8.3 Galvanic Cell
As discussed earlier, redox reactions cause a movement of charge (e.g. electrons). Consider the following reaction between copper and silver.
\[\mathrm{2Ag^+}(aq) + \mathrm{Cu}(s) \longrightarrow \mathrm{Cu^{2+}}(aq) + \mathrm{2Ag}(s)\]
Oxidation half-reaction
\[\mathrm{Cu}(s) \longrightarrow \mathrm{Cu^{2+}}(aq) + 2e^-\]
Reduction half-reaction
\[\mathrm{2Ag^+}(aq) + 2e^- \longrightarrow \mathrm{2Ag}(s)\]
Two moles of electrons are being transferred in the reaction. Two moles of silver gain two moles of electrons. One mole of copper produces two moles of electrons.
This redox reaction could be performed by dipping a copper wire into an aqueous silver nitrate solution (shown below).
](files/ch17/silver-copper-redox.jpg)
Figure 8.1: A silver/copper redox reaction. Source
Electrons flow from the copper wire directly into the silver nitrate solution. A force drives this flow of electrons, offering tremendous utility for electronic devices which rely on this flow. Perhaps we could separate the two half-reactions and channel the electron flow through such a device (through the wires, circuit boards, transistors, capacitors etc.). This is the main function of a battery!
A galvanic cell is an electrochemical cell that uses spontaneous redox reactions to produce electrical energy. The redox reactions are separated and the electrical energy is travels from one electrode to the other. The anode half-cell is the negative electrode whereas the cathode half-cell is the positive electrode. Additionally, an electrolyte is present to charge-balance the redox system. The electrolyte travels between the half-cells via a salt bridge, a component that selectively allows certain particles to pass through (via a semi-permeable membrane, for example).
](files/ch17/galvanic-cell-anatomy.png)
Figure 8.2: A galvanic cell. Adapted from openStax
Visualize a galvanic cell.
8.3.1 Cell Notation
A galvanic cell (or redox system) can be written in cell notation, a shorthand way to communicate important aspects of a galvanic cell. The following cell notation corresponds to the galvanic cell in Figure 8.2.
\[\color{black}{\mathrm{Cu}(s) ~|~ \mathrm{Cu^{2+}}(aq)} ~||~ \color{red}{\mathrm{Ag^+}(aq) ~|~ \mathrm{Ag}(s)}\]
The anode is located on the left (in black) and corresponds to the negative lead on a battery. The cathode is located on the right (in red) and corresponds to the positive lead on a battery. The solid, vertical lines represent a phase boundary and the double vertical lines in the middle represent the permeable membrane accessible by all ions separating the half-cells.
The spectator ions (NO3–) are generally omitted but can be included as shown below.
\[\color{black}{\mathrm{Cu}(s) ~|~ \mathrm{Cu(NO_3)_2}(aq)} ~||~ \color{red}{\mathrm{AgNO_3}(aq) ~|~ \mathrm{Ag}(s)}\] Cell notation can also include the concentrations of the solutions in the half-cells.
\[\color{black}{\mathrm{Cu}(s) ~|~ \mathrm{Cu^{2+}}(aq, 0.15~M)} ~||~ \color{red}{\mathrm{Ag^+}(aq, 0.5~M) ~|~ \mathrm{Ag}(s)}\]
Additionally, cell notations can include inert electrodes (leads that are not chemically involved in the redox reaction). Below is a redox system that uses solid platinum leads to shuttle electrons from iron to chlorine.
\[\color{black}{ \mathrm{Pt}(s) ~ | ~ \mathrm{Fe^{2+}}(aq, 0.5~M),~\mathrm{Fe^{3+}}(aq, 0.5~M) } ~||~ \color{red}{\mathrm{Cl_2}(g) ~ | ~ \mathrm{Cl^{-}}(aq, 0.50~M) ~|~ \mathrm{Pt}(s)} \]