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Chemistry 2 Unit 5: Buffers, titrations and solubility

Buffers resist pH changes, titrations measure unknown concentrations, and Ksp describes how much of a slightly soluble salt dissolves.

Covers: Buffers and the Henderson–Hasselbalch equation, buffer capacity, acid–base titrations and curves, indicators, solubility equilibria and Ksp, the common ion effect, precipitation

Study guide

A buffer is a solution of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH. When acid is added, the conjugate base neutralizes it; when base is added, the weak acid neutralizes it. The Henderson–Hasselbalch equation gives buffer pH: pH = pKa + log([A⁻] ÷ [HA]). When [A⁻] = [HA], pH = pKa. Buffer capacity is the amount of acid or base a buffer can absorb before its pH changes much; it's greatest when the concentrations are high and similar. A good buffer has a pKa close to the target pH. Blood is buffered by the carbonic acid–bicarbonate system.

A titration adds a solution of known concentration (the titrant) to an unknown until the equivalence point, where moles of acid equal moles of base. A titration curve plots pH versus volume. For a strong acid with a strong base, the equivalence point is at pH 7. For a weak acid with a strong base, the equivalence point is above 7 because the conjugate base is basic, and at the half-equivalence point pH = pKa. For a weak base with a strong acid, the equivalence point is below 7. Indicators change color over a pH range and should change near the equivalence point; the color change marks the end point.

Slightly soluble ionic compounds reach a solubility equilibrium described by the solubility product constant, Ksp. For AgCl(s) ⇌ Ag⁺ + Cl⁻, Ksp = [Ag⁺][Cl⁻] = s². For a salt like CaF₂, Ksp = [Ca²⁺][F⁻]² = 4s³, where s is the molar solubility. The common ion effect lowers solubility when a solution already contains one of the salt's ions. To predict precipitation, compare the ion product Q with Ksp: if Q > Ksp, a precipitate forms. Solubility of salts with basic anions, like hydroxides and carbonates, increases in acidic solutions.

When a solution is diluted or mixed with another, the moles of solute stay the same but spread through the new total volume: M₁V₁ = M₂V₂. Always calculate concentrations after mixing before comparing Q with Ksp or K. If the pH is below a weak acid's pKa, the protonated form (HA) is the main species. If the pH is above the pKa, the deprotonated form (A⁻) dominates. At pH = pKa the two forms are equal. An indicator works best when its pKa is close to the pH at the equivalence point.

To prepare a buffer, weigh the solid component on a balance, dissolve it in distilled water, add the measured volume of the conjugate acid or base, and dilute to volume in a volumetric flask. Diluting a buffer leaves its pH nearly unchanged, because the ratio of base to acid stays the same. A polyprotic acid's titration curve has one equivalence point for each acidic proton, with a buffer region and a half-equivalence point (pH = pKa) before each. Doubling the titrant's concentration halves the volume needed to reach equivalence but doesn't change the starting pH.

Key ideas

• Buffers contain a weak acid and its conjugate base.
• Henderson–Hasselbalch: pH = pKa + log([A⁻]/[HA]).
• At the half-equivalence point of a weak acid titration, pH = pKa.
• Ksp measures the solubility of slightly soluble salts.
• A common ion lowers solubility; Q > Ksp means precipitation.

Formulas and rules

Common mistakes

• Thinking every equivalence point is at pH 7; only strong acid–strong base titrations are.
• Forgetting to raise concentrations to their coefficients in Ksp.

Words to know

Buffer

A solution that resists changes in pH.

Henderson–Hasselbalch equation

pH = pKa + log([A⁻]/[HA]).

Buffer capacity

How much acid or base a buffer can absorb.

Titration

Adding a known solution to find an unknown concentration.

Titrant

The solution of known concentration in a titration.

Equivalence point

When moles of acid equal moles of base.

Half-equivalence point

Point where half the weak acid is neutralized; pH = pKa.

Titration curve

A plot of pH versus volume of titrant.

Indicator

A substance that changes color over a pH range.

End point

When an indicator changes color.

Solubility product (Ksp)

Equilibrium constant for a dissolving ionic solid.

Molar solubility

Moles of solute that dissolve per liter.

Common ion effect

Reduced solubility due to an ion already in solution.

Precipitate

A solid that forms from a solution.

Sample practice questions

13 more questions like these, from level 1 to 5, are in the practice set.