AOT - Voltaic Cells (Lesson)

Voltaic Cells

In a previous module, types of chemical reactions were discussed including redox reactions. The redox reactions studied there were comprised mainly of single replacement reactions where a metal was placed in contact with an aqueous solution of a different metal ion. An example of this would be the reaction of metallic copper with a solution of silver nitrate as illustrated below.

The copper metal transfers electrons to the silver ions in solution which results in the solution turning blue due to the production of Cu ions and silver metal depositing onto the metallic copper. In this example, the two materials are in direct contact with one another, and the energy produced from the reaction is given off in the form of heat which is rarely, if ever, useful. However, if the two substances are separated in something called a voltaic cell, then this energy can be utilized to perform electrical work.

Voltaic Cells

A voltaic cell is a device that harnesses the chemical energy resulting from the transfer of electrons in a redox reaction; then, the energy is used to power small devices. A well known and common example of this is alkaline batteries that are used practically every day in a wide variety of applications. Instead of placing the two substances in direct contact with one another, a voltaic cell separates the two half-reactions in an electrical circuit. By doing this, the electrons lost by the substance being oxidized must flow through the circuit via a wire to reach the other substance being reduced. This flow of electrons can thereby be used for some useful purposes such as powering small devices like remote controls, hearing aids, calculators, and much more. A diagram of a voltaic cell from the reaction previously mentioned is shown below.

The first thing that must be noticed from this diagram is that the two substances involved in the chemical reaction have been separated into two electrodes.  These electrodes consist of the two half reactions of the chemical reaction. For example, the overall chemical reaction for the process depicted is:

Cu (s) + 2 Ag⁺ (aq)  $\longrightarrow$  Cu⁺² (aq) + 2 Ag (s)

The overall reaction can then be broken into two half-reactions:  an oxidation half-reaction, referred to as the anode, and a reduction half-reaction, referred to as the cathode.

• Oxidation half-reaction (anode):  Cu(s)  $\longrightarrow$  Cu⁺² (aq) + 2 e^-
• Reduction half-reaction (cathode):  2 Ag⁺ (aq) + 2 e^-   $\longrightarrow$  2 Ag (s) 

As the reaction proceeds, copper metal loses electrons which are transferred through the circuit to the silver solution which results in the addition of silver to the silver electrode as indicated on the diagram.

In order for this apparatus to function as a voltaic cell, a final component is necessary. One criterion for any electrical circuit is that it must be closed for electrical current to be conducted. The final component that accomplishes this is referred to in the diagram as a salt bridge. The salt bridge serves to complete this circuit, but, in addition to this, it also serves another important role. 

As this voltaic cell operates, the copper metal is losing electrons. As these electrons flow from the anode to the cathode, a buildup of positive charge would be expected. In a similar fashion, as the negatively charged electrons flow toward the cathode a buildup of negative charge would be observed. The longer this reaction proceeds, the more difficult it would be for electrons to flow, as they would be forced to become separated from an ever growing positive charge only to be sent toward an ever increasing negative charge. Essentially, the reaction would become less and less favorable.

To prevent this from happening a voltaic cell also contains a component called an electrolyte that flows through the salt bridge. The electrolyte in the example shown above is NaNO₃. Notice that this substance does not appear in the overall chemical reaction and instead serves to prevent the build up of charges associated with each electrode. As the reaction proceeds, the positively charged sodium ions (Na⁺) migrate toward the cathode to balance out the increasing negative charge accompanied by this half-reaction. Similarly, nitrate ions (NO₃^-) migrate across the salt bridge to balance out the positive charges associated with electrons leaving that half-cell.  

Other Types of Electrodes

In the case of the reaction above, it is quite easy to visualize how a physical electrode can be constructed out of metallic copper and silver.  However, what about cases involving substances that are not solids under standard conditions (e.g. Cl₂) or half-reactions where both the reactant and product are ions (e.g. Sn⁺⁴ (aq) + 2 e^-   $\longrightarrow$  Sn⁺² (aq))?  In these instances, it is still possible to construct voltaic cells with these substances; however, a surface is still required in order for the redox reaction to occur. In other words, as the electrons are being transferred from the anode to the cathode, the electrons are physically transferred to the ions in the solution on the surface of a physical electrode.

In these cases, an inert electrode is utilized for this role. Take, for example, the depiction of the voltaic cell shown below. The reaction described in this diagram is:

Mg (s) + 2 Fe⁺³ (aq)   $\longrightarrow$  Mg⁺² (aq) + 2 Fe⁺² (aq)

In this voltaic cell, the Mg serves as the anode, and a piece of metallic Mg serves as the electrode. However, the cathode for this reaction is Fe⁺³.  While this ion can certainly exist in aqueous solution, it is not possible to construct a physical electrode from it. The inert electrode utilized in this case is the piece of Pt metal. The platinum does not take part in the reaction and is not consumed as the voltaic cell operates. Its presence simply serves as the physical surface required for this process to take place.

You Try It!

In the following self-assessment activity, use the diagram to help you answer the questions. Click on the plus sign to check your answers!

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