oxygen, KClO3. It occurs as transparent colourless plates, is soluble in water, especially in hot water, and resembles common table salt in some of its reactions and physical properties; it melts on heating, and in melting begins to decompose, evolving oxygen gas. This decomposition ends in all the oxygen being evolved from the potassium chlorate, potassium chloride being left as a residue, according to the equation KClO3 = KCl + O3.13 This decomposition proceeds at a temperature which allows of its being conducted in a glass vessel. However, in decomposing, the molten potassium chlorate swells up and boils, and gradually solidifies, so the evolution of the oxygen is not regular, and the glass vessel may crack. In order to overcome this inconvenience, the potassium chlorate is crushed and mixed with a powder of a substance which is infusible, incapable of combining with the oxygen evolved, and is a good conductor of heat. Usually it is mixed with manganese peroxide.14 The decomposition of the potassium chlorate is then considerably facilitated, and proceeds at a lower temperature (because the entire mass is then better heated, both externally and internally), without swelling up, and this method is therefore more convenient than the decomposition of the salt alone. This method for the preparation of oxygen is very convenient; it is generally employed when a small quantity of oxygen is required. Further, potassium chlorate is easily obtained pure, and it evolves much oxygen. 100 grams of the salt give as much as 39 grams, or 30 litres, of oxygen. This method is so simple and easy,15 that a course of practical chemistry is often commenced by the preparation of oxygen by this method, and of hydrogen by the aid of zinc and sulphuric acid, since by means of these gases many interesting and striking experiments may be performed.16
A solution of bleaching powder, which contains calcium hypochlorite, CaCl2O2, evolves oxygen on gently heating when a small quantity of certain oxides is added—for instance, cobalt oxide, which in this case acts by contact (see Introduction). When heated by itself, a solution of bleaching powder does not evolve oxygen, but it oxidises the cobalt oxide to a higher degree of oxidation; this higher oxide of cobalt in contact with the bleaching powder decomposes into oxygen and lower oxidation products, and the resultant lower oxide of cobalt with bleaching powder again gives the higher oxide, which again gives up its oxygen, and so on.17 The calcium hypochlorite is here decomposed according to the equation CaCl2O2 = CaCl2 + O2. In this manner a small quantity of cobalt oxide18 is sufficient for the decomposition of an indefinitely large quantity of bleaching powder.
The properties of oxygen.19—It is a permanent gas—that is, it cannot be liquefied by pressure at the ordinary temperature, and further, is only liquefied with difficulty (although more easily than hydrogen) at temperatures below -120°, because this is its absolute boiling point. As its critical pressure20 is about 50 atmospheres, it can be easily liquefied under pressures greater than 50 atmospheres at temperatures below -120°. According to Dewar, the density of oxygen in a critical state is 0·65 (water = 1), but, like all other substances in this state,21 it varies considerably in density with a change of pressure and temperature, and therefore many investigators who made their observations under high pressures give a greater density, as much as 1·1. Liquefied oxygen is an exceedingly mobile transparent liquid, with a faint blue tint and boiling (tension = 1 atmosphere) about -180°. Oxygen, like all gases, is transparent, and like the majority of gases, colourless. It has no smell or taste, which is evident from the fact of its being a component of air. The weight of one litre of oxygen gas at 0° and 760 mm. pressure is 1·4298 gram; it is therefore slightly denser than air. Its density in respect to air = 1·1056 and in respect to hydrogen = 16.22