Additional notes
Of the oxides of sodium, that easiest to form is the peroxide, NaO or Na2O2; this is obtained when sodium is burnt in an excess of oxygen. If NaNO3 be melted, it gives Na2O2 with metallic Na. In a fused state the peroxide is reddish yellow, but it becomes almost colourless when cold. When heated with iodine vapour, it loses oxygen: Na2O2 +I2 = Na2OI2 + O. The compound Na2OI2 is akin to the compound Cu2OCl2 obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na2OI2 is soluble in water, and when acidified gives free iodine and a sodium salt. Carbonic oxide is absorbed by heated sodium peroxide with formation of sodium carbonate: Na2CO3 = Na2O2 + CO, whilst carbonic anhydride liberates oxygen from it. With nitrous oxide it reacts thus: Na2O2 +2N2O = 2NaNO2 +N2; with nitric oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO2.
Sodium peroxide, when treated with water, does not give hydrogen peroxide, because the latter in the presence of the alkali formed (Na2O2+ 2H2O = 2NaHO + H2O2) decomposes into water and oxygen. In the presence of dilute sulphuric acid it forms H2O2 (Na2O2 + H2SO4 = Na2SO4 + H2O2). Peroxide of sodium is now prepared on a large scale (by the action of air upon Na at 300°) for bleaching wool, silk &c. (when it acts in virtue of the H2O2 formed). The oxidising properties of Na2O2 under the action of heat are seen, for instance, in the fact that when heated with I it forms sodium iodate; with PbO, Na2PbO3; with pyrites, sulphates, &c. When peroxide of sodium comes into contact with water, it evolves much heat, forming H2O2, and decomposing with the disengagement of oxygen; but, as a rule, there is no explosion. But if Na2O2 be placed in contact with organic matter, such as sawdust, cotton, &c., it gives a violent explosion when heated, ignited, or acted on by water.
Peroxide of sodium forms an excellent oxidising agent for the preparation of the higher product of oxidation of Mn, Cr, W, &c., and also for oxidising the metallic sulphides. It should therefore find many applications in chemical analysis. To prepare Na2O2 on a large scale, Castner melts Na in an aluminium vessel, and at 300° passes first air deprived of a portion of its oxygen (having been already once used), and then ordinary dry air over it.
Chlorous acid, HClO2 (according to the data given by Millon, Brandau, and others) in many respects resembles hypochlorous acid, HClO, whilst they both differ from chloric and perchloric acids in their degree of stability, which is expressed, for instance, in their bleaching properties; the two higher acids do not bleach, but both the lower ones do so (oxidise at the ordinary temperature). On the other hand, chlorous acid is analogous to nitrous acid, HNO2. The anhydride of chlorous acid, Cl2O3, is not known in a pure state, but it probably occurs in admixture with chlorine dioxide, ClO2, which is obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate with such reducing substances as nitric oxide, arsenious oxide, sugar, &c. All that is at present known is that pure chlorine dioxide ClO2 (see Notes 670–674) is gradually converted into a mixture of hypochlorous and chlorous acids under the action of water (and alkalis); that is, it acts like nitric peroxide, NO2 (giving HNO3 and HNO2), or as a mixed anhydride, 2ClO2 + H2O = HClO3 + HClO2. The silver salt, AgClO2, is sparingly soluble in water.