A buffer is a mixture of a weak acid and its salt (for example acetic acid and sodium acetate) or a weak base and its salt (ammonia and ammonium chloride). The weak component neutralises added base, the salt neutralises added acid, and together they hold the pH near a value set by the dissociation constant and the ratio of the two components.
For acid buffers the pH equals the pKa plus the logarithm of the salt-to-acid ratio. For base buffers you calculate pOH from the pKb in the same way and subtract it from pKw (14.00 at 25 °C). Because only the ratio matters, you can work with concentrations in the final buffer, with moles, or with the volumes of two stock solutions you plan to mix – the total volume cancels out of the ratio, although it still sets the final concentrations and therefore the buffer capacity.
The equation assumes both components are present at concentrations much larger than [H⁺] or [OH⁻] and that activity effects are small, so it is most accurate for buffers between about 0.01 M and 0.5 M with a ratio between 1:10 and 10:1. Outside that range, results drift from measured values by a few hundredths of a pH unit or more.
How to use the Buffer pH Calculator
- Buffer type
- Buffer system
- Buffer system
- I know the
- I know the
- pKa (or pKb for a base buffer)
- Ka (or Kb for a base buffer)
- pKb
- Kb
- Amounts given as
- Weak acid / weak base concentration
- Volume of weak acid / base stock
- Salt (conjugate) concentration
- Volume of salt stock
- Moles of weak acid / base
- Moles of salt
- pKw (14.00 at 25 °C)
Results update as you type, every field has its own unit menu, and everything is calculated in your browser.
How it works
Choose an acid or base buffer and a system (or enter the constant). The calculator turns your amounts into the salt-to-weak-component ratio, applies the Henderson–Hasselbalch equation for pH or pOH, and reports [H⁺], the final concentrations and whether the buffer lies inside its useful range.
Worked example
You mix 40 mL of 0.10 M acetic acid with 60 mL of 0.15 M sodium acetate. Moles of acid = 0.10 × 0.040 = 0.0040 mol; moles of acetate = 0.15 × 0.060 = 0.0090 mol. The ratio is 0.0090/0.0040 = 2.25, so pH = 4.76 + log₁₀ 2.25 = 4.76 + 0.352 = 5.11. In the 100 mL mixture the concentrations are 0.040 M acid and 0.090 M acetate.
For a base buffer of 0.20 M NH₃ and 0.10 M NH₄Cl: pOH = 4.75 + log₁₀(0.10/0.20) = 4.449, and pH = 14.00 − 4.449 = 9.55.
Common buffer systems at 25 °C
| Buffer pair | pKa | Useful pH range |
|---|---|---|
| Formic acid / formate | 3.75 | 2.8–4.8 |
| Lactic acid / lactate | 3.86 | 2.9–4.9 |
| Acetic acid / acetate | 4.76 | 3.8–5.8 |
| Carbonic acid / bicarbonate | 6.35 | 5.4–7.4 |
| H₂PO₄⁻ / HPO₄²⁻ | 7.20 | 6.2–8.2 |
| Tris-H⁺ / Tris | 8.07 | 7.1–9.1 |
| Boric acid / borate | 9.24 | 8.2–10.2 |
| NH₄⁺ / NH₃ | 9.25 | 8.3–10.3 |
| HCO₃⁻ / CO₃²⁻ | 10.33 | 9.3–11.3 |
Tips
- Use the final volume of the solution, not the volume of solvent added, for molarity.
- For dilutions, C₁V₁ = C₂V₂ only holds when both concentrations use the same unit.
- Add acid to water, never water to acid.
- Everything is calculated in your browser.
Common mistakes to avoid
- Mixing w/w, w/v and v/v percentages.
- Forgetting that pH is logarithmic – one unit is a tenfold change in [H⁺].
- Using pKa where Ka is asked for (or the other way round).
Inputs explained
| Field | What to enter |
|---|---|
| Buffer type | Choose one: Weak acid + its salt (e.g. CH₃COOH / CH₃COONa), Weak base + its salt (e.g. NH₃ / NH₄Cl) |
| Buffer system | Choose one: Acetic acid / acetate (pKa 4.76), Formic acid / formate (pKa 3.75), Lactic acid / lactate (pKa 3.86), Benzoic acid / benzoate (pKa 4.20) … |
| Buffer system | Choose one: Ammonia / ammonium (pKb 4.75), Methylamine / methylammonium (pKb 3.36), Tris / Tris-H⁺ (pKb 5.93), Pyridine / pyridinium (pKb 8.77) … |
| I know the | Choose one: pK value, Dissociation constant K |
| I know the | Choose one: pK value, Dissociation constant K |
| pKa (or pKb for a base buffer) | A number |
| Ka (or Kb for a base buffer) | A number |
| pKb | A number |
| Kb | A number |
| Amounts given as | Choose one: Concentrations in the buffer, Moles of each component, Volumes of two stock solutions |
| Weak acid / weak base concentration | Units: mol/L (M), mmol/L (mM), µmol/L (µM), nmol/L (nM), mol/m³ |
| Volume of weak acid / base stock | Units: L, mL, µL, m³, dm³, cm³ … |
| Salt (conjugate) concentration | Units: mol/L (M), mmol/L (mM), µmol/L (µM), nmol/L (nM), mol/m³ |
| Volume of salt stock | Units: L, mL, µL, m³, dm³, cm³ … |
| Moles of weak acid / base | Units: mol, mmol, µmol, nmol, kmol |
| Moles of salt | Units: mol, mmol, µmol, nmol, kmol |
| pKw (14.00 at 25 °C) | A number |
What each result means
- Salt : weak component ratio = [salt]÷[weak]
- pOH = [pKw (14.00 at 25 °C)] – [Buffer pH]
- Hydrogen ion concentration [H⁺] = 10^(-[Buffer pH])
- Weak component concentration after mixing = [weak]÷([Volume of weak acid ÷ base stock] + [Volume of salt stock])
- Salt concentration after mixing = [salt]÷([Volume of weak acid ÷ base stock] + [Volume of salt stock])
- Total volume of the mixture = [Volume of weak acid ÷ base stock] + [Volume of salt stock]
Formula
Acid buffer: pH = pKa + log₁₀([salt] ÷ [acid]) · Base buffer: pOH = pKb + log₁₀([salt] ÷ [base]), pH = pKw − pOH · pKa = −log₁₀ Ka · mixed stocks: [X] = C·V ÷ (V₁ + V₂)Worked example
Inputs — Buffer type: Weak acid + its salt (e.g. CH₃COOH / CH₃COONa); Buffer system: Acetic acid / acetate (pKa 4.76); Amounts given as: Concentrations in the buffer; Weak acid / weak base concentration: 0.1 mol/L (M); Salt (conjugate) concentration: 0.15 mol/L (M); pKw (14.00 at 25 °C): 14.
Results — pK used = 4.76; Salt : weak component ratio = 1.5; Buffer pH = 4.936; pOH = 9.064; Hydrogen ion concentration [H⁺] = 1.1585e-5 mol/L (M); Range check = Good buffer – ratio between 1:10 and 10:1.
Key terms
- Buffer
- A solution of a weak acid or base with its conjugate that resists changes in pH.
- pKa
- −log₁₀ of the acid dissociation constant; the pH at which the acid is half dissociated.
- pKb
- −log₁₀ of the base dissociation constant; pKa + pKb = pKw for a conjugate pair.
- Conjugate pair
- Two species that differ by one proton, such as NH₄⁺ / NH₃.
Frequently asked questions
How do you calculate the pH of a buffer?
For an acid buffer pH = pKa + log₁₀([salt]/[acid]). For 0.10 M acetic acid with 0.15 M sodium acetate: 4.76 + log₁₀(1.5) = 4.94.
How do you calculate the pH of a basic buffer?
Find pOH = pKb + log₁₀([salt]/[base]) and then pH = 14 − pOH. For 0.20 M NH₃ and 0.10 M NH₄Cl: pOH = 4.75 + log₁₀(0.5) = 4.45, so pH = 9.55.
What if I mix two stock solutions?
Multiply each stock concentration by its volume to get moles. 40 mL of 0.1 M acetic acid plus 60 mL of 0.15 M acetate gives a ratio of 9 : 4 and a pH of 4.76 + log₁₀(2.25) = 5.11.
What happens when the acid and salt concentrations are equal?
The log term is zero, so pH = pKa (or pOH = pKb). This is also where buffer capacity is highest.
Does diluting a buffer change its pH?
Hardly – the ratio stays the same, so the calculated pH does not change. In practice it shifts slightly because activity coefficients change, and the capacity falls in proportion to the dilution.
How do I choose a buffer for a target pH?
Pick a weak acid whose pKa is within 1 unit of the target: acetate for pH 3.8–5.8, phosphate for 6.2–8.2, Tris for 7.1–9.1, ammonia or borate for 8.2–10.2.
Does temperature affect buffer pH?
Yes. pKa values are quoted at 25 °C; Tris drops by about 0.03 pH per °C, while acetate and phosphate change much less. pKw also falls with temperature (13.26 at 50 °C).