What is pKa?
The pKa is a negative logarithmic scale describing how strongly an acid retains its proton. A lower pKa means the acid releases hydrogen ions more readily—it is a stronger acid. Conversely, a higher pKa indicates a weaker acid that holds its proton more tightly.
- pKa < 0: Very strong acids that dissociate completely in aqueous solution.
- pKa 0–7: Weak to moderately strong acids; partial dissociation in water.
- pKa > 7: Extremely weak acids; negligible ionisation at neutral pH.
This measure is invaluable when selecting buffer components, predicting protonation states of functional groups in biomolecules, or gauging reactivity in organic synthesis. By knowing the pKa of an acid and the pH of the solution, you can estimate what fraction exists in protonated versus deprotonated form.
Calculating pKa from pH and Concentration
The Henderson-Hasselbalch equation relates pH, pKa, and the ratio of conjugate base to weak acid concentrations. Rearranging this relationship lets you solve for pKa when you have measured pH and know the molar amounts of the acid and its conjugate base.
pKa = pH − log₁₀([A⁻] ÷ [HA])
pKa = −log₁₀(Ka)
pH— The measured acidity or alkalinity of the solution.[A⁻]— Molar concentration of the conjugate base (deprotonated form).[HA]— Molar concentration of the weak acid (protonated form).Ka— The acid dissociation constant, measured at equilibrium.
Understanding Ka and Its Relationship to pKa
The acid dissociation constant Ka quantifies how completely an acid dissociates at equilibrium:
HA ⇌ H⁺ + A⁻
A larger Ka indicates stronger dissociation and a stronger acid. Since pKa is simply the negative logarithm of Ka, they are inversely related: as Ka increases, pKa decreases. This logarithmic conversion compresses the wide range of Ka values into a more manageable scale, making comparison between acids straightforward.
For example, acetic acid has a Ka of 1.8 × 10⁻⁵, yielding a pKa of approximately 4.74. Formic acid, with Ka = 1.8 × 10⁻⁴, has pKa ≈ 3.74—half a unit lower, reflecting its greater acidity.
Using pKa Tables for Quick Reference
Standard pKa tables list values for common acids, bases, and functional groups under specified conditions (usually 25 °C in aqueous solution). These tables are invaluable when designing reactions or selecting buffer systems from first principles, eliminating the need to calculate pKa every time.
Typical ranges include:
- Strong organic acids (carboxylic acids in water): pKa 3–5
- Phenols: pKa 8–11
- Thiols: pKa 8–10
- Ammonium ions: pKa 8–11
When constructing a buffer, choose an acid whose pKa lies within 1 unit of your target pH for maximum buffering capacity.
Common Pitfalls and Practical Considerations
Avoid these frequent errors when working with pKa values and calculations.
- Temperature and Solvent Effects — pKa values shift with temperature and solvent. Standard tables assume aqueous solution at 25 °C. Using a pKa value from an organic solvent reference in water, or vice versa, will introduce significant error. Always note the experimental conditions.
- Concentration Units and Equilibrium — The Henderson-Hasselbalch equation requires molar concentrations at equilibrium. If you have only initial concentrations, you must first calculate equilibrium concentrations using an ICE table or similar approach. Neglecting dissociation of the weak acid or hydrolysis of the conjugate base can skew results.
- pH vs. pKa Near Equivalence Point — When [A⁻] = [HA], the log term equals zero and pH = pKa. This is the point of maximum buffering capacity. Away from this ratio, small changes in concentration cause larger pH shifts; the buffer becomes less effective.
- Ionic Strength and Activity Coefficients — For precise work in solutions of high ionic strength, use activities rather than concentrations. The apparent pKa can shift due to ion pairing and activity coefficient effects, especially important in cellular or high-salt environments.