Which of the following is not correct?

We need to identify which statement about Gibbs free energy ($$\Delta G$$) is NOT correct.

Recall the significance of $$\Delta G$$.

The Gibbs free energy change ($$\Delta G$$) is the criterion for spontaneity of a process at constant temperature and pressure:

$$\Delta G = \Delta H - T\Delta S$$

Evaluate each statement.

Option (1): "$$\Delta G$$ is negative for a spontaneous reaction."

This is CORRECT. A negative $$\Delta G$$ means the reaction decreases the free energy of the system, which is the thermodynamic condition for spontaneity.

Option (2): "$$\Delta G$$ is positive for a spontaneous reaction."

This is INCORRECT. A positive $$\Delta G$$ means the reaction is non-spontaneous - it requires external energy input to proceed. This directly contradicts the fundamental criterion for spontaneity.

Option (3): "$$\Delta G$$ is zero for a reversible reaction."

This is CORRECT. At equilibrium (which is the condition for a reversible process), $$\Delta G = 0$$. The system has reached the minimum of free energy and has no driving force in either direction.

Option (4): "$$\Delta G$$ is positive for a non-spontaneous reaction."

This is CORRECT. A positive $$\Delta G$$ indicates that the reaction requires energy input and will not proceed on its own.

The incorrect statement is Option (2): $$\Delta G$$ is positive for a spontaneous reaction.