Chemistry > Electrochemistry > 15.0 Thermodynamics of the Cells
Electrochemistry
1.0 Introduction
2.0 Conductors and Non-Conductors
3.0 Electrochemical Cells
4.0 Electrolysis and electrode Reactions
5.0 Electrochemical Cell
6.0 Electrode Potential
7.0 Nature of Electrodes
8.0 IUPAC Cell Representation and Convention
9.0 Standard Cell EMF and Standard Reduction Potential
10.0 Electropositive Character of Metals
11.0 Difference between EMF and potential difference
12.0 Nernst Equation
13.0 Laws of Electrolysis
14.0 Electromotive Force
15.0 Thermodynamics of the Cells
16.0 Concentration Cells
17.0 Battery
18.0 Fuel Cell
15.2 Spontaneity of the Reaction
$(i)$ When $\Delta {G_{cell}}^0 < 0$ so ${E_{cell}}^0 > 0$, reaction is spontaneous.
$(ii)$ When $\Delta {G_{cell}}^0 > 0$ so ${E_{cell}}^0 < 0$, reaction is non-spontaneous.
$(iii)$ When $\Delta {G_{cell}}^0 = 0$ so ${E_{cell}}^0 =0$, reaction is in equilibrium.
| Type of reaction | $\Delta {G^0}$ | ${E_{cell}}^0$ | Type of cell |
| Spontaneous | -ve | +ve | Galvanic |
| Non-spontaneous | +ve | -ve | Electrolytic |
| Equilibrium | $0$ | $0$ | Dead Battery |
Consider these two
$$\begin{equation} \begin{aligned} {E_{cell}}^0 = 2.303\frac{{RT}}{{nF}}{\log _{10}}{K_{eq}}...(i) \\ \Delta {G^0} = - nF{E_{cell}}^0...(ii) \\\end{aligned} \end{equation} $$
We can write as $$\Delta {G^0} = 2.303RT\ {\log _{10}}{K_{eq}}$$
