Last updated: July 3, 2026
pKa Calculator
Creators
Dharmendra SinghReviewers
Creators
Dharmendra SinghReviewers
Quick Answer
pKa is the negative base-10 logarithm of the acid dissociation constant Ka: pKa = −log₁₀(Ka). A lower pKa indicates a stronger acid. For a weak acid at concentration C, the pH is approximately ½(pKa − log C), and percent dissociation equals √(Ka/C) × 100. For example, acetic acid (Ka = 1.8 × 10⁻⁵) has pKa ≈ 4.74; a 0.1 M solution has pH ≈ 2.87 and ~1.34 % dissociation.
pKa equals the negative base-10 logarithm of the acid dissociation constant Ka. Lower pKa means stronger acid. For a weak acid at concentration C, pH is approximately one half of pKa minus log C.
Key Takeaways
- pKa = −log₁₀(Ka): a lower pKa means a stronger acid (larger Ka).
- Ka = 10^(−pKa) converts pKa back to Ka — the two are perfectly interconvertible.
- For a weak acid at concentration C: pH ≈ ½(pKa − log₁₀ C) from the square-root approximation.
- At the half-equivalence point of a titration, pH = pKa exactly.
- Buffers are most effective within ±1 pH unit of the acid's pKa (Henderson–Hasselbalch equation).
Creators
Dharmendra SinghReviewers
Creators
Dharmendra SinghReviewers
Formula
pKa = −log₁₀(Ka); Ka = 10^(−pKa); [H⁺] ≈ √(Ka·C); pH = −log₁₀([H⁺]); %diss = ([H⁺]/C)×100
Where:
- Ka=Acid dissociation constant(dimensionless)
- pKa=Negative log₁₀ of Ka(dimensionless)
- C=Initial acid concentration(mol/L)
- [H⁺]=Equilibrium hydrogen-ion conc.(mol/L)
- pH=Measure of acidity(dimensionless)
Worked Examples
Acetic acid — pKa and pH at 0.1 M
Acetic acid (CH₃COOH) is the classic weak-acid benchmark used in every analytical chemistry course.
- 1Calculate pKa: pKa = −log₁₀(1.8 × 10⁻⁵) = −(−4.745) ≈ 4.74.
- 2Estimate [H⁺] ≈ √(Ka × C) = √(1.8 × 10⁻⁵ × 0.1) = √(1.8 × 10⁻⁶) ≈ 1.342 × 10⁻³ mol/L.
- 3pH = −log₁₀(1.342 × 10⁻³) ≈ 2.87.
- 4% dissociation = (1.342 × 10⁻³ / 0.1) × 100 ≈ 1.34 %.
- 5The acid is classified as 'Weak acid' (3 ≤ pKa < 7).
Hydrofluoric acid — moderately strong
HF has Ka ≈ 6.8 × 10⁻⁴, making it a moderately strong acid despite being monoprotic.
- 1pKa = −log₁₀(6.8 × 10⁻⁴) ≈ 3.17.
- 2[H⁺] ≈ √(6.8 × 10⁻⁴ × 0.05) = √(3.4 × 10⁻⁵) ≈ 5.83 × 10⁻³ mol/L.
- 3pH = −log₁₀(5.83 × 10⁻³) ≈ 2.23.
- 4% dissociation ≈ (5.83 × 10⁻³ / 0.05) × 100 ≈ 11.7 %.
- 5The acid is classified as 'Moderately strong acid' (pKa ≈ 3.17).
Ammonium ion — very weak acid
NH₄⁺ acts as a weak acid donating a proton to form NH₃. Ka ≈ 5.6 × 10⁻¹⁰.
- 1pKa = −log₁₀(5.6 × 10⁻¹⁰) ≈ 9.25.
- 2[H⁺] ≈ √(5.6 × 10⁻¹⁰ × 0.25) = √(1.4 × 10⁻¹⁰) ≈ 1.18 × 10⁻⁵ mol/L.
- 3pH = −log₁₀(1.18 × 10⁻⁵) ≈ 4.93.
- 4% dissociation ≈ (1.18 × 10⁻⁵ / 0.25) × 100 ≈ 0.0047 %.
- 5The acid is classified as 'Very weak acid' (7 ≤ pKa < 12).
Introduction
The acid dissociation constant Ka measures the tendency of an acid HA to donate a proton in water: HA ⇌ H⁺ + A⁻, with Ka = [H⁺][A⁻] / [HA]. Because Ka values range over many orders of magnitude — from ~10 for strong acids to 10⁻¹⁶ for very weak ones — chemists take the negative base-10 logarithm to obtain a compact, intuitive scale: pKa = −log₁₀(Ka). A lower pKa means a stronger acid; a higher pKa means a weaker acid. This calculator converts Ka to pKa instantly, and — when you provide an initial concentration — also estimates the solution pH and percent dissociation.
The pKa formula explained
The definition pKa = −log₁₀(Ka) is a logarithmic compression of the equilibrium constant. It mirrors the pH scale: just as pH = −log₁₀([H⁺]), pKa = −log₁₀(Ka). The inverse relation Ka = 10^(−pKa) allows exact round-trip conversion. For example, acetic acid has Ka = 1.8 × 10⁻⁵, so pKa = −log₁₀(1.8 × 10⁻⁵) = 4.74.
Every one-unit increase in pKa corresponds to a 10-fold decrease in Ka (acid becomes 10× weaker).
- pKa < 0:
strong acid (fully or near-fully dissociated in water).
- pKa 3–7:
typical weak acids (acetic, lactic, formic, benzoic).
- pKa > 10:
very weak acids (ammonium ion, water itself at 15.7).
Use the pH calculator if you need to convert pH directly.
Estimating pH from pKa and concentration
For a weak monoprotic acid HA at initial concentration C, the equilibrium hydrogen-ion concentration is approximated by [H⁺] ≈ √(Ka · C), which is valid when percent dissociation is small (< ~5 %). Substituting into the pH definition gives pH = ½(pKa − log₁₀ C). This calculator applies the square-root formula directly for accuracy. For highly dissociated weak acids the exact quadratic x² + Ka·x − Ka·C = 0 is needed, but the approximation is sufficient for most laboratory work.
The approximation H⁺] ≈ √(Ka·C) overestimates [H⁺] slightly, so the computed pH will be marginally lower than the exact value. Use the [hydrogen ion concentration calculator for the exact quadratic solution.
pKa and the Henderson–Hasselbalch equation
In buffer solutions both the acid HA and its conjugate base A⁻ are present simultaneously. The Henderson–Hasselbalch equation relates pH to pKa: pH = pKa + log₁₀([A⁻]/[HA]). When the ratio A⁻]/[HA] = 1, pH = pKa exactly — meaning the buffer is at half-neutralisation. This is why pKa is the pivotal quantity for [buffer design: buffers work best within ±1 pH unit of the pKa. See also the Henderson–Hasselbalch calculator for full buffer pH calculations.
To design a buffer at a target pH, choose an acid whose pKa is within 1 unit of that pH, then adjust the [A⁻]/[HA] ratio accordingly.
Percent dissociation and Le Chatelier's principle
Percent dissociation = (H⁺] / C) × 100 tells you what fraction of the acid has actually ionised at equilibrium. For acetic acid at 0.1 M, only ~1.34 % dissociates; at 0.01 M the value rises to ~4.2 %, illustrating Le Chatelier's principle: dilution shifts the equilibrium toward more dissociation. This is closely related to the [titration calculator, where you track the ongoing ratio of HA to A⁻ as base is added.
Percent dissociation < 5 %: the square-root approximation is valid.
Percent dissociation > 5 %: solve the full quadratic for precise results.
Diluting the acid always increases percent dissociation, never decreases it.
Percent dissociation is independent of the volume — it only depends on Ka and C.
Common pKa values reference table
The table below lists representative pKa values at 25 °C. Data from NIST and standard analytical chemistry textbooks.
| Acid | Formula | Ka | pKa | Strength |
|---|---|---|---|---|
| Hydrochloric acid | HCl | ≫1 | ≪0 | Strong |
| Trichloroacetic acid | CCl₃COOH | 2.0 × 10⁻¹ | 0.70 | Moderately strong |
| Chloroacetic acid | ClCH₂COOH | 1.4 × 10⁻³ | 2.86 | Moderately strong |
| Hydrofluoric acid | HF | 6.8 × 10⁻⁴ | 3.17 | Moderately strong |
| Formic acid | HCOOH | 1.8 × 10⁻⁴ | 3.74 | Weak |
| Acetic acid | CH₃COOH | 1.8 × 10⁻⁵ | 4.74 | Weak |
| Carbonic acid (pKa1) | H₂CO₃ | 4.5 × 10⁻⁷ | 6.35 | Weak |
| Dihydrogen phosphate | H₂PO₄⁻ | 6.2 × 10⁻⁸ | 7.21 | Very weak |
| Ammonium ion | NH₄⁺ | 5.6 × 10⁻¹⁰ | 9.25 | Very weak |
| Water | H₂O | 1.8 × 10⁻¹⁶ | 15.74 | Extremely weak |
Applications in pharmacy, biology, and analytical chemistry
pKa governs how a molecule behaves in solution across different pH environments — critical for drug design, protein chemistry, and analytical separations. A drug molecule's ionisation state at physiological pH (7.4) determines membrane permeability: neutral forms cross lipid bilayers far more readily than charged forms. The buffer capacity calculator shows how the buffering power peaks sharply at pH = pKa. In HPLC method development, mobile-phase pH is typically set at least 2 units away from analyte pKa values to ensure a stable ionisation state and consistent retention time. The IUPAC Gold Book provides the formal definition of acid dissociation constants used throughout the literature.
Acid–base pharmacology: drugs are absorbed best when uncharged (Lipinski's rule).
Isoelectric point of amino acids depends on pKa of their side chains.
Buffer selection for biological assays always targets pH near the buffer's pKa.
pKa measurement by potentiometric titration (Bjerrum method) is the gold standard.
For a diprotic acid such as H₂CO₃ or H₃PO₄, there are two pKa values (pKa1 and pKa2). This calculator handles each ionisation step independently. See NIST Standard Reference Database 46 for tabulated pKa values of thousands of compounds.
Quick Reference Card
pKa — Quick Reference
Quick reference • pKa Calculator
pKa = −log₁₀(Ka) | Ka = 10^(−pKa) | pH ≈ ½(pKa − log C)Valid range: Practical range: pKa ≈ −12 (strong acids) to +20 (extremely weak acids in water)
Common Values
⚠ Watch Out
- •The approximation [H⁺] ≈ √(Ka·C) is invalid when percent dissociation exceeds ~5 %.
- •pKa values are temperature-dependent; standard values are reported at 25 °C.
- •For polyprotic acids (H₂SO₄, H₃PO₄) each ionisation step has its own pKa.
- •Ionic strength affects apparent pKa — use concentration corrections for precise work.
Pro Tips
- →If pKa ≈ target pH, the acid is an excellent buffer candidate (Henderson–Hasselbalch).
- →A one-unit difference in pKa corresponds to a 10-fold difference in Ka.
- →Measure pKa by potentiometric titration: the inflection point of the titration curve occurs at pH = pKa.
- →Use the relationship pKa + pKb = 14 (at 25 °C) to relate a weak acid's pKa to its conjugate base pKb.
FAQs
What does pKa mean?
pKa is the negative base-10 logarithm of the acid dissociation constant Ka: pKa = −log₁₀(Ka). It is a dimensionless number on the same logarithmic scale as pH. A lower pKa means a stronger acid (larger Ka); a higher pKa means a weaker acid (smaller Ka). For example, acetic acid has pKa ≈ 4.74, while hydrochloric acid has pKa ≪ 0.
How do I calculate pKa from Ka?
Apply the formula pKa = −log₁₀(Ka). On a calculator: enter Ka, press log (or log₁₀), then negate the result. For Ka = 1.8 × 10⁻⁵: log₁₀(1.8 × 10⁻⁵) = −4.745, so pKa = 4.745 ≈ 4.74. Conversely, Ka = 10^(−pKa) converts pKa back to Ka.
What is the relationship between pKa and pH?
For a pure weak-acid solution, pH ≈ ½(pKa − log₁₀ C), where C is the initial acid concentration. For a buffer (mixture of weak acid and conjugate base), the Henderson–Hasselbalch equation applies: pH = pKa + log₁₀([A⁻]/[HA]). When the buffer is at half-neutralisation ([A⁻] = [HA]), pH = pKa exactly.
What pKa value separates strong acids from weak acids?
Conventionally, strong acids have pKa < 0 (they are completely or nearly completely dissociated in water). Weak acids have pKa > 0. In practice, the boundary is fuzzy: acids with pKa 0–3 are often called 'moderately strong', while pKa > 3 are 'weak'. The IUPAC definition is thermodynamic: an acid is strong if it dissociates fully in the solvent of interest.
Why does percent dissociation increase with dilution?
The equilibrium expression Ka = [H⁺][A⁻]/[HA] must remain constant. When you dilute the solution (lower C), the denominator [HA] falls faster than the numerator product, so the equilibrium shifts right to restore Ka — producing a higher fraction of dissociated ions. This is Le Chatelier's principle applied to acid–base equilibria.
When is the approximation [H⁺] ≈ √(Ka·C) valid?
The approximation is valid when less than ~5 % of the acid dissociates, i.e., [H⁺] ≪ C. In practice this holds for most weak acids at laboratory concentrations. For stronger weak acids (pKa < 3) or very dilute solutions, the exact quadratic x² + Ka·x − Ka·C = 0 should be solved for x = [H⁺].