pKa Calculator
Convert between Ka and pKa, or find pKa from a measured pH and the concentrations of acid and conjugate base using the Henderson-Hasselbalch equation. Choose a mode below, enter the values you know, and get the result with a full show-your-work breakdown.
Formula
Worked example
Acetic acid has Ka = 1.8 x 10^-5. Then pKa = -log10(1.8 x 10^-5) = 4.74. Alternatively, if a buffer of acetic acid has pH 5.0 with [A-] = 0.18 mol/L and [HA] = 0.10 mol/L, Henderson-Hasselbalch gives pKa = 5.0 - log10(1.8) = 5.0 - 0.255 = 4.74.
What Ka and pKa measure
The acid dissociation constant Ka is the equilibrium constant for an acid releasing a proton in water: HA + H2O ⇌ H3O+ + A-. Ka equals the product of the hydronium-ion and conjugate-base concentrations divided by the undissociated acid concentration. A larger Ka means the equilibrium lies further toward the dissociated products, so the acid is stronger. Because Ka values span many orders of magnitude, from above 10^6 for strong acids like HCl to below 10^-14 for extremely weak ones, chemists usually report the more compact pKa, defined as the negative base-10 logarithm of Ka. The two forms carry identical information; pKa is simply easier to tabulate, compare, and use in the Henderson-Hasselbalch equation.
Reading the pKa scale
Because pKa is a negative logarithm, the relationship runs in reverse: a small pKa corresponds to a large Ka and a strong acid, while a large pKa corresponds to a tiny Ka and a very weak acid. The logarithmic nature means each whole-number change in pKa reflects a tenfold change in Ka, so an acid with pKa 3 is one hundred times stronger than one with pKa 5. Strong mineral acids like hydrochloric acid (pKa around -6) and hydrobromic acid (pKa around -9) sit far below zero. Common weak acids like acetic acid cluster near 4 to 5. Very weak acids like the ammonium ion (pKa 9.25) barely donate their proton at all, and water itself has pKa around 15.7.
Henderson-Hasselbalch and buffer design
The Henderson-Hasselbalch equation, pH = pKa + log([A-]/[HA]), links pKa to the pH of a buffer containing both the acid and its conjugate base. When the pH equals the pKa, the ratio [A-]/[HA] equals 1, meaning the acid is exactly 50% dissociated. This is the half-equivalence point, and it is the point of maximum buffer capacity: the solution resists pH change most effectively within about one unit of the pKa. Rearranging the equation lets you solve for pKa from a measured pH and known concentrations, which is the Henderson-Hasselbalch mode in this calculator. Drug formulation chemists, biochemists managing enzyme assays, and students designing titration buffers all rely on this relationship routinely.
Estimating Ka from pH and initial acid concentration
For a pure monoprotic weak acid dissolved in water with no buffer components, the measured solution pH gives another route to Ka. At equilibrium, every proton donated by the acid becomes one hydronium ion and one conjugate base ion, so [H+] = [A-] = 10^-pH. The remaining undissociated acid is [HA] = [HA]0 - [H+]. Substituting into the Ka expression gives Ka = [H+]^2 / ([HA]0 - [H+]). This method requires [HA]0 to be substantially larger than [H+], the weak acid assumption; if they are similar in magnitude the approximation breaks down and an iterative calculation is needed. The advanced mode in this calculator applies this formula directly.
Functional groups and typical pKa ranges
The pKa of a compound depends on its functional group and any electron-withdrawing or donating substituents nearby. Strong mineral acids like HCl and H2SO4 have large negative pKa values because they dissociate almost completely. Carboxylic acids (-COOH) cluster around pKa 4 to 5 in the absence of strong substituents. Phenols (-OH on a benzene ring) are weaker, typically 9 to 10. Protonated amines (-NH3+) have pKa values in the range 9 to 11. Thiols (-SH) sit near 10 to 12. Electron-withdrawing groups like halogens or nitro groups adjacent to an acidic site lower the pKa (strengthen the acid) by stabilizing the conjugate base; electron-donating groups raise it.
pKa and Ka of common acids and functional groups
| Acid / Group | Representative formula | Ka (approx) | pKa | Strength |
|---|---|---|---|---|
| Hydroiodic acid | HI | ~10^10 | -10 | Strong |
| Hydrobromic acid | HBr | ~10^9 | -9 | Strong |
| Hydrochloric acid | HCl | ~10^6 | -6 | Strong |
| Sulfuric acid (pKa1) | H2SO4 | ~10^3 | -3 | Strong |
| Hydrofluoric acid | HF | 6.3 x 10^-4 | 3.2 | Moderate |
| Formic acid | HCOOH | 1.77 x 10^-4 | 3.75 | Weak |
| Benzoic acid | C6H5COOH | 6.3 x 10^-5 | 4.2 | Weak |
| Acetic acid | CH3COOH | 1.8 x 10^-5 | 4.74 | Weak |
| Carbonic acid (pKa1) | H2CO3 | 4.47 x 10^-7 | 6.35 | Weak |
| Phosphoric acid (pKa2) | H2PO4- | 6.2 x 10^-8 | 7.2 | Very weak |
| Hydrocyanic acid | HCN | 6.2 x 10^-10 | 9.21 | Very weak |
| Ammonium ion | NH4+ | 5.6 x 10^-10 | 9.25 | Very weak |
| Phenol | C6H5OH | 1.0 x 10^-10 | 10 | Very weak |
| Carbonic acid (pKa2) | HCO3- | 4.7 x 10^-11 | 10.33 | Very weak |
| Thiol (generic) | R-SH | ~10^-13 | 13 | Extremely weak |
| Water | H2O | ~2 x 10^-16 | 15.7 | Extremely weak |
| Alcohol (ethanol) | CH3CH2OH | ~10^-17 | 17 | Non-acid |
| Alkyne (terminal) | R-C≡C-H | ~10^-25 | 25 | Non-acid |
Approximate values in water at 25 C. Use these for quick reference; consult primary literature for precise analytical work.
Frequently asked questions
How do I convert Ka to pKa?
Take the negative base-10 logarithm of Ka: pKa = -log10(Ka). For example, an acid with Ka = 1.0 x 10^-5 has pKa = -log10(1.0 x 10^-5) = 5.0. To go the other way, raise ten to the negative pKa: Ka = 10^(-pKa). Because the operation is a logarithm, Ka must be a positive number; a negative or zero Ka has no defined pKa.
Does a higher pKa mean a stronger or weaker acid?
A higher pKa means a weaker acid. Since pKa is the negative logarithm of Ka, the two run in opposite directions: as Ka gets smaller the acid dissociates less and the pKa gets larger. An acid with pKa 9 is far weaker than one with pKa 2. A one-unit increase in pKa corresponds to a tenfold decrease in Ka.
What is the Henderson-Hasselbalch equation and when do I use it?
The Henderson-Hasselbalch equation is pH = pKa + log([A-]/[HA]), where [A-] is the conjugate base concentration and [HA] is the undissociated acid concentration. You use it to design buffers (choose an acid whose pKa is close to your target pH), to predict the pH of a buffer from known concentrations, or to back-calculate pKa from a measured pH when you know the concentrations of both forms. It is most accurate when both [A-] and [HA] are well above the [H+] concentration.
How do I find pKa from pH alone?
pH alone is not enough; you also need concentration information. If you have [A-] and [HA], rearrange Henderson-Hasselbalch: pKa = pH - log([A-]/[HA]). If you have a pure weak acid solution with known initial concentration [HA]0 and you measure the pH, use the equilibrium formula Ka = [H+]^2 / ([HA]0 - [H+]) and then pKa = -log10(Ka). This calculator supports both routes.
What is the relationship between pKa and pH?
pKa is a fixed property of the acid, while pH describes the solution it sits in. When the pH equals the pKa, the acid is exactly half-dissociated, that is, [A-] = [HA]. This is the most useful operating point for a buffer because the solution resists pH changes most effectively within about one pKa unit of the target pH. The Henderson-Hasselbalch equation connects the two: pH = pKa + log([A-]/[HA]).
What factors affect the pKa of an acid?
Several structural factors shift pKa. Electronegativity of the atom bearing the proton matters: O-H and N-H bonds are more acidic than C-H bonds because oxygen and nitrogen stabilize the negative charge on the conjugate base. Electron-withdrawing substituents (halogens, nitro groups, carbonyls) adjacent to the acidic site delocalize the negative charge and lower the pKa (strengthen the acid). Resonance stabilization of the conjugate base lowers pKa further, which is why carboxylic acids (pKa 4 to 5) are far more acidic than simple alcohols (pKa 16 to 17). Solvent and temperature also shift pKa, so tabulated values typically refer to water at 25 C.