Thermodynamics check: If the current reactant concentration exceeds the equilibrium condition, what happens to Gibbs free energy and reaction direction?

Introduction / Context:Chemical and biochemical reactions are driven by the Gibbs free energy change ΔG, which depends on both intrinsic standard free energy (ΔG°) and the instantaneous mass-action ratio (reaction quotient, Q). Understanding how concentrations relative to equilibrium determine spontaneity is central to metabolism.

Given Data / Assumptions:

• Reactant concentration is higher than the equilibrium composition (i.e., there is an excess of reactants relative to products).
• Temperature is constant; ΔG = ΔG° + R*T*ln(Q).
• Keq = exp(−ΔG°/(R*T)).

Concept / Approach:When there are more reactants than at equilibrium, Q = [products]/[reactants] is smaller than Keq. Because ΔG can also be written as R*T*ln(Q/Keq), if Q < Keq then ln(Q/Keq) is negative, making ΔG negative. A negative ΔG indicates the forward reaction is spontaneous and will convert reactants into products until Q approaches Keq.

Step-by-Step Solution:Define Q = [products]/[reactants]. Excess reactants → Q decreases.Compute ΔG = R*T*ln(Q/Keq). With Q < Keq → ln term is negative.Therefore ΔG < 0 (thermodynamically favorable forward direction).As the system evolves, product concentration rises and ΔG approaches 0 at equilibrium.

Verification / Alternative check:Le Chatelier’s principle predicts that adding reactants drives the reaction toward products to re-establish equilibrium, consistent with ΔG becoming negative.

Why Other Options Are Wrong:

• Positive ΔG: Would require Q > Keq (excess products), not the case here.
• Reaction must stop: Q = Keq only at equilibrium; the prompt states a deviation from equilibrium.

Common Pitfalls:Confusing ΔG° (a constant at given conditions) with ΔG (state-dependent). In cells, ΔG often differs markedly from ΔG° due to metabolite concentrations.

Final Answer:Both (b) and (c).