Investigate the structure and properties of buffer systems, including their composition, how they resist pH change, and their importance in natural systems such as blood

What this dot point is asking

NESA wants you to describe what a buffer is, explain how the weak acid / conjugate base equilibrium resists pH change, use the Henderson-Hasselbalch equation for buffer pH calculations, and apply this to natural buffer systems (especially the carbonic acid / bicarbonate buffer in blood). This is a popular extended-response topic combining Brønsted-Lowry acid-base theory with Le Chatelier reasoning.

The answer

What a buffer is

A buffer solution resists changes in pH when small amounts of acid or base are added (or when the solution is moderately diluted). A buffer contains comparable concentrations of:

• a weak acid (HA) and
• its conjugate base ($A^-$).

Or equivalently, a weak base and its conjugate acid.

The buffer equilibrium

For an acetic acid / acetate buffer:

Both species are present in substantial concentration. The equilibrium has "reserves" in both directions.

Adding $H^+$. The conjugate base neutralises it: $CH_3COO^- + H^+ \rightarrow CH_3COOH$. Equilibrium shifts left. The pH drops only slightly.

Adding $OH^-$. The weak acid donates a proton: $CH_3COOH + OH^- \rightarrow CH_3COO^- + H_2O$. Equilibrium shifts right. The pH rises only slightly.

Both responses are applications of Le Chatelier's principle to the buffer equilibrium.

The Henderson-Hasselbalch equation

For a buffer of weak acid HA with conjugate base $A^-$:

Useful results:

• When $[A^-] = [HA]$, $\log(1) = 0$ and **pH = $pK_a$**.
• A buffer is most effective within ±1 pH unit of its $pK_a$.
• The buffer pH depends on the ratio of $[A^-]$ to $[HA]$, so moderate dilution does not change pH much.

Buffer capacity

A buffer has a finite capacity. Once enough acid is added to consume all the conjugate base (or enough base to consume all the weak acid), the buffer fails and the pH changes rapidly. Capacity is maximised when $[HA]$ and $[A^-]$ are equal and both large.

The blood buffer system

Blood is buffered between pH 7.35 and 7.45 by the carbonic acid / bicarbonate system:

The buffer pair is $H_2CO_3$ ($pK_a$ ≈ 6.4) and $HCO_3^-$. At blood pH 7.4, the Henderson-Hasselbalch ratio $[HCO_3^-] / [H_2CO_3]$ is about 20:1.

The system is biologically powerful because both ends are open:

• Lungs regulate $CO_2$. Hyperventilating expels $CO_2$, shifting equilibrium left and raising pH.
• Kidneys regulate $HCO_3^-$. They can secrete or retain bicarbonate to compensate over hours to days.

When the buffer fails, the body experiences acidosis (low pH) or alkalosis (high pH), with consequences for enzyme activity, oxygen transport, and ionic balance.

Worked example

Calculate the pH of a buffer prepared by mixing 0.20 mol/L $CH_3COOH$ and 0.30 mol/L $CH_3COONa$. $K_a$ of acetic acid is $1.8 \times 10^{-5}$.

Step 1: Find $pK_a$.

Step 2: Apply Henderson-Hasselbalch.

The buffer pH is 4.92, close to the $pK_a$ of acetic acid (as expected when the two concentrations are similar).

What if you add 0.010 mol of HCl to 1.00 L of this buffer? The added $H^+$ converts 0.010 mol of $CH_3COO^-$ into 0.010 mol of $CH_3COOH$.

New $[CH_3COO^-] = 0.30 - 0.010 = 0.29$ mol/L.
New $[CH_3COOH] = 0.20 + 0.010 = 0.21$ mol/L.

The pH drops only by 0.04 units. Compare this to adding 0.010 mol of HCl to 1.00 L of pure water (pH falls from 7 to 2). The buffer is effective.

Worked example 2: designing a buffer

Calculate the mole ratio of $CH_3COO^-$ to $CH_3COOH$ required to prepare a buffer with pH 5.20. $K_a$ of acetic acid is $1.8 \times 10^{-5}$ ($pK_a = 4.74$).

Step 1: Apply Henderson-Hasselbalch.

Step 2: Check suitability. The target pH 5.20 is within ±1 pH unit of the pKa (4.74), so acetic acid is a suitable buffer choice. For a buffer at pH 5.20, mix acetate and acetic acid in a 2.88:1 mole ratio. For 1.00 L of buffer, one recipe is 0.10 mol $CH_3COOH$ and 0.288 mol $CH_3COONa$ dissolved in water.

Step 3: Confirm by quick mental check. Above pKa means more conjugate base than acid (ratio > 1). Below pKa means more acid than base (ratio < 1). 5.20 > 4.74, so the ratio should exceed 1. 2.88 > 1. Sensible.

Common traps

Calling a strong acid solution a buffer. A strong acid is not a buffer because there is no significant conjugate base. Buffers need a weak acid plus its conjugate base (or weak base plus conjugate acid).

Forgetting that the buffer pH depends on the ratio, not the absolute concentrations. Diluting the buffer 10-fold barely changes the pH, but does reduce its capacity.

Saying buffers stop pH change. They resist change. Adding acid or base does shift pH a little; if you exceed capacity, pH shifts a lot.

Confusing $K_a$ with $K_b$. When given $K_b$ for a weak base buffer, use $pK_a = 14 - pK_b$ and Henderson-Hasselbalch with the conjugate acid as HA.

Missing the importance angle for blood. Many students write only the chemistry. Markers want a sentence on physiological consequence (enzymes, oxygen transport, acidosis/alkalosis).

In one sentence

A buffer contains comparable concentrations of a weak acid and its conjugate base, resists pH change because the equilibrium shifts (by Le Chatelier) to consume added acid or base, has pH given by the Henderson-Hasselbalch equation $pH = pK_a + \log([A^-]/[HA])$, and the carbonic acid / bicarbonate buffer keeps blood pH between 7.35 and 7.45.