Investigate the application of buffer systems in natural and industrial contexts, including the bicarbonate buffer in blood and the Henderson-Hasselbalch description of buffer pH

What this dot point is asking

NESA wants you to consolidate buffer chemistry by applying it to natural and industrial contexts, especially the bicarbonate buffer in blood, and to use the Henderson-Hasselbalch equation quantitatively. You should be able to write the buffer equilibria, explain the response to added strong acid or base in equation form, and link the chemistry to physiological situations like exercise, hyperventilation, and acidosis. This builds on conjugate acid-base pairs and the Module 5 page on buffer systems.

The answer

Buffer composition and action (recap)

A buffer is a solution containing significant amounts of both a weak acid and its conjugate base (or a weak base and its conjugate acid). The two species sit in equilibrium:

Response to added strong acid (extra $H^+$). The conjugate base consumes it:

Response to added strong base (extra $OH^-$). The weak acid consumes it:

In each case, the strong reagent is converted into the corresponding member of the conjugate pair, so $[H^+]$ moves only slightly. The buffer fails when one component is essentially exhausted.

The Henderson-Hasselbalch equation

Take logarithms of the $K_a$ expression:

This is the Henderson-Hasselbalch equation. Three quick consequences:

• When $[A^-] = [HA]$ (equimolar buffer), $pH = pK_a$. Buffer capacity is maximised here.
• The useful range of a buffer is roughly $pK_a \pm 1$ (corresponding to a 10:1 ratio of components on either side).
• To make a buffer of a target pH, choose a weak acid with $pK_a$ within one unit of the target, then set the ratio.

The bicarbonate buffer in blood

Arterial blood is maintained at pH $7.40 \pm 0.05$ by a coupled buffer-respiratory-renal system. The dominant chemical buffer is the bicarbonate pair:

For this system $pK_{a1} = 6.10$ at body temperature, with typical concentrations $[HCO_3^-] \approx 24$ mmol/L and $[H_2CO_3] \approx 1.2$ mmol/L (in equilibrium with dissolved $CO_2$).

The ratio is far from 1:1, so chemically this is a poor buffer in isolation. What makes it physiologically powerful is that both components are continuously regulated:

• IMATH_22 is controlled by breathing rate (the lungs expel $CO_2$).
• IMATH_24 is controlled by the kidneys (which excrete or retain bicarbonate).

This "open" buffer can therefore reset its components in response to disturbances, something a closed buffer in a beaker cannot do.

Acid-base disturbances

In each case the body shifts the bicarbonate equilibrium to restore the 20:1 ratio.

Other physiological buffers

• Phosphate buffer ($H_2PO_4^-$/$HPO_4^{2-}$): $pK_a = 7.20$. Important inside cells where bicarbonate is less effective. Also the basis for laboratory buffers (PBS).
• Protein buffers (haemoglobin in particular): histidine side chains have $pK_a$ near 6, so they buffer near physiological pH. Haemoglobin doubles as the oxygen carrier and a major intracellular buffer.

Industrial and laboratory buffers

Worked example

You need 500 mL of a pH 5.00 buffer using ethanoic acid ($pK_a = 4.74$) and sodium ethanoate. Calculate the ratio $[CH_3COO^-]/[CH_3COOH]$ required, and a workable pair of concentrations.

Step 1: Apply Henderson-Hasselbalch.

Step 2: Choose concentrations. Pick total buffer concentration ~0.20 mol/L for good capacity. With ratio 1.82:1, set $[HA] = 0.071$ mol/L and $[A^-] = 0.129$ mol/L (sum 0.20 mol/L).

Step 3: Check.

Dissolve $0.071 \times 0.500 = 0.0355$ mol $CH_3COOH$ and $0.129 \times 0.500 = 0.0645$ mol $CH_3COONa$ in water and make up to 500 mL.

Common traps

Calling any weak acid a buffer. A weak acid alone is not a buffer. A buffer needs both the weak acid and its conjugate base in comparable amounts.

Forgetting the 20:1 ratio in blood. Most textbook buffers operate near 1:1 ($pH = pK_a$). Blood is held far from 1:1 because the components are open (lungs and kidneys), not closed.

Confusing buffer capacity with buffer pH. Capacity is how much acid or base the buffer can absorb before pH shifts significantly; pH is the operating point. A dilute buffer has low capacity but the same pH as a concentrated one of the same ratio.

Using strong acid in a buffer recipe. Henderson-Hasselbalch assumes the weak acid only partly ionises. A strong acid + strong base mixture is not a buffer.

Misnaming acidosis and alkalosis. Acidosis is pH below 7.35 (more acidic); alkalosis is pH above 7.45 (more basic). Adding "respiratory" or "metabolic" specifies the origin.

Ignoring the sign on log inside Henderson-Hasselbalch. $\log_{10}(ratio < 1)$ is negative, so pH is below $pK_a$ when the acid form dominates.

In one sentence

A buffer is a weak acid plus its conjugate base in comparable amounts, with $pH = pK_a + \log_{10}([A^-]/[HA])$ (Henderson-Hasselbalch); the bicarbonate buffer in blood is held at pH 7.4 not by a 1:1 ratio but by an open 20:1 ratio of $HCO_3^-$ to $H_2CO_3$, where breathing rate regulates $H_2CO_3$ (via $CO_2$) and the kidneys regulate $HCO_3^-$, allowing the body to compensate for both respiratory and metabolic acid-base disturbances.