Question

The equilibrium concentrations of the species in the reaction A+B$\to$C+D are 2, 3, 10 and 6 mol L–1, respectively at 300 K. ΔGº for the reaction is (R = 2 cal/mol K)

• A. –137.26 cal
• B. –1381.80 cal
• C. –13.73 cal
• D. –1372.60 cal

Answer 1

Answer: (B)

–1381.80 cal

Answer 2

The standard Gibbs free energy change Δ G ∘=− RT ln(K)
Where
R = 2 cal/mol K
T = 300 K
K = equilibrium constant
Given the equilibrium concentrations:
[A]=2 mol/L
[B]=3 mol/L
[C]=10 mol/L
[D]=6 mol/L
K = $\frac {[C][D]}{[A][B]}$
K = $\frac {(10)(6)}{(2) (3)}$
k = $\frac {60}{6}$
K = 10
Now Δ G ∘=− RT ln(K)
Δ G ∘=− 2 x 300 x ln (10)
Δ G ∘=− 2 x 300 x 2.30259
Δ G ∘=−1381.80 cal

Therefore, The correct option is (B): –1381.80 cal