AP Chemistry Solubility Equilibria — AP Chemistry
1. Fundamentals of Solubility Equilibrium ★★☆☆☆ ⏱ 3 min
Solubility equilibria describes the dynamic heterogeneous equilibrium that forms when a sparingly soluble ionic compound is added to water: undissolved solid is in equilibrium with its dissolved ions in a saturated solution. Unlike fully soluble compounds that dissociate completely, sparingly soluble ionic solids only dissolve a small amount, so we use equilibrium tools to quantify their solubility. Per AP Chemistry CED, this topic makes up ~15-20% of Unit 7, corresponding to ~2-4% of total AP exam score, appearing in both MCQ and FRQ, often combined with other topics like acid-base equilibria.
Exam tip: $K_{sp}$ is always defined for the dissolution reaction (solid as reactant), never for precipitation.
2. $K_{sp}$ and Molar Solubility ★★★☆☆ ⏱ 4 min
For the general dissociation of a salt $A_xB_y$: $A_xB_y(s) \rightleftharpoons xA^{y+}(aq) + yB^{x-}(aq)$, the $K_{sp}$ expression omits the solid, so relating to molar solubility $s$:
K_{sp} = (xs)^x (ys)^y = x^x y^y s^{x+y}
A key intuition: smaller $K_{sp}$ means lower solubility *only for salts with the same ion stoichiometry*. For salts with different numbers of ions, you must calculate molar solubility first to compare.
Exam tip: When comparing solubility of different salts, always calculate molar solubility explicitly. Do not rely solely on $K_{sp}$ values for salts with different stoichiometry.
3. The Common Ion Effect ★★★☆☆ ⏱ 3 min
The common ion effect describes the reduction in molar solubility of a sparingly soluble salt when one of its constituent ions is already present in solution from a second, fully soluble compound. This follows Le Chatelier’s principle: adding a product ion shifts the dissolution equilibrium back toward the solid reactant, reducing the amount of solid that can dissolve.
The calculation approach is nearly identical to that in pure water, except the initial concentration of the common ion is not zero. Because $K_{sp}$ is very small for most sparingly soluble salts, the additional concentration of the common ion from the dissolving salt is usually negligible compared to the initial concentration from the soluble compound. This allows the small-$s$ approximation, which is acceptable for AP as long as the change is less than 5% of the initial common ion concentration (the 5% rule).
Exam tip: Never apply the stoichiometric coefficient of the dissolving salt to the initial concentration of the common ion. The coefficient only applies to the change in concentration from the dissolving salt.
4. Predicting Precipitation with $Q$ vs $K_{sp}$ ★★★☆☆ ⏱ 3 min
To predict whether a precipitate will form when two solutions containing ions of a potential sparingly soluble salt are mixed, we use the reaction quotient $Q$, which has the same mathematical form as $K_{sp}$ but uses initial concentrations of the ions (immediately after mixing, before equilibrium is established). The prediction rules are:
- If $Q > K_{sp}$: Solution is supersaturated, precipitation occurs until $Q = K_{sp}$
- If $Q = K_{sp}$: Solution is saturated at equilibrium, no net precipitation forms
- If $Q < K_{sp}$: Solution is unsaturated, no precipitate forms, all ions stay dissolved
This framework is also used for selective precipitation, a separation technique where you add a precipitating ion slowly to precipitate one ion at a time, separating mixtures of ions based on different $K_{sp}$ values.
Exam tip: Always remember to dilute ion concentrations when mixing solutions. This is the most frequently missed step in precipitation prediction problems.
5. pH Effects on Solubility ★★★★☆ ⏱ 3 min
The solubility of a sparingly soluble salt depends on pH if the salt's anion is the conjugate base of a weak acid. This is because the anion will react with $H^+$ at low pH, removing it from solution, shifting the dissolution equilibrium right and increasing solubility. If the anion is the conjugate base of a strong acid, it does not react with $H^+$, so solubility is independent of pH. For hydroxide-containing salts, higher pH (higher $[OH^-]$) decreases solubility via the common ion effect, while lower pH increases solubility.
Exam tip: A common MCQ trick asks which salt's solubility changes with pH. Remember: only salts with basic anions (conjugate of weak acid) are pH-dependent; halide salts like AgCl have pH-independent solubility.
Common Pitfalls
Why: Students habitually include all reactants and products in equilibrium expressions, forgetting pure solids have an activity of 1.
Why: Students associate lower $K_{sp}$ with lower solubility, and incorrectly generalize this to all salts.
Why: Students reuse the original concentrations from before mixing, ignoring the increased total volume.
Why: Students see pH effects in hydroxide examples and generalize to all sparingly soluble salts.
Why: Students reverse the reaction when answering precipitation questions, leading to an inverted K value.
Why: Students confuse the stoichiometry of the dissolution reaction with the source of the common ion.