Final answer:
The false statement between endergonic and exergonic reactions is that endergonic reactions have a negative ΔG. This is false because endergonic reactions absorb energy and have a positive ΔG, while exergonic reactions release energy and have a negative ΔG.
Step-by-step explanation:
The question asks which statement about endergonic and exergonic reactions is false. Endergonic reactions absorb energy and have a positive ΔG, while exergonic reactions release energy and have a negative ΔG. Therefore, the false statement is 'c) Endergonic reactions have a negative ΔG.'
Endergonic reactions require energy to proceed, which is denoted by a positive ΔG, meaning that the products have more free energy than the reactants. In contrast, exergonic reactions release energy, which is indicated by a negative ΔG, signifying that the products have less free energy than the reactants. These reactions can occur with a net release of energy and are considered 'spontaneous,' although they may not happen immediately.
For clarification, the term 'ΔH' in the Gibbs free energy equation represents enthalpy, which is related to heat content of the reacting system. An exergonic reaction, which releases energy, is more likely to occur than an endergonic reaction since it can proceed without additional energy input and can therefore drive cellular work.