pH and Buffer Solutions
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Direct answer
pH = −log[H+] converts tiny hydrogen-ion concentrations into a friendly scale, anchored at 25 °C by the ionic product of water Kw = [H+][OH−] = 1.0 × 10−14, from which pH + pOH = 14 and neutrality sits at pH 7. A buffer is a mixture that resists pH change on adding small amounts of acid or base: an acidic buffer pairs a weak acid with its salt (CH3COOH with CH3COONa), a basic buffer pairs a weak base with its salt (NH4OH with NH4Cl), and their working pH is fixed by the Henderson–Hasselbalch equation pH = pKa + log([salt]/[acid]) or pOH = pKb + log([salt]/[base]). Buffer action peaks when acid and salt concentrations are equal — pH = pKa, the condition for maximum buffer capacity and the design logic behind blood's bicarbonate system.
What you must remember
- Core relations: pH = −log[H+], pOH = −log[OH−], pH + pOH = 14 at 25 °C; Kw = 1 × 10−14 at 25 °C and rises with temperature.
- Henderson equations: pH = pKa + log([salt]/[acid]); pOH = pKb + log([salt]/[base]); the ratio must be within 10:1 to 1:10 for effective buffering (pH = pKa ± 1).
- Equal-concentration rule: when [salt] = [acid], pH = pKa exactly — NEET's shortcut for instant answers.
- Mechanism in one line: added H+ is consumed by the salt's anion (CH3COO− + H+ → CH3COOH), added OH− by the acid, so the ratio barely moves.
- Blood buffer: H2CO3/HCO3− at pKa about 6.1 (or pKb side of the phosphate pair 7.2) holds plasma near pH 7.35–7.45; the bicarbonate pair does the bulk work despite the distance from its pKa, because it is backed by respiratory removal of CO2.
- Dilution immunity: diluting a buffer leaves the salt-to-acid ratio — and hence pH — unchanged, though buffer capacity falls.
- Acidic salts as buffers: a salt hydrolysis solution of CH3COONa with excess acetic acid is the same buffer; pH below 7 means acidic buffer, above 7 basic.
Computing a buffer's pH two ways
Take 0.2 mol acetic acid (pKa 4.74) with 0.05 mol sodium acetate in a litre: pH = 4.74 + log(0.05/0.2) = 4.74 + log 0.25 = 4.74 − 0.60 = 4.14. Now add 0.01 mol of HCl to the same litre. The acetate eats it: salt falls to 0.04, acid rises to 0.21, and pH = 4.74 + log(0.04/0.21) = 4.74 − 0.72 = 4.02 — a drift of just 0.12 units where pure water would have crashed to pH 2. That comparison is the entire meaning of buffer capacity, and reproducing it in the exam is two logs of arithmetic.
For a basic buffer, mirror the logic: 0.1 mol NH4OH (pKb 4.74) with 0.1 mol NH4Cl gives pOH = 4.74, hence pH = 14 − 4.74 = 9.26, and the flat equality case shows why equimolar buffers are the standard exam species.
How NEET frames acid-base questions
The paper rarely integrates; it tests relations. The pH-of-strong-acid question hides dilution logs (10−3 M HCl → pH 3; 10−8 M HCl is not pH 8 — water's own H+ caps it near 6.96, the celebrated trap). The buffer question leans on the equal-ratio shortcut or asks which mixture buffers at all — HCl + NaCl does not, both strong. Temperature items note Kw rising, so pure water at 50 °C has pH below 7 yet remains neutral (both ions equal), a conceptual favourite asserting "pH below 7 always means acidic" as false. Polyprotic and degree-of-ionisation questions belong to the equilibrium chapter but arrive here dressed in pH clothing: for a weak acid, [H+] = √(Ka × C), so pH = ½(pKa − log C), worth memorising since it halves several numericals.
Frequently asked questions
What is the pH of a buffer containing equimolar weak acid and salt?
The log term vanishes, so pH equals pKa of the acid — the maximum-capacity condition and a standard one-step answer.
Why does dilution not change a buffer's pH?
Both salt and acid concentrations fall by the same factor, leaving their ratio — and therefore the Henderson term — unchanged, though capacity decreases.
Why is the pH of 10^-8 M HCl not 8?
Water's own ionisation supplies about 10−7 M H+, dominating the negligible acid, so the solution stays mildly acidic near pH 6.96, never basic.
Which buffer system regulates human blood pH?
The carbonic acid–bicarbonate pair (H2CO3/HCO3−) is the primary regulator, assisted by phosphate and protein buffers and respiratory CO2 removal.
How does a buffer neutralise added strong acid?
The salt's conjugate base consumes the added proton (CH3COO− + H+ → CH3COOH), so [H+] barely rises and the pH shift stays under a tenth of a unit for small additions.