Galvanic Cell: Definition, Applications, Examples, Solutions & Notes

A positive ${\mathrm{E}}_{\text{cell}}^{\circ }$  indicates a spontaneous reaction. For the Daniell cell, ${\mathrm{E}}_{\text{cell}}^{\circ }=1.10$ V confirms spontaneity. The Gibbs free energy change is related by: $\mathrm{\Delta }{G}^{\circ }=-nF{E}_{\text{cell}}^{\circ }$

where $n$  is the number of electrons transferred, and $F=96485\mathrm{C}/\mathrm{m}\mathrm{o}\mathrm{l}$ . A negative $\mathrm{\Delta }{G}^{\circ }$ confirms spontaneity.

The Nernst equation adjusts ${\mathrm{E}}_{\text{cell}}$  for non-standard conditions: $E_{\text{cell}}=E_{\text{cell}}^{\circ }-\frac{0.0591}{n}\mathrm{l}\mathrm{o}\mathrm{g}Q$

where $Q$  is the reaction quotient. For $\mathrm{Z}\mathrm{n}\left(\mathrm{s}\right)+{\mathrm{C}\mathrm{u}}^{2+}\left(\mathrm{a}\mathrm{q}\right)\to {\mathrm{Z}\mathrm{n}}^{2+}\left(\mathrm{a}\mathrm{q}\right)+\mathrm{C}\mathrm{u}\left(\mathrm{s}\right),Q=$ $\frac{\left[{\mathrm{Z}\mathrm{n}}^{2+}\right]}{\left[{\mathrm{C}\mathrm{u}}^{2+}\right]}$ .

Concentration cells are galvanic cells where both electrodes are identical but immersed in electrolytes of different concentrations. Example: $\mathrm{Z}\mathrm{n}\left|{\mathrm{Z}\mathrm{n}}^{2+}\left(0.1\mathrm{M}\right)\right|\mid {\mathrm{Z}\mathrm{n}}^{2+}(1$ $\mathrm{M})\mid \mathrm{Z}\mathrm{n}$  . The ${\mathrm{E}}_{\text{cell}}$ is: $E_{\text{cell}}=\frac{0.0591}{n}\mathrm{l}\mathrm{o}\mathrm{g}\frac{C_2}{C_1}\mathrm{}\left(C_2{C}_1\right)$

Electrons flow from the lower concentration (anode) to the higher concentration (cathode).

Galvanic cells have wide applications:

The major applications of the galvanic cell are mentioned below. 

1. Batteries: Galvanic cells are the basis of batteries. They are used in rechargeable batteries.

2. Electrochemical Sensors: Used in devices where electrical signals are produced by chemical reactions, such as glucose meters and pH meters.

3. Corrosion Prevention: To prevent rusting in pipelines and ships, cathodic protection is done. This is done with the help of galvanic cells.
4. Fuel Cells: These cells produce electricity by reacting hydrogen with oxygen. It is used in spacecraft.

Galvanic Cell is an electrochemical cell where chemical energy is converted to electrical energy due to spontaneous redox reaction. While in electrolytic cells, electrical energy cells are converted to chemical energy. Below are the differences between galvanic cells and electrolytic cells. 

Problem: Calculate the ${\mathrm{E}}_{\text{cell}}$  for the cell $\mathrm{Z}\mathrm{n}\left(\mathrm{s}\right)\left|{\mathrm{Z}\mathrm{n}}^{2+}\left(0.01\mathrm{M}\right)\right|\left|{\mathrm{A}\mathrm{g}}^{+}\left(0.1\mathrm{M}\right)\right|\mathrm{A}\mathrm{g}\left(\mathrm{s}\right)$  at ${25}^{\circ }\mathrm{C}$ . Given ${\mathrm{E}}_{{\mathrm{Z}\mathrm{n}}^{2+}/\mathrm{Z}\mathrm{n}}^{\circ }=-0.76\mathrm{V},{\mathrm{E}}_{{\mathrm{A}\mathrm{g}}^{+}/\mathrm{A}\mathrm{g}}^{\circ }=+0.80\mathrm{V}$ .
Solution: $E_{\text{cell}}^{\circ }=E_{\text{cathode}}^{\circ }-E_{\text{anode}}^{\circ }=0.80-(-0.76)=1.56\mathrm{V}$

Cell reaction: $\mathrm{Z}\mathrm{n}\left(\mathrm{s}\right)+2{\mathrm{A}\mathrm{g}}^{+}\left(\mathrm{a}\mathrm{q}\right)\to {\mathrm{Z}\mathrm{n}}^{2+}\left(\mathrm{a}\mathrm{q}\right)+2\mathrm{A}\mathrm{g}\left(\mathrm{s}\right),n=2$ .

$\begin{array}{c}Q=\frac{\left[{\mathrm{Z}\mathrm{n}}^{2+}\right]}{{\left[{\mathrm{A}\mathrm{g}}^{+}\right]}^2}=\frac{0.01}{(0.1{)}^2}=1\\ E_{\text{cell}}=1.56-\frac{0.0591}{2}\mathrm{l}\mathrm{o}\mathrm{g}1=1.56\mathrm{V}\end{array}$

This tests the Nernst equation and cell potential, common in JEE Main.