Which statement is true for a spontaneous reaction in terms of ∆G?

A spontaneous reaction is defined as one that occurs without needing to be driven by an external force. In thermodynamic terms, the spontaneity of a reaction is determined by the change in Gibbs free energy (∆G). For a reaction to be spontaneous, the Gibbs free energy must decrease, which is represented by a negative value of ∆G.

When ∆G is negative, this indicates that the products of the reaction have lower free energy than the reactants, which drives the reaction forward. This concept is essential in understanding not only chemical reactions but also biochemical processes and systems at equilibrium.

The other statements do not accurately describe conditions for spontaneity. A positive ∆G would indicate that the reaction is non-spontaneous, requiring external energy input to proceed. The notion that ∆G is independent of temperature is incorrect, as temperature can significantly impact the free energy of a system and thus influence the spontaneity of a reaction. Lastly, if ∆G were always zero, it would indicate a system in equilibrium, which is a state where no net change occurs, and therefore does not imply spontaneity.