Reaction 1 has a $\Delta$G° of -12.3 kJ/mol, and Reaction 2 has a $\Delta$G° of 23.4 kJ/mol. Which statement is TRUE of these two reactions?
A) Reaction 1 occurs faster.
B) Reaction 2 occurs faster.
C) Both reactions occur at the same rate.
D) Reaction 2 will not occur.
E) It is impossible to know which reaction occurs faster with this information.

Question

Quizplus · Accepted Answer

The given ΔG° values indicate the spontaneity of the reactions, not their rates. A negative ΔG° suggests that Reaction 1 is spontaneous, while a positive ΔG° indicates that Reaction 2 is non-spontaneous under standard conditions. However, the rate of a reaction depends on its kinetics, which cannot be determined from ΔG° alone.

Reaction 1 Has A Δ\DeltaΔ G° of -123 KJ/mol, and Reaction 2 Has A -12.3

Reaction 1 Has A $\Delta$ G° of -123 KJ/mol, and Reaction 2 Has A -12.3