The formula
pH is defined as the negative base-10 logarithm of the hydrogen ion activity in solution. In practice, for dilute solutions:
pH = −log₁₀[H⁺]
Pure water at 25 °C has an H⁺ concentration of 1 × 10⁻⁷ mol/L, giving pH = 7 (neutral). A strong acid at 1 mol/L gives pH = 0. A strong base at 1 mol/L gives pH = 14 (since pOH = −log₁₀[OH⁻] and pH + pOH = 14 at 25 °C).
The logarithmic nature means a one-unit pH change represents a tenfold change in H⁺ concentration. For the theory of why acids and bases affect H⁺ concentration, see our acids and bases guide.
Common pH values
Buffer solutions
A buffer resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (usually from its salt) in equilibrium:
HA ⇌ H⁺ + A⁻
- If H⁺ is added, the equilibrium shifts left (A⁻ absorbs H⁺).
- If OH⁻ is added, it reacts with H⁺, and the equilibrium shifts right to replenish H⁺.
In both cases, the change in pH is minimised. The Henderson–Hasselbalch equation relates buffer pH to the ratio of conjugate base to acid:
pH = pKa + log₁₀([A⁻] / [HA])
Human blood is buffered at pH 7.35–7.45 by the carbonic acid– bicarbonate system. Enzyme reactions require stable pH — see our enzyme glossary entry for how pH affects the active site shape.
Importance in biology and industry
pH affects nearly every biological process. Most enzymes have a narrow pH optimum — stomach pepsin works best at pH 2; most intracellular enzymes at pH 7–8. Blood pH outside 7.35–7.45 (acidosis or alkalosis) is a medical emergency.
In industry, pH control is critical in:
- Fermentation (brewing, cheese-making, yoghurt production)
- Water treatment (drinking water is maintained at pH 6.5–8.5)
- Agricultural soil management (most crops grow best at pH 6–7)
- Pharmaceutical manufacturing (drug stability depends on pH)
For catalysis in acidic or basic environments, see our catalyst glossary entry.