Buffer solution
Buffer solutions resist pH change via weak acid–conjugate base equilibrium.
A buffer solution maintains a relatively stable pH, even when diluted or when small amounts of strong acid or base are added, provided the temperature stays constant. Its ability to resist pH change is measured by buffer capacity. These solutions are widely used in chemistry to keep pH nearly constant. In nature, many living organisms rely on buffering for pH regulation; for instance, the bicarbonate buffering system helps regulate blood pH, and bicarbonate also serves as a buffer in the ocean.
The resistance to pH change comes from a chemical equilibrium between a weak acid (HA) and its conjugate base (A⁻). Adding a strong acid introduces extra hydrogen ions (H⁺), which shifts the equilibrium to the left (Le Chatelier's principle), so the hydrogen ion concentration rises less than expected. Similarly, adding a strong base reduces hydrogen ion concentration less than expected, because most of the added hydroxide is consumed by the weak acid. This effect is most pronounced in the buffer region, where pH is within one unit of the acid's pKa. Once the acid is over 95% deprotonated, the pH rises quickly because most added alkali then neutralizes the remaining acid.
Buffer capacity (β) quantifies this resistance. It is defined as the infinitesimal change in added base concentration (dCb) divided by the infinitesimal change in pH (d(pH)), or equivalently as the negative of the infinitesimal change in added acid concentration (dCa) divided by d(pH). For a weak acid with dissociation constant Ka, buffer capacity can be expressed as β = 2.303([H⁺] + (THA·Ka·[H⁺])/(Ka + [H⁺])² + Kw/[H⁺]), where [H⁺] is the hydrogen ion concentration, THA is the total concentration of added acid, and Kw is the water self-ionization constant.
- definition
- Solution resisting pH change on dilution or addition of small amounts of strong acid or base
- key principle
- Chemical equilibrium between weak acid HA and its conjugate base A−
- buffer capacity formula
- β = dCb/d(pH) or β = −dCa/d(pH)
- useful pH range
- pKa ± 1
- peak buffer capacity
- At pH = pKa
- self-ionization constant of water
- Kw = 1.0×10−14
Lore & Background
Buffer solutions resist pH change because of a chemical equilibrium between the weak acid HA and its conjugate base A−. When strong acid is added, hydrogen ions shift the equilibrium to the left, per Le Chatelier's principle, so the hydrogen ion concentration increases less than expected. Similarly, adding strong alkali decreases hydrogen ion concentration less than expected, as most added hydroxide is consumed in a reaction with the weak acid. The effect is illustrated by simulated titration of a weak acid with pKa = 4.7, where pH changes slowly in the buffer region pH = pKa ± 1, centered at pH = 4.7, where [HA] = [A−]. Once the acid is more than 95% deprotonated, pH rises rapidly.
Reader's Guide
Buffer capacity is a quantitative measure of resistance to pH change, defined as β = dCb/d(pH) or β = −dCa/d(pH), where dCb and dCa are infinitesimal amounts of added base or acid. This equation shows three regions of raised buffer capacity. In the central region (pH near pKa), the second term dominates, and buffer capacity peaks at pH = pKa, falling to 33% at pH = pKa ± 1, 10% at pH = pKa ± 1.5, and 1% at pH = pKa ± 2. The most useful range is approximately pKa ± 1. For strongly acidic solutions (pH < 2), the first term dominates and buffer capacity rises exponentially with decreasing pH. For strongly alkaline solutions (pH > 12), the third term dominates and buffer capacity rises exponentially with increasing pH. Buffer capacity is negligible when the concentration of buffering agent is very small and increases with its concentration.
Did You Know?
- Buffer capacity is defined as β = dCb/d(pH) or β = −dCa/d(pH).
- Buffer capacity peaks at pH = pKa and falls to 33% at pH = pKa ± 1.
- The bicarbonate buffering system regulates pH of blood and acts as a buffer in the ocean.
- Buffer capacity rises exponentially with decreasing pH below about 2, independent of buffering agent.
Frequently Asked Questions
What is a buffer solution?
A buffer solution is a mixture that barely bounces its pH when you dilute it or drop in a small dose of strong acid or base, provided the temperature stays fixed. Think of it as the pH shock absorber of the chemistry world.
How does a buffer solution actually resist pH change?
It leans on the reversible equilibrium between a weak acid (HA) and its conjugate base (A−), so any extra H⁺ or OH⁻ you introduce gets mopped up by one side of the pair instead of freely swinging the pH.
What pH range does a buffer solution work best in?
You get the most protection within about one pH unit above or below the pKa of the weak acid you chose, and the buffer hits its absolute peak capacity right at pH = pKa.
Where do buffer solutions show up in nature?
The bicarbonate buffering system in human blood and in seawater are the go-to natural examples, holding pH steady enough for enzymes, cells, and marine ecosystems to keep functioning.
How is buffer capacity defined and measured?
Buffer capacity (β) quantifies how much strong acid or base you must add to shift the pH by one unit, and it's expressed mathematically as β = dCb/d(pH) or, equivalently, β = −dCa/d(pH).
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