For a spontaneous reaction the $$\Delta G$$, equilibrium constant ($$K$$) and $$E^\circ_{cell}$$ will be respectively

For any chemical process the three quantities are related by the well-known thermodynamic equations:

1. $$\Delta G^\circ = -\,n\,F\,E^\circ_{cell}$$
2. $$\Delta G^\circ = -\,RT \ln K$$

Here $$\Delta G^\circ$$ is the standard Gibbs free-energy change, $$E^\circ_{cell}$$ is the standard cell potential, and $$K$$ is the equilibrium constant.

For a reaction to be spontaneous under standard conditions we must have $$\Delta G^\circ \lt 0$$. Substituting this negative value in the two equations above:

• In equation (1) a negative $$\Delta G^\circ$$ implies $$E^\circ_{cell}$$ must be positive because $$nF$$ (number of electrons × Faraday constant) is always positive.
• In equation (2) a negative $$\Delta G^\circ$$ gives $$\ln K \gt 0$$, which means $$K \gt 1$$.

Thus, for a spontaneous reaction:

$$\Delta G^\circ$$ is negative (−ve), $$K \gt 1$$, and $$E^\circ_{cell}$$ is positive (+ve).

Option A which is: −ve, >1, +ve