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“Ethanoic acid is weaker than methanoic acid.” How weak? How much weaker? Words like “strong” and “weak” only get you so far. To compare acids properly, chemists use a number: the acid dissociation constant, Ka, and its logarithmic cousin, pKa.
Once you’re comfortable with pKa, a lot of chemistry becomes predictable: which way a proton transfer will go, which indicator to use in a titration, how to make a buffer at a chosen pH, even how a drug behaves in the stomach versus the blood.
What Ka is
A weak acid HA ionises partly in water:
HA ⇌ H⁺ + A⁻
At equilibrium, the ratio of products to reactants is constant at a given temperature:
Ka = [H⁺][A⁻] ÷ [HA]
Ka is just an equilibrium constant, specialised for acids. Its unit is mol/dm³.
- A large Ka means the products are favoured: the acid gives up its proton readily. It’s a stronger acid.
- A small Ka means most of the acid stays as HA. It’s a weaker acid.
Strong acids like HCl have such large Ka values (much greater than 1) that they’re effectively 100% ionised. Ka is really useful for weak acids, whose values typically range from about 10⁻² down to 10⁻¹².
Why we use pKa
Ka values span so many powers of ten that they’re awkward to compare. Is 1.8 × 10⁻⁵ bigger or smaller than 6.2 × 10⁻¹⁰? You can work it out, but it’s slow.
Taking a negative logarithm, exactly as with pH, turns them into manageable numbers:
pKa = −log₁₀ Ka and Ka = 10⁻ᵖᴷᵃ
| Ka | pKa |
|---|---|
| 1.8 × 10⁻⁵ | 4.74 |
| 6.2 × 10⁻¹⁰ | 9.21 |
Now the comparison is instant.
The golden rule: the lower the pKa, the stronger the acid. Each unit decrease in pKa means the acid is ten times stronger (its Ka is ten times larger).
A table of common pKa values
Values at 25 °C in water, rounded:
| Acid | Formula | pKa |
|---|---|---|
| Phosphoric acid (1st) | H₃PO₄ | 2.15 |
| Chloroethanoic acid | ClCH₂COOH | 2.87 |
| Hydrofluoric acid | HF | 3.17 |
| Nitrous acid | HNO₂ | 3.3 |
| Methanoic (formic) acid | HCOOH | 3.75 |
| Benzoic acid | C₆H₅COOH | 4.20 |
| Ethanoic (acetic) acid | CH₃COOH | 4.76 |
| Carbonic acid (1st) | H₂CO₃ | 6.35 |
| Dihydrogenphosphate | H₂PO₄⁻ | 7.20 |
| Hypochlorous acid | HOCl | 7.53 |
| Hydrocyanic acid | HCN | 9.21 |
| Ammonium ion | NH₄⁺ | 9.25 |
| Phenol | C₆H₅OH | 9.99 |
| Hydrogencarbonate | HCO₃⁻ | 10.33 |
(You’ll often see ethanoic acid’s Ka quoted as 1.8 × 10⁻⁵, which corresponds to pKa 4.74. The more precise value is about 1.75 × 10⁻⁵, pKa 4.76. Both are used in textbooks; follow whichever your data sheet gives.)
A few patterns jump out:
- Electron-withdrawing groups strengthen acids. Chloroethanoic acid (pKa 2.87) is about 80 times stronger than ethanoic acid (4.76), because the chlorine atom pulls electron density away from the –COO⁻ group and stabilises the negative charge.
- Removing each successive proton gets harder. For phosphoric acid, pKa values are 2.15, 7.20 and 12.35 for the first, second and third protons. It’s harder to pull a positive proton away from an ion that’s already negative.
- Phenol is an acid, but a weak one. Its pKa of about 10 is much higher than that of carboxylic acids.
Converting between Ka and pKa
Example 1. Benzoic acid has Ka = 6.3 × 10⁻⁵. Find pKa. pKa = −log(6.3 × 10⁻⁵) = 4.20
Example 2. Hypochlorous acid has pKa = 7.53. Find Ka. Ka = 10⁻⁷·⁵³ = 2.95 × 10⁻⁸ mol/dm³
Example 3. Which is stronger, an acid with Ka = 4.0 × 10⁻⁴ or one with pKa = 3.8? Convert the first: pKa = −log(4.0 × 10⁻⁴) = 3.40. Lower pKa means stronger, so the first acid (pKa 3.40) is stronger, by a factor of 10⁰·⁴ ≈ 2.5.
What pKa tells you: the half-ionised point
Here’s the most useful single fact about pKa. Rearranging the Ka expression and taking logs gives the Henderson–Hasselbalch equation:
pH = pKa + log([A⁻] ÷ [HA])
When the acid is exactly half ionised, [A⁻] = [HA], the log term is zero, and:
pH = pKa
So pKa is the pH at which an acid is 50% in its protonated form and 50% in its deprotonated form. Below that pH, the acid form (HA) dominates. Above it, the conjugate base (A⁻) dominates. For every pH unit away from pKa, the ratio shifts by a factor of 10.
| pH relative to pKa | Ratio [A⁻] : [HA] | Mostly… |
|---|---|---|
| pKa − 2 | 1 : 100 | HA |
| pKa − 1 | 1 : 10 | HA |
| pKa | 1 : 1 | half and half |
| pKa + 1 | 10 : 1 | A⁻ |
| pKa + 2 | 100 : 1 | A⁻ |
This is how chemists measure pKa experimentally: titrate the weak acid with strong base and read the pH at the half-way point of the titration.
Where pKa matters in practice
Choosing a buffer. A buffer works best when pH ≈ pKa, because both HA and A⁻ are present in large amounts to absorb added base or acid. To buffer near pH 4.8, use ethanoic acid and ethanoate. To buffer near pH 7.2, use the H₂PO₄⁻/HPO₄²⁻ pair, which is exactly what your cells do. See buffers explained.
Choosing an indicator. An indicator is a weak acid whose two forms have different colours. It changes colour over roughly pKa ± 1. Methyl orange (pKa ≈ 3.5) changes in acidic solution; phenolphthalein (pKa ≈ 9.4) changes in alkaline solution.
Predicting proton transfer. In a reaction between an acid and a base, equilibrium favours the side with the weaker acid, which is the acid with the higher pKa. Ethanoic acid (4.76) will give its proton to ammonia (the conjugate acid NH₄⁺ has pKa 9.25), because NH₄⁺ is the weaker acid.
Medicine. Aspirin has a pKa of about 3.5. In the very acidic stomach (pH around 1.5–2), it’s mostly in its neutral HA form, which crosses cell membranes easily. In the blood (pH 7.4), it’s almost entirely ionised. Pharmacologists use pKa values to predict where drugs are absorbed and how they travel through the body.
pKa and pKb
For a conjugate acid–base pair at 25 °C:
pKa + pKb = 14.00
So once you know the pKa of an acid, you instantly know the pKb of its conjugate base. Ethanoic acid (pKa 4.76) has a conjugate base with pKb 9.24. Our article on Ka × Kb = Kw shows where this comes from.
Common mistakes
- “Higher pKa = stronger acid.” It’s the reverse, just as a lower pH means more acidic.
- Confusing pKa with pH. pKa is a property of the acid; pH is a property of a particular solution. The same acid has one pKa but can be in solutions of any pH.
- Forgetting that Ka depends on temperature. Tables are usually quoted at 25 °C.
- Mixing up Ka values of polyprotic acids. Each proton has its own Ka: Ka1, Ka2, Ka3.
Key takeaways
- Ka is the equilibrium constant for an acid’s ionisation; pKa = −log Ka.
- Lower pKa (larger Ka) means a stronger acid; one pKa unit is a factor of 10.
- When pH = pKa, the acid is exactly half ionised.
- pKa decides the best buffer, the right indicator and the direction of proton transfer.
- Use pKa values to calculate pH in pH of weak acids, or try the pH calculator.
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