Buffer Calculator
Calculate the pH of a buffer solution using the Henderson-Hasselbalch equation, with an ionic strength correction for more accurate real-world results. Perfect for biochemistry and analytical chemistry.
About this calculator
Buffer pH is governed by the Henderson-Hasselbalch equation: pH = pKa + log₁₀([A⁻] / [HA]), where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration. This calculator extends that equation with a Debye-Hückel ionic strength correction: pH = pKa + log₁₀(base_conc / acid_conc) − (0.5 × √I / (1 + √I)), where I is ionic strength in mol/L. The correction accounts for charge-charge interactions in non-ideal solutions that shift effective activity away from concentration. Without it, calculated pH can deviate by 0.1–0.3 units at physiological ionic strengths (~0.15 M). Buffers work best when the ratio [A⁻]/[HA] is between 0.1 and 10, meaning within one pH unit of the pKa.
How to use
Prepare an acetate buffer using acetic acid (pKa = 4.76), 0.1 M acetic acid, and 0.1 M sodium acetate, with ionic strength I = 0.1 M. Step 1: log₁₀(0.1 / 0.1) = log₁₀(1) = 0. Step 2: Ionic correction = 0.5 × √0.1 / (1 + √0.1) = 0.5 × 0.3162 / 1.3162 = 0.1201. Step 3: pH = 4.76 + 0 − 0.1201 = 4.64. Without the ionic correction the pH would equal the pKa of 4.76, but real ionic interactions lower it to 4.64.
Frequently asked questions
How does ionic strength affect the pH of a buffer solution?
Ionic strength reflects the total concentration of ions in solution and influences the activity coefficients of charged species. Higher ionic strength lowers the effective activity of ions relative to their molar concentration, shifting the apparent pH. The Debye-Hückel correction used here quantifies this effect: at I = 0.1 M the correction is about 0.12 pH units. For biological buffers like PBS, ignoring ionic strength can lead to meaningful errors when precision matters in enzyme assays or cell culture work.
What is the Henderson-Hasselbalch equation and when should I use it?
The Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA]), relates buffer pH to the ratio of conjugate base to weak acid concentrations. It is valid when the weak acid and its conjugate base are both present in appreciable amounts and the total solute concentrations are not extremely dilute. It is widely used to design biological buffers (phosphate, acetate, HEPES, TRIS) and to predict pH shifts on adding acid or base. The equation becomes less accurate at very high or very low concentration ratios or in highly ionic solutions without a correction factor.
Why does the buffer capacity decrease far from the pKa value?
Buffer capacity is the ability of a solution to resist pH change, and it peaks when [A⁻] = [HA], i.e., when pH = pKa. At this point, equal amounts of acid and base are available to neutralize added base or acid, respectively. Moving more than one pH unit away from the pKa means one component is vastly depleted, so small additions of acid or base cause large pH swings. Practical buffering range is generally accepted as pKa ± 1 pH unit, where the ratio stays between 0.1 and 10.