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 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 solution, ethanoic acid partially ionises:

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

ethanoic acid        ⇌                ethanoate

high conc              ⇌                 low conc

• Sodium ethanoate is a salt which fully dissociates in solution:

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

sodium ethanoate      →                 ethanoate ion

low conc.                  →                   high conc.

• This results in a buffer solution with high concentrations of CH₃COOH and CH₃COO^-

  • The establishes the equilibrium with hydrogen and ethanoate ions:

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

high conc.                               high conc.

When H⁺ ions (acid) are added:

• The extra H⁺ ions react with CH₃COO^- to form CH₃COOH

• The equilibrium shifts left to reduce [H⁺]

  • This limits any drop in pH

• Since there's a large supply of CH₃COO^-, its concentration doesn’t change much

• So, the pH remains almost constant

When OH^- ions (base) are added:

• The OH^- ions react with H⁺ to form water:

OH^- (aq) + H⁺  (aq) → H₂O (l)

• The [H⁺] decreases, shifting the equilibrium right

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

• More CH₃COOH dissociates to replace H⁺ ions

• So, the pH stays nearly constant due to the buffer action

• The buffer works because it contains a reservoir of both the weak acid and its conjugate base, which can react with added H⁺ or OH⁻ to restore balance.

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:

K_a =

• This can be rearranged to give:

[H⁺] = K_a x

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

-log₁₀[H⁺] = -log₁₀K_a x -log₁₀

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

pH = pK_a + log₁₀

• 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