Quick questions on Buffer systems explained: HSC Chemistry Module 5

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:

For an acetic acid / acetate buffer:

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

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.

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

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

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

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.

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.

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).

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

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

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