Final answer:
Prevailing conditions, such as changes in temperature and concentration, can make the Gibbs free energy (ΔG) for a reaction negative, implying spontaneity, even when the standard free energy change (ΔG°') is positive. This can be achieved, for instance, by increasing temperature to make an endothermic reaction with positive entropy change spontaneous. These reactions are called exergonic and may occur slowly over time.
Step-by-step explanation:
The question concerns the concept of Gibbs free energy (ΔG) in the context of chemical reactions, which is a key topic in thermodynamics. The statement suggests that prevailing conditions can make the change in Gibbs free energy negative (ΔG < 0), indicating spontaneity, even if the standard free energy change (ΔG°') is positive, as might be suggested by non-favorable thermodynamics under standard conditions. To make ΔG negative, and thus the reaction spontaneous, factors such as concentration, temperature, or coupling to a spontaneous reaction can be altered, effectively 'tipping the balance' in favor of product formation. For example, increasing the temperature can cause an endothermic reaction with a positive ΔH and a positive ΔS (increase in randomness) to become spontaneous because the TΔS term may become large enough to override the ΔH, resulting in a negative ΔG according to the equation ΔG = ΔH - TΔS.
An exergonic reaction is a chemical reaction that releases energy, featuring a negative change in free energy (ΔG). Important to note is that the term 'spontaneous' in this context does not necessarily mean that the reaction will occur rapidly, but rather that it will proceed without additional energy input over time.
In summary, a negative ΔG indicates a spontaneous reaction in the forward direction, whereas a positive ΔG indicates non-spontaneity or a spontaneous reverse reaction. The reaction quotient (K), electrochemical potential (E), and other factors like temperature can influence the sign and magnitude of ΔG.