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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
(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.
(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.
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 otherNames. 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.
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.
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.
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.
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.