in the relation of the halogens to hydrogen. In a gaseous state they all enter, with more or less ease, into direct combination with gaseous hydrogen—for example, in the presence of spongy platinum, forming halogen acids, HX—but the latter are far from being equally stable; hydrogen chloride is the most stable, hydrogen iodide the least so, and hydrogen bromide occupies an intermediate position. A very high temperature is required to decompose hydrogen chloride even partially, whilst hydrogen iodide is decomposed by light even at the ordinary temperature and very easily by a red heat. Hence the reaction I2 + H2 = HI + HI is very easily reversible, and consequently has a limit, and hydrogen iodide easily dissociates.73 Judging by the direct measurement of the heat evolved (22,000 heat units) in the formation of HCl, the conversion of 2HCl into H2 + Cl2 requires the expenditure of 44,000 heat units. The decomposition of 2HBr into H2 + Br2 only requires, if the bromine be obtained in a gaseous state, a consumption of about 24,000 units, whilst in the decomposition of 2HI into H2 + I2 as vapour about 3,000 heat units are evolved;74 these facts, without doubt, stand in causal connection with the great stability of hydrogen chloride, the easy decomposability of hydrogen iodide, and the intermediate properties of hydrogen bromide. From this it would be expected that chlorine is capable of decomposing water with the evolution of oxygen, whilst iodine has not the energy to produce this disengagement,75 although it is able to liberate the oxygen from the oxides of potassium and sodium, the affinity of these metals for the halogens being very considerable. For this reason oxygen, especially in compounds from which it can be evolved readily (for instance, ClHO, CrO3, &c.), easily decomposes hydrogen iodide. A mixture of hydrogen iodide and oxygen burns in the presence of an ignited substance, forming water and iodine. Drops of nitric acid in an atmosphere of hydrogen iodide cause the disengagement of violet fumes of iodine and brown fumes of nitric peroxide. In the presence of alkalis and an excess of water, however, iodine is able to effect oxidation like chlorine—that is, it decomposes water; the action is here aided by the affinity of hydrogen iodide for the alkali and water, just as sulphuric acid helps zinc to decompose water. But the relative instability of hydriodic acid is best seen in comparing the acids in a gaseous state. If the halogen acids be dissolved in water, they evolve so much heat that they approach much nearer to each other in properties. This is seen from thermochemical data, for in the formation of HX in solution (in a large excess of water) from the gaseous elements there is evolved for HCl 39,000, for HBr 32,000, and for HI 18,000 heat units.76 But it is especially evident from the fact that solutions of hydrogen bromide and iodide in water have many points in common with solutions of hydrogen chloride, both in their capacity to form hydrates and fuming solutions of constant boiling point, and in their capacity to form haloid salts, &c. by reacting on bases.
In consequence of what has been said above, it follows that hydrobromic and hydriodic acids , being substances which are but slightly stable, cannot be evolved in a gaseous state under many of those conditions under which hydrochloric acid is formed. Thus if sulphuric acid in solution acts on sodium iodide, all the same phenomena take place as with sodium chloride (a portion of the sodium iodide gives hydriodic acid, and all remains in solution), but if sodium iodide be mixed with strong sulphuric acid, then the oxygen of the latter decomposes the hydriodic acid set free, with liberation of iodine, H 2 SO 4 + 2HI = 2H 2 O + SO 2 + I 2 . This reaction takes place in the reverse direction in the presence of a large quantity of water (2,000 parts of water per 1 part of SO 2 ), in which case