Question

Increased gaseous CO2 concentrations are leading to decrease pH of ocean water which in turn causes harm to shellfish populations. Write all the appropriate equilibria (solubility and acid ionization) relevant to these phenomena. In complete sentences briefly, explain the phenomena using Le Chatliers Principle.

Answer 1

Gaseous CO₂ dissolves in water as per the equation:

CO₂ (g) <====> CO₂ (aq)

The dissolved CO₂ in water reacts with water to produce carbonic acid (H₂CO₃) as per the equation:

CO₂ (aq) + H₂O (l) <====> H₂CO₃ (aq)

The equilibrium constant for the reaction is given by

K_eq = [H₂CO₃]/[CO₂] …..(1) (the concentration of H₂O, i.e, [H₂O] is constant since the ocean contains a large excess of H₂O).

Now, carbonic acid is a weak diprotic acid and dissociates in water as below:

H₂CO₃ (aq) <====> H⁺ (aq) + HCO₃^- (aq)

The acid ionization constant is

K_a = [H⁺][HCO₃^-]/[H₂CO₃]

The pH of the carbonate/bicarbonate system is given by

pH = pK_a + log [HCO₃^-]/[H₂CO₃] …..(2)

Now, when the CO₂ concentration in sea-water increases, more CO₂ gets dissolved in water and this increases [CO₂]_aq. Look at expression (1) above; the equilibrium constant K_eq must remain constant as the temperature remains constant (K_eq depends only on temperature). Therefore, as [CO₂]_aq increases, the denominator of expression (1) increases and hence the numerator must increase to counter balance the effect of increased CO₂. Therefore, the equilibrium shifts to the right and the product is favoured.

Now, as [H₂CO₃] increases, the pH of sea-water must decrease. The reason can be explained using expression (2). pK_a is the acid ionization constant and must remain constant. [H₂CO₃] appears in the denominator of the logarithm in expression (2) and as [H₂CO₃] increases, the logarithm becomes more negative and hence, the pH value decreases.