Buffers (AQA A Level Chemistry): Revision Note

Acid & Basic Buffers

• A buffer solution is a solution that 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 consist of a weak acid - conjugate base or a weak base - conjugate acid

Ethanoic acid and sodium ethanoate as a buffer

• A common acidic 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

CH₃COOH (aq)$\rightleftharpoons$ H⁺ (aq) + CH₃COO^- (aq)

• Sodium ethanoate is a salt that fully ionises in solution

CH₃COONa (s) → 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 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

CH₃COOH (aq)$\rightleftharpoons$ 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

Basic Buffers

• A basic buffer can be made from ammonia (NH₃) and ammonium chloride (NH₄Cl)

• Ammonia is a weak base. It reacts with water only partially:

NH₃ (aq) + H₂O(aq) $\rightleftharpoons$ NH₄⁺ (aq) + OH^-(aq)

• Ammonium chloride dissolves completely in water to produce ammonium ions:

NH₄Cl(s) $\to$ NH₄⁺ (aq) + OH^-(aq)

• The buffer therefore contains a significant amount of both NH₃ and NH₄⁺

• When acid is added, the added hydrogen ions (H⁺) react with the ammonia:

NH₃ (aq) + H⁺(aq) $\rightleftharpoons$ NH₄⁺ (aq)

• This removes most of the added H⁺ before it can significantly increase the hydrogen ion concentration. As a result, the pH decreases only slightly

• When alkali is added, the added hydroxide ions (OH⁻) react with the ammonium ions:

NH₄⁺ (aq) + OH^-(aq) $\to$NH₃ (aq) + H₂O(aq)

• This removes most of the added OH⁻ by converting it into water. Again, the pH changes only slightly

Buffer Calculations

• 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 = $\frac{\left[salt\right]\left[H^{+}\right]}{\left[acid\right]}$ which can be rearranged to [H⁺] = K_a x $\frac{\left[acid\right]}{\left[salt\right]}$

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

-log₁₀[H⁺] = -log₁₀K_a + (-log₁₀ $\frac{\left[acid\right]}{\left[salt\right]}$)

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

pH = pK_a + log₁₀ $\frac{\left[salt\right]}{\left[acid\right]}$

• This is known as the Henderson-Hasselbalch equation

Dilution

• On dilution, the ratio [salt]/[acid] is unchanged, so [H⁺] and hence pH stay the same

Related topics

• Acid Dissociation Constant

• pH Curves and Choosing an Indicator

Applications of Buffers

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 malfunction and eventually lead to a 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