the heat evolution is more rapid, and the product of combustion is a solid (cf. § 24) which does not dilute the oxygen remaining, hence burning goes on until all the oxygen is used up.

A moderate draft of air is favorable to combustion, because it brings fresh oxygen and removes the nitrogen and the products of burning. But a too rapid current will cool the burning body below the kindling temperature. Thus, a flame may be "blown out" by a gust of air.

30. The Safety Lamp. The lowering of the temperature below the kindling point is admirably illustrated in the safety lamp invented by Sir Humphry Davy to prevent explosions of "firedamp" a mixture of marsh gas (292) and air in mines. The lamp (Fig. 7) consists of an ordinary lantern entirely surrounded by wire gauze. When it is brought into an explosive gaseous mixture, the gases diffuse through the wire gauze, and burn inside the lamp, but the heat produced is conducted away by the gauze, so that the gases outside do not reach the ignition temperature. A series of small explosions inside the lamp warns the miner of his danger.

FIG. 7.

REVERSED COMBUSTION.

25

31. Spontaneous Combustion. Even in slow oxidation the kindling temperature may be reached if the heat generated is not dispersed. This fact explains "spontaneous combustions," in which oily rags, waste, etc., take fire without apparent cause. The linseed oil of paint becomes hard by absorbing oxygen from the air. Ordinarily the heat produced by the oxidation is dissipated, but when the rags are piled in heaps this is not possible, hence the temperature rises to the point of ignition. Coal packed closely in poorly ventilated bunkers often takes fire because of slow oxidation.

Spontaneous combustion may be illustrated by means of a solution of phosphorus (only a small amount must be used) in carbon disulphide. If this solution is poured upon a filter paper supported on a ring stand, the phosphorus will soon take fire "spontaneously." The explanation of the phenomenon is that the evaporation of the carbon disulphide leaves the phosphorus in the pores of the paper, where it is oxidized. The heat generated, being prevented from escaping by the nonconducting filter paper, soon raises the temperature of some part of the paper to the ignition temperature of the phosphorus.

32. Reversed Combustion. If both the combustible and the supporter of its combustion are gaseous, the combustion may be reversed. Thus, oxygen may become the burning body and illuminating gas the supporter of combustion. This reversal may be shown by a very simple experiment.

FIG. 8.

A bottle (Fig. 8) is supported, mouth downward, and filled with illuminating gas by displacing the air. The gas at the mouth of the bottle is then lighted, and while it is burning a jet of oxygen is brought up into the bottle. The oxygen takes fire at the bottle's mouth and burns in the atmosphere of illuminating gas. The oxygen jet may be replaced by a combustion spoon of potassium chlorate which has been heated so that it gives off oxygen.

33. Importance and Uses of Oxygen. As the table in § 9 shows, oxygen makes up 47% of the earth's solid crust and nearly 86% of the ocean. Eight ninths of water is oxygen (§ 64), and about 23%, by weight, of the atmosphere.

Practically all living things require oxygen for their respiration. Water animals get it from the air dissolved in the water. The object of respiration is to get oxygen in contact with worn-out tissues, so that they may be converted into products which the organism can readily eliminate. Among these is the gas carbon dioxide (cf. § 281). The heat generated in these oxidations keeps the bodies of the higher animals warm (§ 25). In certain diseases the lungs are not able to supply the body with oxygen rapidly enough from the air, so pure oxygen is used.

While plants use oxygen, and give off carbon dioxide, just as animals do, yet they carry on the reverse action, i. e., use up carbon dioxide, and give

off oxygen, in building up their tissues. The carbon dioxide and water used by plants for this purpose together contain too much oxygen, hence the excess is given off to the air. Animals and plants are thus interdependent, the one giving off what the other As a result the quantity of oxygen in the air is kept practically constant.

uses.

34. Exercises.

1. How many grams of mercury and of oxygen will be formed by the decomposition of 43.2 grams of mercuric oxide?

2. How many grams of manganese dioxide are needed to give, when decomposed by heat, 12 grams of oxygen? How much manganous-manganic oxide will be formed?

3. Calculate the weight of potassium chlorate that will produce, when heated, 74.5 grams of potassium chloride. What weight of oxygen is formed at the same time?

4. How many cubic centimeters of oxygen can be made from 1.2 grams of potassium chlorate if 1 c.c of the gas weighs .0014 gram?

5. How many grams of potassium chlorate must be decomposed to give 36 liters of oxygen when 1 liter of oxygen weighs 1.25 grams?

6. What weight of oxygen can be obtained from 50 grams of water? What volume will it have at 0° C. and 760 mm.?

7. Name the materials used in starting a coal fire. In lighting a gas stove. What are the causes of the difference?

8. Name the oxides used as sources of oxygen in the order of their stability toward heat, beginning with the least stable one. Where would you place potassium chlorate?

9. Devise an experiment to prove that part of the air disappears in the rusting of iron.

CHAPTER III.

MEASUREMENT OF GASES.

35. The Volumes of Gases. Because it is difficult to weigh gases accurately their quantity is usually given, not in weight units, but in units of volume, such as the cubic centimeter, or the liter (cf. Appendix i). Hence it is important that the volumes of gases be determined with great accuracy. Now, experiments with gases must often stand for hours, or over night. Meanwhile the temperature and pressure, and, consequently, the volume, undergo change. Cooling will contract the volume; heating expands it. Increase of pressure will make the volume smaller; a lowering of pressure will increase the volume. So the chemist must be able to calculate what the volume of a gas, measured at any temperature and pressure, will be at any other temperature and pressure.

36. Measurement of Pressure. — Pressure of gases is measured in atmospheres, or in millimeters of mercury. The instrument used for measuring atmospheric pressure is the barometer (Fig. 9). The simple barometer consists of a glass tube about a meter long, closed at the upper end, and partly