Buffer pH Calculator
Find the pH of an acidic or basic buffer with the Henderson-Hasselbalch equation. Enter pKa or pKb and the two concentrations to get pH, pOH, and the ratio.
🧪 What is a Buffer pH Calculation?
A buffer pH calculation works out the pH of a buffer solution using the Henderson-Hasselbalch equation. A buffer is a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resists changes in pH when small amounts of acid or base are added. The equation pH = pKa + log([A⁻] / [HA]) links the pH directly to the acid dissociation constant and the ratio of the two components, so you can predict the pH of any buffer from its recipe.
Buffers are essential across chemistry and biology. Blood is held near pH 7.4 by a bicarbonate buffer so that enzymes work correctly. Laboratory buffers keep reactions, electrophoresis gels, and cell cultures at a controlled pH. Shampoos, foods, and medicines are buffered to stay stable and safe. In every case, chemists use the Henderson-Hasselbalch equation to choose the right acid and the right ratio to hit a target pH.
A common misconception is that adding more of the buffer components changes the pH. Because the equation depends on the ratio of base to acid, not the absolute amounts, doubling both leaves the pH unchanged while increasing the buffer capacity. Another misconception is that a buffer holds any pH equally well. In reality a buffer is only effective within about one unit of its pKa, so the acid must be chosen to match the desired pH.
This calculator applies the Henderson-Hasselbalch equation in two forms. The acidic mode uses pKa with a weak acid and its conjugate base, and the basic mode uses pKb with a weak base and its conjugate acid, converting through pOH. It reports the pH, the pOH, the base-to-acid ratio, and whether the solution is acidic, neutral, or basic, with the full working shown.
📐 Formula
📖 How to Use This Calculator
Steps
💡 Example Calculations
Example 1 — Acetate Buffer at Half-Equivalence
Equal 0.1 M acetic acid and 0.1 M acetate, pKa 4.76
Example 2 — Acetate Buffer with More Base
0.2 M acetate and 0.1 M acetic acid, pKa 4.76
Example 3 — Ammonia Basic Buffer
0.1 M ammonium (NH₄⁺) and 0.2 M ammonia (NH₃), pKb 4.75
❓ Frequently Asked Questions
🔗 Related Calculators
What is the Henderson-Hasselbalch equation?
The Henderson-Hasselbalch equation relates the pH of a buffer to the pKa of the weak acid and the ratio of conjugate base to acid: pH = pKa + log([A⁻] / [HA]). It lets you calculate buffer pH directly from concentrations, and shows that when the two are equal the pH equals the pKa.
How do you calculate the pH of a buffer?
For an acidic buffer, use pH = pKa + log([A⁻] / [HA]), where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration. For example, an acetate buffer with pKa 4.76 and equal 0.1 M concentrations gives pH = 4.76 + log(1) = 4.76. For a basic buffer, find pOH from pKb first, then pH = 14 − pOH.
What is a buffer solution?
A buffer solution resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable amounts. When acid is added the base neutralises it, and when base is added the acid neutralises it, keeping pH nearly constant.
When does buffer pH equal pKa?
Buffer pH equals pKa when the concentrations of the weak acid and its conjugate base are equal. The ratio [A⁻] / [HA] is then 1, and log(1) = 0, so pH = pKa + 0 = pKa. This point, called the half-equivalence point, is where the buffer has its maximum capacity to resist pH change.
What is the useful pH range of a buffer?
A buffer works effectively within about one pH unit of its pKa, that is from pKa − 1 to pKa + 1. Within this range the concentration ratio stays between roughly 1:10 and 10:1, so both components are present in useful amounts. Outside it, one species is nearly depleted and buffering becomes weak.
How do you calculate the pH of a basic buffer?
For a basic buffer of a weak base and its conjugate acid, first find pOH = pKb + log([BH⁺] / [B]), where [BH⁺] is the conjugate acid and [B] is the weak base. Then convert with pH = 14 − pOH at 25°C. An ammonia buffer with pKb 4.75, 0.1 M NH₄⁺ and 0.2 M NH₃ gives pOH 4.45 and pH 9.55.
What is the difference between pKa and pKb?
pKa measures the strength of a weak acid (a lower pKa means a stronger acid), while pKb measures the strength of a weak base. For a conjugate acid-base pair at 25°C they are linked by pKa + pKb = 14, so knowing one gives the other. Use pKa for acidic buffers and pKb for basic buffers.
Does diluting a buffer change its pH?
Very little. The Henderson-Hasselbalch equation depends on the ratio of concentrations, not their absolute values. Adding water dilutes both the acid and the base by the same factor, leaving the ratio and therefore the pH almost unchanged. Extreme dilution eventually shifts pH toward neutral, but moderate dilution has minimal effect.
What are common buffer systems and their pKa values?
Common buffers include acetate (pKa 4.76), phosphate (pKa 7.21 for the second dissociation), carbonate (pKa 6.35 and 10.33), and Tris (pKa 8.07). Blood is buffered mainly by the bicarbonate system near pH 7.4. Choose a buffer whose pKa is close to the pH you need to maintain.
Why is the Henderson-Hasselbalch equation an approximation?
It assumes the equilibrium concentrations of acid and base equal the amounts you added, ignoring the small shift from dissociation, and it assumes ideal behaviour with activity equal to concentration. For dilute buffers within the pKa ± 1 range these assumptions hold well, so the equation is accurate enough for most laboratory and teaching purposes.