Calculate pH, pOH, [H+], and [OH−] for strong acids, weak acids, strong bases, weak bases, and buffers using the Henderson-Hasselbalch equation.
pH = pKa + log([A−]/[HA]) or pOH = pKb + log([BH+]/[B])
pH is a measure of the acidity or basicity of a solution, defined as pH = -log[H+]. The pH scale ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral at 25°C.
1) Strong Acid: 0.01 M HCl
HCl completely dissociates → [H+] = 0.01 M
pH = -log(0.01) = 2.0
2) Weak Acid: 0.1 M Acetic Acid (Ka = 1.8×10−5)
Using approximation: [H+] ≈ √(Ka × C) = √(1.8×10−5 × 0.1) = 1.34×10−3 M
pH = -log(1.34×10−3) = 2.87
3) Buffer: 0.1 M CH₃COOH + 0.1 M CH₃COO− (pKa = 4.76)
Henderson-Hasselbalch: pH = 4.76 + log(0.1/0.1) = 4.76 + 0 = 4.76
When [acid] = [base], pH = pKa
| pH Range | Classification | Examples |
|---|---|---|
| 0-3 | Strongly acidic | Stomach acid (1-2), Lemon juice (2) |
| 3-6 | Weakly acidic | Coffee (5), Rainwater (5.5) |
| 6-8 | Neutral | Pure water (7), Blood (7.4) |
| 8-11 | Weakly basic | Baking soda (9), Milk of magnesia (10) |
| 11-14 | Strongly basic | Ammonia (11), Bleach (13) |
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