pH Calculator
Convert between pH, pOH, hydrogen ion and hydroxide ion concentration, plus buffer pH.
What do you know?
Scientific notation works: 1e-3.
pH
3.000
This solution is acidic.
- pOH
- 11.000
- [H⁺]
- 1.000e-3 M
- [OH⁻]
- 1.000e-11 M
Working
- pH = −log₁₀[H⁺] = 3.0000
- pOH = 14 − pH = 11.0000 (at 25 °C)
- [H⁺][OH⁻] = 1.00e-14 ≈ 1.0 × 10⁻¹⁴
Buffer pH (Henderson–Hasselbalch)
pH = pKa + log([A⁻]/[HA]) = 4.760
Why the scale is logarithmic
Hydrogen ion concentrations in ordinary solutions span about fourteen orders of magnitude, from roughly 1 M down to 10⁻¹⁴ M. Writing those as decimals is unusable; taking the logarithm compresses them into a scale that fits on a strip of indicator paper. The cost is that the scale is not linear — pH 4 is ten times as acidic as pH 5, and a hundred times pH 6.
Strong and weak
For a strong acid, dissociation is complete, so the hydrogen ion concentration equals the acid concentration and pH follows directly. For a weak acid it does not: only a fraction dissociates, and you need the acid dissociation constant to get there. Treating a weak acid as strong is the usual reason a calculated pH comes out too low.
Related: molar mass calculator, chemical equation balancer, periodic table.
How it works
pH is a logarithm, which is what lets a range of hydrogen ion concentrations spanning fourteen orders of magnitude fit on a scale from 0 to 14. The cost is that the scale is not linear: one pH unit is a tenfold change in acidity.
pH, pOH and the Henderson–Hasselbalch equation
pH = −log₁₀[H⁺] [H⁺] = 10⁻ᵖᴴ pH + pOH = 14 [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 25 °C) pH = pKa + log₁₀([A⁻] ÷ [HA])
- [H⁺]
- hydrogen ion concentration, mol/L
- [OH⁻]
- hydroxide ion concentration, mol/L
- pKa
- acid dissociation constant of the weak acid, as −log₁₀Ka
- [A⁻] / [HA]
- ratio of conjugate base to weak acid in a buffer
A 0.001 M solution of a strong acid
- [H⁺] = 1 × 10⁻³ mol/L (dissociation is complete)
- pH = −log₁₀(1 × 10⁻³) = 3
- pOH = 14 − 3 = 11
- [OH⁻] = 10⁻¹¹ mol/L
pH 3 — acidic, and a thousand times more acidic than pH 6.
Good to know
- The 14 in pH + pOH = 14 is a 25 °C value, not a constant. At 50 °C neutral water sits nearer pH 6.6 — still neutral, because both ion concentrations are still equal.
- pH below 0 and above 14 is chemically possible for concentrated solutions. The 0–14 range is a convention covering ordinary dilute ones.
- For a strong acid, [H⁺] equals the acid concentration. For a weak acid it does not — only a fraction dissociates, and you need Ka. Treating a weak acid as strong is the usual reason a calculated pH comes out too low.
- A buffer works best near its pKa: when [A⁻] = [HA] the log term is zero and pH equals pKa exactly.
Related Calculators
Frequently Asked Questions
What is pH?
pH is the negative base-10 logarithm of the hydrogen ion concentration in mol/L. A concentration of 10⁻³ M gives a pH of 3. The logarithm is why a change of one pH unit means a tenfold change in acidity.
How are pH and pOH related?
At 25 °C they add to 14, because the product of the hydrogen and hydroxide concentrations in water is fixed at 1.0 × 10⁻¹⁴. Both relations shift with temperature, so 14 is a room-temperature figure, not a universal constant.
What pH is neutral?
Neutral means equal hydrogen and hydroxide concentrations, which at 25 °C is pH 7. At 50 °C neutral water sits nearer pH 6.6 — still neutral, because both ion concentrations are still equal, just larger.
Can pH be negative or above 14?
Yes. The 0–14 range is a convention covering ordinary dilute solutions, not a limit. Concentrated strong acids can have a negative pH. At those concentrations the simple logarithmic relation becomes unreliable, which is why this calculator flags values outside the range.
What does the Henderson–Hasselbalch equation do?
It gives the pH of a buffer from the pKa of the weak acid and the ratio of conjugate base to acid. When the two are equal the ratio is 1, the logarithm is 0, and the pH equals the pKa — which is why buffers work best near their pKa.