The Ka Calculator determines the acid dissociation constant and pKa of a weak acid from pH and concentration data, or converts between Ka and pKa directly. Essential for chemistry students and laboratory scientists working with buffer design, pH prediction, and acid strength comparisons.
0.0000101
4.9956
-4.9956
0.001
M
1
%
0.099
M
0.0000101
4.9956
-4.9956
0.001
M
1
%
0.099
M
The calculator for the acid dissociation constant (Ka) determines the equilibrium constant for the ionization of a weak acid in aqueous solution. Ka quantifies acid strength — a larger Ka means stronger dissociation, more H⁺ ions released, and lower pH at a given concentration. This calculator derives Ka from experimental pH and concentration measurements, or converts between Ka and its logarithmic form pKa.
When a weak acid HA dissolves in water, it partially ionizes according to:
HA ⇌ H⁺ + A⁻
The equilibrium expression is: Ka = [H⁺][A⁻] / [HA]
For a monoprotic weak acid at initial concentration C with measured pH, the hydrogen ion concentration [H⁺] = 10⁻ᵖᴴ. Using the ICE table approximation (valid when dissociation is below 5%): Ka ≈ [H⁺]² / (C − [H⁺]). The ICE table calculator provides the full equilibrium treatment for cases where the approximation breaks down. The Kb calculator handles the complementary base dissociation constant.
Because Ka values span many orders of magnitude (from 10⁻² for strong weak acids to 10⁻¹⁵ for very weak acids), chemists use the logarithmic pKa scale:
pKa = −log₁₀(Ka) and Ka = 10⁻ᵖᴴ
Lower pKa means stronger acid: acetic acid (pKa 4.76) is much weaker than hydrofluoric acid (pKa 3.17), which is weaker than trichloroacetic acid (pKa 0.66). The relationship between pKa and pKb for a conjugate acid-base pair is: pKa + pKb = 14 at 25°C. Use this online calculator to interconvert Ka and pKa values for any acid. The equilibrium constant calculator handles the general case for any reversible reaction.
Ka is the foundation of buffer chemistry. The Henderson-Hasselbalch equation relates pH to pKa and the ratio of conjugate base to acid concentrations:
pH = pKa + log₁₀([A⁻] / [HA])
A buffer is most effective within ±1 pH unit of the acid's pKa. To prepare a pH 7.4 phosphate buffer, choose an acid with pKa near 7.4 — monobasic phosphate (pKa 7.20) is ideal. The ratio [A⁻]/[HA] = 10^(7.4−7.2) = 1.58, meaning 1.58 parts dibasic to 1 part monobasic phosphate. The Ksp calculator and equilibrium calculators category provide the full range of solution equilibrium tools.
The acid dissociation reaction for a weak acid HA is:
$$HA \rightleftharpoons H^+ + A^-$$
The acid dissociation constant is defined as:
$$K_a = \frac{[H^+][A^-]}{[HA]}$$
Mode 1 — pH to Ka: From pH, calculate [H⁺] = 10−pH. For a monoprotic acid, [A⁻] = [H⁺] (from stoichiometry), and [HA] = Cacid − [H⁺]. Then:
$$K_a = \frac{[H^+]^2}{C_{acid} - [H^+]}$$
Mode 2 — Concentrations to Ka: Directly use provided [H⁺] and [A⁻] values with [HA] = Cacid − [H⁺].
The pKa is calculated as:
$$pK_a = -\log_{10}(K_a)$$
Percent dissociation = ([H⁺] / Cacid) × 100%.
Ka > 1 indicates a strong acid (nearly complete dissociation). Ka << 1 indicates a weak acid. The pKa scale is more intuitive: lower pKa = stronger acid. Common weak acids have pKa values between 2 and 12. For example, acetic acid has pKa ≈ 4.76. Percent dissociation tells you the efficiency of ionization — useful for comparing acids at the same concentration.
Inputs
Results
[H⁺] = 10^(−2.87) = 0.001349 M. Ka = (0.001349)²/(0.1 − 0.001349) = 1.82 × 10⁻⁵/0.09865 = 1.84 × 10⁻⁵. pKa = 4.73. Only 1.35% of acetic acid dissociates at 0.1 M.
Inputs
Results
[H⁺] = 0.003 M, [A⁻] = 0.003 M, [HA] = 0.05 − 0.003 = 0.047 M. Ka = (0.003 × 0.003)/0.047 = 1.91 × 10⁻⁴. pKa = 3.72.
Ka (acid dissociation constant) measures the extent to which a weak acid dissociates into H⁺ and its conjugate base A⁻ in aqueous solution. Larger Ka values indicate stronger acids.
pKa = −log₁₀(Ka). It is the negative logarithm of Ka, providing a more convenient scale. Lower pKa values correspond to stronger acids.
pH measures the hydrogen ion concentration in a specific solution, while Ka is a fundamental property of the acid itself. The same acid will have different pH values at different concentrations but the same Ka (at a given temperature).
Yes. Strong acids like HCl have very large Ka values (effectively infinite). However, Ka is most useful for weak acids where Ka << 1. When Ka > 1, the acid is considered strong.
Diluting a weak acid increases its percent dissociation. At lower concentrations, a larger fraction of acid molecules dissociate, although the total [H⁺] decreases.
For a conjugate acid-base pair: Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25°C. A strong acid (large Ka) has a weak conjugate base (small Kb) and vice versa.
The Henderson-Hasselbalch equation uses pKa: pH = pKa + log([A⁻]/[HA]). A buffer works best when pH ≈ pKa, meaning equal amounts of acid and conjugate base.
Yes. Ka generally increases with temperature for endothermic dissociation processes. The van 't Hoff equation describes this temperature dependence quantitatively.
Polyprotic acids (like H₂SO₄, H₃PO₄) have multiple dissociation steps, each with its own Ka (Ka1, Ka2, Ka3). Each successive Ka is smaller than the previous one.
Ka can be determined by measuring the pH of a solution with known acid concentration, by half-equivalence point titration (pH = pKa at half-equivalence), or by conductivity measurements.
How helpful was this calculator?
Be the first to rate!
