Buffer Action (Oxford AQA International A Level Chemistry)

Acidic & Basic Buffers

• 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 by maintaining an almost constant concentration of hydrogen and hydroxide ions in a solution

Acidic Buffers

• Acidic buffers are made from a weak acid and a soluble salt of the acid

• A common example is an aqueous mixture of ethanoic acid and sodium ethanoate

• In a solution:

  • Ethanoic acid partially ionises to form a relatively low concentration of ethanoate ions

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

ethanoic acid        ⇌                ethanoate

high conc              ⇌                 low conc

• Sodium ethanoate is a salt which fully ionises in solution

CH₃COONa + aq → Na⁺ (aq) + CH₃COO^- (aq) 

sodium ethanoate      →                 ethanoate ion

low conc.                  →                   high conc.

• The buffer solution there contains relatively high concentrations of CH₃COOH and CH₃COO^-

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

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

high conc.                               high conc.

• When H⁺ ions (acid) 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 (base) 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

Basic Buffers

• A basic buffer keeps the pH of a solution at above 7

• It is made by mixing a solution of a weak base with its salt

• An example is a mixture of aqueous ammonia, NH₃ (aq), and ammonium chloride, NH₄Cl (aq)

• In a solution:

NH₃ (aq) + H₂O (l) NH₄⁺ (aq) + OH^– (aq)

The equilibrium lies to the left as NH₃ is a weak base

NH₄Cl (aq) → NH₄⁺ (aq) + Cl^– (aq)

NH₄Cl is a soluble salt so fully dissociates in solution

• The mixture therefore contains high concentrations of NH₃ (aq) and NH₄⁺ (aq) which will be able to react with any H⁺ and OH^– added

Adding acid to an basic buffer

• If H⁺ (acid) is added:

  • NH₃ (aq) + H⁺ (aq) NH₄⁺ (aq) 

  • H⁺ will combine with NH₃ to form NH₄⁺ so removing any added H⁺

Adding base to a basic buffer

• If OH^–(base) is added:

  • NH₄⁺ (aq) + OH^– (aq) NH₃ (aq) + H₂O (l)

  • OH^– will combine with the acid NH₄⁺ and form NH₃ and H₂O so removing any added OH^–

• Therefore there is no overall change is pH if there are small amounts of acid or base are added

Applications of Buffers

• 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

  • 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

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

• Buffers are also used in every day products such as shampoos and detergents

Buffer Calculations

• The pH of an acidic 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