Water Ionization Constant pKw Normal and Heavy Water vs Temperature

Autoprotolysis describes the reaction of water with itself — two water molecules exchange a proton:

2 H₂O ↔ H₃O⁺ + OH⁻

An equilibrium constant can be defined for this reaction:

K_eq = a(H₃O⁺) × a(OH⁻) / a(H₂O)²

where a is the chemical activity of each species. Because most acid–base solutions are very dilute, the activity of water is approximated as unity, and the activities of the solute ions are approximately equal to their molar concentrations.

Under the dilute-solution approximation, the water ionization constant (also called the dissociation constant, ion product, or autoprotolysis constant) simplifies to:

K_w = [H₃O⁺] [OH⁻]

A commonly cited value is K_w = 1.00×10⁻¹⁴ at 25 °C (77 °F). It is expressed as a negative base-10 logarithm:

pK_w = −log₁₀(K_w)

Because H₃O⁺ and OH⁻ are produced in a 1:1 molar ratio, [H₃O⁺] = [OH⁻] = √K_w. This is the basis for defining neutral pH as pK_w/2 — at 25 °C this gives pH 7.

pK_w is not constant — it decreases as temperature rises, meaning water ionizes more at higher temperatures. At 0 °C, pK_w ≈ 14.95, giving a neutral pH of about 7.47. At 100 °C, pK_w ≈ 12.25, giving a neutral pH of about 6.13. Above 250 °C, ionization decreases slightly and pK_w rises again as the dielectric constant of water drops near its critical point, reducing the ability to stabilize separated ions.

Heavy water self-ionizes less than normal water because the O–D bond is stronger than O–H — deuterium is heavier than protium, so the zero-point vibrational energy is lower and the bond requires more energy to break. This produces pK_w values roughly 0.9–1.0 units higher for D₂O compared to H₂O across 0–100 °C. Published D₂O data is available only to 100 °C / 212 °F.