pH Calculator

Enter pH, pOH, hydrogen ion concentration, or hydroxide concentration.

pOH

[H⁺] (mol/L)

[OH⁻] (mol/L)

pH —

pOH —

[H⁺] (mol/L) —

[OH⁻] (mol/L) —

Nature —

pH = −log₁₀[H⁺] | pOH = −log₁₀[OH⁻] | pH + pOH = 14 (at 25°C) | [H⁺][OH⁻] = 10⁻¹⁴

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Understanding pH, pOH, and Ion Concentrations

pH measures the acidity of a solution on a logarithmic scale from 0 to 14. A pH of 7 is neutral, below 7 is acidic, and above 7 is basic (alkaline).

The relationships are: pH = −log[H⁺], pOH = −log[OH⁻], and pH + pOH = 14 at 25°C. These hold for aqueous solutions at standard temperature.

pH = −log₁₀[H⁺], where [H⁺] is hydrogen ion concentration in mol/L. pH 7 is neutral (pure water at 25°C); below 7 is acidic; above 7 is basic. Each whole number change represents a 10-fold change in H⁺ concentration. pH 3 is 10× more acidic than pH 4, and 10,000× more acidic than pH 7.

For strong acids (HCl, HNO₃, H₂SO₄) that fully dissociate: pH = −log[acid concentration]. A 0.01 M HCl solution: pH = −log(0.01) = −log(10⁻²) = 2. For 0.1 M H₂SO₄ (diprotic, 2 H⁺ per molecule): pH = −log(0.2) = −log(2 × 10⁻¹) = 0.70.

For weak acids, the calculation requires Ka (acid dissociation constant). pH ≈ ½(pKa − log[HA]) for buffer-free weak acid solutions. Acetic acid (pKa = 4.75) at 0.1 M: pH ≈ ½(4.75 − log(0.1)) = ½(4.75 + 1) = 2.88. The actual value from the quadratic formula is 2.87 — the approximation holds well when [HA] >> Ka.

Buffer systems resist pH change: the Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) shows that a buffer is most effective within ±1 pH unit of the pKa. Biological buffers (bicarbonate/carbonic acid at pKa 6.35 in blood plasma) are selected to operate near their target pH.