67. An Acid is a substance composed of H and a negative element or radical. It has usually a sour taste, and turns blue litmus red. Litmus is a vegetable extract obtained from lichens in Southern Europe. Acids have the same action on many other vegetable pigments. Are the following acid formulae, and why? H2SO3, HBr, HNO2, H3PO3, H4SiO4. Most acids have O as well as H. Complete the symbols for acids in the following list, and name them, from the type given:—



HCl, chlorhydric acid. HN03, nitric acid. ?Br, ? ?Cl? ? ?I, ? ?Br? ? ?F, ? ?I? ? H3PO4, phosphoric acid. H3PO3, phosphorous acid. ?As? ? ?As? ?

Complete these equations:—

H2SO3 - H2O = ? 2 HN03 - H2O = ? H2SO4 - H2O = ? 2 HNO2 - H2O = ? H2CO3 - H2O = ? 2 H3AsO4 - 3 H2O = ?

Are the products in each case metallic or non-metallic oxides? They are called anhydrides. Notice that each is formed by the withdrawal of water from an acid. Reverse the equations; as, SO3 + H2O = ?

68. An Anhydride is what remains after water has been removed from an acid; or, it is the oxide of a non- metallic element, which, united with water, forms an acid. SO2 is sulphurous anhydride, SO2 sulphuric anhydride, the ending ic meaning more O, or negative element, than ous. Name the others above.

Anhydrides were formerly called acids,—anhydrous acids, in distinction from hydrated ones, as CO2 even now is often called carbonic acid.

Experiment 40.—Hold a piece of wet blue litmus paper in the fumes of SO2, and note the acid test. Try the same with dry litmus paper.

Experiment 41.—Burn a little S in a receiver of air containing 10 cc. H2O, and loosely covered, as in the O experiment. Then shake to dissolve the SO2. H2O + SO2 = H2SO3. Apply test paper.

69. Naming Acids.—Compare formulae H2SO3 and H2SO4. Of two acids having the same elements, the name of the one with least O, or negative element, ends in ous, the other in ic. H2SO3 is sulphurous acid, H2SO4, sulphuric acid. Name H3PO4 and H3PO3; H3AsO3 and H3ASO4; HNO2 and HNO3.

If there are more than two acids in a series, the prefixes hypo, less, and per, more, are used. The following is such a series: HClO, HClO2, HClO3, HClO4.

HClO3 is chloric acid; HClO2, chlorous; HClO, hypochlorous; HClO4 perchloric. Hypo means less of the negative element than ous; per means more of the negative element than ic. Name: H3PO4 (ic), H3PO3, H3PO2. Also HBrO (HBrO2 does not exist), HBrO3 (ic), HBrO4.

What are the three most negative elements? Note their occurrence in the three strongest and most common acids. Hereafter note the names and symbols of all the acids you see.

70. What Bases Are.

Experiment 42.—Put a few drops of NH4OH into an evaporating- dish. Add 5 cc. H2O, and stir. Taste a drop. Dip into it a piece of red litmus paper, noting the effect. Cleanse the dish, and treat in the same way a few drops NaOH solution, recording the result. Do the same with KOH. Acid stains on the clothing, with the exception of those made by HNO3, maybe removed by NH4OH. H2SO4, however, rapidly destroys the fiber of the cloth.

Name two characteristics of a base. In the formulae of those bases, what two common elements? Name the radical. Compare those symbols with the symbol for water, HOH. Is (OH) positive or negative? Is the other part of each formula positive or negative? What are two constituents, then, of a base? Bases are called hydrates. Write in a vertical line five positive elements. Note the valence of each, and complete the formula for its base. Affix the names. Can you see any reason why the three bases above given are the strongest?

Taking the valences of Cr and Fe, write symbols for two sets of hydrates, and name them. Try to recognize and name every base hereafter met with.

A Base is a substance which is composed of a metal, or positive radical, and OH. It generally turns red litmus blue, and often has an acrid taste.

An Alkali is a base which is readily soluble in water. The three principal alkalies are NH4OH, KOH, and NaOH.

Alkali Metals are those which form alkalies. Name three.

An Alkaline Reaction is the turning of red litmus blue.

An Acid Reaction is the turning of blue litmus red.

Experiment 43.—Pour 5 cc. of a solution of litmus in water, into a clean t.t. or small beaker. Pour 2 or 3 cc. of HCl into an evaporating-dish, and the same quantity of NH4OH into another dish. Take a drop of the HCl on a stirring-rod and stir the litmus solution with it. Note the acid reaction. Clean the rod, and with it take a drop (or more if necessary) of NH4OH, and add this to the red litmus solution, noting the alkaline reaction. Experiment in the same way with the two other principal acids and the two other alkalies.

Litmus paper is commonly used to test these reactions, and hereafter whenever the term LITMUS is employed in that sense, the test-paper should be understood. This paper can be prepared by dipping unglazed paper into a strong aqueous solution of litmus.

CHAPTER XVI.

SALTS.

71. Acids and Bases are usually Opposite in Character.—When two forces act in opposition they tend to neutralize each other. We may see an analogy to this in the union of the two opposite classes of compounds, acids and bases, to form salts.

72. Neutralization.

Experiment 44.—Put into an evaporating-dish 5 cc. of NaOH solution. Add HCl to this from a t.t., a few drops at a time, stirring the mixture with a glass rod (Fig. 20), and testing it with litmus paper, until the liquid is neutral, i.e. will not turn the test paper from blue to red, or red to blue. Test with both colors. If it turns blue to red, too much acid has been added; if red to blue, too much base. When it is very nearly neutral, add the reagent, HCl or NaOH, a drop at a time with the stirring-rod. It must be absolutely neutral to both colors. Evaporate the water by heating the dish over asbestus paper, wire gauze, or sand, in an iron plate (Fig. 21) till the residue becomes dry and white. Cool the residue, taste, and name it. The equation is: HCl + NaOH = NaCl + HOH or H2O. Note which elements, positive or negative, change places. Why was the liquid boiled? The residue is a type of a large class of compounds, called salts.

(Fig. 20) (Fig. 21)

Experiment 45. — Experiment in the same way with KOH solution and H2SO4, applying the same tests. H2SO4 + 2 KOH = K2SO4 + 2 HOH. What is the solid product?

Experiment 46.—Neutralize NH4OH with HNO3, evaporate, apply the tests, and write the equation. Write equations for the combination of NaOH and H2SO4; NaOH and HNO3; KOH and HCl; KOH and HNO3; NH4OH and HCl; NH4OH and H2SO4. Describe the experiment represented by each equation, and be sure you can perform it if asked to do so. What is the usual action of a salt on litmus? How is a salt made? What else is formed at the same time? Have all salts a saline taste? Does every salt contain a positive element or radical? A negative?

73. A Salt is the product of the union of a positive and a negative element or radical; it may be made by mixing a base and an acid.

The salt KI represents what acid? What base, or hydrate? Write the equation for making KI from its acid and base. Describe the experiment in full. Classify, as to acids, bases, or salts: KBr, Fe(OH)2, HI, NaBr, HNO2, Al2(OH)6, KClO3, HClO3, H2S, K2S, H2S03, K2SO4, Ca(OH)2, CaCO3, NaBr03, CaSO4, H2CO3, K2CO3, Cu(OH)2, Cu(NO3)2, PbSO4, H3P04, Na2P04. In the SALTS above, draw a light vertical line, separating the positive from the negative part of the symbol. Now state what acid each represents. What base. Write the reaction in the preparation of each salt above from its acid and base; then state the experiment for producing it.

74. Naming Salts.—(NO3) is the nitrate radical; KNO3 is potassium nitrate. From what acid? (NO2) is the nitrite radical; KN02 is potassium nitrite. From what acid? Note that the endings of the acids are OUS and IC; also that the names of their salts end in ITE and ATE. From which acid—IC or OUS—is the salt ending in ATE derived? That ending in ITE?

Name these salts, the acids from which they are derived, and the endings of both acids and salts: NaNO3, NaNO2, K2SO4, K2SO3, CaSO4, CaSO3, KClO3, KClO2, KClO, KClO4 (use prefixes HYPO and PER, as with acids), Ca3(PO4)2, Ca3(P03)2, CuSO4, CuSO3, AgNO3, Cu(NO3)2. FeS, FeS2, are respectively FERROUS SULPHIDE and FERRIC SULPHIDE. Name: HgCl, HgCl2, FeCl2, Fe2Cl6, FeSO4, Fe2(SO4)3.75. Acid Salts.—Write symbols for nitric, sulphuric, phosphoric acids. How many H atoms in each? Replace all the H in the symbol of each with Na, and name the products. Again, in sulphuric acid replace one atom of H with Na; then in phosphoric replace first one, then two, and finally three H atoms with Na. HNaSO4 is hydrogen sodium sulphate; HNa2P04 is hydrogen di-sodium phosphate. Name the other salts symbolized. Name HNaNH4P04. Though these products are all salts, some contain replaceable H, and are called acid salts. Those which have all the H replaced by a metal are normal salts. Name and classify, as to normal or acid salts: Na2CO3, HNaCO3, K2SO4, HKSO4, (NH4)2SO4, HNH4SO4, Na3P04, HNa2P04, H2NaP04.

The BASICITY of an acid is determined by the number of replaceable H atoms in its molecule. It is called MONOBASIC if it has one; DIBASIC if two; TRI- if three, etc. Note the basicity of each acid named above. How many possible salts of H2SO4 with Na? Of H3P04 with Na? Which are normal and which acid? What is the basicity of H4Si04?

Some normal, as well as acid, salts change litmus. Na2CO3, representing a strong base and a weak acid, turns it blue. There are other modes of obtaining salts, but this is the only one which we sball consider.

76. Salts Occur Abundantly in Nature, such as NaCl, MgSO4, CaCO3. Acids and bases are found in small quantities only. Why is this? Why are there not springs of H2SO4 and NH4OH? We have seen that acids and bases are extremely active, have opposite characters, and combine to form relatively inactive salts. If they existed in the free state, they would soon combine by reason of their strong affinities. This is what in all ages of the world has taken place, and this is why salts are common, acids and bases rare. Active agents rarely exist in the free state in large quantities. Oxygen seems to be an exception, but this is because there is a superabundance of it. While vast quantities are locked up in compounds in rocks, water, and salts of the earth, much remains with which there is nothing to combine.

CHAPTER XVII.

CHLORHYDRIC ACID.

77. We have seen that salts are made by the union of acids and bases. Can these last be obtained from salts?

78. Preparation of HCl.

Experiment 47.—Into a flask put 10 g. coarse NaCl, and add 20 cc. H2SO4. Connect with Woulff bottles [Woulff bottles may be made by fitting to wide-mouthed bottles corks with three holes, through which pass two delivery tubes, and a central safety tube dipping into the liquid, as in Figures 22 and 23.] partly filled with water, as in Figure 22. One bottle is enough to collect the HCl; but in that case it is less pure, since some H2SO4 and other impurities are carried over. Several may be connected, as in Figure 23. The water in the first bottle must be nearly saturated before much gas will pass into the second. Heat the mixture 15 or 20 minutes, not very strongly, to prevent too much foaming. Notice any current in the first bottle. NaCl + H2SO4 = HNaSO4 + HCl. Intense heat would have given: 2NaCl + H2SO4 = Na2SO4 + 2HCl. Compare these equations with those for HNO3. In which equation above is H2SO4 used most economically? Both reactions take place when HCl is made on the large scale.

(Fig. 22)

79. Tests. Experiment 48.—(1) Test with litmus the liquid in each Woulffbottle. (2) Put a piece of Zn into a t.t. and cover it with liquid from the first bottle. Write the reaction, and test the gas. (3) To 2 cc.solution AgNO3 in a t.t. add 2 cc.of the acid. Describe, and write the reaction. Is AgCl soluble in water? (4) Into a t.t. pour 5 cc.Pb(NO3)2 solution, and add the same amount of prepared acid. Give the description and the reaction. (5) In the same way test the acid with Hg2(NO3)2 solution, giving the reaction. (6) Drake a little HCl in a t.t., and bring the gas escaping from the d.t. in contact with a burning stick. Does it support the combustion of C? (7) Hold a piece of dry litmus paper against it. [figure 23] (8) Hold it over 2 cc.of NH4OH in an evaporating-dish. Describe, name the product, and write the reaction. (3), (4), (5), (8), are characteristic tests for this acid.

80. Chlorhydric, Hydrochloric or Muriatic, Acid is a Gas.—As used, it is dissolved, in water, for which it has great affinity. Water will hold, according to temperature, from 400 to 500 times its volume of HCl. Hundreds of thousands of tons of the acid are annually made, mostly in Europe, as a bye-product in Na2CO3 manufacture. The gas is passed into towers through which a spray of water falls; this absorbs it. The yellow color in most commercial HCl indicates impurities, some of which are Fe, S, As, and organic matter. As, S, etc., come from the pyrites used in making H2SO4. Chemically pure (C.P.) acid is freed from these, and is without color. The gas may be dried by passing it through a glass tube holding CaCl2 (Fig. 16) and collecting it over mercury.

The muriatic acid of commerce consists of about two- thirds water by weight. HCl can also be made by direct union of its constituents.81. Uses.—HCl is used to make Cl, and also bleaching- powder. Its use as a reagent in the laboratory is illustrated by the following experiment:— Experiment 49.—Put into a t.t. 2 cc. AgNO3 solution, add 5 cc. H2O, then add slowly HCl so long as a ppt. (precipitate) is formed. This ppt. is AgCl. Now in another t.t. put 2 cc. Cu(NO3)2, solution, add 5 cc. H2O, then a little HCl. No ppt. is formed. Now if a solution of AgNO3 and a solution of Cu(NO3)2 were mixed, and HCl added, it is evident that the silver would be precipitated as chloride of silver, while the copper would remain in solution. If now this be filtered, the silver will remain on the filter paper, while in the filtrate will be the copper. Thus we shall have performed an analysis, or separated one metal from another. Perform it. Note, however, that any soluble chloride, as NaCl, would produce the same result as HCl.

BROMHYDRIC AND IODIHYDRIC ACIDS.

82. NaCl, being the most abundant compound of Cl, is the source of commercial HCl. KCl treated in the same way would give a like product. Theoretically HBr and HI might be made in the same way from NaBr and NaI, but the affinity of H for Br and I is weak, and the acids separate into their elements, when thus prepared.

83. To make HI.

Experiment 50.—Drop into a t.t. three or four crystals of I, and add 10 cc. H2O. Hold in the water the end of a d.t. from which H2S gas is escaping. Observe any deposit, and write the reaction.

FLUORHYDRIC ACID.

84. Preparation and Action.

Experiment 51.—Put 3 or 4 g. powdered CaF2, i.e. fluor spar or fluorite, into a shallow lead tray, e.g. 4x5 cm, and pour over it 4 or 5 cc. H2SO4. A piece of glass large enough to cover this should previously be warmed and covered on one side with a very thin coat of beeswax. To distribute itevenly, warm the other side of the glass over a flame. When cool, scratch a design (Fig. 24) through the wax with a sharp metallic point. Lay the glass, film side down, over the lead tray. Warm this five minutes or more by placing it high over a small flame (Fig. 25) to avoid melting the wax. Do not inhale the fumes. Take away the lamp, and leave the tray and glass where it is not cold, for half an hour or more. Then remove the wax and clean the glass with naphtha or benzine. Look for the etching.

Two things should have occurred: (1) the generation of HF. Write the equation for it. (2) Its etching action on glass. In this last process HF acts on SiO2 of the glass, forming H2O and SiF4. Why cannot HF be kept in glass bottles?

A dilute solution of HF, which is a gas, may be kept in gutta percha bottles, the anhydrous acid in platinum only; but for the most part, it is used as soon as made, its chief use being to etch designs on glass-ware. Glass is also often etched by a blast of sand (SiO2).

Notice the absence of O in the acids HF, HCI, HBr, HI, and that each is a gas. HF is the only acid that will dissolve or act appreciably on glass.

Chapter XVIII.

NITRIC ACID.

85. Preparation. Experiment 52.—To 10 g. KNO3 or NaNO3, in a flask, add 15 cc. H2SO4. Securely fasten the cork of the d.t., as HNO3 is likely to loosen it, and pass the other end to the bottom of a t.t. held deep in a bottle of water (Fig. 26). Apply heat, and collect 4 or 5 cc.of the liquid. The usual reaction is: KNO3 + H2SO4 = HKSO4 + HNO3. With greater heat, 2 KNO3 + H2SO4 = K2SO4 + 2HNO3. Which is most economical of KNO3? Of H2SO4? Instead of a flask, a t.t. may be used if desired (Fig. 27).

86. Properties and Tests.

Experiment 53.—(1) Note the color of the prepared liquid. (2) Put a drop on the finger; then wash it off at once. (3) Dip a quill or piece of white silk into it; then wash off the acid. What color is imparted to animal substances? (4) Add a little to a few bits of Cu turnings, or to a Cu coin. Write the equation. (5) To 2 cc.indigo solution, add 2 cc. HNO3. State the leading properties of HNO3, from these tests.

87. Chemically Pure HNO3 is a Colorless Liquid.— The yellow color of that prepared in Experiment 52 is due to liquid NO2 dissolved in it. It is then called fuming HNO3, and is very strong. NO2 is formed at a high temperature.

Commercial or ordinary HNO3, is made from NaNO3, this being cheaper than KNO3; it is about half water.

88. Uses. HNO3 is the basis of many nitrates, as AgNO3, used for photography, Ba(NO3)2 and Sr(NO3)2 for fire-works, and others for dyeing and printing calico; it is employed in making aqua regia, sulphuric acid, nitro-glycerine, gun-cotton, aniline colors, zylonite, etc.

Enough experiments have been performed to answer the question whether some acids can be prepared from their salts. H2SO4 is not so made, because no acid is strong enough to act on its salts. In making HCl, HNO3, etc., sulphuric acid was used, being the strongest.

AQUA REGIA.

89. Preparation and Action. Experiment 54.—Into a t.t. put 2 cc. HNO3, and 14 qcm. of either Au leaf or Pt. Warm in a flame. If the metal is pure, no action takes place. Into another tube put 6 cc. HCl and add a similar leaf. Heat this also. There should be no action. Pour the contents of one t.t. into the other. Note the effect. Which is stronger, one of the acids, or the combination of the two? Note the odor. It is that of Cl. 3HCl + HNO3 = NOCl + 2H2O + Cl2. This reaction is approximate only. The strength is owing to nascent chlorine, which unites with Au. Au + 3Cl = AuCl3. If Pt be used, PtCl4 is produced. No other acid except nitro-hydrochloric will dissolve Au or Pt; hence the ancients called it aqua regia, or king of liquids. It must be made as wanted, since it cannot be kept and retain its strength.

CHAPTER XIX.

SULPHURIC ACID.

90. Preparation.

Experiment 55.—Having fitted a cork with four or five perforations to a large t.t., pass a d.t. from three of these to three smaller t.t., leaving the others open to the air, as in Figure 28. Into one t.t. put 5 cc. H2O, into another 5 g. Cu turnings and 10 cc. H2SO4, into the third 5 g. Cu turnings and 10 cc. dilute HNO3, half water. Hang on a ring stand, and slowly heat the tubes containing H2O and H2SO4. Notice the fumes that pass into the large t.t.

Trace out and apply to Figure 28 these reactions:—

(1) Cu + 2 H2SO4 = CuSO4 + 2 H2O + SO2.

(2) 3 Cu + 8 HNO3 = 3 Cu(NO3)2+ 4 H2O + 2 NO.

(3) NO + O = NO2.

(4) SO2 + H2O + NO2 =H2SO4 + NO.

(4) comes from combining the gaseous products in (1), (2), (3). In (3), NO takes an atom of O from the air, becoming NO2, and at once gives it up, to the H2SO3 (H2O + SO2), making H2SO4, and again goes through the same operation of taking up O and passing it along. NO is thus called a carrier of O. It is a reducing agent, while NO2 is an oxidizing agent. This is a continuous process, and very important, since it changes useless H2SO3 into valuable H2SO4. If exposed to the air, H2SO3 would very slowly take up O and become H2SO4.

Instead of the last experiment, this may be employed if preferred: Burn a little S in a receiver. Put into an evaporating-dish, 5 cc. HNO3, and dip a paper or piece of cloth into it. Hang the paper in the receiver of SO2, letting no HNO3 drop from it. Continue this operation till a small quantity of liquid is found in the bottle. The fumes show that HNO3 has lost O. 2 HNO3 + SO2 = H2SO4 + 2 NO2.

91. Tests for H2SO4.

Experiment 56.—(1) Test the liquid with litmus. (2) Transfer it to a t.t., and add an equal volume of BaCl2 solution. H2SO4 + BaCl2 = ? Is BaSO4 soluble? (3) Put one drop H2SO4 from the reagent bottle in 10 cc. H2O in a clean t.t., and add 1 cc. BaCl2 solution. Look for any cloudiness. This is the characteristic test for H2SO4 and soluble sulphates, and so delicate that one drop in a liter of H2O can be detected. (4) Instead of H2SO4, try a little Na2SO4 solution. (5) Put two or three drops of strong H2SO4 on writing-paper, and evaporate, high over a flame, so as not to burn the paper. Examine it when dry. (6) Put a stick into a t.t. containing 2 cc. H2SO4, and note the effect. (7) Review Experiment 5. (8) Into an e.d. pour 5 cc. H2O, and then 15 cc. H2SO4. Stir it meantime with a small t.t. containing 2 or 3 cc. NH4OH, and notice what takes place in the latter; also note the heat of the e.d.

The effects of (5), (6), (7), and (8) are due to the intense affinity which H2SO4 has for H2O. So thirsty is it that it even abstracts H and O from oxalic acid in the right proportion to form H2O, combines them, and then absorbs the water.

92. Affinity for Water.—This acid is a desiccator or dryer, and is used to take moisture from the air and prevent metallic substances from rusting. In this way it dilutes itself, and may increase its weight threefold. In diluting, the acid must always be poured into the water slowly and with stirring, not water into the acid, since, as H2O is lighter than H2SO4, heat enough may be set free at the surface of contact to cause an explosion. Contraction also takes place, as may be shown by accurately measuring each liquid in a graduate, before mixing, and again when cold. The mixture occupies less volume than the sum of the two volumes. For the best results the volume of the acid should be about three times that of the water.

93. Sulphuric Acid made on a Large Scale involves the same principles as shown in Experiment 55, excepting that S02 is obtained by burning S or roasting FeS2 (pyrite),

[Fig. 29.]

and HNO3 is made on the spot from NaNO3 and H2SO4. SO2 enters a large leaden chamber, often 100 to 300 feet long, and jets of steam and small portions of HNO3 are also forced in. The "chamber acid" thus formed is very dilute, and must be evaporated first in leaden pans, and finally in glass or platinum retorts, since strong H2SO4, especially if hot, dissolves lead. See Experiment 124. Study Figure 29, and write the reactions. 2 HNO3 breaks up into 2 NO2, H2O, and O. 94. Importance.—Sulphuric acid has been called, next to human food, the most indispensable article known. There is hardly a product of modern civilization in the manufacture of which it is not directly or indirectly used. Nearly a million tons are made yearly in Great Britain alone. It is the basis of all acids, as Na2CO3 is of alkalies. It is the life of chemical industry, and the quantity of it consumed is an index of a people's civilization. Only a few of its uses can be stated here. The two leading ones are the reduction of Ca3(PO4)2 for artificial manures and the sodium carbonate manufacture. Foods depend on the productiveness of soils and on fertilizers, and thus indirectly our daily bread is supplied by means of this acid; and from sodium carbonate glass, soap, saleratus, baking- powders, and most alkalies are made directly or indirectly. H2SO4 is employed in bleaching, dyeing, printing, telegraphy, electroplating, galvanizing iron and wire, cleaning metals, refining Au and Ag, making alum, blacking, vitriols, glucose, mineral waters, ether, indigo, madder, nitroglycerine, gun- cotton, parchment, celluloid, etc., etc.

FUMING SULPHURIC ACID.

95. Nordhausen or Fuming Sulphuric Acid, H2S207 used in dissolving indigo and preparing coal-tar pigments, is made by distilling FeSO4. 4FeSO4 + H2O = H2S207 + 2Fe203 + 2S02. This was the original sulphuric acid. It is also formed when S03 is dissolved in H2SO4. When exposed to the air, S03 escapes with fuming.

CHAPTER XX.

AMMONIUM HYDRATE.

96. Preparation of Bases.—We have seen that many acids are made by acting on a salt of the acid required, with a stronger acid. This is the direct way. The following experiments will show that bases may be prepared in a similar way by acting on salts of the base required with other bases, which we may regard as stronger than the ones to be obtained.

97. Preparation of NH4OH and NH3.

Experiment 57.—Powder 10 g. ammonium chloride, NH4Cl, in a mortar and mix with 10 g. calcium hydrate, Ca(OH)2; recently slaked lime is the best. Cover with water in a flask, and connect with Woulff bottles, as for making HCl (Fig. 22); heat the flask for fifteen minutes or more. The experiment may be tried on a smaller scale with a t.t. if desired.

The reaction is: 2NH4Cl + Ca(OH)2 = CaCl2 + 2NH4OH. NH4OH is broken up into NH3, ammonia gas, and water. NH4OH = NH3 + H2O. These pass over into the first bottle, where the water takes up the NH3, for which it has great affinity. One volume of water at 0 will absorb more than 1000 volumes of NH3. Thus NH4OH may be called a solution of NH3, in H2O. Write the reaction.

Experiment 58.—Powder and mix 2 or 3 g. each of ammonium nitrate, NH4NO3, and Ca(OH)2; put them into a t.t., and heat slowly. Note the odor. 2NH4NO3 + Ca(OH)2 = ?

98. Tests.

Experiment 59.—(1) Generate a little of the gas in a t.t., and note the odor. (2) Test the gas with wet red litmus paper. (3) Put a little HCl into an e.d., and pass over it the fumes of NH3 from a d.t. Note the result, and write the equation. (4) Fill a small t.t. with the gas by upward displacement; then, while still inverted, put the mouth of the t.t. into water. Explain the rise of the water. (5) How might NH4Cl be obtained from the NH4OH in the Woulff bottles? (6) Test the liquid in each bottle with red litmus paper. (7) Add some from the first bottle to 5 or 10 cc. of a solution of FeSO4 or FeCl2, and look for a ppt. State the reaction.

99. Formation.—Ammonia, hartshorn, exists in animal and vegetable compounds, in salts, and, in small quantities, in the atmosphere. Rain washes it from the atmosphere into the soil; plants take it from the soil; animals extract it from plants. Coal, bones, horns, etc., are the chief sources of it, and from them it is obtained by distillation. It results also from decomposing animal matter. NH3 can be produced by the direct union of N and H, only by an electric discharge or by ozone. It may be collected over Hg like other gases that are very soluble in water.

100. Uses. —Ammonium hydrate, NH4OH, and ammonia, NH3, are used in chemical operations, in making artificial ice, and to some extent in medicine; from them also may be obtained ammonium salts. State what you would put with NH4OH to obtain (NH4)2SO4. To obtain NH4NO3. The use of NH4OH in the laboratory may be illustrated by the following experiment:—

Experiment 60.—Into a t.t. put 10 cc. of a solution of ferrous sulphate, FeSO4. Into another put 10 cc. of sodium sulphate solution, Na2SO4. Add a little NH4OH to each. Notice a ppt. in the one case but none in the other. If solutions of these two compounds were mixed, the metals Fe and Na could be separated by the addition of NH4OH, similar to the separation of Ag and Cu by HCl. Try the experiment.

CHAPTER XXI.

SODIUM HYDRATE.

101. Preparation.

Experiment 61.—Dissolve 3 g. sodium carbonate, Na2CO3, in 10 or 15 cc. H2O in an e.d., and bring it to the boiling-point. Then add to this a mixture of 1 or 2 g. calcium hydrate, Ca(OH)2, in 5 or 10cc. H2O. It will not dissolve. Boil the whole for five minutes. Then pour off the liquid which holds NaOH in solution. Evaporate if desired. This is the usual mode of preparing NaOH.

The reaction is Na2CO3 + Ca(OH)2 = 2NaOH + CaCO3. The residue is Ca(OH)2 and CaCO3; the solution contains NaOH, which can be solidified by evaporating the water. Sodium hydrate is an ingredient in the manufacture of hard soap, and for this use thousands of tons are made annually, mostly in Europe. It is an important laboratory reagent, its use being similar to that of ammonium hydrate. Exposed to the air, it takes up water and CO2, forming a mixture of NaOH and Na2CO3. It is one of the strongest alkalies, and corrodes the skin.

Experiment 62.—Put 20 cc. of H2O in a receiver. With the forceps take a piece of Na, not larger than half a pea, from the naphtha in which it is kept, drop it into the H2O, and at once cover the receiver loosely with paper or cardboard. Watch the action, as the Na decomposes H2O. HOH + Na = NaOH + H. If the water be hot the action is so rapid that enough heat is produced to set the H on fire. That the gas is H can be shown by putting the Na under the mouth of a small inverted t.t., filled with cold water, in a water-pan. Na rises to the top, and the t.t. fills with H, which can be tested. NaOH dissolves in the water.102. Properties.

Experiment 63.—(1) Test with red litmus paper the solutions obtained in the last two experiments. (2) To 5cc.of alum solution, K2A12(SO4)4, add 2cc.of the liquid, and notice the color and form of the ppt.

POTASSIUM HYDRATE.

103. KOH is made in the Same Way as NaOH.

Describe the process in full (Experiment 61), and give the equation.

Experiment 64.—Drop a small piece of K into a receiver of H2O, as in Experiment 62. The K must be very small, and the experiment should not be watched at too close a range. The receiver should not be covered with glass, but with paper. The H burns, uniting with O of the air. The purple color is imparted by the burning, or oxidation of small particles of K. Write the equation for the combustion of each.

H2O might be considered the symbol of an acid, since it is the union of H and a negative element; or, if written HOH, it might be called a base, since it has a positive element and the (OH) radical. It is neutral to litmus, and on this account might be called a salt. It is better, however, to call it simply an oxide.

Potassium hydrate, caustic potash, is employed for the manufacture of soft soap. As a chemical reagent its action is almost precisely like that of caustic soda, though it is usually considered a stronger base, as K is a more electro-positive element than Na.

CALCIUM HYDRATE.

104. Calcium Hydrate, the Most Common of the Bases, is nearly as important to them as H2SO4 is to acids. Since it is used to make the other bases, it might be called the strongest base; as H2SO4 is often called the strongest acid. The strength of an acid or base, however,depends on the substance to which it is applied, as well as on itself, and for most purposes this one is classified as a weaker base than the three previously described.

Sulphuric acid, the most useful of the acids, is not made directly from its salts, but has to be synthesized. Calcium hydrate is also made by an indirect process, as follows:

CaCO3, i.e. limestone, marble, etc., is burnt in kilns with C, a process which separates the gas, CO2, according to the reaction: CaCO3 = CaO + CO2. CaO is unslaked lime, or quick-lime. On treating this with water, slaked lime, Ca(OH)2 is formed, with generation of great heat. CaO + H2O = Ca(OH)2. Its affinity for H2O is so great that it takes the latter from the air, if exposed.

Experiment 65.—Saturate some unslaked lime with water, in an e.d., and look for the results stated above, leaving it as long as may be necessary.

105. Resume.—From the experiments in the last few chapters on the three divisions of chemical compounds, acids, bases and salts, we have seen (1) that acids and bases are the chemical opposites of each other; (2) that salts are formed by the union of acids and bases; (3) that some acids can be obtained from their salts by the action of a stronger acid; (4) that some bases can be got from salts by the similar action of other bases; (5) that the strongest acids and bases, as well as others, may be obtained in an indirect way by synthesis.

CHAPTER XXII.

OXIDES OF NITROGEN.

106. There are five oxides of N, only two of which are important.

NITROGEN MONOXIDE (N2O).

107. Preparation.

Experiment 66.—Put into a flask, holding 200cc, lOg of ammonium nitrate, NH4NO3; heat it over wire gauze or asbestus in an iron plate, having a d.t. connected with a large t.t., which is held in a receiver of water, and from this t.t., another d.t. passing into a pneumatic trough, so as to collect the gas over water (Fig. 30). Have all the bearings tight. The reaction is NH4NO3 = 2H2O + N2O. The t.t. is for collecting the H2O.

[Fig. 30.]

Note the color of the liquid in the t.t.; taste a drop, and test it with litmus. If the flask is heated too fast, some NO is formed, and this taking O from the air makes NO2, which liquefies and gives an acid reaction and a red color. Some NH4NO3 is also liable to be carried over.

108. Properties.

Experiment 67.—Test the gas in the receiver with a burning stick and a glowing one, and compare the combustion with that in O. N20may also be tested with S and P, if desired. N is set free in each case. Write the reactions.

Nitrogen monoxide or protoxide, the nitrous oxide of dentists, when inhaled, produces insensibility to pain,— anaesthesia,— and, if continued, death from suffocation. Birds die in half a minute from breathing it. Mixed with one-fourth O, and inhaled for a minute or two, it produces intoxication and laughter, and hence is called laughing gas. As made in Experiment 66, it contains Cl and NO, as impurities, and should not be breathed.

NITROGEN DIOXIDE (NO, OR N2O2).

109. Preparation.

Experiment 68.—Into a t.t. or receiver put 5g Cu turnings, add 5 cc. H2O and 5 cc. HNO3. Collect the gas like H, over water. 3Cu + 8HNO3 = ? What two products will be left in the generator? Notice the color of the liquid. This color is characteristic of Cu salts. Notice also the red fumes of NO2.

110. Properties.

Experiment 69.—Test the gas with a burning stick, admitting as little air as possible. Test it with burning S. NO is not a supporter of C and S combustion. Put a small bit of P in a deflagrating-spoon, and when it is vigorously burning, lower it into the gas. It should continue to burn. State the reaction. What combustion will NO support? Note that NO is half N, while N2O is two-thirds N, and account for the difference in supporting combustion.

NITROGEN TETROXIDE (NO2 or N2O4).

111. Preparation.

Experiment 70.—Lift from the water-pan a receiver of NO, and note the colored fumes. They are NO2, or N2O4, nitrogen tetroxide. NO + O = NO2. Is NO combustible? What is the source of O in the experiment?OXIDES OF NITROGEN.

NITROGEN TRIOXIDE (N2O3).

112. Preparation.

Experiment 71.—Put into a t.t. 1 g. of starch and 1 cc. of HNO3. Heat the mixture for a minute. The red fumes are N2O3 and NO2.

Nitrogen pentoxide, N2O5, is an unimportant solid. United with water it forms HNO3. N2O5 + H2O = 2HNO3.

CHAPTER XXIII.

LAWS OF DEFINITE AND OF MULTIPLE PROPORTION.

113. Weight and Volume.—We have seen that water contains two parts of H by volume to one part of O; or, by weight, two parts of H to sixteen of O. These proportions are invariable, or no symbol for water would be possible. Every compound in the same way has an unvarying proportion of elements.

114. Law of Definite Proportion.—In a given compound the proportion of any element by weight, or, if a gas, by volume is always constant. Apply the law, by weight and by volume, to these: HCl, NH3, H2S, N2O.

There is another law of equal importance in chemistry, which the compounds of N and O well illustrate.

Weight. Volume. N. O. N. O. Nitrogen protoxide N2O 28 16 2 1 Nitrogen dioxide N2O2 28 32 2 2 Nitrogen trioxide. N2O3 28 48 2 3 Nitrogen tetroxide N2O4 28 64 2 4 Nitrogen pentoxide N2O5 28 80 2 5

Note that the proportion of O by weight is in each case a multiple of the first, 16. Also that the proportion by volume of O is a multiple of that in the first compound. In this example the N remains the same. If that had varied in the different compounds, it would also havevaried by a multiple of the smallest proportion. This is true in all compounds.

115. Law of Multiple Proportion.—Whenever one element combines with another in more than one proportion, it always combines in some multiple, one or more, of its least combining weight, or, if a gas, of its least combining volume.

The least combining weight of an element is its atomic weight; and it is this fact of a least combining weight that leads us to believe the atom to be indivisible.

Apply the law in the case of P2O, P2O3, P2O5; in HClO, HClO2, HClO3, HClO4, arranging the symbols, weights, and volumes in a table, as above.

The volumetric proportions of each element in the oxides of nitrogen are exhibited below.

+ + = N + N + O = N2O

_ + _ + _ + _ = _ N + N + O + O = N2O2

+ + + + = N + N + O + O + O = N2O3

_ + _ + _ + _ + _ + _ = _ N + N + O + O + O + O = N2O4

+ + + + + + = N + N + O + O + O + O + O = N2O5

CHAPTER XXIV.

CARBON PROTOXIDE.

116. Preparation.

Experiment 72.—Put into a flask, of 200 cc., 5 g. of oxalic acid crystals, H2C2O4, and 25 cc. H2SO4. Have the d.t. pass into a solution of NaOH in a Woulff bottle (Fig. 31), and collect the gas over water. Heat the flask slowly, and avoid inhaling the gas.

117. Tests.

Experiment 73.—Remove a receiver of the gas, and try to light the latter with a splinter. Is it combustible, or a supporter of (C) combustion? What is the color of the flame? When the combustion ceases, shake up a little lime water with the gas left in the receiver. What gas has been formed by the combustion, as shown by the test? See page 80. Give the reaction for the combustion.

We have seen that H2SO4 has great affinity for H2O. Oxalic acid consists of H, C, O in the right proportion to form H2O, CO2, and CO. H2SO4 withdraws H and O in the right proportion to form water, unites them, and then absorbs the water, leaving the C and O to combine and form CO2 and CO. NaOH solution removes CO2 from the mixture, forming Na2CO3, and leaves CO. Write both reactions.

118. Carbon Protoxide, called also carbon monoxide, carbonic oxide, etc., is a gas, having no color or taste, butpossessing a faint odor. It is very poisonous. Being the lesser oxide of C, it is formed when C is burned in a limited supply of O, whereas CO2 is always produced when O is abundant. The formation of each is well shown by tracing the combustion in a coal fire. Air enters at the bottom, and CO2 is first formed. C + 2O = CO2. As this gas passes up, the white-hot coal removes one atom of O, leaving CO. CO2 + C - 2CO. At the top, if the draft be open, a blue flame shows the combustion of CO. CO + O = CO2. The same reduction of CO2 takes place in the iron furnace, and whenever there is not enough oxygen to form CO2, the product is CO.

Great care should be taken that this gas does not escape into the room, as one per cent has proved fatal. Not all of it is burned at the top of the coal; and when the stove door is open, the upper drafts should be open also. It is the most poisonous of the gases from coal; hence the danger from sleeping in a room having a coal fire.

119. Water Gas.—CO is one of the constituents of "water gas," which, by reason of its cheapness, is supplanting gas made from coal, as an illuminator, in some cities. It is made by passing superheated steam over red-hot charcoal or coke. C unites with the O of H2O, forming CO, and sets H free, thus producing two inflammable gases. C + H2O —? As neither of these gives much light, naphtha is distilled and mixed with them in small quantities to furnish illuminating power See page 183.

CHAPTER XXV.

CARBON DIOXIDE.

120. Preparation.

Experiment 74.—Put into a t.t., or a bottle with a d.t. and a thistle-tube, 10 or 20 g. CaCO3, marble in lumps; add as many cubic centimeters of H2O, and half as much HCl, and collect the gas by downward displacement (Fig. 39). Add more acid as needed. CaCO3 + 2 HCl = CaCl2 + H2CO3. H2CO3 = H2O + CO2. H2CO3 is a very weak compound, and at once breaks up. By some, its existence as a compound is doubted.

121. Tests.

Experiment 75.—(1) Put a burning and a glowing stick into the t.t. or bottle. (2) Hold the end of the d.t. directly against the flame of a small burning stick. Does the gas support combustion? (3) Pour a receiver of the gas over a candle flame. What does this show of the weight of the gas? (4) Pass a little CO2 into some H2O (Fig. 32), and test it with litmus. Give the reaction for the solution of CO2 in H2O.

Experiment 76.—Put into a t.t. 51 cc. of clear Ca(OH)2 solution, i.e. lime water; insert in this the end of a d.t. from a CO2 generator (Fig. 32). Notice any ppt. formed. It is CaCO3. Let the action continue until the ppt. disappears and the liquid is clear. Then remove the d.t., boil the clear liquid for a minute, and notice whether the ppt. reappears.

122. Explanation.

Ca(OH)2 + CO2 = CaCO3 + H2O. The curious phenomena of this experiment are explained by the solubility of CaCO3 in water containing CO2, and its insolu-bility in water, having no CO2. When all the Ca(OH)3 is combined, or changed to CaCO3, the excess of CO2 unites with H2O, forming the weak acid H2CO3, which dissolves the precipitate, CaCO3, and gives a clear liquid. On heating this, H2CO3 gives up its CO2, and CaCO3 is reprecipitated, not being soluble in pure water.

Lime water, Ca(OH)2 solution, is therefore a test for the presence of CO2. To show that carbon dioxide is formed in breathing, and in the combustion of C, and that it is present in the air, perform the following experiment:

Experiment 77.—(1) Put a little lime water into a t.t., and blow into it through a piece of glass tubing. Any turbidity shows what? (2) Burn a candle for a few minutes in a receiver of air, then take out the candle and shake up lime water with the gas. (3) Expose some lime water in an e.d. to the air for some time.

133. Oxidation in the Human System.—Carbon dioxide, or carbonic anhydride, carbonic acid, etc., CO2, is a heavy gas, without color or odor. It has a sharp, prickly taste, and is commonly reckoned as poisonous if inhaled in large quantities, though it does not chemically combine with the blood as CO does. Ten per cent in the air will sometimes produce death, and five per cent produces drowsiness. It exists in minute portions in the atmosphere, and often accumulates at the bottom of old wells and caverns, owing to its slow diffusive power. Before going down into one of these, the air should always be tested by lowering a lighted candle. If this is extinguished, there is danger. CO2 is the deadly "choke damp" after a mine explosion, CH4 being converted into CO2 and H2O; a great deal is liberated during volcanic eruptions, and it is formed in breathing by the union of O in the air with C in the system. This union of C and O takes place in the lungs and in all the tissues of the body, even on the surface. Oxygen is taken into the lungs, passes through the thin membrane into the blood, forms a weak chemical union with the red corpuscles, and is conveyed by them to all parts of the system. Throughout the body, wherever necessary, C and H are supplied for the O, and unite with it to form CO2 and H2O. These are taken up by the blood though they do not form a chemical union with it, are carried to the lungs, and pass out, together with the unused N and surplus O. The system is thus purified, and the waste must be supplied by food. The process also keeps up the heat of the body as really as the combustion of C or P in O produces heat. The temperature of the body does not vary much from 99 degrees F., any excess of heat passing off through perspiration, and being changed into other forms of energy.

If, as in some fevers, the temperature rises above about 105 degrees F., the blood corpuscles are killed, and the person dies. During violent exercise much material is consumed, circulation is rapid, and quick breathing ensues. Oxygen is necessary for life. A healthy person inhales plentifully; and this element is one of nature's best remedies for disease. Deep and continued inhalations in cold weather are better than furnace fires to heat the system. All animals breathe O and exhale CO2. Fishes and other aquatic animals obtain it, not by decomposing H2O, but from air dissolved in water. Being cold-blooded, they need relatively little; but if no fresh water is supplied to those in captivity, they soon die of O starvation.

124. Oxidation in Water.—Swift-running streams are clear and comparatively pure, because their organic impurities are constantly brought to the surface and oxidized, whereas in stagnant pools these impurities accumulate. Reservoirs of water for city supply have sometimes been freed from impurities by aeration, i.e. by forcing air into the water.

125. Deoxidation in Plants.—Since CO2 is so constantly poured into the atmosphere, why does it not accumulate there in large quantity? Why is there not less free O in the air to-day than there was a thousand years ago? The answer to these questions is found in the growth of vegetation. In the leaf of every plant are thousands of little chemical laboratories; CO2 diffused in small quantities in the air passes, together with a very little H2O, into the leaf, usually from its under side, and is decomposed by the radiant energy of the sun. The C is built into the woody fiber of the tree, and the O is ready to be re-breathed or burned again. CO2 contributes to the growth of plants, O to that of animals; and the constituents of the atmosphere vary little from one age to another. The compensation of nature is here well shown. Plants feed upon what animals discard, transforming it into material for the sustenance of the latter, while animals prepare food for plants. All the C in plants is supposed to come from the CO2 in the atmosphere. Animals obtain their supply from plants. The utility of the small percentage of CO2 in the air is thus seen.

126. Uses.—CO2 is used in making "soda-water," and in chemical engines to put out fires in their early stages. In either case it may be prepared by treating Na2CO3 or CaCO3 with H2SO4. Give the reactions. On a small scale CO2 is made from HNaCO3. CO2 has a very weak affinity for water, but probably forms with it H2CO3. Much carbon dioxide can be forced into water under pressure. This forms soda-water, which really contains no soda. The justification for the name is the material from which it is sometimes made. Salts from H2CO3, called carbonates, are numerous, Na2CO3 and CaCO3 being the most important.

Chapter XXVI.

OZONE.

127. Preparation.

Experiment 78.—Scrape off the oxide from the surface of a piece of phosphorus 2 cm long, put it into a wide-mouthed bottle, half cover the P with water, cover the bottle with a glass, and leave it for half an hour or more.

128. Tests.

Experiment 79.—Remove the glass cover, smell the gas, and hold in it some wet iodo-starch paper. Look for any blue color. Iodine has been set free, according to the reaction, 2 KI + 03= K20 + O2 + I2, and has imparted a blue color to the starch, and ordinary oxygen has been formed. Why will not oxygen set iodine free from KI?. What besides ozone will liberate it?

129. Ozone, oxidized oxygen, active oxygen, etc., is an allotropic form of O. Its molecule is 03, while that of ordinary oxygen is 02.

Three atoms of oxygen are condensed into the space of two atoms of ozone, or three molecules of O are condensed into two molecules of ozone, or three liters of O are condensed into two liters of ozone. Ozone is thus formed by oxidizing ordinary oxygen. 02 + O = 03. This takes place during thunder storms and in artificial electrical discharges. The quantity of ozone produced is small, five per cent being the maximum, and the usual quantity is far less than that.

Ozone is a powerful oxidizing agent, and will change S, P, and As into their ic acids. Cotton cloth was formerly bleached, and linen is now bleached, by spreading it on the grass and leaving it for weeks to be acted on by ozone, which is usually present in the air in small quantities, especially in the country. Ozone is a disinfectant, like other bleaching agents, and serves to clear the air of noxious gases and germs of infectious diseases. So much ozone is reduced in this way that the air of cities contains less of it than country air. A third is consumed in uniting with the substance which it oxidizes, while two-thirds are changed into oxygen, as in Experiment 79.

It is unhealthful to breathe much ozone, but a little in the air is desirable for disinfection.

Ozone will cause the inert N of the air to unite with H, to form ammonia. No other agent capable of doing this is known, so that all the NH3 in the air, in fact all ammonium compounds taken up by plants from soils and fertilizers, may have been made originally through the agency of ozone. At a low temperature ozone has been liquefied. It is then distinctly blue.

Electrolysis of water is the best mode of preparing this substance in quantity. When prepared from P it is mixed with P2O3.

Chapter XXVII.

CHEMISTRY OF THE ATMOSPHERE.

130. Constituents.—The four chief constituents of the atmosphere are N, O, H2O, CO2, in the order of their abundance. What experiments show the presence of N, O, and CO2 in the air? Set a pitcher of ice water in a warm room, and the moisture that collects on the outside is deposited from the air. This shows the presence of H2O. Rain, clouds, fog, and dew prove the same. H2SO4 and CaCl2, on exposure to air, take up water. Experiment 18 shows that there is not far from four times as much N as O by volume in air. Hence if the atmosphere were a compound of N and O, and the proportion of four to one were exact, its symbol would be N4O.

131. Air not a Compound.—The following facts show that air is not a compound, but rather a mixture of these gases.

1. The proportion of N and O in the air, though it does not vary much, is not always exactly the same. This could not be true if it were a compound. Why?

2. If N4O were dissolved in water, the N would be four times the O in volume; but when air is dissolved, less than twice as much N as O is taken up.

3. No heat or condensation takes place when four measures of N are brought in contact with one of O. It cannot then be N4O, for the vapor density of N4O would be 36—i.e. (14 x 4 + 16) / 2; but that of air is 14 1/2 nearly —i.e. (14 x 4 + 16) / 5. Analysis shows about 79 parts of N to 21 parts of O by volume in air.

132. Water.—The volume of H2O, watery vapor, in the atmosphere is very variable. Warm air will hold more than cold, and at any temperature air may be near saturation, i.e. having all it will hold at that temperature, or it may have little. But some is always present; though the hot desert winds of North Africa are not more than 1/15 saturated. A cubic meter of air at 25 degrees, when saturated, contains more than 22 g. of water.

133. Carbon Dioxide.—Carbon dioxide does not make up more than three or four parts in ten thousand of the air; but, in the whole of the atmosphere, this gives a very large aggregate. Why does not CO2 form a layer below the O and N?

134. Other Ingredients.—Other substances are found in the air in minute portions, e.g. NH3 constitutes nearly one-millionth. Air is also impregnated with living and dead germs, dust particles, unburned carbon, etc., but these for the most part are confined to the portion near the earth's surface. In pestilential regions the germs of disease are said sometimes to contaminate the air for miles around.

Chapter XXVIII.

THE CHEMISTRY OF WATER.

135. Pure Water.—Review the experiments for electrolysis, and for burning H. Pure water is obtained by distillation.

Experiment 80.—Provide a glass tube 40 or 50 cm long and 3 or 4 cm in diameter. Fit to each end a cork with two perforations, through one of which a long tube passes the entire length of the larger tube (Fig. 32a). Connect one end of this with a flask of water arranged for heating; pass the other end into an open receptacle for collecting the distilled water. Into the other perforations lead short tubes,— the one for water to flow into the large tube from a jet; the other, for the same to flow out. This condenses the steam by circulating cold water around it. The apparatus is called a Liebig's condenser. Put water into the flask, boil it, and notice the condensed liquid. It is comparatively pure water; for most of the substances in solution have a higher boiling-point than water, and are left behind when it is vaporized.

(Fig. 32a.)

136. Test.

Experiment 81.—Test the purity of distilled water by slowly evaporating a few drops on Pt foil in a room free from dust. There should be no spot or residue left on the foil. Test in the same way undistilled water. 137. Water exists in Three States,— solid, liquid, and vaporous. It freezes at 0 degrees, suddenly expanding considerably as it passes into the solid state. It boils, i.e. overcomes atmospheric pressure and is vaporized, at 100 degrees (760 mm pressure). If the pressure is greater, the boiling-point is raised, i.e. it takes a higher temperature to overcome a greater pressure. If there be less pressure, as on a mountain, the boiling-point is lowered below 100 degrees. Salts dissolved in water raise its boiling-point, and lower its freezing-point to an extent depending on the kind and quantity of the salt. Water, however, evaporates at all temperatures, even from ice.

Pure water has no taste or smell, and, in small quantities, no color. It is rarely if ever found on the earth. What is taken up by the air in evaporation is nearly pure; but when it falls as rain or snow, impurities are absorbed from the atmosphere. Water falling after a long rain, especially in the country, is tolerably free from impurities. Some springs have also nearly pure water; but to separate all foreign matter from it, water must be distilled. Even then it is liable to contain traces of ammonia, or some other substance which vaporizes at a lower temperature than water.

138. Sea-Water.—The ocean is the ultimate source of all water. From it and from lakes, rivers, and soils, water is taken into the atmosphere, falls as rain or snow, and sinks into the ground, reappearing in springs, or flowing off in brooks and rivers to the ocean or inland seas. Ocean water must naturally contain soluble salts; and many salts which are not soluble in pure water are dissolved in sea-water. In fact, there is a probability that all elements exist to some extent in sea-water, but many of them in extremely minute quantities. Sodium and magnesium salts are the two most abundant, and the bitter taste is due to MgSO4 and MgCl2. A liter of sea- water, nearly 1000 g., holds over 37 g. of various salts, 29 of which are NaCl. See Hard Water.

139. River Water.—River water holds fewer salts, but has a great deal of organic matter, living and dead, derived from the regions through which it flows. To render this harmless for drinking, such water should be boiled, or filtered through unglazed porcelain. Carbon filters are now thought to possess but little virtue for separating harmful germs.

140. Spring Water.—The water of springs varies as widely in composition as do the rocks whence it bubbles forth. Sulphur springs contain much H2S; many geysers hold SiO2 in solution; chalybeate waters have compounds of Fe; others have Na2SO4, MgSO4 NaCl, etc.

CHAPTER XXIX.

THE CHEMISTRY OF FLAME.

141. Candle Flame.

Experiment 82.—Examine a candle flame, holding a dark object behind it. Note three distinct portions: (1) a colorless interior about the wick, (2) a yellow light-giving portion beyond that, (3) a thin blue envelope outside of all, and scarcely discernible. Hold a small stick across the flame so that it may lie in all three parts, and observe that no combustion takes place in the inner portion.

142. Explanation.—A candle of paraffine, or tallow, is chiefly composed of compounds of C and H, in the solid state. The burning wick melts the solid; the liquid is then drawn up by the wick till the heat vaporizes and decomposes it, and O of the air comes in contact with the outer heated portion of gas, and burns it completely. Air tends to penetrate the whole body of the flame, but only N can pass through uncombined, for the O that is left after combustion in the outer portion seizes upon the compounds of C and H in the next, or yellow, part. There is not enough O here for complete combustion; at this temperature H burns before C, and the latter is set free. In that state it is of course a solid. Now an incandescent solid, or one glowing with heat, gives light, while the combustion of a gas gives scarcely any light, though it may produce great heat. While C in the middle flame is glowing, during the moment of its dissociation from H, it gives light. In the outer flame the temperature is high enough to burn entirely the gaseous compounds of C and H together, so that no solid C is set free, and hence no light is given except the faint blue. No combustion takes place in the inner blue cone, because no O reaches there.

By packing a wick into a cylindrical tin cup 5 or 10 cm high and 4 cm in diameter, containing alcohol, and lighting it, gunpowder can be held in the middle of the flame in a def. spoon, without burning. This shows the low temperature of that portion. Burning P will also be extinguished, thus showing the exclusion of O.

143. Bunsen Flame.

Experiment 83.—Examine a Bunsen burner. Unscrew the top, and note the orifices for the admission of gas and of air. Make a drawing. Replace the parts; then light the gas at the top, opening the air-holes at the base. Notice that the flame burns with very little color. Try to distinguish the three parts, as in the candle flame. These parts can best be seen by allowing direct sunlight to fall on the flame and observing its shadow on a white ground. Make a drawing of the flame. Hold across it a Pt wire and note at what part the wire glows most. Also press down on the flame for an instant with a cardboard or piece of paper; remove before it takes fire, and notice the charred circle. Put the end of a match into the blue cone, and note that it does not burn. Put the end of a Pt wire into this blue cone, and observe that it glows when near the top of the cone. What do these experiments show? Ascertain whether this inner portion contains a combustible material, by holding in it one end of a small d.t., and trying to ignite any gas escaping at the other end. It should burn. This shows that no combustion takes place in the interior of the flame, because sufficient free O is not present.

Next, close the air-holes, and note that the flame is yellow and gives much light. From this we infer the presence of solid particles in an incandescent state. But these could not come from the air. They must be C particles which have been set free from the C and H compounds of the gas, just as in the candle flame. The smoke that rises proves this. Hold an e.d. in the flame and collect some C. Try the same with the air-holes open. 144. Light and Heat of Flame.—Which of the two flames is hotter, the one with the air-holes open, or that with them closed? Evidently the former; for air is drawn in and mixes with the gas as it rises in the tube, and, on reaching the flame at the top, the two are well mingled, and the gaseous compounds of C and H burn at so high a temperature that solid C is not freed; hence there is little light. On closing the air-holes, no O can reach the flame except from the outside, and the heat is much less intense.

(Fig 33.) (Fig 34.)

The H burns first, and sets the C free, which, while glowing, gives the light. This again illustrates the facts (1) that flame is caused by burning gas; (2) that light is produced by incandescent solids. Charcoal, coke, and anthracite coal burn without flame, or with very little, because of the absence of gases.

145. Temperature of Combustion.

Experiment 84.—Light a Bunsen flame, with the basal orifices open, and hold over it a fine wire gauze. Notice that the flame does not rise above the gauze. Extinguish the light, and try to ignite the gas above the gauze, holding the latter within 5 or 6 cm of the burner tube. Notice that it does not burn below the gauze (Fig. 33).

Gas and O are both present. Evidently, then, the only condition wanting for combustion is a sufficiently high temperature. The gauze cools the gas below its kindling- point.

This principle is made use of in the miner's lamp of Davy (Fig. 34). In coal mines a very inflammable gas, CH4, called fire-damp, issues from the coal. If this collects in large quantities and mixes with O of the air, a kindling-point is all that is needed to make a violent explosion. An ordinary lamp would produce this, but the gauze lamp prevents it; for, though the inside may be filled with burning gas, CH4, the flame cannot communicate with the outside.

(Fig 35.) (Fig 36.) a, reducing flame b, oxidizing flame

146. Oxidizing and Reducing Flames.—The hottest part of a Bunsen flame is just above the inner blue cone (b, Fig. 36). Evidently there is more O at that point. If a reducing agent, i.e. a substance which takes up O, be put into this part of the flame, the latter will remove the O and appropriate it, forming an oxide. Cu heated there would become copper oxide. This part is called the oxidizing flame. The inner blue part of the Bunsen flame is devoid of O. It ought to remove O from an oxidizing agent, i.e. a substance which supplies O. If copper oxide be heated there (a, Fig. 36) by means of a mouth blow-pipe (Fig. 35), the flame will appropriate the O and leave the copper. This is called the reducing flame. Only the upper part of this blue central cone has heat enough to act in this way. By using a prepared piece of metal, to make the flame thin and to shut off the air, and then blowing the flame with a blow-pipe, greater strength can be obtained in both oxidizing and reducing flames (Fig. 36).

147. Combustible and Supporter Interchangeable.— H was found to burn in O. H was the combustible, O the supporter. Would O itself burn in H?—i.e. would the combustible become the supporter, and the supporter the combustible? As illuminating gas consists largely of H, and as air is part O, we may try the experiment with gas and air. Gas will burn in air. Will air burn in gas?

Experiment 85.—Fit a cork with two holes in it to the large end of a lamp chimney. Through each hole pass a short piece of tubing, and connect one of these with a rubber tube leading to a gas-jet. Pass a metallic tube, long enough to reach the top of the chimney, through the other, so that it will move easily up and down. Turn on the gas, and light it at the top of the chimney. Hold the end of the tube passing through the cork in the flame for a minute, then draw it down to the middle of the chimney (Fig. 37, a) and finally slowly remove it (b). Note that O from the air is burning in the gas. Which is the supporter, and which the combustible in this case? O will burn equally well in an atmosphere of H, as can be shown by experiment.

148. Explosive Mixture of Gases.

Experiment 86.—Slowly turn down the burning gas of a Bunsen lamp, having the orifices open, and notice that it suddenly explodes and goes out at the top, but now burns at the base. As the gas was gradually turned off, more air became mixed with it, until there was the right proportion of each gas for an explosion. Figure 38 shows the same thing. Light the gas at the top a, when the tube c covers the jet b. Then gradually raise the tube c. At a certain place there is the same explosion as with the lamp.

149. Generalizations.—These experiments show (1) that three conditions are necessary for combustion,—a combustible, a supporter, and a burning temperature which varies for different substances. Given these, "a fire" always results. The conditions for "spontaneous combustion" do not differ from those of any combustion. See Experiments 34, 112, 113, 114. (2) That combustible and supporter are interchangeable. If H burns in O, O will burn in H, the product, being the same in each case. (3) For any combustion there must be a certain proportion of combustible and of supporter. Twenty per cent of CO2 in the air dilutes the O to such an extent that C will not burn. Hence the utility of the chemical engine for putting out fires. (4) When two

gases, a combustible and a supporter, are mixed in the requisite proportion, they form an explosive mixture, needing only the kindling temperature to unite them.

Chemical combination is always accompanied by disengagement of heat. Chemical dissociation is always accompanied by absorption of heat. The disengagement, or the absorption, is not always evident to the senses.

Combustion is the chemical combination of two or more substances with the self-evident disengagement of great heat, and usually of light.

The temperature of ignition varies greatly with different substances. PH3 burns spontaneously at the usual temperatures of the air. P takes fire at 60 degrees, but even at 10 degrees it oxidizes with rapidity enough to produce phosphorescence. The vapor of CS2 may be set on fire by a glass rod heated to 150 degrees, but a red-hot iron will not ignite illuminating gas.

Spontaneous combustion often takes place in woolen or cotton rags which have been saturated with oil. The oil rapidly absorbs O, and sets fire to the cloth. This is thought to be the origin of some very destructive fires.

CHAPTER XXX.

CHLORINE.

150. Preparation.

Experiment 87.—Put into a t.t. 5 g. of fine granular MnO2 and 10 cc. HCl. Apply heat carefully, and collect the gas by downward displacement in a receiver loosely covered with paper (Fig. 39). Add more HCl if needed. Have a good draft of air, and do not inhale the gas. If you have accidentally breathed it, inhale alcohol vapor from a handkerchief; alcohol has great affinity for Cl. Note the color of the gas, and compare its weight with that of air.

MnO2 + 4 HCl = MnCl2 + 2 H2O + 2 Cl. How much Cl can be separated with 5 g. MnO2?

If preferred, a flask may be used for a generator instead of a t.t. Cl can be obtained directly from NaCl by adding H2SO4 (which produces HCl) and MnO2. 2 NaCl + 2 H2SO4 + MnO2 = MnSO4 + Na2SO4 + 2 H2O + 2 Cl. Try the experiment, using a t.t. and adding water.

151. Cl from Bleaching-Powder.

Experiment 88.—Put a few grams of bleaching- powder into a small beaker, and set this into a larger one. Cover the latter with pasteboard or paper, through which passes a thistle-tube reaching into the small beaker (Fig. 40). Pour through the tube a little H2SO4 dilated with its volume of H2O.

152. Chlorine Water.—A solution of Cl in water is often useful, and may be made as follows:— Experiment 89.—To 3 or 4 crystals of KClO3 add a few drops of HCl. Heat a minute, and when the gas begins to disengage, pour in 10 cc. H2O, which dissolves the gas. 2 KClO3 + 4 HCl = 2 KCl + Cl2O4 + 2 H2O + 2 Cl.

153. Bleaching Properties.

Experiment 90.—Put into a receiver of Cl, preferably before generating it, two pieces of Turkey red cloth, one wet, the other dry; a small piece of printed paper and a written one; also a red rose or a green leaf, each wet. Note from which the color is discharged. If it is not discharged from all, put a little H2O into the receiver, shake it well, and state what ones are bleached.

Experiment 91.—(1) Add 5 cc. of Cl water to 5 cc. of indigo solution. (2) Treat in the same way 5 cc. K2Cr2O7 (potassium dichromate) solution, and record the results.

Indigo, writing-ink, and Turkey red or madder, are vegetable pigments; printer's ink contains C, and K2Cr2O7 is a mineral pigment. State what coloring matters Cl will bleach.

154. Disinfecting Power.

Experiment 92.—Pass a little H2S gas from a generator into a t.t. containing Cl water. Look for a deposit of S. Notice that the odor of H2S disappears. H2S + 2 Cl = 2 HCl + S.

155. A Supporter of Combustion.

Experiment 93.—Sprinkle into a receiver of Cl a very little fine powder or filings of Cu, As, or Sb, and notice the combustion. Observe that here is a case of combustion in which O does not take part. Chlorides of the metals are of course formed. Write the reactions. See whether Cl will support the combustion of paper or of a stick of wood.

Experiment 94.—Warm 2 or 3 cc. of oil of turpentine (C1OH16) in an evaporating-dish; dip a piece of tissue paper into it, and very quickly thrust this into a receiver of Cl. It should take fire and deposit carbon. C1OH16 + 16 Cl = ? Test the moisture on the sides of the receiver with litmus. Clean the receiver with a little petroleum.

Experiment 95.—Prepare a H generator with a lamp-tube bent as in Figure 41. Light the H, observing the cautions in Experiment 23, and when well burning, lower the flame into a receiver of Cl. Observe the change of color which the flame undergoes as it comes in contact with Cl. Give the reaction for the burning. Test with litmus any moisture on the sides of the receiver. A mixture of Cl and H, in direct sunlight combines with explosive violence; whereas in diffused sunlight it combines slowly, and in darkness it does not combine. From these experiments state the chief properties of Cl, and what combustion it will support.

[Figure 41.]

156. Sources and Uses.—The great source of Cl is NaCl, though it is often made from HCl. Its chief use is in making bleaching- powder, one pound of which will bleach 300 to 500 pounds of cloth. Cl is very easily liberated from this powder by a dilute acid, or, slowly, by taking moisture from the air. Hence its use as a disinfectant in destroying noxious gases and the germs of infectious diseases. Cl attacks organic matter and germs as it does the membrane of the throat or lungs, owing to its affinity for H.

Cl is the best bleaching agent for cotton goods. It is not suitable for animal materials, such as silk and wool, as it attacks their fiber. It does not discharge either mineral or carbon colors. The chemistry of bleaching is obscure.

As dry material will not bleach, Cl seems to unite with H in H2O and to set O free. The O then unites with some portion of the coloring matter, oxidizing it, and breaking up its molecule. Colors bleached by Cl cannot be restored.

Chapter XXXI.

BROMINE.

Examine bromine, potassium bromide, sodium bromide, magnesium bromide.

157. Preparation.

Experiment 96.—Pulverize 2 or 3 g. KBr, and mix it with about the same bulk of MnO2. After putting this into a t.t, add as much H2SO4, mix them together by shaking, attach a d.t., and conduct the end of it into a t.t. that is immersed in a bottle of cold water. Slowly heat the contents of the t.t., and notice the color of the escaping vapor, and any liquid that condenses in the receiver. Avoid inhaling the fumes, or getting them into the eyes.

MnO2 + 2 KBr + 2 H2SO4 = ? Compare this with the equation for making Cl from NaCl.

158. Tests.

Experiment 97.—Try the bleaching action of Br vapor as in the case of Cl. Bleach a piece of litmus paper, and try to restore the color with NH4OH. Explain its bleaching and disinfecting action. Try the combustibility of As, Sb, and Cu.

159. Description.—Bromine at usual temperatures is a liquid element; it is the only common one except Hg; it. quickly evaporates on exposure to air. The chemistry of its manufacture is like that of Cl; its bleaching and disinfecting powers are similar to the latter, though they are not quite so strong as those of Cl. Its affinity for H and for metals is also strongly marked. A drop of Br on the skin produces a sore slow to heal. Bromine salts are mainly KBr, NaBr, MgBr2. These in small quantities accompany NaCl, and are most common in brine springs. The world's supply of Br comes chiefly from West Virginia and Ohio, over 300,000 pounds being produced from the salt (NaCl) wells there in 1884. The water taken from these wells is nearly evaporated, after which NaCl crystallizes out, leaving a thick liquid—bittern, or mother liquor—which contains the salts of Br. The bittern is treated with H2SO4 and Mn02, as above.

For transportation in large quantities, Br has to be made into the salts NaBr and KBr, on account of the danger attending leakage or breakage of the receptacles for Br.

160. Uses.—Its chief uses are in photography (page 167), medicine, as KBr, and analytical chemistry.

Chapter XXXII.

IODINE.

Examine iodine, potassium iodide.

161. Preparation of I.

Experiment 98.—Put into a t.t. 2 or 3 g. of powdered KI mixed with an equal bulk of MnO2, add H2SO4 enough to cover well, shake together, complete the apparatus as for making Br, and heat. Notice the color of the vapor, and any sublimate. The direct product of the solidification of a vapor is called a sublimate. The process is sublimation. Observe any crystals formed. Write the reaction, and compare the process with that for making Br and Cl. Compare the vapor density of I with that of Br and of Cl. With that of air. What vapor is heavier than I? What acid and what base are represented by KI?

162. Tests.

Experiment 99.—(1) Put a crystal of I in the palm of the hand and watch it for a minute. (2) Put 2 or 3 crystals into a t.t., and warm it, meanwhile holding a stirring-rod half-way down the tube. Notice the vapor, also a sublimate on the sides of the t.t. and rod. (3) Add to 2 or 3 crystals in a t.t. 5 cc. of alcohol, C2H5OH; warm it, and see whether a solution is formed. If so, add 5 cc. H2O and look for a ppt. of I. Does this show that I is not at all soluble in H2O, or not so soluble as in alcohol?

163. Starch Solution and Iodine Test.

Experiment 100.—Pulverize a gram or two of starch, put it into an evaporating-dish, add 4 or 5 drops of water, and mix; then heat to the boiling-point 10 cc. H2O in a t.t., and pour it over the starch, stirring it meanwhile.

(1) Dip into this starch paste a piece of paper, hold it in the vapor of I, and look for a change of color. (2) Pour a drop of the starch paste into a clean t.t., and add a drop or two of the solution of I in alcohol. Add 5 cc. H2O, note the color, then boil, and finally cool. (3) The presence of starch in a potato or apple can be shown by putting a drop of I solution in alcohol on a slice of either, and observing the color. (4) Try to dissolve a few crystals of I in 5 cc. H2O by boiling. If it does not disappear, see whether any has dissolved, by touching a drop of the water to starch paste. This should show that I is slightly soluble in water.

164. Iodo-Starch Paper.

Experiment 101.—Add to some starch paste that contains no I 5 cc. of a solution of KI, and stir the mixture. Why is it not colored blue? Dip into this several strips of paper, dry them, and save for use. This paper is called iodo-starch paper, and is used as a test for ozone, chlorine, etc. Bring a piece of it in contact with the vapor of chlorine, bromine, or ozone, and notice the blue color.

Experiment 102.—Add a few drops of chlorine water to 2cc. of the starch and KI solution in 10 cc. H2O. This should show the same effect as the previous experiment.

165. Explanation.—Only free I, not compounds of it, will color starch blue. It must first be set free from KI. Ozone, chlorine, etc., have a strong affinity for K, and when brought in contact with KI they unite with K and set free I, which then acts on the starch present. Com- plete the equation: KI + Cl = ?

166. Occurrence.—The ultimate source of I is sea water, of which it constitutes far too small a percentage to be separated artificially. Sea-weeds, or algae, especially those growing in the deep sea, absorb its salts—NaI, KI, etc.—from the water. It thus forms a part of the plant, and from this much of the I of commerce is obtained. Algae are collected in the spring, on the coasts of Ireland, Scotland, and Normandy, where rough weather throws them up. They are dried, and finally burned or distilled; the ashes are leached to dissolve I salts; the water is nearly evaporated, and the residue is treated with H2SO4, and MnO2, as in the case of Br and Cl. I also occurs in Chili, as NaI and NaIO3, mixed with NaNO3. This is an important source of the I supply.

167. Uses.—I is much used in medicine, and was formerly employed in taking daguerreotypes and photographs. Its solution in alcohol or in ether is known as tincture of iodine.

168. Fluorine.—F, Cl, Br, I, are called halogens or haloids, and exist in compounds—salts—in sea water. F is the most active of all elements, combining with every element except O. Until recently it has never been isolated, for as soon as set free from one compound it attacks the nearest substance, and seems to be as much averse to combining with itself, or to existing in the elementary state, as to uniting with O. It is supposed to be a gas, and, as is claimed, has lately been isolated by electrolysis from HF in a Pt U-tube. Fluorite (CaF2) and cryolite (Al2F6 + 6 NaF) are its two principal mineral sources. The enamel of the teeth contains F in composition.

CHAPTER XXXIII.

THE HALOGENS.

169. Halogens Compared.—The elements F, Cl, Br, I, form a natural group. Their properties, as well as those of their compounds, vary in a step-by-step way, as seen below. F is sometimes an exception. They are best remembered by comparing them with one another. Notice:

1. Similarity of name-ending. Each name ends in ine.

2. Similarity of origin. Salt water is the ultimate source of all, except F.

3. Similarity of valence. Each is usually a monad.

4. Similarity of preparation. Cl, Br, I, are obtained from their salts by means of MnO2 end H2SO4.

5. Variation in occurrence. Cl occurs in sea-salt, Br in sea- water, I in sea-weed.

6. Variation in color; F being colorless, Cl green, Br red, I violet.

7. Gradation in sp. gr.; F 19, Cl 35.5, Br 80, I 127.

8. Gradation in state, corresponding to sp. gr.; F being a light gas, Cl a heavy gas, Br a liquid, I a solid.

9. Corresponding gradation in their usual chemical activity; F being most active, then Cl, Br, and I.

10. Corresponding gradation in the strength of the H acids; the strongest being HF, the next, HCl, etc.

11. Corresponding gradation in the explosibility of their N compounds; the strongest NCl3, the next, NBr3, etc.

12. Corresponding gradation in the number of H and O acids; Cl 4, Br 3, I 2.

170. Compounds.—The following are some of the oxides, acids, and salts of the halogens. Name them.

CI2O (+H2O=) 2 HClO. The salts are hypochlorites, as Ca(ClO)2. Cl2O3 (+H20=) 2 HClO2. The salts are chlorites, as KClO2. Cl2O4 — HClO3 The salts are chlorates, as KClO3. — HClO4 The salts are perchlorates, as KClO4, — HBrO The salts are ? KBrO, — — The salts are wanting. — HBrO3. The salts are ? KBrO3, — HBrO4. The salts are ? KBrO4, — — The salts are wanting. — — The salts are wanting. I2O5 (+H2O=) 2 HIO3. The salts are ? KIO3. — HIO4. The salts are ? KIO4.

F forms no oxides, and no acids except HF. HF, HCl, HBr, HI, are striking illustrations of acids with no O. HClO4 is a very strong oxidizing agent. A drop of it will set paper on fire, or with powdered charcoal explode violently. This is owing to the ease with which it gives up 0. Notice why its molecule is broken up more readily than HC103. The higher the molecular tower, or the more atoms it contains, the greater its liability to fall. Some organic compounds contain hundreds of atoms, and hence are easily broken down, or, as we say, are unstable. Inorganic compounds are, as a rule, much more stable than organic ones. It is not always true, however, that the compound with the least number of atoms is the most stable. SO2 is more stable than SO3, but H2SO3 is less so than H2SO4. Chapter XXXIV.

VAPOR DENSITY AND MOLECULAR WEIGHT.

Examine a liter measure, in the form of a cube,—cubic decimeter, —and a cubic centimeter.

171. Gaseous Weights and Volumes.—A liter of H, at 0 degrees and 760 mm., weighs nearly 0.09 g. This weight is called a crith. Find the weight of H in the following, in criths and in grams: 15 1., 0.07 1., 50.3 1., 0.035 1., 0.6 1..

It has been estimated that there are (10) 24. molecules of H in a liter. Does the number vary for different gases? The weight of a molecule of H in parts of a crith is 1/(10) 24.; in parts of a gram .09/(10) 24.. If the H molecule is composed of 2 atoms, what is the weight of its atom in fractions of a crith? What in fractions of a gram? The weight of the H atom is a microcrith. What part of a crith is a microcrith?

172. Vapor Density.—Vapor density, or specific gravity referred to H as the standard, (Physics) is the ratio of the weight of a given volume of a gas or vapor to the weight of the same volume of H. A liter of steam weighs nine times as much as a liter of H. Its vapor density is therefore nine. For convenience, a definite volume of H is usually taken as the standard, viz., the H atom. The volume of the H atom and that of the half-molecule of H2O, or of any gas are identical, each being represented by one square. If, then, the standard of vapor density is the H atom, half the molecular weight of a gas must be its vapor density, since it is evident that we thus compare the weights of equal volumes. The vapor density of H2O, steam, is found from the symbol as follows: (2 + 16) / 2 = 9. To obtain the vapor density of any compound from the formula, we have only to divide its molecular weight by two. Find the vapor density of HCl, N2O, NO, C12H22O11, Cl, CO2, HF, SO2. Explain each case.

The half-molecule, instead of the whole, is taken; because our standard is the hydrogen atom, the smallest portion of matter, by weight, known to science.

How many criths in a liter of HCl? How many grams? Compute the number of criths and of grams in one liter of the compounds whose symbols appear above.