HSC Chemistry – Acid-base Buffers

Last Update: 3 December 2025

This topic is part of the HSC Chemistry course under the topic Quantitative Analysis.

HSC Chemistry Syllabus

conduct a practical investigation to prepare a buffer and demonstrate its properties (ACSCH080)
describe the importance of buffers in natural systems (ACSCH098, ACSCH102)

Buffers

This video aims to help you understand the following

What is a buffer?
How does a buffer work?
Examples of buffers in nature
How to calculate the pH of a buffer mixture using the Henderson-Hasselbach equation.

What is an Acid-Base Buffer?

A buffer is a chemical system that resists changes in pH when a small amount of acid or base is added.

A buffer contains either

comparable amounts of a weak acid and its conjugate base or
comparable amounts of a weak base and its conjugate acid

In simpler words, the amount of weak acid or base must be similar to that of its conjugate base/acid. A buffer's buffering capacity decreases as the two components become more disproportionate.

The strength of acids and bases is important when making a buffer.

Strong acids and strong bases cannot be used to make buffers as they completely ionise in water.

For example, hydrochloric acid is a strong acid that completely dissociates:

$$HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$$

When a base is added, hydroxide ions are quickly neutralised by the presence of hydrogen ions. This will cause the pH to increase markedly as `[H^+]` decreases.

Such a chemical system is unable to minimise the decrease in pH when a small amount of acid is added because the dissociation of HCl is not reversible.

How Does a Buffer Work?

When a small amount of acid is added, a buffer minimises the decrease in pH. When a base is added, a buffer minimises the increase in pH.

Consider a buffer consisting of 0.10 mol/L hydrofluoric acid (HF) and 0.10 mol/L of its conjugate base, fluoride ion. Hydrofluoric acid and fluoride ion will establish an equilibrium as shown by the equation:

$$HF(aq) + H_2O(l) \rightleftharpoons F^-(aq) + H_3O^+(aq)$$

When a small amount of an acid is added to this buffer, `[H_3O^+]` increases. By Le Chatelier's principle (LCP), the equilibrium position will shift to the left to reduce `[H_3O^+]`. Since the concentration of hydronium ions remains relatively constant, the decrease in pH is minimised as `pH = -log[H_3O^+]`.

When a small amount of a base is added, `[OH^-]` increases. Hydroxide ions will neutralise hydronium ions in the buffer equilibrium, causing `[H_3O^+]` to decreases. By LCP, the equilibrium position will shift to the right to increase `[H_3O^+]`. Again, since the concentration of hydronium ions remains relatively constant, the increase in pH is minimised.

Buffering Capacity

Buffering capacity is a buffer's effectiveness in minimising changes in pH when an acid or base is added. Buffering capacity of any buffer decreases with the addition of an acid or base.

For example in the hydrofluoric acid/fluoride ion equilibrium above, when excessive amounts of acid is added, the amount of fluoride ions decreases due to the constant shift in equilibrium position to the left. When there are no more fluoride ions present, the equilibrium position can no longer shift towards the left side, causing the pH to decrease at a greater rate.

Conversely, when excessive amounts of base is added, the amount of HF molecules decreases due to the constant shift in equilibrium position to the right. When there are no more HF molecules present, the equilibrium position can no longer shift towards the right side, causing the pH to increase at a greater rate.

For any buffer, it is most effective (highest buffering capacity) when there are equal quantities or concentrations of the acid and its conjugate base (or base and conjugate acid). Therefore, the buffering capacity is dependent on the ratio between the conjugate acid-base pair. This ratio also affects the pH of the buffer (see below).

Buffers in Natural and Biological Systems

Buffer in Human Blood

Human blood buffer consists of carbonic acid and its conjugate base, hydrogen carbonate (bicarbonate):

$$H_2CO_3(aq)+H_2O(l) \rightleftharpoons H_3O^+(aq)+HCO_3^-(aq)$$

The physiological pH of human blood is maintained between 7.35 and 7.45. If pH falls outside this range, the structure of proteins, concentrations of electrolytes and cellular structure can be affected.

Cellular respiration produces carbon dioxide which readily dissolves in blood to produce carbonic acid. This will lead to acidosis (decrease in pH of blood). However, the presence of the carbonic buffer minimises the decrease in pH until carbon dioxide is gradually removed from blood in the lungs followed by expiration.

Buffer in Cells

The intracellular fluid contains a buffer system composed of dihydrogen phosphate and hydrogen phosphate:

$$H_2PO_4^-(aq)+H_2O(l) \rightleftharpoons HPO_4^{2-}(aq)+H_3O^+(aq)$$

The pH of intracellular fluid is maintained between 7.0 and 7.4. This is required because enzymes in cells can only function properly within this range. When the pH becomes too acidic or too basic, the structure of enzymes can be denatured.

Buffer in Natural Water Bodies

Carbonic acid/hydrogen carbonate buffer system is also present in water bodies (rivers, lakes, oceans).

$$H_2CO_3(aq)+H_2O(l) \rightleftharpoons H_3O^+(aq)+HCO_3^-(aq)$$

The maintenance of pH of water bodies is important for sustaining marine life. Different water bodies have different pH ranges. For example, coral reef prefers a pH range between 8.2 and 8.3. Slightly acidification can hinder coral skeleton growth.

Henderson-HasselBach Equation

The Henderson-HasselBach equation is used to calculate the pH of a buffer mixture when the strength and relative composition of its conjugate acid-base pair are known. Refer to the video for derivation of this equation.

Consider a generic buffer system:

$$HA(aq) + H_2O(l) \rightleftharpoons A^-(aq) + H_3O^+(aq)$$

The pH of this buffer is given by

$$pH = pKa + \log{\frac{[A^-]}{[HA]}}$$

$$pH = pKa - \log{\frac{[HA]}{[A^-]}}$$

where `[HA]` and `[A^-]` are equilibrium concentrations of the weak acid and its conjugate base respectively.

Using this equation, when concentrations of acid and its conjugate base are equal, the pH of the mixture equals to the pKa of the weak acid. This is also the pH at which the weak acid is exactly 50% ionised, the buffering capacity of the buffer is at its maximum.

When acid is added to a buffer, the equilibrium position shifts to the left to counteract this change by reducing the change in `H_3O^+` ions. This in turn increases `[HA]` and decreases `[A^-]`, and hence the ratio between the former and latter increases.

Conversely, when base is added to a buffer, the equilibrium position shifts to the right to counteract this change. This in turn decreases the ratio between `[HA]` and [A^-]`.

Example (Solution in Video)

The pKa of hypochlorous acid, HOCl, is 7.5.

(a) A buffer consists of equal concentrations of HOCl and NaOCl. What is the pH of this buffer?

(b) At equilibrium, a buffer consists of 1.00 mol L^–1 HOCl and 0.75 mol L^–1 OCl^–. Calculate the pH for this buffer.

Check Your Understanding

A buffer was prepared with acetic acid and sodium acetate. A few drops of universal indicator were then added. When small amounts of either 0.1 mol L^–1 HCl(aq) or 0.1 mol L^–1 NaOH(aq) were added, no change in the colour of the solution was observed.

Explain these observations. Support your answer with at least ONE chemical equation.