Hydrogen Ion Concentration Calculator

pH = -log₁₀([H⁺]), and [H⁺] = 10^-pH. A solution with [H⁺] = 1.0 × 10^-7 mol/L has pH = 7.00 (neutral at 25 °C). Each pH unit represents a 10-fold change in [H⁺]: pH 6 has 10x more H⁺ than pH 7; pH 4 has 1000x more H⁺ than pH 7.

pOH = -log₁₀([OH^-]). At 25 °C, pH + pOH = 14.00 (= pK_w). So if pH = 11, then pOH = 3 and [OH^-] = 10^-3 = 0.001 mol/L. This relationship allows conversion between [H⁺] and [OH^-] in any aqueous solution. At 37 °C (body temperature), pK_w ≈ 13.62, so neutral pH is about 6.81, not 7.00.

Neutral means [H⁺] = [OH^-], which occurs at pH = pK_w/2. K_w increases with temperature: at 0 °C K_w ≈ 1.1 × 10^-15 (neutral pH ≈ 7.47), at 25 °C K_w = 1.0 × 10^-14 (neutral pH = 7.00), at 50 °C K_w ≈ 5.5 × 10^-14 (neutral pH ≈ 6.63). A pH of 7.0 at 37 °C is actually slightly basic.

Take the negative base-10 logarithm of the hydrogen ion concentration: pH = -log₁₀([H⁺]). For a clean power of ten this is easy - [H⁺] = 10^-4 mol/L gives pH 4.00. For a value like [H⁺] = 3.2 × 10^-5 mol/L, enter it into a calculator's log key: -log(3.2 × 10^-5) ≈ 4.49. To go the other way, [H⁺] = 10^-pH. The same steps convert [OH^-] to pOH using -log₁₀([OH^-]).

A negative pH just means [H⁺] is greater than 1 mol/L, since pH = -log₁₀([H⁺]). Concentrated strong acids reach it: about 12 mol/L hydrochloric acid has a pH near -1.1. The pH scale has no hard floor of 0 or ceiling of 14 - those are only the values for 1 mol/L acid and base. In practice, at such high concentrations activity differs sharply from concentration, so a measured negative pH is approximate, but the value is chemically real.