Which statement is correct? A) If K < 1, ?G° is negative then the reaction is product favored. B) If K > 1, ?G° is positive then the reaction is product favored C) If K > 1, ?G° is negative then the reaction is product favored. D) If K > 1, ?G° is negative then the reaction is reactant favored.
The Correct Answer and Explanation is:
The correct answer is C) If K > 1, ΔG° is negative then the reaction is product favored.
To understand why this is true, we need to look at the relationship between the equilibrium constant K and the standard Gibbs free energy change ΔG°. This relationship is governed by the equation:
ΔG°=−RTlnKΔG° = -RT \ln K
where:
- ΔG° is the standard Gibbs free energy change,
- R is the gas constant (8.314 J/mol·K),
- T is the temperature in Kelvin,
- K is the equilibrium constant.
If K > 1, then ln K is positive. Because the equation has a negative sign in front of RT ln K, this makes ΔG° negative. A negative ΔG° indicates that the reaction is spontaneous under standard conditions and that the formation of products is favored. Therefore, the equilibrium lies toward the products.
Conversely, if K < 1, then ln K is negative, and the value of ΔG° becomes positive. This suggests the reaction is nonspontaneous under standard conditions, so the equilibrium lies toward the reactants.
Options A, B, and D are incorrect for the following reasons:
- A says that K < 1 and ΔG° is negative, which is inconsistent with the equation. If K < 1, ΔG° must be positive.
- B claims K > 1 and ΔG° is positive, which again contradicts the mathematical relationship.
- D suggests that even with K > 1 and ΔG° negative, the reaction is reactant favored, which is backwards.
In summary, when K > 1, more products are present at equilibrium, and a negative ΔG° confirms that the forward reaction is spontaneous and product favored.
