What Is Ka? The Acid Dissociation Constant Explained
"Strong" and "weak" are useful words right up to the moment someone asks how weak. Is hydrofluoric acid weaker than formic acid? By how much? That is what Ka is for: it turns a vague adjective into a number you can calculate with.
The short answer: Ka, the acid dissociation constant, is the equilibrium constant for an acid giving up its proton to water. A larger Ka means a stronger acid — at the same concentration, more of it is ionised at equilibrium — and because the numbers are tiny, they are usually quoted as pKa = −log Ka, where a smaller pKa means a stronger acid.
What Ka actually is
Every weak acid in water sets up an equilibrium:
HA + H₂O ⇌ H₃O⁺ + A⁻
Write the equilibrium expression for it and you get Ka:
Ka = [H₃O⁺][A⁻] ÷ [HA]
Water does not appear, even though it is a reactant. It is the solvent, present in enormous excess, and its concentration barely changes — so by convention it is folded into the constant rather than written out.
Read the expression as a ratio: products over reactant, or ionised over un-ionised. A big Ka means the top of the fraction wins, so most of the acid has split apart. A tiny Ka means the bottom wins and the acid is mostly sitting there intact.
Some real numbers
| Acid | Ka (25 °C) | pKa | Verdict |
|---|---|---|---|
| Hydrofluoric, HF | 6.8 × 10⁻⁴ | 3.17 | Strongest of the weak acids here |
| Formic, HCOOH | 1.8 × 10⁻⁴ | 3.75 | |
| Acetic, CH₃COOH | 1.8 × 10⁻⁵ | 4.74 | The classic textbook weak acid |
| Carbonic, H₂CO₃ (first H) | 4.5 × 10⁻⁷ | 6.35 | |
| Hydrocyanic, HCN | 6.2 × 10⁻¹⁰ | 9.21 | Barely acidic |
| Ammonium, NH₄⁺ | 5.6 × 10⁻¹⁰ | 9.25 | Barely acidic |
HF is nearly 40 times stronger than acetic acid and roughly a million times stronger than HCN — a gap the words "weak acid" hide completely. So, to answer the question at the top: hydrofluoric acid is not weaker than formic acid, it is about four times stronger — even though both are filed under "weak".
Strong acids such as HCl and HNO₃ ionise essentially completely, so their equilibrium lies so far right that Ka is not a useful number and is rarely tabulated.
Why pKa is more convenient
Comparing 6.8 × 10⁻⁴ with 6.2 × 10⁻¹⁰ in your head is unpleasant. Taking the negative logarithm turns those into 3.17 and 9.21, and the rule inverts:
- Small pKa = strong acid.
- Each 1 unit of pKa is a factor of 10 in Ka.
So an acid with pKa 3 is a thousand times stronger than one with pKa 6. The sign flip catches people out constantly — it is the same flip you already know from pH, where a lower number means more acidic.
Worked example: the pH of 0.10 M acetic acid
Ka = 1.8 × 10⁻⁵. Let x be the concentration of H₃O⁺ formed.
| CH₃COOH | H₃O⁺ | CH₃COO⁻ | |
|---|---|---|---|
| Start | 0.10 | ~0 | 0 |
| Change | −x | +x | +x |
| Equilibrium | 0.10 − x | x | x |
Substitute:
1.8 × 10⁻⁵ = x² ÷ (0.10 − x)
Because the acid is weak, x is small compared with 0.10, so approximate 0.10 − x ≈ 0.10:
x² = 1.8 × 10⁻⁵ × 0.10 = 1.8 × 10⁻⁶ x = 1.34 × 10⁻³ M
pH = −log(1.34 × 10⁻³) = 2.87
Check the approximation: x ÷ 0.10 = 1.3%, comfortably under the 5% rule, so ignoring x in the denominator was fair. That 1.3% is also the percent dissociation — in 0.10 M acetic acid, about 1 molecule in 75 has actually given up its proton.
For comparison, 0.10 M HCl would give pH 1.00. Same concentration, nearly two pH units apart, purely because of Ka.
What Ka is not
Ka describes the acid itself, at a given temperature — it does not change when you dilute the solution. Concentration changes the pH; Ka does not budge. Dilute acetic acid to 0.010 M and the pH rises to about 3.37, but Ka is still 1.8 × 10⁻⁵. That is exactly why chemists prefer it: it is a property of the substance, not of the bottle it came in.
One extra use worth knowing: Ka and the Kb of the acid's conjugate base are locked together by Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25 °C, so a table of Ka values is secretly a table of Kb values too.
Common mistakes to avoid
- Mixing up strength and concentration. A concentrated weak acid and a dilute strong acid can have the same pH. Ka measures the tendency to ionise; molarity measures how much acid is present.
- Getting the pKa direction backwards. Higher pKa means weaker. Write "small pKa, strong acid" in the margin of your data booklet.
- Using the approximation blindly. If the acid is fairly strong or very dilute, x is not negligible and you must solve the quadratic. Always test x ÷ [HA]₀ against 5%.
FAQ
Does Ka change with temperature?
Yes. Ka is an equilibrium constant, so it shifts when temperature does. Values in data books are quoted at 25 °C, which is why exam questions specify it.
What is the difference between Ka and pH?
Ka is a fixed property of the acid. pH is a property of one particular solution and depends on both Ka and concentration.
What counts as a strong acid?
One that ionises essentially completely in water — HCl, HBr, HI, HNO₃, H₂SO₄ (first proton) and HClO₄. Everything else you meet at this level is weak.
Why does water not appear in the Ka expression?
It is the solvent and stays effectively constant, so its concentration is absorbed into the value of the constant itself.
The takeaway
Ka is the equilibrium constant for an acid handing its proton to water, and it puts a number on the vague word "weak". Big Ka, small pKa, strong acid. Combine Ka with the starting concentration and an equilibrium table and you can calculate the pH of a weak acid solution from first principles — which is the whole point of having the constant.
Come from [Conjugate Acid vs Conjugate Base] and pair this with Strong vs Weak Acids: What's the Difference?, What Is Chemical Equilibrium? The Balance Point and What Is pH? The pH Scale Explained Simply. Next: [Strong vs Weak Bases].
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