Short on time? This page is the recap. For explanations and worked examples, use the full study guide.
Recap
- Strong acids: HCl, HBr, HI, HClO₄, HNO₃, H₂SO₄. They ionise completely: HA + H₂O → H₃O⁺ + A⁻.
- Strong bases: group 1 hydroxides (LiOH, NaOH, KOH) and group 2 hydroxides (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂). Whatever dissolves dissociates completely.
- So the solute concentration gives [H₃O⁺] or [OH⁻] directly: no equilibrium table needed.
- In a strong acid solution: [H₃O⁺] = [A⁻] = c, [HA] ≈ 0, and [OH⁻] = K_w ÷ c.
- H₂SO₄: only the first proton ionises completely; take [H₃O⁺] = c unless told otherwise.
Key relationships
| Solution (concentration c) | Ion from solute | Then |
|---|---|---|
| Strong acid HA | [H₃O⁺] = c | pH = −log c; pOH = 14.00 − pH |
| Group 1 hydroxide MOH | [OH⁻] = c | pOH = −log c; pH = 14.00 − pOH |
| Group 2 hydroxide M(OH)₂ | [OH⁻] = 2c | pOH = −log(2c); pH = 14.00 − pOH |
| After dilution | c₂ = c₁V₁ ÷ V₂ | each tenfold dilution moves pH 1 unit towards 7 |
| From a mass | c = (m ÷ M) ÷ V | then as above |
Assumptions behind the numbers
- 25 °C, so K_w = 1.0 × 10⁻¹⁴ and pH + pOH = 14.00.
- Ionisation or dissociation is complete.
- Water’s own H₃O⁺ and OH⁻ are negligible. This fails for strong acids or bases below about 10⁻⁶ M, where the pH is close to 7 (and never on the wrong side of it).
- Volumes of dilute solutions add when mixed.
Mistakes to avoid
- Forgetting the 2 for Ca(OH)₂, Sr(OH)₂ and Ba(OH)₂.
- Calling −log[OH⁻] the pH. It is the pOH.
- Confusing strong with concentrated.
- Writing HA as a major species in a strong acid solution.
- Getting pH 8 for 10⁻⁸ M HCl. An acid cannot make water basic.
- Significant figures. Decimal places in pH = significant figures in the concentration.
Quick self-check
- pH of 0.0010 M HCl? (3.00)
- pH of 0.0050 M Ba(OH)₂? ([OH⁻] = 0.010 M, pOH 2.00, pH 12.00)
- pOH and pH of 0.20 M NaOH? (pOH 0.70, pH 13.30)
- 0.010 M HNO₃ is diluted tenfold. New pH? (3.00, up from 2.00)
Next: practice questions.