Question

In the following hypothetical reaction A + B → C + D, the equilibrium constant, Keq is less than 1.0 at
25°C and decreases by 35% on changing the temperature to 45°C. What must be true according to this information?

A. The ΔH° for the reaction is negative.
B. The ΔS° for the reaction is positive.
C. As temperature increases, the ΔG° for the reaction does not change.
D. The ΔG° for the reaction at 25°C is negative.
E. The ΔG° for the reaction at 45°C is zero.

Answer 1

The correct answer is A, which I got it right, but I was torn between A & B. could you please explain why B is incorrect. Thank you.

Answer 2

I'm sticking my neck out but go through this reasoning and see if you agree.
reactants --> products
k = P/R
For k to be < 1, reactants must be favored. Just try A; if dH is - the reaction would be R ==> P + heat
If we increase T the reaction will be shifted to the left, R increases and P decreases so k decreases and that agrees with the problem.

What about B? If dS is +.
dG = dH - TdS
If dS is + then increases T will make the -TdS term more negative and that added to a constant H (whatever it is) will make dG more negative and that means products are favored so I think no for B.

C? C can't possibly be right because if T changes dG must change since -TdS changes.

D? I don' think so. If dG were - then k would be greater than 1.

E. If dG = 0 it means k = 1; however, if that is true then increase in T would make products favored but the problem says k decreases, not increases from < 1 to 1

so A?

Answer 3

Yes, A! Thank you for making it easy to understand.