Chapter 11. Acids and Bases

Arrhenius: acids give H⁺, bases give OH⁻ in water.

Brønsted–Lowry: acids donate protons, bases accept them; conjugate pairs differ by one H⁺.

Lewis: acids accept electron pairs, bases donate them [2][8].

pH = −log[H₃O⁺], pOH = −log[OH⁻], and pH + pOH = 14.00 at 25 °C (K_w = 1.0 × 10⁻¹⁴).

Strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) ionize completely; weak acids have K_a < 1. For a weak acid, K_a = [H₃O⁺][A⁻]/[HA].

Buffers resist pH change. Henderson–Hasselbalch: pH = pK_a + log([A⁻]/[HA]).

Titration. At the equivalence point, moles of acid equal moles of base (adjusted for stoichiometry); indicators or pH meters detect the endpoint [9].

Review: Find the pH of 0.010 M HCl. (2.00.)