pKa Calculator

Find pKa from pH with the Henderson-Hasselbalch equation or from Ka.

Supports molar concentration units from M to yM for both calculation methods.

Updated August 24, 2026

CreatorFrank Zhao

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Introduction

The acid dissociation constant ( $K_a$ ) tells you how readily a weak acid gives up a proton in water. Because these constants span many orders of magnitude, chemists usually work with their negative base-10 logarithm, pKa:

\mathrm{pK_a}

=

-\log_{10}(K_a)

The lower the pKa, the stronger the acid — it holds its proton more loosely and donates it more easily. This calculator finds pKa in two ways: from a measured pH plus the concentrations of a weak acid and its conjugate base (the Henderson–Hasselbalch route), or straight from a known $K_a$ value.

Who is this for? Chemistry students checking acid strength, lab workers preparing buffer solutions, and anyone who has a pH or $K_a$ measurement and wants the corresponding pKa without working through logarithms by hand.

How to use

Choose the calculation method: pKa from pH when you have a buffer's pH and concentrations, or pKa from Ka when you already know the acid dissociation constant.

In pKa from pH, enter the pH, the conjugate base concentration ( $[\mathrm{A^-}]$ ), and the weak acid concentration ( $[\mathrm{HA}]$ ). Pick the same concentration unit for both — the ratio is what matters.

In pKa from Ka, just enter the acid dissociation constant ( $K_a$ ).

Read the pKa result. You can also type a pKa yourself and the calculator will solve for whichever value is missing — for example, the pH a buffer will reach at a chosen pKa and concentration ratio.

Worked example — pKa from pH

A solution has $\mathrm{pH} = 2$ , a conjugate base concentration of $3\ \mathrm{M}$ , and a weak acid concentration of $2\ \mathrm{M}$ . Enter these three values in pKa from pH mode.

\mathrm{pK_a}

=

\mathrm{pH}

-

\log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

=

2

-

\log_{10}\frac{3}{2}

=

1.824

The calculator reports pKa ≈ 1.82. A pKa below 2 means this acid is fairly strong — it readily gives up its proton.

Worked example — pKa from Ka

Suppose the acid dissociation constant is $K_a = 2$ . Switch to pKa from Ka and enter that value.

\mathrm{pK_a}

=

-\log_{10}(2)

=

-0.301

The result is pKa ≈ −0.30. A negative pKa is normal for strong acids — the larger $K_a$ is, the more negative pKa becomes.

Calculation method

The pKa from pH mode is built on the Henderson–Hasselbalch equation, which links pH, pKa, and the ratio of conjugate base to weak acid:

\mathrm{pH}

=

\mathrm{pK_a}

+

\log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

Rearranging for pKa gives the form the calculator uses:

\mathrm{pK_a}

=

\mathrm{pH}

-

\log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

Where:

$[\mathrm{A^-}]$ — molar concentration of the conjugate base
$[\mathrm{HA}]$ — molar concentration of the weak acid
$K_a$ — acid dissociation constant, used directly in the pKa from Ka mode

The two modes are fully reversible. Instead of only deriving pKa from your inputs, you can type a pKa and let the calculator solve for the missing value — the pH a buffer reaches, the concentration ratio you need, or the $K_a$ that matches a given pKa. This makes the same tool useful for planning a buffer as well as for reading one off a measurement.

Reverse example — find the pH of a buffer: with $\mathrm{pK_a} = 3$ , $[\mathrm{A^-}] = 1\ \mathrm{M}$ , and $[\mathrm{HA}] = 10\ \mathrm{M}$ , enter pKa and the two concentrations, and the calculator returns the pH.

\mathrm{pH}

=

\mathrm{pK_a}

+

\log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

=

3

+

\log_{10}\frac{1}{10}

=

2

Tips & best practices

Only the ratio of concentrations matters

The Henderson–Hasselbalch equation uses [A⁻]/[HA], so doubling both concentrations leaves pKa unchanged. What matters is that both values are in the same unit — mixing M and mM will shift the result by the log of 1000.

When [A⁻] = [HA], pH equals pKa

If the conjugate base and weak acid concentrations are equal, the log term is exactly zero and pH = pKa. This is the midpoint of a buffer — the point where it resists pH changes most strongly.

Use the reverse direction to plan a buffer

To hit a target pH, enter the acid’s pKa and one concentration, then let the calculator solve for the other. The ratio [A⁻]/[HA] tells you how much base to add relative to the acid.

Convert mass concentrations before entering them

The calculator expects molar concentrations (mol/L). If you only know a concentration in g/L, convert it to molarity using the substance’s molar mass first.

If your stock solution is labeled in grams per liter rather than moles per liter, the mass concentration to molar concentration calculator converts it using the compound's molar mass — a handy first step before you enter the concentrations here.

Frequently asked questions

What is the difference between pKa and Ka?

$K_a$ is the acid dissociation constant — a measure of how far a weak acid splits into ions at equilibrium. pKa is simply its negative base-10 logarithm, $\mathrm{pK_a} = -\log_{10}(K_a)$ . Because the logarithm compresses the scale, pKa is easier to compare across acids that differ by many orders of magnitude.

Are pH and pKa the same thing?

No. pH measures the concentration of free $\mathrm{H^+}$ ions in a solution, which depends on what is actually dissolved. pKa is a fixed property of the acid itself — how strongly it holds its proton. They only become equal when the concentrations of the acid and its conjugate base are the same.

Why can pKa be negative?

When $K_a$ is greater than 1, its logarithm is positive and pKa becomes negative. Strong acids such as hydrochloric acid have large $K_a$ values and therefore negative pKa values — the more negative, the stronger the acid.

How do I calculate pKa from pH?

Use the Henderson–Hasselbalch equation: $\mathrm{pK_a} = \mathrm{pH} - \log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}$ . You need the pH and the molar concentrations of the conjugate base and the weak acid. The calculator does this automatically in pKa from pH mode.

Limitations

•The Henderson–Hasselbalch equation assumes an ideal, dilute solution. It works best for weak acids that dissociate only slightly and for concentrations where activity is close to molarity. At very high concentrations, or for strong acids, real behavior can deviate from the estimate.
•The result is only as good as the inputs. pH measurements carry their own uncertainty, and the concentrations must be the equilibrium concentrations of the acid and its conjugate base in the same solution.
•This calculator is intended for education and lab planning. For research-grade values or work where precision matters, verify against a reference standard or an authoritative chemistry source.

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