MFN CHEMISTRY Edmund O'Neill ELEMENTARY CHEMISTRY BY THE SAME AUTHOR. CHEMICAL LECTURE EXPERIMENTS. With 224 Diagrams. Crown 8vo, $3 oo. A TEXT-BOOK OF INORGANIC CHEMISTRY. With 146 Illustrations. Crown 8vo, $1.75. LONGMANS, GREEN, & CO. NEW YORK, LONDON, AND BOMBAY. ELEMENTARY INORGANIC CHEMISTRY BY G. S. NEWTH, F.I.C., F.C.S. DEMONSTRATOR IN THE ROYAL COLLEGE OF SCIENCE, LONDON ASSISTANT EXAMINER IN CHEMISTRY, SCIENCE AND ART DEPARTMENT AUTHOR OF "A TEXT-BOOK OF INORGANIC CHEMISTRY," "CHEMICAL LECTURE EXPERIMENTS," ETC. LONGMANS, GREEN, AND CO LONDON AND BOMBAY 1899 All rights reserved UNIFORM WITH THIS VOLUME. Crown 8vo. ELEMENTARY PHYSICS. By W. WATSON, B.Sc., Demonstrator in Physics in the Royal College of Science, London ; Assistant Examiner in Physics, Science and Art Department. LONGMANS, GREEN, AND CO. NEW YORK, LONDON, AND BOMBAY. IN MEMOR1AM PREFACE THIS book has been written to meet the modern and practical methods of science teaching which are now being universally recognized and adopted in schools and colleges. Formerly students were taught chemistry in the lecture- room, the knowledge so gained being supplemented by a minimum amount of practical work, and that almost ex- clusively analytical. The tendency of the present day is to make the student, from the very beginning, an investigator ; to train and develop his faculties for observation; to make him find out facts and discover truths for himself; in other words, to make him think instead of merely committing to memory what others have thought. I have therefore en- deavoured, as far as it is possible to do so in a text-book, to fall into line with these views. In actual practice the purely inductive method of instruction breaks down. There is so much that the student is required to learn, that life itself is not long enough, and certainly the limited time at the dis- posal of the student is all too short, to admit of his going through the necessarily slow process of gaining this knowledge by his own investigations. Some facts he must take on trust, and the question therefore resolves itself into the judicious 889770 vi Preface. selection on the part of the teacher of the facts he will endeavour to let his students find out for themselves, and those he will teach them, and expect them to commit to memory. In a text-book it is almost inevitable that, in giving such directions as will lead a student on to the discovery of a fact, the fact itself shall be stated* Before introducing the student to the study of any of the elements, I have sought to familiarize him with a number of important common laboratory processes, in a chapter on " Simple Manipulations ; " and this is followed by short chapters on the " Fitting up of Apparatus," and " Simple Glass-blowing Operations." After hydrogen, oxygen, and water have been studied, I have introduced, under the head of "Simple Quantitative Manipulations," a number of experiments or exercises in- volving the operations of weighing and measuring. These experiments have been selected with the object of leading the student on to the discovery of some of the fundamental laws of chemistry, making use of such knowledge of chemical facts as he has already gained. In order that he may do these with an entirely unprejudiced or unbiassed mind, they have intentionally been introduced before he has learnt the use of symbols and formulae, or how to calculate what results he ought to get. For some of these experiments I am indebted to the suggestions of Dr. Tilden, made during the course of a short series of lectures to science teachers, at the Royal College of Science, London, in July, 1895. l 1 Since embodied in a little publication, ' ' Hints on the Teaching of Elementary Chemistry " (Longmans, Green, and Co.). Preface. vii In an elementary practical text-book it would obviously be unwise to take up the study of more than a very limited number of the elements and their compounds. Exactly which elements and which of their compounds are the most suitable for the purpose is probably a point on which teachers will hold different opinions. I have selected those which, in my judgment, are the best adapted for an elementary course, and which I consider are well calculated to give the student a broad, general foundation upon which he can afterwards build. G. S. N. TABLE OF CONTENTS I. STATES OF MATTER SOLID, LIQUID, GAS PHYSICAL AND CHEMICAL CHANGES l II. COMPOUNDS ELEMENTS METALS NON-METALS- MIXTURES CHEMICAL AFFINITY CHEMICAL AC- TION 8 III. SIMPLE MANIPULATIONS SOLUTION DECANTATION FILTRATION PRECIPITATION CRYSTALLIZATION- FUSION DISTILLATION COLLECTING GASES . . 15 IV. FITTING UP APPARATUS 28 V. SIMPLE GLASS-BLOWING 35 VI. HYDROGEN 4 VII. OXYGEN ACIDS ALKALIES BASES SALTS ... 54 VIII. WATER (PEROXIDE OF HYDROGEN) 7 2 IX. MEASURING WEIGHING METRIC SYSTEM THERMO- METERS BAROMETER BALANCE 85 X. RELATION OF VOLUME OF GASES TO HEAT AND TO PRESSURE THE CRITH 97 XI. SIMPLE QUANTITATIVE MANIPULATIONS FIRST AND SECOND LAWS OF CHEMICAL COMBINATIONS . . 104 XII. HYDROCHLORIC ACID CHLORIDES TESTS FOR . . 120 XIII. CHLORINE (THE HALOGENS) 128 XIV. FURTHER QUANTITATIVE MATTERS EQUIVALENCE . 136 XV. ATOMIC THEORY ATOMS MOLECULES 146 XVI. CHEMICAL NOTATION SYMBOLS FORMULA EQUA- TIONS . . 152 xii Table of Contents. CHAPTER PAGE XVII. QUANTITATIVE SIGNIFICANCE OF EQUATIONS COM- BINATIONS BY VOLUME GAY-LUSSAC'S LAW Avo- GADRO'S HYPOTHESIS DENSITIES MOLECULAR WEIGHTS UNIT VOLUMES 158 XVIII. AIR DIFFUSION COMBUSTION 166 XIX. NITROGEN AND ITS COMPOUNDS. NITROGEN, AM- MONIA 176 XX. NITROGEN AND ITS COMPOUNDS (continued}. NITRIC ACID, NITRATES 188 XXI. NITROGEN AND ITS COMPOUNDS (continued). OXIDES OF NITROGEN 197 XXII. OZONE 206 XXIII. CARBON 210 XXIV. CARBON DIOXIDE CARBONATES QUANTITATIVE DE- TERMINATION OF CARBON DIOXIDE 220 XXV. CARBON MONOXIDE 231 XXVI. SULPHUR 236 XXVII. SULPHURETTED HYDROGEN 243 XXVIII. SULPHUR DIOXIDE SULPHUROUS ACID SULPHITES . 250 XXIX. SULPHUR TRIOXIDE SULPHURIC ACID SULPHATES . 255 XXX. SOME COMMON CARBON COMPOUNDS 264 XXXI. SIMPLE QUALITATIVE ANALYSIS 272 LIST OF ELEMENTS Names. Atomic Symbols Atomic Weights Names. Atomic Symbols Atomic Weights. Aluminium . Al 27 Molybdenum Mo 9 6 Antimony (Stibium] Sb 120 Nickel . Ni 59 Argon A (?) Niobrum . Nb 937 Arsenic As 75 Nitrogen . N H Barium Ba 137 Osmium . Os 191 Beryllium * . Be 9 Oxygen . O 16 Bismuth Bi 207-5 Palladium Pd 106 Boron B ii Phosphorus P 31 Bromine Br 80 Platinum . Pt *95 Cadmium . Cd 112 Potassium (Ka- Ccesium Cs 133 lium) . K 39 Calcium Ca 40 Rhodium . Rh 104 Carbon C 12 Rubidium Rb 85 Cerium Ce 141 Rutheni2tm Ru !3'5 Chlorine Cl 35 '5 Samarium Sm 150 Chromium . Cr 5 2 Scandium Sc 44 Cobalt Co 59 Selenium . Se 79 Copper (Cuprum} Cu 63 Silicon Si 28 Didymium . Di H5 Silver (Argentum) Ag 1 08 Erbium Er 1 66 Sodium(lVatrtum) Na 2 3 Fluorine F 19 Strontium Sr 87-3 Gallium Ga 70 Sulphur . S 32 Germanium . . Ge 72 Tantalum Ta 182 Gold (Auruni) Au 197 Tellurium Te I2 5 Hydrogen . H i Thallium . Tl 2037 Indium In "3 Thorium . Th 232 Iodine I 127 Tin (St annum) . Sn 118 Iridium Ir 192-5 Titanium- Ti 48 Iron (Ferrttm) Fe 56 Tungsten . W 184 Lanthanum . La 138-5 Uranium U 239-8 Lead (Plumbum) . Pb 207 Vanadium V 5 1' 1 Lithium Li 7 Ytterbium Yb 173 Magnesium . Mg 24 Yttrium . Y 89-6 Manganese . Mn 55 Zinc Zn 65 Mercury (Hydrar- Zirconium Zr 90-4 gyrum) . TT 200 \ Those printed in italics may be regarded as rare substances. ELEMENTARY PRACTICAL CHEMISTRY. CHAPTER I. STATES OF MATTER PHYSICAL AND CHEMICAL CHANGE. ALL material things with which we 41 e acquainted* exist under ordinary circumstances in one of three states^or conditions either they are solids, like chalk,' iron,: salpWufr^ ic^ ;or liquids, as water, alcohol, mercury ; or they are gases, lik'e air, oxygen, and steam. We know many substances, however, which can easily be made to pass from one of these states to the other ; thus it is familiar to all that when solid water (that is, ice) is gently warmed it changes into liquid wafer, and that when this liquid is more strongly heated it passes off into steam, which is water in the gaseous or vaporous state. We also know that when steam is cooled it changes back again to liquid water, and that when this is further cooled it is turned into ice, or solid water. Sulphur is an example of another common substance that can readily be caused to change from one state to another. Experiment I. Take a piece of sulphur (brimstone) and chip off a fragment about the size of a pea. (Note that sulphur is a pale yellow solid, easily broken, being very brittle.) Place the small piece in a clean dry test-tube, and gently heat it over a small Bunsen flame, in the manner shown in the figure. Notice that the B States of Matter. solid quickly melts, and, if carefully heated, is changed into a pale yellow liquid. Now heat more strongly, and observe that the colour rapidly darkens, and the liquid presently boils. It is now being changed from the liquid to the gaseous state. As the gaseous sulphur reaches the part of the tube a little removed from the flame, it soon becomes cooled again, and consequently changes back again to the liquid state. Notice liquid sulphur collecting on the upper part of the tube in the form of small yellow drops. Continue heating until the whole of the original frag- ment of sulphur has disappeared from the bottom of the test-tube. Allow the tube to cool, and observe that the yellow liquid which had F G. i. , condensed upon the upper part * t /gradually changes back to the solid state. In this experiment, therefore, solid sulphur has been changed , first- ;to the liquid &nrl then to the gaseous state; and gaseous sulphur has been allowed to pass back again, first to the liquid and then to the solid condition. Experiment 2. Gently heat a small quantity of alcohol or methylated spirit in a test-tube fitted with a cork and bent glass tube (as shown in Fig. 2), which is joined to a second test-tube provided with a cork with two bent tubes. This second tube is placed in a glass containing cold water, to keep it cool. The spirit soon begins to boil, and to pass from the liquid to the gaseous, or vaporous state. As this gaseous alcohol passes into the cooled tube, however, it again returns to the liquid condition, and will be seen collecting at the bottom of the tube. We see, from these examples, that with some things it is simply a question of whether they are heated or whether they are cooled that decides the particular state they shall be in. This is also true of a number of other substances ; for instance, the metal mercury (quicksilver) is familiar to us as a liquid, but if we happened to be living in the Arctic regions, we should know it as a hard solid, resembling lead. Again, spectrum analysis has taught us that iron, lead, copper, tin, and many . Physical and Chemical Change. 3 other metals which are familiar to us as solids, are present in the sun, and that, owing to the intense heat, they exist there in the gaseous state, forming a part of the sun's atmosphere. It is not necessary to take these metals to the sun, however, to make them change their state from the solid to the gaseous. When we apply heat to them, some of them, such as lead and tin, melt readily enough ; others, like copper and iron, require a much higher temperature to make them pass into the liquid state ; while all of them can be boiled and made to pass into the gaseous condition by means of the electric furnace. All solids and liquids are visible and tangible ; gases, on FIG. 2. the other hand, cannot be felt, and in most cases they are invisible. It is quite impossible, by merely looking at it, to tell whether a glass bottle is filled with air, or hydrogen, or oxygen, or whether it is entirely empty that is, vacuous ; hence, when a liquid substance passes into the gaseous state, it usually disappears altogether from our sight. For example, if a small quantity cf water be left exposed in a shallow dish or saucer, we know that it gradually diminishes in quantity, and finally disappears entirely. In common language, we say that it has dried up ; in more scientific phraseology, we speak of the process as evaporation that is to say, the water has passed 4 States of Matter. from the liquid to the vaporous or gaseous state, in which condition it is invisible, and mingles with the other equally invisible gases of the air. The visible cloud which appears when steam is allowed to escape from a boiler or locomotive engine, and which is popularly called "steam," consists in reality of minute drops of liquid water, and is not water in the vaporous or gaseous state. The true steam is invisible, but on suddenly coming into contact with cool air the gaseous water changes to liquid water and becomes visible. That gaseous water is invisible may be proved by the following experiment : Experiment 3. Place a small quantity of water in a large glass flask, which is provided with a cork carrying a short glass tube bent at right angles. Boil the water rapidly, until a jet of steam escapes from the tube and produces the familiar cloud of "steam." It is evident that the flask must now be full of gaseous water, but it is perfectly clear and invisible. Now hold immediately beneath the jet of steam another Bunsen flame (as at 2, Fig. 3), and observe that the cloud instantly disappears. Although the same amount of steam is escaping from the tube, it is now invisible. This is because the flame warms the air in the immediate neighbourhood of the jet, so that the gaseous water does not become suddenly cooled on issuing into the air, and therefore does not condense to the liquid state. Some solid forms of matter, when heated, change into the gaseous state without first becoming liquid ; that is to say, they pass at once into gases without melting. For example Experiment 4. Heat in a dry test-tube a fragment of ammo- nium chloride (sal ammoniac} about the size of a grain of wheat. FIG. 3. Physical and Chemical Change. 5 Notice that the white solid does not melt, but at once passes into vapour. The vapour, on reaching the cooler parts of the test-tube, changes back again to the solid state, and collects as a white coating upon the glass. This process is termed sublimation; the ammonium chloride is said to sublime when heated. Any impurities present in the original solid, which do not change into vapour at the temperature employed, will obviously be left behind, hence this process may be used to purify substances like ammonium chloride. Many kinds of matter, when experimented upon in the same way as the sulphur (Exp. i) and the ammonium chloride (Exp. 4), undergo a different kind of change, a change which is more subtle and less simple. This will be seen by the following examples : Experiment 5. Heat in a dry test-tube a small quantity of potassium chlorate, and carefully notice what takes place. The solid quickly melts and changes into the liquid state, and presently appears to be boiling. So far it seems to behave like the sulphur (Exp. i). But is the liquid potassium chlorate being changed into the gaseous state ? In the first place, it will be noticed that practically nothing condenses upon the upper and cooler part of the tube ; this seems to imply that the substance is not passing into the gaseous condition. Apply a lighted taper to the mouth of the tube ; no gas is escaping which will take fire, but notice that the flame of the taper becomes brighter. Dip into the test-tube a match or splinter of wood which has been lighted and blown out, and has still a glowing spark upon it ; the wood will be rekindled and burst into flame. This proves that something gaseous is coming from the boiling liquid. Now fit a cork and bent tube into the test-tube, and connect it to a second tube arranged as in Fig. 2. It will be found that nothing visible, either liquid or solid, collects in the cold tube. As the heating of the potassium chlorate is continued, it will be seen that the liquid becomes less fluid, and finally goes solid again. By this experiment, therefore, we find that when solid potassium chlorate is heated it first becomes liquid, and then is changed into two different things, namely, a colourless gas which does not change either to a liquid or a solid when cooled by cold water, and a solid which is evidently different from the original one, because it is very much more difficult to melt. Experiment 6. Place a few grains of mercuric oxide in a test- tube, and apply heat. In this case the red solid becomes darker in colour, but does not melt. Gradually, however, there will collect 6 States of Matter. upon the cooler part of the tube a sublimate which has the appear- ance of a bright white silvery metal, and upon dipping into the test-tube a glowing splint of wood, we shall obtain the same result as in Exp. 5. By continuing the heat, the whole of the original substance will disappear. We learn from this experiment, there- fore, that when mercuric oxide is heated, it also changes into two different things ; one of them being a colourless gas which passes away (presumably the same gas as was given out by the potassium chlorate in Exp. 5), while the other is a white metal. Examine this metallic sublimate carefully, and notice that it consists of minute globules of liquid metal. This must be mercury, because no other metal is liquid at the ordinary temperature. Certain differences between the kind of change undergone by the mercuric oxide and the potassium chlorate, and that experienced by sulphur and by ice, when these things are heated, will be evident, (i) Neither the sulphur nor the ice change into two different kinds of matter at once. (2) When sulphur or ice are changed by heat into the liquid or gaseous state, the change is only a temporary one ; when cooled again, they change back to their original condition. On the other hand, the mercuric oxide and the potassium chlorate each change into two different forms of matter, and the change in each case is permanent. Changes like those experienced by mercuric oxide and potassium chlorate are called chemical changes, whilst those which the ice and the sulphur under- went are distinguished as physical changes. Many chemical changes are constantly going on around us in the familiar processes of everyday life : thus, when a candle burns, the solid wax is transformed into certain invisible gases which mix with the air, and never return again to the original condition of the wax; the candle undergoes a chemical change. In our ordinary fires, coal is converted into smoke and certain invisible gases, which escape into and pollute the atmosphere, and a small residue of a greyish ash is left ; the change is permanent, and is a chemical change. When an egg is cooked, the clear and almost colourless liquid albumen (white of egg) is converted into a white solid, which does not again return to the liquid state on cooling; the albumen undergoes a chemical change. Physical and Chemical Change. j Human beings eat bread, meat, vegetables, etc. ; these foods undergo chemical changes, whereby they are converted into flesh and bones, into invisible gases which leave the body in the breath, and into waste products which leave the body in the perspiration and evacuations. Chemical and physical changes, however, are very closely associated, and although we may have a physical change without any chemical change, all chemical changes are ac- companied by a physical change ; and in many cases the accompanying physical change is the only outward indication which we have that a chemical change has taken place at all. For example Experiment 7. Take a glass tube about two feet long and half an inch wide, and close one end like the bottom of a test-tube. Fit a cork into the other end, and slip an indiarubber ring upon the tube a short distance from the corked end. About half fill the tube with strong sulphuric acid, and then gently fill the tube up to the ring with cold water. The water, being much lighter than the acid, will float upon it without mixing, if poured in gently. Now tightly cork the tube, and tip it up and down two or three times in order to mix the contents together. A chemical change takes place, but nothing is to be seen : it will soon be found, how- ever, that the tube is getting so hot that it can scarcely be held in the hand. Now hold the tube in an upright position, and notice that the liquid has shrunk, for it no longer reaches to the mark upon the tube. Here we see the chemical change is accompanied by a change of temperature, and a change of volume. CHAPTER II. COMPOUNDS, ELEMEN L'S METALS, METALLOIDS MIXTURES CHEMICAL AFFINITY, CHEMICAL ACTION. Elements and Compounds. When different kinds of matter are experimented upon, it is found that from some of them it is possible to obtain two or more substances totally unlike the original matter, while from others it is impossible to obtain anything essentially different by any process at present known. For example, in Exp. 5 we found that from potassium chlorate we were able to obtain a gas which caused a glowing splint of wood to re-light (a gas called oxygen), and also a white solid residue which required a much stronger heat to melt it than the original substance did. Again, in Exp. 6, from the red mercuric oxide we obtained the same gas oxygen, and a quantity of metallic mercury, two things entirely different from the original. Substances like potassium chlorate and mercuric oxide are called compounds. Substances from which we are unable to obtain anything essentially different, are distinguished as elements. There are only about 70 substances which are believed to be elementary bodies; and nearly the half of these may be considered as rareties. The following list of thirty includes all the most important of the elements (for the complete list see inside the cover) : Aluminium, Al. Calcium, Ca. Gold, Au. Antimony, Sb. Carbon, C. Hydrogen, H. Arsenic, As. Chlorine, Cl. Iodine, I. Bismuth, Bi Copper, Cu. Iron, Fe. Bromine, Br. Fluorine, F. Lead, Pb. Mechanical Mixtures. Magnesium, Mg. Manganese, Mn. Mercury, Hg. Nickel, Ni. Nitrogen, N. Oxygen, O. Phosphorus, P. Platinum, Pt. Potassium, K. Silicon, Si. Silver, Ag. Sodium, Na. Sulphur, S. Tin, Sn. Zinc, Zn. Of these elements two are liquid at the ordinary tempera- ture, namely bromine and mercury ; five are gases, chlorine, fluorine, hydrogen, nitrogen, and oxygen, while the rest are solid. Most of the solid elements are metals, and the above list will be seen to contain the names of such familiar metals as copper, gold, iron, lead, etc. Those of the solid elements in this list that are not metals are arsenic, carbon, iodine, phos- phorus, silicon, and sulphur. Some of these possess, in a greater or less degree, some of those properties which are usually associated with the metals. Thus arsenic and silicon are opaque substances having the power of reflecting light from their surfaces, a property usually known as metallic lustre. Carbon, again, has the power of conducting both heat and electricity, properties which are almost exclusively charac- teristic of metals. The name metalloids is sometimes given to those elements which, while not being true metals, yet closely resemble them in some of their properties. Mechanical Mixtures. When two elements are brought together, either nothing happens, or else a chemical change takes place. In the former case the result is a simple mixture of the two elements ; in the latter it is the formation of a compound. Similarly, when two compounds are brought together, if a chemical change takes place it results in the formation of new compounds, whereas if no chemical change follows we only obtain a mechanical mixture of the two compounds. Experiment 8. Grind to powder in a mortar a small quantity of potassium chlorate, and mix it with about the same quantity of powdered white sugar. No chemical change takes place, thereibre the result is a simple mixture. Experiment 9. Place in a dry test-tube a fragment of phos- phorus about the size of a pea, and drop upon it a few grains of powdered iodine. A chemical change at once takes place, great heat is evolved, and a compound of phosphorus with iodine results. io Mechanical Mixtures. [Phosphorus is a substance which must be handled with great care, as it very easily takes fire. It is always preserved beneath water, and when a small fragment is required for experiment, the larger piece should be taken out of the bottle by means of tongs, placed on a plate with some water, and there cut with a penknife. The piece to be used must be quickly wiped with blotting-paper, and put into the test-tube with the tongs. It should never be taken up by the fingers, as the warmth of the hand might cause it to take fire.] Experiment io. Place in a dry test-tube a few drops of mercury, together with a few particles of powdered iodine, and gently heat the tube. A chemical change follows, and a siiblimate is obtained partly red and partly yellow. This sublimate consists of a compound of mercury and iodine. Each of the ingredients in a simple mixture retains its own individual and characteristic properties. Thus, if the mixture of potassium chlorate and sugar (Exp. 8) be tasted, both the sweet taste of the sugar and the peculiar saline taste of the potassium chlorate will easily be perceived. Owing to the fact that their properties are retained, the ingredients of a mechanical mixture can be separated from each other by processes which are purely mechanical or physical, as distinguished from chemical. Experiment n. Powder some potassium nitrate (nitre), and mix into it a little powdered sulphur (flowers of sulphur) and a small quantity of fine iron filings. The result is a mechanical mixture of these things. If this mixture be now examined through a pocket lens, the separate particles of white nitre, yellow sulphur, and grey iron will be distinctly visible, lying side by side unchanged. Place the mixture upon a sheet of white paper and bring a magnet near to it ; the iron in the mixture will be attracted by the magnet, and may in this way be drawn away and entirely removed. The residue, containing now only the nitre and sulphur, should be removed to a test-tube, a little water added, and the contents of the tube warmed. The nitre being easily dissolved by water is thus separated from the sulphur, which does not dissolve. If the mixture be now poured upon a blotting-paper filter (see p. 17) the watery solution containing the nitre passes through, while the sulphur remains on the paper. If the liquid be then boiled in a little dish, the water will evaporate and the nitre will be left as a Chemical Affinity. 1 1 white solid residue in the dish. In this way the three ingredients have been separated by physical processes. Chemical Affinity. In a compound the various con- stituents stand in a totally different relation to each other than is the case with mixtures ; a relation which is much closer, and more difficult to understand. The elements present in a compound are said to be chemically combined with each other : their union in all cases is controlled by the operation of a particular force, which is spoken of as chemical affinity. Consider some of the cases of chemical combination already referred to. In Exp. 10 the two elements, mercury and iodine, enter into chemical union, and the compound formed is called mercuric iodide. If this scarlet powder be examined by the most powerful microscope it is quite impossible to see either the mercury or the iodine it contains. All the properties belonging to mercury, as well as those belonging to iodine, are gone, and the compound is endowed with an entirely new set of properties which are peculiar to itself. Again, in Exp. 6 we learnt that mercuric oxide was a compound of mercury and oxygen. These two elements, the liquid metal mercury and the colourless gas oxygen, when united together by the influence of chemical affinity, entirely lose their own individuality, and take on altogether new habits the compound is a brick-red crystalline solid, it possesses none of the properties of either mercury or oxygen, and these constituents cannot be separated again by any purely mechanical operations. The familiar substance water, is a compound obtained by the chemical union of two colourless invisible gases ; one of them (hydrogen) a gas which easily burns, and the other (oxygen) a gas which causes ordinary burning things to burn more quickly and brilliantly. The compound, therefore, which these gases give rise to when they chemically combine has none of the properties of the ingredients, its properties are diametrically opposed to theirs it is liquid, it does not burn, 12 Chemical Action. and it is the great antidote for fire. By no mechanical method can we disengage or separate these constituents. Why it is that the compound produced, when two such substances as hydrogen and oxygen enter into chemical union, should have the particular properties with which water is endowed, no one knows. Or why the compound of mercury and iodine should happen to be red, and not blue or green, we cannot tell. And, in the same way, it is quite impossible, from a knowledge of the properties of the in- gredients, to foretell what will be the nature of the compound they will give rise to. For instance, given a certain gas (chlorine), with a greenish-yellow colour, a powerful irritating smell, and extremely poisonous; also a soft white metal (sodium), which if placed upon the tongue would take fire ; now what sort of properties are likely to be possessed by a compound formed by the chemical union of these substances ? No one would predict that the product would be the innocent and necessary article of food, common salt ; but such is actually the case. Chemical Action is the term applied to the actual pro- cesses which take place by the operation of the force called chemical affinity. Thus, when phosphorus unites with iodine (Exp. 9), or when iodine combines with mercury (Exp. 10), the process of combination is termed chemical action, we say that a chemical action takes place between the phosphorus, or the mercury, and the iodine. Chemical action does not take place promiscuously between all the elements. Some are made to combine only with great difficulty; others will not combine together at all As a general rule, those elements which are most unlike each other combine together most readily, they are said to have the greatest affinity for each other. Neither does chemical action take place in all cases under the same conditions ; thus some- times it takes place at once on simply bringing the substances together. Experiment 12. Place a small quantity of sodium peroxide in a dry test-tube, and pour upon the powder a little cold water. Chemical Action. 13 Chemical action at once takes place, the mixture effervesces briskly. Such effervescence signifies that a gas is being set free, as one of the products of the chemical action. If a glowing match or a splint of wood be dipped into the test-tube, the same result will follow as in Exps. 5 and 6, showing that the gas given off is oxygen. See also Exp. 9. In other cases it is necessary to employ some external energy in order to induce chemical action to begin. In a large number of instances the application of heat will bring about chemical action. Thus, in Exp. 10, chemical action between the mercury and iodine was induced by heating the mixture. This may also be exemplified by the following experiments. Experiment 13. Make a small heap of the mechanical mixture of potassium chlorate and sugar (Exp. 8), and apply a lighted match to it. Chemical action is at once set up, and rapidly propagated throughout the heap. Experiment 14. Place a few fragments of copper foil or wire in a test-tube, and pour upon them a small quantity of strong sulphuric acid. While cold no action takes place, but if the mixture be warmed it will first become muddy, and presently give off a gas (sulphur dioxide), which has a choking smell like that produced by burning sulphur. Sometimes chemical action is brought about by the influence of light. Experiment 15. Brush over one side of a half-sheet of note- paper with a dilute solution of silver nitrate, using a clean soft brush, and put the paper away in a drawer until dry. Cut out of a piece of brown paper some figure, either a capital letter, or other design. Lay the prepared paper upon a piece of wood and fasten the brown paper figure down upon it with one or two drawing-pins ; expose it to bright sunshine for a short time. Where the prepared paper has been exposed, the sunlight will cause a chemical change to take place, which will result in the discolouration of the paper, so that on removing the brown paper, an image of the design will be seen on the under paper. This experiment is a primitive photographic process. All photography depends upon the influence of light in promoting chemical action. In a number of cases, chemical action is only able to 14 Chemical Action. proceed in the presence of a third substance, which itself remains unchanged at the conclusion of the action, and which sometimes needs only to be present in the most infinitesimal quantity. These cases are all classed together under the head of catalytic actions, the third substance being spoken of as the catalytic agent. All known instances of chemical action take place accord- ing to one of three general modes. These will be better understood after other matters have been considered (p. 155). CHAPTER III. SIMPLE MANIPULATIONS. Solution. This term is applied both to the act of dissolving, and to the product obtained by dissolving. For example Experiment 16. Throw into some water in a test-tube a little powdered potassium nitrate (nitre} ; and, on shaking, it will soon be entirely dissolved. We say that the potassium nitrate has undergone solution, and we term the resulting liquid a solution of nitre. The liquid in which a substance is dissolved is called the solvent ; thus, in the above example, the solvent was water. If such a solution be heated, or even allowed to stand exposed in an open dish, the solvent will gradually evaporate (see p. 3-), and leave the dissolved sub- stance behind. Experiment 17. Pour the solu- tion of nitre from Exp. 16 into a porcelain evaporating basin, and heat gently by means of a Bunsen with a rose burner, as shown in Fig. 4. Continue the process until the water is all driven off, and a white saline residue is left in the dish. This is called evaporating to dryness. Unfortunately the term solu- tion is employed without dis- tinction, to denote two essentially different processes of dissolv- ing. This will be best understood by the following examples, 1 6 Simple Manipulations. Experiment 18. Place a small quantity of powdered sodium carbonate ('washing' soda) in two separate glasses. To one, add some water, which will dissolve the sodium carbonate, giving a clear solution. We have therefore made a solution of sodium car- bonate in water. Into the second glass pour some dilute hydro- chloric acid, and again the sodium carbonate dissolves, and a clear solution is obtained. In this second case, however, one striking difference will be noticed, which is, that the act of solution of the sodium carbonate in hydrochloric acid is attended by a brisk effervescence. This means that some gas is being disengaged, and that a chemical change is taking place. Now place each solution in a porcelain dish and gently evaporate to dryness. In both dishes a white residue will be left, but let us try and find out whether they are the same things or not. One property of sodium carbonate has been exhibited already during this experiment, namely, that an effervescence takes place when dilute hydrochloric acid is dropped upon it ; let us therefore use this property as a test, and add a few drops of hydrochloric acid to the residue in each dish. In the case of the residue from the watery solution, there will be an effer- vescence, so that we may infer that the sodium carbonate did not undergo any permanent change by being dissolved in water, and is left unaltered when the water is evaporated away. In the other case no effervescence will take place, showing that this residue is something quite different from the original substance. Experiment 19. Place a few scraps of copper in a test-tube, and add water ; the metal does not dissolve. Pour away the water, and add a little strong nitric acid ; effervescence takes place, a brown-coloured offensive smelling gas makes its appearance, and the copper disappears. We say, therefore, that we have made a solution of copper in nitric acid ; or that copper is soluble in nitric acid. It will be noticed that the solution has a blue colour, and if it be evaporated gently to dryness a blue residue will be left, which evidently is not copper. It is a compound of copper, namely, copper nitrate. It will be evident from these experiments that two kinds of solution are recognized one being more of a physical character, in which the substance is dissolved without apparent change, and from which it is again deposited in its original state by evaporating the solvent ; while, in the other case, the act involves a chemical change, a chemical action taking place between the solvent and the dissolved substance, giving rise Decantation ; Filtration. to new compounds entirely different from either of these things. This latter process is sometimes distinguished as chemical solution. \*\ The process of solution affords an important method for separating substances that are mixed together. Thus Experiment 20. Powder a piece of marble, and mix it with powdered sodium carbonate, and throw a little of the mixture into some water in a test-tube, and shake it up. The sodium carbonate dissolves, but the insoluble marble settles to the bottom. In order to separate the solution from the sediment, one of two methods may be used, namely, decantalion we filtration. Experiment 21. Deeantation. Allow the marble to settle, and carefully pour off as much of the clear liquid as possible without disturbing the sediment. Nearly fill the test-tube again with water and, after thoroughly shaking, allow the marble to settle once more, and again decant the clear liquid. By repeating this process the marble will be washed free from the sodium carbonate. Experiment 22. Filtration. The filtering medium, almost exclusively used, is a bibulous paper (like blotting- paper), known as filter-paper. This is usually obtained already cut in circular pieces of various sizes. One of these pieces is made into a cone by being folded first into half, and then at right angles into half again, and is supported in a glass funnel of such a size that the glass will project slightly above the paper (Fig. 5). The cone is then wetted with a little clean water, and the filter is ready for use. The solution to be filtered is poured into the cone (without overflowing the paper), and the in- soluble matter is arrested by the paper, while the solution passes through quite clear. When the whole has run through, the insoluble portion is washed free from any adhering solution by once or twice filling the filter with water, and each time allowing the whole of the wash water to drain through. By conducting such an operation carefully, the exact FIG. 5. 18 Simple Manipulations. FIG. 6. quantities of both the marble and the sodium carbonate present in the mixture can be ascertained. For this purpose the insoluble residue upon the filter must be dried and weighed ; and the filtrate, together with all the wash waters, must be evaporated to dryness, and the residue (consisting of the sodium carbonate) also weighed. It will be obvious that there must be no loss of any of the solution during the whole operation. When pour- ing a liquid from a narrow vessel like a test-tube, there is no risk of it spilling by running down the outside of the tube; but in the case of wide vessels it is very likely to do so, when loss of the solution would arise. This risk is obviated by the device of pouring the liquid down against a glass rod held lightly against the edge of the vessel. Figs. 6 and 7 show the two methods. Again, in order to prevent splashing from the drops of liquid falling from the point of the funnel, the latter should be made to touch the side of the vessel placed below, so that the drops may run down the surface of the glass. To facilitate this, the stems of funnels are usually cut diagonally, as seen in Fig. 5. In the example above given, the separation was made by the physical process of solution, as the sodium carbonate did not undergo any chemical change during the process ; but quite as frequently the method of chemical solution is employed. For example Experiment 23. Mix together powdered marble and white Precipitation. 19 sand. Both substances are insoluble in water. Add to some of the mixture dilute hydrochloric acid : effervescence takes place, showing that a gas is being disengaged. When no further effer- vescence is noticed upon adding more of the dilute acid, the mixture may be filtered. The residue upon the filter is the sand. Sand is not dissolved by hydrochloric acid, but marble is. If the filtrate be evaporated to dryness, a residue wilt be obtained ; but this residue is not marble, because if a drop of hydrochloric acid be added to it no effervescence takes place. The marble and the hydrochloric acid have undergone a chemical reaction, resulting in the formation of new substances, namely, a gas (carbon dioxide), and the residue (calcium chloride). Precipitation. Sometimes when one clear solution is added to another, the resulting mixture is no longer clear ; something is produced which makes the liquid thick or muddy. Experiment 24. Dissolve a small pinch of common salt (sodium chloride) in water in a test-tube, and add a few drops of a solution of silver nitrate. Instantly the mixture becomes milky, owing to the separation of a white solid, which finally settles to the bottom of the tube. The silver of the silver nitrate has a strong chemical affinity for the chlorine of the sodium chloride, and the compound produced when these unite (namely, silver chloride) happens to be insoluble in water, and therefore the moment it is formed it separates out as a solid. This process is called precipitation^ and the solid which is produced is termed the Precipitate. A precipitate may be separated from the liquid in which it is suspended, either by decantation or nitration (see p. 17). By means of the combined processes of precipitation and nitration, it is possible to separate by chemical means sub- stances which are mixed together in solution. For example Experiment 25. Take a crystal of copper nitrate and one of silver nitrate, and dissolve them together in distilled water in a test-tube. The crystal of the copper salt will impart to the solution its own blue colour, but the liquid will be clear. Add to this a few drops of a solution of common salt. Just as in Exp. 24, there is again a white precipitate of silver chloride. Add the salt solution 2O Simple Manipulations. drop by drop until no more of the white precipitate is produced. The whole of the silver originally present in the silver nitrate is now united to chlorine, and is thrown out of the solution as in- soluble silver chloride. Now filter the mixture (see p. 17), and a clear blue liquid, containing all the copper nitrate, will pass through the filter while the silver chloride collects on the filter. This method of separating is the basis of most analytical processes. Crystallization. When we dissolve any substance in cold water, and continue adding more of the substance, a point is ultimately reached when the water will dissolve no more. The water is then said to be saturated with that particular substance. If such a cold saturated solution be warmed, it is then able to dissolve some more of the substance, until again a point is reached when it can dissolve no more. When such a hot saturated solution is again cooled, that quantity of the dissolved substance which the hot solution contained, over and above the amount which the cold water could dissolve, is thrown out of solution, and in many cases it is deposited in the form of crystals. Experiment 26. Add powdered alum in small quantities at a time to some cold water in a test-tube, constantly shaking, and allowing each little portion to dissolve before adding more. Notice that the water gradually dissolves each additional portion of alum more slowly, until at last the solution is saturated. Now heat the solution to boiling, and notice that it will now easily dissolve a further considerable quantity of the alum. Cool the solution by dipping the test-tube into cold water, and almost immediately a quantity of alum will be thrown out of solution and deposited in the form of minute crystals. Once more warm the solution, and observe that these crystals again dissolve. Now stand the test- tube down, and allow it gradually to cool by itself, and then notice that again crystals have been deposited, but that they are much larger, so that their particular shape can easily be seen. The same quantity is deposited in both cases, but when quickly cooled the crystals are smaller. Different substances are soluble in water to a very different extent; thus we find that the same quantity of cold water as is just capable of dissolving eight parts of borax, will dis- solve four times as much nitre, and twenty times as much Fusion. 2 1 zinc sulphate (white vitriol). Owing to this unequal solu- bility of various substances, the process of crystallization is often made use of in order to separate substances from each other ; more especially with a view to removing from one sub- stance any admixed impurities which are soluble, and which, therefore, cannot be got rid of by filtering. Experiment 27. Mix together about equal quantities of powdered potassium chlorate and potassium dichromate ; place the mixture in a beaker, and pour a little boiling water upon it. Place the beaker over a rose burner and boil the solution, and add just enough boiling water to entirely dissolve the mixture. The solution now has the orange-red colour of the potassium dichro- mate. Remove the lamp, and after a few minutes stand the beaker in a basin of cold water, when a crop of crystals is soon deposited. As soon as the solution is cold, decant the clear liquid and drain the crystals. Now pour a little cold water upon the crystals in the beaker, and again drain them. Observe that the solution first decanted and the wash water are coloured, but that the crystals themselves are nearly white. Rinse them once or twice more with small quantities of cold water, and see that each rinse-water be- comes less strongly coloured and the crystals become whiter. This shows that potassium dichromate is more soluble than potassium chlorate : and it will be evident that by repeating the process (that is, by once more dissolving the white salt and crystallizing it again) we can get rid of every trace of the yellow salt. Fusion is the term applied to the process of converting a solid substance into the liquid state by the application of heat. Thus when ice is warmed it enters into a state of fusion, or, in other words, it melts ; and when lead is heated it also fuses or melts. In common language the term melt is often incorrectly employed to denote the process of solution, thus sometimes it is said that sugar melts in warm water. This confusion is to be carefully avoided. When a substance in a state of fusion is allowed to cool and solidify, it very commonly assumes crystalline shapes ; thus, when water passes into the solid state we obtain a crystalline mass of ice. The shape of the crystals of ice is readily seen by allowing a single snow-flake to fall upon the coat sleeve, and looking at it through a pocket lens. The 22 Simple Manipulations. magnificent fern-like crystals of ice upon a window-pane in winter are familiar to all, and very often single star-shaped crystals of great beauty are to be seen. Chemical action often takes place between substances in a state of fusion } which is incapable of taking place when they are only in solution ; for example Experiment 28. Dissolve a small piece of potassium hydroxide (caustic potash} in water, and add to the colourless solution a few grains of powdered manganese dioxide. The black powder simply falls to the bottom of the solution, and no action takes place. Now place another piece of potassium hydroxide in a dry test-tube and heat it : the solid melts to a colourless liquid. While it is in this ftised condition add a few particles of the manganese dioxide, and notice a very different result. The liquid turns to a blue-green colour, and the black oxide of manganese has disappeared. If this be allowed to cool, it again solidifies to a green solid mass ; and when quite cold, if water be added, it dissolves, giving a green solution. A chemical action has taken place between the man- ganese dioxide and the ftised potash forming a compound called potassium manganate, but the solution of potash was incapable of bringing about this reaction. Distillation. The principle of this process has already been explained (p. 2) ; but the process is more conveniently carried on by means of the apparatus in Fig. 8. The flask, a, known as a " Wurtz " flask, contains the liquid to be distilled. Its brancli tube passes through a cork in the end of the Liebigs condenser. This consists simply of a straight tube, T, which is jacketed by a wider tube, W, through which a stream of cold water circulates, the water entering at the lower branch tube, B, and passing out at B'. In this way the inner tube is kept cold. The neck of the Wurtz flask is fitted with a cork, through which a thermometer is fixed, in order to tell the temperature at which the liquid distils. Experiment 29. About half fill the Wurtz flask with water, which has been made dirty by the addition of a small fragment 01 clay, or a little ink, and replace the cork carrying the thermometer, the bulb of which must not reach down into the water. Boil the water by means of a Bunsen lamp, and place a clean dry flask Distillation. 23 to receive the distillate. During the whole experiment a stream of cold water must be circulating through the condenser. As soon as the water in the flask begins to boil, carefully read the thermometer. FIG. 8. Notice that the mercury gradually rises until a point is reached when it remains stationary. Note this temperature. Observe also that the distillate is perfectly colourless, all the impurity which rendered the water in the flask dirty remains behind, and only clean water passes over. The process of distillation therefore purifies the water ; and distilled water is purer than ordinary water. Distillation, also, often enables chemists to identify a liquid. For instance, a colourless liquid is given you, which looks like water; if it be submitted to distillation, and the temperature noted at which the mercury in the thermometer is stationary while the liquid is briskly boiling, you could tell at once whether it was water or not The process of distillation is useful also as a method of separating liquids which are mixed, but which boil at different temperatures. Experiment 30. Place in the Wurtz flask a mixture of about two parts of water and one of alcohol, and have ready three small clean flasks to receive the distillate in. When this mixture is distilled, observe that the mercury in the thermometer quickly rises Simple Manipulations. to a certain point and then remains steady for a time. Presently it begins again to rise ; then change the receiver and collect what passes over separately. At length the mercury is again steady, indicating the temperature at which the water in Exp. 29 distilled ; at this point exchange the second receiver for the third, and continue the process a little longer. In receiver No. i the liquid consists mainly of alcohol, but mixed with a little water. Pour out a little of the liquid and set fire to it, thus proving that it is fairly strong alcohol. Receiver No. 2 contains a mixture of alcohol and a large proportion of water. This liquid will not burn at all. The third receiver contains water free from alcohol. By this process, therefore, we have obtained a portion of one of the liquids, namely, the water, free from the other liquid, but have not completely separated each from the other Collecting Gases. Gases are usually collected by causing them to bubble through water into a bottle or cylinder filled with water, and standing mouth downward in a trough or basin of water. This method is usually described as collecting over wafer, and the basin used is called a pneumatic trough. Experiment 31. Fill a glass cylinder to the brim with water, and slide on to the ground lip a ground glass plate so that no air bubble is included. Grasp the cylinder with the left hand, as shown in Fig. 9, and hold the glass cover in its position with the forefinger of the right hand. Then invert the cylinder, when it will be in the position shown in Fig. 10, and lower its mouth beneath the surface of the water in the basin or trough, as in Fig. ir, and now with- draw the plate. The cylinder then remains rilled with water. In the figure a glass basin is shown, this is in order to render the position of the cylinder visible ; in the laboratory it is more usual to use pneumatic troughs made of metal, having a movable metal FIG. 9. Collecting Gases. 25 shelf with an oblong hole in it through which the bent end of a tube can be introduced. Take a piece of glass tube, bent as shown in Fig. 12, and gently blow through it so that the air from the lungs shall bubble up into the cylinder. As the gas ascends in bubbles, it gradually displaces the water, until at last the cylinder is full of gas. To remove the cylinder in order to examine the gas, slip the ground glass plate beneath its mouth, and. keep it in its place with the finger while the cylinder is being turned over. Now remove the plate, and lower a lighted- candle on a wire, or a FIG- 10. burning taper bent as shown in Fig. 13, into the gas in the cylinder. Notice that the flame is extinguished, showing that the gas which FIG. ii. comes out of the lungs is different from the air which is drawn into them, because it will not allow a candle to burn in it. 1 1 The nature of this gas will be described later on. 26 Simple Manipulations. Some gases cannot be collected over water, because either they dissolve in water, or they enter into chemical combination with water. In such cases the mercurial pneumatic trough may be used : that is to say, the liquid metal mercury is used instead of water, provided the gas is without chemical action upon mercury. When a gas happens to be either much lighter or much heavier than air, we may collect it without using a pneumatic trough at all. Experiment 32. Obtain a ground glass plate with a hole in the centre, 1 and fit the hole with a cork carrying two bent tubes, one FIG. 12. FIG. 13. FIG. 14. long, and the other quite short, as shown in Fig. 14. This is supported on a ring, and the cylinder to be filled is stood over the tubes. Attach the long tube to the ordinary coal-gas supply, and turn on the gas. Coal-gas is very much lighter than air (the fact that it is used for filling balloons shows this), and therefore, by delivering it right up to the top of the cylinder, it accumulates in the upper part of the vessel, and gradually displaces the whole of the air, driving it down through the other tube. We can tell 1 A stout piece of cardboard will answer the purpose, although the glass plate is preferable. Collecting Cases. 27 when the air is all displaced by smelling the gas which will then be escaping at the exit tube. This method of collecting is called upward displacement. Gases that are heavier than air can be collected by down- ward displacement. This is exactly the reverse of upward displacement, and is carried out by simply inverting the apparatus in Fig. 14, so that the cylinder stands mouth upwards. CHAPTER IV. FITTING UP APPARATUS. MUCH of the apparatus required for such experiments as the elementary student will make, he can himself easily put together by means of glass and rubber tubes, corks, and glass bottles. Figs. 15, 16, and 17 show three typical forms of apparatus used for the preparation of various gases, and every student should fit up these for himself. Glass tubes. Glass tubing is manufactured in many FIG. 15. varieties of glass, and of many sizes. For general purposes it is best to obtain soft soda glass ; and a most useful size is shown in section at 0, in Fig. 18. Such tube is cut to any required length by making a slight scratch upon it with the edge of a fine three-cornered file, and Fitting up Apparatus. 29 then breaking it across exactly as one would snap a dry twig of wood. In order to bend a tube, it is made soft by being heated in a common flat gas-flame. Experiment 33. Hold a short piece of tube in the flame in the manner shown in Fig. 19, and slowly rotate it between the thumbs and forefingers in order that it may get equally heated all round. FIG. 16. FIG. 17. As soon as it is quite soft, "withdraw it from the flame, and deliber- ately bend it to a right angle (Fig. 20). When cold, wipe off the soot. (Such a right angle bend will be required at a, a', a" Figs. 1 6 and 17.) Tubes should not be heated for bending in a Bunsen flame, as the glass in this case will become creased at the bend, as in Fig. 21, B. If the tubing employed is too thin in the walls (, Fig. 18), the glass will collapse at the bend, as shown in Fig. 2 1, A. Both of these bends, besides being unsightly, are very easily broken, and they also prevent the free passage of gas through them. When a tube is to be bent into a very acute angle, it should be held in the flame in the manner shown in Fig. 22, with a FIG. 18. Fitting up Apparatus FIG 19 FIG. 20. FIG 21. Bending Glass Tubes. 31 the knuckles downwards instead of upwards, as in Fig. 19, so that the bend may be produced by bringing the hands together FIG 22. FIG. 23. upwards, as in Fig. 23. Such a bend is required in making the delivery tub.es d, d, d", Figs. 16 and 17. The bent tube 32 Fitting up Apparatus. is cut at r, Fig. 24, and the end E is then bent as shown by the dotted lines, by holding it in the flame as explained in Exp. 33. All glass tubes, when bent and cut to shape and length, must be rounded or smoothed at their extremities by holding them in a flame (Fig. 25) until the "raw" or sharp edges have just fused. " Combustion " tubing is made of hard glass, which will stand a high temperature without softening. It is useful when we require to pass a gas over some heated material. c. Fig. 1 8, shows a convenient size for combustion tube. Corks. Select a sound cork which is just too large at its narrow end to fit into the mouth of the flask or FIG. 24. FIG. 25. test- tube, and then squeeze it either in a cork- squeezer or by rolling it beneath the foot on the floor, using a moderate pressure upon it. The cork should then fit comfortably into ihe flask. The cork has next to be bored to take the glass tube which is to pass through it. This is best done by means of a cork-borer, a tool consisting of a brass tube, with a sharpened edge at one end ; it is bored into the cork much in the same way as a gimlet or brad-awl is used to bore a hole in wood, except that the cork-borer is usually wetted. In order to be sure that the borer selected will make a hole the exact size required for the glass tube, a trial hole Fitting up Apparatus. 33 should be bored in a waste scrap of cork. The hole must be of such a size that a little gentle force is required to push the tube through it. When using the cork-borer, the tool should be driven in at the narrow end of the cork, care being taken to make the hole in the centre, and to keep it straight. If the borer is fairly sharp, the hole will be clean and smooth. Sometimes corks are bored by means of a round (or rat- tail) file; but this is not a good plan, as it is much more difficult to make the hole perfectly round ; and if it is not round, the glass tuj)e will not fit properly, and the apparatus will leak at this point. If the bent tube has been rounded at its ends (p. 32), it can be pushed into the cork without cutting or tearing the hole, and a tight fit will be made. The delivery tube is attached to the exit tube by means of a short piece of indiarubber tube. The ends of both glass tubes being smoothed, as already described, they can easily be pushed into the caoutchouc connection without cutting it, especially if the latter be moistened inside by breathing through it immediately before introducing the glass. The apparatus, type No. 2 (Fig. 16), differs only in being fitted with a cork bored with two holes, into the second of which there is fitted a tube known as a thistle Junnel. This reaches nearly to the bottom of the flask, so that it may dip into the liquids which will be present ; and thus, while allowing liquids to be poured in, will prevent gas from escaping through the funnel. The pieces of apparatus of the rirst and second types are intended to be used in certain experiments when heat is applied, hence they have to be supported by suitable clamps or stands at a convenient height, and therefore the delivery tubes must be made of such a length that they will reach into the pneumatic trough. The apparatus of type No. 3 (Fig. 17) is used in the preparation of certain gases, when it is not necessary to apply any heat to the materials. The generating vessel in this case, instead of being a thin glass flask, may be either a two-necked bottle, B, Fig. 17, known as a " Woulfs bottle," or an ordinary D 34 Fitting up Apparatus. wide-mouthed bottle. In the latter case, a single cork with two holes is fitted with two tubes as in A, Fig. 17 ; but when the two-necked bottle is used the thistle funnel is fitted into one neck, and the exit tube to the other. This form of apparatus is preferable to the other, as small corks are more easily fitted so as to be free from leakage than wide ones. When the apparatus has been put together, and before being used, it should be tested to ascertain whether it is tight. In the case of A and B, Fig. 15, this is done by sucking a little of the air out of the apparatus by applying the mouth to the end of the delivery tube, and instantly closing the tube with the tip of the tongue. If the joints are tight the tongue will remain drawn to the end of the tube, and a little effort is felt in pulling it away; whereas if there is any leakage in the apparatus the tongue parts at once away from the tube. To test the other forms of apparatus in Figs. 16 and 17, a quantity of water should first be poured into the flask or bottle, until the end of the thistle tube dips into the liquid. Then by applying the mouth to the end of the delivery tube and blowing gently into the apparatus, the water will be forced up the funnel tube. It should in this way be driven nearly up to the head of the thistle tube, and the end of the delivery tube closed with the tongue. If the apparatus does not leak, the water will remain steady in the thistle tube, otherwise it will gradually sink down. If the apparatus leaks it should be refitted, and on no account should a leaky apparatus be tight by the use of scaling wax or other lutes. CHAPTER V. SIMPLE GLASS-BLOWING OPERATIONS. GLASS-BLOWING operations, as a general rule, require skill, practice, and patience to perform with anything like success ; nevertheless there are a number of smaller and simpler operations which can easily be done by the young student, and which it is most useful that he should know how to perform. The bending of glass tube and the rounding of the ends have already been described. For ihese operations, however, an ordinary gas-flame is employed, but for those now to be described the blow-pipe is to be used. A blow-pipe and some sort of foot-blower usually form a part of the regular fittings of a chemical laboratory, but even in their absence much may be done by means of a small Herapath mouth blow-pipe. General rules. (i) Never bring a piece of cold glass directly into the blow-pipe flame, but first warm it in the smoky flame before admitting wind from the blower. (2) When a tube is being heated, it should (except in special cases) be continuously revolved in the flame, so that the heating may be uniform; and also, as it gets soft, to prevent the glass from falling out of shape. (3) When actually blowing glass, always remove the soft glass from the flame. (4) Always begin the blowing gently, and then regulate the force of the breath as the soft glass gives to the pressure. To open out the end of a glass tube. When a cork is to be fitted into the end of a glass tube, the end should be opened out a little, or " bordered." The simple tool required Simple Glass-blowing Operations. for this is a round stick of charcoal, pointed at one end like a lead pencil. Experiment 34. Take a piece of moderately wide glass tube and warm one end in the smoky flame (Rule i). Then admit wind into the flame, and hold the tube in the position shown in Fig. 26, re- volving it all the time (Rule 2). When the end of the glass is sufficiently soft, remove it from the flame, and push the pointed piece of charcoal into it, giving a screw- ing motion to the charcoal (Fig. 27). When the tube is sufficiently opened, hold it in the smoky flame again, gradually turning down the gas. This anneals the glass, that is, cools it slowly, and makes it less liable FlG - 26 - to crack afterwards. FIG. 27. To draw down a glass tube to a jet. Experiment 35. Heat a piece of tube in the ( blowpipe flame, FIG. 28. holding the glass as in Fig. 19. As soon as it is soft, remove it Joining Glass Tubes. 37 from the flame, and pull gently, revolving each end slowly at the same time. The glass then assumes the form shown at a, Fig. 28, the walls of the tapering and narrow parts being very thin. Now heat another piece, keeping it longer in the flame ; observe that the glass gradually thickens and the walls fall together as the mass gets softer and softer (<, Fig. 28). Keep the tube quickly revolving, or the soft part will drop. Very gently draw the ends apart, still revolving, and the tube will take the shape seen at c, Fig. 28, where the tapering and narrow parts are thick walled. The tube may then be cut at any desired point on the narrow part by means of a file scratch. To seal up the end of a glass tube. Experiment 36. Draw out a piece of glass tube, as in a, Fig. 28. Cut it off at the dotted line e, and heat the narrowed end of d in the blow-pipe flame until it closes up, when it presents the appear- ance shown at f, Fig. 29. Then heat the somewhat thickened end in the blow- pipe, and, when just soft, blow gently into the tube (note Rules 3 and 4), slowly re- volving the glass at the time. It should then appear as seen in gj Fig. 29. If too much pressure was used in blowing, the end will be expanded into a swelling. It may, however, be reduced by revolving it in the blow- pipe flame, when the sides of the enlarged part will fall together again. To join two tubes. Experiment 37. Take two pieces of tube of the same diameter, close one end of one with a cork, and heat the opposite end in the blow-pipe. At the same time heat one end of the other tube. When the ends are soft, bring the two together with a little pressure. This causes them to adhere, and, at the same time, slightly to bulge out at the junction, as at a, Fig. 30. Then, while still soft, blow gently into the tube, drawing 38 Simple Glass -blowing Operations. it out slightly at the same time so as to keep the outer walls parallel. If the blowing operation has not followed the first suffi- ciently quickly, the joint, as it is at a, may be re-heated in a fine- pointed blow-pipe flame, and then blown as described. The joint should have the appearance shown at b, Fig. 30. Experiment 38. Join a wide tube to a narrow one. First draw out the wider tube, and cut it off when it has a diameter equal to the narrow tube. Then heat the drawn-out end of the wide tube, and one end of the narrow tube, and join them in the way de- scribed in Exp. 37, blowing through the narrower tube. To seal platinium wire into glass tubes, (a) In the end of narrow tubes. Experiment 39. Draw the tube out to a point, and cut it oft so that the wire can just pass through the drawn-out end. Then introduce the tip into a flame, when |p ^K the edges of the glass will close to- gether round the wire, as in Fig. 31. Fio ^ While the glass is still soft, the position of the wire can be adjusted, so as to get it quite straight. (b) In either the end, or the side of a wide tube. Experiment 40. Heat the glass, at the point where the wire is to be introduced, with a fine-pointed flame, and when a small spot is soft, stick the end of a platinum wire into it and draw out gently. In this way a tiny branch tube is made, , Fig. 32. This is then cut off short, as at b. The wire is then inserted, and the flame again directed upon it, when, as in Exp. 39, the glass closes up round the wire, c, Fig. 32. To blow a bulb on the end of a tube. Experiment 41. First seal up the tube as shown at /, Fig. 29. Then heat the extreme end where the glass is thick, and gently blow it out so as to obtain the result seen at a, Fig. 33. Next hold the tube in a large blow-pipe flame, heating mostly the part between the dotted lines ; and, as the glass softens, keep quickly revolving it. It then assumes somewhat the shape seen at b, Fig. 33. Withdraw it from the flame, and blow steadily into the tube, holding it in a horizontal position, and revolving the glass all the time, c t Fig. 33. Blowing Bulbs. 39 [Probably the first attempts will either be complete failures, or very remarkably shaped bulbs, but with a little patience, better results will soon follow.] To blow a larger bulb, more glass is necessary, FIG. 32. FIG. 33. and it is better to first join on a piece of larger tube as described in Exp. 38. To blow a bulb on the middle of a tube. Experiment 42. Close one end of the tube with a cork, and heat in a large flame at the spot where the bulb is to be blown. As the glass softens and thickens, gently press it together so as to ac- FlG. 34. cumulate material for the bulb, Fig. 34. Then blow steadily, holding the tube in a horizontal position, and revolve it between the fingers. CHAPTER VI. HYDROGEN. ALTHOUGH solids and liquids are more familiar to us than gases, it will nevertheless be more convenient to begin the study of chemical facts by considering some of the methods of preparing, and a few of the more prominent properties of the gaseous element hydrogen. Let us understand at the very outset, that chemists cannot create; they cannot make hydrogen from either nothing, or from materials which have got no hydrogen in them. The alchemists of old believed in the transmutation of the metals, and they spent their lives endeavouring to change the common metals into gold ; now- adays we know how futile the attempts were, and we no more expect to convert copper into gold, than we expect to " gather figs from thistles." All that the chemist can do is to so experiment upon compounds which contain hydrogen as one of their constituents, as to cause chemical changes to take place which will result in the expulsion of this hydrogen. We must, therefore, select some suitable compounds containing hydrogen from which to obtain this element. Chemists have found out, by numberless experiments, that hydrogen forms one of the constituents of a vast number of compounds; thus it is found to be present in nearly all animal and vegetable substances. It is a constituent of water, and also of all those things chemists term acids. There are three common compounds from which the element is most usually obtained ; these are (i) Water, (2) Sulphuric Acid, (3) Hydrochloric Acid. Hydrogen. 4 1 (i) Hydrogen from Water. As water is the commonest of these three substances, we will first experiment with it. Water is composed of the two elements, hydrogen and oxygen, chemically united together, and in order to separate them by chemical processes, we must find some element which, under suitable conditions, can overcome the force uniting the oxygen and hydrogen, some element which can seize the oxygen and tear it away, so to say, from the grasp of the hydrogen. Many metals are capable of doing this. Experiment 43. Carefully throw a small fragment of sodium, about the size of a pea, upon some cold water in an ordinary dinner-plate. Notice that the metal at once melts, and the globule swims to and fro upon the surface of the water (much in the same manner that a drop of water runs about upon a hot iron), producing a hissing sound. Observe that the globule quickly gets less and less, and finally disappears. Now take a strip of turmeric paper (that is, blotting-paper which has been dyed yellow with turmeric) and dip it first into clean water ; this causes no stain upon it ; now dip it into the water in the plate, and observe that it is strongly stained brown. Also dip the fingers into the water, and notice that it feels slimy, or caustic, to the touch. This shows that there is something in the water after the sodium has been in contact with it which was not previously there. One of the products, there- fore, of the action of the metal sodium upon water is something which dissolves in the excess of water present, yielding a solution which turns yellow turmeric brown, and is caustic to the touch. The other product of the action was hydrogen gas, which escaped unnoticed into the air. If we repeat the experiment, using potassium instead of so- dium, the hydrogen will not escape our observation. Experiment 44. Throw a similar frag- ment of potassium upon some clean water in a plate, and at once cover the whole with a glass bell jar, as in Fig. 35. The potassium appears to take fire the moment it touches FlG the water, but really it is the hydrogen which burns, 1 the hydrogen which is driven out of its combination with oxygen. The action of the potassium upon water develops so 1 A little of the potassium burns also, and this gives to the flame of the burning hydrogen the violet colour* 42 Hydrogen. much heat as to set fire to the hydrogen. Carefully notice that for a few seconds after the flame goes out, a little red-hot globule of something in a melted state remains swimming upon the water, and then suddenly disappears with a little splutter. (// is in order to Prevent this substance from being scattered, and injuring the eyes, that the plate must be covered with the bell glass?) Test the water in the plate with turmeric paper, and the same stain is produced as in the former experiment. Therefore, when potassium acts upon water, hydrogen gas is expelled, and a substance is formed which also stains turmeric. We must now adopt some device for collecting fat hydrogen gas which can thus be expelled from its combination with oxygen, by the action of sodium upon the water. There are various ways of doing this, but only one of them is free from danger ; for sodium, when brought carelessly into contact with water, is liable to give rise to a serious explosion. Experiment 45. Take a piece of lead pipe, 2\ centimetres long (i inch) and centimetre bore, and close up one end by hammering the lead, as seen in Fig. 36. This little pipe is then filled with sodium by first rolling a pellet of the metal (which is about as soft as wax) between the fingers until it will just push into the tube, and then forcing it in by pressing the mouth of the tube firmly down upon the table. The excess of sodium round the mouth is then trimmed off with a knife. When this tube, with its little G< 3 ' charge of sodium, is dropped into water in a dish or trough, it, of course, sinks to the bottom, and a stream of gas- bubbles will be seen rising through the water ; and if the mouth of an inverted glass cylinder, filled with water, be brought over the ascending bubbles, as in Fig. 37, the gas will collect in the upper part. If the sodium is all used up before the cylinder is full, a second leaden tube can be placed in the water. Now remove the cylinder, as described (page 25), and apply a lighted taper to the mouth of the jar. Notice that the gas burns quickly, with rather a yellowish flame, this yellow colour being due to a trace of sodium. If the water in the trough be tested with turmeric paper, it will stain it brown, as in the former case. This other product of the action of sodium on water, which remains dissolved in the large excess of water in the dish, is called Hydrogen from Water. 43 sodium hydroxide, or caustic soda ; while that which was formed in the case of potassium was potassium hydroxide. Many other metals besides sodium and potassium are able to expel the hydrogen from water if the conditions are favourable. Thus Experiment 46. Drop a few fragments of magnesium into cold water in a test tube, observe that no action takes place between FIG. 37. the metal and the water ; in this respect, therefore, magnesium is a contrast to sodium or potassium. Now boil the water, and notice that even at the boiling temperature practically no hydrogen is given off. This experiment shows that magnesium is not able to decompose water appreciably even at the boiling point. Let us, therefore, try the effect of heating the metal much more strongly in a current of steam : that is, of heating it in contact with water in the state of gas. Experiment 47. Fold up a short strip of magnesium ribbon, and place it in a hard glass bulb which is held in a slightly inclined position by a clamp, as in Fig. 38. Attach to one end of 44 Hydrogen. the bulb tube a small empty flask, in the manner shown in the figure, and connect this to a tin can in which water is being boiled. As soon as steam issues from the bulb tube, gradually heat the latter all along, by means of a Bunsen flame, until the inside is quite dry, and the steam no longer condenses as it passes through. Then hold the flame steadily under the magnesium, and heat it until it is nearly red hot, when it will suddenly take fire and burn in the steam, combining with the oxygen of the steam and letting the hydrogen go free. Directly this takes place, light the hydrogen as it escapes from the end of the tube. Notice that it burns with a scarcely visible flame, because magnesium does not impart FIG. 38. any colour to a hydrogen flame like sodium and potassium do. Notice also that the residue in the bulb is white. Shake out some of this and see if it will dissolve in water by adding water to a little in a test-tube. Note that it is not percep- tibly soluble. Take a little more of it and place it on a piece of turmeric paper, and moisten it with a drop of water ; notice that the paper is stained slightly brown, showing a slight resemblance between the behaviour of this substance (magnesium oxide or magnesia) and sodium hydroxide. The common metal iron, at a bright-red heat, will also expel the hydrogen from water when the latter is in the condition of water vapour or steam. This method of obtain- ing hydrogen is often made use of on a large scale. Experiment 48. Take a piece of ordinary iron gas-pipe, suffi- ciently long to project 15 cm. (6 inches) beyond each end of the furnace to be used. Fill the pipe with small iron nails, and fit each end with a cork with a short straight glass tube. To one Hydrogen from Sulphuric Acid. 45 end a delivery tube is attached, while the other is connected to a small steam boiler (conveniently made out of a common tin can, see Fig. 38). The iron pipe is heated to a bright-red heat, either in a gas or coke furnace, and steam is passed through it. The red hot iron seizes the oxygen of the steam, combining with it to form iron oxide, which remains in the tube, while a rapid stream of hydrogen is evolved, and this may be collected over water in the pneumatic trough. (2) Hydrogen from Sulphuric Acid. Although water is the commonest compound of hydrogen, sulphuric acid is the one from which it is most convenient to obtain hydrogen, and when we require this gas for experiments, it is almost always got from sulphuric acid by the action upon it of the metal zinc, Experiment 49. Place some granulated zinc x in a two-necked FIG. 39. Woulf's bottle, arranged as in Fig. 39, with just sufficient water to cover it, and pour upon it by means of the thistle funnel a little strong sulphuric acid. Almost immediately it will be noticed that an effervescence begins, showing that gas is being disengaged. Allow the action to go on for a few minutes until all the air origi- nally present in the bottle has been swept out by the gas, and then collect the gas at the pneumatic trough. Three or four cylinders should be filled for subsequent experiments. In this experiment the metal zinc expels the hydrogen from sulphuric acid, and combines with what is left of the 1 Granulated zinc is made by first melting the metal and pouring it in a thin stream into cold water. 46 Hydrogen. sulphuric acid after the hydrogen is gone. The compound so produced, called zinc sulphate (or white vitriol), is left dissolved in the water in the bottle. It may be obtained from the solution by the following experiment. Experiment 50. When the above experiment is concluded, pour the liquid out of the Woulf 's bottle and filter it. Then gently evaporate it in a dish over a small flame until it is reduced to about half the bulk, and allow it to cool, when long glassy-like colourless crystals will deposit. This is the zinc sulphate. Instead of employing zinc in Exp. 49, the metal iron might have been used ; but the hydrogen obtained in this way is always contaminated with compounds of hydrogen with both carbon and sulphur (both of these elements always being present in ordinary iron), which give to the hydrogen an unpleasant smell. Experirnent 51. Put a small quantity of iron filings into a test- tube, and pour upon them a little dilute sulphuric acid. Notice effervescence, due to the escape of gas. Smell the gas, and note the peculiar and nasty odour. This smell does not belong to hydrogen, but to the impurities present. Bring a lighted taper to the mouth of the tube ; observe that the gas burns. When the action has continued for some time, the liquid may be filtered, and the clear solution evaporated, when small green crystals will be deposited. This substance is the compound of iron with what is left of the sulphuric acid after the hydrogen in it is expelled. It is called ferrous sulphate (or green vitriol). In a similar manner we can expel the hydrogen from sulphuric acid by means of the metal magnesium. Experiment 52. Drop a few fragments of magnesium ribbon into a little dilute sulphuric acid in a test-tube. Notice how briskly the gas is given off. Show that it is hydrogen by lighting it at the mouth of the tube. When the metal is all dissolved, this solution also may be evaporated down and allowed to cool, when colourless crystals of magnesium sulphate will be formed. Compare these crystals with those of zinc sulphate. (3) Hydrogen from Hydrochloric Acid. The hydrogen from this compound is also expelled by the metals, zinc, iron, and magnesium ; therefore, hydrochloric acid may be The Properties of Hydrogen. 47 used instead of sulphuric acid in Exps. 49 to 52. The compounds which remain behind in the solution would, in this case, be compounds of the metals with what is left of hydrochloric acid after the hydrogen has been expelled ; they would be zinc chloride, ferrous chloride, and magnesium chloride respectively. The Properties of Hydrogen. From the various samples collected, we learn that the gas is colourless (there- fore invisible). Also that it does not appreciably dissolve in water, for when left standing in vessels in the pneumatic trough the water does not show any signs of rising up in the cylinders, which it would if the gas were soluble. If we take one of the jars of gas and smell it, we shall find that it is quite odourless. We have also learnt from Exps. 45 and 47 that hydrogen will burn. It is important to remember that if we take any gas which will burn in the air, and previously mix it with a certain proportion of air, and then bring a light to the mixture, an explosion results. Every one knows that it is dangerous to bring a light into a room where there is an escape of coal-gas, that is, where there is a mixture of gas and air. Coal-gas alone burns quietly ; but a mixture of coal-gas and air in certain proportions explodes when lighted. This is particularly true of hydrogen. We have seen by Exp. 47 that, when un- mixed with air, it burns quietly, but if a mixture of hydrogen and air be lighted, the mixture explodes with great violence. Therefore, before bringing a flame to the tube leading from a hydrogen apparatus, it is most important to be sure that all the air originally present in the bottle has been swept out. In order to realize the force of the explosion that would occur under such circumstances, and the danger which would result from neglecting this precaution, the following experiment may be made Experiment 53. Fit an ordinary pear-shaped soda-water bottle with a cork, through which is fixed a short piece of the stem of a clay tobacco-pipe, only just projecting through the cork. Put a little granulated zinc into the bottle, add some dilute sulphuric acid, and insert the cork. Hold a lighted taper to the end of the Hydrogen. tube as though attempting to inflame the hydrogen, and in a few seconds, when the gas has mixed with the air in the bottle in a certain proportion, a loud explosion will result, which will shoot the cork out of the bottle with some violence. There is no fear of the thick bottle bursting, but it should be held in such a position that the cork will not fly in a direction where it can do harm. Experiment 54. Disconnect the bent delivery tube from the hydrogen apparatus, and by means of a rubber tube attach a straight glass tube. Fill a test-tube with the gas by upward displacement, as in Fig. 40. After a few moments withdraw the delivery tube, and ap- ply a light to the mouth of the test- tube. If the hydrogen is mixed with air, a slight pop will result, but if free from air it will burn quietly. If this is the case the gas may be lighted at the end of the de- livery tube. Notice that the flame has a yellowish colour, this is due to the soda present in the glass (recall the appearance of the flame when the hydrogen collected in Exp. 45 was burnt). The true appearance of a flame of hydrogen may be seen by sub- stituting for the glass tube a short piece of lead pipe, into one end of which an ordinary metal gas-burner has been screwed. Notice that the flame now is almost without colour, being slightly bluish, and gives no light. Place an ordinary coal-gas flame by the side of the hydrogen flame and compare them. Depress into each flame a white plate, observe that the bright flame blackens the plate, while no soot is deposited from the other. Hold a clean dry jar over each flame for a few moments, notice that in each case the flames are giving FIG. 40. The Burning of Hydrogen. 49 off steam, because moisture will be deposited on the cold sides of the glass. Cover each of the jars with a glass plate, or piece of card, and pour into each a little clear lime-water. After shaking the lime-water with the air in the jars, it will be seen that in the jar which was held over the gas flame, the lime-water has turned milky, while in that which was over the hydrogen it remains clear. These experiments teach us that water is formed when we burn either hydrogen or coal-gas ; and that from burning coal- gas we also get, besides the water, a gas which will turn lime-water milky. Hydrogen is the lightest substance known to chemists. It is nearly 14^ times lighter than air. On account of its extreme lightness it can be collected in vessels by upward displacement (see Fig. 41). It may also be poured upwards from one vessel to another. Experiment 55. Take a cylinder containing only air, and hold it mouth down- wards in the left hand ; FIG. 41. then take another cylinder filled with hydrogen in the other hand, and bring its mouth just beneath that of the first, in the manner shown in Fig. 41, and gradually empty its contents up into this one by lowering the foot until the position shown in Fig. 42 is reached. Now stand the lower cylinder down, and apply a light to the contents of first one and then the other. Notice that no hydrogen is left in the one which formerly was full, and that the contents of the upper jar burn with the characteristic flame. Owing to the extreme lightness of hydrogen, it is some- times used for filling balloons. On a small scale we can imitate this by filling soap bubbles with the gas. Experiment 56. Attach a common clay tobacco-pipe to the hydrogen apparatus, and proceed to blow a bubble, holding the pipe in the usual position. Notice that the bubble in trying to E 5O Hydrogen. ascend, very soon begins to curl over on to the outside of the pipe, as in Fig. 43. Turn the pipe upside down, and blow another, and observe that as it grows it struggles to tear itself away from the pipe, assuming the shape seen in Fig. 44. When disengaged from the pipe it ascends very quickly. A pretty experiment is to make a small hydrogen bubble carry up a large air bubble. To do this, first blow an ordinary bubble with the breath. Then, while it is still hang- ing on the pipe, bring under it a hydrogen bubble just beginning to form (Fig. 45). As the latter increases in size, gradually invert the two pipes so that the one holding the hydrogen bubble is uppermost. Then, with a slight jerk, detach first the pipe from the air bubble, and next the one con- veying hydrogen. The double bubble then slowly rises (Fig. 46), and generally when it touches the ceiling the little hydrogen one breaks, and allows the other to fall to the ground again. This experiment requires a good soap solution. 1 FIG. 42. FIG. 43 . Although hydrogen will burn, it will extinguish the flame of any ordinary burning substance. This is true of all gases 1 The following is a good receipt. Dissolve 2 grams of sodium oleate in 80 cc. cold distilled water, add 20 cc. glycerine, well shake, and put away in a dark cupboard for two days to settle. Then carefully pour off the clear liquid, add just I drop of ammonia and shake up. The Lightness of Hydrogen. 52 Hydrogen. which burn in the air, that they will put out the flames of all other things which bum in the air. Experiment 57. Thrust a lighted candle (or piece of thick taper), fastened to the end of a wire, into a cylinder of hydrogen held mouth downward. As the burning candle approaches the mouth of the vessel, it there sets fire to the hydrogen ; but as it is pushed up into the gas its own flame is extinguished. Withdraw the candle and ,. v _ . relight it as it passes out through the still / * . burning hydrogen. The same will happen if a paper spill be used instead of a taper, or if an ordinary coal-gas flame be introduced. We therefore say that hydrogen is a combustible gas, but that it is a non-supporter of combustion. EPITOME. The element hydrogen is found free in only very small quantities on the earth ; but it is present on the sun in enormous quantities. It occurs in nature in combination with oxygen in water ; with carbon in marsh gas ; with sulphur in sulphuretted hydrogen. It is present in nearly all animal and vegetable substances, and is a constituent of all acids. It can be obtained from water : (i) by the action of either sodium or potassium at the ordinary temperature ; (2) by the action of magnesium or iron at a red heat. It may be got from either sulphuric acid or from hydrochloric acid by the action of either zinc, iron, or magnesium upon them. The common laboratory method is by acting on sulphuric acid with zinc. Hydrogen is a colourless, tasteless, odourless gas. It burns with a nearly invisible bluish flame. A mixture of hydrogen and air explodes when lighted. It extinguishes the flames of other ordinary burning substances. Hydrogen is the lightest of all known substances, being 14-4 times lighter than air. For this reason it is taken as the standard or unit for comparing the densities of all other gases. Thus we say that the density of air is 14*4, which means that air is 14-4 Hydrogen. 53 times heavier or denser than hydrogen, bulk for bulk. Hydrogen being thus taken as the unit, obviously its density is unity or i. Reactions for hydrogen * (1) From water, by the action of sodium H 2 O + Na = NaHO + 11. (2) ,, ,, ,, magnesium H 2 O + Mg = MgO + H 2 . (3) *, iron4H 2 O + 3Fe = Fe 3 O 4 + 4H 2 . (4) From sulphuric acid, by the action of zinc H 2 SO 4 + Zn = ZnSO., + H 2 . (5) ,, ,, magnesium H 2 SO 4 + Mg = MgSO 4 + H 2 . (6) ,, ,, ,, iron H 2 SO 4 + Fe = FeSO 4 + H 2 . (7) From hydrochloric acid, by the action of zinc 2 HC1 + Zn = ZnCJ 2 + H 2 . (8) The combustion of hydrogen in air or oxygen H 2 + O = H 2 O. 1 These signs and symbols will be explained later on ; the student may pass them over at this stage. CHAPTER VII. OXYGEN. THIS element, like hydrogen, is a gas, but in almost every other respect it presents a complete contrast to that element. There are a great many compounds containing oxygen as one of their constituents, from which the element can easily be obtained, for in combination with others it is the most abundant of all the elements. Unlike hydrogen, it is found in the free or uncombined state in large quan- tities on the earth, for the atmosphere consists essentially of free oxygen, mixed with about four times its volume of nitrogen. We shall consider the methods of obtaining this element from four of its compounds, namely, from (i) Mercuric oxide ; (2) Potassium Chlorate; (3) Sodium peroxide ; (4) Water; and also the way by which it is obtained from the air. (i) Oxygen from Mercuric Oxide. We saw in Exp. 6 that when this compound is simply heated, it is decom- posed into its two constituent elements, oxygen and mercury. In order to collect the gas which is given off we proceed as follows Experiment 58. Heat a small quantity of the red powder in a hard glass tube, arranged as shown in Fig. 15, A, and collect the gas over water. Notice that the evolution of gas is not very rapid ; also that, as in Exp. 6, the metal mercury collects on the cooler part of the tube in small globules. The chief interest in this experiment lies in the fact that Oxygen. 55 it was the very method by which Priestley first discovered oxygen in 1774. He called it dephlogistigated air; the name oxygen was given to it later by Lavoisier. (2) Oxygen from Potassium Chlorate. When this salt is heated, it first melts and then rapidly gives off oxygen. It is decomposed by heat into two things, namely, into oxygen, which passes off as gas, and into potassium chloride, which remains behind as a white solid. Experiment 59. Heat a small quantity of potassium chlorate in a similar apparatus to that used for the last experiment. Notice FIG. 47. that the crystals first crackle, then melt, and that the melted compound then begins to effervesce, owing to the escape of the oxygen gas. Note also that the gas is given off much more rapidly than in the case of the mercuric oxide. It has been found that if the potassium chlorate be previously mixed with manganese dioxide, the chlorate gives up its oxygen much more rapidly, and at a much lower temperature. Experiment 60. Mix about 20 grams of potassium chlorate with about a quarter of its weight of powdered manganese dioxide, and gently heat the mixture in a flask (a common Florence oil flask) as shown in Fig. 47. Notice first that the mixture does not 56 Oxygen. melt ; also that moisture collects in the neck of the flask. This chiefly comes from the manganese dioxide, which is always damp ; and it is in order to prevent this condensed moisture from running back into the heated flask and cracking it, that the apparatus is supported in a horizontal position. Observe how rapidly the gas is evolved, and with what a little heat. It is to allow plenty of passage for the gas that the usual glass delivery tube is here replaced by a wide piece of indiarubber tube. Notice that during the experiment little sparkles occasionally appear in the heated mixture. These are caused by small particles of combustible impurities which are always liable to be present in manganese dioxide. 1 Collect several cylinders or jars with the gas as it is evolved, and keep them standing mouth downward in small plates containing a little water. This is the method usually employed in the laboratory for preparing oxygen. The material left in the flask after the experiment consists of potassium chloride and manganese dioxide. The latter substance therefore comes out of the reaction in exactly the same slate as it was at the beginning. If the mixture is boiled with water, the potassium chloride dissolves and leaves the manganese dioxide ; so that if the mixture is then filtered, the dioxide can be recovered, and used over and over any number of times. The way in which it acts in causing the potassium chlorate to give up its oxygen more readily, involves a series of rather complex changes, which cannot be conveniently considered at this stage. (3) Oxygen from Sodium Peroxide. When this com- pound is brought into contact with water it is at once decom- posed, and oxygen is evolved. Experiment 61. Place a small quantity of sodium peroxide in a dry flask fitted with a delivery tube and a stoppered funnel, as in Fig. 48. Allow water to enter the flask drop by drop by means of the funnel. As each drop falls upon the powder, a brisk action is 1 If a large quantity of such impurity were present, such as would occur if the black oxide of manganese were adulterated with powdered coal, it would give rise to an explosion when heated. Oxygen from Air. 57 noticed, and oxygen is rapidly given off, which may be collected in the usual way. After the experiment, dip a piece of turmeric paper into the liquid in the flask ; notice that the paper is stained. Also dip the fingers into it, and note its caustic nature. Recall the solution obtained oy the action of sodium upon water (see Exp. 43). Here we have the same substance, caustic soda, formed as a second product of the chemi- cal change. FIG. 48. (4) Oxygen from Water. Al- though water is the commonest com- pound of oxygen, we very seldom em- ploy it for obtaining this gas, because water is not so easily decomposed as many other oxygen compounds. There are very few elements which have a sufficiently strong affinity for the hydrogen of the water, for them to tear it away from the oxygen, and let the latter go free. Under certain conditions, however, the element chlorine (also a gas) is able to do this. Thus, if a mixture of chlorine gas and steam be strongly heated by being passed through a red hot tube, the chlorine seizes the hydrogen of the water, unites with it to produce the compound hydrogen chloride (or hydrochloric acid), and the oxygen in the water is set free. (5) Oxygen from the Air. Oxygen is now obtained on a manufacturing scale from the enormous store of it which is present in the air. This is done in two operations; in the first the atmospheric oxygen is made to combine with some substance, and in the second this substance is again decom- posed. The substance employed is barium oxide (baryta). This is heated in iron pipes through which air is pumped under a slightly increased pressure. Under these circumstances the barium oxide combines with oxygen from the air, and gives barium dioxide, while the other chief constituent of the air, namely, the nitrogen, passes away. Presently the pumps are reversed, and a partial vacuum is produced in the heated pipes ; this causes the barium dioxide to decompose, changing back into 5 8 Oxygen. the original barium oxide, and giving up the oxygen it had absorbed from the air. Therefore, by pumping air in and out of these heated pipes containing baryta, it is possible to obtain large quantities of oxygen very rapidly. This method of obtaining oxygen is known as "Erin's process." The Properties of Oxygen. From the various speci- mens of the gas which have been prepared, we see that it is colourless, and that it is so little soluble in water that we observe no loss while collecting it in the pneumatic trough. If we take one of the jars and apply the nose to the gas, we shall find that oxygen has no smell or taste. This is exactly what we might expect, when we remember that the air we breathe contains a large proportion (one-fifth) of free oxygen. Test for Oxygen. Oxygen is generally recognized and distinguished from other gases by thrusting into a jar of it a chip or splinter of wood which has been lighted, and has only a glowing spark upon it. The splinter is instantly rekindled and bursts into flame. There is, however, one other gas (namely, nitrous oxide) which behaves in a similar manner towards a glowing splint. Therefore this test does not dis- tinguish oxygen from nitrous oxide. How these two gases are identified is explained on p. 202. Experiment 62. Test the samples of gas obtained by heating mercuric oxide, and from sodium peroxide, by plunging into them a splinter of wood which has a glowing spark upon the end. The splinter will almost immediately rekindle. Blow it out again, and, while the end is still glowing, thrust it once more into the gas. It again bursts into flame as before. This can be repeated so long as sufficient oxygen is left in the jar. All substances which are capable of burning in the air, will burn more rapidly and with increased brilliancy in pure oxygen. It is entirely on account of the oxygen present in the atmo- sphere, that substances burn in the air at all, and it will be evident that they must burn more quickly in oxygen alone, than in oxygen which is diluted with a large quantity of nitrogen, as it is in the atmosphere. We may prove this by burning a number of substances in the gas already prepared. Properties of Oxygen. 59 Experiment 63. Charcoal in Oxygen. Take a piece of char- coal about the size of a hazel-nut, and fasten it, by means of a piece of thin copper-wire, to a deflagrating spoon. Light one corner of the charcoal in a gas-flame, and then lower the spoon into a jar of oxygen. Notice that the charcoal at once begins to burn much more brightly than it did before. If the charcoal used, happens to be a piece made from the bark of the wood, it will throw off a shower of brilliant sparks or scintillations as it burns in the oxygen. When the charcoal has burnt out, remove the spoon, and pour a little water into the jar. Cover it with a glass plate and keep it for a further experiment. Experiments^. Sulphur in Oxygen. Unscrew the small cup from a deflagrating spoon, and tie a small bundle of asbestos to the end of the wire by means of thin copper wire. Then melt a little sulphur in a test-tube and dip the asbestos into it so as to get it coated over with the sulphur. Now light the sulphur upon the asbestos by means of a lamp flame, allow it to burn in the air for a moment or two and plunge it into a jar of oxygen. Observe the greatly increased brilliancy of the flame the moment it comes into the oxygen. Notice that the burning sulphur produces a little smoke or fume in the jar. When the sulphur has burnt out, pour a little water into the jar, shake it up, cover the jar over with a glass plate, and keep it for a subsequent experiment. Experiment 65. Phosphorus in Oxygen. Take a piece of phosphorus l about the size of a pea, wipe it dry with blotting paper, and place it in a deflagrating spoon. Set fire to the phosphorus by touching it with the end of a wire which has been just warmed in a flame, and lower the spoon into a large flask filled with oxygen. Ob- serve how much more intense is the light of the phosphorus burning in oxygen than burning in the air. Watch the experiment closely, and notice that, as the burning goes on, suddenly the light becomes almost un- bearably dazzling, and then quickly dies down. At that point the phosphorus was rapidly boiling. In preparing to ao tnis experiment, the wire of the deflagrating spoon should be pushed so far through the metal cap that the cup containing the phosphorus reaches nearly ; Remember the precautions as to handling phosphorus given on p. 10. 60 Oxygen. to the bottom of the flask as shown in Fig. 49, so as to prevent the flame from cracking the flask at its shoulder. When the phosphorus is burnt out, remove the spoon, add a little water, and shake it up in the flask. Notice that the dense white fumes in the flask rapidly disappear ; they are dissolved by the water. Experiment 66. Pour a little blue litmus solution into the jars used in Exps. 63, 64, 65. Notice that the blue solution is turned red in each case ; but observe that in the jar in which the carbon was burnt the red colour is less of a scarlet, and more inclined to purple than in the other two cases. Substances which have the power of turning litmus red, are said to be acid ; therefore, when we burn carbon, sulphur, and phosphorus in oxygen, and add water to the products of the burning, we obtain substances which are acids. Experiment 67. Sodium in Oxygen. Put a small fragment of sodium in a clean dry deflagrating spoon, and heat it in a gas flame until the sodium begins to burn in the air, then plunge it into a cylinder of oxygen. The sodium burns very brightly, and produces a white smoke. The product of the combustion is, how- ever, in this instance a solid, and most of it, therefore, remains behind in the spoon. When the sodium has all burnt, remove the spoon, and rinse out the jar with a little water, and pour the liquid cut into two small beakers. To one add a little litmus, and notice that it is not turned red. To the other add a little litmus which has been just turned red by a single drop of dilute acid, and observe that this reddened litmus is turned back to its original blue colour. 1 Substances which restore the blue colour to reddened litmus are said to be alkaline; therefore, when sodium is burnt in oxygen and the product dissolved in water, we obtain an alkali. Many substances which will not burn in the air, will burn readily in pure oxygen. Experiment 68. Iron in Oxygen. In order to burn iron in oxygen, it must, like the phosphorus or the sulphur or the charcoal, be first lighted. To do the experiment it is best to use a jet of 1 If the liquid obtained by rinsing out the jar does not contain enough of the product from the burning sodium to show this result, the main portion which was left in the spoon may be used, by dissolving it off in a little water. Burning Iron in Oxygen. 61 oxygen direct from a store of the gas contained either in a gasholder or better in a metal cylinder, into which the oxygen has been pumped under great pressure. Connect to a gas reservoir by means of a rubber tube, a glass tube, drawn to a jet at one end, and allow a stream of oxygen to blow a spirit lamp flame against the ends of a bundle of fine steel wires, in the manner shown in Fig. 50. Almost immediately the tips of the wires get red hot, and begin to burn. Now remove the lamp and the metal will continue to burn in the jet of oxygen, throwing off a shower of brilliant sparks. Notice that the end of the bundle of wires melts, and drops of molten matter continually fall. These should be allowed to drop into an iron basin of water placed ready to catch them. In the absence of a reservoir of oxygen, fill a common wide- mouthed bottle (such as a pickle bottle or glass marmalade jar) with FIG. 50. oxygen, leaving about two inches of water in the bottle. Stick into a loosely fitting cork a bundle of steel wire or a straightened piece of watchspring, the end of which has been tipped with sulphur. Light the sulphur in a flame and quickly plunge the wire into the bottle of oxygen. The burning sulphur will set fire to the iron, which will go on burning in the oxygen. The melted product of the combustion will be partially quenched by the layer of water ; but, in spite of this, it will probably crack the bottle. Experiment 69. Collect some of the solid product of the burning iron, and place it upon a moistened litmus paper. Notice that the paper is not turned red. Redden another piece of litmus paper by dipping it into water containing a drop or two of acid, and place some of the substance upon this. The blue colour is not restored ; therefore the product obtained by burning iron in oxygen is neither acid nor alkaline. In each case the product obtained when a substance is 62 Oxygen. burnt in oxygen, is a compound of the thing burnt with the oxygen. The process of burning in these instances is there- fore nothing more than the rapid combination of the various substances with oxygen. If we were to burn carbon, sulphur, phosphorus, or sodium in the air, and examine the products obtained, we should find that in all cases they were the same as when these things were burnt in oxygen. This shows that the ordinary process of combustion in the air is also the rapid combination of substances with oxygen, the difference being that the combination is not so rapid as in pure oxygen. We can easily prove that it is the oxygen present in the air which enables substances to burn in air, by removing the oxygen from a quantity of air, and trying the experiment of putting a lighted candle into the air that is left. Experiment 70. Place a fragment of phosphorus (wiped dry) in a little dish, and float the dish, upon water in a pneumatic trough. Set fire to the phosphorus, and cover it with a cylinder a little wider than the small dish as in Fig. 51. The phosphorus burns in the confined air in the cylinder, and combines with the oxygen present, forming the same white fume as when burnt in pure oxygen. When the phosphorus has burnt itself out, allow the apparatus to stand for a short time, so that the jar may cool, and the fumes may become dissolved by the water. Notice that the water has risen in the cylinder, taking the place of the oxygen that has been withdrawn. Now slip a glass plate beneath the mouth of the cylinder and remove it from the trough. Shake the air and water together so as to completely dissolve the remaining fume. Now introduce a lighted taper or candle into the gas in the jar, and notice that the flame is at once extinguished. A lighted spill of paper, or a coal-gas flame, will behave in a similar manner. We have withdrawn all the oxygen from the air in the cylinder, and the gas that remains (namely, nitrogen) will not support the com- bustion of ordinary burning bodies. Oxides. The products obtained when substances are FIG. 51. Oxidation. 63 burnt in oxygen (either pure oxygen, or the oxygen of the air) are called oxides. All the elements, except fluorine, have been made to combine with oxygen either directly or in- directly. We have seen from Exps. 66, 67, and 68, that different oxides have very different properties ; thus, some combine with water to yield acids, 1 while others, under the same cir- cumstances, give compounds which are alkaline. The former of these are called acid-forming oxides, while the latter belong to a class known as basic oxides. (As a general rule the non- metals combine with oxygen to give acid-forming oxides, while the metals yield basic oxides. There are, however, some oxides of both non-metals and of metals which are neither acid-forming nor basic.) Hydroxides. This is the name applied to the com- pounds which are produced when oxides combine with water. The acid-forming oxides yield acid hydroxides, while those derived from basic oxides are called basic hydroxides. The term acid hydroxide, however, is not very often employed, as these compounds are included in the class of substances called acids ; and the name hydroxide alone is more usually used to denote the basic hydroxides. Since the basic hydroxides are derived from oxides of metals, they consist of a metal combined with oxygen and hydrogen : they are the hydroxides of metals. Thus, sodium combined with oxygen and hydrogen (or sodium oxide com- bined with water) is sodium hydroxide ; calcium united with oxygen and hydrogen (or calcium oxide united with water) is calcium hydroxide ; and so on. Oxidation. The process of combining with oxygen is called oxidation, whether the action takes place rapidly, as when substances burn in oxygen, or whether it goes on slowly without any visible signs of heat. We saw in Ex p. 61, the rapid oxidation of the metal sodium, but the same process of oxidation goes on if sodium is merely exposed to oxygen, even the oxygen of the air, without being made to burn. 1 For further development of this subject, see p. 66. 64 Oxygen. Experiment 71. Take a fair sized piece of sodium and quickly cut a slice off it. Notice that for an instant the freshly cut surface of the metal looks bright and silvery, but that almost immediately it becomes dull and tarnished. It quickly gets coated with a white film or crust which consists of the oxide of sodium. The tarnishing is the oxidation of the metal, and if the piece of sodium be left exposed to the air it is soon oxidized right through, the whole piece being changed into the oxide. Other metals behave in a similar way ; thus, when bright iron is exposed to the air, we know that it soon loses its brilliant surface, and becomes coated with a reddish film of rust. This rust is simply the oxide of iron, formed by the slow combination of the metal with the oxygen of the air. Increase of Weight by Burning. If the ordinary process of burning is simply the rapid combination of the burning substance with the oxygen of the air, we ought to find that the products of burning actually weigh more than the material that is burnt. This we can easily put to the test of experiment. Experiment 72. Place a heap of " reduced iron " l on a little iron dish or tray, and balance it upon a small pair of scales. Then, by means of a lighted taper, set the heap alight. Observe that the black powder gradually smoulders right through, and turns to the familiar brownish colour of iron rust ; but note also that as it burns the mass gains in weight, for the scale pan on which it is soon begins to fall. This shows that the rust is heavier than the iron which produced it. In this experiment the only product of the burning is solid and visible, but when we burn a candle we see no products. The candle simply appears to waste away, leaving nothing behind ; by burning it we seem to have completely annihilated it. Is this really the case ? or can it be that the products of the burning of a candle are gases, and escape unobserved into the air ? Let us try the following experiment. 1 Reduced iron is readily prepared by heating oxide of iron in a glass tube, and passing a stream of coal-gas or hydrogen through the tube until the material is black. Respiration. Experiment 73 . Take a piece of wide glass tube (such as a lamp chimney) and fill the upper part with lumps of solid caustic soda, which are kept in position by first hanging a false bottom of wire gauze into the tube by means of wires. Fit a cork into the bottom, on which is fast- ened a short candle, the cork being bored with holes to allow air to enter (Fig. 52). Balance this arrangement upon a pair of scales, and then light the candle and quickly replace .the cork. Again note that as the candle burns the apparatus becomes heavier. The candle in the act of burn- ing is combining with oxygen, and two invisible gaseous pro- ducts are formed, namely, steam and carbon dioxide, both of which are caught or absorbed by the lumps of caustic soda. FIG. 52. Matter is Indestruc- tible. We can destroy a candle by burning it, but not the carbon and hydrogen of which the candle is composed ; these live on, as it were, but in different states of combination. Similarly if we burn a piece of sulphur it disappears from view, but the sulphur is not destroyed, only passed into combination with oxygen, forming a compound which is gaseous and invisible. Respiration. Not only is oxygen necessary to support the combustion of ordinary burning bodies, but it is also indispensable to the process of respiration. If an animal is placed in air from which the oxygen has been removed, or in any mixture of gases which does not contain free oxygen, it quickly dies. All animals require oxygen to breathe. Respira- tion is, in fact, a process of oxidation ; air is drawn into the lungs, and a portion of the oxygen is absorbed by the blood. This oxygen-laden blood (which has a bright red colour) passes throughout the body, and exerts its oxidizing power F 66 Oxygen. upon certain compounds containing carbon which have to be removed from the system, with the result that carbon dioxide is produced (the same compound as is formed when carbon is burnt in oxygen, or when a candle burns in the air) and is exhaled in the breath. When the red blood has thus parted with its oxygen it becomes of a dark colour, and is known as venous blood; this, travelling back to the lungs, is there once more charged with oxygen, again to carry it to all parts of the body. We can, by a simple experiment, show that the air exhaled from the lungs contains the same gas as is formed when a gas flame or a candle burns in the air. Experiment 74. Place some clear lime-water in a test-tube, and by means of a glass tube dipping into it, bubble the breath from the lungs through it. Note that the first portions of breath produce little or no effect, but that presently, as air from the deeper parts of the lungs is expelled, the clear solution quickly becomes milky. And we can also show that the air so expelled has been deprived of some of its oxygen by the following experiment. Experiment 75. Collect in a cylinder over water the breath exhaled by one single long expiration, emptying the lungs as far as possible (see Exp. 31, Fig. 12). Now remove the cylinder and lower into it a lighted candle on a wire. Notice that the flame will be extinguished. [This result is due partly to the presence of the carbon dioxide, as well as to the smaller amount of oxygen.] Acids. There is a class of compounds which chemists call acids. All acids have a sour taste, and they also have the power of reddening a solution of blue litmus (refer back to Exp. 66). It does not follow, however, that all sour substances, or all that will redden litmus, are acids. This is by no means the case, for we know of many things which have a sour taste and which will turn litmus from blue to red, but which do not belong to the class of substances recognised as acids ; they are acid, in the sense of being sour, but are not acids. The question, What is an acid? is one which cannot be answered in a word, and chemists are not agreed as to what is the best definition to be given. At one time it was supposed Acids and Alkalies. 67 that all acids contained oxygen, that this element was the acidifying principle in these substances. Indeed the very name oxygen means the acid producer, and was originally applied to this element by Lavoisier because of this idea. We now know that this belief was wrong, for we have many acids which do not contain any oxygen in their composition. At the present time chemists regard the element hydrogen as a necessary constituent of an acid. According to modern notions, a com- pound which does not contain any hydrogen cannot be an acid. This of course does not mean that all compounds containing hydrogen are acids. We know, for example, that water is a compound of hydrogen, but water is not an acid, it is not sour, and does not redden litmus. We have already learnt that sulphuric acid and hydrochloric acid contain hydrogen, because we obtained this element from both of them by acting upon them with various metals. Can we, therefore, define an acid as a compound from which hydrogen can be expelled by the action of certain metals ? No, because we have also learnt that hydrogen can be expelled from water by means of certain metals, and water is not an acid. Perhaps the best definition that can be given is the following, an acid is a compound from which hydrogen can be displaced by the hydroxide of a metal. The word displaced here only means that the hydrogen is turned out of the compound, and not that it is set free, or liberated as gas. Alkalies. This name is applied to the hydroxides of the alkali metals (of which sodium and potassium are the most important) and to ammonia. When sodium is burnt in air or in oxygen, and the oxide so obtained is dissolved in water, sodium hydroxide is produced (Exp. 67). We have learnt that the alkali produced in this experiment has the property of changing the colour of turmeric from yellow to brown, and also of restoring the blue colour to litmus which has been reddened by an acid. Experiment 76. Make a dilute solution of sodium hydroxide (caustic soda) by dissolving a small piece of the solid in water in a beaker. Dip a turmeric paper into it, and observe the stain. Make a dilute sample of hydrochloric acid by adding a little strong 68 Oxygen. acid to some water in a beaker. Dip a turmeric paper into this ; note that no stain is produced. Add to the acid a few drops of a solution of litmus, which will be at once changed from blue to red. Now add the dilute alkali very gradually to the acid, and watch the effect upon the reddened litmus. Where the alkali meets the acid solution, the litmus shows a blue colour, which, however, dis- appears again on gently shaking or stirring. Presently, however, as more alkali is added, the red liquid turns entirely blue. The solution is now no longer acid, but alkaline. Now add, drop by drop, a little more of the dilute acid, stirring the solution ; stop the moment the liquid turns red. If this is carefully done, one drop of the alkali will now restore the blue colour. Now test this solution with a turmeric paper, and note that no brown stain is produced ; therefore the liquid is neither add nor alkaline. Neutralization. Exp. 76 teaches us that when we mix an acid and an alkali, each one destroys the other, so that a point can be arrived at when the solution has neither the property of an acid nor an alkali. Under these circum- stances we say that the liquid is neutral, that the acid has neutralized the alkali, or the alkali has neutralized the acid. Now we must ask ourselves what has become of the acid and the alkali ? Are both still present, and the properties of each simply masked by those of the other ? In other words, have we simply a mixture of these things so adjusted that the properties of the one just balance those of the other, or has any chemical change taken place between the acid and the alkali resulting in new compounds which happen to have no effect upon either turmeric or litmus ? Let us test this matter by experiment. Experiment 77. Take some of the dilute acid, the alkali, and the neutral mixture, and evaporate each separately to dryness in small dishes. Note first that the acid leaves no residue behind : that is, it is a volatile acid ; therefore, if the neutral solution is simply a mixture of the two substances, when it is evaporated down, the acid ought to volatilize along with the steam and leave the alkali, in which case the residue in the dish containing the alkali, and that in the one containing the neutral solution would be the same. Are they the same? Observe that they appear different. Touch each with a moist turmeric paper, the one is Salts. 69 strongly alkaline, while the other is still perfectly neutral, therefore they are different. The residue being neutral shows that by evaporating the solution, the acid, which had been added, has not been driven off ; it cannot, therefore, have been simply mixed with the alkali, but must have entered into chemical union with it. Now, with the little finger bring a little of the neutral residue on to the tip of the tongue ; the familiar taste of the substance will also prove that we have here a compound which is quite different from either sodium hydroxide or hydrochloric acid. We learn, therefore, that when acids and alkalies neutralize one another they enter into chemical combination with each other, forming new compounds. Bases. A large number of other substances besides the alkalies, are capable of neutralizing acids. Many of these substances resemble the alkalies in having the power of restoring the blue colour to reddened litmus, and in imparting a brown colour to turmeric; they are, therefore, said to be alkaline in character, or to possess an alkaline reaction. Some chemists, indeed, extend the name alkali so as to include many of these compounds. The term base has long been in use to denote any substance which is capable of neutralizing an acid, and it therefore includes the alkalies. For the most part bases are compounds of metals; being either the oxides or hydroxides of metals. Ammonia, however, which contains no metal, being only a compound of nitrogen and hydrogen, is also included in this class, and so are many organic com- pounds which are not compounds of metals. For our present purpose, however, we may define bases as certain oxides and hydroxides of metals, which are able to neutralize acids, and include ammonia. Salts. The compounds that are produced when acids and bases combine together are called salts. We have seen by Exp. 77 that when hydrochloric acid and sodium hydroxide neutralize each other the substance we familiarly call "salt" is produced. "Salt" is one of the commonest and most important of all salts, as well as being one that has been longest known to man ; and the compounds belonging to this class are called salts, because of a general similarity many 70 Oxygen. of them bear to "salt," or common salf, as it is termed. No satisfactory definition of a salt can be given, because chemists are not all agreed as to exactly what shall be included in this class of compounds. Some regard acids themselves, as being salts of hydrogen. At this stage we will consider salts as being the compounds that are produced by the interaction of acids with bases. EPITOME. Oxygen was discovered in 1774 by Priestley, by heating mercuric 1 oxide. It is the most abundant of all the elements. Occurs un- combined in the air. It is usually obtained by heating potassium chlorate : more readily if the chlorate be mixed with manganese dioxide. Oxygen is also given off when certain peroxides, such as manganese dioxide, or barium dioxide, are strongly heated ; the latter substance is used in Erin's process. Oxygen can be obtained from water by heating steam and chlorine together. Oxygen is a colourless, tasteless, and inodorous gas ; supports combustion energetically, and is necessary to life. Oxygen is slightly soluble in water ; 100 volumes of water can dissolve 4 volumes of oxygen. Fish depend upon this dissolved oxygen for their supply of this gas for respiration ; they cannot breathe free air, neither can they use the oxygen which is a chemical constituent of water. The products obtained by burning things in oxygen, or in air, are oxides. Some of these are acid-forming, others are basic oxides. Similar compounds are formed when the same elements combine slowly with oxygen, the process then being called oxidation. The products obtained by burning or by oxidation are heavier than the original substance that is burnt ; burning is only matter under- going change, and not matter being destroyed. Acids are compounds containing hydrogen which can be dis- placed from the compound by the hydroxide of a metal. They are sour or acid to the taste, and will redden blue litmus. Bases are compounds which will neutralize acids. They are mostly certain oxides and hydroxides of metals. Ammonia is also a base. Those that are soluble in water are alkaline ; they restore the blue colour to reddened litmus, and give a brown stain to turmeric. Salts are the compounds produced by the union of acids and bases. Oxygen. 7 1 Reactions for oxygen x (1) From mercuric oxide HgO = Hg + O (2) potassium chlorate KC1O 3 = KC1 + 30 (3) sodium peroxide and water Na 2 O 2 4- H 2 O = 2NaHO + (4) water by the action of chlorine H 2 O + C1 2 = 2HC1 + O (5) ,-, air, by Erin's process (0) BaO + O = BaO 2 (&) BaO 2 = BaO + O Combustions in oxygen Carbon C + O 2 = CO 2 Sulphur S + O 2 = SO 2 Phosphorus aP + $O = P 2 O 5 Iron 3Fe + 40 = Fe 3 O 4 1 These will be explained later on'; the student can pass them over at this stage. CHAPTER VIII. WATER. WE have learnt by Exp. 54 that when hydrogen burns in the air, water is formed ; and as other elements on burning in the air yield their oxides, so we might expect that hydrogen would do the same, and that water is an oxide of hydrogen. First let us make an experiment in order to collect the liquid which is formed, and examine it. Experiment 78. Burn a small flame of hydrogen under the open end of a bent glass tube, arranged as in Fig. 53. A liquid quickly condenses on the long neck and runs down into the flask, and in about half an hour a considerable quantity will be collected. One property by which we can recognize water from other liquids we have learnt from Exp. 44, and that is its behaviour towards the metal potassium. Experiment 79. Pour the liquid collected in Exp. 78 into a test-tube, and drop into it a small bit of potassium ; if the metal takes fire, as it did in Exp. 44, we conclude that the liquid is water. [Other tests by which we can recognize water we shall learn later on.] As the air does not consist of oxygen alone, we cannot say certainly from this experiment that water is a compound of hydrogen and oxygen only. This point might be settled by burning the hydrogen in pure oxygen. It was indeed first settled for us by Cavendish, in 1781. Cavendish used a strong glass vessel, similar to E, Fig. 54, which had a stopcock at the bottom, and was closed at the top with a stopper through which two wires were fixed, so that Water. 73 he could send an electric spark into the gases it contained. He pumped all the air out of this vessel with an air pump, and then attached it to the bell jar, B, containing a mixture of hydrogen and oxygen. On opening the taps the gases of course entered into the vacuous tube. The stopcocks were then closed, and a spark from an electrical machine was made to pass between the two wires. This of course exploded the r FIG. 53. FIG. 54. mixture of hydrogen and oxygen, but the vessel being very strong did not burst. After the two gases had thus combined, there was a minute quantity of water in the tube. By again filling the vessel with more of the gas and ex- ploding the mixture, the quantity of water increased, until by repeating the experiment several times, enough of the liquid was collected to prove that it was actually pure water. As there was nothing besides hydrogen and oxygen present in the gases, this proved that water was composed simply of these two elements. But Cavendish did more than merely prove this, he also 74 Water. taught us the proportions in which these two gases combined together to produce water. If, for instance, the mixture of gases in the bell jar contained equal measures of hydrogen and oxygen, he found that after the explosion there was always some oxygen left over in the vessel E. But if he mixed the gases exactly in the proportion of two measures of hydrogen to one of oxygen, then there was no gas left in E after the ex- plosion ; it entirely disappeared, leaving a vacuum in the vessel. Therefore Cavendish made the discovery that water was a compound of hydrogen and oxygen only, and that these two elements combined in the proportion of tivo volumes of hydrogen to om volume of oxygen. We can make hydrogen and oxygen combine together by what are called indirect methods, and thus get additional proof that water is composed of these two elements only. FIG. 55. Experiment 80. Roll a piece of fine copper gauze into short compact cylinders, c, Fig. 55, and place them in a piece of hard glass tube (combustion tube}. To one end of this tube connect a small Wurtz flask, in the manner shown at w, Fig. 55, and to the other attach a small apparatus for generating oxygen from potas- sium chlorate. Heat the combustion tube with a long flat flame Formation by Synthesis. 75 until the copper is red hot, and then send a slow stream of oxygen over it by gently heating the potassium chlorate. Notice that as the oxygen passes over the copper, the latter becomes black. It is gradually combining with the oxygen and is being converted into copper oxide, which is black. After some little time disconnect the oxygen generator, and replace it by an apparatus for generating hydrogen from zinc and sulphuric acid, and pass a slow stream of hydrogen through the tube. 1 Observe that almost at once some moisture condenses in the Wurtz flask, and as the experiment goes on, more and more water collects in the little receiver. Where is this water coming from ? Notice that the black copper oxide is gradually changing back again into bright metallic copper. The hydrogen has taken away the oxygen which the copper had combined with, and has united with it to form water. In this experiment therefore we have caused hydrogen and oxygen to unite, by first combining the oxygen with copper, and then allowing hydrogen to deprive the copper oxide of the oxygen, whereby the copper oxide was again reduced (that is, the copper was deprived of the element with which it had combined, and was restored to its former state of metallic copper), and water was formed. This experiment is a very important one, and we shall return to it again later. In Cavendish's experiment, and in Exp. 78, water was obtained by the direct combination of its constituent elements; by Exp. 80, it was formed by the indirect union of oxygen and hydrogen ; but whether directly or indirectly, the compound was built up, so to speak, from the elements of which it is composed. This method of proceeding is called synthesis. There is another way by which we can find out the com- position of a compound like water, which is exactly the opposite of synthesis. Instead of taking the constituents of the com- pound and making them combine, we can take the compound and decompose it into its constituents. This method is known as analysis. 1 A convenient piece of apparatus, easy to make, by means of which a regulated stream of hydrogen can be obtained at will, is described on p. 246, Fig. 106. Put granulated zinc in one tube and dilute sulphuric acid in the other. Water. For instance, in Exps. 45, 48, we decomposed water by sodium, and by iron, and got out of the water one of its constituents, namely, the hydrogen. By the action of chlorine upon water (as described on p. 57), the water was again decomposed, and the other con- stituent, namely, the oxygen, was obtained. These are processes of analysis. Water can also be decomposed by means of an electric current from a battery. Experiment 8 1. Take an ordinary wide mouthed bottle, and cut it in half. (This may be done by first making a small scratch on the glass with a file, and then touching the spot with a red hot wire. The glass will then crack at the file mark, and the crack can be made to travel right round the bottle by slowly drawing the hot wire along just in front of the crack.) Fit into the neck a cork through which two short pieces of platinum wire have been pushed. The apparatus is then supported as shown in Fig. 56, and nearly filled with water, to which a little sulphuric acid has been added. 1 Now attach the two wires from the battery to the two platinum wires, and notice that bubbles of gas at once make their appearance on each of the wires in the acid water. Fill a short stout test-tube with the dilute acid and invert it over the two wires in the basin, and collect the gas that is evolved. When the tube is full, remove it and apply a lighted taper to the gas. If it is hydrogen it will be recognized by the characteristic flame ; if oxygen, we shall see the taper flame burn more brightly. Note, however, that the gas does not answer to either of these tests, but that it explodes with a sharp crack. This shows that we have both oxygen and hydrogen mixed together. 1 Pure water will not conduct electricity, hence sulphuric acid is added to it. FIG. 56. The Volume Composition of Water. 77 Experiment 82. Now collect the gas from each wire separately by inverting a separate tube over each wire, using two test-tubes of the same size. Notice that gas collects more quickly in one tube than in the other. As soon as the tube which fills quickest is full down to the water level, stop the experiment by disconnecting one wire from the battery. Note that the other tube is only half full. Now test the gas in each tube, and find that the gas in smallest quantity will rekindle a glowing splint of wood ; it is therefore oxygen ; while the other burns with the characteristic flame of hydrogen. These two experiments prove by analysis that water consists of oxygen and hydrogen ; and also that it contains these two gases in the proportion of two volumes of hydrogen to one volume of oxygen. They therefore confirm Cavendish's synthetical experiments. When oxygen and hydrogen unite (as in Cavendish's ex- periment), the volume occupied by the water that is formed is so extremely minute, that it is scarcely measurable unless very large volumes of the two gases are used. We cannot, therefore, make any comparison between the volume of the gas and that of the liquid water that is produced. But if, instead of letting the water condense to the liquid state, we were to make the experiment at a high temperature, so that the water was kept in the condition of steam, then we could find out the relation between the volume of the mixed gases, oxygen and hydrogen, and the volume of the gaseous compound, steam. This experiment is made in the following way. Two measures of hydrogen and one of oxygen are put in the closed limb of the U tube (Fig. 57), which is divided into three equal measures by rings on the glass. The tube is then heated by boiling some amyl alcohol in a flask and sending the hot vapour through the outer glass tube which surrounds the tube containing the gases. The vapour leaves this "jacket tube" by the small pipe at the bottom, and is again condensed and collected. When the whole apparatus is hot, the mercury is made level in both limbs, and the gas exactly occupies the three measures. Now an electric spark from an electrical machine is passed between two platinum wires which are 78 Water. sealed into the closed end of the tube containing the gases This causes the oxygen and hydrogen to combine (which the} do with some violence, so that precautions are taken to prevent the mercury from being blown out at the open end of the U tube) and form water, which, however, cannot condense to the liquid state, but remains as steam, on account of the high temperature of the vapour surrounding the tube. The first thing noticed after the explosion of the gases, is that the volume is altered, for the mercury has risen in the tube. FIG. 57. Once more, therefore, the mercury must be made level in both limbs, by pouring more into the open side ; and when this is done, it is seen that the gas in the tube now occupies exactly two of the measures. That is to say, two volumes of hydrogen and one of oxygen, three volumes of the mixed gases, form two volumes of steam or gaseous water. The Properties of Water. Water is familiar to us as a colourless and tasteless liquid. When we look through a considerable depth of it, however, it is seen to be possessed of a bluish-green colour. We know that when water is cooled Point of Maximum Density of Water. 79 to a certain temperature it solidifies or freezes ; and when heated to a particular temperature it boils. One method, therefore, by which we are able to distinguish water from any other colourless liquid is to ascertain its freezing and boiling points. The freezing point of water is o on the centigrade scale, while its boiling point is 100 (see p. 88). As water is gradually cooled down, like other substances it shrinks in volume. This shrinking or contraction goes on steadily until the temperature is just within 4 of the freezing point. After this temperature is passed, the water, instead of continuing to contract, actually expands again, just as though it were being warmed. Therefore, if we take some water having a temperature of 4 C., it will expand whether we heat it or cool it. At this particular temperature, water is more dense than at any other point ; 4 C. is, therefore, called its point of maximum density. The expansion that water suffers on being cooled from 4 to o (without freezing) is extremely slight; 1000 cc. of water measured at 4, when cooled to o, will become ex- panded to 1000-13 cc - B ut tn i s expansion, although appa- rently so trifling, plays an important part in nature. When lakes and ponds are exposed to the cold winter winds, the surface water becomes cooled, and consequently contracts and becomes denser. It therefore sinks to the bottom, and fresh layers come to the surface to be cooled in their turn, and also to sink. In this way a circulating movement is set up, which continues until the temperature of the whole mass of water has fallen to 4 C. When this state is reached, any further cooling of the surface expands the water on the top. It there- fore becomes less dense, and so remains as a colder layer floating upon the surface, until at last it solidifies to a thin film of ice, which then protects the water beneath from contact with the cold wind. If water continued steadily to contract with cold until the freezing point, this circulating movement would not stop at 4 C., but would go on until the whole body of water in lakes and ponds had reached o, when ice would begin to form at the bottom as well as the top, and the entire mass of water would verv soon become solidified throughout. 8o Water. When water actually freezes, it suddenly expands very con- siderably, 10 volumes of water becoming n (nearly) volumes of ice. Ice is therefore lighter than water. 10 cc. of water weigh 10 grams, but n cc. of ice only weigh 10 grams. Con- sequently ice floats on water. The force which is exerted by the expansion when water freezes is very great. Thus, if a strong iron bottle be filled entirely with water, and the mouth securely closed, and then the bottle be exposed to a freezing temperature, the expansion will cause the bottle to burst. It is owing to this that water- pipes burst in frosty weather. The pipe splits, or bursts, at the moment the water freezes, but it is only when the ice melts again that the damage to the pipe is revealed ; hence the mistake is often made of supposing that the thaw caused the pipe to burst. We say that water boils at a temperature of 100 C. ; but, to be exact, we must add under a pressure 0/760 mm. (see p. 92). The actual temperature at which water, or any ether liquid, boils, depends entirely upon the pressure. This is very easy to prove. If we take a quantity of water in a beaker, scarcely hotter than can be borne by the hand, therefore far short of boiling, and place it under the receiver of an air- pump, and reduce the pressure by quickly pumping out some of the air, we shall see the water begins to boil violently, although obviously it cannot be getting any hotter. In a vacuum water can be made to boil even at the freezing tem- perature, therefore boiling has not necessarily anything to do with being hot. In ordinary life we are in the habit of associating the two ideas together, so that the very word boiling is almost synonymous with being hot; hence the common expression "boiling hot." In reality the expression "boiling cold" is quite as correct, for many boiling liquids are colder than ice, and colder even than the coldest arctic regions. The Solvent Power of Water. Water is able to dis- solve a great many substances, and chemists make use of this property in many different ways. Thus substances, which in the solid state are incapable of acting chemically upon each other, will often combine if either one or both of them are The Hardness of Water. 8 1 first dissolved in water, hence a large number of chemical operations (especially analytical) are conducted with solutions in water. The solvent power of water is also made use of for separating substances which are more soluble, from others which are either slightly soluble or are altogether insoluble (see Exps. 20 and 27). Many gases are soluble in water; some to a very great extent, others only very slightly. Thus, i cc. of water at o will dissolve 1148 cc. of ammonia gas, but only 0*048 cc. of oxygen. In all cases, gases are less soluble in hot than in cold water ; thus at 20 i cc. of water can only dissolve 680 cc. of ammonia, and 0*028 cc. of oxygen. Therefore, if a solution of a gas in water be boiled, all the gas is expelled. Natural Waters. Owing to its great solvent powers, perfectly pure water is never found in nature, and, indeed, can only be obtained in the laboratory by taking great precautions. The purest form of natural water is rain water, but even this dissolves the gases from the air, and is also contaminated with impurities that are present in the air. The moment rain touches the ground, it begins to dissolve the soil or rocks upon which it falls and through which it percolates, and it gradually becomes more and more impure as it finds its way to river and ocean. Absolutely pure water when evaporated to dryness in a platinum vessel leaves no residue. If boiled down in a glass vessel it slightly dissolves the glass. Sea-water contains more dissolved impurities in the form of various salts than other natural waters. Thus, in 1000 grams (i litre) of the water of the British Channel, there are 35*25 grams of dissolved salts, which remain as a residue when this volume of the water is evaporated to dryness. The greater part of this residue, viz. 27 grams, is common salt. The solid dissolved sub- stances in good drinking water average from 0*437 to 0*03 grams in i litre. Hardness of Water. Certain of the dissolved solids in natural waters impart the property known as hardness. These salts are chiefly the carbonate and sulphate of lime. When hard water is boiled, if its hardness is due to the presence of G 82 Water. dissolved carbonate of lime, it becomes soft, because the carbonate of lime is precipitated as a solid, which in time collects on the sides of the vessel (boiler or kettle), and pro- duces the furring which is so common. Sulphate of lime is not thrown out of solution by boiling ; therefore, if the hardness is due to this salt, boiling the water will not soften it. On this account, hardness caused by carbonate of lime is called temporary hardness ; and that due to sulphate of lime is termed permanent hardness. If a sample of water contains both these compounds, then it loses only a part of its hard- ness, viz. the temporary hardness, on boiling, while the permanent hardness remains. Water of Crystallization. When salts are crystallized from solution in water (as in Exp. 26), it often happens that some of the water solidifies along with the salt. It is as though the particles of the salt were unable to build them- selves up into the regular shapes, we call crystals, without the aid of some particles of water to hold them together ; just as bricks cannot be built up into an edifice without the aid of mortar. If the mortar were to be removed from a building, the bricks would all tumble down in a heap ; and if we remove the solidified water from such a crystal, it will fall down to a powder. Water which is thus held by a crystal, and which is necessary to its existence as a crystal, is called water of crystallization. Experiment 83. Take a few crystals of copper sulphate (blue vitriol} and gently heat them in a dish, either over a small flame or better in an oven. Notice that gradually the crystals lose their blue colour, and become white, while at the same time they lose their crystalline nature and are changed to a powder. Healing the crystals has driven off their water of crystallization, and conse- quently they no longer retain their shape as crystals. In the case of the copper sulphate it is easy to see when the water of crystallization is removed, because the salt itself (the anhydrous salt, as we call it) is nearly white, while the crystallized salt is deep blue. If a quantity of the white de- hydrated copper sulphate be moistened with water, it is at once re-hydrated, that is, it takes up water of crystallization Water of Crystallization. 83 again and consequently turns blue. Many other salts change colour when they lose either some or all of their water of crystallization. Experiment 84. Dissolve a crystal of cobalt chloride in water. Notice the pink colour of the salt and also of the solution. With a brush, or a clean pen, write on paper, using this solution instead of ink. The writing will be invisible, because the pink colour is so faint. Now warm the paper in front of a fire, or over a gas flame, and observe that the writing begins to show, and appears blue. The pink salt has now lost water of crystallization, and turns blue. If the paper be left exposed to the moisture in the air, or more quickly if breathed upon, the salt will re-hydrate itself and turn pink again. Some salts lose their water of crystallization on mere exposure to the air. Common washing soda is an example. If crystals of this are left on the table they soon lose their clear appearance and begin to fall to powder. This process is called efflorescence. Other salts do just the opposite. They absorb moisture from the air ; sometimes enough to cause the salt to liquify. Such salts are said to deliquesce. Substances which have this power are very useful to chemists for removing moisture from gases which are required to be dry. HYDROGEN PEROXIDE. Besides water, hydrogen and oxygen form another compound called hydrogen peroxide. It is obtained by acting on either sodium peroxide or barium peroxide with dilute sulphuric acid. It is a very unstable substance, being easily converted into oxygen and water. The readiness with which it parts with oxygen enables it to oxidize other substances. Thus, if it be added to lead sulphide (a black compound), it converts it into lead sulphate (a white substance). It is used on this account to clean old oil paintings which have become black by the white lead in the paint being changed to lead sulphide. Hydrogen peroxide also has bleaching properties, and is sometimes employed to bleach the hair. EPITOME. Water is a compound of oxygen and hydrogen. This is proved synthetically (i) by burning hydrogen in oxygen, and collecting 84 Wafer. and testing the liquid product that is formed. (2) By exploding oxygen and hydrogen in a closed vessel (Cavendish's method). (3) By reducing certain metallic oxides, such as copper oxide, in a stream of hydrogen (Duma's method). It may be proved analytically (i) by decomposing water by certain metals, as sodium, potassium, magnesium, or iron: whereby oxides of the metals are formed and hydrogen liberated. (2) By heating steam with chlorine, whereby hydrogen chloride (hydrochloric acid) is formed and oxygen set free. (3) By the electrolytic decomposition of water, whereby both oxygen and hydrogen are liberated. Both by synthesis and analysis, we learn that water contains its two constituents combined in the proportion of one part by weight of hydrogen, to eight parts by weight of oxygen ; and also that the proportion by volume is two volumes of hydrogen to one volume of oxygen. When two volumes of hydrogen and one volume of oxygen combine, they give two volumes of steam. When seen in mass, water appears bluish-green. Its point of maximum density is 4 C. When it freezes it expands ^ of its volume. CHAPTER IX. WEIGHING AND MEASURING. As soon as chemists began to study chemical changes from a quantitative point of view; that is to say, as soon as they realized that certain relations existed between the quantities of the substances which take part in chemical changes, it then became important to be able to measure and to weigh with great accuracy ; and as chemistry has become more and more an exact science, the various methods of measuring and weighing have developed in accuracy and delicacy. The Metric System. The measures and weights used in all scientific work are those of the French metric system. It is called the metric system because the standard unit of length is the metre. This corresponds roughly to the English yard, being 39*37 inches, and to this measure the standards ot capacity and of weight are very simply related. (i) Measures of Length. The metre is divided into tenths, hundredihs, and thousandths, which are called respectively, decimetres, centimetres, and millimetres. There are therefore 10 millimetres in i centimetre, and 10 centimetres in i deci- metre. The relation between these subdivisions and the English inch will be seen by Fig. 58, which shows side by side a deci- metre scale, divided into ten centimetres, and each of these again into ten millimetres ; and a four-inch scale, divided into sixteenths. Roughly speaking, 2 '5 cm., or 25 mm., are equal to i inch. For the most part, the chemist uses only one measure of 86 Weighing and Measuring. I length, viz, the millimetre ; he records a length as 20 mm., or 760 mm., as the case may be, instead of 2 cm., or 7 dm. 6 cm. 1000 metres, or i kilometre, is a little over half a mile (0*62 miles), and is the unit employed on the continent of <>> Europe for measuring distances. Instead of milestones, kilometre stones are employed upon the roads. (2) Measures of Capacity, or Volume. The . space occupied by a cube whose sides measure one centimetre (see dotted cube, Fig. 58); that is to say, the volume of one cubic centi- metre is taken as the unit. As there are 10 centimetres in a decimetre, there will be 10X10X10= 1000 cubic centimetres in a cubic decimetre. A cubic decimetre is called a litre; a litre, therefore, contains 1000 cubic centi- metres. (The litre is equal to about if pints, and is the common continental unit of mea- ts sure for liquids.) (3) Measures of Weight. The weight of one cubic centimetre (i cc.) of pure water (measured at a temperature when water is most dense, see p. 79) is taken as the standard unit of weight, and is called a gramme (spelt in English, gram). Fig. 59 represents a gram weight (in brass), exact size; it is equal to about 15^ grains ; and 31*1 grams are equal to i oz. Troy weight. The gram is subdivided into tenths, hundredths, and thousandths, called decigrams, centigrams, and milligrams respectively. The kilogram (frequently called simply a "kilo") is 1000 grams. It is equal to about 2\ Ibs., and is the common continental unit. Since i cc. of water weighs i gram, it is obvious that we at once know the weight of any volume of water expressed in cubic centimetres ; or, vice versa, we know the volume of any 10 C\| tai Thermometers. 87 weight of water given in grams. For instance, 1000 cc. of water that is, i litre weighs 1000 grams, or i kilo. Similarly, if we know the specific gravity of a liquid (that is, how many times lighter or heavier than water it is) we can easily find out the weight in grams of any volume of it expressed in cubic centimetres, or vice versd. For example, suppose we have 200 cc. of a liquid whose specific gravity is 0-5, and we wish to know how much it weighs : Then, as i : 0*5 : : 200 : x, x = 200 X 0-5 = 100 grams . Or, again, the specific gravity of mercury is 13*6, what will be the volume of 408 grams of this liquid ? Then, as 13*6 : i : : 408 : x, = cc 13*6 Instruments for Weighing and Measuring. (i) Thermometers. These are instruments for measuring tem- perature. Their use depends upon the fact that liquids expand when warmed, and contract when cooled. The liquid most commonly employed is either mercury or alcohol. Water would not be suitable, because when moderately cooled it freezes. 1 The graduations, or degrees, upon the stem of a thermometer are purely arbitrary divisions. The instrument is first placed in a vessel filled with broken ice, and the position of the liquid in the stem is marked upon the glass. It is then placed in steam, from water which is kept briskly boiling. The liquid in the thermometer at once expands and rises in the stem, and the point to which it reaches is also marked upon the glass. These, then, are the two fixed points of the scale, namely, the temperature of melting ice and of steam from boiling water. The space upon the stem between these two fixed points is afterwards divided into a certain 1 The student must refer to text-books on physics for details of the method of making and graduating thermometers. 88 Weighing and Measuring. number of equal parts, and the divisions continued both above and below. It is quite optional how many equal divisions the space between the two fixed points shall be divided into. For instance, we might decide to divide it into ten equal parts, calling the lower starting point o. Then, on our scale, the melting point of ice would be o ; the boiling point of water would be 10; blood heat would be 37; and the ordinary temperature of a warm room would be 1-5; but it will be R c F obvious that it would be best to adopt one recognized scale. Unfor- tunately there are three scales in common use. In one, the space be- tween the fixed points is divided into 80 equal parts, the lower starting point being the zero. In the second, it is divided into 100 equal parts, the gzero being the same. In the third, lit is divided into 108 equal parts, and ^the zero is placed 32 divisions below the lower fixed point. The first of these, called the Reaumur thermometer (R. Fig. 60), is the one in common use on the con- tinent of Europe. The third is known as the Fahrenheit thermometer (F.), and is commonly used in England for ordinary purposes. As measured by this scale, the melting point of ice (or the freezing point of water) is 32, and the boiling point of water The other is the Celsius, or more commonly known as the Centigrade (C.) thermometer, and is always used for scientific purposes. The relation in which these three scales stand to each other will be evident from the figure. It is obvious that the divisions on the C. scale are nearly twice as long as those of the F. scale, 100 C. divisions being equal to 180 F. degrees, or iC. = i 8 F. FIG. 60. IS 212 The Barometer. 89 In order, therefore, to translate temperatures given upon the centigrade scale to degrees Fahrenheit, it is only necessary to multiply by r8, and add 32, because the zero of the latter scale is placed 32 divisions below the point at which the centi- grade scale starts. And vice versd, to convert degrees F. into degrees C, we must first subtract 32 and then divide by i'8 : ( C. x 1-8) + 32 = F. ; and ?LA.- 3? = C. I'o Example (i). 1 The temperature of a bath is 95 C. What would this be on the F. scale ? 95 x r8 = 171-6 ; 171-6 + 32 = 203-6. /. 95 C. = 203-6 F. Example (2). The temperature recorded on a hot summer day was 95 F. What is this on the C. scale ? _ A 3 = .'. 95 F. = 35 C. (2) The Barometer is an instrument for measuring the pressure of the atmosphere. In its simplest form it consists of a long straight glass tube closed at one end, which has been filled completely with mercury, and inverted with its open end in a dish of mercury. Experiment 8$. Take a stout glass tube about a metre long and seal up one end in the blow-pipe (see p. 37). Pour mercury into the tube until it is nearly full, firmly close the open end with the thumb, and tip the tube so as to allow a large air bubble to travel to the other end. This will sweep out all the small bubbles which were sticking to the glass. Now completely fill up the tube with mercury, once more close it with the thumb and invert it, and lower the end into a dish or trough containing mercury and care- fully withdraw the thumb. Notice that the mercury falls in the tube to a certain point, and there remains stationary (a, Fig. 61). 1 The student should set himself a few such examples to work out. For instance, he may verify the various temperatures which are shown to correspond on the scales in Fig. 60. go Weighing and Measuring. Now, why does the mercury fall in the tube at all ? and why, if it falls at all, does it not drop down altogether? In order to answer these questions let us first inquire what there is in the clear space above the mercury. Can it be air? Unless a little air was accidentally allowed to get in when the tube was inverted in the mercury dish, it can hardly be air, because the tube was entirely filled with mercury, and the thumb was not removed until after the mouth of the tube was actually under the surface of the liquid. Try the following experiment : Experiment 86. Slowly tilt the tube, and notice that the mercury goes up nearer and nearer to the top, until at last the tube is as completely full of mercury as it was before it was put into the trough. This proves that there is no air or any other gas in the tube. Raise it once more to the perpendicular position and the mercury again falls in the tube. This experiment shows that there is nothing in the space above the mercury, but that it is a vacuum. 1 This being the case, it follows that if we were to connect one end of a long tube to the most perfect air-pump, and dip the lower end in mercury, it would be impossible to pump the mercury up the tube to a greater height than it stands in the tube in the barometer, because in that tube there is a perfect vacuum in the space above the mercury. Experiment 87. Measure the height of the column of mercury in the tube, measuring from the surface of the liquid in the dish to the top of the arched surface (called the mamscus\ a, Fig. 61, of that in the tube. It will be found to be about 760 mm. [If no millimetre measure is at hand, use a foot rule, which will give nearly 30 inches as the height.] Now stand the apparatus against a wall, and make a mark opposite the top of the mercury in the tube, and rule a horizontal line along the wall at that height. Then gradually tilt the tube into the positions /, /', Fig. 61, and notice that all the time, the top of the mercury keeps level with the horizontal line, that is to say, it keeps at exactly the same actual vertical height. 1 This experiment was first made by the Italian physicist Torricelli, and the space is called the Torricellian Vacuum. The Barometer. The mercury is kept up in the tube, because the atmosphere pressing down upon the mercury in the dish, pushes it up until the weight of liquid in the tube exactly balances the pressure of the atmosphere. If from any causes the pressure of the atmosphere be- comes greater, it will push the mercury higher up in the tube ; and, in the same way, if the at- mospheric pressure be- comes less, it cannot sup- port so long a column of mercury, and therefore the liquid sinks a little in the tube. It is per- fectly easy to show by experiment, that it is the pressure of the air upon the liquid in the dish which keeps the mercury up in the tube. If, for instance, we surround the apparatus with a tall glass shade, and with an air-pump gradually pump the air out of the shade, we should notice that the mercury would gradually sink lower and lower in the tube. By the time we had pumped out half the air (that is, reduced the pressure of the air on the mercury in the dish to one half), we should find that the mercury was now standing at only half its original height in the tube, namely, 380 mm. instead of 760. Again, if we were to break open the top of the long tube, the mercury would of course instantly sink down to the same level as that in the dish, because the atmospheric pressure would then be the same on the liquid inside and outside the tube. The pressure of the atmosphere is different at different parts of the earth's surface ; and is also liable to vary in the same locality, from hour to hour. Consequently the height of the column of mercury which it is able to support also FIG. 61. 92 Weighing and Measuring. varies; therefore the barometer enables us to tell at any moment what is the actual atmospheric pressure at the time. What is called the standard or normal pressure is that which is able to support a column of mercury 760 mm. high. (3) The Balance. In order to weigh any substance accurately, or to detect minute differences in the weights of different things, it is necessary to employ a delicate balance and exact weights. For instance, a few grains of sand sprinkled upon the scale of an ordinary kitchen balance will FIG. 62. make absolutely no difference, whereas a single particle of sand placed upon one pan of a chemical balance would com- pletely weigh it down. The chemical balance is usually enclosed in a glass case, partly to protect it from dust and dirt, and partly in order that it may not be exposed to the slightest draught when being used. Such a balance will readily weigh to a fraction of a milligram. For the purposes of the elementary student, an instru- ment of the extremest delicacy is not necessary or desirable. Fig 62 shows a balance suitable for his requirements. By turning the handle or lever H, the beam is liberated from its support, and is then free to swing. This balance The Balance. 93 will turn with 2 milligrams, and is able to carry as much as 100 grams. It must not be used for heavier weights than this. The set of weights consists of the following, 50, 20, 10, TO, 5, 2, i, i, i, grams, making up 100 grams; and 0-5, 0-2, O'l, O'l, 0*05, O'02, O'OI, O'OI. In order to make a weighing, we proceed as follows Experiment 88. Place the object to be weighed, say a clean porcelain crucible with its lid, upon the left pan, and put on the other pan a weight which is roughly judged to be equal to it, say the 20-gram weight. Use the forceps (not the fingers) to lift the weights, and place them gently upon the scale pan. Now release the beam, by means of the lever, and observe which scale pan falls ; we will suppose the 2o-gram weight is too much. Raise the beam again, lift off the weight, return it to its place in the box, and put the lo-gram weight on the pan. Suppose this is too little, put on the 5-gram in addition. The weights must never be removed or added while the beam is free, but only when it is supported in the rest. If 15 grams is too much, remove the 5, return it to its place, and put on the 2-gram : if this is too little put on i gram, and if this is too much then the crucible weighs more than 12 but less than 13 grams. Now add 0*5 gram ; suppose this to be too heavy, remove it and try o'2. If this is too little add 0*1, if now too much remove the OT and put on 0-05. Suppose this to be the exact weight, then, as the beam oscillates, the pointer will swing to the same distance upon the scale on both sides, and the weight of the crucible is 12*25 grams. [NOTE. When not in use, the balance and the weights should not be left exposed to the laboratory atmosphere, but shoidd be either covered ivith a glass shade, or put away into a clipboard^ (4) For measuring liquids, graduated glass vessels are employed. For moderate volumes, a graduated 250 cc. cylinder may be used (Fig. 63). While for small quantities a burette is more useful (Fig. 64). When small definite volumes are required, say exactly 10 cc. or 25 cc., a pipette is employed. This is dipped into the liquid, which is then sucked up the tube, nearly to the top, by the mouth, and the top quickly closed with the finger (Fig 65). By cautiously releasing the finger, the liquid is 94 Weighing- and Measuring. allowed to drop out until it sinks exactly to the mark on the narrow stem. When larger definite volumes are required, flasks holding i, , or i litre are used. Experiment 89. Select a flask with rather a narrow neck, and of such a size that when 500 cc. of water have been measured into it (by twice filling the 250 cc. cylinder) the water stands in the neck of the flask. Now gum a label round the neck in such a position that the top edge of the label is exactly level with the TOM FIG. 63. FIG. 65. FIG. 64. water. If now, by means of a file, a slight scratch be made upon the glass, just where the upper edge of the label is, a half litre measuring flask will have been made. The volume which a given weight of a liquid occupies, however, depends upon the temperature of the liquid. Experiment 90. Fill the half litre flask exactly to the mark with cold water ; place the flask upon a piece of wire gauze upon a tripod and gradually warm it with a gas flame. Notice that the water rises in the neck of the flask some distance above the mark. Measuring Gases. 95 The water has expanded^ or increased in bulk. Now, by means of a pipette, withdraw some of the water, until exactly the half litre is left, and let the flask cool again. When it is cold, observe that it no longer contains as much as half a litre of water. This shows that, if we measure a volume of liquid at a low temperature, we have a greater actual weight of it than if we measure it at a high temperature, and therefore we ought to measure always at the same temperature, or else make a proper allowance for the expansion or contraction. In ordinary experiments, when extreme accuracy is not necessary, and the quantities are only small, measurements made at the ordinary temperature of the room are regarded as being made at the same temperature, and no correction is made. The wholesale spirit dealer, on the other hand, takes careful note of the temperature at which his spirit is measured, else if he were to buy large volumes in hot weather, and sell it in cold weather, he would be a considerable loser by the transaction. (5) Measuring Gases. We may measure the volume of a gas obtained in a chemical process, by collecting it over water or mercury in a vessel graduated into cubic centimetres ; or we can collect it in an urigraduated vessel, mark the volume by gumming a label upon the glass, and afterwards seeing what volume of water has to be poured into the vessel from the 250 cc. measure, to fill it up to the mark. This is a compara- tively rough method ; in more exact work it is usual to measure gases in glass tubes over mercury by means of .a millimetre scale. Vessels used for measuring gases are called eudiometers. Gases expand when warmed, and contract on being cooled just as liquids do, but to a much greater extent. Experiment 91. Fit a cork and delivery tube into a dry half- litre flask, and arrange it as shown in Fig. 66. Gently heat the empty flask with a lamp, and notice that the air within the flask quickly expands, and escaping through the delivery tube, is col- lected in the cylinder over water. Observe that the volume of the gas which is driven out of the flask is much greater than the volume of liquid which was removed from the flask of water in Exp. 90 by the pipette. Remove the flame, and notice that, as the air in the 96 Weighing and Measuring. flask cools, and therefore contracts, water is drawn back from the trough to take its place. When measuring the volume of gases, therefore, it is always FIG. 66. necessary to take into account the temperature of the gas at the time, and then to calculate what the observed volume would be if the temperature were to be lowered to o C, which is taken as the standard or normal temperature. CHAPTER X. SOME GENERAL PROPERTIES OF GASES. Relation of the Volume of Gases to Heat. Nearly a century ago it was discovered by Charles and Gay-Lussac that all gases expand and contract to the same extent under the same changes of temperature, provided there is no alteration in the pressure. This is known as the Law of Charles. For every degree of the centigrade scale that a gas is heated, it expands - 5 -f 3- part of the volume it occupies at o C. Thus, i vol. at o C. becomes i -f ^ at i, and i + ^ at 2, and so on. Or 273 vols. at o become 273 + 1, or 274 at i 273 + 2, 275 2 273 + 3, 276 3, etc. v > , 2 73 + /, f If V stands for the volume at o C., and V for the volume at / temperature, then we get the simple proportion 273:273 + '::^: V, 273 EXAMPLE i. Suppose we have a quantity of gas which measures 320 cc., while the temperature is 15 C. What volume will this occupy at o ? Then V = 32 X27 3 = 303-3 cc. 273 + 15 EXAMPLE 2. 250 cc. of gas measured at - 10 C. What will be the volume at the normal temperature ? Then V = Hl x 2 73 = 2$ $ 273 - 10 Some General Properties of Gases. EXAMPLE 3. A quantity of gas occupies 100 cc. when measured . What will it measure when heated to 30 ? Then V = _ = cc . 273 EXAMPLE 4. 560 volumes of gas measured at 10, are heated to 20. How many volumes will they then occupy ? Then, 273 + 10 : 273 + 20 : : 560 : V or, V = 56o x (273 +.20) = 5?6 volumes 273 + 10 Relation of the Volume of Gases to Pressure. If we squeeze a solid or a liquid, we observe no change in their volumes ; but if we put pressure upon a gas, it at once becomes greatly reduced in volume. Robert Boyle (1661) dis- covered that the volume occupied by a given weight of any gas is inversely as the pressure. This is known as Boyle's Law. It means that if we double the pressure on a gas we reduce the volume to one- half; while if we diminish the pressure to one-half, we increase the volume to double. It is easy to put this law to the test. Suppose a quantity of air is en- closed in the tube T (Fig. 67), which is connected by means of a thick india- rubber tube to a small reservoir of mer- cury, capable of being hoisted up and down. Let us first place the reservoir in such a position that the mercury in it is exactly level with that in the tube. Under these circumstances the enclosed air is under the ordinary atmospheric call it i vol., and mark it upon the tube. Next we will place a mark at an equal distance below, and another at one-half the distance above it Now suppose we raise the mercury reservoir, we shall see the volume of the gas diminishing until presently the mercury FIG. 67. pressure. Let us The Relation of Volume to Pressure. 99 stands at the i-mark. When it reaches this point we shall find, on measuring, that the height of the mercury in the reservoir above that in the tube is about 760 mm. We have already learnt that 760 m;n. of mercury represents a pressure equal to that of the atmosphere, hence we have subjected the gas within the tube to a pressure of an additional atmosphere ; that is, there is now twice the pressure upon it that there was at first, and its volume is reduced to one-half. In the same way, if the reservoir be lowered down until the gas in the tube has expanded to twice its original vol., that is. down to the 2 mark, we should find, on measuring, that the mercury in the reservoir was standing 380 mm. below that in the tube. Instead of being under the ordinary atmospheric pressure, the gas is now under reduced pressure. It is sub- jected to a pressure of 760 mm. 380 mm. = 380 mm. That is, to a pressure of only half the ordinary atmospheric pressure, and we see, as demanded by Boyle's law, that its volume is doubled. Experiment 92. Take a glass tube about half a metre long, and closed at one end, and completely fill it with mercury. Now pour out this mercury into a cc. measure, and in this way ascertain the exact capacity of the tube. Then pour back into the tube one-half the mercury, and mark upon the tube where it comes to. This mark will therefore indicate exactly half the capacity of the tube. Now close the tube with the thumb and invert it in a dish of mercury, and stand it in a vertical position. Notice that the mercury does not stand at the mark, but at a point some distance below it. Why is this, for the tube when inverted was half filled with air and half with mercury ? The reason is because the gas is under reduced pressure, and has therefore expanded. It is under a pressure equal to that of the atmosphere minus that of the column of mercury in the tube. We must, therefore, first ascertain the actual height of the barometer at the time, and then carefully measure the height of the column of mercury, from the surface of that in the dish to the top of that in the tube. Suppose the barometer happens to be low, say 740 mm., and the length of the column in the tube to be 200 mm., then 740 - 200 = 540 mm. is the pressure to which the gas is exposed. Experiment 93. Take the tube used above, fill it with water up to the mark and invert it in a dish of water. Notice where the ioo Some General Properties of Gases. water stands ; it only falls a very little way below the mark. Why; is this ? It is because, water being so much lighter than mercury (13^ times lighter), the short column of it in the tube scarcely reduces the pressure on the gas at all. Seeing that the volume of a gas is so closely dependent upon the pressure to which the gas is exposed, it will be evident that if we wish to compare volumes of gases, we must know what they measure under the same conditions of pres- sure. Now it is not often possible, and is seldom convenient, to actually put gases under one regular pressure before mea- suring their volumes, so the plan always adopted is to measure the volume of the gas at the particular pressure it happens to be under, and then to calculate what volume it would occupy at a pressure of 760 mm., this being the standard or normal pressure. This is called correcting for pressure, and the calcu- lation is a simple proportion, based on Boyle's law. For example, ioo cc. of a gas measured at 380 mm., what volume would it occupy at 760 mm. ? Then, as 760 : 380 : : ioo : x = ioo X 380 = ca 760 Again, a quantity of gas measured 500 volumes at 780 mm. What will it occupy at the standard pressure ? As 760 : 780 : : 500 : x * = 552*-Z?? = 5 13 volumes. 700 In practice, the corrections for both temperature (p. 97) and pressure are usually made together. For example, 1 a sample of gas measured 360 cc, its temperature was 15 C, and it was at atmospheric pressure ; but the barometer at the time was standing at 750 mm. Then 3g, x 2 73 _ tne temperature correction alone, 273+15 and ^ *5 _ t ne pressure correction alone. 760 1 The student should make himself perfectly familiar with this method of reducing volumes of gas to the normal temperature and pressure, by working out a number of examples. The Crith. 101 Putting the two together, we get 360 x 273 x 750 _. 73710000 = -- 6 . 7 cc (273 + 15) X 760 218880 Therefore, 360 cc. of gas at i5"G wJ 750 mite '^'336*7 cc. at the normal temperature and 'ptessure. v ' The Crith. One litre 'oV hydiogf.n gas, sri^asiJr6d<' at o C. and 760 mm., weighs 0*0896 gianib. This is an 'ex'ttemely im- portant figure to remember, for by means of it we can calculate the weight of any volume of any gas, so long as we know the density of that gas; that is, how many times heavier than hydrogen it is. For instance, if we learnt that oxygen was sixteen times as heavy a gas as hydrogen, bulk for bulk, then i litre of oxygen will, of course, weigh sixteen times as much as i litre of hydrogen; that is, 0-0896 x 16, or 1*42 grams, at normal temperature and pressure (N.T.P.). This number, 0*0896, is so important that a name has been given to it (just as in mathematics the name of the Greek letter TT is used to denote the number 3-1416, the ratio between the diameter and the circumference of a circle). The name adopted for the weight of a litre of hydrogen measured at N.T.P. is the Greek word signifying a barleycorn, crith. It is used symbolically for a little weight, and does not mean that a litre of hydrogen weighs as much as a barleycorn. We say that a litre of oxygen weighs 16 criths ; that is, simply sixteen times o f o896. Another important number to be remembered in this connec- tion is the volume of i gram of hydrogen, measured at N.T.P. i gram of hydrogen measures 11*165 litres at N.T.P. From this number, just as from the crith^ we are able to calculate the weight of any volume of any gas whose density is known. Thus, if oxygen is sixteen times as heavy as hydrogen, then, since i gram of hydrogen measures 11*165 litres, obviously, 16 grams of oxygen will measure in65 litres. In other words, 11*165 litres is the volume which will be occupied by the same number of grams of a gas, as expresses the density of that gas. Suppose we have three gases whose densities are respectively 14, i8, and 22 (that is, one is 14 times, one i8 times, and the other 22 times as heavy as hydrogen), then 14 IO2 Some General Properties of Gases. grams of one, i8 grams of the second, and 22 grams of the third will measure 11*165 litres. The liquefaction of gases. We have learnt (p. 97) that when a gas is heated from o to i it expands by o^ ; j of its bulk ; if it be cooled frorrup^tO i'ij: ,&*####? '^} 5 of its volume; from o to 2, - 3, - 4, etc"., 'ft contracts' ^f j, ^>f g , and %$ s respectively. Now, if this law."hal(ls;'gp*dd., 'however, m^ch we cool a gas, it would follow that if 4 a"*qua*ntpfy' Of* gas : be' cooled 'from o to -273, it would contract f | of its volume, that is, it would occupy no volume at all. Now, a temperature as low as - 273 has never yet been reached (this temperature is sometimes called the absolute zero). The lowest degree of cold which has been obtained as yet, is about -220, but before this point is reached all the known gases except one, and that one is hydrogen, pass into the liquid condition. Just as steam, when cooled, changes from the gaseous state to the liquid, so all other gases, when sufficiently strongly cooled, change from the gaseous to the liquid condition. Some gases require very little cooling to make them do this, while others require to be exposed to the lowest possible temperature in order to make them change their state. Among the latter class are oxygen and nitrogen, and it is because it is only recently that chemists have been able to obtain the necessary degree of cold, that these gases have only of late years been obtained in the liquid state. For example, oxygen requires to be cooled down to the extremely low temperature of - 181 to cause it to pass from gas to liquid. There is little doubt but that before long a sufficiently low temperature will have been reached in order to produce liquid hydrogen. As a gas gets near to the temperature at which it turns into a liquid, it begins to depart from the law of Charles and Gay-Lussac, and, of course, when it liquefies, and therefore ceases to be a gas, it is no longer subject to the laws which govern gases. Again, it is found that when gases are subjected to pressure, they sooner or later begin to depart from Boyle's law, and finally to change their state from gases to liquids. A most remarkable point, however, about the effect of pressure in causing the liquefaction of a gas, is, that the gas must be below a certain temperature. If it be above this temperature, no amount of pressure will squeeze it into the liquid state. This particular temperature is different for each gas, and is called the critical temperature of the gas. For example, the critical temperature of chlorine is 141. This means that if chlorine is heated above this point, no amount of pressure will Liquefaction of Gases. 103 make it pass into the liquid state ; but at all ordinary temperatures of the air, chlorine is far below its critical temperature, therefore it can be easily condensed to the liquid state by mere pressure. Again, the critical temperature of ethylene (see p. 184) is about 10, that is, just a trifle below the ordinary temperature of a room ; hence, in order to compress this gas into the liquid state, it must be slightly cooled, so as to bring its temperature below its critical point. In the case of some gases the critical temperature is extremely low, thus in the case of oxygen it is 118*8. Therefore, in order to compress oxygen into a liquid, it is absolutely necessary to cool it down to this intense degree of cold before liquefaction will take place. At - u8'8 oxygen requires a pressure of fifty atmospheres to liquefy it. The more it is cooled below this, the less pressure is needed to liquefy it, until at - 181 it passes into the liquid condition at the ordinary atmospheric pressure. The critical temperature of hydrogen is lower still, and up to the present no artificial cold has been obtained low enough to cool hydrogen in quantity down to this point ; hence hydrogen has not yet been obtained as a coherent liquid, although by special devices momentary indications of liquidity have been observed, when the hydrogen appeared as a froth or spray. CHAPTER XL SIMPLE QUANTITATIVE MANIPULATIONS. 1 EVERY one knows that if we mix two substances together, say salt and sugar, we can have the two ingredients present in any proportion we please. We could, for instance, make a mixture containing so large a proportion of salt, that the sugar present could not be tasted ; or the sugar might so preponderate that we could not taste the salt in the mixture. Now a most important question arises, viz. Can substances enter into chemical combination in this manner ? That is to say, when two elements unite, can they do so in any proportion ? does the composition of the compound they produce vary ? If two substances can combine together in any proportion, just as they can be mixed together, then obviously the composition of the compound produced will depend upon the quantities of the ingredients that were employed to produce it. If, on the other hand, it should be that there is some fixed proportion in which elements combine, that if they unite it must be in some particular proportions, or not at all, then it would follow that any given compound would always have absolutely the same composition. This is a point which it is of the very greatest importance to settle, and in order to do so we must make experiments upon weighed or measured quantities of sub- stances, and carefully weigh or measure the resulting products. 1 The quantitative experiments have been arranged as far as possible so as to make use only of such knowledge of chemical facts as the student has already gained from the earlier chapters. The teacher will do well to substitute or a and to speak of the air as a supporter of combustion. Combustion. 17 1 We have learnt by numerous experiments that it is because of the oxygen present, that the air supports the combustion of burning bodies ; oxygen, as we have seen, being such a good supporter of combustion. In ordinary language we call a gas a supporter of combustion, if it behaves towards common com- bustibles in the same way that air does ; but in reality there is no distinction between a combustible and a supporter of com- bustion. For instance, we have seen sulphur burning in oxygen, and we therefore call sulphur a combustible, and oxygen the supporter of combustion. But let us modify the conditions of the experiment. Experiment 138. Heat some sulphur in a wide test-tube until it boils and the dark vapour takes fire at the mouth. Then lower into the test-tube a bent glass tube (so bent that it can enter the test tube) through which a gentle stream of oxygen is passing (Fig. 79). As the jet is passed through the flame of sulphur burning at the mouth, the oxygen is ignited, and will continue burn- ing in the sulphur vapour when pushed down into the tube. Under these conditions the oxygen is the combustible, and sulphur vapour is the supporter of combustion. Again, we have seen hydrogen burn in oxygen, but we can reverse the condi- tions and make oxygen burn in hydrogen. Experiment 139. Fill a jar with hydrogen. Hold it mouth downwards, and apply a light to ~~~FIG the gas. While the hydrogen is burning, thrust a jet, from which oxygen is slowly issuing, up into the jar (Fig. 80). As the jet passes through the hydrogen flame it is ignited, and goes on burning just as it did in the sulphur vapour. 172 The Atmosphere. Here oxygen is the combustible, and hydrogen the supporter of combustion. Combustion is merely the term applied to describe any chemical action which takes place with so much energy as to produce light and heat. All the more common cases of com- bustion, are the active combination of substances with oxygen, that is, they are rapid processes of oxidation ; but oxygen is not necessary to combustion ; we can have cases of combustion in which oxygen does not participate, for we know of many instances in which chemical action takes place between other substances with sufficient energy to give rise to light. For instance Experiment 140. Take a jar of chlorine, and lower into it a jet from which ammonia gas is slowly escaping (the ammonia being obtained by gently warming a strong solution of ammonia, in a little flask; see Exp. 145). Notice that as soon as the jet of ammonia enters the chlorine it at once takes fire, without being lighted, and goes on burning in the chlorine, (Fig. 81). There is here no oxygen present ; the ammonia is the combustible, and the chlorine is the supporter of combustion. We have also seen (Exp. 119) that many metals take fire and burn when brought into chlorine. Temperature of Combustion. The actual temperature which is produced by any particular process of combustion depends partly upon how quickly the combustion proceeds. For instance, if we burn some FIG. 80. . . - . - . substance m oxygen it burns much faster than when burnt in air, and consequently the temperature is hotter. Experiment 141. Burn a jet of hydrogen from an oxyhydrogen blowpipe (this is merely a very fine-pointed metal tube with a still smaller one passing down the inside, exactly like an ordinary Herapath blowpipe, only smaller). Hold a little piece of platinum wire in the flame, and note that although the wire becomes very hot, it does not melt. Next hold a little block of hard lime against Temperature of Combustion. 173 FIG. 81. the flame ; the lime does not seem to get very hot. Now gently turn on the oxygen, so that the hydrogen flame is fed with oxygen. Notice that the flame does not show any more light, but if the plati- num wire is held in the flame it is instantly melted, be- cause the flame is now so much hotter ; and if the lime is brought into the flame it at once gets so hot as to emit a very bright light. This is the oxyhydrogen lime-light. In everyday life when ^ we wish to increase the rate of combustion, and conse- quently raise the temperature of combustion, we increase the draught of air (by the use of bellows, for instance) so that more oxygen is driven against the burning body in a given time. Ignition point. The particular temperature at which a substance begins to burn, or to " take fire," is called its ignition point. Some- times this is lower than the ordinary temperature of the room, in which case the substance takes fire by it- self when brought into the air. Obviously such things as these must be kept so that they do not come into contact with the air. Experiment 142. Place :f| a small quantity of caustic soda solution in a test-tube, ^ and put in it a piece of phos- phorus about the size of a pea. Attach a cork with two tubes, arranged as in Fig. 82. First pass a stream of ordinary coal gas 1/4 The Atmosphere. through the little apparatus by means of the indiarubber pipe con- nected with the tube T. This is in order to sweep out the air from the apparatus. Now gently boil the liquid, and a gas is given off called phosphoretted hydrogen, which will bubble through the water in the little basin. In a minute or two, when the coal gas has been expelled, and only the phosphoretted hydrogen is bubbling out, each bubble will take fire as it conies into the air. This gas has a very low igniting point. In all the familiar processes of combustion, it is neces- sary to first heat the combustible substance in order to start the combustion ; that is to say, the igniting point is above the common temperature, and therefore the substance must be heated up to that point before active chemical combination can take place. If after a substance is ignited, it will continue burning by itself, like a candle, or piece of paper, this shows that the temperature of combustion is higher than the igniting point; because, as each particle burns, the heat given out is able to set fire to the next particle, and so on. Heat of combustion is the amount of heat as dis- tinguished from the temperature^ produced by combustion. If we draw a pint of hot water and a gallon of water from the same boiler, the temperature of each sample will be the same ; a thermometer placed in each will show the same temperature. But it is obvious that the amount of heat in the gallon of water is greater than in the pint; there is, of course, eight times as much heat in the one as in the other. From this illustration it will be evident that the amount of heat cannot be ascertained by the thermometer. It is really measured by finding out how much water it is capable of heating from o to i. That amount of heat which will raise the temperature of i gram of water from o to i is taken as the unit, and is called the thermal unit (or calorie). Now the amount of heat produced in any process of com- bustion, is exactly the same whether the process be slow or quick, therefore it is quite independent of the temperature that is produced. For example, the amount of heat given out Heat of Combustion. 175 when a definite quantity of hydrogen is burnt in the air, is exactly the same as that produced when the same quantity of hydrogen is burnt in oxygen, although, as we have seen, the temperature in the latter case is very much higher. Indeed, this is also true if the process of oxidation is so slow that there is no active combustion. Thus, when a quantity of iron slowly rusts, heat is produced ; but the process is spread over such a long time, that the iron never even gets warm, the heat being conducted away as fast as it is produced. When the same quantity of iron is burnt in oxygen, the temperature rises enormously, because the process is complete in a few moments, but the actual amount of heat is the same in both cases. CHAPTER XIX. NITROGEN AND ITS COMPOUNDS. Nitrogen. The element nitrogen, as we have seen, is present in the free state in the air ; f of which is nitrogen. It can be obtained from the air by removing the oxygen either with phos- phorus (Exp. 133) or by means of heated copper (Exp. 135). Nitrogen, combined with other elements, is present in a number of compounds, from some of which the element is readily expelled. One of the commonest of the compounds of nitrogen is ammonia ; a compound of nitrogen with hydrogen. If we act on this compound with chlorine, the chlorine takes the hydrogen (forming hydrochloric acid) and the nitrogen is set free, according to the equation, 2 NH 3 + 3 C1. 2 = 6HC1 + N 2 Experiment 143. Fit a wide-mouthed bottle with a cork carry- ing two wide glass tubes, as shown in Fig. 83. Half fill the bottle with strong ammonia solution, and pass a stream of chlorine (prepared as in Exp. no) through the liquid. Notice that, as each bubble of chlorine enters the am- monia, there is a flash of fire in the liquid, so energetic is the chemical action that goes on. Observe also the white fumes ; these consist of ammo- nium chloride ; for, as soon as hydro- chloric is formed (as shown in the equation), it combines with some of the ammonia in the bottle and forms ammonium chloride. Collect the gas which passes out, in the pneumatic trough. The object FIG. 83. Properties of Nitrogen. 177 of the very wide tubes is that they may not get stopped up by the ammonium chloride. [Caution do not let the experiment go on too long ; as soon as one or two jars of gas have been collected, stop the operation.] Another compound from which nitrogen can be readily obtained is ammonium nitrite, NT 4 NO 2 . When a strong solution of this salt is gently heated, it splits up into nitrogen and water. NH 4 N0 2 =N 2 -f- 2 H 2 0. Experiment 144. Put about equal quantities of ammonium chloride and sodium nitrite into a flask, fitted with a cork and delivery tube, and about one-third fill the flask with water. Gently heat the mixture and collect the gas over water. As soon as the action begins, remove the lamp, and allow it to continue by itself. If the liquid begins to boil up too much, cool the flask by bringing a dish of cold water under it. When a mixture of ammonium chloride and sodium nitrite is heated, the two salts interact on each other and give sodium chloride and ammonium nitrite, thus NH 4 C1 + NaNO 2 - NaCl + NH 4 NO 2 , and the ammonium nitrite then decomposes according to the first equation. The final result, therefore, of heating these two salts together is expressed thus NH 4 C1 + NaNO 2 = NaCl + 2H 2 O + N 2 . Properties of Nitrogen. Nitrogen is very different from any of the gases we have a*s yet studied. When a taper is put into it we see that the gas is not like hydrogen, for it will not burn. It is not like oxygen, for it will not allow the taper to burn, but at once extinguishes it. It is not like chlorine or hydrochloric acid. The only property of this gas that can easily be shown is, that it seems to have no properties. It does not burn. It does not support combustion. It is not acid ; does not bleach ; does not act on metals ; is not poisonous. Indeed nitrogen is one of the most inactive or inert substances we know. It will not support animal life, not because it is in any way injurious, but simply because 178 Nitrogen and its Compounds. animals must have free oxygen to breathe ; an animal placed in nitrogen dies from suffocation, just as it would if immersed in water. Nitrogen is slowly absorbed by red hot magnesium, forming a compound of nitrogen and magnesium. This is one method by which nitrogen is separated from argon, which is even more inert a substance than nitrogen itself. Although nitrogen will not burn, it will combine with oxygen slowly, when electric sparks are passed through a mixture of the two gases. In this respect again it differs from the still more inactive gas argon, which does not unite with oxygen when sparked with that gas. EPITOME. Nitrogen occurs uncombined in the air, to the extent of about four-fifths. It is obtained from air by withdrawing the oxygen, either by burning phosphorus or red hot copper. These combine with the oxygen and leave the nitrogen. The chemical compound from which nitrogen is prepared, is ammonium nitrite. This when heated gives only water and nitrogen. Nitrogen is characterised by great inertness. It is a colourless, odourless, tasteless gas ; does not burn, nor support combustion or respiration ; is not poisonous. It does not easily unite directly with other elements. At a high temperature it combines with a few metals, and also with oxygen. Just as chlorine is a member of a certain little family of elements (see The Halogens, p. 133), so nitrogen is also the representative of another group or family, which consists of the five elements nitrogen, phosphorus, arsenic, antimony, and bismuth. These have very little likeness to each other in their outward appearance, but they are closely related in their chemical behaviour. Nitrogen and phosphorus are true non-metals, one being a gas, the other a wax-like solid ; they have, therefore, no properties which belong to metals, no metallic lustre, no power of conducting heat or electricity ; and they yield oxides which are acid-forming Antimony and bismuth are metals ; they have metallic lustre, conduct heat and electricity, and form oxides which are basic. Arsenic stands on the border line between the metals and non- metals, and is called a metalloid. It is a black shiny solid, with Compounds of Nitrogen. 179 about as much lustre as graphite. It conducts heat and electricity, but its oxides are acid-forming compounds. The detailed study of phosphorus, arsenic, antimony, and bismuth does not come within the scope of this elementary book. Compounds of Nitrogen. Although nitrogen in the free state is such an inert element, it has very strong chemical affinities when in combination with other elements. It forms a number of compounds with oxygen, with hydrogen, and with both oxygen and hydrogen together. It is also one of the constituents of a large number of animal and vegetable substances, where it is associated with carbon, hydrogen, and oxygen. When such animal substances decay or "go bad," one of the first products of the decomposition is ammonia. The nitrogen in the compound combines with some of the hydrogen and forms this compound. Hence there is generally a strong smell of ammonia in stables and in urinals, where nitrogenous animal matter is undergoing decomposition. This natural process of decomposition is imitated artificially, when we heat such a compound so as not to allow air to get to it, so that it does not take fire. For example, when coal is heated in retorts, as in the manufacture of ordinary coal-gas, air does not get to the coal, and therefore it does not burn as it would in an open fireplace, but it is decomposed into a great variety of compounds, some solid, some liquid, and some gaseous. The process of heating substances in this manner is called destructive distillation, to distinguish it from the ordinary operation of distilling where the original substance is not decomposed. Now, although coal is very largely composed of carbon, it also contains amongst other constituents some nitrogen, and some hydrogen, and therefore, when it is destructively distilled, one of the gaseous products formed is ammonia. This dissolves in the water, which is another product of the decomposition, yielding the so-called "ammoniacal liquor " of the gas-works. This is the chief source from which ammonia is now obtained. Another substance which is formed when animal matter containing nitrogen is allowed to decompose slowly by itself 180 Nitrogen and its Compounds. in the presence of air, is nitric acid. Sometimes in the neigh- bourhood of ill-drained stables or dwellings, especially in hot climates, where decomposing animal matter soaks into the earth, crystals are to be seen on the soil or lower parts of the walls. These crystals consist of nitre. The nitric acid formed by the decomposition of the organic matter, combines with potash present in the soil and forms potassium nitrate (saltpetre or nitre}. At one time all our supplies of nitre were obtained by this process, which was carried on by purposely mixing manure and such decomposing refuse with wood ashes (which contain a large quantity of potash) and earth, and allowing the heaps to remain exposed to the air, occasionally moistening them with drainage from manure. Nowadays this process is not much used, because enormous beds of sodium nitrate (called Chili saltpetre] have been found, from which potassium nitrate can easily be made. Ammonia and nitre are two of the most important com- pounds of nitrogen. Ammonia. This substance is a gas, but long before it was known to be a gas, a solution of it in water was known, and was called spirits of hartshorn. It got this name because it was obtained by the destructive distillation of horns and hoofs of animals. At the present day the source of all our ammonia is the ammoniacal liquor of the gas-works. This is practically a solution of the gas in tarry water. The liquid used in the laboratory, and called ammonia, is simply a solution of the gas in pure water. If we heat such a solution the gas is all expelled, and can be collected. Experiment 145. Gently heat a little strong solution of ammonia in a flask provided with a cork and short exit tube bent at right angles, and attach this to the apparatus for collecting gas by upward displacement (Fig. 84). Notice that the ammonia solution gives off its gas so rapidly that it appears to be boiling, although scarcely warmed to the temperature of the hand. In a few minutes the cylinder will be rilled with ammonia gas. Now remove it and place it mouth downwards in a trough of water. Notice that the water quickly rises in the cylinder, showing that the gas is far A mmonia. 181 FIG. 84. too soluble in water to allow of its being collected over that liquid. Experiment 146. Collect another cylinder of the gas, and as the glass is filling, hold a piece of litmus paper which has been reddened (by being dipped into very dilute acid) against the exit tube ; ob- serve that the gas is strongly alkaline. Also let the gas blow against a moist- ened piece of turmeric paper, and obtain the reddish brown stain due to the action of an alkali upon it. In the old days when all gases were called airs, ammonia was distinguished as the alkaline air. Cautiously smell the gas, not by applying the nose to the ttibe this would give too strong a sniff of it and would be dangerous but by gently wafting the escaping gas towards the face with the hand. Gently thrust a lighted taper up into the gas. Carefully note the behaviour of this gas towards combustibles. When the taper is plunged right into the gas, the flame is extinguished, and the ammonia does not take fire. But when the taper flame is cautiously brought into the gas, at first the gas seems to be trying 'to burn ; a curious brownish- yellow flame appears to surround the taper flame for a moment before the latter is put out. Ammonia, therefore, will not bum in ordinary air, although it seems very nearly to do so. But if we add a little oxygen to the air, then the ammonia will bum quite easily. Experiment 147. Close one end of a wide glass tube (a gas- lamp chimney) with a cork through which two tubes pass : a moderately wide one reaching to the top of the chimney, and a narrow one passing just through the cork, as in Fig. 85. The wider of these tubes is attached to a small flask in which strong ammonia solution is gently heated ; the other is connected to a supply of oxygen. The apparatus is supported in a clamp on a retort stand, not shown in the figure. A little plug of cotton wool should be pushed down to the bottom of the chimney, so as to cover the open end of the narrow tube through which the oxygen enters : this distributes the oxygen all round the centre pipe. 1 82 Nitrogen and its Compounds. First regulate the little gas flame, so that a gentle stream of ammonia escapes up the centre tube. Then bring a lighted taper to the end of this tube and again notice the appearance of flame. Now gently admit oxygen through the narrow tube, still holding the taper to the escaping jet of am- monia. As soon as a little oxygen reaches the top of the chimney the ammonia will light, and will con- tinue burning without the taper, giving a curious yellow-brown coloured flame. Now stop the oxygen, and notice the ammonia flame gradually languish and go out. Combination of Ammo- nia with Acids. Experiment 148. -Pour a little dilute hydrochloric acid, dilute ni- tric acid, and dilute sulphuric acid into three separate little beakers, and add two or three drops of lit- mus solution to each. Now heat some strong ammonia solution in a flask, and pass a stream of the gas into each of the acids until the red colour of the litmus changes to blue, and then evaporate each solution to dryness in separate dishes heated gently over small rose burners. In each dish a residue is left, which but for the slight colour due to the litmus would be white. These residues are salts of ammonia ; they consist of ammonium chloride (sometimes known by its ancient name of sal-ammoniac}^ ammonium nitrate, and ammonium sulphate. Ammonium chloride and ammonium sulphate are com- mercially obtained by driving off the ammonia from the " ammoniacal liquor " of the gas-works, and passing the gas into either hydrochloric or sulphuric acid, just as in Exp. 148. The equations which represent the combination of am- monia with these three acids are as follows : (1) NH 3 4- HC1 = NH 4 C1. (2) NH 3 + HN0 3 = NH 4 N0 3 . (3) 2 NH 3 + H 2 S0 4 = (NH 4 ) 2 S0 4 . Ammonia. 183 How to get Ammonia out of its Salts. Experiment 149. Heat a small quantity of the ammonium chloride obtained in Exp. 148, in a dry test-tube. Note that the salt sublimes. Apply a lighted taper to the mouth of the tube, and see that it shows no signs of a flame of ammonia. Smell the tube, and hold reddened litmus paper to it, and observe not even a trace of ammonia is given off. Therefore we cannot obtain ammonia from this compound simply by heating it. Experiment 150. Heat a little ammonium nitrate in the same way. Note at once a great difference in the behaviour of this salt. It melts ; the chloride did not. Presently it effervesces ; evidently gas is coming off. Is this gas ammonia? Test with litmus ; note no blue effect. Smell the gas ; there is no smell of ammonia. Bring a lighted taper into the gas ; notice that the gas behaves like oxygen. Test with a glowing splint of wood ; it relights in this gas. Can it be oxygen ? We must examine this point later ; in the mean time, the experiment shows that we do not get ammonia by simply heating this salt. Experiment 151. Heat a small quantity of ammonium sulphate in a similar manner. Notice that, as with the nitrate, this salt also melts and gives off gas. Test with litmus and turmeric, and observe alkalinity. Smell the gas, and note that when ammonium sulphate is heated alone, ammonia is given off. Experiment 152. Now take a small quantity of each of the three ammonium salts in three test-tubes, and add to each about the same quantity of powdered lime and apply heat gently. Notice that in each case ammonia is given off, as indicated by litmus or turmeric paper, as well as by the smell. These four experiments show that it is only certain ammonium salts which give off ammonia when heated alone, but that all of them yield ammonia when heated with lime. We can, therefore, use this latter method as a test as to whether a particular salt is an ammonium compound or not. In practice, when it is required to make such a test, we usually employ a solution of sodium hydroxide (caustic soda) instead of the lime, as it is rather more convenient to use. Experiment 153. Add a little caustic soda solution to a small quantity of say ammonium chloride in a test-tube, and gently warm 184 Nitrogen and its Compounds. the mixture. Smell the escaping gas, and hold moistened litmus or turmeric paper to the mouth of the tube. Note abundance of ammonia. Ammonia Solution is made by dissolving ammonia gas in water. Experiment 154. Place some ammonium chloride in a flask, and add about twice as much dry slaked lime and mix the two together. Arrange the flask in connection with a bottle containing water as shown in Fig. 86, and gently heat the flask. Air is first expelled, which bubbles through the water, but presently, as only pure ammonia escapes from the flask, the whole of the gas is absorbed by the water, no bubbles pass- ing through the water. Notice that, after a while, the liquid in the bottle begins to get perceptibly warm to the hand. If, there- fore, we wish to make a very strong solu- tion of the gas, we must prevent this by immersing the bottle in cold water, for we have learnt by Exp. 145, that by warming a solution of ammonia the gas is expelled again. [Note. By arranging the apparatus so that the tube leading into the water is a good length, there is no fear of the liquid in the bottle being sucked back into the flask ; because if it begins to ascend far up the tube (which might happen if the evolution of gas were interrupted by a draught blowing the flame away for a moment), then air is at once drawn into the flask through the small quantity of mercury placed in the bend of the funnel tube.] The chemical changes taking place when ammonium chloride is heated with caustic soda and with slaked lime, are expressed by the following equations : NH 4 C1 -f NaHO = NH 3 + NaCl + H 2 O and 2NH 4 C1 + CaH 2 O 2 - 2NH 3 + CaCl 2 + 2H 2 O. The nature of the change is the same in both cases, which The Composition of Ammonia. 185 will be more easily understood if the equations be dissected in the following way. (i) With the sodium hydroxide NH 3 H + JHO = H 2 Cl (Na NaCl (NH \ V NH 4 C1+ NaHO - NaCl + H 2 O + NH 8 and (2) with slaked lime (or calcium hydroxide) NH 3 NH 3 NH 3 H + (HO HO = H 2 H 2 O Cl 1 Ca CaCl. (NH, H I Cl 2 NH 4 C1 -f CaH 2 O 2 = CaCl 2 The Composition of Ammonia. We have learnt that hydrogen and oxygen unite together very readily, forming water; also that hydrogen and chlorine combine with great ease, giving hydrochloric acid, so that we were able to gain knowledge of the composition of both water and hydrochloric acid by synthesis. But hydrogen only combines directly with nitrogen with great difficulty, and even then under exceptional conditions, therefore we cannot find the composition of ammonia by the direct union of its elements. We can, how- ever, decompose ammonia, and find the proportional volumes of its constituents. We do this by means of chlorine, which has been shown (p. 176) to be able to decompose this com- pound, combining with the hydrogen and letting the nitrogen go free. Experiment 155. Take a long glass tube, closed up at one end, and divide it into three equal divisions with indiarubber bands. Fit to the tube a cork carrying a small stoppered dropping funnel, as shown in Fig. 87. Collect this long tube full of chlorine (using strong brine in the pneumatic trough) and insert the cork. Pour a little strong ammonia solution into the funnel, and allow it slowly to enter the tube, one drop at a time. The first few drops take fire as they go in (as in Exp. 143). When about a dozen drops have been let in, the action is finished. Fill up the funnel with dilute sulphuric acid, so as to neutralize the 1 86 Nitrogen and its Compounds. excess of ammonia, and fit a bent tube into the mouth of it, and let the long end dip into a beaker of water as arranged in the figure. Then open the tap. Water at once runs in, showing that some gas has disappeared, and it will fill up exactly two of the measures, leaving one measure of gas. Test this gas by removing the cork and dipping a lighted taper into the tube. The gas is nitrogen. We therefore have got one measure of nitrogen, which has been expelled from combination with just so much hydrogen as there was chlorine ori- ginally present in the tube, namely, three measures (because chlorine and hydrogen combine in equal volumes. See Exp. 103), Therefore, in ammonia, the nitrogen and hydrogen are combined in the proportion one volume of nitrogen to three volumes of hydrogen. In order to prove that the formula is NH 3 and not N 2 H 6 (both of which contain the two elements in the same FIG. 8 7 . relative proportion), it is necessary to find out the density of the gas by weighing a known volume of it. When this is done, it is found that ammonia gas is 8*5 times as heavy as hydrogen. Its molecular weight, there- fore, is twice this, namely, 17. This proves that the composi- tion is expressed by the formula NH 3 . N = 14. 3H = 3. .'. 14 + 3 = 17. EPITOME. Ammonia is formed by the putrefaction of organic matter con- taining nitrogen, such as manure. Minute quantities of it are found in the air. Ammonia is produced during the destructive distillation of coal (as in the manufacture of ordinary coal gas), when it collects in the watery liquid known as the " ammoniac al liquor." Ammonia is obtained from its salts by heating them with slaked lime or with sodium hydroxide (caustic soda) Ammonia is a colourless gas, with a powerful pungent smell, Ammonia. 187 and a strong alkaline reaction. It is very soluble in water, and therefore cannot be collected in the ordinary way. One litre of water at the common temperature dissolves about 800 litres of ammonia gas, while at o it will absorb as much as 1148 litres. All the gas is driven out of solution when the liquid is heated. Ammonia will not burn in air, but it will burn in oxygen. It will not support the combustion of a taper. Ammonia is a light gas ; it is just over one half as heavy as air, and is 8*5 times as heavy as hydrogen. Ammonia is easily condensed to the liquid state. At the ordinary temperatures, about 7 atmospheres pressure will squeeze this gas into the liquid form. And again, if the gas, without being squeezed, is simply cooled below - 34 C. then it also turns into liquid ammonia. [Note. Liquid ammonia is not the same as a solution of ammonia in water.] Liquid ammonia boils at 337, and has been largely used for the artificial production of ice. When passed through a red-hot tube, or submitted to electric sparks, ammonia gas is decomposed into nitrogen and hydrogen. When a measured volume of the gas is thus decomposed, the volume is doubled; that is to say, two volumes of ammonia give one volume of nitrogen and three volumes of hydrogen. The volume composition of ammonia is also established by decomposing it with chlorine, when it is found that three volumes of chlorine take three volumes of hydrogen to form hydrochloric acid, and leave one volume of nitrogen. CHAPTER XX. NITROGEN AND ITS COMPOUNDS (continued). Nitric Acid, HNO 3 . This substance is one of the most important of all the nitrogen compounds, and from it, either directly or indirectly, we obtain all the other compounds of nitrogen which it will be necessary for us to study. Nitric acid is composed of nitrogen, hydrogen, and oxygen, but we do not make it by causing these three elements to unite directly. It can be produced synthetically, however, for if we add a little nitrogen to a mixture of oxygen and hydrogen, and ignite the mixture, the water which results from the union of the hydrogen and oxygen will be found to be acid, from the presence of a little nitric acid. Cavendish noticed this when making his experiments on the composition of water (p. 73), and was for a long time puzzled to account for the acidity of the water he got, but discovered that nitrogen from the air had found its way into the apparatus. If a rapid stream of electric sparks is passed between platinum wires in a confined space of air, the oxygen and nitrogen begin to combine, forming an oxide of nitrogen. This gas has a brownish colour, and if a little water be then added the brown gas disappears ; it dissolves in the water and yields nitric acid. In this manner nitric acid is produced during thunderstorms. The lightning flashes (which are simply enormous electric sparks) passing through the air cause the combination of some of the nitrogen and oxygen, and the oxide so formed is washed out by the rain. Preparation. Nitric acid is made from either potassium Nitric Acid. 1 89 FIG. 88. nitrate (nitre or saltpetre} or from sodium nitrate (Chili saltpetre). Experiment 156. Place 50 grams of nitre in a glass retort with a stopper, and pour upon it the same weight of strong sulphuric acid. Gently heat the mixture and collect the distillate in a clean flask, which is kept cool by being placed in a dish of cold water, as shown in Fig. 88. Now and then turn the flask round, so as to cool it all over. Notice that reddish fumes appear in the retort, and that a liquid collects in the flask which has a pale yellowish colour. The equation which takes place in this experiment is the following KNO 3 + H 2 SO 4 - HKS0 4 + HNO 3 . The potassium in the nitre changes places with one-half of the hydrogen in the sulphuric acid, giving a salt called hydrogen potassium sulphate, and nitric acid. The salt remains behind in the retort. If sodium nitrate had been used the reaction would have been quite similar. The manufacturer of nitric acid always employs the Chili saltpetre, because it is a cheaper article than the potassium salt. He also carries on the operation at a higher temperature, using large cast-iron vessels in the place of a glass retort, which enables him to get all the hydrogen in his sulphuric acid ex- changed for sodium, according to the equation 2NaNO 3 + H 2 SO 4 = Na 2 SO 4 + 2HNO 3 . The same quantity of sulphuric acid, therefore, is made to decompose twice as much of the nitrate as in the laboratory experiment. At the high temperature, however, required to Nitrogen and its Compounds. completely bring about this second reaction, a little of the nitric acid itself is decomposed, and therefore wasted. Properties of Nitric Acid. When quite pure, the acid is colourless. The sample obtained in Exp. 156 is coloured slightly yellow, because it has dissolved some of the coloured fumes which were produced, Nitric acid fumes strongly when exposed to moist air. It can be mixed with water in any proportions. The strong acid is highly corrosive, and must be handled with great care ; a few drops spilt upon the skin will cause bad wounds. Even when moderately diluted with water it will burn the clothes, and stain the skin yellow. If strong nitric acid is boiled, it begins to FIG. 89. decompose, giving oxygen and nitrogen peroxide (a reddish- brown gas). When heated strongly this decomposition is very rapid. Experiment 157. Arrange a tobacco-pipe as shown in the figure, the mouthpiece just dipping beneath the water in the trough. When the stem is red-hot, pour a few drops of strong nitric acid into the bowl. As the acid passes the heated place it is decomposed into the two gases, oxygen and nitrogen peroxide. The latter gas, however, is very soluble in water, and therefore dissolves in the trough, while oxygen alone is collected. Notice that, as the bubbles first appear in the water, they have a dark reddish colour, while the gas which actually collects is colourless. Test the gas for oxygen. The equation is the following 2HNO 3 = H 2 O + 2NO 2 + O. Nitric acid is therefore a powerful oxidizing material, and it acts on many substances with great energy. Nitric Acid. 191 Experiment 158. Gently heat a small quantity of sawdust in a small porcelain dish or crucible, until the wood has just become charred. Then cautiously let two or three drops of strong nitric acid fall upon the charred mass, and notice that it instantly takes fire. The charcoal burns at the expense of the oxygen supplied by the nitric acid, while fumes of the brown gas are produced at the same time. Experiment 159. Add a little powdered sulphur to a small quantity of strong nitric acid in a test-tube, and gently heat the mixture. Notice that the brown-coloured oxide of nitrogen is again given off in quantity. The acid is therefore being reduced; in other words, it is giving up some of its oxygen to the sulphur, which is gradually being converted into sulphuric acid. Let the action con- tinue for a few minutes, and then test the liquid for sulphuric acid in the following way Add one or two drops of sulphuric acid to some water in a test- tube, and then add some solution of barium chloride. Notice a white precipitate. The reaction here is H 2 SO 4 + BaCl 2 = BaSO 4 + 2HC1. The barium and hydrogen change places, and barium sulphate is produced. This Substance, being insoluble in water, separates out as a white solid. There are, however, other things which would give a white precipitate with barium chloride, but barium sulphate may be distinguished from the other precipitates by the fact that it cannot be dissolved by acids. Therefore, as a confirmation of the test, pour half the white precipitate into a second test-tube, and to one portion add some strong hydrochloric acid, and to the other some nitric acid, and note that the precipitate is riot dissolved in either case. Now apply this test in the case of the nitric acid in which sulphur has been heated. Dilute the acid with water, and add a few drops of barium chloride. A white precipitate proves the presence of sulphuric acid, which requires no confirmatory test, as the liquid already contains nitric acid. Besides oxidizing such elements as carbon, sulphur, phos- phorus, iodine, etc., it attacks a great many metals, converting them into compounds which contain oxygen, and are, therefore, oxidation products. Experiment 1 60. Drop a few fragments of copper wire or foil into a little nitric acid in a test-tube. Notice that violent 192 Nitrogen and its Compounds. action at once sets in. Torrents of the red gas are given off, and the copper quickly disappears, while the liquid becomes blue. The copper is converted into copper nitrate (a blue salt), which remains dissolved in the liquid. A number of other- metals behave in a similar way, such as silver, mercury, iron, lead, zinc. They are converted into nitrates, while the nitric acid is reduced to one of the oxides of nitrogen. We may suppose that the first action of the acid on such metals is that the metal changes place with the hydrogen, thus Cu+ 2HN0 3 = H 2 + Cu(N0 3 ) 2 , and that immediately this nascent 1 hydrogen attacks another portion of nitric acid, taking away some of its oxygen, whereby water is formed and an oxide of nitrogen left. Which oxide of nitrogen is left depends on a number of circumstances, for hydrogen in this state can gradually take all the oxygen from nitric acid, leaving at last only nitrogen. In no case by the action of nitric acid on metals is any hydrogen given off so as to be collected. On account of its powerful action in dissolving metals, nitric acid used to be called aqua-fortis (the strong water) ; but even this acid cannot dissolve the " noble " metals, gold and platinum. Aqua Regia. Both nitric acid and hydrochloric acid are without any action on gold or platinum, but if the two acids are mixed together, then the mixture will easily dissolve either of these metals. Experiment 161. Place a gold leaf in a wide test-tube and pour strong nitric acid upon it. The leaf breaks up into small particles, but notice that it is not dissolved, and may be left a long time in the acid. Do the same with hydrochloric acid in another test- tube, and notice that, in like manner, the gold leaf is not acted upon. Now mix the contents of the two tubes, and in a few moments the gold will be entirely dissolved. The mixture of these two acids is known as aqua regia (the royal water), just because it is able to dissolve the noble metals. When metals dissolve in aqua regia, it is always the chloride, and not the nitrate of the metal, that is formed. 1 Nascent means just born ; and elements at the moment of their liberation from combination, are said to be in the nascent state ; at this particular moment the element is much more chemically active. Nitrates. 193 Impurities in Nitric Acid. Common nitric acid generally contains a number of impurities, the most usual of which are sulphuric acid (derived from the sulphuric acid used in its manufacture) and iron (from the vessels in which the acid is made). The sulphuric acid may be tested for by the method described in Exp. 159. The presence of iron in the acid may be found out by the following test. Experiment 162. Dilute a little of the suspected acid, and add to it a few drops of a solution of potassium ferrocyanide (yellow prussiate of potash}. If much iron is present, a deep blue precipi- tate will be seen (" Prussian blue ") ; but if there is only a little of this impurity in the acid, the ferrocyanide will only produce a bluish or greenish colour. Nitrates. The salts which nitric acid forms, when its hydrogen is replaced by metals, are called nitrates ; such as potassium nitrate, copper nitrate, etc. Like hydrochloric acid, nitric acid has only one atom of hydrogen in it ; it is on this account called a mono-basic acid. Nitrates are not only produced by dissolving metals in the acid, but also by acting on metallic oxides, hydroxides, or carbonates, with nitric acid. Experiment 163. Place in three separate dishes a little copper oxide, potassium hydroxide, and sodium carbonate. Add a little strong nitric acid to the copper oxide and gently warm it. Notice that the oxide dissolves, giving the same blue solution of copper nitrate as in Exp. 160, but that no brown gas is given off. Dilute some nitric acid, and add it gradually to the potassium hydroxide and the sodium carbonate until each is dissolved. Now slowly evaporate all three solutions down to dryness. Blue crystals of copper nitrate, and white crystals of potassium and sodium nitrates will be obtained. The following are the three equations for the formation of these nitrates (1) CuO + 2HNO 3 = Cu(NO 3 ) 2 + H 2 O (2) KHO + HNO 3 = KNO 3 + H 2 O (3) Na 2 CO 3 + 2HNO 3 = 2NaNO 3 + CO 2 + H 2 O. Experiment 164. Spread a drop or two of the solutions of sodium nitrate and potassium nitrate obtained in the last experi- ment, upon separate pieces of clean, flat glass, and allow the O 194 Nitrogen and its Compounds solution to evaporate by itself. Carefully compare the crystals of the one salt with those of the other, using, if necessary, a pocket lens. The shapes of the crystals are quite different. Those of potassium nitrate are long thin prisms, while the sodium nitrate crystals are more like little cubes ; although not exactly cubes. [Sodium nitrate is sometimes called cubical nitre for this reason.] All nitrates, when strongly heated, are decomposed, and, in most cases, they give off oxygen. They are, therefore, like nitric acid itself, powerful oxidizing agents, and will readily give up oxygen to substances capable of taking it. Experiment 165. Heat a few crystals of potassium nitrate in a test-tube. They first melt, and presently give off gas. Now drop into the test-tube a fragment of charcoal, about the size of a grain of corn. The charcoal takes fire when it touches the hot nitre, and the little fragment dances about on the molten salt. Next drop in a piece of sulphur about the same size, and note how readily it burns in contact with the melted nitre. These three substances, nitre, charcoal, and sulphur, are what gunpowder is composed of. The nitre affords a supply of oxygen sufficient to burn up the carbon and sulphur, so that these materials can burn in places where they cannot get oxygen from the air, such as in the breech of a gun, or even underneath water. Experiment 166.- Weigh out 15 grams of finely powdered nitre, 3 grams of charcoal (also finely powdered), and 2 grams of flowers of sulphur. Mix these together most thoroughly, not with a pestle and mortar, but by putting them all into a sieve, and shaking them through together (a sieve may easily be made by tying a piece of muslin over the mouth of a beaker, the bottom of which, has been broken). Make a little heap with a portion of this mixture, upon a piece of wood, and set fire to it with a match or taper. Notice how it burns not so suddenly as proper gunpowder burns, because the manufacturer is able to get the ingredients much more intimately mixed than in this case. Experiment 167. Make a little paper tube about 7 or 8 centi- metres long (3 inches) by rolling a piece of writing paper closely round a lead-pencil about half a dozen turns. Fasten the edge down with a little strong gum. Slip the pencil out, and stop up one end Tests for Nitrates. 195 of the paper tube with a tiny cork and a little sealing-wax. Now proceed to 'fill the tube with the mixture prepared in the last experiment, putting in a little at a time, and packing it tightly down, using the lead-pencil as a ram-rod. It is now much the same thing as an ordinary " squib," but without the " bang." Now light the open end, and while it is burning plunge it under water, notice that it continues burning, and gives off a quantity of gas which bubbles up through the water. Tests for Nitrates. All nitrates are soluble in water, and the presence of such a salt in a solution can easily be detected by either of the two following tests. Experiment 168. Dissolve a crystal of potassium nitrate in a little water, and to a part of the solution so obtained add some strong sulphuric acid. The nitrate is converted into free nitric acid (just as in the preparation of nitric acid. Exp. 1 56). Then drop into the mixture a fragment or two of metallic copper. As we have learnt, nitric acid acts on copper, and red fumes are evolved ; so that if the liquid containing the copper be now gently warmed, red fumes will be given off, which proves that a nitrate was originally present. The second test is more delicate, and will detect much smaller quantities of a nitrate. Take another portion of the solution con- taining the nitrate, and, as before, add sulphuric acid. Now pour gently down into the test-tube a solution of ferrous sulphate, so that this solution floats on the top of the other. Notice that where the two liquids meet in the tube, there is a dark brown layer formed ; which, if the ferrous sulphate has been added quite carefully so as not to get mixed with the other solution, will appear as a well defined ring. [The nitric acid produced by the addition of sulphuric acid to the nitrate, gives up some of its oxygen to the ferrous sulphate, and a small quantity of one of the oxides of nitrogen is formed. This dissolves in a further portion of ferrous sulphate, forming a dark brown compound.] EPITOME. Nitric acid (aqua fortis) is made by the action of sulphuric acid on potassium or sodium nitrate. On the manufacturing scale sodium nitrate is used (Chili saltpetre), as this salt is cheaper. Nitric acid is produced when electric sparks pass through air in the presence of water, hence it is formed during thunderstorms, and gets washed into the ground with the rain. 195 Nitrogen and its Compounds, Pure nitric acid is a colourless fuming corrosive liquid. It begins to decompose when boiled, giving red fumes of nitrogen peroxide, and oxygen. Nitric acid attacks most metals, but has no action on gold or platinum. The result of the action of nitric acid on metals is either a nitrate or an oxide of the metal, and one or more of the oxides of nitrogen : but hydrogen is never evolved. We can at once distinguish gold from any imitations of this metal, by touching the surface with a drop of nitric acid. If there is no action, the material is gold. Nitric acid is a powerful oxidizing substance, and converts the elements sulphur, phosphorus, and iodine into sulphuric, phosphoric, and iodic acids respectively, with evolution of oxides of nitrogen. The salts of nitric acid are nitrates. Potassium nitrate (nitre, saltpetre] is one of the most important. It is a constituent of gun- powder. When heated, nitrates decompose and give up oxygen ; sometimes forming first of all a nitrite, thus KNO 3 = O + KNO 2 Nitrates of certain metals, when heated, leave an oxide of the metal, and give off nitrogen peroxide and oxygen, thus Pb(NO 3 ) 2 = PbO + 2 NO 2 -f O. CHAPTER XXI. OXIDES OF NITROGEN. THERE are five oxides of nitrogen ; their names and formulae are as follows : (r) Nitrous oxide (laughing gas) ... ... N 2 O (2) Nitric oxide NO (3) Nitrogen trioxide N 2 O 3 (4) Nitrogen peroxide ... ... ... ... N(X (5) Nitrogen pentoxide N 2 O 5 Numbers i, 2, and 5 are acidic oxides. The acid derived from nitrous oxide (called hyponitrous acid) and that from nitrogen trioxide (nitrous acid) are both very unstable com- pounds, and have never been obtained in the free state. The acid derived from nitrogen pentoxide is nitric add. Nitric Oxide, NO. It will be most convenient to study this oxide first. When nitric acid acts on metals, as explained on page 192, the hydrogen which is first displaced from the acid by ihe metal, at once attacks a further portion of nitric acid, depriving it of more or less of its oxygen, and thereby reducing it to one or other of the oxides of nitrogen. The gradual reduction of nitric acid by hydrogen will be seen by the following equations : (1) 2 HNO 3 + 2H = 2 H 2 O + 2NO 2 (2) 2 HN0 3 + 4H = 3H 2 + N 2 O 3 (3) 2 HN0 3 -f 6H = 4H,0 + 2NO (4) 2HNO 3 -f 8H = sH 2 O -f N,O (5) 2HNO 3 + loH = 6H 2 -f N a . 198 Oxides of Nitrogen. The reducing action may even go a step further, and re- sult in the formation of ammonia, thus HN0 3 + 8H = 3 H 2 + NH 3 . Now it depends partly on what metal is being acted on by the nitric acid, partly on the strength of the acid used, partly on the temperature, and partly on the amount of the nitrate of the metal which is produced during the action, which particular oxide will be produced. For instance, when nitric acid is poured upon copper, at first some nitrogen peroxide is formed, then nitric oxide is evolved. After a time nitrous oxide begins to come off. and later on nitrogen is formed. Probably at no single moment is any one oxide alone produced, so that we do not get any one quite free from the others by such an experiment. In many cases, however, we know the particular conditions which will give us the oxide we want in a state which is sufficiently free from the others for ordinary experiment. Thus, in order to get nitric oxide, we use copper and nitric acid, doing the experiment in the following way. Experiment 169. Place a quantity of copper clippings in a two- necked bottle, arranged as in Fig. 39, p. 45. Pour upon the copper a considerable quantity of a mixture consisting of one part of strong nitric acid and two parts of water (by measure). In a few minutes a brisk action sets in, and for some little time the gas which is evolved consists cMefly of nitric oxide. Collect three or four jars of the gas as quickly as possible, and then empty out the apparatus. Avoid breathing the gas, as it is injurious. The reaction which expresses the production of nitric oxide, is that given in equation No. 3 above, where six atoms of hydrogen act on two molecules of nitric acid, the six atoms of hydrogen being displaced from six molecules of nitric acid. The two steps or stages in the reaction may be expressed by the two equations 3 Cu + 6HN0 3 = 6H + 3 Cu(N0 3 ) 2 HNO 3 + 6H = 4 H 2 O + 2NO; Or to express the final results by a single equation 3 Cu 4- 8HNO, = 3 Cu(N0 3 ) 2 + 4H 2 O + 2 NO. Nitric Oxide. 199 As already stated, other reactions resulting in the formation of other oxides of nitrogen are liable to go on at the same time as this one ; hence when chemists want perfectly pure nitric oxide, they ma^e it by other methods. Properties of Nitric Oxide. The samples collected show that it is a colourless gas, which is not appreciably soluble in water. Its most characteristic property is its be- haviour when it comes in contact with the air. Experiment 170. Take one of the jars of gas out of the trough, and for a moment uncover its mouth. Notice that where the gas and air meet, dark red fumes are formed. [This jar may be re- turned to the trough, as it will do for another experiment.] Which of the constituents of the air is the cause of these red fumes ? Experiment 171. Decant some of the gas from one of the jars into a narrow cylinder filled with water and standing in the trough, until the cylinder is about three parts filled with gas. Now bubble up into this a little oxygen [either from another cylinder, or better from a bottle of compressed gas]. Notice that as each bubble of oxygen comes into the nitric oxide, it forms the same red gas as before. Therefore, it was the oxygen in the air which caused it in the former experiment. Notice that, as each addition of oxygen is made, there is a momentary expansion, due to the heat produced by the combination of the two gases ; and also that, after a moment, the volume diminishes and the red gas disappears. The compound that is formed is nitrogen peroxide, NO 2 . This is the red gas, and it is a gas which is quickly dissolved by water. If, therefore, the sample of nitric oxide is quite pure, and oxygen is added cautiously with frequent shaking up with the water, the entire quantity of gas ought to dis- appear altogether, being wholly converted into nitrogen per- oxide, which will dissolve in the water. Since nitric oxide combines with atmospheric oxygen the moment it mixes with the air, of course we cannot tell whether the gas has any smell, as before we could smell it, it is no longer nitric oxide, but nitrogen peroxide; and this gas has a most unpleasant smell and is poisonous. 2OO Oxides of Nitrogen. Next try the action of combustibles on nitric oxide. Experiment 172. Introduce a lighted taper ; note that the gas does not burn, and that it at once puts out the taper. Place a small piece of phosphorus in a deflagrating spoon ; set fire to it, and before it has time to burn up, plunge it into a jar of the gas. The phosphorus will be extinguished. Now withdraw it, and allow it to burn up brightly before putting it into the gas ; note that now it burns in the gas with a bright flame, and forms white fumes, like those produced when it burns in air or oxygen. This shows that nitric oxide itself does not support the combustion of common combustibles, but that if the burning substance is hot enough to decompose the gas into nitrogen and oxygen, then combustion goes on in the oxygen so liberated. Therefore, the product of burning is the same as in the air, but as there is a larger proportion of oxygen in nitric oxide than in air, the combustion is more rapid and brilliant. The composition of nitric oxide is found by taking a measured volume of the gas, and strongly heating in it some metal, such as iron, which can combine with the oxygen, and leave the nitrogen. When this is done, it is found that the volume of the nitrogen left is exactly half the volume of the original gas. This, however, does not tell us anything about the volume of the oxygen in the compound unless we know the density of nitric oxide. When the gas is weighed, it is found to be 15 times as heavy as hydrogen. Two litres of it therefore weigh 30 criths. But as nitric oxide contains half its volume of nitrogen, in these two litres there must be one litre of nitrogen- Now one litre of nitrogen weighs 14 criths, therefore we get 2 litres of nitric oxide, weighing 30 criths contain i litre of nitrogen, weighing 14 leaving 16 criths as the weight of the oxygen ; but 1 6 criths is the weight of i litre of oxygen, therefore 2 litres of nitric oxide contain i litre of nitrogen and i litre of Nitrous Oxide 201 oxygen ; or, in other words, 2 volumes of nitric oxide contain i volume of nitrogen and i volume of oxygen. Nitrogen Peroxide, NO 2 . This is the red-brown gas which is formed when nitric oxide comes in contact with air or oxygen NO + O = NO 2 It is also given off when many nitrates are heated. Experiment 173. Heat a little lead nitrate in a test-tube ; notice that the tube is soon filled with the dark, red-brown gas. We have already seen that this gas is very soluble in water, therefore we cannot collect it over water. Attach a cork and delivery tube to the test-tube, and see if any gas is coming off besides the nitrogen peroxide, by collecting it over water. Note that a colourless gas is collected. Test this gas with a glowing splint ; the gas is oxygen. Therefore, oxygen and nitrogen peroxide are both pro- duced. The equation is Pb(NO 3 ) 2 = PbO + 2NO 2 + O This gas plays an important part in the manufacture of sulphuric acid, as will be explained later on. Nitrous Oxide, N 2 O. This compound is easily pre- pared by simply heating the salt ammonium nitrate. Experiment 174. Place some crystals of ammonium nitrate in a small flask fitted with a cork and delivery tube, and gently heat over a small flame. The crystals quickly melt and give off gas. Collect the gas over water in the pneumatic trough, using water as warm as can be comfortably borne by the hands. Collect several jars full. The equation is the following NH 4 NO 3 = 2H 2 O+ N 2 O. [Compare this reaction with that by which nitrogen was pro- duced by heating ammonium nitrite.] Properties of Nitrous Oxide. From the examples collected, it is seen to be a colourless gas. It has a faint, rather pleasant smell, and a sweetish taste. The gas is moderately soluble in cold water, but less so in warm ; there- fore there is not so much loss of gas if warm water is used for its collection. When nitrous oxide is inhaled for a short time, it causes a kind of intoxication, often accompanied by boisterous 2O2 Oxides of Nitrogen. laughter. On this account it is often called laughing gas. If the inhalation is continued it produces a state of insensibility to pain, and for this reason it is largely used by dentists. 1 Experiment 175. Bring a lighted taper into a jar of nitrous oxide. Notice that the gas behaves like oxygen, for it does not burn, but causes the taper to burn brightly. Dip a glowing splint of wood also into the gas ; the splint is rekindled ; so that from these experiments we could not distinguish this gas from oxygen. Burn a piece of phosphorus in the gas, using a deflagrating spoon as in Exp. 65. The phosphorus burns just as though the gas were oxygen. How, then, can we distinguish between nitrous oxide and oxygen ? Experiment 176. Place a jar of oxygen and one of nitrous oxide side by side. Take a fragment of sulphur on a deflagrating spoon, or on a bundle of asbestos (see Exp. 64), light one corner of it in a gas flame, and before it has time to burn up at all, dip it into the oxygen. Notice that instantly it burns up vividly and continues burning. Now do the same in the other gas. Notice what happens ; the sulphur is extinguished. Repeat this once or twice in the same jar, to be quite sure of it. Now let the sulphur burn up well before plunging it into the gas, and notice that it continues burning when thrust into the nitrous oxide. Nitrous oxide, then, differs from oxygen in that it will extinguish burning sulphur unless the sulphur is thoroughly hot when brought into it. Another and very certain way of distinguishing between nitrous oxide and oxygen, is to mix each of them with nitric oxide. We have already seen what happens when nitric oxide is mixed with oxygen (Exp. 171). But when nitric oxide is mixed with nitrous oxide there is no production of brown fumes. Nitrous oxide and oxygen can also be distinguished by mixing each of them wi'th hydrogen, and exploding them. For instance, if a mixture of equal volumes of oxygen and hydrogen is exploded, we know from what we have learnt of these gases, 1 Students are advised not to experiment upon themselves, or each other, by inhaling the gas. Nitrous Acid and Nitrites. 203 that the hydrogen will combine with half the oxygen to form water, and that half the original volume of oxygen will remain over. But when a mixture consisting of equal volumes of nitrous oxide and hydrogen is exploded, the volume of gas left is the same as the original volume of nitrous oxide, and this gas is found to be nitrogen. This last experiment also teaches us the composition of nitrous oxide. The volume of nitrogen present is equal to the volume of the nitrous oxide, and the oxygen present was just enough to unite with a quantity of hydrogen equal in volume to the nitrous oxide. We know that this amount of oxygen is exactly half the volume of hydrogen, therefore half the volume of the nitrous oxide. In other words, two volumes of nitrous oxide contain two volumes of nitrogen and one volume of oxygen. Nitrous Acid and Nitrites. Nitrous acid has never been obtained in a pure state, and even very dilute solutions of it quickly decompose. The formula for the acid is HNOa } it has one atom less oxygen than nitric acid, and it can readily take up oxygen from compounds which are rich in oxygen, and pass into nitric acid. Experiment 177. Dissolve a few particles of sodium nitrite in half a test-tube of water, and add to it a few drops of sulphuric acid. This decomposes the sodium nitrite, forming sodium sulphate and nitrous acid, which can exist for a short time in the dilute solution. Now add to this some solution of potassium perman- ganate (a salt very rich in oxygen), and notice how the purple colour of the permanganate is instantly destroyed, owing to the nitrous acid depriving it of some of its oxygen. On the other hand, nitrous acid will give up some of its own oxygen to many substances which are ready to unite with oxygen. The salts of this acid are called nitrites, as sodium nitrite, potassium nitrite, etc. Potassium nitrite can easily be made by gently melting potassium nitrate, when the nitrate parts with some oxygen and leaves the nitrite KNO, = KNO 2 + O. We can easily tell a nitrite from a nitrate by adding dilute sulphuric acid to each. The nitrite at once evolves brown fumes, the nitrate does not. 2O4 Oxides of Nitrogen. EPITOME. Five oxides of nitrogen are known. 1. Nitrons oxide (laughing gas), N 2 O, is obtained by heating ammonium nitrate. It is a colourless, slightly sweet tasting gas. Does not burn, but supports combustion almost as well as oxygen. Distinguished from oxygen by (a) much greater solubility in water, (<$) gives no red fumes when mixed with nitric oxide. Nitrous oxide is used by dentists to produce insensibility to pain. It is called laughing gas, because when a little of it mixed with air is inhaled it produces hysterical laughter. 2. Nitric oxide, NO, is prepared by acting on copper with nitric acid. Other oxides of nitrogen are produced at the same time, although at a certain stage of the action the bulk of the gas given off is nitric oxide. It is a colourless gas which instantly combines with free oxygen to form nitrogen peroxide. Therefore, when it comes into the air it at once forms brown fumes of the peroxide. Nitric oxide extinguishes a taper, because the flame is not hot enough to decompose the gas. Even phosphorus, unless strongly burning, is extinguished ; but if burning vigorously when introduced into the gas, the flame is hot enough to decompose nitric oxide into oxygen and nitrogen, when, of course, the phosphorus will burn. Nitric oxide is distinguished from all other gases by forming red fumes in contact with air or oxygen. 3. Nitrogen trioxide, N 2 O 3 , does not exist as a gas. When a mixture of nitric oxide, NO, and nitrogen peroxide, NO 2 , in equal volumes, is passed through a strongly cooled tube, a blue liquid condenses, which is believed to be nitrogen trioxide. 4. Nitrogen peroxide, NO 2 . This gas is formed by heating lead nitrate. When the gas is passed through a cooled tube, it condenses to a yellowish liquid, which is condensed nitrogen peroxide. The liquid gives off the familiar reddish brown gas. When the gas is heated, it rapidly darkens in colour. This change is due to the molecules breaking up into simpler ones. At low temperatures nitrogen peroxide has the composition N 2 O 4 , its density being 46 ; but as the temperature gradually rises, the gas breaks up into molecules having the composition NO 2 , and at 140 its density is 23. Therefore nitrogen peroxide has two formulas, N 2 O 4 at low temperatures, and NO 2 at high temperatures. Nitrogen peroxide dissolves in water, therefore can only be collected by displacement. Oxides of Nitrogen. 205 5. Nitrogen pentoxide, N^O^, is obtained by withdrawing the elements of water from pure nitric acid by means of phosphorus pentoxide. 2HNO 3 - H 2 O = N 2 O 5 . It is a white crystalline substance which cannot be preserved. It combines with water with great eagerness, forming nitric acid. Nitrous acid, HNO 2 , is not known in a pure state. Its salts are nitrites. CHAPTER XXII. OZONE. WHEN we hear of " Smith alias Sampson " getting into trouble, we know at once that the two names stand for the same person. Very often when a man assumes an alias he alters his outward appearance as far as he can in order to make it more difficult to recognize him. Perhaps from being a fair and beardless person, he now appears as a dark man with a black beard ; having a scar on his face, and perhaps even a squint. Now some of the elements are capable of doing something of the same kind. Under certain circumstances they are able to assume an alias t as it were. They adopt new outward appearances and fresh properties so different from what they usually have, that they are then altogether unlike their original selves, and are quite likely to be mistaken for some entirely different element Oxygen is one of the elements which can do this. With the usual properties of this element we are now quite familiar; but when it assumes its alias, we find first that it adopts a powerful and unpleasant smell. Also it develops a habit of attacking organic substances in a violent manner, so that it cannot even be passed through a piece of indiarubber pipe without instantly destroying it. Also, it attacks metals like silver and mercury, which it took no notice of before ; and if it comes into contact with potassium iodide it instantly seizes it, steals the potassium from it, and turns the iodine adrift. We see, therefore, that when oxygen adopts these new hab*its and properties, it seems like an entirely different sub- stance, although all the time it is nothing else but oxygen. Ozone. 207 Phosphorus is another element which can do the same thing. As usually seen, phosphorus is a wax-like solid, very poisonous, easily cut with a knife, melts when warmed by the hand, and takes fire so easily that it has to be kept beneath water, and care has to be taken in handling it. When it assumes its alias, however, it becomes a dark red substance, which looks like chocolate; it is no longer poisonous, does not melt, and requires to be strongly heated to make it burn. All the time it is phosphorus, and nothing else but phosphorus, but yet so unlike the element in its ordinary state. The word that is used for this curious property is allotropy (meaning " other form "). We say that the element when appearing in its alias, or its more unusual character, is an allotropic modification, or more shortly, an allotrope of that element. Thus, when phosphorus appears as the red substance, we call this the allotropic modification of phosphorus, and when oxygen is in the condition in which it shows such active properties, we speak of it as allotropic oxygen, or ozone (ozone means *' a smell "). We can make oxygen assume this allotropic state in several ways. The best method, and the one which gives the largest amount of ozone, is to expose oxygen to a particular kind of electrification, known as the silent discharge. Experiment 178. Take a piece of narrow glass tube about 30 centimetres (12 inches) long and wind a spiral of thin copper wire round the entire length outside : about three turns of wire to each centimetre. Attach a piece of indiarubber tube to one end, so as to conduct a stream of oxygen through. Pierce a hole with a pin through the FIG. 90. rubber just beyond the end of the glass tube, and push a thin copper wire through it, so that the wire reaches nearly to the other end of the glass tube, as shown in Fig. 90. The inside and outside wires are then connected to a Ruhm- korfs coil. When the coil is set in action, there will be no visible 208 Ozone. spark jumping across between the wires, but a vast number of tiny sparks passing from the whole length of one wire to the other. If now a slow stream of oxygen be passed through the tube, it has to " run the gauntlet " of this host of little sparks, and this causes a part of the oxygen to change itself into ozone. Why it does this, nobody knows. Test for Ozone. We test for ozone by making use of the property it has of setting free the iodine contained in potassium iodide ; and in order that we may see when the iodine is so liberated, we have starch present. The instant the iodine is set free it unites with the starch and produces the blue com- pound described on page 134. Experiment 179. Make a little thin starch paste (see p. 134), and add to it a little potassium iodide solution. Cut strips of white paper and dip them into this mixture, and hang them up to dry. These are called ozone test papers. Moisten such a test paper and hold it at the end of the tube, where ozone is being formed in the last experiment. Notice that the paper is instantly turned blue. This means that ozone has de- composed some of the potassium iodide, and that the free iodine has then united with the starch. Notice the curious smell of the ozone which passes out of the tube. Ozone is quickly transformed back into ordinary oxygen by- being heated. Experiment 180. Attach to the end of the "ozone tube" a short straight tube, by means of a rather wider tube and two corks, FIG. 91. as shown in Fig. 91. (The usual indiarubber joint cannot be used, for the reason stated above.) Pass the ozonized oxygen through, Ozone. 209 and at the same time heat the tube. The ozone will be changed into ordinary oxygen, and if the gas passing out is tested with a piece of ozone test paper, there will be no blue colour produced. Chemists now know that the change which oxygen undergoes in passing from oxygen into ozone, is that the molecule of oxygen takes up another atom of oxygen. Molecules of oxygen consist of two atoms, while molecules of ozone are composed of three atoms of the same element. Therefore, when oxygen changes into ozone there is a contraction in the volume, and vice versA, when ozone passes back into ordinary oxygen there is an expansion in the volume. Ozone is on this account sometimes spoken of as " condensed oxygen." This is quite true in one sense, because the molecule of ozone is more condensed than the molecule of oxygen ; but it must be remembered that this is quite a different thing from condensing or compressing oxygen. We cannot condense oxygen into ozone merely by compressing the former so as to reduce the volume. Ozone is present in small quantities in country air. It is produced by lightning discharges. It is also formed in small quantities when certain substances combine with oxygen without actually burning. Thus, if a piece of phosphorus is placed on the table, it is seen to " smoke." It is really oxidizing rather quickly, but is not actually burning. Now the phosphorus, in thus combining with atmospheric oxygen, causes a little of the oxygen to go into the allotropic form. How and why it does so, is not known. Experiment 181. Take a short stick of phosphorus (if it is coated over with a white or greyish film, scrape it clean, underneath water] and place it in a good large bottle, having a small layer of water on the bottom just enough to half bury the phosphorus. Cover the bottle with a piece of cardboard and leave it for about ten minutes. Then dip into the bottle a strip of moist " ozone paper," and note that it shows the presence of ozone. CHAPTER XXIII. CARBON. Occurrence. Carbon is an example of an element which can assume three allotropic forms. In the first it appears as a soft dull-black solid, which has no crystalline shape. 1 Charcoal is the most familiar example of this form of carbon. The second variety is also soft and black, but is bright and shiny, almost like steel, and has a crystalline form. Graphite, or " black-lead," is the name of this variety of carbon. The third modification is extremely hard ; sometimes per- fectly colourless and transparent, and highly crystalline. The name of this allotropic form of carbon is Diamond. Each of these three allotropes of carbon is found in nature. Carbon also occurs in nature in a state of chemical com- bination with other elements. For example, one of its com- pounds with oxygen, namely, the gas carbon dioxide, is present in the air, and is sent out in large volumes from rents in the rocks in volcanic districts. Again, one of its compounds with hydrogen, the gas known as marsh gas, or fire-damp, is found in large quantities in coal mines, and is given out by rotting vegetable matter in marshy places (hence its name, " marsh gas"). Carbon is also a constituent of all animal and vegetable matters ; therefore meat, wool, bread, sugar, alcohol, wood, all contain carbon, and in many cases the carbon is associated simply with hydrogen and oxygen. 1 Substances which show no crystalline form are called amorphous bodies that is, without form. Carbon. 211 Carbon is also present as a constituent of a large number of the common minerals which help to make up the solid earth. Thus, chalk, marble, limestone, dolomite, are all extremely common and abundant minerals, and they all contain carbon, associated with oxygen and lime, as carbonate of lime. How to obtain Carbon from its Compounds. When a piece of meat or bread or wood is partly burnt, we say that it is charred. This means that some of the hydrogen and oxygen have been expelled as water, and a portion of the carbon has been set free. Experiment 182. Place a little dry powdered starch in a hard glass tube, the end of which has been sealed up (see p. 37), then draw out the tube and bend it as shown in Fig. 92. This constitutes a small retort. Now heat the starch, and no- tice that it chars or blackens, and at the same time a liquid is expelled. This condenses in the drawn-out part of the tube and can be collected in a small test-tube. This liquid is chiefly water, resulting from the decomposition of the starch ; and the blackened mass contains free carbon. To identify the liquid as water, it will be sufficient here to FlG> 92< drop upon it a small fragment of potassium. If the metal takes fire (as in Exp. 44) we may safely conclude that the liquid is water. The process here illustrated is that of destructive distillation (see p. 179), and whenever compounds containing carbon are submitted to this treatment, carbon is set free in a more or less pure state, depending on circumstances. Another way by which we can get carbon out of some of its compounds, is by withdrawing the other elements with which it is associated, not by fire, but by the use of some chemical reagent. 212 Carbon For example, sugar is a compound of carbon with oxygen and hydrogen, and we can throw the carbon out of combina- tion by acting on the sugar with sulphuric acid. Experiment 183. Roughly weigh out 12 grams of loaf sugar ; place it in a good-sized beaker and pour over it 10 cc. of warm water. This will dissolve the sugar in a little while, especially if gently warmed, and give a strong syrup. Now pour into this, all at once, 12 cc. of strong sulphuric acid ; the whole mixture at once froths right up. The acid takes from the sugar the hydrogen and oxygen, and leaves the carbon, which will appear as a black spongy mass. Alcohol, or "spirits of wine," is another compound of carbon with hydrogen and oxygen, but it has a great deal more hydrogen in proportion to carbon than sugar has. Let us try and get the carbon out of some alcohol. Experiment 184. Pour 15 cc. of strong sulphuric acid into a little flask, and add 5 cc. of water. Cool it by dipping the flask into cold water. Then add 5 cc. of pure alcohol. [If pure alcohol is not available, use methylated spirit.] Provide a cork and delivery- tube to the flask, and arrange to collect gas in the water trough. Gently heat the flask, gas soon begins to come off. Collect first a small jar full, and allow the rest to collect in a tall cylinder until it is about one-third filled with the gas. [If pure alcohol has been used, there will be very little blackening of the mixture in the flask, but methylated spirit contains other carbon compounds which very quickly char, or carbonize when treated in this way.] The gas we have collected is called ethylene. It contains some of the carbon from the alcohol, combined with some of the hydrogen ; so that by this experiment we have not yet got free carbon, but have only expelled it from the alcohol still combined with hydrogen. The gas is, therefore, a hydro- carbon, as it is called ; that is, simply a compound of carbon and hydrogen. Hydro-carbons, as a rule, burn with a flame which gives a good light. Experiment 185. Remove the small jar of gas and bring a lighted taper to it. Notice the kind of flame it burns with. Con- trast this flame with that of burning hydrogen. Diamond. 2 1 3 We have learnt that chlorine has the power of taking hydrogen away from compounds of carbon and hydrogen, for in Exp. 1 1 6, when turpentine (which is also a hydro-carbon) was brought into chlorine, carbon was set free. Let us, there- fore, try and get the carbon out of the gas we have obtained from alcohol, by acting on it with chlorine. Experiment 1 86. Collect a quantity of chlorine in the long cylinder containing the gas obtained in Exp. 184, until the cylinder is nearly but not quite full. Slip a glass plate over the mouth of the cylinder and shake up the remaining water in it so as to mix the gases as well as possible. Then apply a lighted taper. The mixture burns with a curious flame, producing a dense black smoke, and a black deposit of carbon all down the sides of the vessel. The chlorine has combined with the hydrogen and set the carbon free. This carbon came, therefore, originally out of the alcohol. In these experiments it is to be noted that the carbon is always obtained in the first allotropic modification, that is, as a soft black non-crystalline substance, like charcoal or soot. It is very much more difficult to obtain carbon in the second form, as graphite ; and to make it pass into the third or diamond variety is a task which has baffled almost every attempt. Diamond. This form of carbon, although the most valuable in one sense, is by far the most useless so far as the chemist is concerned. It is found in gravel deposits in India, Africa, and Australia. When found, diamonds are not at all like the gems. They more often look like common little rough stones, scarcely transparent, and with little appearance of being crystals. To obtain them as usually seen in the gem, they are ground or u cut" to the desired shape, so as to " sparkle" in the light. Diamond is the hardest known substance, and will scratch all other stones. It is, therefore, employed for cutting glass, and for giving a hard edge to drills and rock-borers. Some diamonds are brownish, and even black. These are valueless for gems, and are used for drilling and also for grinding and polishing the clear ones. When strongly heated in oxygen, the diamond first turns black and then burns, giving carbon 2 1 4 Carbon. dioxide as the only product This proves that it is pure carbon. Rock crystal or quartz is sometimes cut to look like diamonds. It is easy to distinguish between them by the fact that diamond burns in oxygen and gives carbon dioxide, whereas quartz does not burn at all. Quartz is not nearly so hard as diamond. Graphite. This form of carbon is not nearly so rare as diamond. It is found in great quantity in California. It can be made by dissolving charcoal in melted iron (just as salt is dissolved in water) and then allowing the iron to cool As it cools carbon crystallizes out, and deposits in the form of graphite. Long before it was known that graphite was simply carbon, it was employed for writing purposes. It is so soft that when drawn across paper it leaves a black mark. Hence it got the name plumbago, or " the writing lead," and it is also known under the common name of black-lead. It must be remembered, however, that these names were given to it simply because it looked rather likz lead. Graphite is used for making " lead " pencils, and for " black-leading " iron work. When strongly heated in oxygen, graphite, like diamond, burns, and gives the same oxide of carbon. Amorphous Carbon. This includes common wood- charcoal, coke, gas-carbon, lampblack, soot, animal- charcoal or boneblack. All these substances consist of non-crystalline carbon, more or less impure. Charcoal is made by subjecting billets of wood to the process of destructive distillation (just as the starch was treated in Exp. 182) in ovens or retorts, so as to collect the volatile products as well as to secure the charcoal : or else it is made by piling the wood into stacks, and setting fire to the interior ; the outside being so covered over with turf or earth as to prevent much air getting to the smouldering heap. By this method everything is lost except the charcoal. Care has to be taken to regulate the supply of air, for if too much gets into the heap the charcoal begins to burn away and is also lost. Charcoal floats when thrown on water, but it is not really Charcoal. 215 lighter than water ; it only floats because its pores are full of air,' and this buoys it up. Experiment 187. Tie a little weight to a piece of charcoal so as to sink it, and throw it into some water in a test-tube, and then heat the water. Notice bubbles of air coming out of the charcoal. After a few minutes' boiling, allow the water and charcoal to cool, and it will be found that it will no longer float when the weight is taken off, but at once sinks in the water. Charcoal, like both the other forms of carbon, burns in oxygen (see Exp. 59), and produces exactly the same oxide of carbon. Charcoal is very much easier to burn, however, than either diamond or graphite, and can be used to make an FIG. 93. ordinary fire with. It would be quite impossible to light a fire with graphite or diamonds. Charcoal has a wonderful power of absorbing gases, and is on this account employed to arrest the bad smelling gases arising from drains and other places. Experiment 188. Fit a large bottle with a cork and two glass tubes, as shown in Fig. 93, B. Break up some charcoal into little pieces and after heating them for a few minutes in a metal dish or tray, fill a piece of combustion tube with the charcoal and fit a cork and exit tube into each end. Now place a few small particles of ferrous sulphide in a test- tube, and add a little dilute sulphuric acid. Remove the cork 216 Carbon. from the bottle B, and, by means of a piece of string, lower the test-tube for a few seconds into the bottle ; then draw it out and replace the cork. Attach the tube W to a water tap, and place a pinch-cock on G. In this way a small quantity of a gas, whose offensive smell will have been noticed, has been put into the air in the bottle B. This gas is called sulphuretted hydrogen, and although its smell is quite enough to recognize it by, we can use a more convenient test. Dip a piece of paper into a solution of lead acetate (sometimes called sugar of lead}, and let a little of the air in the bottle blow against it, by allowing water to enter through tube W, and opening the pinch-cock on tube G. Notice that the paper is stained black by the gas. Now connect G to one end of the tube containing the charcoal, and gently open the pinch-cock so as to drive the air in the bottle slowly through the charcoal. Test the air as it passes out of the tube with another piece of paper moistened with acetate of lead. It no longer is blackened : neither will any smell of sulphuretted hydrogen be detected. The charcoal has absorbed all the bad smelling gas. No variety of wood charcoal is quite pure carbon. For instance, if we burn a piece of charcoal there is always a certain amount of white ash left. This of course is not carbon ; but besides this, there is always a certain amount of some compounds of carbon with hydrogen. Coke is made by treating coal very much as wood is treated in making charcoal. Coke stands very much in the same relation to coal as charcoal does to wood. Large quantities of it are produced in the manufacture of ordinary coal gas, when coal is heated in large fireclay retorts. Coke is often manufactured specially for the coke itself, either by burning coal in great stacks (like the charcoal stacks), when all the gases and liquids produced at the same time are lost ; or in special ovens or kilns, when these are caught. Coke, like charcoal, contains a certain amount of mineral matter, which is left as ash when the coke is burnt; and it also contains a little hydrogen. It is much harder and heavier than charcoal, and not so easily lighted. A nimal- Charcoal. 2 1 7 Gas-Carbon is a still harder sort of coke, which is found lining the retorts in which coal-gas is manufactured. It is even harder and heavier than ordinary coke, and conducts electricity extremely well. This is the form of carbon which is employed for making the carbon rods used in electric lights. Lampblack is the soot obtained by burning substances like petroleum or tar, which give off a large quantity of smoke. It is chiefly used for printers' ink, and for black paint. Animal- Charcoal is the name given to the substance obtained by charring bones in iron retorts (just as wood or coal is treated). It is a very impure product, containing only about 10 parts of carbon in 100 parts, the rest being the mineral matters of the bone (chiefly phosphate of lime). When this is ground up fine it is called boneblack. This substance is used chiefly for removing the dirty colour from sugar syrup, in making white sugar. It has a great power of absorbing colouring matter, and is better in this respect than any other kind of charcoal. Experiment 189. Half fill a wide-mouthed stoppered bottle with water, which has been tinted with either magenta, aniline blue, indigo, or some other such colouring matter. Add to this a spoonful of fine boneblack, and shake it well for a moment or two. Then filter the liquid, and note that what passes through the filter paper is quite free from colour. All the varieties of amorphous carbon burn in oxygen and give carbon dioxide, the same product as graphite or diamond yield. Carbon is an element which does not readily enter into chemical union with other elements. Thus, at ordinary tem- peratures, carbon is not acted on by hydrogen, oxygen, chlorine, or nitrogen. On this account, wood which has been charred on the surface is not so soon rotted when buried in the ground, as wood which has not been so protected. Hence it is usual to char the end of stakes or posts before putting them into the ground. Coal is the product of a process of natural decomposition of vegetable matter which has taken long ages to complete. 218 Carbon. The coal we burn to-day was once living vegetation. It has long been buried, owing to geological changes having taken place, and has been subjected to pressure. Coal is a very impure form of carbon, and contains compounds of carbon with hydrogen and oxygen. Roughly speaking, the proportion of carbon in soft or bituminous coal is about 80 in 100 parts; while in the hard or anthracite varieties it is about 90 parts per 100. Although carbon is not acted on by oxygen at ordinary temperatures, it unites with it very readily when heated. Not only will it combine with free oxygen, but it will easily take oxygen away from certain oxides. On this account carbon is a most useful element to the metallurgist, for by heating oxides of metals with carbon, the carbon takes the oxygen and leaves the metal in the uncombined state. Experiment 190. Take a small quantity of red-lead, Pb 3 O 4 , and mix it thoroughly with about one quarter as much finely powdered charcoal. Put the mixture in a small porcelain crucible and strongly heat it. After a short time throw out the contents of the crucible, and there will be a globule of bright metallic lead. This process is called reduction. We say that the lead has been reduced by the charcoal. This simply means that the oxygen with which the lead was combined has been taken from it by the carbon, and the lead, therefore, was left by itself. This is how many of the metals are obtained from their ores. Iron ores, for instance, which are used for getting iron from, are oxides of iron. These when very strongly heated with carbon (usually coal or coke) are deprived of their oxygen, and the iron is set free. All metals cannot be reduced from their oxides in this way. EPITOME. Carbon is an element known in three allotropic forms, (i) Diamond. (2) Graphite. (3) Charcoal. Diamond is the hardest known substance ; highly crystalline, and often nearly or quite colourless. Used as a gem, and for cutting purposes. Carbon, 219 Graphite is soft, black, shining, and crystalline. Used for " lead " pencils, for " black-leading." Charcoal is soft, black, dull, and non-crystalline. Used for fuel and for making gunpowder. Coke, lampblack, soot, bone black or animal-charcoal are all more or less impure forms of amorphous carbon. All the varieties of carbon burn in oxygen and give the same compound, namely, carbon dioxide. In this way diamond is easily distinguished from quartz (which is silicon dioxide, and will not burn) or other imitations. Carbon in combination occurs in all organic substances, and when these are charred, or heated so that air does not get to them to cause them to burn away altogether, the carbon is left in a more or less pure state. Wood-charcoal, animal-charcoal, and coke, are obtained by heating wood, bones, and coal in this manner. Charcoal absorbs gases readily, and is therefore used to remove bad smelling gases ; wood-charcoal does it best. Charcoal also absorbs colouring matter from liquids which are filtered through it or shaken up with it. Animal-charcoal does this best ; it is therefore used for filters. Coal is an impure form of carbon, containing a number of carbon compounds. Charcoal, coke, and coal are used as reducing agents, for taking oxygen away from oxides of certain metals like iron, copper, lead, etc. When such oxides are strongly heated with carbon, the metal they contain is reduced. This is the principle of the smelting of many of the ores of metals. CHAPTER XXIV. CARBON DIOXIDE, CARBONIC ACID, CARBONATES. Carbon Dioxide, CO 2 . When a piece of charcoal is burnt in oxygen, as in Exp. 63, the carbon and oxygen combine and form the compound called carbon dioxide. Let us repeat that experiment in a slightly different way, so as to collect the compound and examine it. Experiment 191. Place a few little pieces of charcoal in a short piece of combustion tube, through which a stream of oxygen from FIG. 94 . a gasholder can be passed ; attach a cork and delivery tube to one end, and arrange to collect the gas over water, as shown in Fig. 94. Now heat the charcoal by means of a Bunsen flame and allow it to burn in the stream of oxygen, which should be regulated so as just to keep the charcoal burning brightly. Collect two jars of the gas. Into one jar dip a lighted taper. Notice that the flame is instantly put out, and also that the gas does not burn. In these two respects it is like nitrogen. Carbon Dioxide. 221 Now pour into another jar some clear lime water, 1 and shake it up. Note that the lime water at once becomes milky. This action of carbon dioxide on lime water is a test by which we can distinguish this gas from all others, as it is the only gas which can produce this effect. The equation is CaH 2 O 2 Lime water. co c = CaCO 3 -f Calcium carbonate. FIG. 95. Not only is carbon dioxide produced when charcoal is burnt in oxygen, but when any carbon compound is burnt in the air. Experiment 192. Set fire to a piece of paper, and drop it while burning into a dry bottle, and cover the mouth with a piece of card or glass plate. Then pour in a little lime water, and shake it up. Experiment 193. Hold a dry jar over a candle flame for a minute or two, so as to catch some of the invisible products of its burning, as in Fig. 95. Then add lime water, and shake up. Do the same with a spirit lamp flame, and with an ordinary coal-gas 1 To make lime water, place a handful of powdered lime in a big bottle, and fill it up with water, cork it up, and after shaking for a minute, leave it to settle. Then pour off the clear liquid into a separate bottle for use. The other bottle can be filled up with water over and over again, 222 Oxides of Carbon. flame. Notice that in every case there is abundance of carbon dioxide being produced. Roughly speaking, every ton of coal that is burnt, produces about three tons of carbon dioxide, all of which escapes into the atmosphere. We have already learnt that respiration is a sort of com- bustion, let us see if carbon dioxide is produced. Experiment 194. Place some lime water in a large test-tube, and by means of a glass tube make the breath bubble through the solution. Notice that the first portions of breath produce hardly any milkiness, but presently, when air that has been further into the lungs is blown through, the liquid quickly shows that there is carbon dioxide. (See also Exp. 75.) Carbon dioxide is also given off during many processes of decay and fermentation. Thus, when vegetable matter (such as the leaves of trees which fall in autumn) gradually rots away, carbon dioxide is given off. Experiment 195. Dissolve a handful of sugar (common brown moist sugar is best) in some tepid water in a good-sized flask, and add some yeast. Fit a cork and delivery-tube to the flask. The sugar soon begins to ferment and gas escapes. Collect some of the gas at the water trough and test it with lime water. In breweries enormous quantities of carbon dioxide are in this way produced. Since these different operations, namely, the burning of ordinary fuel, respiration of man and animals, decay and fermentation, besides others which are constantly going on, all result in sending carbon dioxide into the air, it is no wonder that this compound is present in the atmosphere. The wonder rather seems that the air does not become too impure to breathe. But, strange to say, if we test a small sample of ordinary air by means of lime water, we find that we can scarcely even detect the presence of this gas at all. Experiment 196. Take a bottle (such as has been used for testing in the above experiments) which contains simply ordinary air. Pour some lime water in and shake it up. No turbidity of the lime water is noticeable. Of course, if this experiment is made in a small room where a number of persons have been present for Carbon Dioxide. 223 some time, and several gas-lamps have been burning, most likely the air will contain enough carbon dioxide to cause a turbidity in lime water when a small sample of it is tested in this way. Since in the open air there is, after all, so small an amount of carbon dioxide (only about 3 parts in 10,000 parts of air, by measure), in spite of the enormous quantities which are every day being sent into the atmosphere think of the millions of people in London alone, and the hundreds of tons of coal daily burnt ! it is evident that there must be some natural process at work which constantly removes this compound from the atmosphere. Nature, as usual, has made a beautiful provision for preventing the accumulation of carbon dioxide, and her secret for doing this has long been found out. She has endowed the green parts of plants with the power of decomposing carbon dioxide, by the aid of sunlight, into its two component elements, carbon and oxygen. Every green leaf and every blade of grass is a tiny chemical laboratory, where oxygen is being made from carbon dioxide, and is ever being returned to the atmosphere with all its life-support- ing properties. The carbon which comes out of the carbon dioxide is utilized by the plant. Vegetation, therefore, removes dioxide from the atmosphere, and at the same time gives back the oxygen which had been taken from it by the carbon. To prepare Carbon Dioxide for experiments, we usually adopt quite a different method. Experiment 197. Take the apparatus shown in Fig. 39, p. 48, and put into the bottle a quantity of broken-up marble and some water. Now pour a little strong hydrochloric acid down the funnel, and note that effervescence at once begins. After a minute or two, when the air is all swept out, collect the gas which is given off. Marble is a variety of calcium carbonate, and the change here going on is the following CaCO 3 + 2HC1 = CaCl 2 + H 2 O + CO* Calcium carbonate. Calcium chloride. Marble is used because it is one of the most convenient of 224 Oxides of Carbon. all the carbonates ; but we can equally well obtain the gas by acting on any carbonate with almost any common acid. Experiment 198. Break up some sodium carbonate (common " washing soda ") and put a little into three small beakers. Stand each of the beakers in a larger beaker, or jar, and then by means of a pipette, add to one a few drops of dilute sulphuric acid ; to the next, dilute hydrochloric acid ; and to the third, dilute nitric acid. Cover the outer jars with pieces of paper. The little beakers can now be lifted out with tongs, and lime water poured into each jar. In each case the lime water will become turbid. In practice we cannot use sulphuric acid with calcium carbonate, unless the carbonate is first powdered up very fine and made into a paste with water. Experiment 199. Put a few lumps of marble in a test-tube with some water, and add a little strong sulphuric acid. Notice at first there is an effervescence, but that it very soon leaves off, and no more gas comes off. The first action, when it is effervescing, is this CaC0 3 + H 2 SO 4 = CaS0 4 + H 2 O + CO 2 . Calcium sulphate. Calcium sulphate is formed; this is the same as "plaster of Paris," and it at once coats over the lump of marble and prevents any more acid getting to it. Carbon dioxide is a very heavy gas. It is about i times heavier than air, so that it is quite easy to collect a jar of it by " downward displacement." More than this, the gas can even be poured from one jar to another like a liquid ; and if poured into a vessel suspended on a balance, it will weigh down that end of the beam. Experiment 200. Take a long light wooden rod, and hang from one end of it a light cardboard box (or an old hat), and place upon the other end a piece of bent lead rather lighter than the box. Balance this upon the edge of a paper-knife or other convenient metal edge, in the manner shown in Fig. 96. Then bring a good large jar of carbon dioxide, and pour it into the box or hat, and notice that it weighs it quite down. Carbon dioxide is soluble to a small extent in water, but not to such a degree as to make any difference when we Carbon Dioxide. 225 collect the gas at the water trough. All natural waters contain a little carbon dioxide dissolved in them, and some waters contain quite a large quantity. For instance, Seltzer water and Apollinaris water contain so much of this gas dissolved in them that they actually effervesce or " sparkle," owing to the escape of the carbon dioxide. Under ordinary conditions water dissolves about its own volume of this gas ; but under increased pressure it can take up more. When such extra pressure is again removed, the extra gas that was dissolved is given off. This is illustrated in the ordinary aerated waters used for drinking. Gas is pumped into the water under great FIG. 96. pressure so that the water dissolves a large quantity, but the moment the pressure is released by drawing the cork, then the gas rapidly escapes with the familiar effervescence. Experiment 201. Colour some ordinary water with a few drops of litmus, and let carbon dioxide bubble through it for a few minutes from the apparatus used in Exp. 197. Notice that the litmus turns red, showing the presence of an acid. Note that the colour is not quite so bright red as when a drop of hydrochloric or dilute sulphuric acid is added. It is because the solution of this gas in water is acid that the gas is sometimes called " carbonic acid gas." The acid so obtained is a very feeole acid, and easily decomposes. It is called carbonic acid, H 2 CO 3 . If it is gently heated it is decomposed again into water and carbon dioxide. H 2 CO 3 = H 2 + C0 2 . Q 226 Oxides of Carbon. Experiment 202. Take some of the solution used in the last experiment and heat it in a test-tube. Notice that the reddish colour of the litmus soon changes back again to the original blue, as the feeble acid is being decomposed. Carbonates. Although real carbonic acid is only a feeble acid, and so easily decomposed that it cannot be obtained by itself, it forms important salts called carbonates. Calcium carbonate, CaCO 3 , occurs as marble, limestone, chalk. This is the compound that is formed when lime water is brought into carbon dioxide. Experiment 203. Pass some carbon dioxide from the gene- rating apparatus (Exp. 197) into some lime water until there is a thick milkiness. Filter the liquid through a small filter. The white deposit will hardly be visible on the paper, but we can test it in the following way. Put the paper into a small beaker, and pour upon it a few drops of hydrochloric acid. Notice effervescence. Put a little lime water in a wide test-tube, and pour some of the gas out of the beaker into it and shake it up. The equation which shows what takes place when carbon dioxide is passed into lime water is Ca(HO) 2 + CO 2 = CaCO 3 4- H 2 O. Lime water or calcium Calcium carbonate hydroxide. or chalk. so that we get back the same compound, calcium carbonate, which was used to make the carbon dioxide from. Experiinent 204. Pass carbon dioxide through a solution of caustic soda. Notice that there is no turbidity ; also observe that the gas is eagerly absorbed, because bubbles hardly rise to the top of the liquid. Continue bubbling the gas into the solution until no more is absorbed, and then gently evaporate the solution. Take a little of the residue and add hydrochloric acid to it. Observe the effervescence. Test the gas with lime water. Take another portion of the residue and add water to it ; it dissolves easily. This is the reason why there was no precipitate formed when the gas was passed into caustic soda ; the sodium carbonate that was formed is soluble in water. The equation in this case is 2 NaHO + CO 2 = Na,CO 3 + H a O. Caustic soda or Sodium carbonate, sodium hydroxide Carbonates. 227 There are two sodium carbonates. One is common "washing soda," and the other is usually called sodium bi- carbonate. The difference in these salts is that the first contains twice as much sodium as the last ; or, in other words, the last contains twice as much carbonic acid in proportion to sodium as the first. Their formulae are Na 2 CO 3 , sodium carbonate, or normal sodium carbonate. HNaCO 3 , sodium bi-carbonate, or hydrogen sodium carbonate. Experiment 205. Dilute some lime water with about half as much distilled water, and pass carbon dioxide through it. As in the former experiment there is a precipitate of calcium carbonate. FIG. 97. But continue passing the gas, stirring up the liquid, and notice that in a few minutes it becomes perfectly clear again. The calcium carbonate has dissolved in the solution of carbon dioxide. Take a little of this clear solution and boil it. It again becomes turbid, because the carbon dioxide is expelled, and the calcium carbonate cannot remain in solution. The presence of calcium carbonate dissolved in this way is what causes the temporary hard- ness of natural waters (see pp. 82 and 269). Many carbonates part with carbon dioxide when they are heated, leaving an oxide. Experiment 206. Heat a little magnesium carbonate in a test- tube, and allow the gas to flow down into another test-tube, con- taining lime water, in the manner indicated in Fig. 97, MgCO 3 = CO 3 + MgO. 228 Oxides of Carbon. Calcium carbonate undergoes the same change, but re- quires a higher temperature. The process of lime burning, carried on in lime kilns, illustrates this. Limestone (that is, calcium carbonate) is heated in the kiln, when carbon dioxide escapes into the air, and lime (that is, calcium oxide) remains. CaCO 3 = CO, + CaO. Limestone. Lime. The " setting " of mortar is partly due to the absorption of carbon dioxide out of the air by the lime in the mortar. Old mortar, therefore, contains calcium carbonate. Experiment 207. Pour a little hydrochloric acid upon some fragments of old mortar notice the effervescence, and pass the gas into lime water. To find the Weight of Carbon Dioxide in Marble. Experiment 208. Select a small thin flask with a wide mouth, and fit it with a cork having two holes. Through one hole insert a short straight glass tube, the top of which can be closed with a tiny cork. Into the other hole fit a bent glass tube with a small bulb blown on it, as shown in Fig. 98. Make a hole in the bottom of a small test-tube, by heating the end with a fine-pointed blow- pipe flame, and blowing down the test-tube. Then cut off the tube about 3^ cm. (\\ inches) from the bottom, and " border " the end (p. 36) so as to obtain the little appa- ratus shown at A, Fig. 98. Fasten a thin copper wire to the neck, and hang the little tube inside the flask in the manner shown. \ *:*-" ~ - Place a little strong sulphuric acid in the bulb tube B, so as to just fill the bend, and close the open end with a little cap, C (p. in). Put about 10 or 12 cc. of a mixture of equal parts of hydrochloric acid and water in the flask, but do not let it touch the little hanging tube. Now carefully weigh the whole apparatus. Then remove the cork and drop into the hanging tube one or two little fragments of marble (not so small as to drop through the hole Analysis of a Carbonate. 229 in the bottotri), and weigh again. The increase gives the weight of marble, which should be from i to 2 grams. Remove the cap C, and then lower the tube containing the marble so that it dips into the hydrochloric acid, and replace the little cork d. The acid dissolves the marble, and the carbon dioxide which is evolved makes its escape by bubbling through the sulphuric acid in B. This prevents the gas from carrying away any vapour of water with it. When the marble has all dissolved, the apparatus of course is full of carbon dioxide, and as this is much heavier than air it must be removed. Take out cork d, and by means of a short piece of rubber tube attached to the bulb-tube, slowly suck the gas out of the flask, making it bubble quite slowly through the sulphuric acid in the tube. When the gas is all out, and the apparatus is full of air as it was at first, replace the cork d, and cap C, and weigh again. The loss of weight represents the carbon dioxide which has passed out. Two experiments should be made, so as to confirm the result. EXAMPLE. Weight of marble used = 1*5 grams Weight of carbon dioxide = O'6 as 1*5 : 100 :: 0*6 = 40 therefore this sample of marble contains 40 per cent, of carbon dioxide. EPITOME. Carbon dioxide is produced when all carbon compounds burn. Coal, gas, candles, oil, wood, etc., when burnt give carbon dioxide amongst other things. The breathing of animals and man pro- duces carbon dioxide, and it is given off during processes of decay and fermentation. It is also formed when limestone (that is, calcium carbonate) is heated, as in the process of lime-making in lime kilns. Carbon dioxide is prepared by acting on calcium carbonate (marble, chalk, or limestone) with hydrochloric acid. Carbon dioxide is a colourless gas, slightly soluble in water forming a feebly acid solution, which is decomposed again when heated. The gas does not burn, and puts out ordinary flames. Carbon dioxide is not exactly poisonous, because we are always breathing small quantities of it in the air, but if animals are placed in the gas they quickly die for want of oxygen. Even a moderate 1 The student should calculate what weight of carbon dioxide would have been obtained if the marble had been pure calcium carbonate. 230 Oxides of Carbon. quantity of this gas in the air, over and above the usual amount, is injurious to life. Carbon dioxide is a heavy gas, which can be collected' by down- ward displacement, and can be poured from one vessel to another. When passed into lime water it unites with the lime and forms a white precipitate of chalk. It is quickly absorbed by either caustic soda or potash, forming sodium or potassium carbonates. The test for a carbonate is to add an acid to it, and to prove that the gas which is given off is carbon dioxide by passing it through lime water and obtaining the precipitate of chalk. CHAPTER XXV, OXIDES OF CARBON continued. Carbon Monoxide, CO. It has been shown that when carbon burns in air or oxygen under ordinary circumstances it gives carbon dioxide, but under particular conditions the other oxide is formed. For instance, if we fill a long piece of combustion tube with fragments of charcoal, and heat the tube to a good red heat in a furnace, and then pass a moderate stream of oxygen in at one end, instead of getting carbon dioxide coming out at the other, we should find that it was quite a different gas ; for we could set fire to it, and should see it burn with a beautiful blue flame. What takes place is this. As the oxygen first meets the hot charcoal, the charcoal burns and greedily takes all the oxygen it can get, and produces carbon dioxide. But as this passes along the red hot tube it gives up half of its oxygen to a further portion of carbon, so that they go equal shares, as it were, with the oxygen, and the result is carbon monoxide. C0 2 + C = 2CO. We can prove that this is the true explanation in the follow- ing way. Experiment 209. Take a piece of ordinary iron gas-pipe and fill it with fragments of charcoal and make it red hot in a furnace (Fig. 99). (If more convenient, the iron pipe may be bent in the middle into a sort of elbow shape, and the bent part pushed into an ordinary fire, between the bars.) Now we must not pass oxygen direct into this pipe, or the iron itself would burn up (refer back to oxygen) ; so we must connect to one end of it a short piece of glass combustion tube containing 232 Oxides of Carbon. a few pieces of charcoal. Now heat the charcoal in the glass tube and pass a gentle stream of oxygen through. The carbon burns, and the carbon dioxide passes along over the red hot charcoal in the iron pipe. Set fire to the gas which escapes at the other end, and observe its flame. This gas is carbon monoxide. Carbon monoxide is always being produced by exactly this method in our ordinary fireplaces. The oxygen of the air which enters the front and bottom of the grate is taken up by the first portion of burning coal, and forms with the carbon carbon dioxide. This, in passing through the fire, meets with red-hot carbon, and gives up half its oxygen, forming carbon FIG. 99. monoxide. This operation always goes on in the fire, but to a greater extent if the fire is hot and clear, and especially in a coke fire. A good deal of the carbon monoxide thus formed burns on the top of the fire, and the familiar bluish flame seen flickering about on the top of a clear bright fire is this gas burning. But some of it escapes unburnt up the chimney, because there is often so much carbon dioxide passing off from the fire as to prevent the other gas from burning. If instead of passing carbon dioxide through the tube of red-hot charcoal, we were to send steam through it, we should then get a mixture of carbon monoxide and hydrogen, C + H 2 O = CO -f H* This mixture is sometimes called water-gas, and is made Carbon Monoxide. 233 on a large scale by passing steam over strongly heated coal or coke. Preparation of Carbon Monoxide. Experiment 210. Put a few crystals of oxalic acid into a test- tube, and pour upon them a little strong sulphuric acid. Fit a cork and delivery tube to the test-tube and gently heat the mixture. Notice that a brisk effervescence soon begins to take place. Collect two jars of the gas, over water. Test the gas in one jar with a lighted taper, and note that it burns with a blue flame, but does not seem to burn very well. Into the other jar pour some lime water ; notice that milkiness is at once produced. This shows that carbon dioxide is also present, because no other gas pro- duces this turbidity with lime water. To find out how much carbon dioxide is present, we can make use of the fact that this gas is quickly absorbed by caustic soda. Experiment 211. Collect some more of the gas given off by the oxalic acid in a burette (fill the burette with water, and invert it in a basin or trough, just as an ordinary gas-collecting jar). When it is full of gas, attach a small funnel to the top with a piece of rubber tube, as shown in Fig. ioo. Pour some caustic soda solution into the funnel, and then gradually turn the tap so as to allow the liquid slowly to enter the tube and trickle down the side. Close the tap before the funnel is empty, or else air will be drawn in. FIG i< Notice that as the caustic soda enters, the gas is quickly absorbed by it, and the water rises in the tube. When it stops rising, note how much gas remains ; almost exactly one half. Close the tube with the thumb, remove it from the trough and invert it. Then bring a lighted taper to the gas. Notice that it burns with the blue flame of carbon monoxide. This shows that, when oxalic acid is acted on by sulphuric acid, carbon dioxide and monoxide are given off in equal volumes. The equation expressing the change is this C 2 H 2 O 4 = C0 2 + CO + H 2 0. Oxalic acid. 234 Oxides of Carbon. The water is taken up by the sulphuric acid, which has a powerful affinity for water. When we want carbon monoxide pure, we must pass the gas obtained from oxalic acid through bottles containing caustic soda, so as to get the carbon dioxide removed. Experiment 212. Fit up the apparatus shown in Fig. 101. Place in each of the two bottles some solution of caustic soda, and heat the mixture of oxalic acid and sulphuric acid in the small flask. As the gases bubble through the caustic soda, the carbon dioxide is absorbed, and the monoxide passes on. Collect three jars. Dip a lighted taper into FIG IQI one. Notice that the gas burns with a much stronger flame than before. ' Also note that the taper itself is extinguished if thrust into the gas. Add lime water to the second. Is there any turbidity produced ? If so, it proves that even bubbling through two bottles has not quite removed all the carbon dioxide. Take the third jar, and quickly pour into it a little caustic soda and cover it again imme- diately. Shake the liquid up with the gas, and then replace the jar in the trough for a few minutes, in order to let the caustic soda go out into the water. Now take the jar out again, and pour some lime water into it and shake up. This time there should be no turbidity at all. Now light the gas, and as the flame passes down into the jar, cover it with the glass plate. Again shake up the lime water that is in the jar, and notice that now it is instantly made milky. This shows that when carbon monoxide burns, it gives carbon dioxide. CO + O = C0 2 . Carbon monoxide takes oxygen away from many metallic oxides when they are strongly heated in this gas. It is, there- fore, like carbon, a reducing agent. Thus, if carbon monoxide is passed over heated oxide of iron, the oxide is deprived of Carbon Monoxide. 235 its oxygen and becomes reduced to the metallic state, and carbon dioxide is formed. Fe 2 O 3 + SCO = 3CO 2 + 2Fe. This process goes on in the blast furnace where iron ores are smelted. Carbon monoxide is extremely poisonous. Many people have been killed by the gas escaping from coke or charcoal fires, their deaths being due to the poisonous nature of the carbon monoxide which such fires give off. EPITOME. Carbon monoxide is produced when carbon burns in an insuffi- cient supply of oxygen, or when carbon dioxide is passed over red hot charcoal. It is prepared from oxalic acid by the action of sulphuric acid, the gas so obtained being passed through caustic soda to remove the carbon dioxide. Carbon monoxide is a colourless poisonous gas, which burns with a beautiful blue flame. When it burns it produces carbon dioxide. Carbon monoxide does not dissolve in water ; does not make lime water turbid ; is not absorbed by caustic potash. It is not an acid forming oxide like carbon dioxide, and, therefore, forms no salts. It is readily distinguished from all other gases by burning with a blue flame, and forming carbon dioxide. CHAPTER XXVI. SULPHUR. THIS element is found chiefly in volcanic regions, such as Sicily and Iceland, where it exists in the free state ; that is, not in chemical combination with other elements. The sulphur as thus found is called native sulphur. Besides occurring in this uncombined state, sulphur is also a constituent of a large number of important ores, in which it is combined with various metals. Some of the commonest of these sulphides are iron pyrites (sulphur combined with iron, FeS 2 ) ; copper pyrites (sulphur with copper and iron, CuFeS 2 ) ; zinc blende (sulphur and zinc, ZnS) ; galena (sulphur and lead, PbS). In combination with metals and with oxygen together, it is found in heavy spar (barium sulphate, BaSO 4 ) ; and the very common mineral gypsum (calcium sulphate, CaSO 4 + H 2 O). Modes of obtaining Sulphur. (i) "Native sulphur" is always mixed up with more or less earthy and mineral matters. In order to separate the sulphur, the crude material is piled up into heaps on a slanting hearth, and the heaps set on fire. Some of the sulphur burns away, but the heat it gives out melts the remainder, which runs away from the impurities down the sloping floor. Of course the supply of air to the heap requires to be regulated, for if it had free access, the whole of the sulphur would be burnt away. (2) Sulphur can also be got from iron pyrites, by heating the ore strongly, without letting air get to it. Experiment 213. Heat a little powdered iron pyrites in a test- Purification of Sulphur. 237 tube. Notice the sublimate which collects on the cooler part of the tube, as it forms little drops of melted sulphur. After a time, just touch the hot end of the test-tube with a drop of water so as to crack it off. Now heat the sulphur on the side of the tube, holding the tube in an inclined position. Notice the sublimed sulphur takes fire and burns, and the gas which escapes at the- top of the tube has the characteristic choking smell produced by burning sulphur. The pyrites does not part with all its sulphur when heated in this way. The change is expressed by the equation 3 FeS 2 = Fe 3 S 4 + 28. If the iron pyrites is roasted in a free current of air, it loses all its sulphur, and both the iron and the sulphur are con- verted into oxides, thus 2FeS 3 -f- nO = Fe 2 O 3 A large quantity of sulphur is now recovered from waste products which contain this element, and which used to be thrown away. Such a refuse sub- stance as that known as Alkali- waste (obtained in the process of manufacturing sodium carbonate) is now utilized in this way. Purification of Sulphur. Experiment 214. Bend a piece of wide glass tube, one end of which is closed up, and attach it to a retort in the manner shown in Fig. 102. Place a few pieces of sulphur in the little bent tube, and apply heat to it. The sulphur melts ; then gets very dark in colour, and presently boils, and the vapour passes into the body of the retort. The sulphur is here being distilled, and the impurities it may contain remain behind. Notice that, as the vapour enters the large retort, some of it condenses upon the glass as a yellowish powder, while some collects on the lower part of the vessel in the liquid state and quickly solidifies. FIG. 102. 238 Sulphur. On a large scale this process is carried out by boiling the sulphur in earthenware retorts, and sending the vapour into large brickwork chambers. At first it condenses as a fine light yellow powder on the walls. In this condition it is called flowers of sulphur. After a time the walls of the chamber get warm, and the sulphur melts and collects on the floor, and is then run out into wooden moulds so as to cast it in the form of sticks. This is called roll sulphur, or brimstone. Properties of Sulphur. Experiment 215. Take a piece of common roll sulphur, and strike it gently on the table, or with a pestle. Notice how brittle it is. Examine the freshly broken surfaces and see that it is highly crystalline. Powder a little of it in a mortar, and try if it dissolves in water. After shaking it up with water in a test-tube for a little time, decant some of the water into a small dish and evaporate gently to dry ness. If there is nothing left it shows that sulphur is not dissolved by water. When sulphur is heated it behaves in a rather striking manner. Experiment 216. Carefully heat a little of the powdered sulphur in a test-tube ; it easily melts, and, if not over-heated, gives a pale amber-coloured liquid which runs about the tube like oil. Now heat more strongly, and notice that the sulphur rapidly deepens in colour, becoming like dark treacle, and gets so thick and sticky that if the tube is turned upside down it does not run out at all. Heat it still more, and note that it becomes quite liquid again, although remaining dark coloured, and presently boils. It is difficult to see the colour of the vapour, because of the almost black appearance of the liquid ; but by looking through the tube at any part where the glass is clear, it will be seen that the vapour has a pale yellow colour. Let the test-tube cool, and the sulphur goes through the same changes in the opposite order. Notice that as it solidifies it forms crystals on the sides of the tube. Sulphur melts at 114*5; at a temperature about 230 it passes into the thick condition, and at 448 it boils. Allo tropic Modifications. Sulphur, like carbon, exists in three allotropic forms; and, like those of carbon, two are Allotropic Forms of Sulphur. 239 crystalline, and one non-crystalline or amorphous. But here the similarity ends. Sulphur is quite easily made to assume either of its three forms (not so, carbon), but it will only remain in one of them, for both the other varieties gradually pass back into the first. (Carbon is stable in each of its allotropic forms.) Experiment 217. Put some fragments of roll sulphur into a test-tube, and add to them a small quantity of a liquid called carbon disulphide (CS 2 ) just to cover them. Notice that the sulphur quickly dissolves. Pour the solution into a small dish, cover it .with a piece of cardboard, and leave it for some time to* slowly evaporate. Examine the residue with a pocket lens, and carefully note the shapes of the crystals. If this experiment is made on a larger scale, and with certain precautions, more perfectly shaped crystals will be obtained, like the one shown in Fig. 103. The form of the crystals of sulphur obtained in this way is what is known as Rhombic Octahedral. They have a beautiful amber-like and are very appearance, brittle. The sulphur which is found "native" is in this form. FIG. 103. Experiment 2 1 8. Carefully melt some roll sulphur in a small beaker; the beaker being about three-quarters full. Allow it to cool, carefully watching it, and as soon as a thin crust has formed on the top, pour out what remains of the liquid into a dish or plate. On cutting away the crust, the interior of the beaker will be found to be lined with long needle-shaped crystals, like those shown in Fig. 104. Examine these crystals, and note how entirely different they are in shape to the others. They have also a transparent appearance. FIG. 104. 240 Sulphur. This is the second allotropic form of sulphur, and is known by the name Prismatic Sulphur, because the crystals are in the shape of long thin prisms. If these crystals are kept for a day or two they lose their transparent appearance, and become exactly the colour of ordinary brimstone. In fact, they change back again from the prismatic modification to the rhombic form; and although the crystals retain the out- ward form of the prism, they crumble on the slightest touch to a number of minute crystals, having the shape of rhombic octahedrons. Experiment 219. Heat a quantity of sulphur (either "flowers" or powdered lump) in a common oil flask until it gets into a boiling condition. Then pour the hot sulphur in a thin stream into cold water in a beaker. (A funnel may be stood in the water, and the stream of sulphur poured round and round it, Fig. 105.) Now lift out the funnel with the congealed sulphur, and notice how en- tirely different it is from either of the other forms. It is no longer brittle, but seems almost like indiarubber, and can be stretched "FIG! 105. an d pulled into threads. This is the third allotrope of sulphur, and is called Plastic Sulphur. It has no crystalline character at all. Like the prismatic form, this plastic sulphur when left to itself for a few days changes back again into the rhombic variety. It gradually loses its curious elasticity, and becomes the ordinary brittle, yellow, crystalline sulphur. If it is stretched about, or slightly warmed, it changes from its plastic state to the ordinary condition much more quickly. Milk of Sulphur is an old-fashioned name for sulphur obtained in the following way. Experiment 220. Throw a small handful of flowers of sulphur into a saucepan half full of hot water, and add about twice as much lime. Boil the mixture for five or ten minutes, and then allow it to settle. Take some of the clear yellowish liquid and add to it a little strong hydrochloric acid. A white precipitate of sulphur in the condition of very fine particles is produced, which Sulphides 241 makes the liquid look almost like milk ; hence the name " Milk of Sulphur." This is a much finer powder than flowers of sulphur, and is on this account used in medicine. When sulphur, in any of its allotropic forms, is heated in air, it burns with a blue flame (see Exps. 64, 138), and gives an oxide of sulphur called sulphur dioxide, SO 2 . Combination of Sulphur with Metals. Sulphur combines with many metals, and forms compounds called sulphides ; just as oxygen unites with metals and gives oxides. Experiment 22 1 . Heat a small quantity of flowers of sulphur in a test-tube until it boils, and drop in a fragment of sodium about the size of a large pin's head. The sodium instantly takes fire and burns brilliantly, producing sodium sulphide, Na 2 S. Experiment 222. Drop into a similar small quantity of boiling sulphur some reduced iron (that is, iron obtained by heating iron oxide in a stream of hydrogen or coal gas). The black powder at once takes fire as it comes into the sulphur vapour, and gives ferrous sulphide, FeS. If iron filings are used instead of the "reduced" iron, they combine with the sulphur, but being so much coarser they do not take fire ; but if a fine iron wire is made red hot in a Bunsen flame, and then plunged into the sulphur vapour, the wire burns, and melted drops of iron sulphide fall to the bottom of the tube. Thin copper wire also readily combines with sulphur vapour, and takes fire without being heated in a lamp, producing copper sulphide, CuS. EPITOME. Sulphur occurs in volcanic regions, such as Sicily and Iceland, in the free or elementary state. In combination with many metals, in some of the commonest ores of these metals, as sulphides ; also in combination with metals and oxygen, as sulphates. " Native sulphur " is burnt in heaps, with a limited air supply, so as to melt the sulphur and separate it from rocky matter with which it is mixed. Sulphur is got from iron pyrites by simply heating it, when it gives off one-third of its sulphur. Sulphur is purified by distillation ; and is obtained either as flowers, or is melted and cast into sticks, known as roll sulphur. R 242 Sulphur. The three allotropic forms of sulphur are (1) Rhombic (octahedral). Permanent, soluble in carbon disulphide. Specific gravity, 2^05 . Lemon yellow colour, very brittle. (2) Prismatic. Not permanent, slowly changes to No. I. Translucent amber yellow crystals. Specific gravity, 1*98. (3) Plastic. Not permanent, slowly changes to No. i. Soft, non-crystalline, indiarubber-like, translucent yellow. Not dis- solved by carbon disulphide. Specific gravity, 1-95. Sulphur (any variety) burns in the air with a blue flame, producing sulphur dioxide. Sulphur is not soluble in water, but is oxidized by nitric acid into sulphuric acid (see p. 191). It unites directly with metals, forming sulphides. Sulphur belongs to a family of elements, the members of which are oxygen, sulphur, selenium, and tellurium. The last two are rather rare elements. The chemical relationship between sulphur and oxygen will be seen by comparing some of the compounds which each forms with other elements. Water, H 2 O Sulphuretted hydrogen, H 2 S. Potassium hydroxide, KHO Potassium hydrosulphide, KHS. Calcium hydroxide, Ca(HO) 2 Calcium hydrosulphide, Ca(HS) 2 . Carbon dioxide, CO 2 Carbon disulphide, CS 2 . Potassium, oxide, K 2 O Potassium sulphide, K 2 S. Copper oxide, CuO Copper sulphide, CuS. CHAPTER XXVII. SULPHURETTED HYDROGEN, H 2 S. Occurrence. This compound is a gas, and is present in volcanic gases. It is met with, dissolved in water, in certain sulphur springs, such as those at Harrowgate. When animal substances containing sulphur become putrid, sulphuretted hydrogen is formed; bad eggs owe their disagreeable smell to the presence of this gas. Ordinary coal-gas, as it leaves the retorts, contains a considerable amount of sulphuretted hydrogen. This is removed in the " purifiers " before the gas is sent out from the gasworks, and the sulphur it contains is extracted and sold. Modes of Formation. Sulphur does not very easily combine with hydrogen ; but if hydrogen is passed over boiling sulphur, a portion of the sulphur unites with hydrogen and forms sulphuretted hydrogen. Experiment 225. Heat a few fragments of sulphur in a horizontal bulb tube, until the sulphur boils, and then allow hydrogen to pass through the tube. Smell the gas escaping at the end. Also test it by holding a piece of paper moistened with a solution of lead acetate in the gas. The paper becomes black, showing the presence of sulphuretted hydrogen. In the laboratory, the gas is always prepared by another method, namely, by acting on ferrous sulphide with either sulphuric or hydrochloric acid. Ferrous sulphide is made by heating iron and sulphur together. Its composition is expressed by the formula FeS. This substance must not be confounded with iron pyrites, FeS 2 . 244 Sulphuretted Hydrogen. Experiment 224. Place a quantity of ferrous sulphide in a two-necked bottle (see Fig. 39), cover it with water, and pour a small quantity of strong sulphuric acid through the thistle funnel. Notice that effervescence immediately begins. After a few minutes, during which the air in the bottle is being gradually swept out, collect two jars of the gas, using water in the trough as warm as the hands can comfortably bear. The reaction in this case is FeS + H 2 SO 4 = FeSO 4 + H 2 S. Ferrous sulphate. If hydrochloric acid is used instead of sulphuric acid the equation is FeS + 2HC1 = FeCl 2 + H 2 S. Ferrous chloride. Then remove the delivery tube from the apparatus and light the gas as it escapes from the exit tube. Notice that the flame is bluer than that of hydrogen, but not so blue as that of burning sulphur. Gently smell the products of the burning gas ; note that there is the same choking smell of sulphur dioxide, as when sulphur burns. Hold a cold tumbler over the flame, and observe the moisture collecting. When sulphuretted hydrogen burns from a jet with a free supply of air, it gives sulphur dioxide and water - H 2 S + 30 = S0 2 + H 2 0. Experiment 22$. Depress a piece of glass down on to the flame, and note that there is a deposit of sulphur. Experiment 226. Test the gas in one of the jars with a lighted taper, note that the taper is extinguished when thrust into the gas. Observe also that, as the gas burns, there is a deposit of sulphur formed on the sides of the jar. This is from the same cause 'as in the last experiment, namely, because the supply of air is not sufficient to completely burn the gas. We may express it by this equation H 2 S + O = H 2 O + S. Experiment 227. Transfer the second jar of gas to a trough of cold water, and leave it standing mouth downwards. Notice that the water rises a little in the jar. Shake it up as much as possible without lifting the mouth out of the water, and in a short Sulphuretted Hydrogen. 245 time the water will have absorbed nearly all the gas. Smell the water ; note that it smells like the gas. Sulphuretted hydrogen is considerably soluble in cold water. At the common temperature, water dissolves about three times its own bulk of this gas. The solution is called sulphuretted hydrogen water. Warm water dissolves much less of the gas, therefore we usually collect it over hot water, as described in Exp. 224. Sulphuretted hydrogen made from ferrous sulphide is always mixed with free hydrogen, because the ferrous sulphide always contains some iron which is not combined with sulphur ; and when the acid comes in contact with this free iron, hydrogen is evolved (see Hydrogen, p. 46). Therefore, if sulphuretted hydrogen is required quite pure, we generally use antimony sulphide, Sb 2 S 3 . For the purposes for which the gas is generally required, however, the presence of a little hydrogen does not matter, so that in practice it is nearly always made from the iron compound. For a number of experiments we want just a few bubbles of sulphuretted hydrogen, and therefore it is convenient to have an apparatus so arranged for making the gas that we can stop the action when we please, and start it again. This cannot be done with the form of apparatus used for Exp. 224, without emptying the bottle each time. A very simple form of constant apparatus can be made in the following way. Experiment 228. Obtain two large test-tubes (" boiling tubes ") and draw them out at one end, as shown in Fig. 106. Secure one of them with wire or thread to a retort stand, and join their drawn- out ends with a piece of indiarubber pipe in the manner shown. Half fill the fixed one with small broken pieces of ferrous sulphide, and close the tube with a cork and exit tube, the latter carrying a short piece of rubber tube, /, with a screw clamp, s, upon it. Suspend the other boiling tube to a ring by means of a wire hook. Close the clamp, and pour dilute sulphuric acid into the open tube until it is about three-quarters full. Now gently open the clamp, when the acid will gradually enter the other tube, and, coming in contact with the ferrous sulphide, will cause the evolution 246 Sulphuretted Hydrogen. of sulphuretted hydrogen. So long as the clamp is open, the gas will escape from the tube, but as soon as it is closed again, the gas, which is still being produced, not being able to get out, begins to drive the acid back into the open tube, and then, of course, the action stops. In this way, by opening and closing the clamp, we can produce a little gas and then stop the supply at will. By means of this little apparatus for getting sulphuretted hydrogen, do the following experiments Action of Sulphuretted Hy- drogen on Metals. Experiment 229. Open the screw clump, s, and let the gas blow against a clean silver coin for a moment. Notice that the silver is at once turned black. [A silver spoon is stained in the same way by a bad egg.] The black substance is silver sulphide, Ag 2 S. Experiment 230. Place a small frag- FlG Io6 ment of potassium in a horizontal bulb tube, and pass sulphuretted hydrogen through the tube. Heat the potassium, when it will burn brightly in the gas. Sulphuretted hydrogen, therefore, will support the combustion of potassium. The reaction is H 2 S + K = KHS + H. Potassium hydrosulphide. Compare this with the action of potassium upon water H 2 O + K= KHO + H. Action of Sulphuretted Hydrogen on Metallic Compounds. Experiment 231. Let sulphuretted hydrogen blow against a piece of ordinary canvas, used by artists. Notice that it is quickly blackened. The canvas is painted with white lead (a lead car- bonate), this is at once converted into lead sulphide, which is black. Metallic Sulphides. 247 [There is always a little sulphuretted hydrogen in the air of towns, and this makes oil paintings gradually turn black.] Experiment 232. Put a little oxide of iron in a horizontal glass tube, and gently warm it. Then pass sulphuretted hydrogen through the tube. Notice that the reddish oxide is turned black, and glows with the heat produced by its combination with the sulphur Fe a O s + 3H 2 S = 2FeS + S + 3H a O. Similarly if sulphuretted hydrogen is passed over slaked lime (calcium hydroxide), we get a sulphur compound of calcium produced. CaH 2 O 2 + 2H 2 S - 2H 2 O + CaH 2 S 2 . Calcium hydroxide. Calcium hydrosulphide. These two substances, namely, slaked lime and iron oxide, are the materials used in the " purifiers " of the gas-works, to absorb the sulphuretted hydrogen from the coal-gas. Action of Sulphuretted Hydrogen on Metallic Solutions. Experiment 233. Prepare the three following solutions : (1) Dissolve a small particle of lead acetate in water in a test-tube. (2) Dissolve a similar quantity of white arsenic (arsenious oxide) in a few drops of hydrochloric acid, and add water to half fill the test- tube. (3) Dissolve a small quantity of tartar emetic (an antimony salt) in water. Now pass a few bubbles of sulphuretted hydrogen through each of these, by dipping a tube from the apparatus (Fig. 106) into the solutions in turn. Notice what happens in each case. We get a black, a yellow, and a red precipitate. In other words, the three elements, lead, arsenic, and antimony, form sulphides ; lead sulphide being black, arsenic sulphide yellow, and antimony sulphide red. Sulphuretted hydrogen, therefore, affords a test by which we can easily distinguish between compounds of these elements. Experiment 234. In three separate test-tubes put a little dilute solutions of (i) copper sulphate, (2) ferrous sulphate, and (3) potas- sium nitrate ; add to each, one or two drops of hydrochloric acid. 248 Sulphuretted Hydrogen. and pass into each a few bubbles of sulphuretted hydrogen. Note that there is a precipitate only in the copper solution. This means that copper sulphide is precipitated in an acid solution ; whereas no sulphide of either iron or potassium is produced in a solution containing hydrochloric acid. In acid solution, CuSO 4 + H 2 S = CuS + H 2 SO 4 . Now add a few drops of ammonia to the two solutions which were not precipitated, and note that in the case of the iron salt, a black precipitate is produced, but nothing happens in the solution of the potassium salt. In alkaline solution, FeSO 4 + H 2 S = FeS +.H 2 SO 4 . This black precipitate is ferrous sulphide ; and this experiment shows that in an alkaline solution iron is precipitated in the form of sulphide, while no potassium sulphide is produced either in the acid or alkaline solutions. The following experiment shows what use we can make of these facts Experiment 235. Pour a little of the solutions of copper sulphate, ferrous sulphate, and potassium nitrate into the same test-tube. Add a drop or two of hydrochloric acid, and bubble sulphuretted hydrogen through the solution for a few minutes. We know by the former experiment that only the copper will be precipitated as sulphide. Now filter the liquid. The copper sulphide remains on the filter, while the solution which passes through contains the iron and potassium salts. Next add to the filtrate some ammonia. This we have seen causes the precipitation of ferrous sulphide (there being sulphuretted hydrogen dissolved in the solution). Pass this through another filter. The ferrous sulphide remains on the filter and the potassium salt passes through. In this way we have separated the three metals, copper, iron, and potassium, whose salts were originally mixed in the solution. Sulphuretted hydrogen is, therefore, a most important agent in analysis; for we find (i) a certain number of metals which are precipitated as sulphides in acid solution, (2) others which are not precipitated as sulphides in acid, but only in alkaline liquids, and (3) others which are not precipitated as sulphides in either acid or alkaline solutions. Test for Sulphuretted Hydrogen. 249 Besides this, as we have seen, some sulphides have cha- racteristic colours, by which they are easily distinguished. Test for Sulphuretted Hydrogen. The smell of this gas is sufficiently characteristic to distinguish it from all others. An additional test is its action on a salt of lead, such as lead acetate. Paper moistened with a solution of lead acetate is turned black by this gas owing to the formation of lead sulphide. EPITOME. Sulphuretted hydrogen is prepared by the action of sulphuric or hydrochloric acid on ferrous sulphide. The gas has a dis- agreeable smell, like bad eggs, and is poisonous. It is moderately soluble in cold water, giving a solution having the smell of the gas. It must be collected over hot water. The gas burns with a bluish flame, giving water and sulphur dioxide if excess of air is present, but depositing some of its sulphur if the supply of air is limited. The gas combines directly with many metals, as lead, copper, silver, forming sulphides. The brown " tarnish " which comes on silver articles exposed to the air of towns, is due to the formation of silver sulphide. Sulphuretted hydrogen also acts on metallic compounds, both in the solid state (as in the case of " white lead," oxide of iron, lime) or in solution. On account of its behaviour towards metallic salts in solution, it is used in analysis, for separating and detecting the various metals whose compounds are present. Thus, of the six metals, lead, copper, iron, zinc, calcium, potassium, the sulphides of lead and copper 'are precipitated in acid solutions, whilst those of the others are not. The sulphides of iron and zinc are precipitated in alkaline solutions, whilst sulphides of calcium and potassium' are soluble in both acids and alkalies. The action of sulphuretted hydrogen on solutions of salts of these metals is represented by the following equations Pb(NO 3 ) 2 + H 2 S = PbS + 2HNO S ) , CuS0 4 + H 2 S = CuS + H 2 S0 4 } In acid solutlOns ' FeSO 4 + H 2 S = FeS + H 2 SO 4 ) T ZnS0 4 + H 2 S = ZnS + H 2 SO 4 In akalme solutlons - CaCl 2 + H 2 S. No action. KC1 + H 2 S. No action. CHAPTER XXVIII. SULPHUR DIOXIDE SULPHUROUS ACID SULPHITES. THE two most important oxides of sulphur are sulphur dioxide, SO 2 , and sulphur trioxide, SO 3 . Both of these compounds are acid-forming oxides. The first, when dissolved in water, gives sulphurous acid, H 2 SO 3 ; whilst the second gives sulphuric acid, H 2 SO 4 . Sulphur Dioxide, SO 2 . This compound is always produced when sulphur is burnt in the air or in oxygen, so that we can make it this way S + O 2 = SO* Experiment 236. Place a piece of sulphur in a short horizontal piece of combustion tube, attach a delivery tube to one end, and arrange to pass oxygen in at the other. Heat the sulphur until it begins to burn, and let oxygen slowly stream through. Collect the gaseous product by downward displacement, covering the mouth of the cylinder with a piece of paper or card. [Note that the gas is not clear. This is because there is always produced at the same time a small quantity of sulphur trioxide, along with the dioxide. ,] Test the gas with a lighted taper, notice that the gas does not burn, and that it puts out the taper. Fan a little of the gas towards the face, so as to get a slight whiff of it. Note its choking smell, familiarly known as " the smell of burning sulphur." Enormous quantities of sulphur dioxide are made in this way for the manufacture of sulphuric acid. Sometimes sulphur itself, and sometimes iron pyrites, is used ; and it is burnt, not in glass tubes in a stream of pure oxygen, but in special furnaces and in ordinary air. Sulphur Dioxide. 251 Preparation. When we want sulphur dioxide for experiments in the laboratory, we always make it by acting on copper with sulphuric acid. Experiment 237. Place a quantity of scrap copper in a flask, fitted with a cork and exit tube, and pour upon it enough strong sulphuric acid to well cover it. Heat the acid carefully with a rose burner. Notice that, as the acid gets warm, it becomes muddy, owing to the formation of a black powder, which collects in the flask in considerable quantity. Presently effervescence sets in, and sulphur dioxide is rapidly evolved. Collect four jars of the gas by displacement. The final result of the action of sulphuric acid on copper is expressed by the equation Cu + 2 H 2 S0 4 = CuS0 4 + 2 H 2 O + SO** Properties of Sulphur Dioxide. The specimens col- lected show that it is a colourless gas : and its smell will have been perceived during the preparation of it. We have also seen, by Exp. 236, that the gas does not burn, nor support the combustion of a taper. Experiment 238. Place one of the jars of gas mouth downwards in water. Notice that the gas is absorbed fairly rapidly. Add a few drops of litmus to the water and note that the solution of this gas is strongly acid. At the common temperature, water dissolves about fifty times its own volume of sulphur dioxide. Experiment 239. Pour into a second jar of the gas 3 or 4 cc. of litmus solution, and shake it up. As before, the litmus is instantly reddened, but after a time the colour gets fainter and finally almost entirely disappears. The gas has bleaching properties. Now add to the bleached (or nearly bleached) liquid a few drops of strong sulphuric acid. Notice that the red colour is restored. Sulphur dioxide is used for bleaching straw, flannel, sponges, and other articles which would be injured by chlorine. Its * But this does not explain everything that goes on, because it only tells us that copper sulphate, water, and sulphur dioxide are formed, and takes no notice of the black substance which is produced in the flask. Copper sulphate is not black. The equation, therefore, only gives us the final products. We do not know exactly what intermediate products are formed, and therefore cannot express their formation by an equation. 252 Sulphur Dioxide. bleaching power is not due to the oxidation of the colouring matter, as in the case of chlorine, and is not always permanent. Sponges, flannels, and straw articles, gradually return to their unbleached state. Flowers, especially those of a violet or purplish tint, are quickly bleached if placed in sulphur dioxide. Although sulphur dioxide does not support the combustion of ordinary combustibles, some things will burn in this gas. Experiment 240. Place a small heap of lead dioxide on a deflagrating spoon, gently warm it and then lower it into a jar of sulphur dioxide. The two dioxides combine with so much energy that the lead dioxide becomes red hot, and forms lead sulphate. PbO 2 + SO 2 = PbSO 4 . Experiment 241. Sprinkle a few particles of sodium dioxide (sodium peroxide) into a jar of sulphur dioxide. The sodium dioxide takes fire and burns brilliantly, forming sodium sulphate, Na 2 O 2 + SO 2 == Na 2 SO 4 . Sulphur dioxide is used as a disinfectant, and for this purpose is generally obtained by burning sulphur. FIG. 107. When sulphur dioxide is cooled below 8 it condenses to a liquid. Experiment 242. Cut a test-tube about 4 cm. (nearly 2 inches) from the closed end, and border the end. Into this short tube fit Sulphites. 253 a cork, carrying two long thin tubes, as shown in Fig. 107. Place this in a vessel filled with a mixture of powdered ice and salt, and connect it to the apparatus for making sulphur dioxide, first drying the gas by bubbling it through strong sulphuric acid in the bottle w. The ice and salt mixture cools the tubes below 8, so that the gas which goes through will be condensed to the liquid state, and will collect in the little test-tube. In case any gas passes through with- out condensing, attach a delivery tube to the apparatus, and let it dip into water. This will absorb any escaping gas. When enough liquid is collected, lift the tube out of the freezing mixture, and pour the liquid upon a little water in a small dish. The sulphur dioxide boiling at 8 at once freezes some of the water. Do not inhale the gas, as it is very irritating to the lungs. Liquid sulphur dioxide is used for the artificial production of ice on a large scale. Sulphurous Acid, H 2 SO 3 . A solution of this acid is produced when sulphur dioxide is dissolved in water. The solution is strongly acid towards litmus, and smells like the gas. When it is warmed the gas is driven off again, and the whole of it is expelled by boiling the solution. Sulphurous acid does not keep. It slowly absorbs oxygen from the air, and changes into sulphuric acid. This change is expressed by the equation 2 S0 3 + O = H 2 S0 4 . On account of the readiness with which it takes up oxygen from other compounds, sulphurous acid is a powerful reducing substance. Experiment 243. Dissolve a crystal of potassium permanganate in water. This substance is very rich in oxygen (KMnO 4 ). Pour this deep violet solution into some sulphurous acid (made by passing sulphur dioxide into water). Notice that the colour is instantly discharged. The sulphurous acid takes some of the oxygen away from the permanganate, and changes into sulphuric acid. Sulphites are salts of sulphurous acid, obtained by neutralizing the acid with a base. Experiment 244. Add caustic soda solution cautiously to some sulphurous acid, until a drop of the liquid on a glass rod just turns reddened litmus paper blue. Then add one or two drops more of 254 Sulphur Dioxide. the acid so as to make the solution just acid. Evaporate the solution to dryness, and obtain a white salt. This is sodium sulphite. Add a few drops of sulphuric acid to it. Notice that there is effervescence. Smell the gas, and note that it is sulphur dioxide. Sulphuric acid, therefore, decomposes sulphites, expelling sulphur dioxide, and forming sulphates. The reactions here are (1) H 2 S0 3 + 2NaHO = Na 2 SO 3 + 2H 2 O. (2) Na 2 SO 3 + H 2 SO 4 = Na 2 S0 4 + H 2 O + SO 2 . Sulphurous acid contains two atoms of hydrogen ; it is, therefore, called a dibasic acid. It can form two classes of salts (just as carbonic acid does, see p. 227), depending on whether both or only one of these hydrogen atoms are dis- placed. Thus, there are two sodium sulphites (1) Normal sodium sulphite, Na 2 SO 3 . (2) Hydrogen sodium sulphite (or sodium bi-sulphite), HNaSO 3 . Salts, like this hydrogen sodium sulphite, in which only a part of the hydrogen originally present in the acid has been exchanged, are sometimes called acid salts. It does not follow, however, that they have an acid reaction towards litmus, although in many cases they do happen to possess this property. It must be remembered that an acid salt simply means a salt in which the whole of the hydrogen of the acid has not been displaced by the metal. Sulphites are all decomposed by stronger acids, such as hydrochloric or sulphuric. The action of sulphuric acid is shown by the equation above. With hydrochloric acid the only difference is that a chloride of the metal is formed, thus K 2 S0 3 + 2HC1 - 2KCI + H 2 O + SO 2 . Note that when lead dioxide and sodium dioxide were burnt in sulphur dioxide, the sulphates and not the sulphites of the metals were produced. CHAPTER XXIX. SULPHUR TRIOXIDE SULPHURIC ACID SULPHATES. Sulphur Trioxide, SO 3 . When sulphur burns in oxygen, however much oxygen there may be, the compound produced is always sulphur dioxide, and only a minute trace of the trioxide is formed at the same time. But if a mixture of sulphur dioxide and oxygen is passed through a heated tube containing very finely divided platinum, then the sulphur dioxide combines with the oxygen and gives sulphur trioxide. The way in which the heated platinum causes these two gases to unite together is not clearly known, but the platinum itself is not altered. Experiment 245. Dip some asbestos fibres into a solution of platinum chloride, and then hold them by means of a small pair of tongs in a bunsen flame until they are quite hot. The platinum chloride first dries and then decomposes, leaving the asbestos coated over with very finely divided platinum. Now pack a quantity of this " platinized asbestos " into a bulb tube, which is supported as shown in Fig. 38. Remove the bottle w from the sulphur dioxide apparatus (Fig. 107), and replace it by a similar bottle having a third tube, which dips into the acid. Connect this tube with a supply of oxygen, so as to let both sulphur dioxide and oxygen bubble through the same bottle and become mixed. Notice that so long as the platinum is cold, no fumes of sulphur trioxide escape from the bulb tube, but as soon as it is heated, white fumes make their appearance. If these fumes are passed through a U-shaped tube, kept cold by being placed in a freezing mixture (powdered ice and salt), white silky -looking crystals will condense in the cold tube. This is the sulphur trioxide. Sulphur trioxide has a most powerful affinity for water. If exposed to the air, it soon takes enough moisture from the 256 Sulphuric Acid. air to convert itself into sulphuric acid. If the crystals of sulphur trioxide are dropped into water they combine with great energy, making a hissing sound like a red hot iron going into water. If placed upon the skin it produces painful burns. Sulphuric Acid, H 2 SQ 4 , is the most important of all the sulphur compounds, and its manufacture is carried on on an enormous scale. The frequency with which we have used this powerful acid substance in order to bring about chemical decompositions cannot fail to have been noticed. It was used in the prepara- tion of hydrogen, chlorine, hydrochloric acid, nitric acid, carbon monoxide, and sulphur dioxide. Modes of Formation. Sulphuric acid is produced when sulphur trioxide is dissolved in water. S0 3 + H 2 = H 2 S0 4 . It is also produced slowly, by the gradual absorption of oxygen by a solution of sulphurous acid (see Sulphurous Acid). The process by which it is prepared on a manufacturing scale, consists in making sulphur dioxide combine with oxygen (from the air) in the presence of water (steam). S0 2 + O + H 2 = H 2 S0 4 . We have seen already that sulphur dioxide and oxygen do not combine very readily, but require help in order to make them unite. Therefore, if sulphur dioxide, oxygen, and steam were simply mixed together, they would be a very long time in uniting to form sulphuric acid. The compound employed to cause the sulphur dioxide to combine with oxygen is nitric oxide, NO. We have learnt already (p. 199) that when nitric oxide comes in contact with the air, it unites with another atom of oxygen, and forms nitrogen peroxide, NO^ a reddish gas. Now nitrogen peroxide easily gives up this extra atom of oxygen to sulphur dioxide in the presence of steam, and goes back again to nitric oxide. Therefore, when steam, sulphur Manufacture of Sulphuric Acid. 257 dioxide, and nitrogen peroxide are mixed, the following change takes place H 2 + S0 2 + N0 2 = H 2 S0 4 + NO. The nitric oxide that is thus formed instantly seizes another atom of oxygen from the air, again forming nitrogen peroxide, NO 2 , NO + O = NO 2 and this again hands on this extra oxygen to another portion of sulphur dioxide. The nitric oxide (NO) is, therefore, a sort of " middleman," who takes oxygen from the air and passes it on to the sulphur dioxide ; and the same quantity of nitric oxide can keep on doing this, and can convert an unlimited amount of sulphur dioxide into sulphuric acid. The Manufacture of Sulphuric Acid is carried on in enormous rooms made of sheet lead. These great rooms are called leaden chambers, and they are often of such a size that 250 people could sit down to dine in one. Generally several are placed in a row. Into these chambers are sent sulphur dioxide, air, nitrogen peroxide, and steam. (1) The sulphur dioxide is produced either by burning sulphur, or roasting iron pyrites, in special furnaces called sulphur burners or pyrites burners. In either case sulphur dioxide is formed, which, with the excess of air passing through the furnace, is drawn into the " chambers." S + O 2 = SO 2 . Sometimes the sulphur dioxide is obtained by burning sulphuretted hydrogen. (2) The nitrogen peroxide is produced by placing inside one of the sulphur burners a pot containing a little Chili saltpetre and sulphuric acid. This mixture when heated gives nitric acid (p. 189), and the fumes of nitric acid coming in contact with the sulphur dioxide are decomposed, yielding nitrogen peroxide, which passes on into the chambers. 2 HNO 3 -f SO 2 = H 2 SO 4 4- 2 NO,. s 258 Sulphuric Acid. (3) The air that is admitted into the chambers is what is allowed to pass through the pyrites burners, and its amount is regulated. (4) The steam is blown into the chambers in jets from steam boilers. When these gases mix in the chambers, the chief reactions which go on are represented by the equations given above, and sulphuric acid (moderately strong) collects on the floors and is drawn off. Care is taken to adjust the proportion of the various gases. As the oxides of nitrogen go on transferring oxygen from the air to the sulphur dioxide over and over again, it is only necessary to add a very small amount of the nitrogen peroxide from time to time, to make up for the slight loss of this gas which always takes place. 1 The acid which is drawn from the chambers (chamber acid} is boiled down, either in glass or platinum vessels, so as to drive off the water, and so get strong sulphuric acid. Properties. Pure sulphuric acid is a heavy, colourless, oily liquid (hence the common name oil of vitriol}. It is a powerfully corrosive substance, and if spilt upon the skin produces bad burns. Therefore some care must be taken in handling this acid. It has a strong affinity for water, and if mixed with water the mixture gets nearly boiling hot. On account of its power of combining witli water, it is constantly used for withdrawing water vapour from gases. Thus, when we require to dry a gas, that is, to remove the vapour of water from it, we bubble the gas through sulphuric acid, provided it is a gas which has no action on the acid. If this acid is exposed to the air, or is not kept in well-stoppered bottles, it quickly absorbs water vapour from the air, and, of course, by so doing gets more and more dilute. Its affinity for water is so great that it decomposes many compounds containing hydrogen and oxygen, and takes these elements away from the compound in the proportion to yield 1 For further details of the manufacturing process, see "Newth's Inorganic Chemistry." Sulphates. 259 water. Thus, in the case of oxalic acid (p. 233), C 2 H 2 O 4 . This is decomposed by sulphuric acid, which in its eagerness for water, abstracts from the oxalic acid the elements which yield water, H 2 and O, leaving just enough oxygen for the carbon to form carbon monoxide and carbon dioxide. Again in the case of sugar (p. 212), C^H^Ou. When sulphuric acid acts on this, it abstracts all the hydrogen and oxygen (which are present in exactly the proportion to give nH 2 O) and leaves the carbon in the free or uncombined state. Its power of charring organic matter may be shown by the following experiment. Experiment 246. Take some very dilute sulphuric acid, and with the finger write a word on a piece of paper. Now gently dry the paper by holding it at some distance above a gas flame. Notice that the paper is charred where the letters were drawn upon it. Sulphates. Just as carbonic acid and sulphurous acid form two classes of salts, so, for the same reason, there are two classes of sulphates. The reason being that sulphuric acid, like these others, contains two atoms of hydrogen which can be replaced by metals. Thus, by replacing the hydrogen atoms by potassium we get either normal potassium sulphate (or potassium sulphate), K 2 SO 4 ; or hydrogen potassium sulphate (potassium bi-sulphate), HKSO 4 . Certain of the sulphates were known to the very early chemists, and were called vitriols (because they had rather a vitreous or glassy appearance), such as blue vitriol (copper sulphate), green vitriol (iron sulphate), white vitriol (zinc sulphate). The name " oil of vitriol " is derived from the fact that the acid was formerly obtained by distilling green vitriol. The sulphates, like all other salts, are formed when the acid is neutralized with a metallic hydroxide. 2KHO + H 2 S0 4 = K 2 S0 4 + 2 H 2 0. They are also produced by the action of the acid on oxides and carbonates ZnO + H 2 SO 4 = ZnSO 4 + H 2 O; Na 2 CO 3 -f H 2 SO 4 = Na 2 SO 4 + CO 2 -f H a O. 260 Sulphuric Acid. Being such a powerful acid, sulphuric acid is capable of taking metals away from the salts of almost any other acid ; for example, from sodium chloride or potassium nitrate it takes the sodium or potassium, and gives its hydrogen in exchange. Thus 2 NaCl + H 2 SO 4 = Na 2 SO 4 + 2HC1 (p. 127); 2 KNO 3 + H 2 SO 4 = K 2 SO 4 + 2HNO 3 (p. 189). Test for Sulphates. (i) Sulphates which are soluble in water are recognized and distinguished by giving a white precipitate with barium chloride, consisting of barium sulphate. This white precipitate is insoluble in either hydrochloric or nitric acid. Experiment 247. Dissolve in three separate test-tubes a small particle of (i) potassium sulphate, (2) sodium sulphite, and (3) sodium carbonate. Add to each a few drops of barium chloride. Notice that a very similar precipitate is formed in each case, but in reality they are totally different one is barium sulphate, the next is barium sulphite, and the third is barium carbonate. To each add a few drops of strong hydrochloric acid. Notice that the barium sulphate is unaffected ; the barium carbonate quickly dis- solves with effervescence, giving off carbon dioxide ; while the barium sulphite also dissolves with effervescence, and gives off sulphur dioxide (which can be detected by the smell). [Probably in this case the precipitate will not wholly dissolve, because the sodium sulphite originally used is likely to contain a little sodium sulphate mixed with it, so that the precipitate obtained when barium chloride was added, consists partly of barium sulphite (which will dissolve in the acid) and barium sulphate which will not dissolve.] The reactions in this experiment are the following. () When barium chloride is added to the three solutions K 2 SO 4 + BaCl 2 = 2KC1 + BaSO 4 white precipitate Na 2 SO 3 + BaCl 2 = 2NaCl + BaSO 3 Na 2 C0 3 + BaCl 2 = 2 NaCl + BaCO 3 (b) The action of hydrochloric acid on the three precipitates BaSO 4 + 2HC1 no action. BaSO 3 + 2HCI = BaCl 2 + H 2 O + SO 2 BaCO 3 + 2HC1 = BaCl 2 + H 2 O + CO 2 . Tests for Sulphates. 261 (2) Sulphates which are not soluble in water are tested for in a different way. Experiment 248. Take a pinch of plaster of Paris (calcium sulphate) and mix it with about three times as much powdered sodium carbonate, and heat the mixture on a little piece of platinum foil, bent into a sort of spoon (or on the lid of a platinum crucible) until it has completely melted. At this high temperature the two compounds make a mutual exchange, resulting in the formation of sodium sulphate and calcium carbonate. CaSO 4 + Na 2 CO 3 = Na 2 SO 4 + CaCO 3 . The sodium sulphate is soluble in water, while the calcium car- bonate is not ; therefore, put the platinum spoon into water in a test-tube and boil it for a minute or two, and then filter it. The solution contains sodium sul- phate, and any excess of sodium carbonate which may have been used. Now add hydrochloric acid until the solution is quite acid, and then add barium chloride. The white precipitate of barium sulphate is at once formed. [In- stead of fusing the insoluble sul- phate with sodium carbonate, we may mix it with a strong solution of sodium carbonate and boil it for some time, when there will be enough of the calcium sulphate decomposed and sodium sulphate produced to give the test with barium chloride.] Experiment 249. Heat a little heap of a mixture of plaster of Paris and sodium carbonate on a piece of charcoal by means of a small blow-pipe flame, as shown in Fig. 108. [Select a good piece of charcoal, which is not full of cracks, and scoop a small hollow upon it to hold the substance which is being heated.] The result of this operation, is that the sodium of the sodium carbonate combines with the sulphur of the calcium sulphate to give sodium sulphide. When it is cold, place the fused residue on a silver coin and touch it with a drop of water. The sodium sulphide at once acts on the silver and causes a black stain upon it. FIG. 108. 262 Sulphuric Add. EPITOME. The two commonest oxides of sulphur are sulphur dioxide, SO 2 , and sulphur trioxide, SO 3 . Sulphur dioxide is a colourless choking gas, obtained when sulphur burns either in air or in oxygen. It is prepared by heating sulphuric acid and copper. It is more than twice as heavy as air, and therefore can be collected by downward displacement. It is soluble in water, and therefore cannot be collected at the pneumatic trough. Sulphur dioxide does not burn, nor support the combustion of ordinary burning bodies. Lead dioxide and sodium dioxide take fire in the gas and produce sulphates of the metals. Sulphur dioxide bleaches, but not always permanently ; the colour often being restored either by stronger acids or by alkalies. The gas is used for disinfecting purposes ; but its chief use is in the manufacture of sulphuric acid. Sulphur dioxide is easily condensed to a liquid by cooling it. The liquefied gas is colourless, and boils at 8. The solution of sulphur dioxide in water is acid, and contains sulphurous acid, H 2 S0 3 . This acid has never been obtained except as a solution in water. When the solution is boiled, the acid is decomposed and sulphur dioxide escapes. The salts of sulphurous acid are called sulphites. The acid is dibasic, and therefore forms two classes of salts ; those in which all the hydrogen has been replaced by metals, and those in which only half the hydrogen is so replaced. Thus : Na 2 SO 3 , di-sodium sulphite, or normal sodium sulphite ; and HNaSO 3 , hydrogen sodium sulphite (sometimes called add sodium sulphite, or sodium bisulphite) . The sulphites are decomposed by hydrochloric or sulphuric acid, with the evolution of sulphur dioxide. Sulphur trioxide is a white solid, forming silky crystals. It is produced when sulphur dioxide and oxygen are passed over heated spongy platinum. This oxide has a powerful affinity for water, with which it combines to form sulphuric acid. Sulphuric acid (or oil of vitriol) is a colourless oily liquid, strongly corrosive, and a powerful acid. Its manufacture is the most important of all chemical industries. It is made by oxidizing sulphur dioxide by means of nitrogen peroxide in the presence of steam. Sulphur dioxide combines with atmospheric oxygen, only with extreme slowness ; but, by means of nitric oxide, oxygen is Sulphuric Acid. 263 taken from the air and handed on to the sulphur dioxide. Nitric oxide unites with oxygen of the air, forming nitrogen peroxide, and this gives oxygen to the sulphur dioxide, and is again reduced to nitric oxide. The operation is carried on in enormous chambers built of lead. Sulphuric acid combines with water with the production of great heat. On account of its eagerness to unite with water it is used for drying gases. It also decomposes many organic compounds con- taining oxygen and hydrogen, withdrawing these two elements in the proportion required to form water. Sulphuric acid is dibasic, and forms two classes of salts, accord- ing as to whether all, or only half, the hydrogen of the acid is replaced by metals. CHAPTER XXX. SOME COMMON CARBON COMPOUNDS. CARBON forms such an enormous number of compounds that a simple list of they: names alone would more than fill this little book. The study of these compounds is generally called " organic chemistry," because, in the early days of chemistry, it was thought that these compounds could only be produced as the result of living organized bodies, like animals and plants. In order to study this vast host of compounds, they are divided and subdivided into classes and families much in the same way as animals or plants are classified. Four very important classes are : 1. Hydrocarbons. 2. Acids. 3. Alcohols. 4. Carbohydrates. 1. Hydrocarbons. These, as the name implies, are com- pounds of carbon with hydrogen only. Important amongst these are the following : Marsh Gas, CH 4 . Found in coal mines, and called fire-damp. Also in marshy places where vegetable matter is rotting. In the laboratory it is made by strongly heating a mixture of sodium acetate and caustic soda, when sodium carbonate is left behind. NaC 2 H 3 O 2 + NaHO = Na 2 CO 3 + CH 4 . Marsh gas is colourless, and has no smell. It burns easily, but gives very little light, although rather more than hydrogen. Common Carbon Compounds. 265 If mixed with air, or with pure oxygen, and fired, the mixture explodes. This is the cause of coal-mine explosions. Water and carbon dioxide are produced, and the latter gas is called choke-damp by the miners. CH 4 + 3O = 2H 2 O + COo. Ethylene, C 2 K 4 (olefiant gas), is obtained from common alcohol (spirits of wine) by heating it with sulphuric acid (see Exp. 184). C 2 H 6 - H 2 = C 2 H 4 Ethylene is a colourless gas with a faint pleasant smell. It burns easily, with a very bright flame, producing carbon dioxide and water, the same products as are formed when any hydro- carbon burns. Acetylene, C S H 2 , is formed whenever coal gas burns without a sufficient supply of air. Thus, when a Bunsen lamp gets alight down at the little jet at the base of the chimney, some acetylene is produced. This it is which causes the bad smell we notice when the lamp so burns. Acetylene is best prepared in the laboratory by acting on calcium carbide with water. A little of the carbide, in small lumps, is put into a test-tube, and a few drops of water added The gas is at once evolved, and can be lighted at the mouth of the tube. CaC 2 + 2H 2 O - CaH 2 O 2 + C 2 H 2 . Acetylene burns with a very bright and smoky flame. Marsh gas, Ethylene, and Acetylene are all present in ordinary coal gas. Other important hydrocarbons are the various mineral oils (paraffin oils) used for illuminating purposes. Paraffin wax, used for candles. Turpentine. 2. The Acids. This is a large and important class, divided into many families. Amongst the most important of these is the one known as the acetic, or fatty series of acids. These include acids present in many fats. Formic Acid (CH 2 O 2 ) is found in ants, and in the hairs 266 Common Carbon Compounds. of stinging-nettles. When we are stung by ants or nettles, it is because a minute drop of this formic acid has been injected into the skin. If a piece of litmus paper is placed upon an ants' nest which has been just disturbed, it is instantly reddened by a shower of formic acid being squirted against it by the irritated ants. Formic acid is decomposed by strong sulphuric acid into water and carbon monoxide. Acetic Acid i(C 2 H 4 O 2 ) is the acid of vinegar. When beer goes sour it is owing to the formation of acetic acid, and this change is brought about by the agency of a living organism familiarly known as mother of vinegar. Acetic acid is also produced when wood is destructively distilled, the material so obtained being called pyroligneous arid, that is, the fire-wood acid. Pure acetic acid is liquid at the temperature of a warm room ; but in winter it freezes to an ice-like solid. It is on this account called " glacial " acetic acid. This strong acid stings the skin as formic acid does. The salts are called acetates ; one of the commonest is lead acetate, familiarly known as sugar of lead. Butyric Acid is the name of the acid which is present in rancid butter, and which gives to it the disagreeable smell. Palmitic Acid and Stearic Acid are important con- stituents of most solid animal fats, such as beef and mutton suet, butter, and human fat. Palmitic acid is also one of the chief constituents of palm oil (hence its name). These two acids are extensively used for making candles. 3. Alcohols. This class contains a number of very important compounds. The following may be taken as examples. Methyl Alcohol (wood spirit}, CH 4 O or CH 3 HO, is present in the watery liquid obtained when wood is distilled. When it is extracted from this liquid and purified, it is a colourless liquid, which burns with a flame without light, and without any smoke. Ethyl Alcohol, C 2 H 6 O or C 2 H 5 HO, is the familiar spirits of wine. It is the best known of all the alcohols, and is therefore called simply " alcohol." Methylated Spirit. 267 It is obtained by the fermentation of ordinary sugar or of grape sugar (glucose) by means of yeast. The yeast organism transforms the sugar into alcohol and carbon dioxide (see Exp. 195). C 6 H 12 6 = 2C 2 H 6 + 2 C0 2 . Glucose. Alcohol. The liquid obtained contains only a small proportion of alcohol mixed with a large quantity of water. It has, there- fore, to be distilled, or "rectified," in order to separate the alcohol from the water. Ethyl alcohol is present in all fermented liquors, and it is the presence of this compound which gives them their intoxi- cating properties. Ordinary beer contains from 3 to 6 per cent, of alcohol. [If some beer is boiled in a flask with a long upright tube fastened into the neck with a cork, the alcohol which first boils off can actually be lighted as it escapes from the tube.] Light wines, such as claret, etc., contain from 8 to 14 per cent.; port and sherry 15 to 25 per cent, while brandy and other " spirits " contain from 50 to 60 per cent of alcohol. Alcohol is capable of dissolving things like resins, gums, oils, etc., and is therefore a most useful substance for the manufacture of varnishes, and for other purposes. Methylated Spirit. On account of the high duty upon "alcohol," it is too expensive for most of the manufacturing and chemical purposes for which alcohol is required. There- fore a mixture consisting of 90 per cent. " spirits of wine " and 10 per cent. " wood spirit " (impure methyl alcohol) is used instead of pure " alcohol." This mixture called methylated spirit is quite unfit for drinking, and, being duty free, is quite cheap. Phenol or Phenyl Alcohol, C 6 H 5 HO. This substance is familiarly known by the name carbolic arid, although in reality it is not an acid, but belongs to the class of alcohols. It is formed when coal is distilled for making coal-gas, and is extracted from the coal-tar oil. Pure phenol forms long needle-shaped crystals, which melt a little above the temperature of the hand (42). It has 268 Common Carbon Compounds. a sharp burning taste and is poisonous. It is a powerful disinfectant and antiseptic, and is largely used in surgery. The common " carbolic acid powders " sold in tins for disinfecting purposes, consist of some powder such as lime or gypsum, impregnated with about 10 per cent, of very crude carbolic acid. Glycerin, C 3 H 5 (HOJ 3 . This familiar substance is also an alcohol. Chemists call it glycerol. It is an important constituent of fats. Fats are compounds of glycerin with such acids as palmitic and stearic. Mutton suet, for instance, consists chiefly of a compound of stearic acid and glycerin. This compound is called stearin. Palmitin is the name of the compound of palmitic acid with glycerin. Glycerin is obtained from fats by heating them in very hot steam, or by boiling them with a caustic alkali such as sodium hydroxide. The alkali combines with the stearic or palmitic acid forming a salt (soap), which separates out when the liquid cools ; while the glycerin which is set free remains dissolved in the watery liquid. Glycerin is a thick syrupy liquid, with a very sweet taste, almost like sugar syrup. Soap. When fats are boiled with caustic alkalies, the fat (which is a compound of fatty acids with glycerin) is decom- posed ; glycerin is set free and the fatty acid combines with the alkali. This process is called saponification ; and the compound of the fatty acids (chiefly stearic and palmitic acids) with the alkali, is known as a soap. When caustic potash is the alkali used soft soap is pro- duced ; while if caustic soda is employed the soap is harder, as in the ordinary forms of soap used for washing purposes. In some out of the way parts of the world, people often make a crude kind of soap by collecting the ashes of burnt wood (these contain potash, the name potash simply mean- ing pot-ashes), mixing them with water, and boiling the liquid so obtained with mutton or beef fat The Action of Hard Waters upon Soap. The property of hardness in water is chiefly due to the presence of either carbonate of lime or sulphate of lime (see p. 81). Soap. 269 When water containing these salts in solution comes in contact with a solution of soap (say sodium stearate), a chemical action takes place, thus Sodium stearate 4- calcium carbonate = calcium stearate -f (Insoluble.) sodium carbonate. The calcium stearate, being insoluble, separates out as a greasy scum, which is always seen when hard water is used for washing. So long as any of the lime salts remain in the water, the soap is used up in bringing about this double decomposition, and, therefore, is wasted so far as its powers of cleansing are concerned ; for we cannot wash with soap until the water has become softened by the removal or destruction of the hardening lime salts. As soon as ever all the calcium carbonate or sulphate has been decomposed by the soap, then, and not till then, will the soap give a lather ; and not until 'a lather can be produced is the soap of any service for washing. Experiment 250. Take two good-sized stoppered bottles, and about half fill one with some hard water (common tap-water will generally do, but if this should happen to be a soft water, a sample of hard water can readily be made by adding a little lime-water, and then bubbling carbon dioxide through it). At first, calcium carbonate CaCO 3 is precipitated, but presently this dissolves in the carbonic acid, forming the bi-carbonate of calcium H 2 Ca(CO 3 ) 2 , or CaCO 3 ,H 2 CO 3 . Into the other bottle put an equal quantity of distilled water, or rain water. Now make a solution of soap by shaking up a few thin shavings of soap with a little water in a bottle. Add a small quantity of this soap solution at a time to the soft water, until on shaking the bottle a lather is raised which does not disappear again. Now add the soap solution to the hard water. Note that at first no lather is produced, although much more of the soap is added. Also observe that the solution turns muddy, and a scum is produced. This is the calcium stearate being precipitated ; continue adding the soap until at last a per- manent lather is obtained. If a little more hard water or a few drops of a solution of any lime salt be now added, the lather instantly disappears. 270 Common Carbon Compounds. 4. Carbohydrates. This family of carbon compounds includes the sugars, starches, etc. The following are some common examples of carbohydrates. Cane-sugar (Saccharose), C^H^On. This is ordinary sugar, and is obtained chiefly from sugar-cane and from beet-root. It is present in almost all sweet fruits. When cane-sugar is brought into contact with yeast, it is converted into dextrose and laevulose. The same change takes place when it is warmed with dilute sulphuric acid. C 12 H 22 O n + H 2 - C 6 H 12 6 -fC 6 H 12 6 . Cane-sugar. Dextrose. Lsevulose. Grape-sugar (Glucose or Dextrose], C 6 H 12 O 6 , is obtained from the juice of sweet grapes. It is also present in honey. It is not so sweet as cane-sugar, and is less easily dissolved by water. Grape-sugar is distinguished from cane-sugar by the following test : Experimental. Take a crystal of copper sulphate, and one about the same size of tartaric acid. Dissolve them together in a little water, and add a solution of caustic potash until the liquid is strongly alkaline. Add a little of this mixture to a solution of grape-sugar in a test-tube, and warm the liquid. A red precipitate of cuprous oxide is formed, and the solution loses its blue colour. Do the same with a solution of common sugar ; this does not give the red precipitate, and the liquid remains blue. Starch, C 6 H 10 O 5 , occurs in many parts of plants, such as the seeds, stems, roots, and tubers, and it may be obtained from them by crushing them with water, and separating the bruised fibre or pulp by means of a sieve. Experiment 252. Take a raw potato and rub it down on a coarse grater, and collect all the pulp on a piece of muslin. Screw it up into a sort of bag, and squeeze it into a small basin of water ; dipping it once or twice into the water and squeezing again. In this way the starch passes through the muslin into the water, which is thereby made milky. If this is allowed to stand, the clean white starch settles to the bottom. If wheat meal or flour is treated in a similar manner, starch is also separated, and a whitish sticky mass is left. This is called gluten. It is a compound containing Starch. 27 1 nitrogen, and it is the presence of this substance which gives to the wheat its value as a food. Wheat flour contains about ^th of its weight of gluten. When starch is boiled with dilute sulphuric acid, or is acted on by the ferment present in germinating barley (diastase), it is converted into glucose and dextrin. 3C 6 H 10 8 + H 2 = C 6 H 12 6 + 2C 6 H 10 5 . Starch. Glucose. Dextrin. When examined through a microscope, starch is found to consist of minute granules, not crystals. The size and shape of these vary considerably. Those of potato starch and arrowroot are much larger than those from rice. Starch gives a blue colour with iodine (see p. 134). Starch does not dissolve in cold water, but with boiling water the granules swell up and burst. If the amount of hot water is very large, the starch disappears and seems to dis- solve, but with less water it forms a gelatinous or pasty mass. Dextrin. When starch paste is boiled with a little dilute sulphuric acid, it soon becomes thinner and thinner, being changed into glucose and dextrin, both of which are soluble. Dextrin is a sticky gummy substance, often used instead of ordinary gum or paste, for mounting photographs and other similar purposes. It is sold under the name of British gum. CHAPTER XXXI. SIMPLE QUALITATIVE ANALYSIS. THE word "analysis" means the breaking up or separation of a compound into its components or elements (see p. 75). But the word is also used in a broader sense, and is applied to any processes or methods by which the chemist is able to find out what a substance is composed of, in order to identify that substance. For example, we speak of " micro- scopic analysis" and spectrum analysis ; these are not^ pro- cesses in which compounds are split up into their components, but are methods which enable chemists to identify different substances. The following illustration will make this plain. Experiment 253. Take a crystal of nitre (potassium nitrate), and just touch the edge of a Bunsen flame with it. Notice the lilac or violet colour which it imparts to the flame. This colour is characteristic of all potassium compounds (compare Exp. 44). Place the crystal on a glass plate, dissolve it in a drop of warm water, and allow the solution to evaporate. Examine the crystals which are formed with a pocket lens or microscope. Notice the long prismatic shaped crystals, characteristic of nitre. Dissolve a crystal of the salt in water, and apply the test for a nitrate, as explained on p. 195. By these three experiments or tests, we have analysed this substance, and identified it as potassium nitrate, although we have not separated the compound into its constituent elements. Separation Usually Necessary. In a great many instances, the tests which are used to identify a substance are interfered with if certain other substances are present at the same time. In all such cases it is necessary to separate Reagents. 273 the substances before applying the special test. For example Experiment 254. Dissolve a small crystal of potassium chloride in a drop or two of water, and add to it one or two drops of a solution of platinum chloride (PtQ 4 ). A precipitate is produced, consisting of tiny yellow crystals. This is a characteristic test for potassium compounds. Now treat a similar quantity of ammonium chloride in exactly the same way ; notice that a precisely similar looking yellow crystalline precipitate is formed. Therefore, before applying this test for the detection of a potassium salt, it is absolutely necessary to separate it from any ammonium salts. How Separation is made. The method most fre- quently used for bringing about analytical separation, is to cause one or more of the compounds present to ^undergo a double decomposition with certain chosen reagents, whereby the metals in these compounds form fresh compounds which are insoluble, in water, and which are, therefore, thrown down as precipitates. An example of this method is given on p. 20, Exp. 25. The silver nitrate present is caused to enter into double decomposition with sodium chloride. This results in the formation of silver chloride, which is precipitated as an insoluble compound, and sodium nitrate remains in solution along with the copper salt. It must be noted that in this separation it is only the silver from the silver nitrate that is actually separated from the copper compound, for the other part of the compound is left in solution combined with sodium instead of with silver. By this process we have not separated silver nitrate from copper nitrate, but only withdrawn the silver and replaced it by sodium. Reagents. The solutions which are used to separate substances in this way are called reagents. Sometimes the same reagent will form insoluble compounds with a whole group of metals, in which case it may be used to separate an entire family of metals from others not belonging to the group. Such reagents are called group-reagents. 274 Simple Qualitative Analysis. Groups. For convenience the metals are divided into a number of groups, based upon their behaviour towards certain chosen group-reagents, used in a certain order. GROUP I. or HYDROCHLORIC ACID GROUP. Metals whose chlorides are precipitated on the addition of hydrochloric acid. Lead 1 (silver, mercury). GROUP II. or SULPHURETTED HYDROGEN GROUP. Metals whose sulphides are precipitated from acid solutions by sulphuretted hydrogen. Lead, 2 Copper (mercury, bismuth, cadmium, tin, arsenic, antimony). GROUP Ilia, or AMMONIA GROUP. Metals whose hydroxides are precipitated by ammonia, in the presence of ammonium chloride. Iron (chromium, aluminium). GROUP lllb. or AMMONIUM SULPHIDE GROUP. Metals whose sulphides are precipitated by ammonium sulphide, in the presence of ammonia. Zinc (manganese, nickel, cobalt). GROUP IV. or AMMONIUM CARBONATE GROUP. Metals whose carbonates are precipitated by ammonium carbonate, in the presence of ammonium chloride. Calcium (barium, strontium). GROUP V. No group reagent. Potassium, Ammonium (sodium, magnesium). Reactions for the Metals. In order that we may be able to recognize and identify a metal in the various com- pounds it produces with different reagents, it is necessary to make ourselves quite familiar with these compounds. In order to gain this knowledge, the following reactions or tests should be carefully made, and the student should make exact notes of all he does and observes. If any experiment he makes seems to give a different result from that which is 1 Those metals printed in thick type are the only ones the elementary student will be concerned with. 2 The reason why lead is in both groups I. and II. is explained on page 276. Reactions for Lead. 275 indicated in the book, he should not pass it over, but should repeat it more carefully. Lead, Pb. (Group I.) Use lead nitrate, Pb(NO 3 ) 2 . Take a few crystals of the salt ; note their dry, hard, milk-white appearance. Dissolve a little in water in a test-tube. Note that it is not very readily dissolved in cold water, but more quickly if the water is warmed. Pour some of this solution into five separate test-tubes. I. To one, add a few drops of dilute hydrochloric acid (group-reagent). Note the white, curdy precipitate of lead chloride. Shake the tube, and, when the precipitate has settled, add more of the reagent, until no further precipitation takes place. Pb(NO 3 ) 2 4- 2HC1 = PbCl 2 + 2HN0 3 . Now heat the mixture, and observe how the precipitate disappears. Lead chloride is soluble in hot water. This distinguishes lead from the other metals of this group. Cool the test-tube again, and the lead chloride is again precipitated ; but note that it comes down in the form of shining white needle-shaped crystals. II. Take the second portion of the lead nitrate solution, and pass sulphuretted hydrogen through it (see p. 246), or add sulphuretted hydrogen water. A black precipitate of lead sulphide is formed. Pb(N0 3 ) 2 + H 2 S = PbS + 2 HN0 3 . Filter the liquid, scrape a little of the precipitate off the filter into a test-tube, add a few drops of strong nitric acid, and boil. Notice that the black precipitate turns white. The lead sulphide is oxidized by the nitric acid into lead sulphate. III. Treat the third portion as the first, and filter the mixture. Then pass sulphuretted hydrogen into the clear filtrate. Note a black precipitate, as in II. This shows that the group reagent does not entirely separate lead from group II. ; that is to say, lead chloride is slightly soluble 276 Simple Qualitative Analysis. even in cold water, and, therefore, a portion of it passes through along with the metals of group II. (see p. 274). IV. To the fourth portion add potassium chromate. Note the yellow precipitate of lead chromate. Pb(NO 3 ) 2 + K 2 Cr0 4 = PbCrO 4 + 2 KNO 3 . V. To the fifth portion add dilute sulphuric acid ; a white granular precipitate of lead sulphate is produced. Pb(NO 3 ) 2 + H 2 SO 4 = PbS0 4 + 2 HNO 3 . Dry Reaction. Powder a small crystal of lead nitrate, mix about twice as much sodium carbonate with it, and place the mixture in a shallow cavity scooped out on a piece of charcoal. Heat the mixture by means of a blow-pipe flame (see Fig. 108, p. 261), holding the charcoal in such a position that the middle or inner part of the flame, and not the tip of it, plays upon the mixture. [The outer part of the flame, where the oxygen of the air is in excess, is called the oxidizing flame ; while the inner portion, where heated coal-gas is in excess, is known as the reducing flame^\ The mixture quickly melts, and the lead compound is reduced to the state of metallic lead, which will appear in the form of small brilliant globules upon the charcoal. At the same time, some of the lead is oxidized, and oxide deposits round the cavity as a yellow incrustation. Allow the mass to cool ; pick out one of the globules of metal with a penknife, and show that it is soft and malleable by hammering it ; also that if rubbed across paper it leaves a black mark. Copper, Cu. (Group II.) Use copper sulphate, CuSO 4 ,5H 2 O. Dissolve a little of the salt in water. Note that it is readily soluble. I. Take a small portion of the solution and pass sulphu- retted hydrogen (the group-reagent). Notice the brownish- black precipitate of copper sulphide. CuSO 4 -f H 2 S = CuS + H 2 SO 4 . Filter the mixture, and show that the precipitate dissolves in boiling nitric acid, giving a bluish solution. (Compare the Reactions for Iron. 277 behaviour of lead sulphide.) Cautiously add ammonia to this solution, and note the deep azure-blue colour of the liquid. II. To a second portion of the copper sulphate solution add ammonia, drop by drop. Notice the pale greenish-blue precipitate. Add more ammonia, and the precipitate quickly dissolves, forming a deep blue solution, characteristic of copper compounds. Iron, Fe. (Group 1 110.) This element forms two classes of compounds, which behave quite differently towards reagents. One of these classes is derived from ferrous oxide, FeO, and the other from ferric oxide, Fe 2 O 3 , in which the iron is combined with a larger proportion of oxygen, or is in a higher state of oxidation, as we say. These two classes of iron compounds are therefore distinguished as ferrous and ferric salts. The former are mostly pale green (Exp. 51, p. 46) or white, while the ferric salts are generally yellow. Ferrous are readily converted into ferric salts by the action of oxidizing agents ; while reducing agents change ferric back to ferrous compounds. Thus, if sulphuretted hydrogen is passed through a neutral or acid solution of ferric chloride, ferrous chloride is produced and sulphur is precipitated. FeoCl 6 + H 2 S = 2FeCl 2 + 2HC1 + S. Use ferrous sulphate, FeSO 4 ,7H 2 O, and ferric chloride, Fe 2 Cl 6 . Make a solution of ferrous sulphate by dissolving two or three crystals of the salt in cold water. I. Add ammonia (group-reagent), and obtain a dirty greenish precipitate of ferrous hydrate, which on exposure to the air, gradually becomes oxidized into ferric hydrate (brown). FeS0 4 + 2 NH 4 HO - Fe(HO) 2 + (NH 4 ) 2 SO 4 . Repeat with a solution of ferric chloride ; a reddish-brown precipitate of ferric hydrate is formed. Fe 2 Cl 6 -I- 6NH 4 HO = Fe 2 (HO) 6 + 6NH 4 C1. 278 Simple Qualitative Analysis. II. Add ammonium sulphide to another portion of the ferrous sulphate solution. Black ferrous sulphide is pre- cipitated. FeSO 4 + (NH 4 ) 2 S = FeS + (NH 4 ) 2 SO 4 . Repeat with ferric chloride. The same precipitate is obtained, mixed with sulphur. Fe 2 Cl 6 + 3(NH 4 ) 2 S = 2FeS + S + 6NH 4 C1. III. Add a few drops of ammonium thiocyanate (frequently wrongly called ammonium sulphocyanide) to the ferrous sul- phate. No change takes place. Repeat with ferric chloride. An intense blood-red coloured solution is produced. IV. Add a few drops of potassium ferri cyanide to the ferrous sulphate. A dark blue precipitate is formed (called Turnbull's blue). Repeat with ferric chloride. No blue colour, but the mixture becomes brownish. V. Add potassium ferrocyanide to ferrous sulphate, a light blue precipitate is formed, which on exposure to air becomes darker blue. Repeat with ferric chloride. A dark blue precipitate re- sults (called Prussian blue). Zinc, Zn. (GROUP III.) Use zinc sulphate, ZnSO 4 ,7H 2 O. Dissolve a few crystals in water. Note the appearance of the crystals and their ready solubility. I. Add ammonium sulphide (group-reagent). A white precipitate of zinc sulphide. ZnSO 4 + (NH 4 ) 2 S = ZnS + (NH 4 ) 2 SO 4 . II. Add ammonia drop by drop to another portion. Note that a white precipitate is at first produced ; which, however, readily dissolves as more ammonia is added. (This reaction enables us to separate zinc from iron.) III. Dry Reaction. Heat a little solid zinc sulphate with sodium carbonate upon charcoal in the inner blowpipe- Reactions for Calcium. 279 flame. No metallic beads are formed, because zinc oxidizes too easily ; but an incrustation of zinc oxide is formed on the charcoal, which appears canary yellow while it is hot, but turns white on cooling. Touch the incrustation with a drop of cobalt nitrate solution, and again heat it in the extreme tip of the flame. T he mass becomes green. Calcium, Ca. (GROUP IV.) Use calcium chloride, CaCl 2 ,6H 2 O. Dissolve some of the salt in water. Note how quickly and easily it dissolves. I. Add ammonium carbonate (group-reagent) to a portion of the solution. A white precipitate of calcium car- bonate is obtained. CaCl 2 + (NH 4 ) 2 C0 3 - CaC0 3 + 2NH 4 CL Filter the mixture, and pour a few drops of acetic acid upon the filter. Note effervescence as the precipitate dissolves. II. Add ammonium oxalate to another portion. A white precipitate of calcium oxalate is obtained. CaCl 2 + (NH 4 ) 2 C 2 4 - CaC 2 4 + 2 NH 4 C1. Pour half this mixture into a second test-tube. To one portion add hydrochloric acid. Note that the precipitate dissolves ; to the other add acetic acid. The precipitate does not dissolve. Dry Reaction. Dip a clean platinum wire into the calcium chloride solution, and bring it against the edge of a Bunsen flame. Note the reddish colour of the flame. Potassium, K. (GROUP V.) Use potassium chloride, KC1. Dissolve a few crystals in a small quantity of water so as to obtain a strong solution. I. Add platinum chloride (one or tw$ drops) to a small quantity of this solution. A yellow crystalline precipitate is formed, consisting of the compound K 2 PtCl 6 . Cautiously add water drop by drop, and gently warm the mixture. Notice that the precipitate dissolves. Hence this test can only be used when the solution is strong. Now add 280 Simple Qualitative Analysis. a little strong hydrochloric acid. This causes the reformation of the precipitate. Therefore, before applying this test, it is best to add a drop of strong hydrochloric acid, as this promotes the formation of the precipitate. II. Add to a second portion of potassium chloride a little strong solution of hydrogen sodium tartrate, a white precipitate of hydrogen potassium tartrate is formed. Add a little water, and note that the precipitate quickly dissolves. Therefore this test also can only be made with strong solutions. Dry Reactions. Dip a clean platinum wire into the potassium chloride solution and bring it into the Bunsen flame. A lilac colour is given to the flame. This colour, however, is entirely masked by a yellow colour if any sodium compounds are present, even in the minutest quantities. If the flame be looked at through deep blue glass (or better a glass prism filled with indigo) the colour given by the potassium salt will appear crimson-red in spite of the sodium impurity. AMMONIUM COMPOUNDS. These all evolve ammonia, when heated with caustic soda or potash, which is detected by its smell, and by its action on turmeric or litmus paper. See p. 183. Reactions for Acids. The four commonest acids are hydrochloric, nitric, carbonic, and sulphuric. The reactions by which these are distinguished have been already described. Chlorides, p. 125 ; Nitrates, p. 195 ; Carbonates, p. 230, and sulphates, p. 261. Method of analysis of a simple mixture. 1 If the substance given is solid, its general appearance (such as its colour, whether crystalline or not, etc.) should be carefully observed aod noted down. Then proceed to make the following tests directly upon the solid. I. Put a little of the substance into a short, narrow test- tube and heat it. Observe closely what happens. 1 Containing only chlorides, nitrates, carbonates, or sulphates of ammonium, potassium, calcium, zinc, iron, copper, or lead. Preliminary Tests. 281 (a) If a white sublimate is formed, suspect ammonium salts. Confirm this by heating a little of the salt with caustic soda. Ammonia given off, detected by its smell, and its action on turmeric paper, proves the presence of ammonium salts. (b) If a brown-coloured vapour is given off, suspect a nitrate, probably lead nitrate. Confirm a nitrate by applying the test for nitric acid, given on p. 195. II. Heat a little of the substance (powdered) on charcoal in the blow- pipe flame (as shown on p. 261). (a) If the substance deflagrates, appearing to go on fire on the charcoal, suspect a nitrate. (U) If a bright white residue is left, suspect calcium or zinc compounds. Place a portion of the residue upon turmeric paper and moisten with water. If a brown stain is produced, suspect calcium. Moisten a part of the residue on the charcoal with cobalt nitrate, and again heat. If a green residue remains, suspect zinc. (c) If metallic beads are formed, suspect lead. Confirm by mixing the substance with sodium carbonate, and heating on charcoal in the inner blow-pipe flame. Mal- leable bead, which marks paper, confirms lead. III. Heat a little of the substance on platinum wire in a Bunsen flame. (#). If reddish flame, suspect calcium. I These must be con- ($) If lilac flame, suspect potassium. \ firmed later. IV. Add a little dilute hydrochloric acid to some of the original substance, in a test-tube. (a) If effervescence takes place, suspect a carbonate. Confirm by allowing the gas to enter a tube containing lime water (as shown on p. 227). V. Gently heat a little of the substance with a few drops of strong sulphuric acid. (a) If acid fumes are evolved, suspect a chloride or a nitrate. Dip a glass rod, moistened with silver nitrate, into the mouth of the test-tube. 282 Simple Qualitative Analysis. A white precipitate on the rod indicates hydrochloric acid. Drop a small fragment of copper into the tube. Brown fumes of oxides of nitrogen indicate nitric acid. (Hydrochloric acid may also be confirmed by heating the substance with manganese dioxide and sulphuric acid, when chlorine will be evolved, which is recognized by its colour and its bleaching properties.) Prepare a solution of the substance. Place a little of the powder in a test-tube, add water, and gently warm the mixture. If it does not dissolve, add hydrochloric acid, 1 and again boil. When the solution is obtained, proceed according to the following Table ; except that if hydrochloric acid has been used to dissolve the substance the first step is omitted. 1 The only compounds of the prescribed metals and acids (see note, p. 280) which are not soluble in water are the carbonates of lead, copper, iron, and calcium ; and the sulphates of lead and calcium. The carbonates dissolve readily in dilute hydrochloric acid, with evolution of carbon dioxide : in the case of lead carbonate, lead chloride is formed, which dissolves on boiling, but crystallizes out again on cooling (see Reaction I., p. 275). Lead sulphate and calcium sulphate are partially dissolved by boiling water. ( 283 ) ph S s f IS o to pa o "S 2 CQ ^ y^ o o "H, 1 | '5 o imo Iter 'd ss nj a; XI HflH^clil" llljjlilltaii ^ *- oo c W but exclusive of platinum wire. i Winchester Ammonia (sp. g. 0-880). ., Hydrochloric acid (pure). Nitric acid (1-42 pure). Sulphuric acid (pure com.). i-lb. bottle Alcohol (pure abs.). Ammonium sul- phide, lib. Ammonium chloride (pure). Lime. Manganese dioxide. Marble. Potassium chlorate. nitrate. Soda, caustic, sticks. Sodium carbonate (pure). Sulphur, roll. Zinc, granulated. I Ib. Alum. Ammonium carbonate. nitrate. sulphate. Barium chloride. Calcium chloride (cryst.). Copper foil. gauze. sulphate. Ferrous sulphate. Ferric oxide. ,, Glycerin. Iron pyrites. filings. Lead acetate. dioxide. nitrate. Mercury. Mercuric oxide (red). lib. Oxalic acid. Potash, caustic. Potassium chloride. dichromate. Red lead. ,, Sodium nitrate. nitrite. Sulphur, flowers. Zinc sulphate, i oz. Ammonium thiocyanate. Aniline blue (or magenta). Arsenious oxide. Bromine. Calcium carbide. Cobalt chloride. Copper chloride. nitrate. oxide. Ferric chloride. Iodine. Litmus. Magnesium carbonate. Mercuric chloride. ,,, Phosphorus. Potassium iodide. ferricyanide. ferrocyanide. permanganate. Sodium. oleate. peroxide. sulphite. Tartar emetic. Tartaric acid. oz. Silver nitrate. I oz. Magnesium. Potassium. i drachm Platinum chloride, i book Gold leaf. 75 grains Platinum wire (No. 29 standard wire gauge). INDEX ABSOLUTE zero, 102 Atmosphere, composition of, 168 Absorption of gases by charcoal, 215 Atomic symbols, 152 Acetylene, 265 weights, 148 Acid, acetic, 266 Atoms, 149 , carbonic, 225 Avogadro's hypothesis, 162 , formic, 265 , hydriodic, 134 BALANCE, 92 VXT/1-rtV\1-r\TV*l/^ T 1 A Barometer, 89 "j iiyuruuruiiiiCj 1,^4. , hydrochloric, 120 Bases, 69 , hydrofluoric, 134 Basic oxides, 63 nifvir* yQQ Bending glass, 30 , mine, ioo , nitrous, 203 Black-lead, 214 , oxalic, 233 Bleaching powder, 132 , stearic, 266 Blue vitriol, 82 , sulphuric, 256 Boiling point of water, 80 , sulphurous, 253 Bone black, 217 Acid-forming oxides, 63 Bordering glass tubes, 35 Acids, monobasic, 193 Boring corks, 33 Affinity, chemical, II Boyle's law, 98 Alcohols, 266 Erin's process (oxygen), 57 Alkalies, 60, 66 Bromine, 133 Allotropy, 207 Burette, 93 Ammonia, 180 Ammoniacal liquor, 179 CALORIE, 175 Ammonium chloride, 182 Cane sugar, 270 nitrate, 182 Carbohydrates, 270 nitrite, 177 Carbon, 210 Analysis, simple, 272 A f\* Animal charcoal, 217 dioxide, 220 Aqua fortis, 192 monoxide, 231 Aqua regia, 192 Carbonates, 226 Argon, 1 66 Carbonic acid, 22^ Atmosphere, 166 Catalysis, 14 286 Index. Cavendish's experiment, 72 Chalk, 211 Chamber acid, 258 Charcoal, 214 , absorption of gases by, 215 , animal, 217 Charles' law, 97 Chemical action, 12 , modes of, 155 affinity, 1 1 change, 6 combination, laws of, 117 equations, 154 formulae, 153 symbols, 152 Chili saltpetre, 180 Chlorides, tests for, 126 Chlorine, 128 Coal, 217 Coke, 216 Collecting gases, 24 Combustion, 170 , gain in weight by, 64 , heat of, 174 , supporter of, 52 , temperature of, 172 Compounds, 8 Copper nitrate, 16 pyrites, 236 , tests for, 276 Crith, 101 Critical temperature, 103 Crystallisation, 120 Cubical nitre, 194 DECANTATION, 17 Deliquescence, 83 Densities of gases, 163 Dephlogistigated air, 55 Destructive distillation, 179 Dextrin, 271 Diamond, 213 Diffusion of gases, 169 Distillation, 23 Dolomite, 211 Double decomposition, 156 Downward displacement, 27 Dumas' apparatus, 112 Dutch metal in chlorine, 131 EFFLORESCENCE, 83 Electrolysis of water, 76 Elements and compounds, 8 Equations, chemical, 154 Equivalents, chemical, 139 Ethylene, 212, 265 Evaporation, 15 FERMENTATION, 222 Filtration, 17 Fire damp, 264 Fitting up apparatus, 28 Fluorine, 133 Fluor-spar, 133 Fusion, 21 GALENA, 236 Gases, densities of, 163 , diffusion of, 169 , general properties, 97 , relation of volume to heat, 97 , pressure, 98 Glass blowing, 35 Glass tubes, bending, 30 Glucose, 270 Glycerin, 268 Graham's law, 170 Graphite, 214 Green vitriol, 46 Gunpowder, 194 Gypsum, 236 HALOGENS, 133 Hardness of water, 81 Hydrocarbons, 264 Hydrogen, 40 chloride, 125 . peroxide, 83 Hydroxides, 63 ICE, 79 Ignition point, 173 Index. 287 Indestructibility of matter, 65 Iodine, 134 Iron, tests for, 277 Iron pyrites, 236 LAMP-BLACK, 217 Laughing gas, 202 Law of Boyle, 98 Charles, 97 constant composition, 117 Gay-Lussac, 162 multiple proportions, 118 Lead, tests for, 275 Lime-light, 173 Limestone, 211 Lime water, 221 Liquefaction of gases, 102 MAGNESIUM carbonate, 227 combustion in air, 105 oxide, composition of, 105 Marble, 211, 228 Marsh gas, 210, 264 Matter, indestructibility of, 65 Mechanical mixtures, 9 Mercuric oxide, 54, 115 Metalloids, 9, 178 Metals and non-metals, 9 Metric system, 85 Molecular formulae, 153 weights of gases, 163 Molecules, 149 Monobasic acids, 193 Mortar, 228 NASCENT state, 130 Natural waters, 81 Neutralization, 68, 143 Nitrates, detection of, 195 Nitre, 10, 180 Nitric oxide, 197 Nitrites, 203 Nitrogen, 176 , oxides of, 197 Nitrous oxide, 201 OLEFIANT gas, 265 Oxidation, 63 Oxides, 62 Oxygen, 54 Oxyhydrogen flame, 173 Ozone, 206 , tests for, 208 PERCENTAGE composition, 114 Permanent hardness of water, 82 Phosphorus, combustion in oxygen, 59 Pipette, 93 Plaster of Paris, 224 Plumbago, 214 Potassium, tests for, 279 permanganate, 203 Precipitation, 19 Prussian blue, 193 Pyrites burners, 257 QUANTITATIVE chemical equations, 158 Quartz, 214 Quicksilver, 2 RAIN water, 81 Reagents, 273 Red lead, 218 Respiration, 65 Rock crystal, 214, 219 Rust, 64 SAL ammoniac, 4, 182 Saltpetre, 180 Salts, 69 Saturated solutions, 20 Silver displaced by zinc, 140 Simple manipulations, 15 Soap, 268 Soap-bubbles with hydrogen, 50 Sodium, action on water, 42 carbonate, 227 peroxide, 56 Solubility of gases in water, 81 of salts in water, 2 1 288 Index. Solution, 15 Solvent power of water, So Spirits of hartshorn, 180 Starch, 211, 270 compound with iodine, 134 States of matter, i Steam, 4 Sugar of lead, 216 Sublimation, 5 Sulphates, 259 Sulphides, 241, 247 Sulphites, 253 Sulphur, 236 , allotropic forms of, 239 dioxide, 250 Sulphuretted hydrogen, 243 TEMPORARY hardness, water, 82 Thermal unit, 175 Thermometers, 87 Torricellian vacuum, 90 UNIT volume, 164 Upward displacement, 27 VITRIOL chambers, 257 Vitriols, 259 WASHING soda, 224 Water, 72 , colour of, 78 , electrolysis of, 76 , freezing of, 80 , gaseous, 4 , gravimetric composition of, 112 , hardness of, 81 , maximum density of, 79 , of crystallization, 82 , sea, 8 1 , volume composition of, 77 Water-gas, 232 Waters, natural, 81 Weighing and measuring, 85 ZINC blende, 236 , granulated, 45 , tests for, 278 PRINTED BY WILLIAM CLOWES AND SONS, LIMITED, LONDON AND BECCLES. 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