AP Chemistry

Properties of Buffers

Learning Objective 8.8.A Explain the relationship between the ability of a buffer to stabilize pH and the reactions that occur when an acid or a base is added to a buffered solution.

Quick Notes

Buffers resist changes in pH when small amounts of acid or base are added.
A buffer contains a conjugate acid-base pair: a weak acid (HA) and its conjugate base (A⁻), or a weak base (B) and its conjugate acid (BH⁺).
The acid neutralizes added base, and the base neutralizes added acid.

Full Notes

What is a Buffer?

A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers work by having both a weak acid and its conjugate base (or a weak base and its conjugate acid) in equilibrium.

Types of Buffer Systems

1. Acidic Buffers: Made from a weak acid and its salt (that contains the acid’s conjugate base).

For Example: The weak acid ethanoic acid (CH₃COOH) and its salt sodium ethanoate (CH₃COONa).

AP Chemistry diagram showing preparation of an acidic buffer using ethanoic acid and sodium ethanoate.

You can also prepare a buffer by reacting a weak acid with a limited amount of strong base, such as:

CH₃COOH + NaOH → CH₃COONa + H₂O

This produces both HA and A⁻ in the same solution.

2. Basic Buffers: Made from a weak base and its salt (that contains the conjugate acid of the base).

For example The weak base ammonia (NH₃) and its salt ammonium chloride (NH₄Cl).

AP Chemistry diagram showing preparation of a basic buffer using ammonia and ammonium chloride.

When added to a solution of ammonia, the NH₄Cl would dissociate and release NH₄⁺ ions, which are the conjugate acid ions of the ammonia.

How Do Acidic Buffers Work?

An equilibrium is established in the buffer system between HA, A⁻ and H⁺.

AP Chemistry diagram showing equilibrium between HA, A− and H+ in a buffer system.

The concentration of HA and A⁻ in the mixture must be much greater than the concentration of H⁺. This ensures the equilibrium is sensitive to H⁺ changes more than HA or A⁻. Equilibrium can shift to keep [H⁺] nearly constant.

Example: Ethanoic Acid/Sodium Ethanoate Buffer (CH₃COOH/CH₃COO⁻)

Equilibrium reaction:
CH₃COOH ⇌ H⁺ + CH₃COO⁻

AP Chemistry diagram of ethanoic acid/sodium ethanoate buffer equilibrium.

When an acid (H⁺) is added:

CH₃COO⁻ combines with added H⁺ to form CH₃COOH.
Equilibrium shifts left, reducing the increase in H⁺.
[HA] increases and [A⁻] decreases.

When a base (OH⁻) is added:

AP Chemistry diagram of buffer action when hydroxide ions are added.

Added OH⁻ reacts with H⁺ to form H₂O.
CH₃COOH dissociates more to replace lost H⁺.
Equilibrium shifts right, resisting pH increase.
[HA] decreases and [A⁻] increases.

As long as there are significant amounts of HA and A⁻ present, this system can buffer against pH changes.

Matt’s exam tip

Remember that the concentration of HA and A⁻ will change when H⁺ or OH⁻ are added. When H⁺ ions are added then moles of HA increase and moles of A⁻ decrease. When OH⁻ ions are added then moles of HA decrease and moles of A⁻ increase.

Worked Example

Q: What happens to the pH when a small amount of NaOH is added to a buffer made of 0.10 M CH₃COOH and 0.10 M CH₃COO⁻?

Set up & assumptions
Take 1.00 L of buffer so initial moles are CH₃COOH (HA) = 0.10 mol and CH₃COO⁻ (A⁻) = 0.10 mol. Add 0.010 mol NaOH (small amount). Assume volume change is negligible.
Neutralization reaction (stoichiometry)
OH⁻ + HA → A⁻ + H₂O
New moles: HA = 0.10 − 0.010 = 0.090 mol; A⁻ = 0.10 + 0.010 = 0.110 mol; OH⁻ is fully consumed.
pH before addition (Henderson–Hasselbalch)
pH = pK_a + log([A⁻]/[HA]) = 4.76 + log(0.10/0.10) = 4.76.
pH after addition
[HA] ≈ 0.090 M; [A⁻] ≈ 0.110 M (1.00 L total).
pH = 4.76 + log(0.110/0.090) = 4.76 + log(1.22) ≈ 4.76 + 0.087 = 4.85.

A: The added OH⁻ reacts with CH₃COOH:
CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O.
CH₃COOH decreases to 0.090 M and CH₃COO⁻ increases to 0.110 M. Using Henderson–Hasselbalch, the pH changes only slightly: 4.76 → 4.85 (ΔpH ≈ +0.09), demonstrating buffer action.

Why Weak Acids and Bases Are Needed

Weak acids and bases only partially dissociate, setting up an equilibrium with their conjugates. This equilibrium helps them resist changes in pH by reacting with added H⁺ or OH⁻.

Strong acids and bases fully dissociate, so they can’t maintain this balance — meaning they can’t act as buffers.

Buffer Capacity

Buffer capacity is the amount of acid or base the buffer can absorb without a significant pH change. It depends on the absolute concentrations of HA and A⁻ and therefore the ratio [A⁻]/[HA]

Buffers are most effective when [HA] ≈ [A⁻] and this happens when pH ≈ pKa.

Summary

Buffers stabilize pH through neutralization reactions.
They rely on conjugate acid-base pairs to absorb added H⁺ or OH⁻.
Their effectiveness is greatest when pH ≈ pKa.