Chlorates

Potassium and sodium chlorates are usually made by Liebig's .process,* in which double decomposition between calcium chlorate and a chloride, sulphate, or carbonate of the alkali metal is accomplished. If the chlorine is passed into a hot potash 01'soda solution, and the liquid evaporated, a very small yield of chlorate, with a large quantity of chloride, is obtained: ~
3 K2COa + 6 ql = 5 KCl +KCIOa +3 CO2,
There is also difficulty in separating the chloride and chlorate. In Liebig's process chlorine is passed into milk of lime at or above a temperature of 1000 C.; the apparent reaction being:-

a) 6 Ca (OH)2 +6 Cl2= 5 CaCl2+Ca (CIOa)2+6 H20.

But this may comprise two minor reactions, viz. :-

2 Ca (OH)2 +2 Cl2= CaCl2+ Ca (OCl)2+2 H20;
3 Ca (OCI)2= 2 CaCl2+ Ca (CIOa)2'

It is possible, however, that the hypochlorite may decompose thus: -

Ca (OCI)2= CaCl2+ O2,

causing waste of chlorine. To prevent this, an excess of chlorine must always be present. Theoretically, only one molecule of calcium chlorate is obtained from 12 atoms of ch10rine. This would yield two molecules of potassium chlorate, according to the reaction: -

b) Ca (CIOa)2+2 KCl = CaCl2+2 KCJOa·

But the actual yield is only about 70 per cent of the theoretical, since much of the potassium chlorate is lost in the mother-liquor. provided with agitators, and when the liquor has a density of 25° to 30° Tw., it is run off and settled. The clear solution is then mixed with the calculated quantity of potassium chloride (which should Le purified, since sodium or magnesium chloride is difficult to separate from the product); the resulting solution of potassium chlorate is evaporated in wrought iron pans, to a density of 70° Tw. tested in the hot liquor. On cooling, the chlorate crystallizes nearly pure, and the mother-liquor, containing about 20 per cent K0103, goes to waste. The crude chlorate is purified by recrystallizing from water, and the purified crystals are chained and washed carefully in a centrifugal machine, and may be sold as coarse crystals; or they are ground to a fine powder in buhrstone mills, care being taken that no organic matter, dirt, or metal (iron, etc.) gets into the mill, lest an explosion result. :K0 fire should be permitted in the building, and heating should be by steam, and lighting by electricity. The grinding mill should be at a little distance from the main building.
Magnesia * is sometimes substituted for lime, in order to increase the yield of potassium chlorate, since the latter is much less soluble in a magnesium chloride solution than in one of calcium chloride. Thus a yield of DO per cent can be obtained, owing to more complete separation when crystallizing. In this process chlorine is passed into a "milk" of powdered magnesia in water, forming a solution of magnesium chloride and chlorate, which is then concentrated until crystals of magnesium chloride (l\fgCI2• GH20) separate on cooling. 'fhese are removed, and the proportion of chlorate to chloride in the mother-liquor is about 1l\fg(CI03)2 to 2.8 l\fgCI2• The theoretical quantity of potassium ehlOl'ide is then added, and potassium chlorate separates, leaving the magnesium chloride in solution. Any excess of potassium chloride must be avoided, since it would combine with the magnesium chloride to form a crystalline precipitate of a double salt (l\fgCI2• KCl . () H20 -artificial carnallite), which would contaminate the product. The mother-liquors might be worked for chlorine, according to the "\Velclon- Pechiney process (p. 102).
Sodium chlorate is much more soluble than potassium chlorate, and is more difficult to crystallize. Pechiney devised a method for its production by which a solution of calcium chloride and chlorate, made by treating milk of lime with chlorine, is emporatecl to a density of 100° Tw. Then, by cooling the solution to exactly 12° C., a part (t) of the calcium chloride crystallizes as CaCI2 • 2 H"O, leav-Sodium sulphate is then added, which precipitates calcium sulphate, and leaves sodium chloride and chlorate in solution. On concentrating, sodium chloride crystallizes and is "fished" out of the warm li<Iuor, which, on cooling, deposits crystals of sodium chlorate; these are recrystallized.
The process depends on the fact * that a saturated solution of the mixed salts at 12° C. contains 2J.4 grams NaCl and 50.75 grams NaCI03 in 100 c.c., but at the boiling point (122° C.) a saturated solution contains only 11.5 grams NaCl, with 2·W.G grams NaCI03; hence, on cooling to 12° C., all the NaCl, and only 58.G grams of XaCI03, remain in solution, the remaining 181 grams XaUI03 crystallizing. This process has been largely used in France. Chlorate is formed by the electrolysis of a potassium chloride solution, when conducted in such a manner that the alkali formed around the cathode comes into contact with the chlorine set free at the anode, the temperature being kept above 50° C. No hypochlorite can exist at this temperature, and if the liquid is concentrated, the chlorate crystallizes.
In the Gall and Montlaur process, a 25 per cent solution of potassium chloride is decomposed by a current-density of 50 amperes per square decimeter, the tension of each bath being [) volts. About 45 per cent of the theoretical yield is obtained, the mother-liquor being again saturated with chloride and returned to the process. The reactions are probably as follows: -
2KOH + CI = KCIO + KCI + I-I~O;
3 KCIO = KCI03 +2 KCl,

the hypochlorite being instantly decomposed by the temperature of the bath. The hydrogen set free in the bath may cause part of the loss: -

KCI03 + G II = KCI + 3 II~O.

Sodium chlorate may be formed in the same way, but being more soluble, does not precipitate as crystals, and hence a larger proportion of it is destroyed by the reducing action of the hydrogen.