LIBRARY OF THE UNIVERSITY OF CALIFORNIA. Deceived |^ ^ ^93 l %9 Accessions A/o.(4ear in the next experiment. Ex. 36. Preparation of Chlorine Gas. Fill a flask (50 cubic centimetres capacity), fitted with a perforated rubber stopper and outlet tube, to three fourths of its capacity with lumps of black oxide of manganese. Pour upon these lumps strong liquid hydrochloric acid so as to fill the interstices only. Allow the flask to stand for a short time, and then apply a gentle heat. A yel- lowish gas comes off in abundance, which is much heavier than the air, and can be collected by dis- placement if the outlet tube reache's quite to the bottom of the jars used for the purpose, and the mouth is covered with a disk of cardboard, as already described. Chlorine gas is very suffocat- ing, and the smallest puff, if inhaled, may pro- duce serious results. This experiment should therefore be performed with extreme caution, * Use a glass or porcelain dish to hold the water. Not the pneumatic trough, which would be corroded by the mercury falling into it. 72 LABORATORY PRACTICE. either in the open air or under a hood with a strong draught. The colour of the gas shows when a jar is full, and three jars thus filled are required for the following experiments : 1. In the first of these jars plunge some tinsel or other metal leaf hanging from the end of a long stick, and almost every metal will at once enter into direct union with the chlorine, often with ignition. 2. Place mouth to mouth a jar of hydrogen over a jar of chlorine, and, holding the open mouths together confined by a rubber band, invert the two, and in a few moments, when the gases have mixed, loosely cover (but on no account seal) both of the jars. If now one of the jars is exposed to bright daylight (not direct sunlight), a gradual union be- tween the chlorine and the hydrogen gases will take place, and after the yellow colour has disap- peared the product can easily be recognized as hy- drochloric acid gas. The second jar, kept in the dark, will undergo no change, and can be used for comparison. If now this jar is exposed to direct sunlight, the same combination will suddenly take place with explosive violence.* This experiment is a dangerous one, and should only be made with the greatest caution, the mouth of the jar being * There is more or less danger in all stages of this experiment, not only from the violence of the explosion, but also from the risk of breathing a puff of chlorine. It should never be intrusted to care- less hands, nor indeed to the hands of any student before all the necessary precautions have been pointed out and enforced. CARBON. 73 loosely covered with a pasteboard disk. The com- position of hydrochloric acid is thus fully estab- lished. 3. Into the third jar of chlorine, prepared as above, pour 200 cubic centimetres of water, and, after closing the jar, shake the water up with the gas, which will be almost completely absorbed and impart its colour to the solution. Soak then in the water a strip of calico printed with madder, and notice how rapidly the colour is discharged. In this connection the use of chlorine as a bleach- ing agent should be explained. (1) What proof has been given that hydrochloric acid gas consists solely of hydrogen and chlorine ? Do our ex- periments show in what proportions hydrogen gas and chlo- rine gas combine by volume ? (2) What is the composition of liquid hydrochloric acid? Have you noticed any difference between the union of hydrochloric-acid gas with water and ordinary solution ? Does the composition of liquid hydrochloric acid conform to the general scheme exhibited under Ex. 33 ? Point out agreements and differences. (3) Compare sulphuric acid and liquid hydrochloric acid, using the ordinary laboratory acids diluted with four or five times their volume of water. Try taste, litmus paper, and action on zinc clippings. Also evaporate a few drops of each on watch glasses and observe effects. 7. Carbon. Ex. 37. (a) Preparation of Charcoal. Take small billets of three or four diiferent kinds of wood, including the densest and lightest that can be procured. Cover with sand in an iron crucible, 74: LABORATORY PRACTICE. and heat to redness until the smoking stops. It is best to light the gas thus given off above the sand to prevent it from escaping into the room. (b) Heat to redness over a lamp in a porcelain crucible some lumps of sugar as long as any vapor is evolved. In this connection the general facts in regard to the composition of organic substances should be briefly stated, and the relations of carbon as the non-volatile skeleton of organized matter should be explained. Also, the relations of diamond and graphite to charcoal should be stated, and compared with those between the dif- erent states of sulphur. Ex. 38. Specific Characters of Charcoal. The student should be asked to study the distinguish- ing characters of the charcoal prepared in the last experiment, and to compare these with those of sulphur already studied. If the work is judicious- ly directed and criticised this exercise will be very instructive ; and, for the very reason that the two substances are so different, it will be a good prepa- ration for comparing hereafter two substances which are closely alike. As charcoal is a porous body whose external volume depends on that of the organized material from which it was made, the density must obviously be left out of consid- eration in this comparison. Ex. 39. (a) Preparation of Carbonic Dioxide (Carbonic-Acid Gas). This gas, as we have seen, CARBONIC DIOXIDE. 75 is formed by the burning of charcoal (Ex. 24, c) ; it is also easily made by the action of liquid hydro- chloric acid on marble (calcic carbonate). Half fill a glass flask (250 cubic centimetres capacity) with small lumps of marble. Pour upon the mar- ble common muriatic acid (the commercial name of crude liquid hydrochloric acid) mixed with three times its volume of water. Connect with a pneumatic trough and collect in the usual way. Half fill one of the jars, and, after closing, shake the gas up with the remaining water. Open from time to time to admit air until the absorption ceases. At the ordinary pressure of the air water will absorb its own volume of carbonic-dioxide gas, and soda water is the same solution under pressure. This aqueous solution may be called carbonic acid. Dip into it a strip of blue litmus paper and notice the difference between the effect of this weak volatile acid and that of a strong fixed acid like sulphuric acid when both are equally dilute. (&) WJiat was the Source of Carbonic Dioxide in Last Experiment f To answer this question, prepare some lime water from quicklime slaked as in Ex. 15. Fill a quart fruit jar somewhat over one half (four sevenths) with lime water and the rest with carbonic dioxide (pouring in the heavy gas just as you would a liquid). Close, and shake the gas and water well together, admitting air from 76 LABORATORY PRACTICE, time to time as the absorption goes on. Allow the precipitate to settle ; pour off the clear water, collect the precipitate on a filter, wash, and dry. Transfer next the powder to a test tube and pour on a few cubic centimetres of dilute hydrochloric acid. Notice the effervescence and recognize as carbonic dioxide the gas evolved. To make the demonstration complete the student must be told that marble is one of the many mineral forms of carbonate of lime and is chemically the same substance as the white powder thus .prepared. The proof of this will appear later. In this con- nection the important relations of carbonic di- oxide in nature should be discussed, its pres- ence in the breath should be shown, and its association with alcohol as a product of fermen- tation and its presence in beer, sparkling wine, and other effervescing drinks, should be ex- plained. (c) Repeat Ex. 24 (c), and after the jar has cooled open the mouth under water. There will be no expansion of the aeriform product, and no contraction except that due to the slow solution of the gas in water. Obviously, then, a given volume of oxygen gas yields the same volume of carbonic-dioxide gas, and the last must weigh more than the first by the weight of the charcoal which the oxygen gas absorbs in the process of burning, CAKBONIC OXIDE. 77 (1) Is the volume of the atmosphere altered by the smoke which our fires pour into it ? Ex. 40. (a) Production of Carbonic Oxide. Provide two rubber gas bags, holding about one litre each. Connect these with the ends of a length of hard Bohemian glass tubing, which should be filled, but not tightly packed, with fine- ly pulverized charcoal that has been thoroughly burned. Having filled one of the bags to abont one half of its capacity with carbonic dioxide, heat the tube over two or more Bunsen lamps (best a gas tube furnace) to a full red heat,* and pass the gas slowly backwards and forwards so long as any increase of volume is perceptible, and at last it will be found that the volume has doubled. Remove now the full bag and transfer a portion of the product to a small glass jar over the pneumatic trough. Lift the jar from the water with the mouth down and the gas will not at once escape, because the new product is even lighter than air. It may now be lighted at the mouth of the jar and the peculiar color of the flame noticed and the product of the combustion shown to be carbonic dioxide. In this experi- ment it is evident that the carbonic dioxide must * The kerosene stove does not give a sufficiently high tempera- ture. By burning alcohol in the stove, however, the requisite heat can be obtained ; but care should be taken to avoid explosions, and it will be safer to reserve this experiment for the lecture table. 6 78 LABORATORY PRACTICE. have united with the material of the charcoal, and the new product, called carbonic oxide, must dif- fer from carbonic dioxide only in containing more carbon or, what amounts to the same thing, pro- portionally less oxygen. The same inference may be drawn from the fact that in burning carbonic oxide changes back to carbonic dioxide. (b) Add ten grammes of oxalic acid to a small flask, corked and fitted with an evolution tube leading to a pneumatic trough. Pour over this crystalline solid five or six times its weight of strong sulphuric acid (oil of vitriol). Support (on a retort stand) the flask, protected by a square of asbestos paper, and apply gentle heat. The gas which comes over copiously is a mixture of car- bonic oxide and carbonic dioxide. The last will be slowly absorbed by the water, and very rapidly if some caustic soda is added to the pneumatic trough. Collect in fruit jar, and, after allowing to stand until absorption is ended, compare this product with that of last experiment. Ex. 41. Etliylene (Oleflant Gas). Pour into a fifty-cubic-centimetre flask five cubic centimetres of high-proof alcohol, and then add slowly twenty cubic centimetres of strong sulphuric acid. Con- nect with a pneumatic trough and heat carefully, protecting the glass by interposing asbestos paper between the flask and the lamp. Ethylene, the gas thus obtained, burns with a brilliant flame OLEFIANT GAS. 79 and is one of the constituents of illuminating gas. The sole products of its combustion are carbonic dioxide and water, and it must therefore contain both carbon and hydrogen. It is here selected as an example of a very large class of substances called hydrocarbons. The phenomena attending their combustion have already been discussed (Ex. 28). In illuminating gas, a very complex product, ethylene is mixed, among other things, with a very large amount of another hydrocarbon, called me than (or marsh gas), which contains only half as much carbon and has far less illuminating power. The petroleums and the products ob- tained from them, known as benzine, kerosene, astral oil, paraffine, etc., are chiefly mixtures of a great number of hydrocarbons (gases, liquids, and solids), resembling in their chemical relations marsh gas, and classed with it under the general designation of the paraffines. Olefiant gas pre- pared as above is so called because it unites di- rectly either with chlorine or bromine to form a liquid which has an oily aspect, and there are several other hydrocarbons formed in the distilla- tion of coal which resemble olefiant gas in this re- spect and are classed with it under the name of the olefines. Then there is a hydrocarbon con- taining only one half as much hydrogen as olefi- ant gas, which is formed abundantly when a Bunsen lamp burns at the base^ and is at once 80 LABORATORY PRACTICE. recognized by its unpleasant odour. This hydro- carbon is also one of a class known as the acety- lenes, and is itself usually called by the same name. Lastly, there is a very remarkable class of hydrocarbons obtained by the distillation of coal tar, of which benzol and toluol are the chief mem- bers and from which the aniline dyes are pro- duced. These three classes, although the most important groups, by no means include all the known hydrocarbon compounds, while the possi- bilities of multiplication are unlimited, and from the great family of hydrocarbons the almost end- less products of organic chemistry may be de- rived. The points here suggested the teacher will expand as he sees fit. (1) Charcoal graphite or diamond when burnt in oxy- gen gas all yield the same product (carbonic dioxide). Are they the same substance ? (2) Does a solution of carbonic acid in water conform to your general conception of an acid ? (3) Compare the composition of carbonic oxide and car- bonic dioxide with that of sulphurous oxide and sulphuric oxide. Do all these oxides form acids by uniting with water ? Regarding the intensity of the acid taste and of the acid reaction as some measure of the strength of the acids, what would be your estimate of the relative strength of the acids thus formed, and how does this degree of strength com- pare with the apparent attraction of the oxides for water ? (4) Is the proof here given that olefiant gas is com- posed of carbon and hydrogen satisfactory ? Is it equally clear that the gas consists only of carbon and hydrogen ? PREPARATION OF NITRIC ACID. 81 8. Nitrogen. Ex. 42. Preparation of Nitrogen Gas. The aeriform product left in the jar from Ex. 24 (b) is nitrogen gas. The student should test the gas by immersing in it a lighted match ; but the element- ary student can not be expected to learn through actual experiments the complex relations of this remarkable substance. It should, however, be made evident by the teacher that the inability to support combustion is in entire harmony with the general inert relations of nitrogen gas, and that this is its most striking characteristics. Never- theless, when the necessary conditions are fulfilled, nitrogen readily enters into combination, espe- cially with oxygen and hydrogen, forming a nu- merous and important class of products. In illus- tration of this last point the formation of nitre should be explained ; and, starting from this nat- ural product, the student may prepare a few well- marked nitrogen compounds as follows : Ex. 43. (a) Preparation of Nitric Acid. Mix in the body of a small glass retort 30 grammes of nitre with the same weight of sulphuric acid. Allow the mixture to stand for several hours, and then distil over 15 grammes of a yellowish liquid, which is nitric acid (a very important chemical agent, consisting of nitrogen combined with both oxygen and hydrogen). The yellow colour of the 82 LABORATORY PRACTICE. product is due to an admixture of another nitro- genized product, which is very volatile, and may be driven off by gently heating the acid in a flask. (b) Nitric Acid contains Oxygen. Take in a test tube four or five cubic centimetres of the acid just made. Drop into it in small portions at a time not over one decigramme of coarsely pulver- ized roll brimstone.* Cautiously heat to boiling and maintain gentle ebullition for some time. Largely dilute with water and decant into a clean test tube from the remaining sulphur. Test with solution of baric chloride as described in Ex. 33. It will thus appear that sulphuric acid is formed by action of nitric acid on sulphur. Remembering now that sulphuric acid consists of sulphur, oxy- gen, and water, draw your own inferences. Indeed, nitric acid contains so much oxygen, held feebly in combination, that it furnishes an efficient means of uniting oxygen to other bodies. It may sustain combustion like the atmosphere, and it is for these reasons said to be a powerful oxidizing agent. The teacher can effectively illus- trate this point by pouring from a long- handled glass or porcelain spoon a few cubic centimetres of the strongest acid on finely pulverized charcoal. The powder should be well dried by heating it to * The action of very strong nitric acid on combustible matter is often violent, and this experiment should be made with caution. COMPOSITION OF NITRIC ACID. 33 incipient redness in a porcelain dish, and while still hot the acid poured upon it (a few drops at a time). The charcoal will then flash almost like gunpowder. In this connection the teacher may add that in the explosion of gunpowder charcoal burns at the expense of the oxygen stored in the nitre. (c) Nitric Acid contains Nitrogen and Water. This can be shown by passing the vapor of the acid carried by a current of carbonic dioxide over copper clippings heated to redness in a combustion tube. The metal will take up all the oxygen in the acid except that belonging to the water present, while the water set free may be collected by pass- ing the current after leaving the combustion tube through a U tube packed in ice. If, lastly, the current is passed on to a small pneumatic trough filled with water holding caustic alkali in solu- tion the carbonic dioxide will be absorbed and nitrogen gas collected. A regulated current of carbonic dioxide is easily obtained with the gen- erator described in note to Ex. 25, filling the bell with broken marble and the beaker with common muriatic acid diluted with an equal volume of water. The few cubic centimetres of nitric acid required are best held in a bulb tube so support- ed that it may be warmed with a lamp, and con- nected at one end with the generator and at the other with the combustion tube by a rubber con- 84 LABORATORY PRACTICE. nector (corks would be instantly corroded). The combustion tube may be arranged as in Ex. 33, but should be somewhat larger, and is best filled with finely pulverized copper, such as is obtained by the reduction of copper oxide. The small pneumatic trough is easily extemporized out of a glass or porcelain dish and a large test tube. Other details of the apparatus may now be left to the ingenuity of the student ; but the experiment should not be intrusted except to skilful manipu- lators, and in most cases will best be shown on the lecture table. In all cases, however, he should make a sketch of the apparatus in his note book and point out the use of each part and justify the conclusion that Nitrogen, oxygen, and water yield nitric acid. (1) Is the proof which has been given of the composi- tion of nitric acid synthetical or analytical ? The synthesis of nitric acid can not be readily made because, in conformity to its great inertness, nitrogen does not combine directly with oxygen. Nevertheless, by indirect means, the oxide of nitrogen corresponding to nitric acid has been prepared. It is a white crystalline solid which eagerly unites with water. Compare sulphuric oxide. (2) You have now handled the most important acids used in a chemical laboratory. Make a list of them with the composition of each so far as you have discovered it. Make clear that you can recognize all of them whether con- centrated or diluted. In every case a specific test has not been given, but by inquiry of your teacher or elsewhere seek the necessary knowledge until you are perfectly certain of your ability to distinguish all these substances. NITRIC OXIDE. 85 Ex. 41. (a) Preparation of Nitric Oxide. Place fifty grammes of copper clippings in a glass flask fitted with a cork, through which passes a tube funnel as well as an evolution tube. Drench the clippings with water, and then pour on through the funnel in successive portions the product of the last experiment mixed with three times its volume of water. After the action has started add a fresh portion from time to time as the effervescence slackens. Connect with a pneu- matic trough and collect the gas in a quick-sealing jar. The colorless gas is a compound of nitrogen and oxygen, called nitric oxide. The deep -red fumes which appear in the generating flask, and whenever the nitric oxide mixes with the oxygen gas of the air, is another compound of nitrogen and oxygen containing more oxygen and called nitric peroxide ; and it is chiefly this adventitious product which imparts the yellow tint to the crude nitric acid. () Analysis of Nitric Oxide. In a pint jar of nitric oxide collected in the last experiment burn a small bit of phosphorus, not exceeding one fourth of a gramme in weight, with all the pre- cautions stated in Ex. 24. Notice that the same white product is formed as when phosphorus burns in pure oxygen gas or in air. Hence we may infer that nitric oxide contains oxygen. After the jar is cold open the mouth under water. 86 LABORATORY PRACTICE. Notice that the residual gas fills only one half of the volume of the jar, and, further, that it has the characteristic inertness of nitrogen. Hence the additional conclusion that nitric oxide consists of equal volumes of nitrogen gas and oxygen gas. (1) Compare the action of copper on nitric acid in the last two experiments. Notice that while in the first it re- duces the acid to water and nitrogen gas, in the second it does not remove so much oxygen and leaves nitric oxide. Observe also that nitric oxide shows no tendency to unite with water. Ex. 45. Preparation of Ammonia. Mix in a gasometer* one volume of nitric oxide with two and a half volumes of hydrogen, and pass a slow * A very useful and inexpensive gasometer may be made of a large glass bottle of the capacity of two quarts or one gallon, as re- quired. Fit tightly in the neck a rubber cork with three perfora- tions. Through these perforations pass glass tubes, all bent at right angles a short distance above the cork. Two of the tubes should reach the bottom of the bottle, the third only pass through the stop- per. Connect one of the longer tubes with a sink by a length of rubber hose. In the same way connect the second of the longer tubes with a water tap ; and, lastly, slip on to the shorter tube a third length of rubber hose to serve as an outlet for the gas. Guard all three of the tubes with pressure taps. Begin by filling the ga- someter with water, allowing the liquid to flow in from the tap and the air to escape from the outlet. When full, close the outlet, open the overflow into the sink, remove the rubber hose from the water tap, and connect it with the gas generator. The gas will then flow in, the displaced water running off by the overflow. When a suffi- cient amount of gas has been collected the overflow must be closed and the rubber hose removed from the generator and replaced on the water tap. Then, on opening the tap, water will flow in again and drive out the gas by the outlet tube through any apparatus with which it may be connected. AMMONIA. 87 stream of this mixture through a short tube filled with platinized asbestos and heated to a low red heat. There will be an abundant formation of aqueous vapours, indicating that a combination has taken place between the hydrogen gas and the oxygen of the nitric oxide ; and at the same time there will be developed a strong pungent odour, familiar to every one as the odour of am- monia, which must evidently be formed by the union of nitrogen with hydrogen. This pungent product is a gas, and common aqua ammonia is a solution of the gas in water. To obtain a more familiar acquaintance with this important chem- ical agent, half fill a small glass flask with con- centrated aqua ammonia and close the neck with a rubber stopper, through which passes an evolu- tion tube leading to the top of a quick- sealing glass jar, arranged as in Ex. 26. Heat the liquid in the flask to boiling, when a large volume of colourless gas will pass over and nray be collected by displacement. When the jar is full, the over- flow will be at once recognized by the strong odour. The process may then be stopped, the jar closed, and the gas preserved for another experi- ment. Ex. 46. (a) Ammonia Salts. Mix five cubic centimetres of strong aqua ammonia with twice its volume of water. In the same way mix five cubic centimetres of strong nitric acid with twice its 88 LABORATORY PRACTICE. volume of water. Study the effects of these solu- tions on test papers, both litmus and turmeric. Next add slowly the ammonia to the nitric acid until the opposite effects exactly neutralize each other. Lastly, evaporate the mixture at a low heat until a drop taken out on a rod solidifies, and then when the dish is set on one side a white salt, called ammonic nitrate, will crystallize out. (b) Make the same experiment with sulphuric acid, with phosphoric acid, with hydrochloric acid, and with carbonic acid. In the last case add the fifteen cubic centimetres of diluted aqua ammonia to a jar of carbonic-acid gas and shake together. Evaporate each solution to dryness and collect the crystalline salt. Obviously the solution of ammonia gas (aqua ammonia) sustains relations which are the very opposite to those of the acids, and belongs to an equally important class of chemical products variously called alkalis (when the solutions are caustic) or bases (when they are not). The acids exhibit no tendency to unite with each other, but they eagerly unite with the anti-acids (as in the above experiments) to form a class of bodies, for the most part crystalline, called salts. The anti- acids are formed, as a rule, by the union of met- als with oxygen and water, while the acids (as we have seen) result as generally from a similar union of bodies, like phosphorus, sulphur, and MAGNESIUM AND ZINC. 9 carbon, which do not exhibit metallic characters. Having studied the relations of a few of these so- called metalloids we pass next to study the rela- tions of several metals. 9. Magnesium and Zinc. Ex. 47. Specific CJiaracters. The student should be given a gramme of magnesium ribbon, and with this he should study the characters of the metal (colour, lustre, tenacity, specific gravity) and compare these with the corresponding prop- erties of metallic zinc, also given to him rolled out into ribbon of about the same size. Reserving a short length of the magnesium ribbon for burn- ing, the student should dissolve the rest in a few cubic centimetres of dilute sulphuric acid. Col- lect and examine the gas evolved, and compare the reaction with that in Ex. 25. Lastly, let him evaporate the solution of magnesium thus ob- tained on a watch glass, and compare the crystal- line residue with that obtained from the solution of zinc formed in the above-cited experiment. Ex. 48. Burning of Magnesium. Let the student burn the reserved piece of magnesium ribbon, holding it by pincers and lighting it like a match, and let him compare the combustibility and colour of the flame with that produced with a similar ribbon of zinc. In order that he may fur- 90 LABORATORY PRACTICE. ther study the nature of the product formed, fur- nish the student with half a gramme of magne- sium* powder. Let him place this on a small square of asbestos paper (previously ignited) and weigh the amount accurately on the pan of a balance. Let him now ignite the powder with a match, and when burned out and cooled reweigh it. What means the increase of weight, and what must be the composition of the white powder left? Transfer the powder to a small evaporat- ing dish. Thoroughly drench with water. Place a small bit of the wet powder on red litmus paper. Compare the effect with that of an acid. Lastly, dissolve the residue in the smallest pos- sible amount of dilute sulphuric acid, adding the acid drop by drop. Evaporate the solu- tion till a crust appears and leave to crystallize. Can you recognize the saline product by the taste ? Try the same experiments with zinc ; but its powder does not burn so readily, and it is more difficult to recognize the increase of weight. Nev- ertheless, it is easily burned by sifting the powder on to a sheet of paper through the flame of a Bunsen burner held obliquely, or by spreading * A few years ago magnesium would have been too expensive for general experimenting, but, as a result of its application in the arts, it can now be purchased for a moderate price. At the price quoted by German dealers in chemicals the two grammes required for each student would cost less than three cents. MAGNESIUM AND ZINC. 91 the powder over a square of asbestos paper and playing on it with the same flame. (1) The comparison of two metals closely resembling each other, like magnesium and zinc, affords excellent prac- tice, and may be used to test the student's skill in observa- tion and deduction. If further practice is thought necessary a comparison may be made between two metals resembling each other still more closely; as, for example, iron and nickel. In all such cases the student should be required to work out the results unaided, and make a full and clear statement in his note book of what he observes and what he infers. His work should then be carefully criticised, and, if necessary, the experiments repeated, after fresh directions or suggestions from the teacher. The student should be led to appreciate the fact that although the distinctions between substances are usually broad and clear, they are also at times narrow and indefinite, and that the identification or differ- entiation of a newly found substance often turns on minute observations and delicate discriminations. (2) How can you prove that magnesic oxide combines with water ? Does zinc oxide combine in like manner ? Compare Magnesium, oxygen, and water yield magnesic hydrate basic. Sulphur, oxygen, and water yield sulphuric acid acid. Further, it appears that Magnesic hydrate and sulphuric acid yield Epsom salts salt. Would magnesic oxide yield the same product as mag- nesic hydrate ? (3) Compare the product of the action of magnesium on dilute sulphuric acid with that obtained by dissolving mag- nesic oxide in the same solvent. Why is hydrogen gas not evolved in the second process ? 92 LABORATORY PRACTICE. 10. Sodium. Ex. 49. Specific Characters. The specific char- acters of this interesting alkaline metal should be shown as far as possible to the student ; but it will seldom be advisable to intrust the material to inexperienced hands, and equally good practice in studying specific characters can be had with cheaper and less dangerous substances. The ac- tion of sodium on water illustrates principles so fundamental in the theory of chemistry that the experiment should on no account be omitted. The action of the pure metal is always violent, and frequently dangerous ; still, it is a very inter- esting experiment, which the teacher may make before the class, with proper precautions. The best way is to throw a bit of sodium, not larger than a pea, on some sheets of porous paper thor- oughly soaked and running with water. The melted globule is thus prevented from swimming round, and the heat developed by the chemical change accumulates to such an extent as to in- flame the escaping hydrogen, which burns with a flame that is coloured yellow by the presence of sodium vapour. For use of students, an amalgam of sodium one part of sodium to about ten parts of mercury should be prepared ; and this may be used with entire safety. The action of sodium on mercury is violent, but the amalgam can be easily ALKALIES. 93 made by heating the mercury to about 200 in a Hessian crucible of eight or ten times the capacity required to hold the metal, and then adding the sodium in one large bar. Assuming this amal- gam to have been previously prepared or pur- chased, the student should make the following experiment, having, of course, been previously told the object of using the mercury, and that it plays no part in the chemical change : Take a small flask fitted for the evolution of gas, and place in it about twenty- five grammes of water ; add now a few lumps of the amalgam, and collect the gas over the pneumatic trough. As the evolution lessens, hasten it by heating the flask with a lamp. Burn the gas that is collected, and recognize that it is hydrogen. Pour off now the solution left in the flask from the mercury, and, in the first place, test it with litmus and turmeric paper which have previously been dipped in a very weak acid. It will thus be seen that the product is a sub- stance which, like the solution of ammonia, re- verses the effect of an acid on vegetable dyes ; in other words, that it is basic. Such soluble and caustic bases are called alkalies. Rub a few drops of the liquid between the fingers, and notice the effect, which is termed caustic. Evaporate now the liquid, and compare the residue with caustic soda. Redissolve this residue in a very small amount of water, and divide the solution between 7 94 LABORATORY PRACTICE. three watch glasses ; neutralize the solution in the first glass with a few drops of hydrochloric acid, that in the second glass with a few drops of nitric acid, and add the contents of the third watch glass to about twenty -five cubic centimetres of carbonic acid (soda water). Allow the solutions to evapo- rate, and examine the crystals formed with a lens; also attempt to recognize the products by tasting the residue in each case ; they will be discovered to be common salt, sodic nitrate, and sodic car- bonate, respectively. As caustic soda has thus been made solely with sodium and water, a prob- able inference in regard to its composition may at once be drawn, and the three familiar products last obtained will be recognized as salts of sodi- um. In this connection the student should be told about the sources of these substances, and their uses in daily life and in the arts. Ex. 50. (a) Using a small amount of magnesi- um or zinc powder and boiling with water in a test tube, compare the action of these metals on water with that of sodium. It will appear that Sodium acts violently on water and yields so- dic hydrate and hydrogen gas. Magnesium acts slowly on water and yields magnesic hydrate and hydrogen gas. Zinc acts very feebly on water and yields zinc oxide and hydrogen gas. (b) The teacher should burn sodium in dry METALLIC OXIDES. 95 oxygen, first melting the metal in an iron spoon. Dissolve the oxide formed in water, evaporate to dryness, and compare it with product of the direct action of sodium on water. Care must be taken to separate the white powder from unburnt metal before adding water, and the experiment should not be intrusted to unskilful students. It will now further appear that Sodium unites with oxygen to form sodic oxide (white powder). Magnesium unites with oxygen to form mag- nesic oxide (white powder). Zinc unites with oxygen to form zinc oxide (white powder). Also that Sodium oxide unites with water eagerly to form sodic hydrate. Magnesium oxide unites with water feebly to form magnesic hydrate. Zinc oxide will not unite with water. The striking differences depend on the fact that sodic hydrate is very soluble in water. In- terpret all the phenomena. (c) The teacher should burn sodium in dry chlorine. It burns readily when heated above melting point in an iron spoon, but the experi- ment should be made under a hood with powerful draught. A considerable part of the product vola- tilizes. This is readily dissolved in water and crys- 96 LABORATORY PRACTICE. tallized, when the cubic form of the crystal and taste show it to be common salt. It thus appears that sodium and chlorine yield common salt. So- dic hydrate (or sodic oxide) and hydrochloric acid also yield common salt. Consider in what respect the formation of common salt differs from that of other salts for example, magnesium sulphate. Ex. 51. Volatile and Fixed Alkali. Compare the formation of the ammonia with that of the sodium salts. Ammonia, gas, water, and sulphuric acid yield ammonic sulphate. Ammonia, gas, water, and nitric acid yield ammonic nitrate. Ammonia, gas, water, and hydrochloric acid yield ammonic chloride. Sodic hydrate and sulphuric acid yield sodic sulphate. Sodic hydrate and nitric acid yield sodic ni- trate. Sodic hydrate and hydrochloric acid yield sodic chloride. Compare the two salts of each acid and deter- mine whether volatile or not. Heat solutions of each of the ammonia salts with a solution of caus- tic soda and test the gas given off with red litmus paper and with great caution by the smell. REDUCTION AND OXIDATION. 97 11. Copper. Ex. 52. Distinctive Characters. The metal is best used in a very thin, flexible sheet, easily cut with a pair of scissors. Let the student first study and describe the properties of the metal, deter- mining its specific gravity in the usual way, and comparing its colour, hardness, toughness, mallea- bility, etc., with the similar qualities of zinc. Let him harden on an anvil and anneal with heat. Let him next heat in separate test tubes a few bits of the metal with nitric, hydrochloric, and sulphuric acids, using both strong and weak acid. Com- pare with zinc. Observe behaviour, to be inter- preted beyond. Ex. 53. (a) Reduce Oxide wiili Hydrogen Gas and reoxidize in the Air. Introduce into a corn- bastion tube about twenty grammes of black ox- ide of copper. Take the tare on the balance. Mount on tube furnace, connecting one end with a hydrogen generator and the other end with a small U tube kept cool with ice. Pass now a slow current of hydrogen gas over the powder, and after the gas has expelled the air heat the com- bustion tube to low redness. Observe that the powder takes on the colour and lustre of metallic copper and that water collects in the tube. After the reduction is complete, dismount the apparatus and reweigh the tube. 98 LABORATORY PRACTICE. (b) Kemount on the furnace the combustion tube with its contents as left in the last experi- ment. Leave one end open to the air and con- nect the other with the gasometer before de- scribed (note to Ex. 45), which, having been pre- viously filled with water and the overflow opened into a sink at a lower level, will act as an aspira- tor. Regulate by pressure tap so that air will be drawn through the tube not faster than about two bubbles a second, and then heat the reduced cop- per to redness. After collecting one or two litres of gas close and dismount the gasometer, while at the same time connecting the combustion tube with an aspirator pump to hasten the process. Continue heating the tube until the powder has acquired a uniform black color. Then allow to cool, dismount, and reweigh. Meanwhile transfer the gas collected to gas bottles over a pneumatic trough and seek to recognize the substance. In- terpret all the phenomena observed. Compare Ex. 24 (b) and Ex. 42. Keep the oxide of copper in the tube for another experiment. (1) Water consists of hydrogen and oxygen. The pro- duction of water from hydrogen gas implies what ? The formation of water is attended with the change of the black powder to metallic copper. Of what must that powder con- sist? (2) Whence comes the nitrogen collected in the gasome- ter ? How may free oxygen, or substances holding oxygen loosely united, be expected to act on free copper ? How may SALTS OF COPPER. 99 oxide of copper be expected to act on substances containing hydrogen or carbon ? Ex. 54. (a) Salts of Copper. Take three test tubes holding a few cubic centimetres of dilute sulphuric, hydrochloric, and nitric acid respect- ively. Dissolve in each black oxide of copper, adding in very small quantities so long as solution is obtained on boiling. Evaporate on watch glasses and crystallize the products. Oxide of copper and sulphuric acid give cop- per sulphate (salt). Oxide of copper and nitric acid give copper nitrate (salt). Oxide of copper and hydrochloric acid give copper chloride (salt). (b) Take a few cubic centimetres of the solution of copper in nitric acid Ex. 52 or 44 (a). Secure the ready solution of copper in sulphuric and hy- drochloric acids by adding a few drops of nitric acid in each case ; evaporate to .dry ness (not -over 100), dissolve the residue in the smallest possible amount of water, and allow to crystallize in a warm place. Compare products with those ob- tained in (a). Lastly, heat copper nitrate in a porcelain crucible to low redness, until fumes cease, then to bright redness. When cold pulver- ize the residue in a mortar and examine the black powder left. Seek to interpret now the phenomena observed when attempting to dissolve the metals. 100 LABORATORY PRACTICE. (1) What is the difference between the action of copper and zinc on dilute sulphuric acid ? This difference can be explained by the thermal relations of the metals, but must here be accepted as a fact. Why should you expect that the addition of nitric acid would secure the ready solu- tion of copper in both sulphuric and hydrochloric acids ? What double part does nitric acid play in dissolving copper ? Compare Ex. 44 (a) and also the additional fact that when in this preparation of nitric oxide the temperature is allowed to rise too high and the action to become violent the prod- uct is chiefly, or even altogether, nitrogen gas. 12. Iron. Ex. 55. Distinguish Ing Properties. The met- al is best used in the shape of wrought-iron-wire nails or tacks of various sizes, also in fine powder u iron by hydrogen." Let the student first study and describe the properties of the metal, as in the case of copper (Ex. 52). He should then find its specific gravity, using a method applicable in many cases. Take a small vial or flask (ten to twenty cubic centimetres capacity) and make a mark on the narrow neck. Fill the bottle to the mark with water and tare it on the balance. Select some iron nails as large as will conveniently be held by the bottle. Place about twenty grammes of these nails at the side of the bottle, and take the exact weight. Remove the bottle from the pan and drop in the nails one by one, taking care to avoid entangling bubbles of air. Wipe off the water which runs over, and with a small roll of porous OXIDATION OF IKON. 101 paper reduce the level in the neck to the mark. Replace on the pan and reweigh, when the loss of weight is obviously the weight of water displaced by the nails. In this connection the qualities and relations of the different kinds of iron wrought iron, cast iron, and steel should be explained by the teacher, and the prominent features of the metallurgy of iron might appropri- ately be discussed. Ex. 56. (a) Burning of Iron. If the iron pow- der is sufficiently fine it will burn on the pan of a balance like magnesium (Ex. 48), and the increase of weight may be found. If too coarse to burn in this way, the iron powder will burn brilliantly by sprinkling it through the flame of a Bunsen burn- er held obliquely. The residue may be collected on a large sheet of paper held obliquely. A more striking experiment of burning a watch spring in oxygen gas should be made by the teacher. After removing the temper of the steel .by heating in a lamp flame, the spring can be coiled into a spiral. Tip one end with sulphur like a match, and hang the other from a wire, which should be slid through a cork, closing the tubulature of the jar as fast as the oxygen is exhausted. Protect the exposed face of the cork with metallic foil, and then light the iron match and plunge it into the gas. The experiment is best made in a tubulated bell jar standing over water into which the oxide 102 LABORATORY PRACTICE. of iron melted by the flame falls in drops. These melted globules would crack the bottom of a glass jar unless protected by sand. (b) Weigh out in a porcelain crucible five grammes of the iron powder. Ignite with fre- quent stirring (use iron wire) until completely oxidized. Reweigh when cold, and interpret the result. (c) Weigh out five grammes of iron powder in a shallow porcelain dish. Keep the powder moist with water until it is wholly converted into rust. ,Then allow to dry in the air and weigh again. Save a part of the residue for further use. (1) Why is the weight of the rust in (c) greater than that of the red powder formed in (6) ? Devise an experiment for testing your inference. Ex. 57. (a) Iron and Sulphuric Acid. Iron dissolves in dilute sulphuric acid like zinc or mag- nesium, with rapid evolution of hydrogen gas. Repeat Ex. 25, using 5 grammes wrought-iron tacks (instead of zinc clippings), also 10 grammes strong sulphuric acid diluted with 30 grammes of water. Use a lOO-cubic-centimetre flask. Col- lect the gas as before, and compare with the pre- vious product. Observe precautions previously given (note to Ex. 27) in regard to the preparation of hydrogen gas. (b) After the evolution of gas has ceased, re- move the outlet tube from the flask and provide a IRON SALTS. 103 tightly fitting cork. Boil down the residual solu- tion to about one half, cork the flask while still full of steam, set aside and allow to cool. Exam- ine the crystals which form. They are ferrous sulphate. Interpret the. phenomena observed. What is the source of the hydrogen ? Compare Exs. 47 and 48. As evidence bearing on this point, it should here be stated that iron forms with oxy- gen an oxide containing less oxygen than that obtained in Ex. 56 (&), but so difficult to prepare and keep as to be unsuited for class experiments. This oxide dissolves in dilute sulphuric acid with- out evolution of hydrogen, yielding the same ferrous sulphate. In this connection, and as a step towards the answer to the above question, let the student review the relations of magnesium, magnesium oxide, and magnesium hydrate to- wards sulphuric acid. Ex. 58. Two Classes of Iron Salts. It is by no means true of all metallic oxides that they will dissolve directly in acids to form salts without the intervention of any other agent or the forma- tion of any other product. The oxides that do sustain this relation to acids, like most of those we have studied, have been called for distinction sake salifiable oxides. Most of the metals, like sodium, magnesium, and zinc, form but one sali- fiable oxide ; but there are a number of metals which yield two such, and it is a remarkable fact 104: LABORATORY PRACTICE. that the two classes of salts thus formed differ widely from each other in their properties and re lations. Iron is an example in point, and the two classes of salts thus formed are distinguished as ferrous and ferric salts. The green transparent crystals (green vitriol or ferrous sulphate) formed by dissolving iron in dilute sulphuric acid is a ferrous salt, and the same product, as has been said, also results when the above-mentioned fer- rous oxide is dissolved in the same acid. The oxide formed in Ex. 56 (&), called ferric oxide, is also salifiable, but when once ignited, as in that experiment, dissolves in acids with difficulty. The hydrate formed by slow oxidation in the air in contact with water dissolves readily. To substan- tiate that the ferric salts thus formed are essen- tially different from the ferrous salts, dissolve a portion of the residue from Ex. 56 (c) in dilute sulphuric acid, evaporate nearly to dryness on a watch glass, and try to crystallize the residue. Compare this ferric sulphate with the ferrous sul- phate before made. Ex. 59. (a) Iron and Sulphur. Thoroughly mix 5 grammes of iron powder with 2 '8 grammes of flowers of sulphur. Save a small portion of the mixture for comparison ; heat the rest in a small flask (150 cubic centimetres) until the mass glows. After cooling remove with a rod a few grains of the product and compare with the mixture. (Use HYDROGEN SULPHIDE. 105 microscope and magnet.) The black product is called iron sulphide. (&) Hydrogen Sulphide. Leaving residue in the flask, to which has been fitted cork and out- let tube, pour in 10 grammes of strong sulphuric acid mixed with 50 grammes of water. After air has been driven from the flask, collect the first portions of the escaping gas in large test tubes over hot water. Ignite the gas at the open mouth of these test tubes, and observe the colour and odour produced by the flame. Let the rest of the escaping gas bubble up through cold water in a glass - stoppered bottle, and when action is ex- hausted withdraw outlet tube, stopper the bottle, and set aside for future use. Boil now contents of flask as in Ex. 57 (&), and set aside to crystallize. These crystals will at once be recognized as fer- rous sulphate, and the nauseous - smelling gas, which is, to a limited extent, soluble in water, is called hydrogen sulphide. It is ene of the most important of chemical reagents. As this gas is not only nauseous, but also to some extent poi- sonous, this experiment must be made under a hood or in the open air.' Interpret all the phases. (1) Iron sulphide obviously consists of iron and sulphur. Iron dissolved in dilute sulphuric acid yields ferrous sulphate and hydrogen gas. Ferrous oxide dissolved in the same acid yields also ferrous sulphate, but no free gas, because the hy- drogen, otherwise formed, unites with the oxygen of the 106 LABORATORY PRACTICE. oxide to form water. Ferrous sulphide dissolved in the same acid yields, again, ferrous sulphate and a nauseous- smelling gas. What must be the composition of this gas ? Do the phenomena observed when the gas burns confirm your inference ? CHAPTER II. GENERAL PRINCIPLES. 13. Province of Chemistry. AT tliis stage of Ms study the student, having become familiar with the distinctions implied by the word " substance," and having acquired some knowledge of chemical phenomena, is prepared to understand what is the province of the science of chemistry. Chemistry comprises and classi- fies our knowledge of those phenomena which imply a change of substance. The science of physics, on the other hand, deals with phenomena which do not necessarily imply a change of sub- stance ; and hence the distinction .between chemi- cal and physical changes. This distinction should be illustrated by the teacher from the experiments already made by the student. In every chemical change one or more sub- stances, called the factors, change into one or more other substances, called the products; and it is a primary object in the study of chemistry to learn what are the factors and what are the prod- ucts of every process that comes under notice. 108 LABORATORY PRACTICE. Substances may be mixed with one another, like the ingredients of gunpowder, or one substance may be dissolved in another, like salt in water, without undergoing chemical change ; but in all such cases the qualities of the original substances may be recognized in the mixture or solution. Hence the very broad distinction between a mixt- ure, or a solution, and a chemical combination, which the following experiments will illustrate. Ex. 60. Mixture and Chemical Compound. Mix together in a mortar as intimately as possible 3'26 grammes of zinc dust with 1*60 gramme of flowers of sulphur. First examine a small amount of this powder under a microscope of sufficient power to show the yellow grains of sulphur lying side by side with the metallic grains of zinc. Make with the rest of the mixture a conical pile on a square of asbestos paper, and apply the flame of a match. A chemical change ensues, marked by a brilliant deflagration ; and as the result there is left on the paper a white powder, which is a com- pound of zinc with sulphur, and is called sulphide of zinc. Here there were two factors of the chem- ical change, zinc and sulphur, and one product, sul- phide of zinc. Examine with the microscope this product, and no traces can be seen either of zinc or of sulphur. The deflagration was a manifestation of the heat evolved by the chemical change ; and in every chemical change there is either a setting PHYSICAL AND CHEMICAL SOLUTION. 109 free or an absorption of energy, usually as heat. But this last feature of a chemical process will be considered later by itself. Compare Ex. 59 (a). Ex. 61. Physical and Chemical Solution. Take two portions of one gramme each of sodic carbonate. Dissolve one portion in three cubic centimetres of water, evaporate to dryness slowly, and compare the residue with the original salt in appearance, crystalline form, and taste. Dissolve the second portion in dilute hydrochloric acid, evaporate, and compare. What are the factors and what are the products of the chemical change in the second case? Notice that water is the medium of the chemical process, and dissolves one of the products ; so that we have here both a chemical change and a simple solution. The same is true when zinc dissolves in dilute sulphuric acid (Ex. 47), or when copper dissolves in dilute nitric acid Ex. 44 (a) and the double use of the term " solution" must be made clear. What are the factors and what are the products in the two cases of chemical solution last cited? In the same way the teacher should review the experiments which the student has made, and point out what are the factors and what are the products in each case. 110 LABORATORY PRACTICE. 14. Fundamental Laws. In every well-marked chemical change three fundamental laws are observed, and these are called the law of conservation of mass, the law of definite proportions by weight, and the law of definite proportions by volume. Ex. 62. Law of Conservation of Mass. Re- peat Ex. 24, but after adjusting the apparatus with the bit of phosphorus in the spoon and fast- ening the cover, balance the jar on the pan of the balance with a second jar of the same volume and such additional tare as may be needed. Ignite now the phosphorus with a burning-glass, and after the chemical action, when the jar is cold, replace it on the balance. If the jar was tight there will have been no change of weight. Hence it must be that Tlie sum of the weights of the products of a chemical change is exactly equal to the sum of the weights of the factors. We may conceive of any chemical process as taking place in an hermetically sealed space in- deed, the earth is essentially such a space and hence this law must be universally true. The re- sult of this experiment might be anticipated, and it may therefore be thought unnecessary ; but its very form will make evident to the student that the law of conservation of mass is in harmony FUNDAMENTAL LAWS. with general principles which, he already recog- nizes. Ex. 63. Law of Definite Proportions ~by Weight. Take five grammes of sal-soda (crystal- lized sodic carbonate), selecting material that has not effloresced ; dissolve in dilute hydrochloric acid, as directed in Ex. 61, taking care to avoid loss during the effervescence ; evaporate to dry- ness, and weigh the residual salt; calculate the ratio of the sal-soda used to the salt produced ; repeat the same determination with ten grammes of sal-soda, and within the limits of experimental error the ratio will be the same as obtained be- fore, and so would it be whatever the amount of sal- soda employed. In this chemical change the factors are sal-soda and hydrochloric acid, while the products are common salt, carbonic dioxide, and water. The last two, being volatile, escape during the effervescence and subsequent evapora- tion. By this experiment we have proved that the proportion between the weight of the sal-soda and the weight of the common salt is definite, and it could readily be shown experimentally that the proportion between any two of the five sub- stances involved in this chemical change was equally definite. So of any other well-marked chemical change, and hence the general law that In any well-marked chemical change the rela- 112 LABORATORY PRACTICE. tive weights of the several factors and products are definite and invariable. Here, again, the result might have been antici- pated, for it only amounts to finding that if we use twice as much sal-soda we shall obtain twice as much common salt, which might seem self- evident; and this consideration will show that the law of definite proportions by weight is in entire harmony with principles universally rec- ognized. Ex. 64. Law of Definite Proportions ~by Vol- ume. This law, sometimes called the law of Gay- Lussac, may be thus stated : In any well-marked chemical change the rela- tive volumes of the aeriform factors or products, if measured under the same conditions, bear to each other a simple numerical ratio. It has already been illustrated by several ex- periments, which it is unnecessary to repeat. Thus it was shown by Ex. 32 that when oxygen combines with sulphur to form sulphurous oxide the volume of this sole product is the same as the volume of the oxygen gas used. A similar rela- tion appeared when oxygen united with carbon to form carbonic dioxide in Ex. 24 (c). Again, when, in Ex. 40 (a), carbonic dioxide united with more carbon to form carbonic oxide the volume of the gas was doubled. A still more striking il- lustration of the law is to be found in the fact ELEMENTARY SUBSTANCES. H3 that two volumes of hydrogen gas combine with one volume of oxygen gas to form two volumes of vapour of water, all measured, of course, under the same pressure and at a temperature above the boiling point of water. The experiment is easily made with a form of eudiometer invented for the purpose by Hof mann and sold by all the dealers in chemical apparatus, and it should be shown to the class if possible. 15. Compounds and Elements. The student can not have performed the ex- periments heretofore described without himself drawing the inference in certain cases that the products have been formed by the union of two or more factors, and in other cases that the prod- ucts have resulted from the breaking up of a factor into simpler parts. Hence come the fun- damental conceptions of composition and decom- position, of synthesis and of analysis, as we have previously called them. Our judgment in any case depends not only on the circumstances of the experiment, but also on a comparison of the weights of the products with those of the factors from which they were formed. Thus, in Ex. 24 (a\ it is perfectly evident from the conditions of the experiment that the white product results from the union of phosphorus and oxygen. If LABORATORY PRACTICE. now in addition we could weigh the white prod- uct and find that its weight was exactly equal to that of the phosphorus and oxygen used, the proof of its composition would be complete. So also when, in Ex. 16, we pass a current of elec- tricity through water and see oxygen and hydro- gen gases escaping from the platinum poles of the apparatus, and notice that everything else re- mains unchanged, we conclude that the two gases must come from the water and are the products of its decomposition ; but we do not have absolute proof until, as in Ex. 53 (a), we pass hydrogen over oxide of copper and find that the weight of the water formed is exactly equal to that of the hydrogen and oxygen which have disappeared. In like manner, our knowledge of the composition of other substances is the result of our knowledge of chemical processes, which has been accumu- lated during long years of study. As the total result of this study, we may say that while the larger number of substances which we handle may be decomposed or analyzed, there are about sev- enty known substances which can not be broken up into simpler parts, and these we call ele- mentary substances. An elementary substance differs from other substances only in this, that it enters into all chemical changes as a whole, and we know of no chemical process in which it be- comes divided. It does, however, enter into ELEMENTARY SUBSTANCES. H5 union with other substances ; and, speaking in general, we may by the combination of the ele- mentary substances reproduce all the forms of matter with which we are acquainted. The sys- tematized knowledge of the methods, whether analytical or synthetical, by which the composi- tion of bodies has been determined is a very im- portant branch of chemical science, known under the name of chemical analysis ; and the subject is subdivided into qualitative and quantitative analysis, according as the object in view is to de- termine solely the nature or the proportion of the ingredients. In either case the analysis may be either ultimate or approximate. It is ultimate when we seek for the elementary substances of which the compound consists. It is approximate when we look for the simpler products (for the most part acids and metallic oxides) into which the complex material may be primarily divided. The following experiments wi-ll give a general, but necessarily a very imperfect idea of the man- ner in which the results are reached : (1) A list of the elementary substances will be found in the table at the end of this book, and this should be care- fully examined by the student in reference to the substances he has met with in the course of his experiments. Which of these are elements 2 The student should make a list of the elementary substances with which he has become famil- iar. Can an elementary substance be told by its external characters ? Is there not a class of b^dJ^wJaich are uni- formly elementary ? TJHIVEKSITT 116 LABORATORY PRACTICE. 16. Qualitative Analysis. Ex. 65. Analysis of a Silver Com. Dissolve a ten-cent coin in 5 cubic centimetres of pure, strong nitric acid, diluted with its own volume of water. Dilute to 50 cubic centimetres. Add hydrochloric acid to hot solution so long as a precipitate is pro- duced. Filter, wash thoroughly (three times) with water, dry precipitate. Transfer to glass combustion tube connected with hydrogen genera- tor as in Ex. 53 (a). Interpose chloride-of-calci- um tube between generator and combustion tube. Allow hydrogen to flow through the apparatus long enough to expel the air ; then heat the com- bustion tube, and continue until reduction is com- plete. Test gas evolved from outlet, which should be bent downwards. Preserve silver. Add a few drops of sulphuric acid to blue filtrate and evapo- rate (under hood) to get rid of the volatile acids ; dilute to 10 cubic centimetres and insert a strip of zinc. (1) What is the gas evolved during the reduction? What is its composition ? What must be the composition of the precipitate ? Does the formation of this precipitate con- form to any general principle Ex. 14 (3). On what circum- stances does the separation of copper form silver in this ex- periment ? How can you be sure that the copper obtained came from the coin and not from any of the accessory ma- terials employed ? Ex. 66. (a) Analysis of Marble. Heat two grammes of marble dust in a small iron crucible ANALYSIS OF MARBLE. over a blast lamp, so long as the material contin- ues to lose in weight. The residue easily recog- nized as lime must be one of the constituents. Add water, test with litmus paper, and compare with a mixture of marble dust and water. What inference would you draw from Ex. 48 in regard to the probable constitution of such white pow- ders? We know that magnesium has a very strong attraction for oxygen (Ex. 48), and there- fore, to test this inference heat over a lamp in a small ignition tube two decigrammes of lime in powder mixed with one decigramme of mag- nesium powder. When cold add water, and shake up with residue. Note the evolution of hydrogen gas and the production of lime water. The metal liberated, which decomposes water, is calcium. Hence lime consists of calcium and oxygen. Wliat was driven off from the marble dust ~by Tieat f To show this, let the teacher procure half a metre of small iron gas pipe. Close one end by welding. Drop in ten grammes of marble dust and shake down to the closed end. Mount in a Fletcher gas furnace and connect the open end with a pneumatic trough. Heat to a full white heat, avoid excess of illuminating gas at the burner lest it diffuse through the tube. Col- lect and examine gas evolved. It will not sup- port combustion, it dissolves in water, and feebly 118 LABORATORY PRACTICE. reddens litmus paper. It is obviously carbonic dioxide. Of what does carbonic dioxide consist ? Take a length of magnesium ribbon. Ignite and plunge the burning end in a jar of carbonic diox- ide. Observe the separation of carbon. What is the necessary inference in regard to the composi- tion of carbonic dioxide ? Discuss all points of this evidence. (b) Synthetical Confirmation. Review in this connection Ex. 39 (&), and repeat on a small scale with the lime water obtained by the action of cal- cium on water. Discuss the phenomena as syn- thetical evidence of the composition of marble. Explain the action of hydrochloric acid on marble, bringing in contrast the two facts Marble and hydrochloric acid yields calcic chloride and carbonic dioxide. Lime and hydrochloric acid yield calcic chlo- ride. Confirm these facts experimentally and draw your own inferences. As will hereafter appear, the facts as above stated are not complete state- ments, since in both processes water is also formed, which in qualitative experiments escapes notice by mixing with the mass of liquid, acting as the medium of the chemical change. Still in the pres- ent case the fact overlooked was not material, and, since in experimental science we can fre- quently draw correct conclusions from similar in- CHEMICAL TESTS. 119 complete evidence, this experience may teach a valuable lesson. It was from exactly this evi- dence that the proximate composition of marble was first inferred by Dr. Black, of Edinburgh, a century ago. Synthetical processes are often of great value in confirming analytical results, and give fre- quently the most direct and efficient means of finding out the composition of a material. As commonly used the term chemical analysis in- cludes all methods of establishing the chemical constitution of substances, whether synthetical or strictly analytical. Ex. 67. Chemical Tests. In the examples of analysis given above, we have actually separated the elementary substances of which two familiar bodies consist. But such a separation is not al- ways practicable or necessary, and we can gen- erally discover the constituents of a substance by applying certain characteristic tests. Take four short lengths, not over 50 millimetres long, of platinum wire and make a loop at the end of each ; melt into the first of these loops chloride of sodium ; into the second, chloride of potassium ; into the third, chloride of strontium ; and into the fourth, chloride of barium. Hold the loops suc- cessively in the flame of a Bunsen lamp, and no- tice the colours which they impart to it ; and if a spectroscope is accessible, examine the coloured 120 LABORATORY PRACTICE. flames with this instrument. Almost any prepa- ration of sodium, potassium, strontium, or bari- um would produce the same effect ; and these characteristic colors, or still better the correspond- ing bands seen in the spectroscope, are indica- tions, or, as we usually say, tests of these metals. For another illustration, take five test tubes ; in the first dissolve a small amount of zinc dust in acetic acid ; in the second, a bit of white arsenic in hydrochloric acid ; in the third, a bit of anti- mony in hydrochloric acid to which has been added a drop of nitric acid ; in the fourth, a bit of iron in hydrochloric acid ; in the last, a bit of lead in weak nitric acid. In each case use a bit of metal no larger than a pin's head, and dissolve in the least amount of acid possible ; half fill the test tubes with water ; add to the first three the solution of sulphide of hydrogen obtained in Ex. 59 (#) ; to the fourth, a few drops of a solution of ferrocyanide of potassium ; and to the last, a few drops of a solution of potassic chromate. The characteristically coloured precipitates obtained under the conditions present are in each case tests of the several metals. So, in general, it is not necessary to isolate an acid, a metallic oxide, or an elementary substance, in order to prove that it is present or absent in a given case, but only to use the proper tests in the right way ; and the works on qualitative analysis teach us QUANTITATIVE ANALYSIS. 121 in what order and under what conditions the proper tests should be applied. The practice of qualitative analysis affords a most admirable train- ing in the methods of inductive reasoning. 17. Quantitative Analysis. Ex. 68. Analysis of Potassium Bromide. The analysis made in Ex. 65 may be made quanti- tative by first weighing the coin and afterwards weighing the silver and copper obtained. Of course if the coin consists of nothing else, the sum of the weights of the two metals ought to exactly equal the weight of the coin, and such a coinci- dence would go far to establish the accuracy of our work. In the analysis as above made only a very rough approximation to equality could be expected, but by more accurate methods such a confirmation of the work could be almost abso- lutely secured. In general, however, in order to determine the relative proportions of the different constituents in a compound, it is rarely practica- ble to separate the ingredients and weigh the sev- eral amounts. The method is to transfer each ingredient to some new combination which can be formed without loss, weighed with accuracy, and the composition of which through previous an- alyses is absolutely known. Take the simplest case. We wish to analyze common salt, which is 122 LABORATORY PRACTICE. known to consist wholly of chlorine and sodium. Neither of these are ingredients which can be ac- curately separated and weighed. So, with a care- fully weighed quantity of salt, we prepare chlo- ride of silver by a process in which we are sure that every particle of chlorine has been trans- ferred from its previous combination with sodium to a new combination with silver. Chloride of silver is a substance which can be collected and weighed with perfect precision. It has previously been accurately analyzed over and over again, and by referring to tables we find what fraction of the weight of chloride of silver thus found consists of chlorine, and then a very simple calculation gives the weight of chlorine sought. If the salt is abso- lutely pure the rest of the weight taken consists of sodium, and the amount of sodium could not be determined so accurately in any other way. Obviously, the analysis of common salt rests back on the analyses of chloride of silver previously made and recorded ; and so in most cases our an- alyses of to-day rests back on the work of those who have gone before us. After these relations have been explained, let the student weigh in a small beaker glass exactly one gramme of pure potassic bromide,* and dissolve the salt in about * With the rude manipulation here expected, potassium bromide will give more precise results than common salt and illustrates equally well the general methods of quantitative analysis. Potas- QUANTITATIVE ANALYSIS. 123 twenty -five cubic centimetres of water. Weigh in a similar beaker one and a half gramme of silver nitrate and dissolve in an equal amount of water. Pour now with constant stirring the first solution into the second, rinse the beaker, wash in the last drops, and allow to stand until the precipitate fully settles. Collect on a tared filter, wash dry, and weigh. It is known from previous work, and can be found by reference to any work on quanti- tative analysis, that every gramme of silver bro- mide contains 0'4255 gramme of bromine ; and practically an analyst would always assume that this value was given and at once calculate the amount of bromine in the weight of the silver bro- mide he had obtained, and this would be the amount in the weight of the bromide of potassium he had taken for analysis. To show the student, however, that results thus obtained rest back on previous analyses, let him make this additional determination, not that he can c.ompete with the old work, which has been often repeated with the greatest care in order to establish fundamental data for just such uses as are here indicated, but to the end that he may realize the actual relations in most problems of quanitative analysis. Weigh out exactly one gramme of pure metallic silver, slum bromide is a familiar medicinal salt, consisting of potassium and bromine, two elementary substances closely allied respectively to sodium and chlorine. 124 LABORATORY PRACTICE. place in a small beaker, and dissolve in about two cubic centimetres of strong nitric acid diluted with four cubic centimetres of water, and add twenty-five cubic centimetres of water. In a sec- ond beaker dissolve in twenty-five cubic centime- tres of water 1/2 gramme of potassic bromide, and then proceed as before. From the result cal- culate the weight of bromine in one gramme of silver bromide, which should be, within the limits of error, the same as the value given above. The student should now calculate from his own results the per cent of bromine and of potassium in po- tassic bromide, and should have practice in simi- lar calculations until he is familiar with the usual manner of stating the results of analysis in per cent. (1) It is known that pure common salt consists wholly of sodium and chlorine, and also that one gramme of silver chloride contains 0*2474 gramme of chlorine. In one deter- mination 0-5723 gramme of salt gave T4038 gramme of sil- ver chloride. Calculate the percentage composition of com- mon salt. Ans. Chlorine, 60 '69 Sodium, 39-31 100-00 (2) It is known that pure crystallized cane sugar con- sists wholly of carbon hydrogen and oxygen. By the usual process of organic analysis, 0'2569 gramme of sugar gave G'3966 gramme of carbonic dioxide and 0'1487 gramme of water. It is known that one gramme of carbonic dioxide contains 0'2727 gramme of carbon, and one gramme of water 0-1111 gramme of hydrogen. What is the percentage com- position of cane sugar ? QUANTITATIVE ANALYSIS. 125 Ans. Carbon, 4211 Hydrogen, 6 '43 Oxygen, 51 '46 100-00 (3) Obviously the above processes assume a qualitative knowledge of the composition of the substance analyzed, and so, in general, quantitative analysis implies a previous qualitative analysis. Indeed, the process of determining the amount of an ingredient present must constantly be varied according as it is associated with different substances, and a large knowledge is required in order to meet the conditions in any case and secure accurate results. Thus quantitative analysis becomes a distinct and widely extended branch of chemical study, and it is the chief work of the practical chemist. CHAPTER III. MOLECULES AND ATOMS. 18. Molecular Theory. THE theory of the new science of thermo-dy- namics assumes that the material of aeriform bod- ies is not continuously distributed through the spaces they seem to fill, but consists of a vast number of exceedingly minute particles in rapid motion to and fro, constantly rebounding from one another, or from the walls of the containing vessel. These minute particles are called mole- cules, and the phenomena of heat are supposed to be manifestations of their moving power. The molecules of the same substance, of water, for ex- ample, are supposed to have the same weight in fact, to be alike in every respect ; while the mole- cules of different substances are as unlike as the substances themselves. This theory has been worked out mathematically with great ability, and the phenomena of nature have been found, in a most remarkable manner, to conform to the de- ductions of mathematical analysis. Of these de- ductions one of the most remarkable is, that MOLECULAR WEIGHT. 127 Equal volumes of all gases or vapour s^ meas- ured under the same conditions, contain the same number of molecules. This deduction is usually called the law of Avogadro ; and if we accept the fundamental con- ception of molecular structure we must also ac- cept this inference which it involves. The mod- ern theory of chemistry accepts the law of Avo- gadro as a fundamental principle, and builds upon it a large superstructure. The law of Avogadro does not absolutely hold except when the material is in a perfectly aeri- form condition. It is only approximately true in the case of dense gases or vapours under the pressure of the air, and near the point of con- densation the deviations are sometimes very marked. It has no reference whatever to liquids or solids. These forms of matter are supposed also to consist of moving particles ; but, if so, the molecules must be variously compacted, and their motions otherwise circumscribed than in the aeri- form state. 19. Physical Method of Determining Molec- ular Weights. If the law of Avogadro is true, the molecular weight of a substance must be proportioned to its specific gravity in the state of gas or vapour. If 128 LABORATORY PRACTICE. we take hydrogen gas as the unit of reference for the specific gravity, and the molecule of hydrogen gas as the unit of reference for molecular weights, then the number which expresses the specific gravity of a substance in the state of gas or vapour would also express the molecular weight of that substance. For considerations which will shortly appear, half the weight of the molecule of hydro- gen has been taken as the unit of molecular weight, so that the molecule of hydrogen gas weighs two of the assumed units ; and hence, on this system, the molecular weight of any substance is found by doubling its specific gravity taken in the aeriform state, and referred to hydrogen gas as the standard. The physical method of determining molecular weights, therefore, reduces itself to finding the specific gravity of a substance when in the condi- tion of gas or vapour with reference to hydrogen. The substance must be in the condition of gas or vapour, and the method is only applicable to those bodies which are naturally aeriform, or which can be volatilized at temperatures within control without undergoing decomposition. Ex. 69. Density of Hydrogen. Use a flask not exceeding 100 cubic centimetres capacity. Cork tightly, and connect through cork a small chloride of calcium tube, so proportioning the parts that the flask will stand on the pan of the MOLECULAR WEIGHT. 129 balance. Add to the flask 20 cubic centimetres of strong hydrochloric acid and an equal volume of water. Weigh out closely 5 grammes of sheet zinc that has been carefully cleaned. Place this at the side of the flask on the scale pan and take the tare as closely as possible. Removing now the flask from the balance, connect, by means of a flexible-rubber connector, the chloride-of-calcium tube with the inlet tube of the gasometer before described (note to Ex. 45), which should be large enough to hold two full litres of gas. When all is ready withdraw the cork, drop in the zinc, and quickly recork the flask. Wait until the evolu- tion of hydrogen has altogether ceased, then shut off the gasometer and disconnect the flask. Be- fore replacing it on the balance pan withdraw the cork for a few minutes to give the hydrogen gas in the interior time to diffuse into the air. The loss of weight is obviously the weight of the hydrogen set free. To find the -volume of this hydrogen transfer the gas from the gasometer in successive portions to a litre measure over a pneu- matic trough. In reading the volumes take care that the level of the water is the same inside and outside the glass and avoid warming with the hands. This volume should be corrected for the tension of aqueous vapour (Ex. 22) and reduced to standard conditions, observing, for the purpose, the height of the barometer and the temperature 130 LABORATORY PRACTICE. of the water in the trough. We have, then, the weight of a measured volume of hydrogen, and can easily calculate the weight of one litre. With such appliances as are here assumed the process is accurate within about 5 per cent. On account of its great lightness the exact determination of the density of hydrogen gas is a difficult problem, and for the purpose of calculating the specific gravity of other aeriform bodies referred to hy- drogen as unity, we will assume the value gener- ally received, 0*0896, and for the specific gravity referred to air, 0'0692. (Compare Ex. 17.) Ex. 70. Specific Gravity of Carbonic Dioxide. Take a quart tin can, with narrow neck fitted with selfsealing stopper, and, having measured its exact contents, as in Ex. 17, carefully clean and dry it. Place it, open, on the balance-pan, and equipoise it with a second can of the same size and pattern, tightly sealed. Observe the thermometer and barometer at the time the equilibrium is estab- lished. It will be obvious now that if w T e calcu- late the weight of air which the open can holds at this temperature and pressure (Ex. 17) and add an equal weight to the pan carrying the open can, we should have what would be the exact tare if the can were sealed with all the air exhausted from the interior. Moreover, since the can and its counterpoise displace the same volume of air, it is also obvious that this equilibrium would not MOLECULAR WEIGHT. 131 be disturbed by any changes in the atmosphere. If, therefore, we fill the can with any gas for example, carbonic dioxide the increased weight will be simply the weight of this gas. Remove then the open can, fill it with carbonic-dioxide gas by displacement, and seal it, observing the thermometer and barometer at the moment the can is closed. Determine the increased weight ; and this is the weight at the last observed tem- perature and pressure of a volume of carbonic- dioxide gas equal to the known capacity of the can. From this value calculate what would be the weight of the same volume if the thermome- ter marked and the barometer stood at 30 inches. The density of carbonic-dioxide gas that is, the weight of a litre under the standard conditions of temperature and pressure is then found by dividing the weight of the gas by the capacity of the can. The density of carbonic dioxide divided by the density .of air gives the specific gravity of carbon dioxide referred to air ; or, if divided by the density of hydrogen, the spe- cific gravity referred to hydrogen ; and these val- ues are the same for all temperatures and press- ures. Why ? (1) What is the molecular weight of carbonic dioxide ? (2) The specific gravity of nitrous-oxide gas referred to hydrogen is 22*04. What is its molecular weight ? (3) The specific gravity of cyanogen gas referred to hy- drogen is 26*06. What is its molecular weight ? 132 LABORATORY PRACTICE. (4) The specific gravity of oxygen gas referred to hydro- gen is 16. What is its molecular weight ? Ex. 71. Specific Gravity of Vapours. Find the molecular weight of alcohol, ether, chloro- form, or ethylene bromide, by determining in either case the specific gravity of the vapour of the substance by the following method : Determine the volume of one of the bulbs provided for the purpose,* and, having thor- oughly dried both the interior and the exteri- or surface, seal the shorter tubulature. Select a second bulb of approximately the same size, and, having sealed both of its tubulatures, use it to equipoise the first, completing the tare as conven- ient. Observe now the thermometer and barome- ter, and calculate the weight of the dry air which fills the open bulb Ex. 17 (1). Add this weight to the pan holding the first bulb ; and, as thus load- ed, the balance would be in equilibrium were the glass vessel completely exhausted. Moreover, this constructive equilibrium will not be disturbed by any atmospheric changes (Ex. 70). Introduce now into the first bulb about 50 cubic centimetres of the volatile liquid under examination. Hang the bulb above the water in an ordinary tea-ket- * These bulbs hold about 400 cubic centimetres, and are blown in a mould so as to secure a uniform size. They have a long, nar- row stem and opposite to the stem a short and still narrower tubu- lature. MOLECULAR WEIGHT. 133 tie sufficiently capacious for the purpose, with the longer tubulature projecting through a cork fit- ting a hole in the cover. Boil the water under the bulb so that the steam surrounding the glass and escaping by the nozzle of the kettle shall maintain a uniform temperature near 100. Ob- serve this temperature with a thermometer pass- ing through a cork fitted to a second hole in the cover, and as soon as the current of vapour from the bulb stops seal the tubulature by melting with a blowpipe the glass at the tip. At the same time note the height of the barometer. When cold, replace the bulb on the balance and deter- mine the increased weight above the equilibrium just described. This value is the weight of the vapour which filled the bulb when in the kettle at the temperature and pressure observed. Find next what must have been the slightly increased volume of the bulb when in the kettle, using the formula V' = V(1 + 0-000024 x t), and calculate what would be the weight of the same volume of dry air at the temperature and pressure in the kettle. Then the weight of the va- pour divided by the weight of the air gives the spe- cific gravity of the vapour referred to air, and this result multiplied by 14 "43 gives the specific gravity of the same vapour referred to hydrogen gas. To make sure that the bulb when sealed was full of 134 LABORATORY PRACTICE. vapour, break off the tip of the tubulature under water (recently boiled to drive out the dissolved air), when the liquid should rush in and complete- ly fill the interior. If any considerable volume of residual air then appears (more than five or ten cubic centimetres) the determination should be re- peated, using more liquid and taking more care to seal the bulb at the right time. If, as is usu- ally the case, the material used is combustible, the right moment is easily caught by lighting the jet of vapour as it issues from the tubulature (after the first violent rush has ceased), and watch- ing the flame as it diminishes. The moment the flame disappears the tubulature should be sealed. In repeating the experiment it is of course unne- cessary to alter the tare or disturb the equilibri- um if only the tips broken off are kept and re- turned to the balance pan with the bulb. (1) The student should now answer the following ques- tions : (1.) What was the density of the vapour which filled the bulb when in the kettle at the temperature and pressure noted ? (2.) According to what laws does the density of a dry vapour vary when it freely partakes of the temperature and pressure of the surrounding medium ? (3.) Does a va- pour confined over the liquid from which it rises conform to the same laws ? (4.) Why does not the specific gravity of a vapour vary with the temperature and pressure ? (5.) Why is the molecular weight of a substance equal to twice its., specific gravity in the aeriform state referred to hydrogen (2) The student should carefully review in this connec- tion Subdivision 2, on Air. MOLECULAR WEIGHT. 135 (3) The above method obviously only applies to such liquids as boil below the boiling point of water. With less volatile liquids the bulb may be sunk, by means of appro priate apparatus, in a bath of melted paraffine maintained at a constant temperature ; and the specific gravity of the va- pours of comparatively fixed bodies has been formed by using globes of porcelain heated in a bath of boiling zinc by means of very powerful furnaces. Quite a different method of experimenting * is better adapted to such cases, but it is beyond the scope of this book. (4) Sharp results conforming to theory can not be ex- pected by the method here described unless the materials used are of a very high degree of purity. The alcohol must be absolute and the ether or chlorform free from all admixt- ures. Obviously a process of fractional distillation takes place in the bulb and the impurities may be concentrated in the residual vapour which is weighed. This objectionable feature of the process is avoided in still a third method, ap- plicable only to comparatively volatile liquids in which the volume of vapour formed by a weighed amount of sub- stance is accurately measured under observed conditions. For a description of this process see loc. cit. in (3). 20. Chemical Method of determining Mo- lecular Weights. From the chemist's point of view, the molecule of a substance is the ultimate particle which pos- sesses the qualities of the substance. Molecules may be divided, but if divided the properties of the substance are lost, new molecules are formed, and new substances with new properties appear. In every chemical process the change must take * See author's Chemical Philosophy, pp. 32-37. 136 LABORATORY PRACTICE. place between molecules ; that is, one or more molecules of one substance must act upon or must yield one or more molecules of another substance. Hence it must be that in any chemical change the weights of the substances involved must be in the proportion of their molecular weights, or in some multiple of this proportion. In other words, assuming the fundamental conception of molecu- lar structure, the law of definite proportion is a necessary deduction of the molecular theory. And not only is the law of definite proportions a fixed principle of nature, but, moreover, the definite proportions shown by chemical anatysis are found to bear a very simple numerical relation to the molecular weights measured by the vapour densi- ties. Thus the facts of chemistry furnish a very remarkable confirmation of the molecular theory of physics. Moreover, the methods of quantita- tive chemical analysis give us another method of determining molecular weights ; for if in any chemical process we can find the quantitative rela- tions between any two of the substances con- cerned, whether as factors or products, the ratio of these weights must be the proportion of the molecular weights, or else some simple multiple of that proportion ; and in most cases we are able to infer from the chemical relations what the multi- ple is. Ex. 72. Molecular Weiglits of Potassic Ohio- MOLECULAR WEIGHT. 137 rate and ofPotassic Chloride. Weigh, in a porce- lain crucible (or better, a small platinum crucible) about two grammes of dry potassic chlorate (pow- dered). Heat to fusion, gradually increasing the temperature as the oxygen gas escapes, taking care to avoid sputtering, and finally heating to low redness for several minutes. Weigh the resi- due. Make the proportion, As the weight of the oxygen driven off is to the weight of potassic chlorate taken, so is 32, the molecular weight of oxygen gas assumed to be known, to the corre- sponding weight of potassic chlorate. The weight thus found is known to be two thirds of the mo- lecular weight. What is the molecular weight? Make also the proportion, As the weight of oxy- gen expelled is to the weight of potassic chloride left, so is 32, as before, to a number which is known to be two thirds of the molecular weight of potassic chloride. What is the molecular weight of potassic chloride ? Ex. 73. Molecular Weight of Oxalic Acid. To a small and light glass fiask fit a cork with two perforations to one of these adapt a small chlo- ride-of-calcium tube, and to the other a short out- let tube. Weigh in the flask about one gramme of crystallized oxalic acid, determining the weight with precision. Dissolve in fifty cubic centime- tres of water and add ten cubic centimetres of strong sulphuric acid; allow to cool, and then 138 LABORATORY PRACTICE. add one gramme of powdered pyrolusite. After the apparatus is thus mounted take the tare. Then, closing the outlet tube with a small bit of wax, gently heat the flask so long as the evolu- tion of carbonic dioxide continues. Again allow to cool, and when cold remove the wax stopper and suck through the chloride-of-calcium tube so long as the taste of carbonic acid is perceptible. Lastly, determine the loss of weight, which is the weight of the carbonic dioxide formed in this somewhat complex chemical process. Still the principle holds that as the weight of the carbonic dioxide formed is to the weight of the oxalic acid used so is 44, the known molecular weight of car- bonic dioxide, to a number which must be a sim- ple multiple or submultiple of the molecular weight of crystallized oxalic acid. In this case the result will be one half of the required molecu- lar weight. Let the student see that in this method it is only necessary to know the relative weights of two of the substances concerned in the chemical process, which may be very complex and in regard to which nothing else need be known. Let him also notice that the chemical method has the advantage over the physical method in that it is applicable to substances which are not volatile. It adopts the same unit as the physical method, and refers the unknown molecu- lar weight to a molecular weight previously deter- ATOMS. 139 mined and in the first instance controlled by the physical method ; but in this way, step by step, it covers the whole ground. It is far more accurate than the physical method, and practically the physical method is only used to control the re- sults of chemical analysis that is, to show whether the definite proportions observed are the relations between single molecules or between multiple molecules. Since, however, the chem- ical method involves the question of multiple ratios, it has its necessary limitations. When the substance under examination is volatile, the re- sults of analysis can, as just said, be controlled by a determination of vapour density. In other cases a study of the chemical process itself gives us the additional information required, but in a way that can not be made intelligible in this con- nection. Not unfrequently, however, all these means fail, and then the chemist is forced to se- lect between the possible multiples the value which he thinks most probable. 21. Conception of Atoms. The ultimate particles of substances, called molecules, although far beyond the range of visi- bility, are by no means inconceivably small, and still less indivisible ; and the teacher should at- tempt to aid the student's imagination by giving 140 LABORATORY PRACTICE. the estimates of physicists in regard to the abso- lute size of these bodies, and showing that the relations between their dimensions and our ordi- nary standards of magnitude are no more extreme than those we meet with in astronomy, in elec- tricity, and in other branches of physical science. So far are the molecules from being indivisible that it is perfectly evident that they must be di- vided in almost every chemical change. For ex- ample, as we have seen, two volumes of hydrogen gas combined with one volume of oxygen gas to form two volumes of aqueous vapour. Here it is evident that if equal gas volumes contain the same number of molecules, it must be that every two molecules of hydrogen gas combine with one molecule of oxygen gas to form two molecules of water that is to say, the molecule of oxygen is divided between two molecules of water ; or, again, every molecule of water contains one half as much oxygen as the molecule of oxygen gas. We reach the same result in the analysis of water. If we calculate from the results of analysis the percentage composition of water, we find that in 100 parts water contains of hydrogen 11 "111 per cent, and of oxygen 88*888 per cent. Further, the specific gravity of aqueous vapour referred to hydrogen is 9, and hence the molecular weight of water is 18. Of this weight, 88 '888 per cent, or 16 parts, consist of oxygen. Again, the specific ATOMIC WEIGHT. gravity of oxygen gas referred to hydrogen is 16, and therefore the molecular weight of oxygen gas is 32. We know then that One molecule of oxygen gas weighs 32 microcriths. One molecule of water weighs 18 microcriths. One molecule of water contains 16 microcriths of oxygen. Hence, as before, every molecule of water con- tains one half as much oxygen as the molecule of oxygen gas. Obviously we can repeat this calcula- tion with every compound of oxygen in regard to which we know the molecular weight and the per cent of oxygen which the compound contains. If now we arrange the results in a table, as below ATOMIC WEIGHT OF OXYGEN. Compounds of Observed Weight of Mole- Weight of Oxygen. Sp. Gr. cule. Oxygen in Molecule. M. c. M. c. Water 9'00 18 16 Carbonic oxide 13 '95 28 16 Nitric oxide 14'97 30 16 Alcohol 23-28 46 16 Ether 37'32 74 16 Carbonic dioxide 22 '06 44 32 Nitric dioxide 24 '82 46 32 Sulphurous dioxide 32 '24 64 32 Acetic acid 30'07 60 48 Sulphuric trioxide 39 '87 80 48 Osmic tetroxide 128'30 263 '2 64 Oxygen gas 15'96 32 32 it will appear that in every case the molecule of an oxygen compound contains either sixteen mi- 10 142 LABORATORY PRACTICE. crocriths of oxygen or a simple multiple of six- teen microcriths. The smallest amount of oxygen in any molecule is sixteen microcriths, and this is the weight of what we call an atom of oxygen. The word "atom" is derived from a Greek word meaning indivisible ; and this smallest known mass of oxygen, weighing sixteen microcriths, has never been divided. The molecule of oxygen gas consists of two atoms, and of course can be broken in two. We can reason in regard to the com- pounds of every other elementary substance in precisely the same way and make a similar table,* and in each case we shall find that the weights of a given element in the molecules of its several compounds are simple multiples of a smallest weight, which we take as the weight of the atom of that element. In the case of the compounds of hydrogen the smallest weight is one microcrith, which is the weight of the atom of hydrogen the smallest mass of matter recognized by science. It is not unreasonable, therefore, that it should be selected as the unit of molecular and atomic weights ; and we call this unit by a definite name, a microcrith, in order that the student may asso- ciate with the name a real if not a tangible mass of matter. The molecule of hydrogen gas, like the molecule of oxygen gas, contains two micro- * See author's Chemical Philosophy, pp. 42-45, or New Chemis- try, pp. 141. ATOMIC WEIGHT. 143 criths, and hence before we attained the concep- tion of the hydrogen atom we described correctly the unit of molecular weight as the half -hydrogen molecule. In this way our conception of atoms and our general knowledge of atomic weights have been reached, and in every work on chemis- try the values of the atomic weights adopted are given in tables opposite to the names of the ele- ments. (See table at end of this book.) The student must seek to make clear to his mind the distinction between the conception of the atom and the conception of the molecule. The ultimate particles which retain the qualities of a substance are molecules, and there are as many kinds of molecules as there are substances. Atoms are the smallest masses of the chemical elements yet known, and there are only as many kinds of atoms as of elements. To speak of an atom of a substance, especially of a compound substance, is a misuse of terms. In the case of elementary substances we have still to distinguish between the molecules of the substances and the atoms of the element. There is but one kind of atom of any element, but there may be several distinct elementary substances. Thus in the case of oxy- gen we have oxygen gas, the molecules of which consist of two atoms of oxygen, and ozone, a wholly different substance, the molecules of which consist of three atoms of oxygen. The 144 LABORAT011Y PRACTICE. molecules of elementary substances are formed by the aggregation of atoms of the same kind ; the molecules of compound substances by the aggre- gation of atoms of different kinds. There are a few cases, like the vapours of mercury and zinc, where the molecule consists of a single atom. In a chemical change the molecules of the substance we call the factors break up into atoms, which group themselves into new associations to form molecules of the products. 22. Determination of Atomic Weights. The exact determination of an atomic weight now resolves itself into a simple question of quantitative analysis. If in any process of quan- titative analysis we can determine the weights of two of the elementary substances involved, the proportion between these quantities will be either the ratio of the atomic weights of the two ele- ments, or else that of some simple multiple of these weights, the multiple in all cases being pre- viously known from the relations of the com- pounds of the element as exhibited in such tables as we have described, or otherwise. Ex. 74. Atomic Weight of Zinc. Adapt to a small flask (100 cubic centimetres) a tightly fitting cork and exit tube leading to a pneumatic trough. Place in the flask 10 cubic centimetres of strong ATOMIC WEIGHT. 145 hydrochloric acid and 20 cubic centimetres of water. Clean scrupulously a strip of the purest sheet zinc, and accurately determine its weight, which should be as nearly 1'25 gramme as pos- sible. Use as a receiver a glass flask of 500 cubic centimetres capacity. When all is ready, the flask filled with water standing inverted on the shelf of the trough and the mouth of the exit tube under its lip, drop the metal into the acid and quickly cork the flask. This amount of metal should yield nearly 500 cubic centimetres of hy- drogen gas at the ordinary temperature of the laboratory. When the apparatus is in equilib- rium notice whether any water has been sucked back towards the flask. If so, make the neces- sary allowance in measuring the volume of gas formed. Observe the thermometer and barome- ter. Sinking now the flask in the water of the trough until the level of the water is the same within and without the glass, place the palm of the hand under the mouth and quickly invert the flask and place it on the pan of the balance with the water it still holds. Take the tare, and, hav- ing exactly filled the flask with water, again weigh. The difference of these weights that is, the weight in grammes of the water required to fill the flask, reduced for temperature if more ac- curacy is required will give the volume of the gas collected at the temperature and pressure ob- 146 LABORATORY PRACTICE. served. Reduce the volume to standard condi- tion, making allowance for the tension of aqueous vapour (Ex. 21 and Ex. 22). From this volume calculate the weight of hydrogen formed. Make then the proportion, As the weight of hydrogen is to the weight of zinc, so is unity (the atomic weight of hydrogen) to a value which we know, from a comparison of the zinc compounds, to be one half of the atomic weight of zinc. Double the value to find the accepted atomic weight. In a similar way the atomic weight of magne- sium may be found ; and by dissolving alumi- num in a solution of caustic soda the atomic weight of aluminum may be determined with great accuracy. As the atomic weights are fun- damental constants in chemical calculations, it is essential that they should be known with the greatest possible precision ; and hence a great- deal of labor has been spent on the analytical processes used in determining their value. These processes admit of very different degrees of accu- racy. There are only a very few which in the most skilful hands yield results accurate to the ten- thousandth part of the quantity estimated ; and even the thousandth part is regarded as a very high degree of accuracy in chemical analy- sis. Most processes do not give results which can be relied upon much within one per cent ; and in many cases we are forced to use methods that are ATOMIC WEIGHT. far less accurate even than this. In fixing the precise value of an atomic weight our choice is usually limited, both as to the material to be analyzed and the analytical process to be used, to one or two methods ; but in almost all cases there are abundant analyses of compounds of the same element which are sufficiently accurate to enable us to interpret the results obtained. (6) Examples illustrating the above points may be multiplied by the teacher. Thus, the produc- tion of silver bromide from silver or the reduction of silver bromide to silver gives the means of con- necting the atomic weight of bromine with that of silver. (Compare Ex. 65 and Ex 68.) So also the reduction of silver nitrate in a porcelain crucible on simple ignition will enable the student to de duce the molecular weight of silver nitrate from the atomic weight of silver, and by inference from what will soon appear he can thus determine also the molecular weight of nitric acid. (1) It will be observed that tlie method of determining an atomic weight is essentially the same as the chemical method of determining molecular weights. In each case the method is based on the law of definite proportions, which applies to elementary substances as well as to compounds, only in one case the definite proportions are theoretically interpreted as the definite relative weights of atoms, and in the other as the equally definite relative weights of mole- cules. In all instances what we can determine with accuracy experimentally is a relative weight. What that relative weight represents is a question of interpretation. The ratio 148 LABORATORY PRACTICE. of the two weights determined forms the first two terms of a proportion of which the third term is some known molecu- lar or atomic weight. We thus can connect one molecular weight with another or one atomic weight with another. Moreover, since we refer both molecular and atomic weights to the same unit, we can connect a molecular weight with an atomic weight, as is constantly, in fact, done. Always OUT results are subject to interpretation so far as regards the question of multiple values (Ex. 73). CHAPTER IV. SYMBOLS AND NOMENCLATURE. 23. Chemical Symbols. THE initial letter, or letters, of its Latin name are used to represent one atom, and therefore the atomic weight, of each chemical element. Thus H (Stands for 1 microcrith of the element hydro- gen ; O, for 16 microcriths of the element oxygen ; C, for 12 microcriths of the element carbon. Sev- eral atoms are represented by means of subnumer- als ; as S a , which stands for 2 X 32 = 64 micro- criths of sulphur ; C1 8 , which stands for 3 X 35*5 = 106 '5 microcriths of chlorine. Molecules are repre- sented by writing together the 'symbols of the atoms of which they consist. Thus, H 2 O stands for a molecule of water because each molecule of water is made up of two atoms of hydrogen and one atom of oxygen ; H a SO 4 stands for a molecule of sulphuric acid, consisting of two atoms of hydrogen, one atom of sulphur, and four atoms of oxygen ; O 2 stands for a molecule of oxygen gas, an aggregate of two molecules of oxygen; while O 3 stands for a molecule of 150 LABORATORY PRACTICE. ozone gas, which, although consisting solely of oxygen, has molecules made of three atoms in- stead of two, and is a wholly different substance. The molecular symbol represents the molecular weight, which is the sum of the weights of the atoms of which the molecule consists. Thus the molecular weight of sulphuric acid is (2xl) + 32 + (16 X 4) = 98. From a molecular symbol we can always deduce the percentage composition of the substance it represents. Thus it is obvious that -g of a molecule of sulphuric acid consists of hydrogen, f f of sulphur, and || of oxygen. The substance, having the same composition as the molecule, must then contain in one hundred parts Hydrogen 2'04 Sulphur 32-65 Oxygen 65 -31 4 100-00 On the principle of Avogadro, all molecular symbols represent the same volume in the state of gas ; thus Hydrogen Gas. Ozone. Carbonic Dioxide. H a = 2 m.c. O 3 = 48 m.c. CO 2 = 44 m.c. Water. Alcohol. H 9 = 18 m.c. C 2 HeO = 46 m.c., all represent the same gas volumes compared un- der the same conditions of temperature and press- ure. It follows, then, that the specific gravity of a gas or vapour referred to hydrogen can be at CHEMICAL SYMBOLS. 151 once deduced from the molecular symbol by divid- ing the molecular weight by 2. The weight of a litre of hydrogen when the barometer stands at 30 inches and the thermometer at is 0*0896 gramme. At 273 under the same pressure the weight would be 0*0448 gramme. At 27 (a very convenient standard)* the weight would be 0*0815 gramme. Hence the weight of a litre of any gas or vapour, under either condition, may also be calculated from, the molecular symbol by multi- plying the specific gravity obtained as above by one of these factors. * If we select 300 absolute temperature that is, 27 C. for our standard temperature and 30 inches of mercury for the stand- ard pressure all reductions of gas volumes can be made with the greatest facility. A variation of one degree from this standard temperature corresponds exactly to a variation of one tenth of an inch in the barometer, and the effect of temperature can at once be eliminated by altering to a corresponding extent the height of the mercury column measuring the pressure. Moreover, as our observa- tions are almost invariably made at temperatures below the stand- ard (27 C., or 80-6 F.), this correction is usually additive. As- sume, for example, that the temperature is 20 and the observed height of the barometer 3O3 inches, and it is desired to reduce the observed gas volume to the assumed standard. Were we to raise the temperature to 27 it is obvious that we should expand the gas ; and to bring it back to its previous volume it would be necessary to increase the pressure by 0*7 inch, which corresponds, as we have stated, to seven degrees. It is thus evident that at 27 (the assumed standard) and at 31 inches the volume of gas would be the same as at 20 and 30-3 inches. The problem is then reduced to this : Given a volume of gas at 31 inches, to find what would be its volume at 30 inches ; and this is obviously a very simple case under Mari- otte's law. In like manner all similar problems may be solved, whether they relate to the volume of a given weight of gas or to the weight of a given volume. 152 LABORATORY PRACTICE. The specific gravity referred to air, or the den- sity under other conditions of temperature and pressure, may now be deduced by the methods before described. Let the student now answer, in regard to each of the molecular symbols last given, this question : What information does the symbol give in regard to the substance it repre- sents? He ought also now to be able, without further assistance, to reverse the reasoning, and, when the percentage composition and vapour density are given, to deduce the symbol. Assume that we have given as the result of analysis that the percentage composition of abso- lute alcohol is as below, and that the specific gravity of the vapour referred to hydrogen is 23. The weight of a molecule of alcohol is then 46, and the amount of each element in a molecule is that given in the second column of figures. Knowing now the weight of the several atoms, we easily find the number of each kind in one molecule Carbon .... 52 '17 24 = 2 x 12 C 3 Hydrogen ... 13 '05 6=6x1 H 6 Oxygen .... 3478 16 = 1 x 16 O 100-00 46 By studying the above scheme it will be seen that we shall reach the same result if we at once divide the percentages by the atomic weights, and CHEMICAL SYMBOLS. 153 then seek the simplest ratio of whole numbers corresponding to the results. Since 52-17 : 13-05 : 3478 = 24 : 6 : 16, it must be that 52-17 13-05 34-78 _ 24 6 16 _ 12 1 16 "12 : 1 : 16~ So when we do not know the molecular weight we can always find the simplest ratio of whole atoms corresponding to the percentage composi- tion, and the true symbol of the compound must be either the symbol thus obtained or some mul- tiple of it ; and the molecular weight must be either the weight represented by this symbol or a multiple of it. In this way we can always find a symbol corresponding to the percentage composi- tion of minerals and of similar inactive and non- volatile chemical products, and we accept the symbol thus obtained until further investigation shows that some multiple of it is more correct. In practice the result is often indefinite, because the percentages, owing to errors of analysis, are not exactly known, and we obtain a proportion which is only approximately a simple ratio of whole numbers. In this case we select what we regard as the most probable ratio, and a good deal of judgment is necessary in inte^pjcalin^the re- sults. * * fc> OF THE 154 LABORATORY PRACTICE. (1) Given the percentage composition of chloroform as follows : Carbon, 10 '04 ; hydrogen, 83 ; chlorine, 89 '13. Required the symbol, knowing that the specific gravity of chloroform vapour equals 5975. Ans. CHC1 3 . (2) The percentage composition of sugar is given Ex. 68 (2). What is the symbol ? Ans. C 12 H 22 On. (3) Calculate the percentage composition of nitro-benzol, CeH 5 N0 2 . Ans. Carbon, 58'53 ; hydrogen, 4'07 ; nitrogen, 11-39 ; oxygen, 26 '01. (4) What is the weight of one litre of alcohol vapour at 273? 24. Chemical Reactions. As a chemical process consists in the breaking up of the molecules of the factors into atoms and the regrouping of the same atoms without loss to form the molecules of the products, it is obvious that every chemical change may be represented by an equation, writing the symbols of the mole- cules of the factors in the first member and the symbols of the molecules of the products in the second member, using figures like coefficients in algebra to indicate the number of molecules in- volved in the process and plus signs to separate the symbols. Thus the chemical change in the preparation of oxygen gas from potassic chlorate (Ex. 23), is represented by the equation Potassic Chlorate. Potassic Chloride. Oxygen Gas. 2KClOs = 2KC1 + 3O a Such an expression is called in chemistry a re- action. It presupposes a knowledge of the sym- CHEMICAL REACTIONS. 155 bols of the products and of the factors and also of the number of molecules which concur in the process. Again, the preparation of hydrogen gas from zinc and dilute sulphuric acid (Ex. 25) is rep- resented thus : Dilute Solution of Hydrogen Zinc. Sulphuric Acid. Zinc Sulphate. Gas. Zn + H 2 SO 4 + A = (ZnSO 4 + A) + H a Here Aq indicates an indefinite amount of water used to dilute the acid or dissolve the salt. So also the formation of ammonia (Ex. 45) is ex- pressed by this reaction : Nitric Oxide. Hydrogen Gas. Water. Ammonic Gas. 2 NO + 5H a = 2H a O + 2 NIL In like manner the student should review all the experiments he has tried, and with the aid of the teacher learn to write the reaction in every case. He should be required to state the full signifi- cance of every reaction he writes until he has ac- quired a complete mastery of this symbolical lan- guage. If correctly written, a chemical reaction illus- trates always the first two, and, when it involves aeriform factors or products, all three of the fun- damental laws of chemistry. The law of con- servation of mass is expressed by the equation sign, the law of definite proportions by weight is indicated by the definite numerical values of the several terms, and the law of definite proportions 156 LABORATORY PRACTICE. by volume is seen in the simple ratios of the co- efficients. In like manner the system of chemical symbols involves most of the principles which have been discussed in this course of experiments. For example, we see at once from the reaction why in determining the atomic weight of zinc in Ex. 74 the observed value was doubled, since what we observed was the ratio H 2 : Zn, and what we required the ratio H : Zn. So also in Ex. 66 (b) we see that water must have been formed by the reactions in question. Indeed, so fully do the symbols embody the general principles of chem- istry that students are apt to infer that these prin- ciples have been deduced from the symbols, just as in mathematics similar general results have been discovered by the working out and interpretation of algebraic formulae ; and when, as in many text books, the phenomena are subordinated to the symbolical expression, this false impression is in- evitable. Chemistry is an inductive, not a de- ductive science, like mathematics, and chemical symbols differ essentially from mathematical for- mulae. In mathematics everything that can be legitimately deduced from an algebraic equation must be true ; but this is far from being the case with chemical reactions. Chemical symbols sim- ply stand for the facts and theories they were de- vised to express and for nothing more. To secure the peculiar discipline of the physical sciences it CHEMICAL REACTIONS. 157 is essential that they should be studied as they have been built up. The student must begin by observing the phenomena, and be led up to the general principles through his own inferences, and this is the order which has been followed in preparing this book. To begin with an abstract statement of these principles, or, what amounts to the same thing, to express at once every phe- nomenon observed in symbolical language which embodies these principles, is to invert the natural order and to abandon the inductive method. Nevertheless, such are the perfection and grasp of this system of symbols that it is of the greatest value in aiding us to realize relations and foresee results which without it we might not have dis- covered. (1) The direction above given, that the student should review all the experiments heretofore described and write all the reactions of the processes described so far as it is possi- ble, can not be too strongly insisted upon. Without such practice the student can not be expected to grasp the sub- ject; but with it many of the relations before observed will become clear and the facts will all appear in a clearer light. To aid him we give below the more important reactions, the figure prefixed being the number of the experiment under which the process is described or indicated : 14. (a) (H,,C4H 4 O 6 +HNaCO s + Aq)= (HNa,C4H 4 6 + H 2 O + Aq) + C0 a 14. (&) 16. 23. 24. (a) 24. (c) 11 158 LABORATORY PRACTICE. -* 25. Zn + (HaS0 4 + Aq) = (ZnSO 4 + Aq) + H 2 27. 2H 2 + O 2 = 2HaO 32. S + Oa = S0 2 (SOa + HaO + Aq) = (HaSOs + Aq) 33. 2SO 2 + O a = 2SO 3 (SO 3 + H a O + Aq) = (HaSO 4 + Aq) (BaCla + HaSOi + Aq) = Ba 2 S0 4 + (2HC1 + Aq) (PaO 6 + 3H 3 O + Aq) = (2H 3 PO. + Aq) 34. (NaCl + H 2 SO 4 + Aq) = (HNaS0 4 + Aq) + 2HC1 35. 2HC1 + Na a = 2NaCl + H a 36. (MnOa + 4HC1 + Aq) = (MnCla + 2H a O + Aq) + Cl, Ha + Cla = 2HC1 39. (a) (CaCO 3 + 2HC1 + Aq) = (CaCla + HaO + Aq) + C0 a 39. (6) (Ca0 2 H 2 + Aq) + CO 2 = (CaCO, + (H 2 O + Aq) 40. (a) CO 2 + C = 2CO 41. CaHaO + HaSO* = (H 2 SO 4 .HaO) + Call* CH 4 = Marsh Gas, CaH 4 =Ethylene, CaH a = Acetylene. 43. (a) KNOs + H 2 SO 4 = HKSO 4 + HNO 3 43. (6) 2HNO 3 + S = HaSO 4 + 2NO 43. (c) 2HNO> + 5Cu = 5CuO + HaO + N a 44. (a) 3Cu + (8HN0 8 + Aq) = (SCuNaOo + 4H 2 + Aq) 4- 2NO 2NO + O s = 2NOa 44. (&) 10NO + P 4 = 2PaO 6 + 5Na 45. 2NO + 5Ha = 2H,0 + 2NH 3 48. 2Mg + Oa = 2MgO MgO + HaO = MgOaHa MgO + (HaSO 4 + Aq) = (MgS0 4 + H 2 O + Aq) MgOaHa + (H 2 SO 4 + Aq) = (MgSO 4 + 2H 2 O + Aq) Mg + (H 2 S0 4 + Aq) = (MgSO 4 + Aq) + Ha 2Zn + Oa = 2ZnO ZnO does not unite directly with water. ZnO + (HaSO 4 4- Aq) = (ZnS0 4 + HaO + Aq) 49. Na a + (2H 2 O + Aq) = (2NaOH + Aq) + H 2 (NaOH + HC1 + Aq) = (NaCl 4- H 2 + Aq) (NaOH + HNO 3 + Aq) = (NaNOi + H 2 O + Aq) (NaOH + CO 2 + Aq) = (HNaCO. + Aq) 50. (c) Na a + Cla = 2NaCl 2Na a + O a = 2Na a O CHEMICAL REACTIONS. 159 Na a O + H S O = 2NaOH (Na a O + 2HC1 + Aq) = (2NaCl + H a O + Aq) (NaOH + HOI + Aq) = (NaCl + H 2 O + Aq) 51. (2NH 3 + H 2 SO 4 + Aq) = ((NH 4 ) a SO4 + Aq) (NH. + HN0 3 + Aq) = (NH 4 NO S + Aq) (NH 3 + HC1 + Aq) = (NH 4 C1 + Aq) (2NaOH + H 2 SO 4 + Aq) = (Na a S0 4 + 2H a O + Aq) (NaOH + HNO 3 + Aq) = (NaNO 3 + H 2 O + Aq) (NaOH + HCl + Aq) = (NaCl + H 2 O + Aq) 53. (a) Cu + = CuO CuO + H 2 = Cu + H.O 54. (a) (CuO + H 2 SO 4 + Aq) = (CuSO 4 + H 2 O + Aq) (CuO + 2HNO 3 + Aq) = (CuN 2 O + H 2 O + Aq) (CuO + 2HC1 + Aq) = (CuCl a + H 2 O + Aq) 57. (a) Fe + H 2 SO 4 + Aq) = (FeS0 4 + Aq) + H a 58. FeO = Ferrous Oxide Fe 2 O 3 = Ferric Oxide FeOSO 3 (or FeSO) = Ferrous Sulphate Fe 2 O 3 .3SO 3 = Ferric Sulphate 59. (a) Fe + S = FeS 59. (6) FeS + (H 2 SO 4 + Aq) = (FeSO 4 + Aq) + H 2 S FeO + (H,SO + Aq) = (FeSO 4 -f Aq + H a O) 65. (AgNO 3 4- HCl + Aq) = AgCl + (HNO 3 + Aq) SAgCl + H. = 2Ag + 2HC1 68. (AgN0 3 + KBr + Aq) = (KNO + Aq) + AgBr 73. (MnO, + H 2 S0 4 + H 2 C a O 4 + Aq) = (MnSO 4 + 2H 2 O + Aq) + 2CO 9 (2) To ensure a full comprehension of the subject the teacher should ask such questions, as the following : Does it appear from 27 that when hydrogen unites with oxygen two volumes of the first combine with one volume of the second to form two volumes ? Does it appear that when either charcoal or sulphur burn in oxygen gas the volume of the product is the same as the volume of the oxygen consumed ? In the ordinary process of preparing oxygen gas, how does the residual salt differ in composition from the salt used ? Both zinc and zinc oxide when dissolved in dilute sul- 160 LABORATORY PRACTICE. phuric acid yield the same zinc sulphate. Why is hydrogen gas evolved in the first case and not in the second ? How do you explain the production of nitric acid from nitre Ex. 43 (a) ? When iron dissolves in nitric acid either nitric oxide or nitrogen gas is evolved (as in case of copper), while from dilute sulphuric acid the same metal liberates hydrogen. Why the difference ? When to a solution of a silver coin we add hydrochloric acid all the silver is precipitated as chloride, but none of the copper. Why this selection ? By the judicious use of such questions the student will be led to think, the deadening effect of mechanical routine will be avoided, and the teaching power of the course greatly increased. 25. Stochiometry. Since in writing a chemical reaction the rela- tive weights of all the factors and products are ne- cessarily implied, it follows that if the total weight of any one substance concerned in the process is given the weight of every other may be calculated. It is only necessary to make the pro- portion As the total molecular weight of the given sub- stance is to the total molecular weight of the re- quired substance, so is the gross weight given to the gross weight required. By total molecular weight is here meant the simple molecular weight of the substance multi- plied by the coefficient with which it appears in the reaction. If in such problems a volume is STOCHIOMETRY. 161 given this volume must be reduced to weight by the simple method already described before ap- plying the rule ; and when the volume of a factor or product is sought the reverse reduction is readily made after the weight is known. The re- lation between any two gas volumes is of course directly seen on inspecting the reaction, and needs no calculation. The student ought to have a great deal of practice in stochiometrical calcula- tion. A very large number of problems of this sort will be found in the author's Chemical Phi- losophy, and there are many works wholly de- voted to the subject. The teacher, however, will add interest to a necessarily dry subject if he constructs problems of his own, based on the ex- periments which the student has actually per- formed. (1) How many grammes of common salt can be made from 25 grammes of sodium bicarbonate ? (2) In the process of making nitric acid, how many grammes of sulphuric acid will be required to every kilo- gramme of nitre, assuming that the acid used contains 95 per cent of HaSO* ? How many cubic centimetres would be required ? (Such an acid has the specific gravity T84.) (3) How many grammes of charcoal and how many of sulphur will burn in one litre of oxygen measured at stand- ard conditions ? What will be the weight of one litre of each of the products under the same conditions ? (4) How many grammes of potassium chlorate are re- quired to yield four litres of dry oxygen (standard condi- tions) ? (5) When dissolved in acid 1*25 gramme of zinc will yield what volume of hydrogen gas collected over water 162 LABORATORY PRACTICE. when temperature of room is 20 and barometer stands at 750 millimetres ? (6) In preparing ammonia gas from nitric oxide and hy- drogen gas in what proportions by volume should the last two be mixed ? (7) In preparing nitric oxide, how many grammes of copper will be required to each litre of gas if the product is only nitric oxide ? How many if the product is wholly nitrogen gas ? (8) A cubic decimetre of marble contains how many times its own volume of carbonic-dioxide gas ? Specific gravity of marble, 2 '75. (9) The reactions involved in these problems are all given in the list under the preceding division The student should not limit his study to the few problems here given as examples, and it must be borne in mind that the same prin- ciples apply to any chemical process, however complex or however so many simple reactions it may involve, provided only that the whole material from one reaction passes for- ward to the next. 26. Nomenclature. Before the present century the names given to chemical products were almost wholly arbitrary, and a few of these, like oil of vitriol, blue vitriol, sugar of lead, calomel, and Epsom salts, still re- main in common use. In 1787 a systematic no- menclature was devised by a committee of the French Academy of Sciences, under the lead of Lavoisier, in which the name of a substance was made to indicate its composition, and at the time of its adoption and for more than fifty years after- wards it was probably the most perfect nomen- NOMENCLATURE. 163 clature which any science ever enjoyed. It was based, however, on the dualistic theory of La- voisier, and when the science outgrew the theory the old names lost much of their significance and appropriateness. Nevertheless, the main features of the Lavoisierian nomenclature are still pre- served, although with some variations of usage as to details and the introduction of many arbitrary names, like carbinol, phenol, pinakone, de- manded by the necessities of a rapidly expanding science. The old nomenclature is so oat of har- mony with our modern conceptions that it would be impossible to explain its full significance with- out entering into details which would be out of place in an elementary course. A few of the rules should be stated by the teacher and the use of the ordinary terminations and prefixes so far ex- plained as to render the usually occurring names intelligible. All that is required may be found in any elementary text book on chemistry. CHAPTER V. MOLECULAR STRUCTURE. 27. Quantivalence. OF all chemical reactions by far the most com- mon is a claso in which, judging from the prod- ucts, the only change that takes place is an inter- change of atoms or groups of atoms between two sets of molecules, leaving all relations otherwise the same as before. Such reactions are described as metathetical, and the process is termed meta- thesis. Our chemical symbols here come to our aid by enabling us to form a clear idea of what is meant by these terms. Thus in the reaction of a solution of silver nitrate or a solution of potas- sium bromide (Ex. 68) (AgNO 3 + KBr + Aq) = (KNOs + Aq) + AgBr it is obvious that Ag changes place with the K. So also in the reaction by which hydrogen gas is made from zinc and dilute sulphuric acid Aq) + Zn = (ZnSO 4 + Aq) + H a it is equally obvious that Zn changes place with METATHESIS. 165 H 9 . Again in the ordinary test for sulphuric acid (H 2 SO 4 + Bad, + Aq) = BaS0 4 + (2 HC1 + Aq) it is evident that Ba has changed place with H,. The following experiment will further illustrate this point : Ex. 75. Metathesis. Pour ten cubic centime- tres of water into a test tube and dissolve in it one gramme of silver nitrate. Immerse in the liquid a small strip of pure copper, whose weight must be accurately determined and must not ex- ceed eighteen centigrammes. After the silver has separated, wash the powder on to a small filter and continue to pour water on to the filter until it runs through tasteless. Dry the filter on the tun- nel. Remove when dry, and after carefully wrap- ping the loose paper round the silver powder place the ball in a tared porcelain crucible and slowly heat to redness until the paper has been burned, when the silver will appear bright. When cold again weigh the crucible and find the weight of the silver, which should be to the weight of the copper approximately as Ag 2 = 216 : Cu = 63 '6. Immerse now in the blue liquid decanted from the silver a strip of zinc. The copper which had passed into solution in the previous process will now be precipitated. The reactions may be writ- ten: 166 LABORATORY PRACTICE. Cu + (2 AgNOs + Aq) = Ag 2 + (Cu(N0 3 ) 2 + Aq). Zn + (Cu(NO 8 ) a + Aq) = Cu + (Zn(NO)> + Aq). Obviously, in the first, one atom of copper re- places two atoms of silver, and in the second one atom of zinc replaces one atom of copper ; and in studying these reactions it must be remembered that the symbols correctly represent the atomic relations. Now, by studying in a similar way a very large number of metathetical reactions, it ap- pears that the atoms of hydrogen, lithium, so- dium, potassium, caesium, rubidium, silver, thali- um, chlorine, bromine, and iodine are alike in this, that while among themselves they can be ex- changed atom for atom, they are replaced by all other atoms in groups of two, three, four, or more. The atoms enumerated appear to have the small- est exchangeable value and are said to be univa- lent, while atoms which will fill the place of two univalent atoms are said to be bivalent, those that can fill the place of three trivalent, those that can fill the place of four quadrivalent, etc. So also the terms quanti valence and multivalence. The facts here stated suggest at once the conception of molecular structure, for it would seem as if the parts of a molecule must be bound together in definite relative positions in order to render such substitutions possible. This conception of struct- ure is greatly widened and strengthened when we compare together the symbols of molecules formed ATOM-FIXING POWER. 167 by the union of atoms having different degrees of quanti valence, as shown in metathetical reactions similar to those just described, for it appears that the combining power corresponds exactly to the replacing power. In making such comparisons it must constantly be borne in mind that the symbol represents in every case our knowledge of the composition and relations of the substance, and that the symbolical language thus enables us to bring before the mind in one view the results of long-continued and laborious investigation. We give below the symbols of four well-known and typical compounds of hydrogen : Hydrochloric Acid. Water. Ammonia Gas. Marsh Gas. HC1 H 2 O H 3 N H 4 C In these compounds the atoms Cl, O, N, and C are united with the same number of univalent atoms, which, under other circumstances, they might replace. So, also, we have Common Salt. Baric Chloride. hiSSd Zirconic Chloride. NaCl BaCl a SbCls ZrCl* Compare now with these the corresponding ox- ides Na 2 O BaO Sb a O 8 ZrO a , and it will be seen how we are led to the conclu- sion that in these molecules the atoms of lower quantivalence are united to those of higher quan- 168 LABORATORY PRACTICE. tivalence, which are, as it were, the nucleus of the molecule, and serve like a clamp to bind the parts together. In order to express this we often write the symbols as below, using dashes to indicate what we call the bonds. Symbols thus written are said to be graphic, while as before written they are not inappropriately spoken of as em- pirical. H H H-C1 H-O-H H-N-H H-C-H H Cl Cl Na-Cl Cl-Ba-Cl Cl-Sb-Cl Cl-Zr-Cl dl Na-O-Na Ba = O O = Sb-O-Sb = O O = Zr = O This is but a very short step in our reasoning, and yet it opens to view at once a very definite struct- ure. Notice that the only inference we have drawn is that the atoms, instead of being indis- criminately piled together, are united each sepa- rately to the multivalent atom or atoms of the group. This inference granted, we can take an- other step. When sodium acts on water we have a simple metathesis (2H-0-H + Aq) + Na a = (2Na-O-H + Aq) + H., and the product ]STa-0-H must have the same structure as H-O-H, containing only Na in place of one of the hydrogen atoms. In a similar way, GRAPHIC SYMBOLS. 169 when magnesium acts on water it must be that H-O-H + M g = MgI lH + Ha - And here it is evident that the multivalent atom Mg binds together two molecules of water, and thus we can conceive how complex molecules may be built up, and we also see that the replacing and combining powers are merely different mani- festations of the same atribute of the atoms which we express by the term " bonds"; and hence the reason that the two powers necessarily cor- respond. The products whose formation and structure we have studied are two members of a very large class of substances, all having similar relations, and which must have a similar structure. The symbols of four other members of the same class are given on the next line : -0-H K-O-H caigii zrigil at:8:l -O-H -O-H -O-H Bodies of this class are called hydrates, and the chief feature in their structure is that they con- tain atoms of hydrogen united with a multivalent nucleus through atoms of oxygen ; and, further, the one chemical relation which marks all such 170 LABORATORY PRACTICE. substances is that the atoms of hydrogen so united are easily replaced by simple metatheti- cal reactions. Ex. 76. Replacement of Hydrogen in Sodic Hydrate. In a small flask (50 cubic centimetres) dissolve 5 grammes of caustic soda in 10 cubic centimetres of water ; add a strip of aluminum weighing less than three decigrammes ; connect with pneumatic trough, and collect the hydrogen gas evolved. Measure the gas volume, observing thermometer and barometer ; correct for tension of aqueous vapor and calculate the weight of hydrogen obtained. Compare this weight with the weight of aluminum, and estimate how many atoms of hydrogen must have been replaced by each double atom of aluminum (54 microcriths) indicated by the barred symbol. If now, in this connection, the student will review the familiar experiment by which hydro- gen gas is usually made, he will see that this pro- cess is also a metathetical reaction in which hy- drogen atoms have been replaced by atoms of zinc ; and if we take such reactions as an indica- tion of the type of structure we have assigned to the hydrates, it will appear on further study that most of the active agents of chemistry, including both acids and alkalies, must be classed in the same group. To class acids and alkalies in the same group of compounds seems at first sight MOLECULAR STRUCTURE. 171 very anomalous, for in most respects their prop- erties are the direct opposites each of the other. Nevertheless, not only do these bodies resemble each other in the one essential relation of a hy- drate, but also they may be grouped in series varying so gradually from strong alkalies at one end to strong acids at the other that no natural dividing line can be found. Moreover, the dis- tinction between an acid and a base is of a rela- tive, and not of an absolute character, as is shown by the fact that in such series as have been men- tioned a given member may act as an acid towards the members at one end of the list, and as a base towards those at the other end. If now we study the symbols of the ordinary acids, and arrange the hydrogen atoms after the pattern of a hy- drate, we shall get such a result as this : Nitric Acid. Sulphuric Acid. Phosphoric Acid. TT O H -\ H-0-N0 9 S~X~ S0 3 H-O^PO H-0 X And it will be seen that though here, as before, the hydrogen atoms are united through oxygen atoms to a multivalent atom which serves as the nucleus of the molecule, this atom of high quan- tivalence acting as an atomic clamp is one of a group of atoms. Such groups are termed in chemistry compound radicals, and the chain H-O- is called hydroxyl. If the reasoning has been fol- 172 LABORATORY PRACTICE. lowed thus far the student will be prepared to admit that the relations of acids and bases which play such an important part in the science of chem- istry depend upon molecular structure. But why the opposition between these bodies ? As otherwise they have the same structure, it is evident that the antagonism must be connected with the atoms or compound radical with which the hydroxyl groups are associated ; and when we compare the nuclei of the respective molecules we find a very manifest difference between the nucleus of a marked alkali and the nucleus of a pronounced acid. In the alkali the nucleus is a metallic atom, like sodium or potassium ; in the acid it is a non- metallic atom, or a group of such atoms, like nitrogen or sulphur. Moreover, we find that while it is very easy to replace the hydrogen atoms either of an alkali or of an acid by atoms unlike the nucleus, it is difficult to replace them by atoms similar to the nucleus. Thus it is diffi- cult to change Na-O-H to Na-O-Na, but very easy to change it to Na-O-Cl or Na-0-NO 3 . The phenomena point to a polarity in the molecule similar to that of a magnet, and only by such analogies can we explain them. The objects of this discussion have been, in the first place, to give an idea of the mode of reason- ing by which our knowledge of molecular struct- ure is reached ; and, in the second place, to ex- MOLECULAR STRUCTURE. 173 plain the distinction between acids, alkalies, and salts. This distinction, however much it may be described, can not be made clear except through the molecular structure on which it is supposed to depend. Towards the first object we have been able to advance only a very few steps, but far enough to point out the way by which the com- plex structure of organic compounds has been un- ravelled and the whole subject of organic chemis- try developed. As towards the second object, we hope we have been able to make clear that acids and alkalies have in their larger relations similar qualities and a similar structure, and that they differ in the character of the nuclei to which the hydroxyl groups are united ; and, further, that when the hydrogen atoms thus united are replaced, in compounds of either class, by atoms opposite in qualities to the atoms of the nuclei, the prod- ucts are salts so called because for the most part they are bodies that can be readily crystallized. In this connection there are one or two other points to be noticed before leaving the subject. As has been shown, an acid most readily com- bines with an alkali to form a salt ; and the reason seems to be that the polarity of tb 3 molecules tends to bring atoms of opposite characters to the two ends, and determines a metathesis thus Na-O-H + H-O-NO, = Na-O-NO a + H-O-H 12 174 LABORATORY PRACTICE. The number of hydroxyl groups in a hydrate, whether acid or alkali, measures what we call its atomicity. Thus sodic hydrate is monatomic, and aluminic hydrate hexatomic. In an acid the atomicity is frequently called basicity, and just as the metallic atoms associated with the acid nuclei in a salt are often spoken of as basic radicals, so the metallic hydrates used for neutralizing the acid, as in the above reaction, are often termed bases, or the base of the salt * Thus nitric acid is monobasic, sulphuric acid is dibasic, and phos- phoric acid is tribasic. Hence, while nitric acid will form only one salt with sodium, sulphuric acid will form two, and phosphoric acid will form three. This point is illustrated by the following symbols Na-O-NO 2 H -O\ Qrk Na-O XQn Na-O /bUa Na-O /toUa H -O x H -O x Na-O. H -O-PO Na-O-PO Na-O^PO Na-O Na-O Na-O And the fact that the several salts which the sym- bols show to be possible can be prepared is in harmony with the molecular structure we have described. The chief features of the type of molecular * The term " base " is used in a broader sense than the word " alkali," and is applied to any hydrate which will unite with an acid. MOLECULAR STRUCTURE. 175 structure we have unfolded are very strikingly illustrated by the formation and decomposition of ammonium nitrate. The molecules of ammonia gas must have (as we have seen) the simple structure H H-N-H; but when the gas dissolves in water the solution acquires properties so closely resembling those of a solution of sodium hydrate that we naturally conclude that new molecules have been formed by combination with water having a structure simi- lar to Na - - H ; that is When next we neutralize ammonium hydrate with nitric acid, as in Ex. 46 (a\ the reaction must be like the reaction of the same acid on sodium nitrate given above, or H 4 N-O-H + H-O-NO, = H 4 N-O-NO a + H-O-H. Here, then, if our reasoning is correct, we have the molecule of a salt having as a nucleus N - - 1ST, but with hydrogen atoms attached to the nitro- gen atom which forms one pole of the molecule, and oxygen atoms united to the nitrogen atom which forms the opposite pole ; and that this is the true structure is indicated by the fact that when we simply heat the salt, the oxygen and hydrogen atoms, thus for a time kept apart, rush 176 LABORATORY PRACTICE. into combination, and form molecules of water, leaving the nuclei free to become the molecules of a well-known substance called nitrous-oxide gas. Ex. 77. Preparation of Nitrous Oxide. Con- nect a small flask (fifty cubic centimetres) by means of perforated corks and glass tubes, first with a test tube and second with a pneumatic trough. Place in the flask twenty -five grammes of ammonium nitrate, and mount the apparatus so that while the flask is held by a retort holder the test tube may stand in a beaker of water and the exit tube may open under the mouth of a glass jar standing full of water and inverted on the shelf of the trough. Cautiously heat the salt until it melts, and then press the heat until decomposition en- sues. Water will distil over and collect in the test tube, while nitrous oxide will bubble up and displace the water in the jar. When the jar is filled, seal it and preserve the gas for comparison. Ex. 78. Composition of Nitrous Oxide. Into a jar of nitric oxide, prepared as in Ex. 44 (a), cautiously pour 100 cubic centimetres of a con- centrated solution of green vitriol acidified with hydrochloric acid, seal the jar, and shake the solu- tion with the gas so long as absorption continues. Open now the mouth of the jar under water, and, after comparing the residual gas volume with the original volume of the nitric oxide, identify the product as the same substance which was formed MOLECULAR STKUCTUKE. 177 in the last experiment. Knowing that the symbol of nitric oxide is NO, and that the effect of the green vitriol is to withdraw oxygen, what infer- ence can you make in regard to the composition of nitrous oxide and as to the nature of the molecu- lar change which has taken place in this experi- ment. (1) The molecular structure of ammonium nitrate thus developed may serve to give some conception of the condi- tions to which modern explosives, like nitre-glycerin and gun cotton, owe their remarkable relations. In the mole- cules of these explosives it is supposed that three or more of the nuclei N - O - N are bound together by groups of hydrogen and carbon atoms at one end of a multiple chain, while oxy- gen atoms, in sufficient numbers to unite with all the car- bon and hydrogen, are attached at the other end. By this structure the intensely powerful affinities between the great fire element and the combustibles are for a time held in abeyance. But when the equilibrium is disturbed the atoms thus held apart rush together and a great volume of aeriform products are suddenly developed whose enormous expansive force produces the destructive effects so well known. CHAPTER VI. THERMAL RELATIONS. 28. Heat of Chemical Action. EVEN the most elementary course, proposing to treat only of the fundamental principles of chemistry, would be incomplete without some dis- cussion of the thermal relations of chemical changes, and the most striking chemical experi- ments will have to the student no more meaning than fireworks if they remain in his mind as mere- ly brilliant phenomena. After what must be here assumed to have been already learned, the student should be prepared to understand the treatment of this subject in the last chapters of the author's New Chemistry, and more fully in the chapter on the " Thermal Relations of Atoms" in his Chemi- cal Philosophy. As an introduction to the sub- ject, the student should try the following experi- ments, using for the purpose the calorimeter, already fully described (Ex. 8), but when corro- sive liquids are used substituting for the inner brass vessel as thin a beaker glass as can be had, of about the same capacity, and filling the space HEAT OP SOLUTION. 179 between the glass and the sides of the chamber with layers of wool wadding, caught together so that the beaker can readily be removed and re- placed. For accurate experiments a dish made of thin platinum plate is always to be preferred. In using a beaker the heat absorbed by the glass becomes a quantity of importance. The glass must therefore be weighed ; and the weight of the glass multiplied by the specific heat of glass (Ex. 8 (4)) gives a value which is called the thermal water equivalent, and this in every experiment is to be added to the weight of the water. Ex. 79. Heat of Hydration and of Solution. Place in the calorimeter about 300 grammes of water, weighing the amount accurately to a gramme. Prepare and pulverize 35 grammes of anhydrous sodic sulphate by driving off the water from the crystallized salt (Glauber's salts). Keep the salt between watch glasses over the beaker of water and under the cover of the calorimeter until a perfect equilibrium of temperature is reached. Then stir the salt into the water with a glass rod, and observe the rise of temperature. Calculate the number of units of heat evolved by 142 grammes of the salt.* This number is the molecu- * In making this and similar calculations it must be remem- bered that is the whole mass of the solution, and not the water merely, whose temperature is changed. To obtain strictly accurate 180 LABORATORY PRACTICE. lar weight in gramme units ; and it is convenient to state results on this basis, as we can then carry our calculations through successive reac- tions without constant reduction. Repeat now the same experiment, but use 79*3 grammes Glau- ber's salts (Na 2 S04 . 10H a O) in fine crystals. Cal- culate as before for one molecule in grammes, and compare the two results. How much heat is lib- erated in the union of JSTa a SO 4 with 10H a O ? Ex. 80. Heat of Neutralization. Weigh in a glass-stoppered vial about twenty -five grammes of strong sulphuric acid, the specific gravity of which has been previously accurately ascertained. Mix this acid with about 250 cubic centimetres of water and give time to cool before pouring into the cal- orimeter. Finding from the tables the amount of H a SO 4 thus taken, calculate the amount of Na-0- H required to neutralize the acid and weigh out about one fifth more than the calculated amount in order to insure an excess. Dissolve the alkali also in about 250 cubic centimetres of water and place the vessels holding the two solutions under cover in a protected place at the side of the cal- results, we should know the specific heat of the solution ; but for all practical purposes it is sufficiently accurate to count one cubic cen- timeter of the solution, measured at 4 C.. as the thermal equiva- lent of one gramme of water. The simplest way is, after the close of the experiment, to weigh the solution and determine its specific gravity with a delicate spindle hydrometer graduated at 4 C. Then the weight, divided by the specific gravity, is the thermal water equivalent in grammes. HEAT OF CHEMICAL ACTION. 181 orimeter. When both solutions have cooled to the same temperature, pour them together into the calorimeter and observe the rise of tem- perature. Calculate for one molecule H a SO 4 in grammes. Ex. 81. Heat of Chemical Action. Mix 250 cubic centimetres of water with fifty cubic centi- metres of strong sulphuric acid, and when the mixture is cold place it in the calorimeter. Scrupu- lously clean a strip of sheet zinc about two inches wide by five inches long. Accurately weigh the zinc. Plunge the strip into the acid and allow the action to continue until the temperature has risen three or four degrees. Then remove the metal, note the rise of temperature, and after washing strip with water and alcohol, dry, and de- termine the loss of weight. Calculate the amount of heat evolved for each molecule in grammes of ZnS0 4 formed. Ex. 82. Heat of Precipitation. Dissolve a weighed amount about twenty grammes of crys- tallized baric chloride (BaCl a . 2 H 3 O) in 250 cubic centimetres of water. Calculate the quantity of sulphuric acid of known specific gravity required to decompose the salt, and mix the acid in slight excess of the calculated amount with 250 cubic centimetres of water. Handle the solutions as in Ex. 80, pouring first the acid and afterwards the salt solution, slowly and with constant stirring, 182 LABORATORY PRACTICE. into the calorimeter. Calculate for each molecule of BaS0 4 (in gramme units) formed. Ex. 83. Crystallized Cupric Sulphate. Dis- solve a weighed amount about twenty grammes of blue vitriol in about 250 cubic centimetres of water and place the solution in the calorimeter. Plunge in the solution a strip of zinc prepared as in Ex. 81 and stir until the copper is wholly pre- cipitated, and then note the rise of temperature. Calculate for every 63*6 or every atom in grammes of copper reduced. Unfortunately the fundamental experiments on this subject such, for example, as those on the heats of combustion of hydrogen, carbon, and sulphur, or the heat of formation of hydrochloric acid are out of the reach of elementary students, and indeed of most teachers, but the general prin- ciples involved can be readily made clear. The chief points to be insisted on are : (1) It follows from the principle of conservation of en- ergy, and has been fully proved by investigation, that in a series of chemical changes the total amount of heat devel- oped depends wholly on the initial and final states of the system, and is not dependent on the intermediate steps. Thus ninety-six grammes of sulphuric acid consists of thirty- two grammes of sulphur, sixty-four grammes of oxygen, and two grammes of hydrogen, and can be prepared by com- bining these relative amounts of roll brimstone, oxygen gas, and hydrogen gas in several different ways. But whatever may be the series of processes employed, the total amount of heat evolved in the production of ninety -eight grammes of this definite compound from the several elementary sub- EXOTHERMOUS AND ENDOTHERMOUS. 183 stances will be 193,100 units. Hence, conversely, if by a se- ries of analytical processes we resolve back this compound into the same elementary substances in the same condition an equal amount of heat will be absorbed. It is only excep- tionally the case that we can prepare compounds by the direct union of elementary substances, and even when we can it is rarely that we can measure the heat thus devel- oped ; but the above principle renders this unnecessary. We can always determine the heat evolved in the production of a compound from elementary substances (or conversely) by measuring the heat evolved at each step of the successive operations by which it may be made, and we are thus able to choose such processes as are adapted to thermal measure- ments ; and in the investigations of thermo - chemistry a great deal of ingenuity has been shown in this selection or in devising new processes which are compatible with the methods of calorimetry. In this manner what is termed the heat of formation of a large part of known compounds has been measured. In making our calculations and in stating results we adopt the system already referred to (Ex. 79), and a table giving the more important data will be found in the author's work on Chemical Philos- ophy. The thermal relations have led to the division of com- pounds into two large classes exothermous bodies (by far the larger class), whose formation from known elementary substances in their familiar state is attended with the evo- lution of heat, and endothermous compounds, of which the reverse is true. It also follows from the principle we have been consid- ering that if we begin with the same substance in the same state and by different processes reach two different products, the difference in the heat evolved in the two cases is that required to pass from one product to the other, or, what is an obvious corollary, if the initial states are different and the final results in all respects the same, then the same relation will hold between the initial states. This deduction gives us a very simple means of determining the heat of combina- tion in a great number of cases where direct union is impos- 184 LABORATORY PRACTICE. sible or where the action is so violent that all thermal meas- urements are impracticable. Thus 28 grammes of olefiant gas, C 2 H 4 , contain 24 grammes of carbon and 4 grammes of hydrogen, and, al- though we can not combine directly charcoal and hydrogen gas, we can determine the heat evolved in the production of the compound in this way. If we burn 28 grammes of olefi- ant gas the products will be 88 grammes of carbonic diox- ide and 36 grammes of water 28 88 36 C 2 H 4 + 30 2 == 2CO a + 2H 2 = 332,024. If we burn 24 grammes of charcoal and 4 grammes of hydro- gen separately we shall obtain the same weights of the same products in the same condition 24 88 20 + 20 3 = 2CO 2 = 193,920 4 36 2H 2 + 2 = 2H 2 = 137,848 331,768 The heat evolved in all these three processes of combus- tion has been measured and is given after the reaction. The total heat evolved in burning the elementary substances is less than that set free in burning the compound by 256 units. Hence in passing from charcoal and hydrogen gas to olefi- ant gas this small amount of heat must have been absorbed. Olefiant gas is therefore an endothermous compound. Again, sulphuric oxide and water when united in the proportions indicated by the reaction S0 3 + H 2 O = H.S04 combine with explosive violence, but we can readily dis- solve both SOs and H 2 SC>4 in an equally large volume of water and determine the heat evolved in each case. The final result in both cases is a weak solution of sulphuric acid, and the difference between the amounts of heat evolved gives the amount which would be given by the above reac- tion if it could bo measured. TENDENCY OF CHEMICAL PROCESSES. 185 (2) "With a table giving the heats of formation of the more important compounds in different conditions (whether solid, liquid, aeriform, or in solution in water) we are in po- sition to calculate the heat evolved in any ordinary chemi- cal process. We have only to compare the sum of the heats of formation of the factors of the reaction with that of the products, paying careful regard to the conditions in which the several materials are present. Thus, in the reaction we have discussed so often (H 2 SO4 + Aq) + Zn = (ZnS04 + Aq) + H 2 , the heat evolved during the process is the difference between the heat of formation of sulphuric acid in aqueous solution and that of zinc sulphate dissolved in an equal amount of water. We need pay no regard to the elementary sub- stances, either the zinc dissolved or the hydrogen gas set free, for they are present in the very condition which our calculations assume, and since the heat of formation of dilute sulphuric acid is 210,000 and that of the solution of zinc sulphate 252,000 there must be set free in the reaction 252,000 - 210,000 = 42,000 units. (3) It has been inferred as a generalization from a great number of facts that, other things being equal, the activity of a chemical process is proportional to the amount of heat evolved, and that where several courses are possible the tend ency is always to form those products which involve the greatest evolution of heat. Often by restraining the reac- tion (as by lowering the temperature, diluting the solution, or restricting the amount of material) other products may result ; but if we give the chemical action full play the tend- ency is as above stated. In the action of nitric acid on cop- per there may be formed either nitric oxide or nitrogen gas, thus : 3Cu + (SHNOs + Aq) = (3Cu(NO 3 ) a + 4H 2 O + Aq) + 2NO, or 5Cu + (12HNO, + Aq) = (5Cu(NO 3 ) a + 6H 3 O + Aq) -f N a . 186 LABORATORY PRACTICE. The last develops the most heat, and nitrogen gas is the chief or sole product formed if the materials are allowed to become heated ; but if the flask is kept cool the chief or only product is nitric oxide, as in Ex. 44 (a). If no heat would be set free by an assumed process the reaction can not take place without some aid. There ap- pears no reason in the form of the reaction why copper should not act on dilute sulphuric acid like zinc, that is Cu + (H 2 SO4 + Aq) yield (CuS04 + Aq) + H 3 But while the heat of formation of an aqueous solution of sulphuric acid is 210,000, as above, that of an aqueous solu- tion of copper sulphate is 199,100 (instead of 252,000, as in the case of zinc sulphate), and heat would be absorbed, not evolved, by the chemical change. The aid required to determine a reaction in such cases may be furnished either by external energy, as the sun's rays acting on the green foliage of the vegetable kingdom or by some simultaneous exothermous process which en- trains the other and supplies the necessary heat. Thus, in the above reaction, if a small amount of nitric acid is added (as shown in Ex. 54 (&)) the copper at once dissolves simply because the nitric acid, by oxidizing the hydrogen evolved (compare Ex. 43 (6)), generates the heat required to render the process, as a whole, exothermous. Thus it is that the formation of endothermous com- pounds becomes possible. Nitrous oxide, N 2 O, is such a com- pound. In its production from nitrogen and oxygen gases 18,000 units of heat are absorbed. Its tendency, therefore (in itself alone), is to fall back into the constituent gases, when the same amount of heat would be evolved. In the reaction by which nitrous oxide is made NH 4 NO 3 = 2H,O + N,O 80,700 118,800 - 18,000, the heats of formation are printed under the symbol, and it will be seen that, as a whole, the process develops 100,800 units of heat. But, as will also be noticed, this heat wholly comes from the oxidation of hydrogen to form water, which CAUSE OF INSTABILITY. 187 is sufficient to furnish all that is required for the production of N 2 O and still have a large excess over what is required to determine the reaction. Indeed, unless restrained (by keeping the temperature at the lowest possible point), this re- action will take the form. NHNOs = 2H 2 O + N 3 + iO., which corresponds to a larger evolution of heat, and to just as much more as was used above in the production of N 2 O. (4) Endothermous compounds are always in a condition of unstable equilibrium, and sometimes highly explosive. This is strikingly true of iodide of nitrogen, which often ex- plodes at the mere touch of a feather and is resolved wholly into elementary substances 2NI 8 = N 2 + 3I a . They may endure, often for a long time, in consequence probably of features of molecular structure such as we en- deavored to illustrate in the case of ammonium nitrate, but sooner or later they fall into more stable conditions. They may be compared to a vaulted cathedral roof of which the stones are firmly locked together and held high in air by buttresses, but when keystone or buttress fail fall in ruin. There are many substances which, although exother- mous to the elementary substances to which their heat of formation is referred, are endothermous in their relations to certain definite products into which .they are more or less readily resolved with evolution of heat, or in their rela- tions to associated material, from uniting with which they are restrained by physical disabilities or conditions of struct- ure. Nitro-glycerin and gun-cotton, already referred to, are examples of the first type, while gunpowder, in which com- bustible charcoal and sulphur is kept apart from the great store of oxygen in the grains of nitre by the inertness of the solid state, is an equally striking example of the second type. All these explosive agents owe their efficiency not only to the heat evolved by the internal combustion, which ensues when they are fired, but also to the circumstance that the products into which they fall are for the most part aeriform 188 LABORATORY PRACTICE. bodies whose molecules acquire an enormous moving power under the influence of the heat thus generated. Of instability arising from association by far the most wonderful example is furnished by the presence on the sur- face of the globe of a large amount of combustible material in contact with the oxygen of the atmosphere. Almost the whole of this material is made up either of the organized structure of plants and animals or else of the remains of such structures, and however multifarious the substances of which this organic matter may consist, the ultimate ele- ments, with unimportant exceptions, are carbon, hydrogen, nitrogen, and oxygen, and the whole of this material was primarily formed from the constituents of air and water, including, of course, carbonic acid and ammonia, always present in the atmosphere and in the water permeating the soil. From these materials the plant obtains almost its sole food, and the animal ultimately at least lives on the plant. In some mysterious way the sun's rays give the plant the power of producing the substances of its tissues from the simple articles of its diet. Of the manner in which the highly complex products are built up we know almost noth- ing. But of this we are sure. The energy of the sun's rays is the power by which carbonic acid and water are decom- posed and materials so unstable in the presence of the atmos- phere constructed. Probably the effect is indeed a result of molecular construction and the structure endures until a conflagration, or the slower processes of decay, destroys the fabric and resolves the organic matter into the elements from which it sprung. Thus there is a constant cycle in nature. The sun's rays are ever building up, and in so do- ing are setting free the very oxygen which, sooner or later, will destroy all this work. We also know that the heat given out in burning is the exact equivalent of the work done in building, and therefore is simply transmuted solar energy. Just as the sun lifts the water whose fall main- tains the great aqueous circulation of the globe, so the same vitalizing energy builds up organic structures, by whose reabsorption into the all-devouring atmosphere the life and activity on the planet is sustained. Moreover, only by util- EXPLOSIVES. 189 izing this same energy though often so indirectly that it es- capes notice are we able to produce endothermous or un- stable compounds in our laboratories. (5) It must be remembered that compounds are endoth- ermous only in relation to the elementary substances in a definite condition (carbon as charcoal sulphur, as brimstone, etc.), from which in our system they are regarded as having been formed. Their peculiar thermal relations do not neces- sarily imply a want of chemical energy between the element- ary atoms of which their molecules are supposed to consist. Thus the oxides of nitrogen are endothermous, and yet there can be no question that the nitrogen atoms have a marked affinity for the atoms of oxygen. If we could deal with ele- mentary atoms all compounds would be, doubtless, exother- mous, but when we deal with the elementary substances in their normal condition the constitution of the molecules of these elementary substances comes into play. As has been shown, the molecules of elementary substances, as well as those of compounds, are aggregates of atoms, only of atoms all of which are of the same kind, and not, as in compound substances, of different kinds. The molecule of oxygen gas (O 2 ) is formed of two atoms of oxygen ; that of ozone (O 3 ), of three atoms of oxygen ; that of nitrogen gas (N 9 ), of two atoms of nitrogen, etc. The atoms of nitrogen, united in a molecule of the gas must have an attraction for each other, else they would not so group themselves in pairs. As yet we have not been able to measure this attraction with confi- dence ; but we have good reason for believing that it is very strong. The endothermous compounds called nitrogen iodide (NI 3 ) and nitrogen chloride (NCL) are so explosive not, as we believe, because the nitrogen atoms have no affin- ity for the iodine and chlorine atoms, but because they have such a strong attraction for each other that they break from the iodine and chlorine atoms and rush together. 2NL = N a + 3L 2NC1 S = N 2 + 3C1 3 We explain the inertness of nitrogen gas in this way. Simply the nitrogen atoms exert a stronger attraction 13 190 LABORATORY PRACTICE. among- themselves than for those of the other chemical ele- ments with only a few exceptions. Hence, also, the general instability of the compounds of nitrogen as a class. Animal structures are for the most part made up of nitrogenized substance, and thus decay and death in nature are closely associated with this striking feature of our chemical phi- losophy. (6) Could we experiment with isolated atoms all chemi- cal relations would unquestionably appear simpler, and modern science has rendered probable that there exists such a condition in the universe. It is well known that heat tends to decompose chemical compounds, and the phe- nomena thus resulting form a very interesting subject of chemical inquiry known as thermolysis or dissociation. Steam passed through metal tubes at a white heat acts in every respect like a mixture of oxygen and hydrogen gases, and, as common experience shows, most compounds even at a red heat suffer more or less fundamental chemical changes. The known facts point to the conclusion that at such high temperatures as must rule at the sun and at the fixed stars all known materials would be resolved into ele- mentary atoms. Spectroscopic observations confirm this inference, and such evidence has even been interpreted as indicating that some of the atoms which we regard as ele- mentary are resolved into still simpler parts at the great focus of solar radiation. If the nebular hypothesis is cor- rect our world must have been primarily in this condition, and the substances which we now find on its crust must have been formed as the elementary atoms came together in the process of cooling and united in accordance with their mutual affinities. According to this hypothesis, the original chaos out of which the present order sprang was a condition of isolated atoms, and the foundations of the globe must have been laid in flames. (7) In this last chapter of our book we have only been able to illustrate a few of the more important relations of thermo-chemistry. Our one object has been to exhibit the scope of this department of chemical science, as we have sought to show in earlier chapters that of qualitative and CONCLUSIONS. 191 of quantitative analysis. The theoretical relations of the science have been previously set forth by us in popular form in the New Chemistry.* It is expected that this will serve as a companion to the former book, and the student who thoughtfully performs all the experiments and dili- gently inquires what each is calculated to teach can not fail to gain clear ideas of the methods of chemical investigation and at the same time will acquire skill in drawing inferences from experimental data. Having thus seen the relations of the broader divisions of the subject, the student will be pre- pared to enter on the professional study of chemistry intelli- gently, or, if he goes no further, will have acquired a clear conception of the aims and methods of the science and of its true position in a scheme of education. The next step in the study of chemistry should be to acquire an adequate knowledge of the scheme of the chemical elements as it is presented by Roscoe and Schorlemmer,f or, still better, as it is illustrated experimentally in an extended course of lect- ures such as is given every year at Harvard or at any one of our principal universities. This is a serious task, since the mass of details is very great and the subject can not be profitably abridged beyond a limited extent. A brief epit- ome will not be of much value. The mind must dwell on the subject in its various relations in order to make the knowledge real or lasting, and unless the student has that object in the acquisition which will lead him to give to the work the requisite time, he had better, especially if it is a question of liberal education, limit his study to the general principles of the science as presented in this or some similar book. * D. Appleton & Co., publishers. New York. Last edition, 1890. f Treatise on Chemistry, D. Appleton & Co. 192 LABORATORY PRACTICE. List of the Elementary Substances, excepting such very Rare or of Doubtful Authenticity. as are Aluminum Al, 27-1 Molybdenum, Mo 96 Antimony Sb 120 Nickel, Ni 59 Arsenic As 75 Nitrogen. N 14 Barium Ba 137 Osmium, Os, 190-9 Bismuth Bi 208 Oxveren, o 16 Boron B 11 Palladium Pd 106*6 Brom.in6 Br 80 Phosphorus p 31 Cadmium Cd 112-2 Platinum, Pt, 194-8 Caesium Cs 133 Potassium K 39'1 Calcium, Carbon Ca, c 40 12 Rhodium, Rubidium, Rh, .... Eb . 103 85 "4 Cerium Ce 140 Rutheniu m Ru 101*8 Chlorine Cl 35'5 Scandium J-VIA, .... Sc 44 Chromium Cr 52'1 Selenium, k->v^, * Se 79-2 Cobalt Co 59 Silicon Si 28'3 Columbium. Cb 94 Silver As" 108 Copper, Didymium, I Erbium, V^kf, .... Cu fd, 141 Pr, Er, 63-6 144 166 Sodium, Strontium, Sulphur, J - x &5 ' Na, .... Sr, .... S, .... 23 87-6 32-1 Fluorine F 19 Tantalum, Ta, .... 182 Gallium Ga 70 Tellurium Te, 125? Germanium, Ge, 72-3 Terbium ? Tr, .... 171? Glucinum Gl 9-1 Thallium, Tl, 204-1 Gold Au . ... 197'2 Thorium, Th, 232 Hvdrocren H 1 Thulium ? Tm . 171 Indium i-*., In 113-7 Tin, Sn, 118 Iodine I 126-9 Titanium Ti 48 Indium Ir 193 Tungsten W 184 Iron Fe 56 Uranium Ur 240 Lanthanum, Lead La, Pb 139 206-9 Vanadium, Ytterbium ? Va, .... Yb, . 51-4 173 Lithium Li 7 Yttrium, Y 90 IVIajrnesium Mff 24 '4 Zinc Zn, 65-2 Manganese, Mercurv. Mn, . ... Hff. . 55 200 Zirconium, Zr, .... 90 THE EXD. KEROSENE STOVE AND TUBE FURNACE. Well adapted for most chemical experiments referred to in this book (Ex. 28). D, APPLE TON & CO, '8 PUBLICATIONS, JOHN TYNDALL'S WORKS. ESSAYS ON THE FLOATING MATTER OF THE AIR, in Relation to Putrefaction and Infection. 12mo. Cloth, $1.50. ON FORMS OF WATER, in Clouds, Rivers, Ice, and Glaciers. With 35 Illustrations. 12mo. Cloth, $1.50. HEAT AS A MODE OF MOTION. New edition. 12mo. Cloth, $2.50. ON SOUND : A Course of Eight Lectures delivered at the Royal Institution of Great Britain. Illustrated. 12rao. 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