Acid-base
Summary
An acid-base reaction is usually treated as proton transfer from a Brønsted-Lowry acid to a base, giving a conjugate base and conjugate acid. The Lewis view instead uses electron-pair acceptor and donor. Testable: equilibrium favors the weaker acid and weaker base.
Detail
Meaning: Three definitions are used. Arrhenius: an acid raises H⁺ (H₃O⁺) in water and a base raises OH⁻. Brønsted-Lowry: an acid donates a proton and a base accepts one. Lewis: an acid accepts an electron pair and a base donates one, which also covers reactions with no proton, such as BF₃ + NH₃. In HA + B ⇌ A⁻ + BH⁺, A⁻ is the conjugate base of HA and BH⁺ is the conjugate acid of B. A stronger acid has a weaker conjugate base, and equilibrium favors the weaker acid and weaker base.
Equation: Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25 °C, where Ka is the acid dissociation constant of an acid, Kb is the base dissociation constant of its conjugate base, and Kw is the water autoionization constant. All three are conventionally written without units; the concentrations inside them are in mol/L.
On the MCAT: Passages give Ka or pKa values and ask you to rank acid strength, label conjugate pairs, decide which side of a reaction is favored, or identify a Lewis acid. Titration curves, buffers and the pH of salt solutions build on the same ideas.
Common confusion: Strong does not mean concentrated. Strength is how completely an acid dissociates (larger Ka, lower pKa), while concentration is how much is dissolved.
Sources
- Chemistry: The Central Science
- General Chemistry: Principles and Modern Applications
Written for the MCAT and checked before it was published. For study reference only. Spotted an error? Let us know.