Meaning of pH and buffer solutions
pH measures how acidic or basic a solution is. It is defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = −log[H⁺]. The scale runs from 0 to 14 at 25 °C:
- pH < 7: acidic; pH = 7: neutral (pure water); pH > 7: basic (alkaline).
- Each unit change means a tenfold change in [H⁺]. A solution of pH 3 has 10 times more H⁺ than one of pH 4.
Worked example: [H⁺] = 10⁻⁵ mol/L, so pH = −log(10⁻⁵) = 5 (acidic).
pH is measured with a pH meter (glass electrode) or with indicators and pH paper. In agriculture, soil pH controls nutrient availability: most crops grow best at about pH 6.5–7.5. Acid soils are corrected with lime, alkaline (sodic) soils with gypsum.
Buffers are solutions that resist change in pH when small amounts of acid or alkali are added. A buffer is usually a weak acid with its salt (acetic acid + sodium acetate) or a weak base with its salt (ammonium hydroxide + ammonium chloride). The pH of an acid buffer is given by the Henderson–Hasselbalch equation: pH = pKa + log([salt]/[acid]).
- When [salt] = [acid], pH = pKa.
- Example: acetic acid has pKa ≈ 4.76. If [salt]/[acid] = 10, pH = 4.76 + 1 = 5.76.
Living systems depend on buffers: blood (pH about 7.4) is buffered mainly by the bicarbonate (H₂CO₃/HCO₃⁻) system; phosphate buffers act inside cells. Soil also has buffering capacity due to clay and organic matter, so a lot of lime is needed to change the pH of a heavy, organic soil.
Key points
- pH = −log[H⁺]; neutral 7 at 25 °C.
- One pH unit = 10× change in [H⁺].
- Buffer = weak acid + its salt (or weak base + its salt).
- pH = pKa + log([salt]/[acid]).
- Blood ≈ pH 7.4, bicarbonate buffer.
Exam tips
Calculations ask for pH from [H⁺] or compare acidity of two pH values. Remember the logarithmic scale.
Common mistakes
- Thinking pH 4 is "twice" as acidic as pH 8; the difference is 10⁴ times.
- Calling a strong acid with its salt a buffer.