Buffers (OCR A Level Chemistry A): Revision Note

Buffer Solutions

• A buffer solution is a solution which resists changes in pH when small amounts of acids or alkalis are added

  • A buffer solution is used to keep the pH almost constant

  • A buffer can consists of weak acid - conjugate base or 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 is a weak acid and partially ionises in solution to form a relatively low concentration of ethanoate ions

• Sodium ethanoate is a salt which fully ionises in solution

• There are reserve supplies of the acid (CH₃COOH) and its conjugate base (CH₃COO^-)

  • The buffer solution contains relatively high concentrations of CH₃COOH (due to partial ionisation of ethanoic acid) and CH₃COO^- (due to full ionisation of sodium ethanoate)

• In the buffer solution, the ethanoic acid is in equilibrium with hydrogen and ethanoate ions

• 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 hydrogen ions are added to the solution the pH of the solution would decrease; However, the ethanoate ions in the buffer solution react with the hydrogen ions to prevent this and keep the pH 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

When hydroxide ions are added to the solution, the hydrogen ions react with them to form water; The decrease in hydrogen ions would mean that the pH would increase however the equilibrium moves to the right to replace the removed hydrogen ions and keep the pH constant

Calculating the pH of a Buffer Solution

• The pH of a buffer solution can be calculated using:

  • The K_a of the weak acid

  • The equilibrium concentration of the weak acid and its conjugate base (salt)

• To determine the pH, the concentration of hydrogen ions is needed which can be found using the equilibrium expression

• To simplify the calculations, logarithms are used such that the expression becomes:

• Since -log₁₀ [H⁺] = pH, the expression can also be rewritten as:

• This is known as the Hendersen-Hasselbalch equation

Uses of Buffers

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 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