Quick questions on Buffer systems and resistance to pH change (QCE Chemistry Unit 3)

Two equivalent compositions both produce a buffer:

Take the ethanoate buffer at equilibrium:

Add a small amount of NaOH. The added OH- is consumed by the weak acid:

Buffer capacity is the amount of strong acid or base that can be added before the pH changes significantly. It is determined by:

A subtle point QCAA examines. Sodium ethanoate provides a high initial concentration of CH3COO- directly. By Le Chatelier, this suppresses the ionisation of the added CH3COOH (common-ion effect): the equilibrium shifts left, leaving CH3COOH largely as undissociated weak acid. So you start with both species at comparable concentrations and both ready to act as absorbers of perturbation.

The textbook real-world example. Blood pH must remain at 7.4 plus or minus 0.05 for human survival. The bicarbonate buffer is the primary mechanism.

H2CO3 produced from acid neutralisation decomposes to CO2 and water; CO2 is exhaled. Removing CO2 keeps [H2CO3] low so the buffer remains effective. Hyperventilation removes CO2 faster, raising pH; hypoventilation retains CO2, lowering pH.

Kidneys excrete or retain HCO3- on longer timescales, providing slower-acting compensation when respiratory adjustment alone is insufficient.

A solution of HCl plus NaCl is not a buffer because Cl- is not the conjugate base of a weak acid. The conjugate of a strong acid is too weak to absorb H+.

A 1:100 ratio is not an effective buffer; the minor component runs out almost immediately.

It changes slightly. The buffer reduces the change, not eliminates it.

The buffering equilibrium chain extends to CO2 gas exchange; QCAA expects you to bring the lungs (and sometimes the kidneys) into the explanation.

When a strong acid or base reacts with the buffer component, the reaction goes essentially to completion (single arrow). The underlying buffer equilibrium has the double arrow.