Thermodynamics for biochemistry — If the Gibbs free energy change (ΔG) for a reaction is negative, what does this indicate about the reaction?

Introduction / Context:
Gibbs free energy (ΔG) predicts spontaneity and energy flow in biochemical reactions. Cells couple exergonic reactions (negative ΔG) to endergonic processes to drive metabolism forward, often via ATP hydrolysis. Understanding the sign of ΔG is fundamental for interpreting catabolic and anabolic pathways.

Given Data / Assumptions:

• ΔG < 0 denotes an exergonic reaction under specified conditions.
• Exergonic reactions can proceed spontaneously (thermodynamically favorable) though rate depends on activation energy.
• We are not considering kinetics; only thermodynamics.

Concept / Approach:
Negative ΔG means the products have lower free energy than reactants, so energy is released to the surroundings (often as heat or captured in high-energy intermediates). Positive ΔG indicates an endergonic reaction that requires energy input. Phrases like “negative/positive direction” are imprecise in thermodynamics and do not define spontaneity or energy flow.

Step-by-Step Solution:

Verification / Alternative check:
Common cellular examples include oxidation of NADH (negative ΔG) and ATP hydrolysis under cellular conditions; both release usable free energy to drive work.

Why Other Options Are Wrong:

• Absorbs energy: describes endergonic (ΔG > 0), not exergonic reactions.
• Negative/positive direction: nonstandard phrasing; direction is defined by equilibrium and ΔG, not sign labels.

Common Pitfalls:
Equating negative ΔG with rapid reaction; catalysts alter rate (activation energy), not ΔG.

Final Answer:
reaction releases energy