4.7: Electrochemical Cells

Consider what happens when a clean piece of copper metal is placed in a solution of silver nitrate (Figure \(\PageIndex{1}\)). As soon as the copper metal is added, silver metal begins to form and copper ions pass into the solution. The blue color of the solution on the far right indicates the presence of copper ions. The reaction may be split into its two half-reactions. Half-reactions separate the oxidation from the reduction, so each can be considered individually.

\[\begin{align}
&\textrm{oxidation: }\ce{Cu}(s)⟶\ce{Cu^2+}(aq)+\ce{2e-}\\
&\underline{\textrm{reduction: }2×(\ce{Ag+}(aq)+\ce{e-}⟶\ce{Ag}(s))\hspace{40px}\ce{or}\hspace{40px}\ce{2Ag+}(aq)+\ce{2e-}⟶\ce{2Ag}(s)}\\
&\textrm{overall: }\ce{2Ag+}(aq)+\ce{Cu}(s)⟶\ce{2Ag}(s)+\ce{Cu^2+}(aq)
\end{align} \nonumber \]

Figure \(\PageIndex{1}\): When a clean piece of copper metal is placed into a clear solution of silver nitrate (a), an oxidation-reduction reaction occurs that results in the exchange of Cu²⁺ for Ag⁺ ions in solution. As the reaction proceeds (b), the solution turns blue (c) because of the copper ions present, and silver metal is deposited on the copper strip as the silver ions are removed from solution. (credit: modification of work by Mark Ott)

Parts of an Electrochemical Cell

The redox reaction between Cu(s) and Ag⁺ (aq) involves the movement of electrons. The two half-reactions can be physically separated to make an electrochemical cell as shown in Figure \(\PageIndex{2}\). The beaker on the left side of the figure is called a half-cell, and contains a 1 M solution of copper(II) nitrate [Cu(NO₃)₂] with a piece of copper metal partially submerged in the solution. The copper metal is an electrode. The copper is undergoing oxidation; therefore, the copper electrode is the anode. The anode is connected to a voltmeter with a wire and the other terminal of the voltmeter is connected to a silver electrode by a wire. Electrons flow from the anode to the cathode. The silver is undergoing reduction; therefore, the silver electrode is the cathode. The half-cell on the right side of the figure consists of the silver electrode in a 1 M solution of silver nitrate (AgNO₃).

Figure \(\PageIndex{2}\): In this standard galvanic cell, the half-cells are separated; electrons can flow through an external wire and become available to do electrical work.

The circuit is closed using a salt bridge, which transmits the current with moving ions. The salt bridge consists of a concentrated, nonreactive, electrolyte solution such as the sodium nitrate (NaNO₃) solution used in this example. As electrons flow from left to right through the electrode and wire, nitrate ions (anions) pass through the porous plug on the left into the copper(II) nitrate solution. This keeps the beaker on the left electrically neutral by neutralizing the charge on the copper(II) ions that are produced in the solution as the copper metal is oxidized. At the same time, the nitrate ions are moving to the left, sodium ions (cations) move to the right, through the porous plug, and into the silver nitrate solution on the right. These added cations “replace” the silver ions that are removed from the solution as they were reduced to silver metal, keeping the beaker on the right electrically neutral. Without the salt bridge, the compartments would not remain electrically neutral and no significant current would flow. However, if the two compartments are in direct contact, a salt bridge is not necessary.

The voltmeter connected in the electrochemical circuit reads +0.46 V; this is called the cell potential. The cell potential is created when the two dissimilar metals are connected, and is a measure of the energy per unit charge available from the oxidation-reduction reaction. The volt is the derived SI unit for electrical potential

\[\mathrm{volt=\mathit{V}=\dfrac{J}{C}} \nonumber \]

In this equation, A is the current in amperes and C the charge in coulombs. Note that volts must be multiplied by the charge in coulombs (C) to obtain the energy in joules (J).

Some oxidation-reduction reactions involve species that are poor conductors of electricity, and so an electrode is used that does not participate in the reactions. Frequently, the electrode is platinum, gold, or graphite, all of which are inert to many chemical reactions. One such system is shown in Figure \(\PageIndex{3}\). Magnesium undergoes oxidation at the anode on the left in the figure and hydrogen ions undergo reduction at the cathode on the right. The reaction may be summarized as

\[\begin{align}
&\textrm{oxidation: }\ce{Mg}(s)⟶\ce{Mg^2+}(aq)+\ce{2e-}\\
&\textrm{reduction: }\ce{2H+}(aq)+\ce{2e-}⟶\ce{H2}(g)\\
&\overline{\textrm{overall: }\ce{Mg}(s)+\ce{2H+}(aq)⟶\ce{Mg^2+}(aq)+\ce{H2}(g)}
\end{align} \nonumber \]

The magnesium electrode is an active electrode because it participates in the oxidation-reduction reaction. Inert electrodes, like the platinum electrode in Figure \(\PageIndex{3}\), do not participate in the oxidation-reduction reaction and are present so that current can flow through the cell. Platinum or gold generally make good inert electrodes because they are chemically unreactive.

Figure \(\PageIndex{3}\): The oxidation of magnesium to magnesium ion occurs in the beaker on the left side in this apparatus; the reduction of hydrogen ions to hydrogen occurs in the beaker on the right. A nonreactive, or inert, platinum wire allows electrons from the left beaker to move into the right beaker. The overall reaction is: \(\ce{Mg + 2H+ ⟶ Mg^2+ + H2}\), which is represented in cell notation as: \(\ce{Mg}(s)│\ce{Mg^2+}(aq)║\ce{H+}(aq)│\ce{H2}(g)│\ce{Pt}(s)\).

Two Types of Electrochemical Cells

Galvanic/Voltaic Cells 

Galvanic cells, also known as voltaic cells, are electrochemical cells in which the oxidation-reduction reactions occur spontaneously and produce electrical energy. The examples shown in Figures \(\PageIndex{1}\) and \(\PageIndex{2}\) are galvanic cells. Galvanic cells have a positive cell potential. These cells are important because they are the basis for the batteries that fuel modern society.

Electrolytic Cells

It is possible to construct a cell that does work on a chemical system by driving an electric current through the system. These cells are called electrolytic cells. Electrolytic cells, like galvanic cells, are composed of two half-cells--one is a reduction half-cell, the other is an oxidation half-cell. The direction of electron flow in electrolytic cells, however, may be reversed from the direction of spontaneous electron flow in galvanic cells, but the definition of both cathode and anode remain the same, where reduction takes place at the cathode and oxidation occurs at the anode. Because the directions of both half-reactions have been reversed, the sign, but not the magnitude, of the cell potential has been reversed.

Electrolytic cells are very similar to voltaic (galvanic) cells in the sense that both require a salt bridge, both have a cathode and anode side, and both have a consistent flow of electrons from the anode to the cathode. However, there are also striking differences between the two cells. The main differences are outlined below:

Figure \(\PageIndex{1}\): Electrochemical Cells. A galvanic cell (left) transforms the energy released by a spontaneous redox reaction into electrical energy that can be used to perform work. The oxidative and reductive half-reactions usually occur in separate compartments that are connected by an external electrical circuit; in addition, a second connection that allows ions to flow between the compartments (shown here as a vertical dashed line to represent a porous barrier) is necessary to maintain electrical neutrality. The potential difference between the electrodes (voltage) causes electrons to flow from the reductant to the oxidant through the external circuit, generating an electric current. In an electrolytic cell (right), an external source of electrical energy is used to generate a potential difference between the electrodes that forces electrons to flow, driving a nonspontaneous redox reaction; only a single compartment is employed in most applications. In both kinds of electrochemical cells, the anode is the electrode at which the oxidation half-reaction occurs, and the cathode is the electrode at which the reduction half-reaction occurs.

Table \(\PageIndex{1}\): Properties of Galvanic and Electrochemical Cells

Electrochemical cell (Galvanic Cell)	Electrolytic cell
A Galvanic cell converts chemical energy into electrical energy.	An electrolytic cell converts electrical energy into chemical energy.
Here, the redox reaction is spontaneous and is responsible for the production of electrical energy.	The redox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction.
The two half-cells are set up in different containers, being connected through the salt bridge or porous partition.	Both the electrodes are placed in a same container in the solution of molten electrolyte.
Here the anode is negative and cathode is the positive electrode. The reaction at the anode is oxidation and that at the cathode is reduction.	Here, the anode is positive and cathode is the negative electrode. The reaction at the anode is oxidation and that at the cathode is reduction.
The electrons are supplied by the species getting oxidized. They move from anode to the cathode in the external circuit.	The external battery supplies the electrons. They enter through the cathode and come out through the anode.