Ka and Kb are essential equilibrium constants in IB Chemistry Topic 8 (Acids and Bases) and HL equilibrium calculations. They quantify the strength of weak acids and weak bases, determine pH, and explain why some species ionize more than others. Understanding Ka and Kb helps you analyze acid–base behavior, buffer systems, and titration curves with clarity.
What Is Ka?
Ka (acid dissociation constant) measures the extent to which a weak acid ionizes in water.**
A generic weak acid, HA, dissociates like this:
HA ⇌ H⁺ + A⁻
The Ka expression is:
Ka = [H⁺][A⁻] / [HA]
Key points:
- Ka applies only to weak acids
- Strong acids do not have Ka values (because they dissociate fully)
- Larger Ka → stronger acid
- Smaller Ka → weaker acid
Ka shows how much the acid “lets go” of its proton.
What Is Kb?
Kb (base dissociation constant) measures the extent to which a weak base accepts a proton from water.**
A generic weak base, B, reacts like this:
B + H₂O ⇌ BH⁺ + OH⁻
The Kb expression is:
Kb = [BH⁺][OH⁻] / [B]
Key points:
- Kb applies to weak bases only
- Larger Kb → stronger base
- Smaller Kb → weaker base
Kb shows how effectively the base attracts a proton.
Ka and Kb Reflect Strength, Not Concentration
This is one of the biggest IB exam misconceptions:
- Strength = how much a substance ionizes
- Concentration = how much substance is present
A dilute strong acid still has a very high H⁺ concentration relative to a concentrated weak acid. Ka and Kb always refer to degree of ionization, not the amount of acid or base added.
Relationship Between Ka and Kb
For any conjugate acid–base pair:
Ka × Kb = Kw
Where Kw = 1.00 × 10⁻¹⁴ at 25°C.
This means:
- A strong acid (large Ka) has a conjugate base with a very small Kb
- A strong base (large Kb) has a conjugate acid with a very small Ka
This reciprocal relationship is frequently tested in Paper 2.
Using Ka and Kb to Determine pH
Ka and Kb allow you to calculate:
- [H⁺] for weak acids
- [OH⁻] for weak bases
- pH or pOH
- Degree of dissociation
- Concentration of ions in buffers
For a weak acid HA, if x is small:
[H⁺] ≈ √(Ka × [HA])
For a weak base B:
[OH⁻] ≈ √(Kb × [B])
These approximations work because weak acids and weak bases ionize only slightly.
Ka and Kb in Buffer Solutions
Buffers rely on conjugate acid–base pairs.
Their pH depends on the ratio of acid to conjugate base.
The Henderson–Hasselbalch equation (HL) uses Ka:
pH = pKa + log([A⁻]/[HA])
Similarly, for basic buffers:
pOH = pKb + log([BH⁺]/[B])
This shows how Ka and Kb control buffer behavior.
pKa and pKb
To simplify calculations, Ka and Kb are often converted to logarithmic form:
pKa = −log(Ka)
pKb = −log(Kb)
Interpretation:
- Smaller pKa → stronger acid
- Smaller pKb → stronger base
These values are easier to compare than Ka and Kb directly.
Trends in Ka and Kb
Stronger acids:
- Higher Ka
- Lower pKa
Stronger bases:
- Higher Kb
- Lower pKb
These trends help identify whether species behave strongly or weakly in water.
Common IB Mistakes
- Mixing up strength with concentration
- Forgetting that Ka and Kb apply only to weak acids and bases
- Using the wrong equilibrium expression
- Neglecting units (Ka and Kb are typically dimensionless in IB)
- Forgetting the relationship Ka × Kb = Kw
Careful setup of expressions prevents calculation errors.
FAQs
Do strong acids have Ka values?
Not in practice. Their ionization is essentially 100%, so Ka would be extremely large and not useful.
Does temperature affect Ka and Kb?
Yes. Both are equilibrium constants and vary with temperature.
Why do weak acids need Ka instead of simple stoichiometry?
Because they only partially dissociate, so equilibrium—not full dissociation—must be considered.
Conclusion
Ka and Kb are equilibrium constants that measure the strength of weak acids and weak bases. Ka shows how much an acid donates protons, while Kb shows how effectively a base accepts protons. These constants determine pH, relate conjugate pairs through Kw, and form the basis for buffer calculations. Mastering Ka and Kb is essential for tackling acid–base questions in IB Chemistry.
