Question

Oxides of sulfur are important in atmospheric pollution, arising particularly from burning coal. Use the thermodynamic data at 25o C given in the appendix to answer the following questions.

(a) In air, the oxidation of SO2 can occur: ½ O2(g) + SO2 (g) → SO3 (g). Calculate ∆Go rxn,298.

(b) Find the equilibrium ratio of partial pressures of SO3 (g) to SO2(g) in air at 25o C. The partial pressure of O2(g) is 0.21 bar.

(c) SO3(g) can react with H2O (g) to form sulfuric acid, H2SO4 (g). Air that is in equilibrium with liquid water at 25o C has a partial pressure of H2O (g) of 0.031 bar. Find the equilibrium ratio of partial pressures of H2SO4 (g) to SO3 (g) in air at 25o C.

Answer 1

In this case I will use the ∆Gf of these species:
∆Gf SO3 = -370.4 kJ/mol
∆Gf SO2 = -300.4 kJ/mol
∆Gf O2 = 0

∆Grxn = (-370.4) - (-300.4) = -70 kJ/mol or 70000 J/mol

To get the ratio, we need the value of Kp:

∆G = -RTlnKc
Kc = exp(-∆G/RT)
Kc = exp(70000 / 8.3144 * 298) = 1.86x10¹²

Kp = Kc(RT)∆n
∆n = 1 - (1-1/2) = 1/2
Kp = 1.86x10¹² (0.0821*298)^1/2 = 9.2x10¹²

The ratio would be:

kp = pSO3 / pO2^1/2 pSO2
kp*pO₂^1/2 = pSO3/pSO2
pSO3/pSO2 = 9.2x10¹² * (0.21)^1/2
pSO3/pSO2 = 4.22x10¹²

Part c) you have to calculate ∆Grxn and kp again. Try to do it yourself, you already have the guidance.