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The equation that describes the equilibrium between solid calcium carbonate and its solvated ions is:
We can establish this equilibrium either by adding solid calcium carbonate to water or by mixing a solution that contains calcium ions with a solution that contains carbonate ions. If we add calcium carbonate to water, the solid will dissolve until the concentrations are such that the value of the reaction quotient is equal to the solubility product ( K sp = 8.7 10 –9 ). If we mix a solution of calcium nitrate, which contains Ca 2+ ions, with a solution of sodium carbonate, which contains ions, the slightly soluble ionic solid CaCO 3 will precipitate, provided that the concentrations of Ca 2+ and ions are such that Q is greater than K sp for the mixture. The reaction shifts to the left and the concentrations of the ions are reduced by formation of the solid until the value of Q equals K sp . A saturated solution in equilibrium with the undissolved solid will result. If the concentrations are such that Q is less than K sp , then the solution is not saturated and no precipitate will form.
We can compare numerical values of Q with K sp to predict whether precipitation will occur, as [link] shows. (Note: Since all forms of equilibrium constants are temperature dependent, we will assume a room temperature environment going forward in this chapter unless a different temperature value is explicitly specified.)
The concentration of Mg 2+ ( aq ) in sea water is 0.0537 M . Will Mg(OH) 2 precipitate when enough Ca(OH) 2 is added to give a [OH – ] of 0.0010 M ?
shifts to the left and forms solid Mg(OH) 2 when [Mg 2+ ] = 0.0537 M and [OH – ] = 0.0010 M . The reaction shifts to the left if Q is greater than K sp . Calculation of the reaction quotient under these conditions is shown here:
Because Q is greater than K sp ( Q = 5.4 10 –8 is larger than K sp = 8.9 10 –12 ), we can expect the reaction to shift to the left and form solid magnesium hydroxide. Mg(OH) 2 ( s ) forms until the concentrations of magnesium ion and hydroxide ion are reduced sufficiently so that the value of Q is equal to K sp .
No precipitation of CaHPO 4 ; Q = 1 10 –7 , which is less than K sp
(Note: The solution also contains Na + and ions, but when referring to solubility rules, one can see that sodium nitrate is very soluble and cannot form a precipitate.)
The solubility product is 1.6 10 –10 (see Appendix J ).
AgCl will precipitate if the reaction quotient calculated from the concentrations in the mixture of AgNO 3 and NaCl is greater than K sp . The volume doubles when we mix equal volumes of AgNO 3 and NaCl solutions, so each concentration is reduced to half its initial value. Consequently, immediately upon mixing, [Ag + ] and [Cl – ] are both equal to:
The reaction quotient, Q , is momentarily greater than K sp for AgCl, so a supersaturated solution is formed:
Since supersaturated solutions are unstable, AgCl will precipitate from the mixture until the solution returns to equilibrium, with Q equal to K sp .
No, Q = 4.0 10 –3 , which is less than K sp = 1.05 10 –2
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