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Chemistry Unit 7: Acids and bases

Acids donate H⁺ and bases accept it. The pH scale measures how acidic a solution is.

Covers: Arrhenius and Brønsted-Lowry acids and bases, pH and pOH, strong vs weak, neutralization, titration, buffers

Study guide

Acids and bases are defined in two main ways. Arrhenius acids produce H⁺ (H₃O⁺) in water, and Arrhenius bases produce OH⁻. The broader Brønsted-Lowry definition says acids donate protons and bases accept them. Every acid-base reaction forms conjugate pairs: after an acid donates a proton, what's left is its conjugate base.

The pH scale measures acidity: pH = −log[H⁺]. A pH below 7 is acidic, 7 is neutral and above 7 is basic at 25°C. Similarly pOH = −log[OH⁻], and pH + pOH = 14. Water's ion product is [H⁺][OH⁻] = 1 × 10⁻¹⁴. Each pH unit is a factor of 10, so pH 3 is ten times more acidic than pH 4.

Strong acids (HCl, HNO₃, H₂SO₄) and strong bases (NaOH, KOH) ionize completely, so [H⁺] equals the acid concentration. Weak acids, like acetic acid, ionize only partly and set up an equilibrium described by Ka; a larger Ka means a stronger acid, and pKa = −log Ka.

Neutralization reactions combine an acid and a base to form water and a salt. In a titration, a solution of known concentration is added until the equivalence point, where moles of acid equal moles of base; an indicator signals the endpoint. Use M_aV_a = M_bV_b (adjusting for the number of H⁺ or OH⁻ per formula) to find an unknown concentration. Titration curves show a sharp pH jump near the equivalence point.

Buffers resist changes in pH. They contain a weak acid and its conjugate base (or a weak base and its conjugate acid), which neutralize added base or acid. Blood is buffered near pH 7.4.

Halfway to the equivalence point, half the weak acid has been converted to its conjugate base, so [HA] = [A⁻] and pH = pKa. More generally, pH = pKa + log([A⁻]/[HA]). A buffer works best within about one pH unit of its pKa. Weak bases have a base-dissociation constant Kb for B + H₂O ⇌ BH⁺ + OH⁻. For a conjugate acid-base pair, Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25°C, so a stronger acid has a weaker conjugate base.

At the equivalence point, a strong acid titrated with a strong base gives pH 7. A weak acid titrated with a strong base gives a pH above 7, because the conjugate base that forms is basic. A weak base titrated with a strong acid gives a pH below 7. A good indicator changes color in a pH range that includes the equivalence point. Phenolphthalein (about pH 8–10) suits weak acid–strong base titrations, and methyl orange (about pH 3–4.5) suits strong acid–weak base titrations. Litmus is red in acid and blue in base.

For a weak acid with initial concentration C, Ka = x²/(C − x), where x = [H⁺]. When Ka is small, C − x ≈ C, so x ≈ √(Ka·C) and pH = −log x.

Key ideas

• Brønsted-Lowry: acids donate protons; bases accept protons.
• pH = −log[H⁺]. Lower pH = more acidic; pH 7 is neutral at 25°C.
• pH + pOH = 14.
• Strong acids (HCl, HNO₃, H₂SO₄) ionize completely; weak acids partially.
• Neutralization: acid + base → salt + water. Buffers resist pH changes.

Formulas and rules

Common mistakes

• Thinking a pH change of 1 is small. It's a tenfold change in [H⁺].
• Calling every acid strong. Most are weak.

Words to know

pH formula

pH = −log[H⁺]

pH of a neutral solution at 25°C

7

Brønsted-Lowry acid

A proton (H⁺) donor.

pH + pOH

Always 14 at 25°C.

Buffer

A solution that resists changes in pH.

Arrhenius acid

Produces H⁺ (H₃O⁺) in water.

Brønsted-Lowry base

A proton (H⁺) acceptor.

Conjugate base

What remains after an acid donates a proton.

pOH

−log[OH⁻]

Strong acid

Ionizes completely, like HCl, HNO₃, H₂SO₄.

Weak acid

Ionizes partly; described by Ka.

Ka

The acid dissociation constant; larger means stronger.

Neutralization

Acid + base → water + salt.

Equivalence point

Moles of acid equal moles of base in a titration.

Indicator

A dye that changes color near the endpoint.

Sample practice questions

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