Question

consider an aqueous solution of calcium carbonate determine all chemical reactions taking place in this solution

Answer 1

The solubility of Calcium carbonate is very poor in pure water (47 mg/L at 1 atm).

Dissolution of calcium carbonate on the right and its solubility constant on the left shown below.

CaCO3 ⇌ Ca2+ + CO32−

Ksp = 3.7×10−9 to 8.7×10−9 at 25 °C

  The complications occur when the equilibrium of carbon dioxide with water takes place(see carbonic acid). CO₃²⁻ combines with H⁺ in the solution according to:

HCO3− ⇌ H+ + CO32−

Ka2 = 5.61×10−11 at 25 °C

HCO₃⁻ ( bicarbonate ion) i.e. Calcium bicarbonate is much more soluble in water than CaCO3, only in solution.

HCO₃⁻ combines with H⁺ in solution according to:

H2CO3 ⇌ H+ + HCO3−

Ka1 = 2.5×10−4 at 25 °C

H₂CO₃ dissociates into water and CO2:

H2O + CO2(dissolved) ⇌ H2CO3

Kh = 1.70×10−3 at 25 °C

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:

• 
  At atmospheric levels, CaCO₃ solubility is 47 mg/L.
• CO₂ partial pressure is below atmospheric levels, the solution becomes more alkaline. Due to dissolved CO₂, bicarbonate ion, and carbonate ion evaporates from the solution, high alkaline solution of Ca(OH)2, more soluble than CaCO₃.
• As CO₂ partial pressure increases above atmospheric, pH drops, the carbonate ion is converted into bicarbonate ion, results in higher solubility of Ca²⁺.

With varying pH, temperature and salinity:

  Addition of HCO₃⁻ will increase CO₃²⁻ concentration at any pH.

[Ca²⁺] = Ksp / [CO₃²⁻], and [CO₃²⁻] = K_a2 × [HCO₃⁻] / [H⁺]. Therefore, when HCO₃⁻concentration is known, the maximum concentration of Ca²⁺ ions :

Ca²⁺_max = (K_sp / K_a2) × ([H⁺] / [HCO₃⁻])

The solubility product for CaCO₃ (K_sp) and the dissociation constants for the DIC species (including K_a2) are all substantially affected by temperature and salinity, with the overall effect that Ca²⁺_max increases from fresh to salt water, and decreases with rising temperature, pH, or added bicarbonate level.

Solubility in a strong and weak acid solution

• the initial state is the acid solution with no Ca²⁺ (not taking into account possible CO₂dissolution) and the final state is the solution with saturated Ca²⁺. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca²⁺ and A⁻ so that the neutrality equation reduces approximately to 2[Ca²⁺] = [A⁻] yielding . When the concentration decreases, [HCO₃⁻] becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of CaCO₃ in pure water.

• For the same total acid concentration, the initial pH of the weak acid is less acid than the one of the strong acid; however, the maximum amount of CaCO₃ which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the pK_A, so that the weak acid is almost completely dissociated, yielding in the end as many H⁺ ions as the strong acid to "dissolve" the calcium carbonate.