Fire damp, 146. Fixing bath, 244. Flame, 58. Flame tests tor barium, etc., 230. Flint glass, 188. Fluorine, 101. compounds of, 102. Fool's gold, 300. Formulae, determination of, 201. meaning of, 66. Franklinite, 250. Galena, 274. Ganister, 303. Gas carbon, 138. Gas generators, 373. Gases, collecting, 362. illuminating, 152. German silver, 253. Glacial phosphoric acid, 194. Glass, 187. annealing, 189. cutting, 358. etching, 102. manufacture of, 187. varieties of, 188. Glauber's salt, 212. Glossary of chemicals and min- erals, 391. Glossary of technical terms, 396. Gold, 246. characteristics of, 248. methods of mining, 246. Graphite, 136. WDEX 437 Greek fire, 175. Green fire, 229. Greenockite, 255. Green vitriol, 308. Guncotton, 89. Gunpowder, 175. separation of, 19. Gypsum, 221. Halogens, 101. comparison of, 122. Hard waters, 226. Harveyized steel, 310. Heavy spar, 229. Hematite, 299. Horn silver, 238. Hydraulic cement, 225. Hydraulic mining, 246. Hydriodic acid, test, 344. Hydrobromic acid, test, 344. Hydrocarbons, 146. Hydrochloric acid, characteristics of, 115. commercial supply, 113. composition of, proof, 354. composition of, estimation, 352. experiments with, 114. history of, 112. preparation of, 112. test for, 344. uses, 115. Hydrogen, 36. characteristics of, 42. experiments with, 42. liquid, 45. methods of preparing, 36. occlusion of, 44, 314. uses, 45. Hydrogen dioxide, 63. Hydrogen sulphide, 175. Hydroxides, 125. Hydroxyl, 66. Iceland spar, 221. Ice machine, 78, 80. Ice manufacture, 79. Illuminating gases, 152. composition of, 155. manufacture of, 153. Iodides, test for, 344. Iodine, 119. characteristics of, 121. experiments with, 121. preparation of, 119. solvents for, 122. uses, 122. Ionic theory, 328. lonization, 329. Ions, 329. Iridium, 314. Iron, cast, 302. compounds of, 305. distribution of, 299. forms of, 305. protoxide, 309. pyrite, 300. reduction of, 302. test for, 337. uses, 305. wrought, 302. Jets, preparation of, 359. Kaolin, 265. Kindling temperature, 57. Laboratory suggestions, 355. Lampblack, 139. Lead, 274. acetate, 279. carbonate, 280. characteristics of, 277. chloride, 279. chromate, 282. nitrate, 279. oxides, 280. reduction of ores, 275. sulphate, 279. sulphide, 282. tests for, 283, 331. uses, 277. 438 INDEX Leblanc's process for soda, 211. Lime, 221. properties of, 222. Limonite, 300. Linde's apparatus, 98. Liquid air, 100. apparatus for, 98. Liter, weight of, 201. Litharge, 280. Lithium, test for, 342. Lunar caustic, 243. Magnesia, 220. Magnesium, 219. compounds, 220. test for, 340. Magnetite, 299. Malachite, 233. Manganese, 323. compounds of, 323. dioxide, 324. dioxide, as a catalytic agent, 50. test for, 338. Manganic acid, 324. Marsh gas, 146. Marsh's test for arsenic, 287. Matte, 233. Matter, 11. theories of, 8. Measurements, 363. Meerschaum, 219. Mercuric chloride, 260. nitrate, 259. oxide, 260. salts, distinguished, 261. sulphide, 260. Mercurous chloride, 260. nitrate, 259. Mercury, 256. characteristics of, 257. reduction of, 257. solvents for, 259. tests for, 331, 333. uses, 259. 1 Metals, 203. displacing power of, 169. Metaphosphoric acid, 194. Meteorites, 299. Methane, 146. Micro-crith, 68. Minium, 280. Mixtures, 18. Molecular weight, 68. Molecules, 11. constitution of, 199. of compound bodies, 12. Monads, 23. Monatomic molecules, 200. Monobasic acids, 194. Mortar, 222. Multiple Proportions, Law 161. Natural gas, 155. Negatives, photographic, 244. Neutralization, 124. Nickel, 309. compounds of, 310. tests for, 311, 338. uses, 310. Nitric acid, characteristics of, 87. in the air, 86. preparation of, 86. test for, 85. uses, 88. Nitric oxide, 82. characteristics of, 83. Nitrogen, 71. characteristics of, 73. oxides of, 81. pentoxide, 86. tetroxide, 85. uses of, 73. Nitrogen group, 297. Nitroglycerine, 89. Nitrous acid, preparation, 84. test for, 84. Nitrous anhydride, 83. oxide, 81, INDEX 439 Occlusion, 44, 314. Oil of vitriol, 179. Olefiant gas, 147. Orpiment, 285. Osmium, 314. Oxidation, 56. Oxides, 132. Oxidizing flame, 361. Oxygen, 47. characteristics of, 54. determination of weight, 55. experiments with, 49. liquid, 54. Motay's method, 52. preparation, 48. preparation from potassium per- manganate, 53. uses, 55. Oxy-hydrogen blowpipe, 58. Ozone, 59. characteristics of, 61. comparison with oxygen, 60. liquid, 61. Palladium, 314. Panning gold, 246. Paris green, 289. Parke's method, 239. Paste, 188. Pattison's method, 239. Pearl ash, 216. Pearl white, 297. Pentads, 23. Periodic Law, 204. Phosphates, 194. Phosphine, 192. Phosphoric acid, test for, 344. Phosphorus, 190. acids of, 194. characteristics of, 191. manufacture of, 190. oxides of, 193. uses of, 192. Photographic papers, 244. plates, 243. Photography, 243. Physical changes, 12. illustration of, 13. experiments, 12, 14. Pig iron, 302. Pintsch gas, 154. Placer mining, 246. Plaster of Paris, 224. Platinum, 314. alloys of, 315. spongy, 314. uses, 316. wires, 366. Polymerism, 62. ' Porcelain, 265. Potassium, 214. bromide, 217. carbonate, 216. chlorate, 216. chromate, 318. dichromate, 318. hydroxide, 215. iodide, 217. nitrate, 217. permanganate, 324, 325. ' tests for, 217, 342. Precipitates, 364. Prefixes, per, pro, etc., 133. Pressure, standard, 93. Puddling furnace, 302. Pyrite, 300. Pyrolusite, 323. Qualitative analysis, 327. Quantitative experiments, carbon dioxide, estimation of, 348. combining weight of copper, 164. combining weight of tin, 165. composition of air, 350. composition of ammonia, 351. composition of hydrochloric acid, 352. definite proportions, proof of law, 158, 159, 160, 161. 440 INDEX displacing power, aluminum, 170. magnesium, 171. zinc, 170. electrolysis of water, 32. manganese dioxide as a catalytic agent, proof of, 51. multiple proportions, proof of law, 162. strength of acid, determination of, 167. strength of alkali, determination of, 167. strength of salt solution, determi- nation of, 169. synthesis of water, 33, 34. water of crystallization, determi- nation of, 349. weight of 1 liter of air, 97. weight of 1 liter of oxygen, 55. Quicksilver, 256. Radicals, 66. Reactions, 67. Realgar, 285. Red fire, 228. Red precipitate, 260. Reducing flame, 361. Reference library, 381. Rose quartz, 186. Ruby, 265. Safety lamp, 147. Saltpeter, 217. Chile, 212. Salts, acid, 127. definition of, 127. formulae of, 128. nomenclature of, 130. normal, 127. Sapphire, 265. Scheele's green, 289. Scheele's test for arsenic, 290. Separation of metals, arsenic, an- timony, tin, 334. barium, strontium, calcium, mag- nesium, 340. bismuth, copper, cadmium, 333. iron, chromium, aluminum, 337. lead, silver, mercury, 330. nickel, cobalt, manganese, zinc, 338 sodium, potassium, lithium, 342. Shot, 278. Siderite, 300. Silica, 186. Silicates, 186. Silicic acid, 187. Silicon, 184. Silver, 238. characteristics of, 241. chloride, 242. chromate, 242. experiments with, 240. nitrate, 241. Parke's process, 239. Pattison's process, 239. reduction of ores, 238. test for, 332. uses, 241. Smalt, 312. Smithsonite, 250. Smoky quartz, 186. Soap, 212. hard and soft, 213. Sodium, 207. carbonate, 210. characteristics of, 208. chloride, 209. effects upon water, 38. experiments with, 38. hydroxide, 209. nitrate, 212. preparation of, 210. sulphate, 212. tests for, 214, 342. Sodors, 145. Solder, 278. Solubility of salts, 346. Solutions, preparation of, 376. Solvay process for soda, 210. INDEX 441 Sparklets, 145. Spathic iron, 300. Spelter, 253. Spiegeleisen, 304. Stalactite, 231. Stannic chloride, 272. oxide, 274. sulphide, 273. Stannous chloride, 272. sulphide, 273. Steel, 303. basic lining process, 304. comparison with cast iron, 305. tempering, 305. Stibine, 292. Stibnite, 290. Strass, 188. Strontianite, 228. Strontium, 228. carbonate, 228. hydroxide, 229. nitrate, 228. separation of, 340. tests for, 340. Sugar of lead, 279. Suint, 214. Sulphides, test for, 304. Sulphur, 171. acids of, 179. allotropic, 174. characteristics of, 173. forms of, 174. oxides of, 177. source of supply, 172. uses, 175. Sulphur dioxide, 177. characteristics of, 178. uses, 179. Sulphuric acid, 179. characteristics of, 182. manufacture of, 181. test for, 181, 343. uses, 183. Sulphurous acid, 183. test for, 344. Supplies needed, 378. Supporters of combustion, 57. Symbols, 65. Sympathetic inks, 312. Synthesis of water, 33. Tables : comparison of oxygen and ozone, 62. comparison of metals and non- metals, 203. composition of cements, 226. compounds of chromium, 318. compounds of manganese, 326. iron salts, distinctions between,306. names of elements, 9, 388. nitrogen group, 297. periodic law, 204. salts of mercury, 261. separation of metals, group I, 331. group II, 336. group III, 339. group IV, 341. group V, 343. tension of aqueous vapor, 374. three forms of iron, 305. valence, 27. Ternary compounds, 131. Test-tube repairing, 360. Tetrads, 23. Thiosulphuric acid, 184. test for, 344. Tin, 270. alloys of, 272. characteristics of, 270. foil, 272. plate, 272. salts, 272. stone, 270. test for, 335. uses, 272. Triads, 23. Tribasic, 194. Type-metal, 278. 442 INDEX Univalent atoms, 23. Upward displacement, 362. Valence, definition of, 21. double, 25. exercise in, 26. illustration of, 22. of radicals, 25. variation of, 23. Vapor density, determination of, 200. Vein mining for gold, 246. Vermilion, 260. Vitriol, oil of, 179. blue, 335. green, 308. white, 253. Wash bottle, preparation of, 359. Water, abundance of, 29. analysis of, 32, 34. characteristics of, 31. composition of, 32. decomposition by sodium, 37. forms of, 29. solvent powers of, 32. synthesis of, 33. Water gas, 154. Water glass, 187. Water of crystallization, 30. estimation of, 349. proof of, 30. Weldon's mud, 105. White arsenic, 288. White lead, 280. Dutch method of preparation, 280 electrolytic method, 281. Milner's method, 281. White vitriol, 253. Witherite, 229. Wrought iron, 302. Zinc, 250. alloys of, 253. blende, 250. characteristics of, 252. chloride, 253. hydroxide, 254. ores, 250. oxide, 254. reaction with acids, 40. reduction of, 250. sulphate, 253. sulphide, 254. test for, 338. uses, 253. white, 254. If I 1 Millimetres X .03937 = inches. Millimetres -*- 25.4 = inches. Centimetres X .3937 = inches. Centimetres -4- 2.54 = inches. Metres X 39.37 = inches. (Act Congress.) 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