Q.26 For exergonic reactions
- ΔG° (standard free energy change) is negative and
Keq (equilibrium constant) is more than one - ΔG° is positive and
Keq is more than one - ΔG° is negative and
Keq is less than one - ΔG° is positive and
Keq is less than oneExergonic reactions have a negative standard free energy change (ΔG°), making them spontaneous, and their equilibrium constant (Keq) exceeds 1, favoring products. The correct option is the first one: ΔG° is negative and Keq is more than one.
Introduction
Exergonic reactions ΔG° negative Keq play a key role in biochemistry, driving spontaneous processes like ATP hydrolysis. This guide explains why ΔG° is negative and Keq exceeds 1 for exergonic reactions, analyzes all MCQ options, and links to the equation ΔG° = -RT ln Keq.
Core Relationship
Exergonic reactions release free energy, so their standard free energy change (ΔG°) under standard conditions (1 M concentrations, 25°C, 1 atm) is negative. This relates to the equilibrium constant (Keq) via ΔG° = -RT ln Keq, where R is the gas constant (8.314 J/mol·K) and T is temperature in Kelvin.
A negative ΔG° means ln Keq > 0, so Keq > 1, indicating products dominate at equilibrium. Positive ΔG° yields Keq < 1, favoring reactants (endergonic).
Option Analysis
Option ΔG° Sign Keq Value Correct for Exergonic? Explanation 1 Negative >1 Yes Matches exergonic: spontaneous, product-favored. E.g., glucose + Pi → glucose-6-P has ΔG° ≈ -17 kJ/mol, Keq >>1. 2 Positive >1 No Positive ΔG° means endergonic (non-spontaneous); Keq >1 impossible per equation. 3 Negative <1 No Negative ΔG° requires Keq >1; Keq <1 implies positive ΔG°. 4 Positive <1 No Describes endergonic reactions, not exergonic. Biochemical Context
In cells (pH 7), use ΔG°’ (biochemical standard state, [H+] = 10^{-7} M). Relationship holds: exergonic like glycolysis steps have ΔG°’ <0, Keq >1. Rate depends on activation energy, not just ΔG°.
