chemical principle/ revi~i ted
Edited by BOYD BISHOP Ciemson Universiw Ciemson, SC 29631
MURIEL
The Pitfalls of Precipitation Reactions Peter W. Slade Fraser Valley College, Abbotsford, BC, Canada and Geoffrey W. Rayner-Canham Sir Wilfred Grenfell College, Corner Brook, NF, Canada When we teach precipitation reactions, we usually present them as a simple cation-plus-anion combination. Yet we should remember that, in reality, chemistry is rarely that simplistic. There are often competing equilibria that we usually ignore more through oversight than through logical thought. From time to time articles about specific equilibria problems appear in the chemical education literature, but we feel it is worthwhile to give an overview of the difficulties.' One particular complication involves the possibility of a competing acid-base equilihrium. For example, when treating the soluhility of silver sulfate:
will cause precipitation of additional solid calcium oxalate. Addition of hydrochloric acid will reduce the hydroxide ion concentration according to: Hence the precipitate will redissolve. A more complicated case is provided by the solubility equilibrium of calcium carbonate:
First, there are the equilibria of the carbonate ion to consider:
we often overlook the fact that the sulfate ion is the conjugate base of the hydrogen sulfate ion:
As we can see from the equilibrium constant, little hydrogen sulfate ion will be produced in pure water. However, the solubility of silver sulfate will increase when the pH is decreased. If we are working with the salt of a weak acid, however. then the hvdrolvsis . . reaction becomes of maior importanci ( I ) . We can illustrate this point experimentally by taking a saturated solution of calcium oxalate. The equilibrium relationship is:
The oxalate ion is in equilihrium with the hydrogen oxalate ion (note that the second equilihrium to give oxalic acid itself has a Kb value of 1.7 X 10-l3 and is negligible):
Addition of ammonia solution will provide a source of hydroxide ion:
This will "shift to the left" the oxdatehydrogen oxalate equilihrium. The increase in concentration of oxalate ion 316
Journal of Chemical Education
The crucial competing equilibrium is that of carbonate to hydrogencarbonate. At first glance, we might think that this equilibrium "lies to the left." However, if we look a t the equilibrium constant relationship, we see that there are two concentration terms divided by one concentration term: As calcium carbonate has a low solubility, of the order of the numerical value of Kbl, then the concentration of the species on the right-hand side of the equation will he significant. In fact, we find that in a saturated solution of pure calcium carbonate, these carbonate and hydrogen carbonate ions are present in almost exactly equal proportions (2).Thus for the solubility of calcium carbonate we cannot "get away with" using only the solubility product relationship. A saturated solution of calcium carbonate in pure water will have a p H of 10.0 (3). If we allow the solution to equilibrate with the carbon dioxide in the air, three times more calcium carbonate will dissolve (3), for we are affecting the equilibrium from "the other end". This can he best appreciated by combining the Kbl, Kbl, and K p equilihria to give:
' In this paper we will be ignoring the additional cornplicatlons of actlvhles and Ion pairing.
