This method is applied in practice only to the preparation of oxygen, and of chlorine, bromine, and iodine; but many other elements may be thus made, where the compound heated does not tend to re-form on cooling. These cases will be considered first.

Ordinary coal-gas consists chiefly of methane, CH4, ethylene, C2H4, carbon monoxide, CO, and hydrogen, the last amounting to nearly 50 per cent, of the volume of the gas. This hydrogen owes its origin, at least in part, to the decomposition of its compounds with carbon, by their coming into contact with the red-hot walls of the retort in which the coal is distilled. Carbon deposits in a dense black mass on the iron, and is removed from time to time with a chisel. Hydrogen escapes and mixes with the coal-gas. This form of carbon is used for the pencils for arc-lights, and for the anodes of Bunsen's and bther forms of cells, and also for anodes in electro-chemical processes.

The compounds of hydrogen with nitrogen (ammonia, NH3), sulphur, selenium, and tellurium (sulphuretted, seleniuretted, or telluretted hydrogen, H2S, H2Se, H2Te), phosphorus, arsenic, and antimony (PH3, AsH3, SbH3) all of which are gases at the ordinary temperature, are decomposed if passed through a red-hot tube, giving hydrogen, which escapes along with nitrogen if ammonia be heated; or a deposit of the sulphur, etc, in the cold part of the tube if one of the other gases mentioned be employed.

The oxides of the metals ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, gold, and mercury are decomposed at a red heat; and the chlorides, bromides, iodides, and sulphides are also decomposed, except those of silver and mercury.

But none of these methods are practical plans of preparing the elements. On the other hand, as already stated, this method is generally used for the production of oxygen. This gas, although it had probably been obtained in an impure state by the older experimenters, was first produced in approximate purity by Priestley and simultaneously by Scheele in 1774. Priestley produced it by heating mercuric oxide, HgO, which decomposes thus: 2HgO = 2Hg4-O2. And Lavoisier showed that it was. possible to produce mercuric oxide by heating mercury to its boiling-point in a confined portion of air, and by separating and weighing the oxide, and subsequently heating it till it decomposed again, he proved that the oxygen had really been extracted from the air.

Certain oxides are not wholly decomposed into oxygen and element when heated, but leave an oxide containing less oxygen than that originally heated. Among these is black manganese dioxide, a mineral named pyrolusite; 3MnO2 = Mn3O4 + O0. Lead dioxide undergoes a similar change: 2PbO2= 2PbO + O2. The most important application of this method, however, is the commercial plan of producing oxygen carried out in the " Brin Company's works. In their process, barium oxide, BaO, is heated in iron tubes under pressure, air being pumped in. The barium oxide absorbs the oxygen of the air, the nitrogen being allowed to escape. After the operation has gone on for about five minutes, a considerable amount of oxygen is absorbed, barium dioxide, BaO2, being formed. The stopcocks of the pipes leading to the pump are then reversed, so that gas is exhausted from the hot iron tubes. When the pressure is reduced, the barium dioxide loses oxygen, and again returns to the state of monoxide: 2BaO2 = 2BaO + O2. The pumping is continued for about five minutes, and the valves are again reversed. The process is thus a continuous one; the oxygen is not pure, for it contains about 7 per cent, of nitrogen; but for medical use in cases of pneumonia, and for the oxy-hydrogen blowpipe, its purity is sufficient.

This method of preparing oxygen is an instance of what is termed " mass-action." The temperature is kept constant, but the pressure is raised when it is desired to cause the oxide to absorb oxygen, and lowered when it is necessary to remove the oxygen. When pressure is raised, the number of molecules of oxygen in unit volume of the space {or the mass) is increased, and hence the number in contact with the absorbing medium, the barium oxide. Combination, therefore, takes place between the two. On reducing pressure, the number per unit volume is reduced, and the compound decomposes. The phenomenon is analogous with the behaviour of a vapour when it is compressed; after a certain pressure has been reached-the vapour pressure-the vapour condenses to a liquid, and if more vapour be compressed into the same space, the pressure does not rise further, but more vapour is condensed: this is analogous to the formation of more BaO2. On pumping out vapour, the pressure does not fall, but the liquid evaporates: this is the analogue of the decomposition of the BaO2 into BaO. This law of mass-action is very generally applicable.

Certain oxides, for instance, pentoxide of iodine, I2O5, and of nitrogen, N0O5, decompose when heated. These oxides form combinations with the oxides of many other elements, such as sodium or potassium oxide, e.g. Na2O.I2O5 or NaIO3, K2O.N2O5 or KNO3; a similar compound is potassium chlorate, KC1O3 or K9O.C12O5, although the simple oxide of chlorine is unknown. Now, potassium and sodium oxides are not decomposed by heat, and when these salts are heated oxygen is evolved from the pentoxide of chlorine or iodine. These elements, however, do not escape, but replace the oxygen combined with the sodium or potassium, forming chloride of the metal, thus: K2O.C12O5 = K2O 4- Cl2 +50, and K2O+ C12 = 2KCI -f O, or, summing up both changes in one equation, 2KC1O3 = 2KCl + 3O2. Nitrate of potassium, on the other hand, loses only one atom of oxygen, leaving nitrite: 2kno0 = 2KNO2 + O2.

Oxygen is a colourless gas, without smell or taste; it can be liquefied, at a high pressure and a low temperature, to a pale blue liquid boiling at -182°. Most elements unite directly with it, often with such a rise of temperature that incandescence is produced; in such a case the phenomenon is termed " combustion." In many instances, for example when iron rusts, the oxidation is not attended by any measurable rise of temperature, although in all cases heat is evolved, but in some cases extremely slowly.

Chlorine, bromine, and iodine are generally prepared by heating together a chloride, bromide, or iodide with manganese dioxide and sulphuric acid diluted with water. Here the first change is the formation of the halogen hydride, HC1, HBr, or HI. The hydride, however, is ionised in water, and the H.Aq. at once reacts with the oxygen of the ionised MnO2, forming non-ionised water and Mn ions, thus : MnO2 + 4HCI. Aq. = MnCl4. Aq. + 2H2O. Tetrad manganese, however, appears not to be able to co-exist with chlorine in solution; hence the manganese gains an electron and becomes -Mn.Aq, the electron coming from one of the charged chlorine ions, which escapes in an electrically neutral state. Even then, however, the -Mn, though capable of existence at low temperature, still gains an electron, and a second chlorine atom is liberated in a non-ionised state. Hence the whole change is: Mn.Aq. H--Cl4.Aq. = = Mn.Aq. + -Cl2.Aq. + Cl2. Summing all these changes in one equation, we have: MnO2 + 2NaCl.Aq. + 2H2SO4.Aq. = MnSO4.Aq. + Na2SO4. Aq. + 2H2O 4- Cl2 ; or, if hydrochloric acid alone be warmed with manganese dioxide, MnO2 + 4HCI. Aq. = MnCl2.Aq. + 2H2O + Cl2.

Other substances which yield oxygen to the hydrogen of hydrochloric acid are potassium permanganate, KMnO4, potassium bichromate, K2Cr2O7, and, in general, peroxides.