HSC Chemistry – Acid/base Buffers

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 discusses everything about buffers including

  • What a buffer is
  • How a buffer works
  • When a buffer doesn't work

 

What is a buffer?

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

What does a buffer contain? 

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.

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 via 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.