Topic 14 · page 17 of 18

Buffers in blood, cells and food

Enzymes and proteins only work in a narrow pH range, and microorganisms spoil food by changing its pH. Buffers keep the pH steady in both.

10 min read14.22(i)14.22(ii)

Read first: What a buffer is and how it works, pH of a buffer

Why living things need buffers 14.22(i)

Enzymes and other proteins hold their shape through interactions between charged groups. A change in pH changes those charges, the shape changes, and the enzyme stops working. So cells and the fluids around them must stay inside a narrow pH range. Reactions in the body keep making acids (for example lactic acid in muscle, and carbon dioxide from respiration), so a buffer is needed to absorb them.

The buffer in blood 14.22(i)

Blood is held at pH 7.35 to 7.45. The buffer is carbonic acid, H₂CO₃, and its conjugate base, the hydrogencarbonate ion, HCO₃⁻. Carbon dioxide from respiration dissolves in the blood and forms the carbonic acid:

COX2(aq)+HX2O(l)⇌HX2COX3(aq)⇌HX+(aq)+HCOX3X−(aq) \ce{CO2(aq) + H2O(l) <=> H2CO3(aq) <=> H+(aq) + HCO3-(aq)}
The carbonic acid and hydrogencarbonate buffer in blood. Three boxes are linked by equilibrium arrows: carbon dioxide plus water, carbonic acid, and hydrogen ions plus hydrogencarbonate ions. A downward arrow from the hydrogencarbonate box says extra H plus, for example from lactic acid, is mopped up by hydrogencarbonate. A downward arrow from the carbonic acid box says that for extra OH minus, carbonic acid gives up H plus to replace what was removed.
The H₂CO₃ / HCO₃⁻ buffer. It works like any acidic buffer.

The pair works just like ethanoic acid and ethanoate. The equilibrium shifts to remove what has been added, and the ratio [HCO₃⁻] ÷ [H₂CO₃] changes only a little.

Worked exampleThe pH of blood

Typical values in blood are [HCO₃⁻] = 24 mmol dm⁻³ and [H₂CO₃] = 1.2 mmol dm⁻³. The pKa\text{p}K_\text{a} of H₂CO₃ in blood is about 6.1. Calculate the pH.

  1. KaK_\text{a} = 10−6.110^{-6.1} = 7.94 × 10⁻⁷ mol dm⁻³
  2. [H⁺] = KaK_\text{a} × [H₂CO₃] ÷ [HCO₃⁻] = 7.94 × 10⁻⁷ × (1.2 ÷ 24) = 3.97 × 10⁻⁸ mol dm⁻³
  3. pH = 7.40

The ratio is 20 : 1 in favour of the base, which is well above 1. That means the blood buffer has a large reserve of HCO₃⁻, and is especially good at absorbing acid. Adding 0.5 mmol dm⁻³ of H⁺ makes [HCO₃⁻] = 23.5 and [H₂CO₃] = 1.7 mmol dm⁻³, and the pH falls from 7.40 only to 7.24. The same amount of acid in pure water would give pH 3.3.

Buffers in food 14.22(ii)

Bacteria and fungi grow on food and release acids as they do. Milk turns sour because bacteria turn lactose into lactic acid. Without a buffer the pH would fall as the acid builds up, and the food would change in taste, texture and colour, and would deteriorate faster.

A weak acid and its salt are added to keep the pH steady. Examples are citric acid with sodium citrate, and ethanoic acid with sodium ethanoate. They work exactly as the buffers on the earlier pages do: the acid half removes added alkali, and the base half removes the acid that the microorganisms produce.

Check yourself

  1. Name the buffer system that keeps the pH of blood steady and write an equation to show how it responds to added acid.
  2. Explain why the pH of blood and cells must be kept in a narrow range.
  3. A sample of blood contains 24.0 mmol dm⁻³ HCO₃⁻ and 1.50 mmol dm⁻³ H₂CO₃. Calculate its pH. The pKa\text{p}K_\text{a} of H₂CO₃ is 6.1.
  4. Explain how bacterial growth can spoil a food, and how a buffer helps to prevent this.
  5. Write an equation to show how the blood buffer removes added alkali.
  6. State why a buffer in the blood would stop working if a large amount of acid were added.

Answers cover these until you have tried the questions

  1. The carbonic acid / hydrogencarbonate pair, H₂CO₃ / HCO₃⁻. Added acid is removed by HCO₃⁻: HCOX3X−+HX+→HX2COX3\ce{HCO3- + H+ -> H2CO3}.
  2. Enzymes and other proteins work only within a narrow pH range. A change in pH alters the charges on a protein, and so its shape, and the enzyme stops working.
  3. Ratio [HCO₃⁻] ÷ [H₂CO₃] = 24.0 ÷ 1.50 = 16.0. pH = pKa\text{p}K_\text{a} + lg 16.0 = 6.1 + 1.20 = 7.30 (using KaK_\text{a}: [H⁺] = 7.94 × 10⁻⁷ × 1.50 ÷ 24.0 = 4.96 × 10⁻⁸, pH 7.30).
  4. Bacteria and fungi produce acids as they grow. These lower the pH, which changes the taste, texture and colour of the food and shortens its life. A buffer (a weak acid and its salt) absorbs the added acid, because the base half reacts with the H⁺, so the pH hardly changes.
  5. HX2COX3+OHX−→HCOX3X−+HX2O\ce{H2CO3 + OH- -> HCO3- + H2O} (the acid half of the buffer removes the OH⁻).
  6. The buffer has a limited capacity. If so much acid is added that nearly all the HCO₃⁻ is used up, nothing is left to react with more H⁺, and the pH then falls quickly.