Buffers (CIE A Level Chemistry)

• A buffer solution is a solution in which the pH does not change a lot when small amounts of acids or alkalis are added
  • A buffer solution is used to keep the pH almost constant
  • A buffer can consist of a weak acid - conjugate base or a weak base - conjugate acid

Ethanoic acid & sodium ethanoate as a buffer

• A common buffer solution is an aqueous mixture of ethanoic acid and sodium ethanoate
• Ethanoic acid, CH₃COOH, is a weak acid and partially ionises in solution to form a relatively low concentration of ethanoate ions, CH₃COO^– 

CH₃COOH (aq)  H⁺ (aq) + CH₃COO^– (aq)

• Sodium ethanoate, CH₃COONa, is a salt which fully ionises in solution to form a relatively high concentration of ethanoate ions, CH₃COO^– 

CH₃COONa (a) + aq → Na⁺ (aq) + CH₃COO^– (aq)

• There are reserve supplies of the acid (CH₃COOH) and its conjugate base (CH₃COO^-)
  • The buffer solution contains relatively high concentrations of CH₃COOH (due to the ionisation of ethanoic acid) and CH₃COO^- (due to the ionisation of sodium ethanoate)
• In the buffer solution, the ethanoic acid is in equilibrium with hydrogen and ethanoate ions

CH₃COOH (aq)  H⁺ (aq) + CH₃COO^– (aq)

When H⁺ ions are added:

• The equilibrium position shifts to the left as H⁺ ions react with CH₃COO^– ions to form more CH₃COOH until equilibrium is re-established
• As there is a large reserve supply of CH₃COO^–, the concentration of CH₃COO^– in solution doesn’t change much as it reacts with the added H⁺ ions
• As there is a large reserve supply of CH₃COOH, the concentration of CH₃COOH in solution doesn’t change much as CH₃COOH is formed from the reaction of CH₃COO^– with H⁺
• As a result, the pH remains reasonably constant

When OH^- ions are added:

• The OH^- reacts with H⁺ to form water

OH^- (aq) + H⁺  (aq) → H₂O (l)

• The H⁺ concentration decreases
• The equilibrium position shifts to the right and more CH₃COOH molecules ionise to form more H⁺ and CH₃COO^- until equilibrium is re-established

CH₃COOH (aq) → H⁺ (aq) + CH₃COO^- (aq)

• As there is a large reserve supply of CH₃COOH the concentration of CH₃COOH in solution doesn’t change much when CH₃COOH dissociates to form more H⁺ ions
• As there is a large reserve supply of CH₃COO^- the concentration of CH₃COO^- in solution doesn’t change much
• As a result, the pH remains reasonably constant

Uses of buffer solutions in controlling the pH of blood

• In humans, HCO₃^- ions act as a buffer to keep the blood pH between 7.35 and 7.45
• Body cells produce CO₂ during aerobic respiration
• This CO₂ will combine with water in the blood to form a solution containing H⁺ ions

CO₂ (g) + H₂O (l) ⇌ H⁺ (aq) + HCO₃^- (aq)

• This equilibrium between CO₂ and HCO₃^- is extremely important
• If the concentration of H⁺ ions is not regulated, the blood pH would drop and cause ‘acidosis’
  • Acidosis refers to a condition in which there is too much acid in the body fluids such as blood
  • This could cause body malfunctioning and eventually lead to coma
• If there is an increase in H⁺ ions
• The equilibrium position shifts to the left until equilibrium is restored

CO₂ (g) + H₂O (l) → H⁺ (aq) + HCO₃^- (aq)

• This reduces the concentration of H⁺ and keeps the pH of the blood constant
• If there is a decrease in H⁺ ions
  • The equilibrium position shifts to the right until equilibrium is restored

CO₂ (g) + H₂O (l) → H⁺ (aq) + HCO₃^- (aq)

• This increases the concentration of H⁺ and keeps the pH of the blood constant