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Acid-Base Equilibrium
Acid-Base Equilibrium
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Acid base equilibrium describes how acids and bases transfer protons and settle into a balance between reactants and products. It matters because pH controls many chemical systems, including blood chemistry, ocean chemistry, food preservation, and laboratory titrations. At equilibrium, reactions have not stopped, but the forward and reverse reaction rates are equal. This makes acid base chemistry a dynamic process rather than a one-way change.
A weak acid such as acetic acid only partially ionizes in water, so both acid molecules and ions remain in solution. The strength of an acid or base is measured by equilibrium constants such as Ka and Kb, which show how far the reaction favors products. Buffers use a weak acid and its conjugate base to resist pH changes when small amounts of acid or base are added. Understanding these relationships helps students predict pH, compare acid strength, and explain titration curves.
Key Facts
- Acid dissociation: HA + H2O ⇌ H3O+ + A-
- Base reaction: B + H2O ⇌ BH+ + OH-
- Acid ionization constant: Ka = [H3O+][A-] / [HA]
- Base ionization constant: Kb = [BH+][OH-] / [B]
- Water ion product at 25°C: Kw = [H3O+][OH-] = 1.0 × 10^-14
- pH = -log[H3O+] and pOH = -log[OH-]
Vocabulary
- Equilibrium
- A state in which the forward and reverse reactions continue at equal rates, so concentrations stay constant.
- Weak acid
- An acid that only partially donates protons in water and forms an equilibrium mixture.
- Conjugate base
- The particle left after an acid donates a proton.
- Ka
- The acid ionization constant that measures how strongly an acid produces H3O+ in water.
- Buffer
- A solution that resists changes in pH because it contains a weak acid and its conjugate base or a weak base and its conjugate acid.
Common Mistakes to Avoid
- Treating equilibrium as a stopped reaction is wrong because particles continue reacting in both directions even when concentrations are constant.
- Assuming all acids fully ionize is wrong because weak acids only partially dissociate and must be described using Ka.
- Forgetting to include water correctly is wrong because water can act as an acid or base, but it is usually omitted from Ka and Kb expressions because it is a pure liquid.
- Using initial concentrations as equilibrium concentrations is wrong because reaction changes must be accounted for, often with an ICE table.
Practice Questions
- 1 A solution has [H3O+] = 2.5 × 10^-4 M. Calculate its pH.
- 2 For the equilibrium HA ⇌ H+ + A-, the equilibrium concentrations are [H+] = 1.0 × 10^-3 M, [A-] = 1.0 × 10^-3 M, and [HA] = 0.20 M. Calculate Ka.
- 3 A weak acid solution is diluted with water. Explain how the equilibrium shifts and why the percent ionization may increase even though the solution becomes less acidic.