PROBLEMS.

(1) A certain volume of H weighs 0.36 g. at standard temperature and pressure. How many liters does it contain? If one liter weighs 0.09 g., to weigh 0.36 g. it will take 0.36 / 0.09 = 4 liters.

(2) How many liters, or criths, of H in 63 g.? 2.7 g.? 1 g.? 5 g.? 250 g.? Explain each.

(3) Suppose the gas to be twice as heavy as H, how many liters in 0.36 g.? A liter of the gas will weigh 0.18 g. (0.09 X 2). In 0.36 g. there will be 0.36 / 0.18 = 2. Answer the question for 63 g., 2.7 g., etc.

(4) How many liters of Cl in each of the above numbers of grams?

(5) How many of HCl? H2O (steam)? CO2? Explain fully every case.

Vapor density is very easily determined from the formula by the method given above. But in practice the formula is obtained from the vapor density, and hence the method there given has to be reversed.

173. Vapor Density of Oxygen.—Suppose we were to obtain the vapor density of O. We should carefully seal and weigh a given volume, say a liter, at a noted temperature and barometric pressure, which are reducedto 0 degrees and 760 mm, and compare it with the weight of the same volume of H. This has been done repeatedly, and O has been found to weigh 16 times as much as H, volume for volume, or, more exactly, 15.96+. Now a liter of each gas has the same number of molecules, therefore the O molecule weighs 16 times the H molecule. The half-molecule of each has the same proportion, and the vapor density of O is 16. Atomic weight is obtained in a very different way.

PROBLEMS.

(1) A liter of Cl is found to weigh 3.195 g. Compute its vapor density, and explain fully.

(2) A liter of Hg vapor, under standard conditions, weighs 9 g. Find its vapor density, and explain.

The vapor density of only a few elements has been satisfactorily determined. See page 12. Some cannot be vaporized; others can be, but only under conditions which prevent weighing them. The vapor density of very many compounds also is unknown.

(3) A liter of CO2 weighs 1.98 g. Find the vapor density, and from that the molecular weight, remembering that the latter is twice the former. See whether it corresponds to that obtained from the formula, CO2. This is,in fact, the way a formula is ascertained, if the atomic weights of its elements are known.

(4) A liter of a compound gas weighs 2.88 g. Analysis shows that its weight is half S and half O. As the atomic weight of S is 32, and that of O is 16, what is the symbol for the gas?

Solution. Its molecular weight is 64, i.e. (2.88=0.09) X 2, of which 32 is S and 32 O. The atomic weight of S is 32, hence there is one atom of S, while of O there are two atoms. The formula is SO2.

(5) A liter of a compound gas, which is found to contain 1 C and 3 O by weight, weighs 1.26 g. What is its formula? Atomic weights are taken from page 12. Prove your answer.

(6) A liter of a compound of N and O weighs 1.98 g. The N is 7/11; and the O 4/11. What is the gas?

(7) A compound of N and H gas weighs 0.765 g. to the liter. The N is 14/17 of the whole, the H 3/17. What gas is it? CHAPTER XXXV.

ATOMIC WEIGHT.

174. Definition.—We have seen that the molecular weight of a compound, as well as of most elements, is obtained from the vapor density by doubling the latter. It remains to explain how atomic weights are obtained. The term is rather misleading. The atomic weight of an element is its least combining weight, the smallest portion that enters into chemical union, which is, of course, the weight of an atom.

175. Atomic Weight of Oxygen.—Suppose we wish to find the atomic weight of oxygen. We must find the smallest proportion by weight in which it occurs in any compound. This can only be done by analyzing all the compounds of O that can be vaporized. As illustrative of these compounds take the six following:—

Wt. of other Names. V. d. Mol. Wt. Wt. of O. Elem. Symbol. Carbon monoxide... 14 28 16 12 ? Carbon dioxide.... 22 44 32 12 ? Hydrogen monoxide... 9 18 16 2 ? Nitrogen monoxide... 22 44 16 28 ? Nitrogen trioxide... 38 76 48 28 ? Nitrogen pentoxide... 54 108 80 28 ?

176. Molecular Symbols.—From the vapor density of the gases— column 2—we obtain their molecular weight— column 3. To find the proportion of O, it must be separated by chemical means from its compounds and separately weighed. These relative weights are given in column 4. Now the smallest weight of O which unites in any case is its atomic weight. If any compound of O should in future be found in which its combining weight is 8 or 4, that would be called its atomic weight. By dividing the numbers in column 4, wt. of O, by 16, the atomic weight of O, we obtain the number of O atoms in the molecule. Subtracting the weights of O from the molecular weights, we have the parts of the other elements, column 5, and dividing these by the atomic weight of the respective elements, we have the number of atoms of those elements, these last, combined with the number of O atoms, give the symbol. In this way complete the last column.

Show how to get the atomic weight of Cl from these compounds, arranging them in tabular form, and completing as above: HCl, KCl, NaCl, ZnCl2, MgCl2; the atomic weight of N in these: N2O, NO, NH3.

177. Molecular and Atomic Volumes.—We thus see that vapor density and atomic weight are obtained in two quite different ways. In the case of elements the two are usually identical, i.e. with the few whose vapor density is known; but this is not always true, and it leads to interesting conclusions regarding atomic volume. In O both vapor density and atomic weight are 16. This gives 2 atoms of O to the molecule, i.e. the molecular weight / the atomic weight. The size of an O atom is therefore half the gaseous molecule, and is represented by one square. S has a vapor density and an atomic weight of 32 each. Compute the number of atoms in the molecule. Compute for I, in which the two are identical, 127. P has an atomic weight of 31, while its vapor density is 62. Its molecule must consist of 4 atoms, each half the size of the H atom, The vapor density of As is 150, the atomic weight 75. Compute the number of atoms in its molecule, and represent their relative size. Hg has an atomic weight of 200, a vapor density of 100. Compute as before, and compare the results with those on page 12. Ozone has an atomic weight of 16, a vapor density 24. Compute.

Chapter XXXVI.

DIFFUSION AND CONDENSATION OF GASES.

178. Diffusion of Gases.—Oxygen is 16 times as heavy as H. If the two gases were mixed, without combining, in a confined space, it might be supposed that O would settle to the bottom and H rise to the top. This would, in fact, take place at first, but only for an instant, for all gases tend to diffuse or become intimately mixed. The lighter the gas the more quickly it diffuses.

179. Law of Diffusion of Gases.—The diffusibility of gases varies inversely as the square roots of their vapor densities. Compare the diffusibility of H with that of O. dif. H:dif. O:: sqrt(16): sqrt(1), or dif: H: dif. O:: 4: 1.

That is to say, if H and O be set free from separate receivers in a room, the H will become intermingled with the atmosphere four times as quickly as the O. Compare the diffusibility of O and N; of Cl and H. Take the atomic weights of these, since they are the same as the vapor densities. In case of a compound gas, half the molecular weight must be taken for the vapor density; e.g. dif. N20: dif. O.:: sqrt(16): sqrt(22).

180. Cause.—Diffusion is due to molecular motion; the lighter the gas the more rapid the vibration of its molecules. Compare the diffusibility of CO2 and that of Cl; of HCl and SO2; of HF and I.

181. Liquefaction and Solidification of Gases.—Water boils at 100 degrees, under standard pressure, though evaporating at all temperatures; it vaporizes at a lower point if the pressure be less, as on a mountain, and at a higher temperature if the pressure be greater, as at points below the sea level. Alcohol boils at 78 degrees, standard pressure, and every liquid has a point of temperature and pressure above which it must pass into the gaseous state. Likewise every gas has a critical temperature above which it cannot be liquefied at any pressure.

This condition was not recognized formerly, and before 1877, O, H, N, C4, CO, NO, etc., had not been liquefied, though put under a pressure of more than 2,000 atmospheres. They were called permanent gases. In 1877 Cailletet and Pictet liquefied and solidified these and others. The lowest temperature, about -225 degrees, was produced by suddenly releasing the pressure from solid N to 4mm, which caused it rapidly to evaporate. Evaporation, especially under diminished pressure, always lowers the temperature by withdrawing heat.

These low degrees are indicated by a H thermometer, or if too low for that, by a "thermo-electric couple" of copper and German silver.

The pupil can easily liquefy SO, by passing it through a U-tube which is surrounded by a mixture of ice and salt in a large receiver. At the meeting of the American Association for the Advancement of Science in 1887, a solid brick of CO2 was seen and handled by the members, Liquid H is steel blue.

A few results obtained under a pressure of one atmosphere are:— Boiling Points: C2H4—102 degrees; CH4—184 degrees; O—181 degrees; N —194 degrees; CO—190 degrees; NO—154 degrees; Air— 191 degrees.

Solidifying Points: Cl -102 degrees; HCl -115 degrees; Ether -129 degrees; Alcohol -130 degrees.

Chapter XXXVII.

SULPHUR.

Examine brimstone, flowers of sulphur, pyrite, chalcopyrite, sphalerite, galenite, gypsum, barite.

182. Separation.

Experiment 103.—To a solution of 2 g. of sodium sulphide,, Na2S2 in 10 cc. H2O add 3 or 4cc. HCl, and look for a ppt. Filter, and examine the residue. It is lac sulphur, or milk of sulphur.

183. Crystals from Fusion.

Experiment 104.—In a beaker of 25 or 50 cc. capacity put 20 g. brimstone. Place this over a flame with asbestos paper interposed, and melt it slowly. Note the color of the liquid, then let it cool, watching for crystals. When partly solidified pour the liquid portion into an evapo- rating-dish of water, and observe the crystals of S forming in the beaker (Fig. 42). The hard mass may be separated from the glass by a little HNO3 and a thin knife-blade, or by CS2.

184. Allotropy.

Experiment 105.—Place in a t.t. 15g of brimstone, then heat slowly till it melts. Notice the thin amber-colored liquid. The temperature is now a little above 100 degrees. As the heat increases, notice that it grows darker till it becomes black and so viscid that it cannot be poured out. It is now above 200 degrees. Still heat, and observe that it changes to a slightly lighter color, and is again a thin liquid. At this time it is above 300 degrees. Now pour a little into an evaporating dish containing water. Examine this, noticing that it can be stretched like rubber. Leave it in the water till it becomes hard. Continue heating thebrimstone in the t.t. till it boils at about 450 degrees, and note the color of the escaping vapor. Just above this point it takes fire. Cool the t.t., holding it in the light meantime, and look for a sublimate of S on the sides.

185. Solution.

Experiment 106.—Place in an evaporating-dish a gram of powdered brimstone, and add 5cc, CS2, carbon disulphide. Stir, and see whether S is dissolved. Put this in a draft of air, and note the evaporation of the liquid CS2, and the deposit of S crystals. These crystals are different in form from those resulting from cooling from fusion.

186. Theory of Allotropy.—The last three experiments well illustrate allotropy. We found S to crystallize in two different ways. Substances can crystallize in seven different systems, and usually a given substance is found in one of these systems only; e.g. galena is invariably cubical. An element having two such forms is said to be dimorphous. If it crystallizes in three systems, it is trimorphous. A crystal has a definite arrangement of its molecules. If without crystalline form, a substance is called amorphous. An illustration of amorphism was S after it had been poured into water. Thus S has at least three allotropic forms, and the gradations between these probably represent others. Allotropy seems to be due to varied molecular structure. We know but little of the molecular condition of solids and liquids, since we have no law to guide us like Avogadro's in gases; but, from the density of S vapor at different temperatures, we infer that liquids and solids have their molecules very differently made up from those of gases. The least combining weight of S is 32. Its vapor density at 1,000 degrees is 32; hence its molecular weight is 64, i.e. vapor density x 2; and there are 2 atoms in its molecule at that temperature, molecular weight / atomic weight. At 500 degrees, however, the vapor density is 96and the molecular weight 192. At this degree the molecule must contain 6 atoms. How many it has in the allotropic forms, as a solid, is beyond our knowledge; but it seems quite likely that allotropy is due to some change of molecular structure.

The above experiments show two modes of obtaining crystals, by fusion and by solution.

187. Occurrence and Purification.—Sulphur occurs both free and combined, and is a very common element. It is found free in all volcanic regions, but Sicily furnishes most of it. Great quantities are thrown up from the interior of the earth during an eruption. The heat of volcanic action probably separates it from its compound, which may be CaSO4. Vast quantities of the poisonous SO2 gas are also liberated during an eruption, this being, in volume of gases evolved, next to H2O. S is crudely separated from its earthy impurities in Sicily by piling it into heaps, covering to prevent access of air, and igniting, when some of the S burns, and the rest melts and is collected. After removal from the island it is further purified by distilling in retorts connected with large chambers where it sublimes on the sides as flowers of sulphur (Fig. 43). This is melted and run into molds, forming roll brimstone. S also occurs as a constituent of animal and vegetable compounds, as in mustard, hair, eggs, etc. The tarnishing of silver spoons by eggs is due to the formation of silver sulphide, Ag2S. The yellow color of eggs, however, is due to oils, not to S.

The main compounds of S are sulphides and sulphates. What acids do they respectively represent? Metallic sulphides are as common as oxides; e.g. FeS2, or pyrite, PbS, or galenite, ZnS, or sphalerite, CuFeS2, or chalcopyrite, etc. The most abundant sulphate is CaSO4, or gypsum. BaSO4, or barite, and Na2SO4, or Glauber's salt, are others.

The only one of these compounds that is utilized for its S is FeS2. In Europe this furnishes a great deal of the S for H2SO4. S is obtained by roasting FeS2. 3 FeS2 = Fe3S4 + 2 S.

188. Uses. -The greatest use of S is in the manufacture of H2SO4. A great deal is used in making gunpowder, matches, vulcanized rubber, and the artificial sulphides, like HgS, H2S, CS2, etc. The last is a very volatile, ill- smelling liquid, made by the combination of two solids, S being passed over red-hot charcoal. It dissolves S, P, rubber, gums, and many other substances insoluble in H2O.

189. Sulphur Dioxide, SO2, has been made in many experiments. It is a bleaching agent, a disinfectant, and a very active compound, having great affinity for water, but it will not support combustion. Like most disinfectants, it is very injurious to the system. It is used to bleach silk and wool—animal substances— and straw goods, which Cl would injure; but the color can be restored, as the coloring molecule seems not to be broken up, but to combine with SO2, which is again separated by reagents. Goods bleached with SO2 often turn yellow after a time.

190. SO2 a Bleacher.

Experiment 107.-Test its bleaching power by burning S under a receiver under which a wet rose or a green leaf is also placed.

Chapter XXXVIII.

HYDROGEN SULPHIDE.

Examine ferrous sulphide, natural and artificial.

191. Preparation.

Experiment 108.—Put a gram of ferrous sulphide (FeS) into a t.t. fitted with a d.t., as in Figure 32. Add 10cc. H2O and 5cc. H2SO4. H2S is formed. Write the equation, omitting H2O. What is left in solution?

192. Tests.

Experiment 109.-(1) Take the odor of the escaping gas. (2) Pour into a t.t. 5cc.solution AgNO3, and place the end of the d.t. from a H2S generator into the solution and note the color of the ppt. What is the ppt.? Write the equation. (3) Experiment in the same way with Pb(NO3)2 solution. Write the equation. (4) Let some H2S bubble into a t.t. of clean water. To see whether H2S is soluble in H2O, put a few drops of the water on a silver coin. Ag2S is formed. Describe, and write the equation. Do the same with a copper coin. (5) Put a drop of lead acetate solution, Pb(C2H3O2)2, on a piece of unglazed paper, and hold this before the d.t. from which H2S is escap- ing. PbS is formed. Write the equation. This is the characteristic test of H2S.

193. Combustion of H2S

Experiment 110.—Attach a philosopher's lamp tube to the H2S generator, and, observing the same precautions as with H, light the gas. What two products must be formed? State the reaction. The color of the flame. Compute the molecular weight and the vapor density of H2S. 194. Uses. -Hydrogen sulphide or sulphuretted hydrogen, H2S, is employed chiefly as a reagent in the chemical laboratory. It forms sulphides with many of the metals, as shown in the last experiment. These are precipitated from solution, and may be separated from other metals which are not so precipitated, as was found in the case of HCl and NH4OH. The subjoined experiment will illustrate this. Suppose we wished to separate Pb from Ba, having salts of the two mixed together, as Pb(NO3)2 and Ba(NO3)2.

195. H2S an Analyzer of Metals.

Experiment 111.—Pass Some H2S gas in to 5cc.solution Ba(NO3)2. No ppt. is formed. Do the same with Pb(NO3)2 solution. A ppt. appears. Now mix 5cc.of each of these solutions in a t.t. and pass the gas from a H2S generator into the liquid. What is precipitated, and what is unchanged? When fully saturated with the gas, as indicated by the smell, filter. Which metal is on the filter and which is in the filtrate? Other reagents, as Na2CO3 solution, would precipitate the latter.

196. Occurrence and Properties. — H2S is an ill-smell- ing, poisonous gas, formed in sewers, rotten eggs, and other decaying albuminous matter. It is formed in the earth, probably from the action of water on sulphides, and issues with water from sulphur springs.

A characteristic property is the formation of metallic sulphides, as above. A skipper one night anchored his newly painted vessel near the Boston gas-house, where the refuse was deposited, with its escaping H2S. In the morning, to his consternation, the craft was found to be black. H2S had come in contact with the lead in the white paint, forming black PbS. This gradually oxidized after reaching the open sea, and the white color reappeared.

Chapter XXXIX.

PHOSPHORUS.

NOTE.—Phosphorus should be kept in water, and handled with forceps, never with the fingers, except under water, as it is liable to burn the flesh and produce ulcerating sores. Pieces not larger than half a pea should be used, and every bit should finally be burned.

197. Solution and Combustion. Experiment 112. -Put 1 or 2 pieces of P into an evaporating- dish, and pour over them 5 or 10cc.CS2 carbon disulphide. This will be enough for a class. When dissolved, dip pieces of unglazed paper into it, and hold these in the air, looking for any combustion as they dry. The P is finely divided in solution, which accounts for its more ready combustion then. Notice that the paper is not destroyed. This is an example of so-called "spontaneous combustion." The burning- point of P, the combustible, in air, the supporter, is about 60 degrees.

198. Combustion under Water.

Experiment 113. -Put a piece of P in a t.t. which rests in a receiver, add a few crystals KClO3 and 5cc. H2O. Now pour in through a thistle-tube 1cc.or more of H2SO4. Look for any flame. H2SO4 acts very strongly on KClO3. What is set free? From this fact explain the combustion in water.

199. Occurrence.—P is very widely disseminated, but not abundant, and is found only in compounds, the chief of which is calcium phosphate Ca3(PO4)2. It occurs in granite and other rocks, as the mineral apatite, in soils, in plants, particularly in seeds and grains, and in the bones, brains, etc., of vertebrates. From the human system it is excreted by the kidneys as microcosmic salt, HNaNH4PO4; and when the brain is hard- worked, more than usual is excreted. Hence brain-workers have been said to "burn phosphorus."

200. Sources.—Rocks are the ultimate source of this element. These, by the action of heat, rain, and frost, are disintegrated and go to make soils. The rootlets of plants are sent through the soil, and, among other things, soluble phosphates in the earth are absorbed, circulated by the sap, and selected by the various tissues. Animals feed on plants, and the phosphates are circulated through the blood, and deposited in the osseous tissue, or wherever needed.

Human bones contain nearly 60 per cent of Ca3(PO4)2; those of some birds over 80 per cent.

The main sources of phosphates and P are the phosphate beds of South Carolina, the apatite beds of Canada, and the bones of animals.

201. Preparation of Phosphates and Phosphorus.—Bone ash, obtained by burning or distilling bones, and grinding the residue, is treated with H1SO4, and forms soluble H4Ca(PO4)2, superphosphate of lime, and insoluble CaSO4.

Ca3(PO4)2 + 2 H2SO4 = H4Ca(PO04)2 + 2 CaSO4. This completes the process for fertilizers. If P is desired, the above is filtered; charcoal, a reducing agent, is added to the filtrate; the substance is evaporated, then very strongly heated and distilled in retorts, the necks of which dip under water. It is then purified from any uncombined C by melting in hot water and passing into molds in cold water.

The work is very dangerous and injurious, on account of the low burning-point of P, and its poisonous properties. While its compounds are necessary to human life, P itself destroys the bones, particularly the jaw bones, of the workers in it.

Between 1,000 and 2,000 tons are made yearly, mostly for matches, but almost all at two factories, one in England, and one in France. 202. Properties.—P is a colorless, transparent solid, when pure; the impure article is yellowish, translucent, and waxy. It is insoluble in water, slightly soluble in alcohol and ether, and it readily dissolves in CS2, oil of turpentine, etc. Fumes, having a garlic odor, rise when it is exposed to the air, and in the dark it is phosphorescent, emitting a greenish light.

203. Uses. -The uses of this element and its compounds are for fertilizers, matches, vermin poisons, and chemical operations.

204. Matches.-The use of P for matches depends on its low burning-point. Prepared wood is dipped into melted S, and the end is then pressed against a stone slab having on it a paste of P, KClO3, and glue. KNO3 is often used instead of KClO3. In either case the object is to furnish O to burn P. Matches containing KClO3 snap on being scratched, while those having KNO3 burn quietly. The friction from scratching a match generates heat enough to ignite the P, that enough to set the S on fire, and the S enough to burn the wood. Give the reaction for each. Paraffine is much used instead of S. Safety matches have no P, and must be scratched on a surface of red P and Sb2S3, or on glass.

205. Red Phosphorus.-Two or three allotropic forms of P are known, the principal one being red. If heated between 230 degrees and 260 degrees, away from air, the yellow variety changes to red, which can be kept at all temperatures below 260 degrees. Above that it changes back. Red P is not poisonous, ignites only at a high temperature, and is not phosphorescent, like the yellow. 206. Spontaneous Combustion of Phosphene, or Hydrogen Phosphide, PH3.

Experiment 114.—Put into a 20cc.flask 1 g. P and 50cc.saturated solution NaOH or KOH. Connect with the p.t. by a long d.t., as in Figure 44, the end of which must be kept under water. Pour 3 or 4cc.of ether into the flask, to drive out the air. It is necessary to exclude all air, as a dangerously explosive mixture is formed with it. Heat the mixture, and as the gas passes over and into the air, it takes fire spontaneously, and rings of smoke successively rise. It will do no harm if, on taking away the lamp, the water is drawn back into the flask; but in that case the flask should be slightly lifted to prevent breakage by the sudden rush of water. On no account let the air be drawn over.

The experiment has no practical value, but is an interesting illustration of the spontaneous combustion of PH3 and of vortex rings. What are the products of the combustion? An admixture of another compound of P and H causes the combustion.

Chapter XL.

ARSENIC.

Examine metallic arsenic, realgar, orpiment, arsenopyrite, arsenic trioxide, copper arsenite.

The compounds of arsenic are very poisonous if taken into the system, and must be handled with care.

207. Separation. Experiment 115.—Draw out into two parts in the Bunsen flame a piece of glass tubing 20cm long and 1 or 2cm in diameter. Into the end of one of the ignition tubes thus formed, when it is cool, put one-fourth of a gram of arsenic trioxide, As2O3, using paper to transfer it. Now put into the tube a piece of charcoal, and press it down to within 2 or 3cm of the AS2O3 (Fig. 45). Next heat the coal red-hot, and then at once heat the As203. Continue this process till you see a metallic sublimate- metallic mirror-on the tube above the coal. Break the tube and examine the sublimate. It is As. Heat vaporizes the As2O;3. Explain the chemical action. What is the agency of C in the experiment? Of As2O3? 2 As2O3 + 3 C = ?

208. Tests.-Experiments 115 and 116 are used as tests for the presence of arsenic.

Experiment 116.—Prepare a H generator, - a flask with a thistle- tube and a philosopher's lamp tube (Fig. 46), put in some granulated Zn, water, and HCl. Test the purity of the escaping gas (Experiment 23), and when pure, light the jet of H. H is now burning in air. To be sure that there is no As in the ingredients used, hold the inside of a porcelain evaporating-dish directly against the flame for a minute. If no silvery-white mirror is found, the chemicals are free from As. Then pour through the thistle-tube, while the lamp is still burning, 1cc.solution of AS2O3 in HCl or H2O a bit of As2O3 not larger than a grain of wheat in 10 cc. HCl.

See whether the color of the flame changes; then hold the evaporating-dish once more in the flame, and notice a metallic deposit of As. Set away the apparatus under the hood and leave the light burning.

This experiment must not be performed unless all the cautions are observed, since the gas in the flask (AsH3) is the most poisonous known, and a single bubble of it inhaled is said to have killed the discoverer. By confining the gas inside the flask there is no danger.

Instead of using As2O3 solution, a little Paris green, wall paper suspected of containing arsenic, green silk, or green paper labels, etc., may be soaked in HCl, and tested.

209. Explanation.—The chemical changes are as follows: The compounds of As, in this case As2O3, in presence of nascent H, are immediately converted into the deadly hydrogen arsenide (arsine, arseniuretted hydrogen), AsH3. As2O3 + 12 H = 2 AsH3 + 3 H2O. The AsH3 mixed with excess of H tends to escape and is burned to As2O3 and H2O, and thus is rendered comparatively harmless as it passes into the air. This is why the flame must be burning when the arsenic compound is introduced. 2 AsH3 + 6 O = As2O3 + 3 H2O.

In the combustion of AsH3, H burns at a lower point than As. The introduction of a cold body like porcelain cools the flame below the kindling-point of As, and this is deposited, while H burns, in exactly the same way as lamp- black was collected in Experiment 26.

210. Expert Analysis.—A modification of this experiment is employed by experts to test for AS2O3 poisoning. The organs.— stomach or liver—are cut into small pieces dissolved by nascent Cl, or HClO, made from KC1O3 and HCl, and the solution is introduced into a H generator, as above. AS2O3 preserves the tissues it comes in contact with, for a long time, and the test can be made years after death. All the chemicals must be pure, since As is found in small quantities in most ores, and the Zn, HCl, and H2SO4 of commerce are very likely to contain it. The above is called Marsh's test, and is so delicate that a mere trace of arsenic can be detected.

211. Properties and Occurrence.—As is a grayish white solid, of metallic luster, while a few of its characters are non-metallic. It is very widely distributed, being sometimes found native, and sometimes combined, as AsS, realgar, As2S8, orpiment, and FeAsS, arsenopyrite. Its chief source is the last, the fine powder of which is strongly heated, when As separates and sublimes. It has the odor of garlic, as may be observed by heating a little on charcoal with the blow-pipe.

212. Atomic Volume.—As is peculiar in that its atomic volume, so far as the volume can be determined, is only half that of the H atom. Its vapor density is 150, which gives 300 for the molecular weight, while its least combining or atomic weight is 75. 300, the molecular weight = 75, the atomic weight =4, the number of atoms in the molecule. All gaseous molecules being of the same size, represented by two squares, the atomic volume of As must be one-fourth of this size, represented by half of one square. Of what other element is this true? 213. Uses of As2O3.-Arsenic is used in shot-manufacture, for hardening the metal. Its most important compound is As2O3, arsenic trioxide, called also arsenious anhydride, arsenious acid, white arsenic, etc. So poisonous is this that enough could be piled on a one-cent piece to kill a dozen persons. Taken in too large quantities it acts as an emetic. The antidote is ferric hydrate Fe2(OH)6 and a mustard emetic, followed by oil or milk.

The vapor density of this compound shows that its symbol should be As4O6, but the improper one, As2O3, is likely to remain in use. Another oxide, As2O5, arsenic pentoxide, exists, but is less important. Show how the respective acid formulae are obtained from these anhydrides. See page 50.

AS2O3 is used in making Paris green; in many green coloring materials, in which it exists as copper arsenite; in coloring wall papers, and in fly and rat poisons. It is employed for preserving skins, etc. Fashionable women sometimes eat it for the purpose of beautifying the complexion, to which it imparts a ghastly white, unhealthy hue. Mountaineers in some parts of Europe eat it for the greater power of endurance which it is supposed to give them. By beginning with small doses these arsenic-eaters finally consume a considerable quantity of the poison with apparent impunity; but as soon as the habit is stopped, all the pangs of arsenic-poisoning set in. Wall paper containing arsenic is said to be injurious to some people, while apparently harmless to others.

Chapter XLI.

SILICON, SILICA, AND SILICATES.

214. Comparison of Si and C.—The element Si resembles carbon in valence and in allotropic forms. It occurs in three forms like C, a diamond form, a graphite, and an amorphous. C forms the basis of the vegetable and animal world; Si, of the mineral. Most soils and rocks, except limestone, are mainly compounds of O, Si, and metals. While O is estimated to make up nearly one- half of the known crust of the earth, Si constitutes fully a third. The two are usually combined, as silica, SiO2, or silicates, SiO2 combined with metallic oxides. This affinity for O is so strong that Si is not found uncombined, and is separated with great difficulty and only at the highest temperatures. No special use has yet been found for it, except as an alloy with Al. Its compounds are very important.

215 Silica.—Examine some specimens of quartz, rock crystal, white and colored sands, agate, jasper, flint, etc.; test their hardness with a knife blade, and see whether they will scratch glass. Notice that quartz crystals are hexagonal or six-sided prisms, terminated by hexagonal pyramids. The coloring matters are impurities, often Fe and Mn, if red or brown. When pure, quartz is transparent as glass, infusible except in the oxy- hydrogen blow- pipe, and harder than glass. Rock crystal is massive Si02. Sand is generally either silica or silicates.

The common variety of Si02 is not soluble in water or in acids, except HF. An amorphous variety is to some extent soluble in water. Most geysers deposit the latter in successive layers about their mouths. Agate, chalcedony, and opal have probably an origin similar to this. A solution of this variety of SiO2 forms a jelly-like masscolloid—which will not diffuse through a membrane of parchment -dialyzer—when suspended in water. Crystalloids will diffuse through such a membrane, if they are in solution. This principle forms the basis of dialysis.

All substances are supposed to be either crystalloids, i.e. susceptible of crystallization, or colloids-jelly-like masses. HCl is the most diffusible in liquids of all known substances; caramel is one of the least so. To separate the two, they would be put into a dialyzer suspended in water, when HCl will diffuse through into the water, and caramel will remain. As2O3, in cases of suspected poisoning, was formerly separated from the stomach in this way, as it is a crystalloid, whereas most of the other contents of the stomach are colloidal.

216. Silicates.—Si is a tetrad. SiO2 + 2 H2O =? Si02 + H2O =? In either case the product is called silicic acid. Replace all the H with Na, and name the product. Replace it with K; Mg; Fe; Ph; Ca. Na4SiO4 and Na2SiO3 are typical silicates of Na, but others exist.

217. Formation of SiO2 from Sodium Silicate. Experiment 117.—To 5cc.Na4SiO4 in au evaporating-dish add 5cc. HCl. Describe the effect. Pour away any extra HCl. Heat the residue gently, above a flame, till it becomes white, then cool it and add water. In a few minutes taste a drop of the water, then pour it off, leaving the residue. Crush a little in the fingers, and compare it with white sand, SiO2. Apply to the experiment these equations: - Na4SiO4 + 4 HCl = 4 NaCl + H4SiO4. H4SiO4 + 2 H2O = Si02. Why was H4Si04 heated? Why was water finally added?

Water glass, sodium or potassium silicate, used somewhat for making artificial stone, is made by fusing SiO2 with Na2CO3 or K2CO3, and dissolving in water. Silicic acid forms the basis of a very important series of compounds, - the silicates. The above two are the only soluble ones, and may be called liquid glass.

Chapter XLII.

GLASS AND POTTERY.

Examine white sand, calcium carbonate, sodium carbonate, smalt; bottle, window, Bohemian and flint glass.

218. Glass is an Artificial Silicate.—Si02 alone is almost infusible, as is also Ca0; but mixed and heated the two readily fuse, forming calcium silicate. Ca0 + SiO2 = ? Notice that Si02 is the basis of an acid, while CaO is essentially a base, and the union of the two forms a salt. There are four principal kinds of glass: (1) Bohemian, a silicate of K and Ca, not easily fused, and hence used for chemical apparatus where high temperatures are required; (2) window or plate glass, a silicate of Na and Ca; (3) bottle glass, a silicate of Na, Ca, Al, Fe, etc., a variety which is impure, and is tinged green by salts of Fe; (4) flint glass, a silicate of K and Pb, used for lenses in optical instruments, cut glass ware, and, with B added, for paste, or imitation diamonds, etc. Pb gives to glass high refracting power, which is a valuable property of diamonds, as well as of lenses.

219. Manufacture.—Pure white sand, Si02, is mixed with CaCO3 and Na2CO3, some old glass - cullet - is added, and the mixture is fused in fire-clay crucibles. For flint glass, Pb304, red lead, is employed. If color is desired, mineral coloring matter is also added, but not always at this stage. CoO, or smalt, gives blue; uranium oxide, green; a mixture of Au and Sn of uncertain composition, called the "purple of Cassius," gives purple. MnO2 is used to correct the green tint caused by FeO, which it is supposed to oxidize. Opacity, or enamel, as in lamp-shades, is produced by adding As2O3, Sb2O3, SnO2, cryolite, etc. The glass- worker dips his blowpipe—a hollow iron rod five or six feet long—into the fused mass of glass, removes a small portion, rolls it on a smooth surface, swings it round in the air, blowing meanwhile through the rod, and thus fashions it as desired, into bottles, flasks, etc. For some wares, e.g. common goblets, the glass is run into molds and stamped; for others it is blown and welded. All glass must be annealed, i.e. cooled slowly, for several days. The molecules thus arrange themselves naturally. If not annealed, it breaks very easily. It may be greatly toughened by dipping, when nearly red-hot, into hot oil. Cut glass is prepared at great expense by subsequent grinding. Glass may be rendered semi-opaque by etching either with HF, or with a blast of sand.

220. Importance.—Few manufactured articles have more importance than glass. Without it the sciences of chemistry, physics, astronomy, microscopic anatomy, zoology, and botany, not to mention its domestic uses, would be almost impossible.

221. Porcelain and Pottery.—Genuine porcelain and china-ware are made of a fine clay, kaolin, which results from the disintegration of feldspathic rocks. Bricks are baked clay. The FeO in common clay is oxidized to Fe2O3, on heating, a process which gives their red color. Some clay, having no Fe, is white; this is used for fire-bricks and clay pipes. That containing Fe is too fusible for fire-clay, which must also have much SiO2. The electric arc, however, will melt even this, and the most refractory vessels are of calcium oxide or of graphite. Pottery is clay, molded, baked, and either glazed, like crockery, or unglazed, like flower-pots. Jugs and coarse earthenware are glazed by volatilizing NaCl in an oven which holds the porous material. This coats the ware with sodium silicate. To glaze china, it is dipped into a powder of feldspar and SiO2 suspended in water and vinegar, and then fused. If the ware and glaze expand uniformly with heat, the latter does not crack.

Chapter XLIII.

METALS AND THEIR ALLOYS.

222. Comparison of Metals and Non-Metals.—The majority of elements are metals, only about a dozen being non-metallic in their properties. The division line between the two classes is not very well defined; e.g. As has certain properties which ally it to metals; it has other properties which are non-metallic. H occupies a place between the two classes. The following are the more marked characteristics of each group: -

METALS.

1. Metals are solid at ordinary temperatures, and usually of high specific gravity.

Exceptions: Hg is liquid above -39.5 degees; Li is the lightest solid known; Na and K will float on water.

2. Metals reflect light in a way peculiar to themselves. They have what is called a metallic luster.

3. They are white or gray. Exceptions: Au, Ca, Sr are yellow; Cu is red.

4. In general they conduct heat and electricity well.

NON-METALS. 1. Non-metals are either gaseous or solid at ordinary temperatures, and of low specific gravity. Exceptions: Br is a liquid; I has the heaviest known vapor.

2. Non-metallic solids have different lusters, as glassy, resinous- silky, etc. Exceptions: I, B, and C have metallic luster.

3. Non-metals have no characteristic color.

4. They are non-conductors of heat and electricity. Exceptions: C and some others are conductors. 5. They are usually malleable and ductile.

6. They form alloys, or "chemical mixtures," with one another, similar to other solutions. Exceptions: Some, as Ph and Zn, will not alloy with one another.

7. Metals are electro-positive elements, and unite with O and H to form bases. Exceptions: Some of the less electro-positive metals, with a large quantity of O, form acids, as Cr, As, etc.

Numbers 2, 6, and 7 are the most characteristic and important properties.

5. They are deficient in malleability and ductility.

6. They often form liquid solutions, similar to alloys in metals.

7. Non-metals are electronegative, and with H, or with H and O, form acids.

Examine brass, bronze, bell-metal, pewter, German silver, solder, type-metal.

223. Alloys.-An alloy is not usually a definite chemical compound, but rather a mixture of two or more metals which are melted together. One metal may be said to dissolve in the other, as sugar dissolves in water. The alloy has, however, different properties from those of its elements. For example, plumber's solder melts at a lower temperature than either Ph or Sn, of which it is composed. Some metals can alloy in any proportions. Solder may have two parts of Sn to one of Pb, two of Pb to one of Sn, or equal parts of each, or the two elements may alloy in other proportions. Not all metals can be thus fused together indefinitely; e.g., Zn and Pb. Nickel and silver coins are alloyed with Cu, gold coins with Cu and Ag.

Gun-metal, bell-metal, and speculum-metal are each alloys of Cu and Sn. Speculum-metal, used for reflectors in telescopes, has relatively more Sn than either of the others; gun-metal has the least. An alloy of Sb and Pb is employed for type-metal as it expands at the instant of solidification. Pewter is composed of Sn and Pb; brass, of Cu and Zn; German silver, of brass and Ni; bronze, of Cu, Sn, and Zn; aluminium bronze, of Cu and Al.

224. Low Fusibility is a feature of many alloys. Wood's metal, composed of Pb eight parts, Bi fifteen, Sn four, Cd three, melts at just above 60 degrees, or far below the boiling-point of water. By varying the proportions, different fusing-points are obtained. This principle is applied in automatic fire alarms, and in safety plugs for boilers and fire extinguishers. Water pipes extend along the ceiling of a building and are fitted with plugs of some fusible alloy, at short distances apart. When, in case of fire, the heat becomes sufficiently intense, these plugs melt and the water flows out.

225. Amalgams.—An amalgam is an alloy of Hg and another metal. Mirrors are "silvered" with an amalgam of Sn. Tin-foil is spread on a smooth surface and covered with Hg, and the glass is pressed thereon.

Various amalgams are employed for filling teeth, a common one being composed of Hg, Ag, and Sn. Au or Ag, with Hg, forms an amalgam used for plating. Articles of gold and silver should never be brought in contact with Hg. If a thin amalgam cover the surface of a gold ring or coin, Hg can be removed with HNO3, as Au is not attacked by it. Would this acid do in case of silver amalgam? Heat will also quickly cause Hg to evaporate from Au.

CHAPTER XLIV.

SODIUM AND ITS COMPOUNDS.

Examine NaCl, Na2SO4, Na2CO3, Na, NaOH, HNaCO3, NaNO3.

226. Order of Derivation.—Though K is more metallic, or electro- positive, than Na, the compounds of Na are more important, and will be considered first. The only two compounds of Na which occur extensively in nature are NaCl and NaNO3. Almost all others are obtained from NaCl, as shown by this table, which should be memorized and frequently recalled.

) Na NaCl ) Na2SO4) Na2CO3) NaOH NaNO3) ) ) HNaCO3



From what is Na2SO4 prepared, as shown by the table? Na2CO3? Na?

227. Occurrence and Preparation of NaCl.—NaCl occurs in sea water, of which it constitutes about three per cent, in salt lakes, whose waters sometimes hold thirty per cent, or are nearly saturated, and, as rock salt, in large masses underground. Poland has a salt area of 10,000 square miles, in some parts of which the pure transparent rock salt is a quarter of a mile thick. In Spain there is a mountain of salt five hundred feet high and three miles in circumference. France obtains much salt from sea water. At high tide it flows into shallow basins, from which the sun evaporates the water, leaving NaCl to crystallize. In Norway it is separated by freezing water, and in Poland it is mined like coal. In New York and Michigan it is obtained by evaporating the brine of salt wells, either by air and the sun's heat, or by fire. Slow evaporation gives large crystals; rapid, small ones.

228. Uses.—The main uses are for domestic purposes and for making the Na and Cl compounds. In the United States the consumption amounts to more than forty pounds per year for every person.

229. Sodium Sulphate.—What acid and what base are represented by Na2SO4? Which is the stronger acid, HCl or H2SO4? Would the latter be apt to act on NaCl? Why?

230. Manufacture.—This comprises two stages shown by the following reactions, in which the first needs moderate heat only; the last, much greater.

(1) 2 NaCl + H2SO4 = HNaSO4 + NaCl + HCl: (2) NaCl + HNaSO4 = Na2S4 + HCl.

The operation is carried on in large furnaces. The gaseous HCl is passed into towers containing falling water in a fine spray, for which it has great affinity. The solution is drawn off at the base of the tower. Thus all commercial HCl is made as a by- product in manufacturing Na2SO4.

When crystalline, sodium sulphate has ten molecules of water of crystallization (Na2SO4, 10 H2O); it is then known as Glauber's salt. This salt readily effloresces; i.e. loses its water of crystallization, and is reduced to a powder. Compute the percentage of water.

231. Uses.—The leading use of Na2SO4 is to make Na2CO3; it is also used to some extent in medicine, and in glass manufacture. 232. Sodium Carbonate.—Note the base and the acid which this salt represents. Test a solution of the salt with red and blue litmus, and notice the alkaline reaction. Do you see any reason for this reaction in the strong base and the weak acid represented by the salt?

233. Manufacture.—Na2CO3 is not made by the union of an acid and a base, nor is H2CO3 strong enough to act on many salts. The process must be indirect. This consists in reducing Na2SO, to Na2S, by taking away the O with C, charcoal, and then changing Na2S to Na2O3 by CaCO3, limestone. The three substances, Na2SO4, C, CaCO3, are mixed together and strongly heated. The reactions should be carefully studied, as the process is one of much importance.

(1) Na2SO4 + 4 C = Na2S + 4 CO. (2) Na2S + CaCO3 = CaS + Na2CO3.

Observe that C is the reducing agent. The gas CO escapes. The solid products Na2CO3 and CaS form black ash, the former being very soluble, the latter only sparingly soluble in water. Na2CO3 is dissolved out by water, and the water is evaporated. This gives commercial soda. CaS, the waste compound in the process, contains the S originally in the H2SO4 used. This can be partially separated and again made into acid. Describe the manufacture of NaCO3 in full, starting with NaCl. This is called the Le Blanc process, but is not the only one now employed to produce this important article.

234. Occurrence.-Sodium carbonate is found native in small quantities. It forms the chief surface deposit of the "alkali belt" in western United States, where it often forms incrustations from an inch to a foot in thickness. It was formerly obtained from sea-weeds, by leaching their ashes, as, by a like process, K2CO3 was obtained from land plants.

235. Uses.—Na2CO3 forms the basis of many alkalies, as H2SO4 does of acids. Of all chemical compounds it is one of the most important, and its manufacture constitutes one of the greatest chemical industries. Its economical manufacture largely depends on the demand for HCl, which is always formed as a by-product. As but little HCl is used in this country, Na2CO3 is mostly manufactured in Europe. The chief uses are for glass and alkalies.

236. Sodium.—Na must always be kept under naphtha, or some other liquid compound containing no O, since it oxidizes at once on exposure to the air. For this reason it never occurs in a free state.

237. Preparation.-By depriving Na2CO3 of C and O, metallic sodium is formed. As usual, heated charcoal is the reducing agent. The end of the retort, which holds the mixture, dips under naphtha.

Na2CO3 + 2 C = 2 Na + 3 CO. The process is a difficult one, and Na brings five dollars per pound, though in its compounds it is a third as common as Fe. K is as abundant as Na, but more difficult of separation, and is worth three dollars per ounce. Notice the position of K and Na at the positive end of the elements.

238. Uses.—Na is used to reduce Al, Ca, Mg, Si, which are the most difficult elements to separate from their compounds. It acts in these cases as a reducing agent.

239. Sodium Hydrate. Review Experiment 62.

Experiment 118.—Put into a t.t. 10cc. H2O and 2 or 3 g. NaOH. Note its easy solubility. Test with litmus. Will it neutralize any acids?

240. Preparation. — Sodium hydrate, caustic soda, or soda by lime, is made by treating a solution of Na2CO3 with milk of lime. CaCO3 is precipitated and al- lowed to settle, the solution is poured off, and NaOH is obtained by evaporating the water and running the residue into molds.

241. Use.—NaOH is a powerful caustic, but its chief use is in making hard soap.

242. Hydrogen Sodium Carbonate.—Hydrogen so- dium carbonate, bicarbonate of sodium, acid sodium carbonate, cooking-soda, etc., HNaCO3, is prepared by passing CO2 into a solution of Na2CO3. Na2CO3 + H2O + CO2 = 2 HNaCO3. Test a solution of it with litmus. Account for the result. Its use in bread-making depends on the ease with which CO2 is liberated. Even a weak acid, as the lactic acid of sour milk, sets this free, and thus causes the dough to rise.

243. Sodium Nitrate.—Sodium nitrate occurs in Chili and Peru. It is the main source of HNO3.

Review Experiments 46 and 52. From NaNO3 is also made KNO3, (NaNO3 + KCl = NaCl + KNO3), one of the ingredients of gunpowder. By reason of its deliqcescence NaNO3 is not suitable for making gunpowder, though it is sometimes used for blasting-powder. The action of the latter is slower than that made from KNO3. NaNO3 is cheaper and more abundant than KNO3; this is true of most Na compounds in comparison with those of K.

Chapter XLV.

POTASSIUM AND AMMONIUM.

POTASSIUM AND ITS COMPOUNDS.

Examine K, KCl, K2SO4, K2CO3, KOH, HKCO3, KCLO3, KCN.

244. Occurrence and Preparation.—Potassium occurs only in combination, chiefly as silicates, in such minerals as feldspar and mica. By their disintegration it forms a part of soils from which such portions as are soluble are taken up by plants. The ashes of land-plants are leached in pots to dissolve K2CO3; hence it is called potash. Sea-plants likewise give rise to Na2CO3. Wood ashes originally formed the main source of K2CO3. From plants this substance is taken into the animal system, and makes a portion of its tissue. Sheep excrete it in sweat, which is then absorbed by their wool. Large quantities are now obtained by washing wool and evaporating the water. K2CO3 and other compounds of K are mainly derived from KCl, beds of which exist in Germany.

In the following list each K compound is prepared like the same Na compound, and the uses of each of the former are similar to those of the latter. K compounds are made in much smaller quantities than those of Na, as KCl is far less common than NaCl.

{ K KCl { K2SO4 { K2CO3 { KOH KNO3 { { HKCO3



Examine specimens of each, side by side with like Na compounds. Describe in full their preparation, giving the reactions. Also, perform theexperiments given under Na, substituting K therefor. From KOH are made KClO3 and KCN.

KOH {KCl03 {KCN

245. Potassium Chlorate.—KCl03 is made by passing Cl into a hot concentrated solution of KOH.

6 KOH + 6 Cl = KCl03 + 5 KCl + 3 H2O

Its uses are making O, and as an oxidizing agent.

246. Potassium Cyanide, KCN, is a salt from HCN—hydrocyanic or prussic acid. Each is about equally poisonous, and more so than any other known substance. A drop of pure HCN on the tongue will produce death quickly by absorption into the system. In examining these compounds take care not to handle them or to inhale the fumes. KCN is used as a solvent for metals in electro-plating, and is the source of many cyanides, i.e. compounds of CN and a metal. KCN is employed to kill insects for cabinet specimens. In a wide-mouthed bottle is placed a little KCN, which is covered with cotton, and over this a perforated paper. The bottle is inverted over the insect, and the fumes destroy life without injuring the delicate parts. HCN is made from KCN and H2SO4.

247. Gunpowder.—Gunpowder is a mixture of KNO3, C, and S. Heat or concussion causes a chemical change, and transforms the solids into gases. These gases at the moment of explosion occupy 1500 or more times the volume of the solids. Hence the great rending power of powder. If not confined, powder burns quietly but quickly. The appended reaction is a part of what takes place, but it by no means represents all the chemical changes.

2KNO3 + S + 3C =K2S + 2N + 3CO2.

From this equation compute the percentage, by weight, of each substance used to make gunpowder economically.

Thoroughly burned charcoal, distilled sulphur, and the purest nitre are powdered and mixed in a revolving drum,made into a paste with water, put under great pressure between sheets of gun metal, granulated, sifted, to separate the coarse and fine grains, and glazed by revolving in a barrel which sometimes contains a little powdered graphite.

Experiment 119.—Pulverize and mix intimately 4 g. KNO3, l/2 g. S, 1/2 g. charcoal. Pile the mixture on a brick, and apply a lighted match. The adhering product can be removed by soaking in water.

AMMONIUM COMPOUNDS.

248. Read the chapter on NH3. Also, review the experiments on bases. Examine NH4Cl, NH4NO3, (NH4)2SO4, (NH4)2CO3.

Ammonium, NH4, is too unstable to exist alone, but it forms salts similar to those of K and Na. NH3 dissolved in water forms NH4OH.

The food of plants, as well as that of animals, must contain N. It has not yet been shown that they can make use of that contained in the air, but they do absorb its compounds from the soil. All fertilizers and manures contain a soluble compound of NH4. All NH4 compounds are now obtained either from coal, in making illuminating-gas, or from bones, by distillation.

Suppose the product obtained from the gas-house to be NH4OH, how would NH4Cl be made? (NH4)2SO4? NH4NO3? Write the reactions. (NH4)2CO3 is made by heating NH4Cl with CaCO3. Give the reaction.

Chapter XLVI.

CALCIUM COMPOUNDS.

Examine CaCO3—marble, limestone, chalk, not crayon,—CaSO4 — gypsum or selenite—CaCl2, CaO.

249. Occurrence.—The above are the chief compounds of Ca. The element itself is not found uncombined, is very difficult to reduce (page 141), is a yellow metal, and has no use. Its most abundant compound is CaCO3. Shells of oysters, clams, snails, etc., are mainly CaCO3, and coral reefs, sometimes extending thousands of miles in the ocean, are the same. CaCO3 dissolves in water holding CO2, and thence these marine animals obtain it and therefrom secrete their bony framework. All mountains were first laid down on the sea bottom layer by layer, and afterwards lifted up by pressure. Rocks and mountains of CaCO3 were formed by marine animals, and all large masses of CaCO3 are thought to have been at one time the framework of animals. Marble is crystallized, transformed limestone. The process, called metamorphism, took place in the depths of the earth, where the heat is greater than at the surface.

250. Lime.—If CaCO3 be roasted with C, CO2 escapes and CaO is left. CaCO3 - CO2 = ? This is called burning lime, and is a large industry in limestone countries. CaO is unslaked lime, quicklime or calcium oxide. It may be slaked either by exposure to the air, air-slaking, when it gradually takes up H2O and CO2; or by mixing with H2O, water-slaking. Ca0 + H2O = Ca(OH)2.

Great heat is generated in the latter case, though not so much as in the formation of KOH and NaOH. Like them, Ca(OH)2 dissolves in water, forming lime-water. Milk of lime, cream of lime, etc., consist of particles of Ca(OH)2 suspended in H2O.

251. Uses of Lime—CaO is infusible at the highest temperatures. If it be introduced into the oxy-hydrogen blow-pipe (page 28), a brilliant light, second only to the electric, is produced. Mortar is made by mixing CaO, H2O, and Si02. It hardens by evaporating the extra H2O, absorbing CO2 from the air, and uniting with Si02 to form calcium silicate. It often continues to absorb CO2 for hundreds or thousands of years before being saturated, as is found in the Egyptian pyramids. Hence the tenacity of old mortar. Hydraulic mortar contains silicates of Al and Ca, and is not affected by water. What are the uses of mortar? Being the important constituent of mortar and plaster, lime is the most useful of the bases.

252. Hard Water.—Review Experiment 76. The solubility of CaCO3 in water that contains CO2 leads to important results. Much dissolves in the waters of all limestone countries; and the water, though perfectly transparent, is hard; i.e. soap has little action on it. See page 187. Such water may be softened by boiling, a deposit of CaCO3 being formed as a crust on the kettle. Such water is called water of temporary hardness. MgCO3 produces a similar effect, and water containing it is softened in the same way. Permanently hard waters contain the sulphates of Ca and Mg, which cannot be removed by boiling, but may be by adding (NH4)2CO3. 253. The Formation of Caves in limestone rocks is due also to the solubility of CaCO3. Water collects on the mountains and trickles down through crevices, dissolving, if it contains CO2, some of the CaCO3, and thus making a wider opening, and forcing its way along fissures and lines of least resistance into the interior of the earth, or out at the base of the mountain. Its channel widens as it dissolves the rock, and the stream enlarges until in the course of ages an immense cavern may be formed, with labyrinths extending for miles, from the entrance of which a river often issues. In the long ages which elapsed during the slow formation of Mammoth Cave its denizens lost many of the characters of their ancestors, and eyeless fish and also eyeless insects now abound there.

254. Reverse Action.—Drops of water on the roofs of these caverns lose their CO2, and deposit CaCO3. Thus long, pendant masses of limestone, called stalactites, are slowly formed on the roofs like icicles. From these, water charged with CaCO3 drops to the bottom, loses CO2 and deposits CaCO3, which forms an upward- growing mass, called stalagmite. In time it may meet the stalactite and form a pillar. Notice that the same action which formed the cave is filling it up; i.e. the solubility of CaCO3 in water charged with CO2.

255. Famous Marbles.—The marble from Carrara, Italy, is most esteemed on account of a pinkish tint given by a trace of oxide of iron. The best of Grecian marble was from Paros, one of the Cyclades. The isles of the Mediterranean are of limestone, or of volcanic, origin, often of both. 256. Calcium Sulphate occurs in two forms, (1) with water of crystallization—gypsum, CaSO4 + 2 H2O, —(2) without it—anhydrite, CaSO4. The former, on being strongly heated, gives up its water, and is reduced to a powder— plaster of Paris. This, on being mixed with water, again takes up 2 H2O, and hardens, or sets, without crystallizing. If once more heated to expel water, it will not again absorb it. When plaster of Paris sets, it expands slightly, and on this account is admirable for taking casts.

257. Uses.—Gypsum finds use as a fertilizer and as an adulterant in coloring-materials, etc. CaSO4 is employed in making casts, molds, statuettes, wall-plaster, crayons, etc.

How can CaCl2 be made? What is its use? See page 27. What else is used for a similar purpose?

Symbolize and name the acid represented by Ca(ClO)2, and name this salt (page 107). It is one of the constituents of bleaching- powder, the symbol of which, though still under discussion, may be considered Ca(ClO)2 + CaCl2. This is made by passing Cl over Ca(OH)2 2 Ca(OH)2 + 4 Cl = Ca(ClO)2 + CaCl2 + 2 H2O.

CHAPTER XLVII.

MAGNESIUM, ALUMINIUM, AND ZINC.

MAGNESIUM AND ITS COMPOUNDS.

Examine magnesite, dolomite, talc, serpentine, hornblende, meerschaum, magnesium ribbon, magnesia alba, Epsom salt.

258. Occurrence and Preparation.—Mg is very widely distributed, but does not occur uncombined. Its salts are found in rocks and soils, in sea water and in the water of some springs, to which they impart a brackish taste.

The most common minerals containing Mg are magnesite, MgCO3, dolomite, MgCO3 + CaCO3, and talc, serpentine, hornblende, and meerschaum. The last four are silicates, and often are unctious to the touch. What proportion of the earth's crust is composed of Mg? See page 173.

259. Metallic Mg is prepared by fusing MgCl2 with Na. Why is the process expensive? Write the reaction.

Experiment 120.—With forceps hold a short strip of Mg ribbon in a flame. Note the brilliancy of the light, and give the reaction. Examine and name the product.

Photographs of the interior of caverns, where sunlight does not penetrate, are taken by Mg light. Gun-cotton sprinkled with powdered Mg has recently been employed for that purpose. Mg tarnishes slightly in moist air. Compounds of Mg.—MgO, magnesia, like CaO, is very infusible, and is used for crucibles. Magnesia alba, a variable mixture of MgCO2 and Mg(OH)2, is employed in medicine, as is also Epsom salt, MgSO4 + 7 H2O.

ALUMINIUM AND ITS COMPOUNDS.

Examine aluminium, aluminium bronze, corundum, emery, feldspar, argillite, clay. Note especially the color, luster, specific gravity and flexibility of Al.

What elements are more common in the earth than Al? What metals? Compare the abundance of Al with that of Fe.

260. Compounds of Al.—Al occurs only in combination with other elements. Feldspar, mica, slate, and clay are silicates of it. It occurs in all rocks except CaCO3 and SiO2, and in nearly 200 minerals. Though found in all soils, its compounds are not taken up by plants, except by a few cryptogams. Corundum, Al2O3, is the richest of its ores. Compute its percent of Al. Compounds of Al are very infusible and difficult of reduction.

261. Reduction.—Like most other metals not easily reducible by C or H, it was originally obtained by electrolysis, but more recently from its chloride, by the reducing action of strongly heated K or Na. Al2Cl6 + 6 Na = 6 NaCl + 2 Al.

What is the chief use of Na? As it takes three pounds of Na to make one pound of Al, the cost of the latter has been fifteen dollars or more per pound. Its use has thus been restricted to light apparatus and aluminium bronze, an alloy of Cu 90, Al 10, which is not unlike gold in appearance.

Al2O3 has lately been reduced by C. Higher temperatures than have heretofore been known are obtained by means of the electric arc and large dynamo machines. Afurnace made of graphite, because fire-clay melts like wax at such a high temperature, is filled with Al2O3—corundum, —C, and Cu. In the midst of this are embedded large carbon terminals, connected with dynamos. The reduction takes several hours.

The following reaction takes place: Al2O3 + 3 C = 2 Al + 3 CO. Cu is also added, and an alloy of Al and Cu is thus formed. This alloy is not easily separable into its elements. Explain the action of the C. CO escapes through perforations in the top of the furnace, burning there to CO2. Only alloys of Al have yet been obtained by this process. This method has not been employed before, simply because the highest temperatures of combustion, 2000 degrees or 2500 degrees, would not effect a reduction. In the same way Si, B, K, Na, Ca, Mg, Cr, have recently been reduced from their oxides; but a process has yet to be found for separating them easily from their alloys.

262. Properties and Uses.—Al is a silvery white metal, lighter than glass, and only one-third the weight of iron. It does not readily rust or oxidize, it fuses at 1000 degrees (compare with Fe), is unaffected by acids, except by HCl and, slightly, by H2SO4, is a good conductor of electricity, can be cast and hammered, and alloys with most metals, forming thus many valuable compounds. Every clay-bank is a mine of this metal, which has so many of the useful properties of metals and has so few defects that, if it could be obtained in sufficient quantities, it might, for many purposes, take the place of iron, steel, tin, and other metals. From its properties state any advantages which it would have over iron in ocean vessels, railroads, and bridges. Why is it better than Sn or Cu for culinary utensils? An alloy of Al, Cu, and Si is used for telephone wires in Europe, and the Bennett-Mackay cable is of the same material. Washington monument, the tallest shaft in the world, is capped with a pyramid of Al,ten inches high.

For the uses of alumina, Al2O3, and its silicates, see page 133.

ZINC AND ITS COMPOUNDS.

Examine zincite, sphalerite, Smithsonite, sheet zinc, galvanized iron, granulated zinc, zinc dust.

263. Compounds.—The compounds of zinc are abundant. Its chief ores are zincite, ZnO, sphalerite or blende, ZnS, Smithsonite, ZnCO3. For their reduction these ores are first roasted, i.e. heated in presence of air. With ZnS this reaction takes place: ZnS + 3 O = Zn0 + S02. The oxide is reduced with C, and then Zn is distilled. State the reaction. Zinc is sublimed-in the form of zinc dust-like flowers of S. Granulated Zn is made by pouring a stream of the molten metal into water.

Experiment 121.—Burn a strip of Zn foil, and note the color of the flame and of the product. State the reaction. The red color of zincite is supposed to be imparted by Mn present in the compound.

264. Uses.—Name any use of Zn in the chemical laboratory. It is employed for coating wire and sheet iron —galvanized iron. This is done by plunging the wire or the sheets of iron into melted Zn. Describe the use of Zn as an alloy. See page 136.

ZnO forms the basis of a white paint called zinc white. White vitriol, ZnSO4 + 7 H2O, is employed in medicine. Name two other vitriols.

CHAPTER XLVIII.

IRON AND ITS COMPOUNDS.

Examine magnetite, hematite, limonite, siderite, pig-iron, wrought-iron, steel.

265. Ores and Irons.—As Fe occurs native only in meteorites and in small quantities of terrestrial origin, it is obtained from its ores. There are four of these ores—magnetite (Fe3O4), hematite (Fe2O3), limonite (2 Fe2O3 + 3 H2O), and siderite (FeCO3). Which is richest in Fe? Compute the proportion. FeCO3 occurs mostly in Europe. The reduction of these ores, as well as of other metallic oxides, consists in removing O by C at a high tempera- ture. As ordinarily classified there are three kinds of iron,—pig- or cast-iron, steel, and wrought-iron.

Study this table, noting the purity, the fusing-point, and the per cent of C in each case.

Per Cent Fe Fusibility. Per Cent (general). C. Pig......... 90 1200 degrees 2-6 Steel........ 99 1400 degrees 0.5-2 Wrought....... 99.7 1500 degrees Fraction.

Pure iron melts at about 1800 degrees. Pig-iron is obtained from the ore by smelting, and from this are made steel and wrought- iron.

266. Pig-Iron.—The ore is reduced in a blast furnace (Fig. 47), in some cases eighty or one hundred feet high, and having a capacity of about 12,000 cubic feet. The reducing agent is either charcoal, anthracite coal, or coke,bituminous coal being too impure. Charcoal is the best agent, and is used in preparing Swedish iron; but it is too expensive for general use.

Fig. 47. Blast furnace. F, entrance of tuyeres, or blast-pipes. E, F, hottest part. C, conductor for gases, which are subsequently used to heat the air going into the tuyeres. G, upper portion, slag, lower portion, melted iron.

Were ores absolutely pure, only C would be needed to reduce them. Complete: Fe3O4 + 4 C ? Fe3O4 + 2C?

Much earthy material—gangue—containing silica and silicates is always found with iron ores. These are infusible, and something must be added to render them fusible. CaO forms with SiO2 just the flux needed. See page 132. Ca0 + Si02 = ? Which of these is the basic, and which the acidic compound? CaO results from heating CaCO3; hence the latter is employed instead of the former. In what case would Si02 be used as the flux?

Into the blast furnace are put, in alternate layers, the fuel, the flux, and the ore. The fire, once kindled, is kept burning for months or years. Hot air is driven in through the tuyeres (tweers). O unites with C of the fuel, forming CO2 and CO. The C also reduces the ore. Fe2O3 + 3 C = ? CO accomplishes the same thing. 3 CO + Fe2O3 = ? The intense heat fuses CaO and SiO2 to a silicate which, with other impurities, forms a slag; this, rising to the surface of the molten mass, is drawn off. The iron is melted, falls in drops to the bottom, and is drawn off into sand molds. See Figure 47. This is pig-iron. It contains as impurities, C, Si, S, P, Mn, etc. If too much S or P is present in an ore, it is worthless. This is why the abundant mineral FeS2 cannot be used as a source of iron. From the top of the furnace N, CO, CO2, H2O, etc., escape. These gases are used to heat the air which is forced through the tuyeres, and to make steam in boilers.

267. Steel.—The manufacture of steel and wrought-iron consists in removing most of the impurities from pig-iron. It will be seen that the most common compounds of C, S, Si, and P, are their oxides, and these are for the most part gases. Hence these elements are removed by oxidation.

Bessemer steel is prepared by melting pig-iron and blowing hot air through it. A converter (Fig. 48) lined with siliceous sand, and holding several tons, is partially filled with the molten metal; blasts of hot air are driven into it, and the C and other impurities, together with a little of the Fe, are oxidized. The exact moment when the process has gone far enough, and most of the impurities have been removed, is indicated by the appearance of the escaping flame. It usually takes from five to ten minutes. The blast is then stopped, and the metal has about the composition of wrought-iron; it contains some uncombined O. A white pig-iron (spiegeleisen), which contains a known quantity of C and of Mn, is at once added. Mn removes part of the extra O, and, though it remains, does not injure the metal. The C is "dissolved" by the Fe, which is then run into molds (ingots). This process, the Bessemer, invented in 1856, has revolutionized steel manufacture. No less than ten tons of iron have been converted into steel, in five minutes, in a single converter.

268. Wrought-Iron.—The chemical principle involved in making wrought-iron is the same as that in making steel, but the process is different. Impurities are burned out from pig-iron in an open reverberatory furnace, by constantly stirring the metal in contact with air. This is called puddling. A reverberatory furnace is one in which the fuel is in one compartment, and the heat is reflected downward into another, that holds the substance to be acted upon (Fig. 49).

Steel may also be made by carburizing wrought-iron. Iron and charcoal are packed together and heated for days, without melting, when it is found that, in some unknown way, solid C has penetrated solid Fe. The finer kinds of steel are made in this way, but they are very expensive.

Wrought-iron may also be made directly from the ore in an open hearth furnace, with charcoal. This was the original mode.

269. Properties.—The varying properties of pig-iron, steel, and wrought-iron are due in part to the proportion of C and of other elements present, either as mixtures or as compounds, and in part to other causes not well understood. Wrought-iron is fibrous, as though composed of fine wires, and hence is ductile, malleable, tough, and soft, and cannot be hardened or tempered, but it is easily welded. Pig-iron is crystalline, and so is not ductile or malleable; it is hard and brittle, and cannot be welded. On account of its low melting-point it is generally employed for castings. Steel is crystalline in structure, and when suddenly cooled from red heat by plunging into cold water, becomes hard and brittle. The tempering can be varied by afterwards heating to any required degree, indicated by the color of the oxide formed on the exterior. The higher temperatures give the softer steel.

270. Salts of Iron.—Examine FeSO4, FeS, FeS2.

Fe has a valence of 2 or 4. This gives rise to two kinds of salts, ferrous and ferric, as in FeCl2 and Fe2Cl6 The valence of Fe in ferric salts is 4. Ferrous sulphate is FeSO4; ferric sulphate, Fe2(SO4)3. Write the symbols for ferrous and ferric hydrate; for the oxides; for the nitrates. Write the graphic symbols for each.

271. Colors.—The characteristic color of ferrous salts is green, as in FeSO4. These salts give the green color to the chlorophyll in leaves and grass, and bottle glass owes its green color to ferrous silicate. Ferric salts are a brownish red, as shown in hematite and limonite, and in some bottles. Red sandstone, and most soils and earths, are illustrations of this coloring action. The blood of vertebrates owes its color to ferric salts. Bricks are made from a greenish blue clay in which iron exists in the ferrous state. On being heated, ferrous salts are oxidized to ferric, and their color is changed to red. Iron rust is hydrated ferric oxide, Fe2O3 and Fe2(OH)6.

272. Change of Valence.

Experiment 122.—Dissolve 2 g. of iron filings in diluted HCl. Filter or pour off the clear liquid, divide it into two parts, and add NH4OH to one part till a ppt. occurs. Notice the greenish color of Fe(OH)2. Oxidize the other part by adding a few drops of HNO3 and boiling a minute. Now add NH4OH, and observe the reddish color of the ppt., Fe2(OH)6.

Solutions of ferrous salts will gradually change to ferric, if allowed to stand, thus showing the greater stability of the latter. In changing from FeCl2 to Fe2Cl6 oxidation does not consist in adding O, but in increasing the negative element or radical. This is possible only by changing the valence of Fe from 2 to 4. Hence oxidation, in its larger sense, means increasing the valence of the positive element. To oxidize FeSO4 is to make it Fe2(SO4)3, changing the valence of Fe as before. Reduction or deoxidation diminishes the valence of the positive element. Illustrate this by the same iron salts. Illustrate it by PbO and Pb02; AuCl and AuCl3; Sb2S3 and Sb2S5. In this sense define an oxidizing agent. A reducing agent.

273. Ferrous Sulphate.

Experiment 123.—Dissolve a few iron filings in dilute H2SO4, and slowly evaporate for a few minutes. Write the equation.

Ferrous sulphate, green vitriol, or copperas, FeSO4 + 7 H2O, is the source of what acid? See page 66. It is also one of the ingredients in many writing inks. On being heated, or exposed to the air, it loses its water of crystallization and becomes a white powder. It is prepared as above, or by oxidizing moistened FeS2 by exposure to the air.

Ferrous sulphide, protosulphide of iron, FeS, is how prepared? See Experiment 6. State its use. See Experiment 108. It also occurs native.

Ferric sulphide, pyrite, FeS2, occurs native in large quantities. What is its use? See page 65.

CHAPTER XLIX.

LEAD AND TIN.

LEAD.

Examine galena, lead protoxide and dioxide, red-lead, lead carbonate, acetate, and nitrate. Note especially the colors of the oxides, the cubical crystallization and cleavage of galena, the specific gravity of the compounds, the softness of Pb, and the tarnish, Pb2O, which covers it,if long exposed.

274. Distribution of Pb.—Pb is widely distributed, occurring as PbS and PbCO3. PbS, galenite or galena, is its main source. By heating it in air, SO2 is formed, and Pb liberated and drawn off.

Pb is but little acted on by cold H2SO4, unless concentrated. Describe its use in making that acid. See page 65. To show that a little Pb has been dissolved, as PbSO4, in the manufacture of that acid, perform this experiment.

Experiment 124.—To 5cc. of water in a clean t.t. add the same volume of H2SO4, not C.P.; shake, and notice any fine powder suspended. PbSO4, being insoluble in water, is precipitated. What is the test for Pb? See Experiment 109.

275. Poisonous Properties.—Ph is very flexible and soft, and is much used for water pipes. In moist air it is soon coated with suboxide, Pb20, as may be seen by exposing a fresh surface. Some portion of this is liable to dissolve in water, and, as all soluble salts of Pb are poisonous, water that has stood in pipes should not be used fordrinking. Lead is employed as an alloy of tin for covering sheet-iron in "terne plate." T his plate is rarely used except for roofing. The "bright plate," used for tin cans and other purposes, scarcely ever contains any lead except the small portion in solder. In soldering, ZnCl2 is employed for a flux. Sn, Pb, and Zn are somewhat soluble in vegetable acids. If citric acid be present, as it usually is, citrates of these metals are formed, and all of them are poisonous. The action is far more rapid after opening the can, since oxidation is hastened. Hence the contents should be taken out directly after opening.

Lead poisons seem to have an affinity for the tissues of the body, and accumulate little by little. Painter's colic results from lead poisoning. Epsom salt, or other soluble sulphate, is an antidote, since with Pb it makes insoluble PbSO4.

276. Some Lead Compounds.—Lead salts form the basis of many paints. White paint is a mixture of PbCO3 and Pb(OH)2 suspended in linseed oil. It is often adulterated with BaSO4, ZnO, CaCO3. Other lead compounds are used for colored paints. The two chief soluble salts are Pb(NO3)2 and lead acetate, Pb(C2H302)2.

Red-lead, Pb3O4, and, to some extent, litharge, PbO, are employed in glass manufacture. Name the kind of glass in which it is used, describe its manufacture, and write a symbol for lead silicate. What is the characteristic of lead glass? See page 132.

Experiment 125.—Put a small fragment of Pb on a piece of charcoal, and blow the oxidizing flame against it for some time with a mouth blow-pipe. Note the color of the coating on the coal. PbO has formed.

Experiment 126.—Dissolve a small piece of lead in dilute HNO3. Pour off the solution into a t.t. and add HCl or other soluble chloride. Pb(NO3)2 + 2 HCl = ? What is the insoluble product?

Experiment 127.—Add to a solution of Pb(C2H3O2)2 some H2SO4. Give the reaction and the explanation. TIN.

Examine cassiterite, tin foil, "terne plate," "bright plate."

277. Sn occurs as the mineral cassiterite, tin stone, Sn02, and is found in only a few localities, as Banca, Malacca, and England. It does not readily tarnish, and is used to cover thin plates of copper and iron. Tin foil is generally an alloy of Pb and Sn.

Sn is sometimes a dyad, at others a tetrad. Write symbols for its two chlorides, stannous and stannic, also for its sulphides and oxides.

CHAPTER L.

COPPER, MERCURY, AND SILVER.

COPPER.

Examine native copper, chalcopyrite, malachite, azurite, copper acetate, copper nitrate, copper sulphate.

278. Occurrence.—Copper occurs both native and in many compounds, being diffused in rocks and, in minute quantities, in soils, waters, plants, and animals. Spain, Chili, and the United States are the chief Cu producing countries. The extensive mines of Michigan yield the native ore. The Calumet and Heela mine alone produces 4,000,000 pounds per month. The most abundant compound of Cu is chalcopyrite, or copper pyrites, CuFeS2. Malachite, which is green, and azurite, which is blue, are carbonates, the former being used for ornamental purposes.

Cu is, next to Ag, the best conductor of electricity and heat among the elements; it is very ductile, malleable, and tenacious.

Cu has two valences, 1 and 2. Symbolize and name its chlorides, iodides, sulphides, and oxides. Cupric compounds, as a rule, are more stable than cuprous.

279. Uses.—Thousands of tons of Cu find use in domestic utensils, ocean vessels, electric wires, batteries, and plating. Name the chief alloys of Cu and their uses. See page 136. How may CuS be obtained? See Experiment 7. Cu2O, cuprous oxide, is used to color glass red. CUSO4 is employed in calico-printing, electric batteries, etc. It is called blue vitriol.

Paris green, used for killing potato-beetles, is composed chiefly of copper arsenite. Write the symbol for this compound. All soluble salts of Cu are poisonous; hence care should be taken not to bring any acid in contact with copper vessels of domestic use. With acetic acid, what would be formed?

MERCURY AND ITS COMPOUNDS.

Examine cinnabar, vermilion, mercury, red oxide, mercurous and mercuric chloride.

280. Cinnabar, HgS, is practically the only source of mercury— quicksilver. Austria, Spain, and California contain nearly all the mines. In these mines the metal also occurs native to a small extent. It is the only commonly occurring metal that is liquid at ordinary temperatures; it solidifies at about -40 degrees. What other common liquid element? See page 12. Hg is reduced from the ore by Fe, Hg being distilled over and collected in water. Heat regularly expands the metal.

281. Uses.—For uses see Reduction of Ag and Au, pages 165 and 170; amalgams, page 137; laboratory work, page 68. It is also employed for thermometers and barometers, and as the source of the red pigment vermilion, which is artificial HgS.

Compare the vapor density and the atomic weight of Hg, and explain. See page 12. Hg is either a monad or a dyad. Symbolize its ous and ic oxides and chlorides. Which of the following are is salts, and which are ous, and why? HgNO3, Hg(NO3)2, HgCl, HgCl2? Calomel, HgCl or Hg2Cl2, used in medicine, and corrosive sublimate, HgCl2, are illustrations of the ous and ic salts. The former is insoluble, the latter soluble. All soluble compounds of Hg are virulent poisons, for which the antidote is the white of egg, albumen. With it they coagulate or form an insoluble mass.

SILVER AND ITS COMPOUNDS.

282. Occurrence and Reduction.—Silver is found uncombined, and combined, as Ag2S, argenite, and AgCl, horn silver. It occurs usually with galena, PbS. It is abundant in the Western States, Mexico, and Peru. Silver is separated from galena by melting the two metals. As they slowly cool, Pb crystallizes, and is removed by asieve, while Ag is left in the liquid mass. The principle is much like crystallizing NaCl from solution and leaving behind the salts of Mg, etc., in the mother liquor. When, by repeating the process, most of the Pb is eliminated, the rest is oxidized by heating in the air. Pb + O = PbO. Ag does not oxidize, and is left in the metallic state.

Another mode of reduction is to change the silver salt to its chloride, and then remove the Cl with Fe. Roasting with NaCl makes the first change, 2 NaCl + Ag2S = Na2S + 2 AgCl, and with Fe the second, 2 AgCl + Fe = FeCl2 + 2 Ag. Ag is separated from the other products by adding Hg, with which it forms an amalgam. By distilling this, Hg passes over and Ag remains. This is the amalgamating process.

283. Salts of Silver are much employed in organic chemistry, and AgCl, AgBr, and AgNO3 are used in photography. AgNO3 is a soluble, colorless crystal, and is the basis of the silver salts. It blackens when in contact with organic matter. Stains on a photographer's hands are due to this substance, and the use of AgNO3 in indelible inks depends on the same property. This may be due to a reduction of AgNO3 to Ag4O. Stains can be removed from the skin or from linen by a solution of Kl, or of CuCl2 followed by sodium hyposulphite. Lunar caustic is made by fusing AgNO3 crystals, and is used for cauterizing (burning) the flesh. Much AgCN finds use in electroplating.

Experiment 128.—Put 5 cc. AgNO3 solution in each of three t.t. To the first add 3 cc. HCl, to the second 3cc.NaCl solution, and to the third 3 cc. KBr solution. Write the reaction for each case, and notice that the first two give the same ppt., as in fact any soluble chloride would. Filter the second and third, on separate filter papers, and expose half the residue to direct sunlight, observing the change of color by occasionally stirring. Solar rays reduce AgCl and AgBr, it is thought, to Ag2Cl and Ag2Br. Try to dissolve the other half in Na2S2O3, sodium thiosulphate solution. This experiment illustrates the main facts of photography.

CHAPTER LI.

PHOTOGRAPHY.

284. Descriptive.—The silver halogens, AgCI, AgBr, AgI, are very sensitive to certain light rays. Red rays do not affect them; hence ruby glass is used in the "dark room."

Photography involves two processes. The negative of the picture is first taken upon a prepared glass plate, and the positive is then printed on prepared paper. The negative shows the lights and shades reversed, while the positive gives objects their true appearance.

Few photographers now make their own plates, these being prepared at large manufactories. The glass is there covered on one side with a white emulsion of gelatine and AgBr, making what are called gelatine-bromide plates. This is done in a room dimly lighted with ruby light. The plates are dried, packed in sealed boxes, and thus sent to photographers. The artist opens them in his dark room, similarly lighted, inserts the plates in holders, film side out, covers with a slide, adjusts to the camera, previously focused, and makes the exposure to light. The time of exposure varies with the kind of plate, the lens, and the light, from several seconds, minutes, or hours, to 1/250 part of a second in some instantaneous work. In the dark room the plates are removed and can be at once developed, or kept for any time away from the light. No change appears in the plate until development, though the light has done its work.

To develop the plate, it is put into a solution of pyrogallic acid, the developer, and carbonate of sodium, the motive power in the process. Other developers are often used. The chemical action here is somewhat obscure, but those parts of the plates which were affected by the light are made visible, a part of the AgzBr being reduced to Ag by the affinity which sodium pyrogallate has for Br. Ag2Br = 2 Ag + Br. Br is dissolved and Ag is deposited. When the rather indistinct image begins to fade out, the plate is dipped for a minute into a solution of alum to harden the gelatine and prevent it from peeling off (frilling). It is finally soaked in a solution of sodium thiosulphate (hyposulphite or hypo), Na2S208. This removes the AgBr that the light has failed to reduce. The processis called fixing, as the plate may thereafter be exposed to the light with impunity. It must be left in this bath till all the white part, best seen on the back of the plate, disappears. 2AgBr + 3Na2S2O3 = Ag2Na4(S2O3) + 2 NaBr. Both products are dissolved. It is then thoroughly washed. Any dark objects become light in the negative, and vice versa. Why?

For the positive, the best linen paper is covered on one side with albumen, soaked in NaCl solution, dried, and the same side laid on a solution of AgNO3. What reaction takes place? What is deposited on the paper, and what is dissolved? This sensitized paper, when dry, is placed over a negative, film to film, and exposed in a printing frame to direct sunlight till much darker than desired in the finished picture. What is dark in the negative will be light in the positive. Why? The reducing action of sunlight is similar to that in the negative. Explain it.

After printing, the picture is toned and fixed. Toning consists in giving it a rich color by replacing part of the Ag2Cl with gold from a neutral solution of AuCl3. 3 Ag2Cl+ AUCl3 = 6AgCI + Au. Fixing removes the unaffected AgCl, as in the negative, the same substance being used. Describe the action. 2 AgCI + 3 Na2S203 = Ag2Na4(S203) + 2 NaCl. Both the positive and the negative must be well washed after each process, particularly after the last. The picture is then ready for mounting. In fine portrait work both the negative and the positive are retouched. This consists in removing blemishes with colored pencils or India ink.

The negative—No. 1. Dissolve: sulphite soda crystals, 2 oz. (57 g) in 8 oz. (236 cc.) water (distilled); citric acid, 60 grains (4 g) in 1/2 oz. (15 cc.) water; bromide ammonium, 25 grains (1 1/2 g) in 1/2 oz. water; pyrogallic acid, 1 oz. (28 g) in 3 oz. (90 cc.) water. After dissolving, mix in the order named, and filter. No. 2. Dissolve: sulphite soda, 2 oz. (57 g) in 4 oz. (118 cc.) water; carbonate potash, 4 oz. (113 g) in 8 oz. (236 cc.) water. Dissolve separately, mix, and filter. To develop plates, mix 1 dram (3 2/3 cc.) of No. 1 and 1 dram of No. 2 with 2 oz. (60 cc.) water. Cover the plate with the mixture, and leave as long as the picture increases in distinctness. Remove, wash, and put it into a saturated solution of alum for a minute or two, then wash and put it into a half-saturated solution of hypo. Leave till no white AgCl is seen through the back of the plate. Wash it well.

The positive.—1. Dissolve 30 grains (2 g.) pure gold chloride in 15 oz. (450 cc.) water. This forms a stock solution. 2. Make a saturated solution of borax. 3. Prepare a toning bath by adding 1/2 oz. (15 cc.) of the gold chloride solution and 1 oz. (30 cc.) of the borax solution to 7 oz. (210 cc.) water. After printing the picture, wash it in 3 or 4 waters, put it into the toning bath, and leave it till considerably darker than desired; wash, and put it for 15 minutes into a hypo solution that has been, after saturation, diluted with 3 or 4 volumes of water. Then wash repeatedly.

CHAPTER LII.

PLATINUM AND GOLD.

PLATINUM.

Examine platinum foil and wire.

285. Platinum is much rarer than gold, and is about two-thirds as costly as the latter. It is found alloyed with other metals, as An, and is obtained from sand, in which it occurs, by washing. Aqua regia is the only acid which dissolves it, and the action is much slower than with Au. Pt is one of the heaviest metals, having a specific gravity three times that of Fe, or twenty-one and a half times that of water. Its fusing-point is about 1600 degrees, or just below the temperature of the oxy-hydrogen flame. Like Au it has little affinity for other elements, but alloys with many metals. Pt is so tenacious that it can be drawn into wire invisible to the naked eye, being drawn out in the center of a silver wire, which is afterwards dissolved away from the Pt by HNO3. Noting its valences, 2 and 4, write the symbols for the ous and ic chlorides and oxides.

286. Uses.—Pt is much used in chemistry in the form of foil, wire, and crucibles. On what properties does this use depend? Describe its use in making H2SO4.

PtCl4 is made by dissolving Pt in aqua regia, and evaporating the liquid. On heating PtCl4, half of its Cl is given up, leaving PtCl2. If it be still more strongly heated, the Cl all passes off, leaving spongy Pt. By fusing this in the oxy-hydrogen flame, ordinary Pt is obtained. Spongy Pt has a remarkable power of absorbing, or occluding, O without uniting with it. This O it gives up to some other substances, and thus becomes indirectly an oxidizing agent. What other element has this property of occluding gases?

GOLD.

Examine auriferous quartz, gold chloride, yellow and ruby glass colored with gold. 287. Gold is rarely found combined, and has small affinity for other elements, though forming alloys with Cu, Ag, and Hg. Its source is usually either quartz rock, called auriferous quartz, or sand in placer mines. The element is widely distributed, occurring in minute quantities in most soils, sea water, etc. California and Australia are the two greatest gold- producing countries. That from California has a light color, due to a slight admixture of Ag. Australian gold is of a reddish hue, due to an alloy of Cu. Gold-bearing quartz is pulverized, and treated with Hg to dissolve the precious metal, which is then separated from the alloy by distillation. Compare this with the preparation of Ag.

Such is the malleability of Au that it has been hammered into sheets not over one-millionth of an inch thick; it is then as transparent as glass. Gold does not tarnish or change below the melting-point. On account of its softness it is usually alloyed with Cu, sometimes with Ag. Pure gold is twenty-four carats fine. Eighteen carat gold has eighteen parts Au and six Cu. Gold coin has nine parts Au to one part Cu. The most important compound is AuCl3. Describe a use of it. This metal is much employed in electroplating, and somewhat in coloring glass.

CHAPTER LIII.

CHEMISTRY OF ROCKS.

288. Classification.—Rocks may be divided, according to their origin, into three classes: (1) Aqueous rocks. These have been formed by deposition of sedimentary material, layer by layer, on the bottoms of ancient oceans, lakes, and rivers, from which they have gradually been raised, to form dry land. (2) Eruptive or volcanic rocks. These have been forced, as hot fluids, through rents and fissures from the interior of the earth. (3) Metamorphic rocks. These, by the combined action of heat, pressure, water, and chemical agents, have been crystallized and chemically altered. The rocks of the first class, such as chalk, limestone, shale, and sandstone, are distinguished by the existence of fossils in them, or by the successive layers of the material which goes to make up their structure and to give them a stratified appearance. The rocks of the second class are recognized by their resemblance to the products of modern volcanoes and their non-stratified appearance. Rocks of the third class are composed of crystals, which, though often very minute, are minerals having a definite chemical composition. Examples of the third class are gneiss, slate, schist, and marble. The last two classes abound on the Eastern sea-board, while the interior of our continent is composed almost exclusively of stratified sedimentary rocks.

289. Composition.—Rocks are not definite compounds, but variable mixtures of minerals. Some, however, are tolerably pure, as limestone (CaCO3) and sand-stone.

Granite is mainly made up of three minerals,—quartz, feldspar, and mica. Quartz, when pure, is SiO2. Feldspar is a mixed silicate of K and Al, and often several other metals, K2Al2Si6O16 (=K2O, Al2O3, 6 SiO2) symbolizing one variety, while a variety of mica is H8Mg5Fe7Al2Si3O18.

The pupil should learn to distinguish the different minerals in granite. Quartz is glassy, mica is in scales, usually white or black, and feldspar is the opaque white or red mineral.

290. Importance of Siliceous Rocks.—Slate and schist are also mixed silicates. Pure sandstone is SiO2, the red variety being colored by iron. Igneous rocks are always siliceous. Obsidian is a glassy silicate. A mountain of very pure glass, obsidian, two hundred feet high, has lately been found in the Yellow-stone region. We see how important Si is, in the compounds Si02 and the silicates, as a constituent of the terrestrial crust. Limestone is the only extensive rock from which it is absent. Always combined with O, it is, next to the latter, the most abundant of elements. Silicates of Al, Fe, Ca, K, Na, and Mg are most common, and these metals, in the order given, rank next in abundance.

291. Soils.—Beds of sand, clay, etc., are disintegrated rock. Sand is chiefly SiO2; clay is decomposed feldspar, slatestone, etc. Soils are composed of these with an added portion of carbonaceous matter from decaying vegetation, which imparts a dark color. The reddish brown hue so often observed in soils and rocks results from ferric salts.

292. Minerals, of which nearly 1000 varieties are now known, may be simple substances, as graphite and sulphur, or compounds, as galena and gypsum. Only seven systems of crystallizations are known, but these are so modified as to give hundreds of forms of crystals. See Physics. A given chemical substance usually occurs in one system only, but we saw in the case of S that this was not always true.

Crystals of some substances deliquesce, or take water from the air, and thus dissolve themselves. Some compounds cannot exist in the crystalline form without a certain percentage of water. This is called "water of crystallization"; if it passes into the air by evaporation, the crystal crumbles to a powder- and is then said to effloresce.

293. The Earth's Interior.—We are ignorant of the chemistry of the earth's interior. The deepest boring is but little more than a mile, and volcanic ejections probably come from but a very few miles below the surface. The specific gravity of the interior is known to be more than twice that of the surface rock. From this it has been imagined that towards the center heavy metals like Fe and Au predominate; but this is by no means certain, since the greater pressure at the interior would cause the specific gravity of any substance to increase.

294. Percentage of Elements.—Compute the percentage of O in the following rocks, which compose a large proportion of the earth's crust: SiO2, Al2SiO4, CaCO3. Find the percentage of O in pure water. In air. Taking cellulose, C16H30O15, as the basis, find the percentage of O in vegetation.

An estimate, based on Bunsen's analysis of rocks, of the chief elements in the earth's crust, is as follows:—

O, 46 per cent Ca, 3 per cent Si, 30 per cent Na, 2 per cent Al, 8 per cent K, 2 per cent Fe, 6 per cent Mg, 1 per cent

More than half the elements are known to exist in sea-water, and the rest are thought to be there, though dissolved in such small quantity as to elude detection. What four are found in the atmosphere?CHAPTER LIV.

ORGANIC CHEMISTRY.

295. General Considerations.—Inorganic chemistry is the chemistry of minerals, or unorganized bodies. Organic chemistry was formerly defined as the chemistry of the compounds found in plants and animals; but of late it has taken a much wider range, and is now defined as the chemistry of the C compounds, since C is the nucleus around which other elements centre, and with which they combine to form the organic substances. New organic compounds are constantly being discovered and synthesized, so that nearly 100,000 are now known. The molecule of organic matter is often very complex, sometimes containing hundreds of atoms.