The hydroxides of the metals of the sodium group, as already mentioned, do not lose water on heating, and the oxides, therefore, cannot be thus obtained. Neither do their carbonates lose carbon dioxide, nor their nitrates oxides of nitrogen, save at an impracticable temperature. But all other basic oxides may be prepared by heating the hydroxides, carbonates, or nitrates of the metals, and a few may be obtained by heating the sulphates. Calcium and strontium oxides are generally prepared from the carbonates, which are found as minerals, named limestone and strontianite respectively. The operation of preparing " quicklime n or calcium oxide is technically, but wrongly, called " burning." Alternate layers of lime and coal are placed in a tower of brick or stone, termed a limekiln ; the coal is set on fire, and its heat expels the carbon dioxide from the carbonate: CaCO3 = CaO+CO2. If calcium carbonate be heated in a closed vessel, however, so that the carbon dioxide does not escape, the dissociation proceeds until the amount of carbon dioxide in the vessel has reached a certain proportion, which is perfectly definite for each temperature, or until the carbon dioxide has attained a certain " concentration." The reaction then stops. But if the carbon dioxide be removed as it is formed, the reaction goes on to the end, until all carbon dioxide has escaped. The draught in the kiln removes the carbon dioxide, hence the product is calcium oxide. Strontium carbonate is causticised in the same way as limestone ; but the temperature for witherite (BaCO3) is inconveniently high ; baryta is consequently prepared by heating the nitrate, Ba(NO3)2 ; it may be supposed to split into BaO and N2O5; the latter, however, decomposes even at moderate temperatures into NO2 and O ; hence the equation is : 2Ba(NO3)2 = 2BaO + 4NO2 + O2. These oxides are whitish-grey solids, volatile at the temperature of the electric arc, and combining with water with great rise of temperature to form hydroxides. The hydroxides are soluble in water-barium hydroxide most, calcium least. An aqueous solution of the former deposits crystals of a hydrate, Ba(OH)2.8H2O.

The sparing solubility of calcium hydroxide makes it possible to precipitate it by the addition of caustic alkali to a soluble salt of calcium, provided too much water is not present : CaCl2.Aq + 2NaOH.Aq = Ca(OH)2 + 2NaCl.Aq. Of course, a saturated solution of calcium hydroxide remains, hence the precipitation is not complete. This plan is applicable to the preparation of all hydroxides which are insoluble in water, unless they dissolve in excess of the caustic alkali; if they do, they are said to display " acid " properties. Beryllium and magnesium hydroxides are thus precipitated : MgCl2.Aq + 2KOH.Aq = Mg(OH)2 + 2KCl.Aq. The hydroxide may be filtered off and dried, and the white mass, on ignition, leaves the oxide as a white powder. By this means, too, the hydroxides of copper (cupric, Cu(OH)2), silver, AgOH, zinc, Zn(OH)2, cadmium, Cd(OH)2, aluminium, Al(OH)3, scandium, yttrium, lanthanum, and ytterbium, gallium, indium, and thallium, with similar formulae, titanium, zirconium, thorium, with formulae OM(OH)2, (where M stands for an element of that group ; germanium, tin (stannous, Sn(OH)2, and stannic, SnO2.nH2O), lead (plumbous, Pb(OH)2), bis-muthous, Bi(OH)3, chromic, Cr(OH)3, and chromous,

Cr(OH)2, manganic and manganous, ferric and ferrous, cobaltous and nickelous : in short, from all elements which form " basic " hydroxides. And from almost all these the oxides may be obtained by heating the hydrates to redness. Excess of the precipitant, however, must be avoided in many cases, for some of these hydroxides display acid properties if in presence of excess of alkali. Thus, for example, if excess of sodium hydroxide or potassium hydroxide be added to the solution of a soluble salt of zinc, such as the chloride, nitrate, or sulphate, the first change, as already shown, is the precipitation of the hydroxide; but on addition of excess of alkali, the precipitate redissolves, forming the compound Zn(OK)2.Aq, of which the ions are K, K, and =ZnO2; this compound is thrown down by alcohol, in which it is insoluble. It is generally named zincate of potassium. Cadmium forms a similar compound, but that of aluminium has the formula OAl(OK); the hydroxide, Al(OH)3, on losing water is transformed into the condensed hydroxide, OAl(OH), which may be termed aluminic acid, of which the hydrogen atom is exchangeable for metals. Stannous and plumbous hydroxides dissolve in excess of alkali, doubtless forming compounds similar to that of zinc ; and chromic hydroxide is soluble in cold solution of caustic alkali, forming, no doubt, a compound analogous to that of aluminium; but it is decomposed on warming, with reprecipitation of the hydroxide. The hydroxides of all these elements may also be precipitated by a solution of ammonium hydroxide, and some of them are redissolved ; but the compounds formed are of a different nature from those described in the case of zinc, etc, and will be afterwards considered.