Question

Pyrite is oxidized at Earth’s surface via the reaction This reaction is what causes “acid mine drainage.” Is the reaction spontaneous? What is the Keq for this reaction?

Answer 1

The reaction given for the oxidation of pyrite is:

4FeS2(s) + 15O2(g) + 14H2O(l) → 8H2SO4(aq) + 4Fe(OH)3(s)

To determine whether the reaction is spontaneous, we need to calculate the Gibbs free energy change (ΔG) for the reaction. If ΔG is negative, the reaction is spontaneous, and if ΔG is positive, the reaction is non-spontaneous.

We can calculate ΔG using the equation:

ΔG = -RTln(K)

where R is the gas constant (8.314 J/mol·K), T is the temperature in Kelvin, ln is the natural logarithm, and K is the equilibrium constant.

At standard conditions (25°C or 298 K and 1 atm), we can look up the standard Gibbs free energy change of formation (ΔG°f) values for the reactants and products and calculate the standard Gibbs free energy change (ΔG°) for the reaction:

ΔG° = ΣnΔG°f(products) - ΣnΔG°f(reactants)
ΔG° = [8(-1017.2) + 4(-1133)] - [4(0) + 15(0) + 14(-237.2)]
ΔG° = -8081.6 J/mol

Since ΔG° is negative, the reaction is spontaneous under standard conditions.

We can then calculate the equilibrium constant (K) using the equation:

ΔG° = -RTln(K)

K = e^(-ΔG°/RT)

K = e^(-(-8081.6)/(8.314*298))

K = 2.17 x 10^39

Therefore, the equilibrium constant (Keq) for the oxidation of pyrite is 2.17 x 10^39.