Free account · save your work, full figure editorSign up free

pH calculator

Q: How do you calculate pH from concentration?

For a strong acid, pH = -log₁₀ of the acid concentration (times the number of protons released). For a weak acid you also need K_a: [H⁺] ≈ √(K_a × C) when the acid is only slightly ionised, or solve the equilibrium exactly, as this calculator does.

Q: What is the Henderson–Hasselbalch equation?

pH = pK_a + log₁₀([A⁻]/[HA]). It gives the pH of a buffer from the ratio of base to acid forms, and buffers work best within one pH unit of the pK_a.

Q: Why is pH + pOH = 14?

Because water’s ion product K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. At other temperatures K_w changes, so the sum is not exactly 14.

Q: Does Tris buffer pH change with temperature?

Yes, strongly: Tris’s pK_a drops by about 0.03 per °C, so set the pH at the temperature you will use it.

pH = −log₁₀[H⁺] and, at 25 °C, pH + pOH = 14. This calculator converts between pH, pOH, [H⁺] and [OH⁻], gives the pH of strong and weak acids and bases (solving the full equilibrium, including water), and calculates buffer pH or the base-to-acid ratio for a target pH with the Henderson–Hasselbalch equation.

What it does: Converts between pH, pOH, [H⁺] and [OH⁻], gives the exact pH of strong and weak acids and bases, and designs buffers with Henderson–Hasselbalch, at 25 °C.
Formula: $\begin{aligned} &\mathrm{pH} = -\log_{10}[\mathrm{H^+}] \\ &K_\mathrm{w} = 1.0\times10^{-14} \\ &\text{weak acid: } [\mathrm{H^+}]^3 + K_\mathrm{a}[\mathrm{H^+}]^2 \\ &\quad {}- (K_\mathrm{a}C + K_\mathrm{w})[\mathrm{H^+}] - K_\mathrm{a}K_\mathrm{w} = 0 \\ &\text{buffer } \mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log_{10}([\mathrm{A^-}]/[\mathrm{HA}]) \end{aligned}$
Example: 0.1 M acetic acid (pKa 4.76) gives [H⁺] = 1.31 × 10⁻³ M, so pH = 2.88 with 1.3 % ionised.

Last updated 8 November 2026

FreeNo sign-upRuns in your browserLab calculators

Loading pH calculator…

At a glance

pH calculator at a glance
Converts	pH, pOH, [H⁺] and [OH⁻] ( $K_\mathrm{w} = 1.0\times10^{-14}$ at 25 °C)
Strong acids and bases	Fully dissociated, with water’s own ions, so very dilute solutions come out right
Weak acids and bases	Solved exactly from the charge balance, from $\mathrm{p}K_\mathrm{a}$
Buffers	$\mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log_{10}([\mathrm{A^-}]/[\mathrm{HA}])$ , and the ratio for a target pH
Limit	Concentrations stand in for activities, which is accurate for dilute solutions
Runs in	Your web browser; nothing is uploaded
Price	Free, no account needed

What it does

How to use the pH calculator

Choose a calculatorConvert a value, or calculate for a strong acid/base, a weak acid/base or a buffer.
Enter the valuesType the concentration and, for weak acids, bases and buffers, the $\mathrm{p}K_\mathrm{a}$ (or pick a common one).
Read the resultGet the pH with pOH, [H⁺] and [OH⁻], plus the fraction ionised or the buffer ratio.

How it works

$\begin{aligned} \mathrm{pH} &= -\log_{10}[\mathrm{H^+}] \\ \mathrm{pOH} &= -\log_{10}[\mathrm{OH^-}] \\ [\mathrm{H^+}][\mathrm{OH^-}] &= K_\mathrm{w} = 1.0\times10^{-14}\ \ (25\ ^\circ\mathrm{C}) \end{aligned}$ $\begin{aligned} &\text{Weak acid:} \\ &[\mathrm{H^+}]^3 + K_\mathrm{a}[\mathrm{H^+}]^2 \\ &\quad {}- (K_\mathrm{a}C + K_\mathrm{w})[\mathrm{H^+}] - K_\mathrm{a}K_\mathrm{w} = 0 \end{aligned}$ $\text{Buffer: } \mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}$

Strong acids are treated as fully dissociated with water’s own ions included, so a $10^{-8}$ M HCl solution correctly gives a pH just below 7. Weak acids and bases are solved exactly from the charge balance rather than with the $\sqrt{K_\mathrm{a}C}$ shortcut. Concentrations stand in for activities, which is accurate for dilute solutions.

Worked example

0.1 M acetic acid ( $\mathrm{p}K_\mathrm{a}$ 4.76): $[\mathrm{H^+}] = 1.31\times10^{-3}\ \mathrm{M}$ , so $\mathbf{pH = 2.88}$ , with 1.3 % of the acid ionised. A phosphate buffer at pH 7.4 ( $\mathrm{p}K_\mathrm{a}$ 7.20) needs $[\ce{HPO4^2-}]/[\ce{H2PO4-}] = 10^{0.2} = 1.58$ .

Frequently asked questions

How do you calculate pH from concentration?

For a strong acid, $\mathrm{pH} = -\log_{10}$ of the acid concentration (times the number of protons released). For a weak acid you also need $K_\mathrm{a}$ : $[\mathrm{H^+}] \approx \sqrt{K_\mathrm{a} \times C}$ when the acid is only slightly ionised, or solve the equilibrium exactly, as this calculator does.

What is the Henderson–Hasselbalch equation?

$\mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log_{10}([\mathrm{A^-}]/[\mathrm{HA}])$ . It gives the pH of a buffer from the ratio of base to acid forms, and buffers work best within one pH unit of the $\mathrm{p}K_\mathrm{a}$ .

Why is pH + pOH = 14?

Because water’s ion product $K_\mathrm{w} = [\mathrm{H^+}][\mathrm{OH^-}] = 1.0\times10^{-14}$ at 25 °C. At other temperatures $K_\mathrm{w}$ changes, so the sum is not exactly 14.

Can pH be negative?

Yes, for very concentrated strong acids (above 1 M), although activities then differ from concentrations and measured values deviate from the simple formula.

Does Tris buffer pH change with temperature?

Yes, strongly: Tris’s $\mathrm{p}K_\mathrm{a}$ drops by about 0.03 per °C, so set the pH at the temperature you will use it.

Make the whole figure in LabCanvas

Vector figure editor, slides, 3D models, protein structures and Prism-style graphs, in your browser.

Get started