pKa Calculator

Background.

Adjust a TRIS buffer to pH 8.07 on a 25 °C bench, carry it into a 37 °C incubator, and its pH falls to about 7.75 with nothing added. Nobody made a mistake. TRIS's acid dissociation constant rises steeply with temperature, so the pH at which it buffers slides down with it. An acetate buffer making the same trip barely moves. The difference comes down to the heat absorbed when each acid gives up its proton, and to the pKa that heat shifts.

pKa is the negative base-ten logarithm of the acid dissociation constant, Ka. Read it as the pH at which an acid is exactly half dissociated: at pH = pKa the acid and its conjugate base are present in equal amounts, which is why a buffer holds its pH best within about one unit either side. The logarithm runs against intuition — a smaller pKa means a stronger acid — and each unit is a factor of ten in Ka. Formic acid at 3.75 is about ten times stronger than acetic acid at 4.756.

You can start from four kinds of data: a pKa from a table, a raw Ka, a buffer whose pH you measured and whose composition you know, or a standard Gibbs energy of ionisation from a thermodynamic compilation. Each gives the same set of results: pKa and Ka, the conjugate base's pKb and Kb, the Gibbs energy, and the pKa carried to your working temperature.

The defaults describe acetic acid as critically evaluated by NIST (Goldberg, Kishore and Lennen, 2002): pK 4.756, ΔrH° −0.41 kJ/mol and ΔrCp −142 J/(K·mol), all at 25 °C and zero ionic strength. That last condition is easy to overlook. Every constant here is a thermodynamic value for an infinitely dilute aqueous solution, and a pKa read from a titration in salty solution is a different, conditional number that typically sits a tenth of a unit or more away from it.

What is pka calculator?

The acid dissociation constant Ka is the equilibrium constant for HA ⇌ H⁺ + A⁻ in water. The IUPAC Gold Book writes it as Ka = [H⁺][A⁻] / ([HA] c°), where c° is the standard concentration of 1 mol/L; dividing by c° is what leaves Ka without units. Because Ka spans dozens of orders of magnitude across common acids, chemists use pKa = −log₁₀ Ka, for the same reason pH is used instead of hydrogen-ion concentration.

The conjugate base A⁻ has its own constant, Kb, for A⁻ + H₂O ⇌ HA + OH⁻. Multiply the two equilibria together and everything cancels except the self-ionisation of water, so Ka × Kb = Kw and pKa + pKb = pKw. A strong acid therefore has a weak conjugate base, and a table of Ka values is all you need to know about the bases too.

pKa is also a free energy. Since ΔrG° = −RT ln K, the standard Gibbs energy of ionisation is ΔrG° = R·T·ln 10·pKa, and at 25 °C each pKa unit is worth 5.708 kJ/mol. A weak acid has a positive ΔrG°: under standard conditions the undissociated form is favoured.

A pKa belongs to an acid in a particular solvent at a particular temperature and ionic strength. Change any of those and the number changes. Values in dimethyl sulfoxide or acetonitrile sit on entirely different scales from aqueous ones — acetic acid's pKa in DMSO is above 12 — and no fixed offset converts between them.

How to use this calculator.

1. 01Under 'What do you already know?', pick the data you actually have: a pKa, a Ka, a buffer pH with its composition, or a ΔrG° of ionisation in kJ/mol.
2. 02Enter that value. Ka can be typed in scientific notation, such as 1.7539e-5, and must be greater than zero. For a buffer, enter the concentrations of the acid form [HA] and the conjugate base [A⁻] in the same unit; only their ratio is used, so millimoles per flask works as well as mol/L.
3. 03Set the working temperature if you need the pKa somewhere other than 25 °C (in the buffer mode, set it to the temperature at which you read the pH), then replace ΔrH° and ΔrCp with the values for your acid. The defaults belong to acetic acid; NIST's buffer compilation gives, for example, ΔrH° = 47.45 kJ/mol and ΔrCp = −59 J/(K·mol) for TRIS, and 20.4 kJ/mol for HEPES. Enter 0 for ΔrCp if you have no value for it.
4. 04Leave pKw at 14.00 to match textbook answers, or enter 13.99, the IAPWS 2024 value for 25 °C. It affects only pKb and Kb.
5. 05Read the outputs together. pKa, Ka, pKb, Kb and ΔrG° all refer to 25 °C; only the last line is moved to your working temperature (in the buffer mode that last line is your measurement itself, and the 25 °C lines are derived from it). Quoting a pKa to two or three decimals matches the precision of the best tabulated data.

The formula.

Four exact relations and one extrapolation produce every result.

Definition: pKa = −log₁₀ Ka, and back again Ka = 10^(−pKa).

Conjugate pair: pKb = pKw − pKa and Kb = 10^(−pKb), from Ka × Kb = Kw.

Gibbs energy: ΔrG° = R × 298.15 K × ln 10 × pKa, with the gas constant R = 8.314462618 J/(mol·K), which has been exact since the 2019 SI revision fixed the Boltzmann and Avogadro constants. The factor R × 298.15 × ln 10 is 5,708.0 J/mol per pKa unit. Running it backwards turns a tabulated ΔrG° into a pKa.

Buffer route: the Henderson–Hasselbalch equation pH = pKa + log₁₀([A⁻]/[HA]) rearranged to pKa = pH − log₁₀([A⁻]/[HA]). When [A⁻] = [HA] the logarithm vanishes and the pKa is simply the measured pH — the half-equivalence point of a titration. This route treats the measured pH as a reading at your working temperature, so pH − log₁₀([A⁻]/[HA]) is the pKa at that temperature and appears unchanged as the pKa at your temperature; the 25 °C pKa and the other 25 °C results are that value moved back to 25 °C with the ΔrH° and ΔrCp you enter, and at 25 °C the two coincide.

Temperature: the van 't Hoff relation, d ln K / dT = ΔrH° / (RT²), integrated with an enthalpy that itself changes linearly with temperature, ΔrH°(T) = ΔrH°(θ) + ΔrCp (T − θ). Written for pKa with θ = 298.15 K:

pKa(T) = pKa(θ) − [ (ΔrH° − ΔrCp θ)(1/θ − 1/T) + ΔrCp ln(T/θ) ] / (R ln 10)

A positive ΔrH° — an ionisation that absorbs heat — makes pKa fall as the solution warms; that is TRIS. With ΔrCp set to zero the expression collapses to the familiar two-point van 't Hoff equation, which is adequate for amine buffers but misses the curvature carboxylic acids show. ΔrCp is held constant, so the further from 25 °C you go, the more the result depends on that assumption.

Some inputs are refused: a Ka, concentration or pKw of zero or less, a temperature at or below absolute zero, and any route that would produce a pKa outside −40 to 60.

A worked example.

Frequently asked questions.

References& sources.

1. [1]Goldberg, R. N.; Kishore, N.; Lennen, R. M. "Thermodynamic Quantities for the Ionization Reactions of Buffers." Journal of Physical and Chemical Reference Data 31(2), 231–370 (2002). NIST Standard Reference Data. Table 7.2 (Acetate): selected values at T = 298.15 K, I = 0, evaluation grade AAA — pK = 4.756, ΔrG° = 27.147 kJ/mol, ΔrH° = −0.41 kJ/mol, ΔrCp = −142 J K⁻¹ mol⁻¹. Also Table 7.7 (ammonia, pK 9.245), Table 7.51 (phosphate, pK 2.148 / 7.198 / 12.35) and Table 7.68 (TRIS, pK 8.072, ΔrH° +47.45 kJ/mol). Open access, NIST-hosted; retrieved 2026-07-29.
2. [2]CODATA 2022 recommended value of the molar gas constant R = 8.314462618… J mol⁻¹ K⁻¹, listed as exact (no standard uncertainty) because R = N_A·k is fixed by the 2019 SI redefinition. NIST Reference on Constants, Units and Uncertainty; retrieved 2026-07-29.
3. [3]International Association for the Properties of Water and Steam, IAPWS R11-24, "Revised Release on the Ionization Constant of H₂O" (Boulder, Colorado, June 2024), Table 3. Calculated pKw at the saturation/0.1 MPa row: 14.95 at 0 °C, 13.99 at 25 °C, 13.26 at 50 °C, 12.70 at 75 °C, 12.25 at 100 °C. Open access; retrieved 2026-07-29.
4. [4]IUPAC, Compendium of Chemical Terminology (the Gold Book), entry "acid dissociation constant", term identifier 15441: the equilibrium constant for the reaction of an acid, Ka = [H⁺][B⁻]/([HB] c°). Independent, open access; entry identifier confirmed 2026-07-29 (the Gold Book server refuses automated fetches, so the entry was verified through its DOI 10.1351/goldbook.15441 and public index record).
5. [5]Harned, H. S.; Ehlers, R. W. "The Dissociation Constant of Acetic Acid from 0 to 35° Centigrade." Journal of the American Chemical Society 54(4), 1350–1357 (1932); and "…from 0 to 60° Centigrade", JACS 55(2), 652–656 (1933). A Harned-cell EMF determination, methodologically independent of the calorimetric and conductometric studies NIST also compiles; Goldberg et al. call it "the definitive set of results for the pK of acetic acid at I = 0". PAYWALLED (ACS); consulted through the abstract page below and through the temperature series reproduced under reference tag 33HAR/EHL in NIST Table 7.2. It is what showed that a constant-enthalpy van 't Hoff correction gets acetic acid's temperature curve qualitatively wrong.
6. [6]Atkins, P.; de Paula, J.; Keeler, J. "Atkins' Physical Chemistry", 12th ed. (Oxford University Press, 2022), Topic 6D "Acids and bases" and Topic 6B "The response of equilibria to temperature" — derivations of Ka·Kb = Kw and of the van 't Hoff equation used for the temperature correction. PRINT / bibliographic reference; no public URL.