Potassium Bicarbonate, KHCO₃

In the aqueous solutions of potassium carbonate, the ion CO₃'' is, as has already been mentioned, partly converted by the action of the water into the ion HCO₃'; the quantity transformed, however, amounts to only a few per cent of the total quantity. If, however, carbon dioxide be passed into the solution, the reaction CO₃'' + CO₂ + H₂O = 2HCO₃' takes place almost completely, and a solution of the acid or primary potassium carbonate, KHCO₃, is formed. If the solution was concentrated, the solubility product of this salt is exceeded, and it is deposited in monoclinic crystals.

The solution reacts fairly neutral, but still not so definitely as that of a salt of a strong acid, and dilute solutions exhibit even a distinctly alkaline reaction. This is due to the fact that the first ion of the dibasic carbonic acid, although much stronger than the second, is, nevertheless, the ion of a very weak acid. Hydrolysis therefore occurs, hydrion from the solvent water uniting with HCO₃' to form undissociated carbonic acid H₂CO₃, or its anhydride CO₂. The presence of the last compound can be easily demonstrated by heating the solution; even before the boiling point has been reached, bubbles of carbon dioxide are evolved. In proportion as carbon dioxide escapes, more is formed. By reason, however, of the increasing concentration of hydroxidion, the equilibrium changes so as to become more and more unfavourable to carbon dioxide, and the evolution of the gas finally sinks practically to zero. The ratio of the concentrations at which this occurs depends on the degree of dilution, more carbon dioxide being evolved the greater the dilution.

In dilute aqueous solution it has an alkaline reaction, owing to hydrolytic dissociation in accordance with the equation

KHCO₃ + H₂O = KOH + H₂O + CO₂.

The solubility in water at 20° C. is stated to be 26.31 grams per 100 grams of water, and also 33.2 grams. On heating, the salt is converted into the normal carbonate:

2KHCO₃ = K₂CO₃ + CO₂ + H₂O.

The dissociation-pressure has been investigated by Caven and Sand.

Although, therefore, acid potassium carbonate is partially decomposed in aqueous solution, the pure salt is obtained by the careful evaporation of such a solution. This is due to the fact that when some of the pure salt has separated from the saturated solution, the decomposition of the solution must become reversed, for by the separation some of the undecomposed salt is removed from participation in the equilibrium, and equilibrium can be re-established only by more of the salt being formed. This goes on until the solution is evaporated to dryness, provided there is no deficiency of one of the components, especially carbon dioxide. In order, therefore, that pure salt may be obtained with certainty, an excess of carbon dioxide must be maintained in the solution, by passing some of this gas into the solution from time to time.

The above relations receive an application in numerous similar cases.

On heating dry potassium bicarbonate, also, decomposition occurs, and, in this case, proceeds further, there remaining ultimately a residue of pure normal carbonate. The process is represented by the equation 2KHCO₃ = K₂CO₃ + H₂O + CO₂. It proceeds in a perfectly similar manner to the evaporation of a volatile liquid, and at each temperature a quite definite pressure is established; on this pressure the amounts of the two solid substances present, normal and acid potassium carbonate, have no influence. If for any given temperature that pressure has been established at which equilibrium exists, and if it be now attempted to diminish the pressure by increasing the volume, more of the acid carbonate decomposes until the pressure has again reached the former value. If, on the other hand, the volume is diminished, the pressure is only temporarily increased; carbon dioxide and aqueous vapour are absorbed until the original pressure is again established.

Such a behaviour follows from an application of the phase law. We are here dealing with three components, and there are three phases present, viz. the two solid salts and the gaseous mixture. Consequently, there still remain two degrees of freedom, and at each temperature we can obtain different pressures by suitably changing the composition of the gas mixture. So long, however, as this mixture is evolved from the bicarbonate itself, its composition is constant, since it consists of equal molecular amounts, and therefore, also, equal volumes of carbon dioxide and aqueous vapour. We have thereby disposed of one of the degrees of freedom, and only one now remains, so that to each temperature there belongs a definite pressure.

If, on the other hand, the composition of the vapour is changed, the pressure can also be changed even at a constant temperature. The law for this follows from an application of the theory of chemical equilibrium. If in the equation 2KHCO₃ = K₂CO₃ + H₂O + CO₂ the concentrations of the four substances are designated respectively by a, b, c, and d, we have the equation a² = k•bcd, where k is the equilibrium constant. Of these magnitudes, a and b are constant, since they refer to solid substances; collecting all the constants, there follows c•d = K, where K is the new constant, which is still dependent on the temperature but is no longer dependent on the concentrations or amounts of the participating substances. Hence it follows that if, at a given temperature, the concentration of the aqueous vapour in the gas mixture is increased, that of the carbon dioxide must decrease, and vice versa, until the product of the two concentrations again assumes its former value. The bicarbonate, therefore, will undergo less decomposition in an atmosphere of aqueous vapour or of carbon dioxide than in a vacuum or in a foreign gas. This rule is a general one for decompositions of this kind.