In each of the following questions, a statement I is given followed by a corresponding statement II just below it. Based on the given statements, mark the correct answer.

Statement I: For every chemical reaction at equilibrium standard Gibb’s energy of reaction is zero.

Statement II: At constant temperature and pressure, chemical reactions are spontaneous in the direction of decreasing Gibb’s energy.

Statement I is false but statement II is true.

The correct answer is option 4. Statement I is false but statement II is true.

Let us delve into the concepts of Gibbs free energy and equilibrium to explain why statement I is false and statement II is true.

Statement I: For every chemical reaction at equilibrium standard Gibb’s energy of reaction is zero.

This statement is false. To understand why, we need to differentiate between standard Gibbs free energy change (\(\Delta G^\circ\)) and the Gibbs free energy change (\(\Delta G\)).

Standard Gibbs Free Energy Change (\(\Delta G^\circ\)): This refers to the Gibbs free energy change when the reactants and products are in their standard states (usually 1 atm pressure and 1 M concentration for solutions). \(\Delta G^\circ\) is a fixed value for a particular reaction at a given temperature and does not depend on the current concentrations of reactants and products. It helps to predict the direction of a reaction under standard conditions.

Gibbs Free Energy Change (\(\Delta G\)): This refers to the Gibbs free energy change at any point in the reaction and depends on the actual concentrations or partial pressures of the reactants and products. At equilibrium, \(\Delta G = 0\), meaning the system has reached a state where there is no net change in the concentration of reactants and products.

At equilibrium: The reaction quotient \(Q\) equals the equilibrium constant \(K\). 

\(\Delta G = \Delta G^\circ + RT \ln Q\).

At equilibrium, \(Q = K\), so:

\(\Delta G = \Delta G^\circ + RT \ln K = 0\)

Therefore, \(\Delta G^\circ = -RT \ln K\), which is generally not zero unless \(K = 1\), which is rare. Thus, the standard Gibbs free energy change (\(\Delta G^\circ\)) is generally not zero at equilibrium; only the actual Gibbs free energy change (\(\Delta G\)) is zero.

Statement II: At constant temperature and pressure, chemical reactions are spontaneous in the direction of decreasing Gibb’s energy.

This statement is true. The spontaneity of a reaction at constant temperature and pressure is determined by the Gibbs free energy change (\(\Delta G\)):

If \(\Delta G < 0\), the reaction is spontaneous in the forward direction.

If \(\Delta G > 0\), the reaction is non-spontaneous in the forward direction but spontaneous in the reverse direction.

If \(\Delta G = 0\), the system is at equilibrium, and there is no net change.

Explanation:

Spontaneity: For a reaction to proceed spontaneously, it must lower the Gibbs free energy of the system. Thus, the direction in which \(\Delta G\) decreases is the spontaneous direction.

Equilibrium: At equilibrium, the system has minimized its Gibbs free energy, and no further net change occurs, which is why \(\Delta G = 0\).

Conclusion:

Statement I is false because the standard Gibbs free energy change (\(\Delta G^\circ\)) is not necessarily zero at equilibrium. Statement II is true because it correctly describes the condition for spontaneity in terms of Gibbs free energy.

Thus, the correct answer is: 4. Statement I is false but statement II is true.