Case 1
- pKa of the weak acid
- 4.76
- Conjugate base [A⁻]
- 0.1mol/L
- Weak acid [HA]
- 0.1mol/L
4.76
Open with these valuesHenderson-Hasselbalch Calculator
Change any value to update the result.
4.76pH
Buffer pH: 4.76 pHExamples
How the result moves
Weak acid [HA] (mol/L) → Buffer pH (pH)Your current valid inputs, with one value varied across the range.
How the result moves
Held fixed: pKa of the weak acid 4.76, Conjugate base [A⁻] 0.1000 mol/L.
| Weak acid [HA] (mol/L) | Buffer pH |
|---|---|
| 0.0500 | 5.06 |
| 0.1000Your value | 4.76 |
| 0.1500 | 4.58 |
The buffer pH is the pKa plus the base-10 log of the base-to-acid ratio: pH = pKa + log₁₀([A⁻] ÷ [HA]). Equal concentrations make the log term zero, so the pH is exactly the pKa — that is where a buffer works best. Only the ratio matters, not the amount.
Worked examples
Case 2
- pKa of the weak acid
- 4.76
- Conjugate base [A⁻]
- 0.2mol/L
- Weak acid [HA]
- 0.1mol/L
5.06
Open with these valuesCase 3
- pKa of the weak acid
- 7.21
- Conjugate base [A⁻]
- 0.3mol/L
- Weak acid [HA]
- 0.1mol/L
7.69
Open with these valuesHow it's calculated
pH = pKa + log₁₀([A⁻] ÷ [HA])
- StepEnter the pKa of the weak acid in your buffer.pKa of the weak acid=4.76
- StepEnter base and acid in the same concentration unit.Conjugate base [A⁻]=0.1mol/L
- ResultRead the pH — equal concentrations put it exactly on the pKa.4.76pH
Reference table
| pKa, [A⁻], [HA] | Base to acid | Buffer pH |
|---|---|---|
| 4.76, 0.1, 0.2 | 1 to 2 | 4.46 |
| 4.76, 0.1, 0.1 | 1 to 1 | 4.76 |
| 4.76, 0.2, 0.1 | 2 to 1 | 5.06 |
| 9.25, 0.05, 0.15 | 1 to 3 | 8.77 |
| 7.21, 0.3, 0.1 | 3 to 1 | 7.69 |
Questions
How do I use the Henderson-Hasselbalch equation?
Add the pKa of the weak acid to the base-10 logarithm of the conjugate base divided by the acid: pH = pKa + log₁₀([A⁻] ÷ [HA]). Use the same concentration unit for both species. With pKa 4.76 and equal concentrations the log term is zero, so the pH is 4.76.
What is a buffer solution?
A buffer is a mixture of a weak acid and its conjugate base that resists pH change when small amounts of acid or base are added. The equation predicts its pH from the ratio of the two species.
Why does the pH equal the pKa at equal concentrations?
Because the ratio is then 1 and log₁₀(1) = 0, which leaves pH = pKa. That is also where the buffer is strongest, since it can absorb acid and base equally well.
What is pKa and where do I get it?
pKa is the negative base-10 logarithm of the acid dissociation constant, and a lower value means a stronger acid. It is a measured property of the particular acid and varies slightly with temperature and ionic strength, so take it from a table for your buffer.
When is the equation inaccurate?
It assumes the weak acid is only partly dissociated and the concentrations are moderate. It loses accuracy for very dilute or very concentrated solutions, for strong acids and bases, and when the ratio is far from 1.