As regards the properties of the hydroxides, that of copper (cupric) is light blue, and of silver, brown ; chromic hydroxide is grey-green, and chromous, yellowish; manganic, brown, and manganous, very pale pink; ferric, rust brown, and ferrous, white when pure, but usually dirty green; cobaltous, dingy red, and nickelous, grass-green. The others are all white amorphous bodies, and they all yield oxides on heating. Cupric oxide is black ; even when boiled with water the hydroxide loses water and changes colour. Argentous oxide is brown; when heated to redness it loses oxygen, leaving a residue of metallic silver. Zinc oxide is yellow when hot and white when cold; cadmium oxide is a brown powder; the oxides of aluminium, scandium, yttrium, lanthanum, ytterbium, gallium, and indium, and of titanium, zirconium, thorium, germanium, and stannic oxide are white powders ; thallium oxide is a yellow powder ; tin monoxide is a black powder; that of lead (litharge, massicot) is a yellow substance, fusible at a red heat; bismuth sesquioxide is a yellowish powder; chromic, ferric, and manganic are respectively green, rust-red, and brown; chromous oxide is unknown, for any attempt to dry it results in the decomposition of water, the absorption of its oxygen by the chromous oxide which becomes chromic oxide, and the evolution of hydrogen. Ferrous hydroxide can be dried, but only with rigid exclusion of air; it is a black powder. Manganous oxide is greyish-green, cobaltous, olive-green, and nickelous also greyish-green. Manganous hydroxide must also be dried in absence of air.
These hydroxides and oxides are named bases. There are some basic oxides, which are precipitated by adding a hydroxide, such as that of sodium, to a soluble salt, and to which there is no corresponding hydroxide. This is the case with the oxides of mercury. On referring to the table of halides on p. 50, it will be seen that the chloride of mercury has the formula HgCl2. This compound, commonly called corrosive sublimate, when treated in solution with sodium hydroxide, gives a precipitate, not of hydroxide, as might be expected, but of oxide: HgCl9.Aq + 2NaOH.Aq = HgO + zNaCl.Aq + H2O. Similarly, a soluble mercurous salt, such as mercurous nitrate, Hg2(NO3)2, on treatment with an alkali gives a precipitate of mercurous oxide: Hg2(NO3)2.Aq +2NaOH.Aq = Hg2O + 2NaNOg.Aq + H2O. These are cases of relative stability; for, as has been already remarked, on boiling a solution from which cupric hydroxide has been precipitated, the blue hydroxide is changed into black oxide; other hydroxides lose their water at a still higher temperature; while those of the alkaline metals may be volatilised without decomposing.
Most of the basic oxides may also be prepared by heating the carbonates, a class of salts afterwards to be discussed. The carbonates of the alkali metals, however, are not thus decomposed; like their hydroxides, they may be volatilised without decomposition. But all other carbonates are decomposed by exposure to a red heat. The process has already been described as a method of manufacturing quicklime. Most carbonates, however, do not require the same high temperature; a dull red heat suffices. And the oxides do not, as a rule, recombine with the carbon dioxide expelled, as does lime; hence there is no danger of re-car-bonating the oxide.
The nitrates, too, of nearly all the basic metals, yield the respective oxides when they are heated to bright redness. The nitrates of the alkali metals in this instance, as in others, do not behave in this way. When heated they lose oxygen, but only at a very high temperature, forming the nitrites, a class of salts afterwards to be described. Thus, potassium nitrate undergoes the decomposition : 2KNO3 = 2KNO2 + O2. The product of heating the other nitrates, however, is the oxide, while a mixture of oxides of nitrogen is evolved. This may be supposed to take place in two stages: first, the nitrate may be imagined to decompose into the oxide and nitrogen pentoxide, thus : Zn(N0g)9 = ZnO + N2O5 ; the last compound is very easily decomposed by heat, and yields a lower oxide of nitrogen : 2N2O5 = 4NO2 + O2 ; while if the temperature is over 6oo°, which is usually exceeded in decomposing the nitrates, the nitric peroxide is partly further decomposed into nitric oxide and oxygen : 2NO2 = 2NO + O2. The products, therefore, are NO2, NO, and O2.
A metal which forms two oxides, one containing more oxygen than the other, if the nitrate of the lower oxide is heated, yields the higher oxide. Cases of this are mercury, tin, and iron. Mercurous nitrate, carefully heated, gives, not mercurous oxide, Hg2O, but mercuric oxide, HgO : HgNO3 = HgO + NO2; similarly Sn(NO3)2 yields SnO2, and not SnO ; and Fe(NO3)2, Fe2O3, and not FeO.
The sulphates require a higher temperature than the nitrates for their decomposition, consequently they are not generally used as a source of oxides. But the equivalents of magnesium, zinc, and some other metals have been determined by estimating the weight of oxide obtainable on heating a weighed amount of sulphate; and ferrous sulphate has been distilled in fireclay retorts for many years past at Nord-hausen, in Saxony, for the purpose of making " Nordhausen sulphuric acid," H2S2O7, and red oxide of iron, Fe2O3, which, made in this way, has a fine colour, and is used as a paint. When a sulphate is heated, the gas which escapes is not entirely SO3, as might be imagined from the equation : MgSO4 = MgO + SO3 ; the high temperature decomposes most of the sulphur trioxide into the dioxide, SO2, and oxygen; and the oxygen, in the case of ferrous sulphate, oxidises the FeO into Fe2O3.