The concept of separating a redox reaction into separate reduction and oxidation parts is not purely a theoretical mechanism to help in this paper-and-pencil task. In real life, one can often actually carry out the two half-reactions in separate compartments or beakers, and couple the two so that the electrons are forced to flow through an external circuit. Such a configuration occurs in the galvanic cell, one class of electrochemical cells.
An electrochemical cell is a contained system in which a redox reaction occurs in conjunction with the passage of electric current. There are two types of electrochemical cells, galvanic cells (also known as voltaic cells) and electrolytic cells. (We shall investigate the differences between the two types shortly.) Both kinds of electrochemical cells contain two electrodes, which are essentially two pieces of metal, that serve as the sites for the oxidation and reduction half-reactions separately. The electrode at which oxidation occurs is called the anode, and the electrode at which reduction occurs is called the cathode. This is true for both galvanic (voltaic) and electrolytic cells.
A redox reaction occurring in a galvanic cell has a negative ∆G and is therefore a spontaneous reaction. Galvanic cell reactions supply energy and are used to do work. This energy can be harnessed by placing the oxidation and reduction half-reactions in separate containers called half-cells. The half-cells are then connected by an apparatus that allows for the flow of electrons. The spontaneous flow of electrons is forced to go through external circuitry in which their potential energy is extracted.
A common example of a galvanic cell is the Daniell cell, shown below:
In the Daniell cell, a zinc bar is placed in an aqueous ZnSO4 solution, and a copper bar is placed in an aqueous CuSO4 solution. The anode of this cell is the zinc bar where Zn (s) is oxidized to Zn2+ (aq). The cathode is the copper bar, and it is the site of the reduction of Cu2+ (aq) to Cu (s). The half-cell reactions are written as follows:
| Zn (s) → Zn2+ (aq) + 2e− | (anode) |
| Cu2+ (aq) + 2e− → Cu (s) | (cathode) |
If the two half-cells were not separated, the Cu2+ ions would react directly with the zinc bar and no useful electrical work would be obtained. To complete the circuit, the two solutions must be connected. Without connection, the electrons from the zinc oxidation half-reaction would not be able to get to the copper ions, so a wire (or other conductor) is necessary. If only a wire were provided for this electron flow, the reaction would soon cease anyway because an excess negative charge would build up in the solution surrounding the cathode and an excess positive charge would build up in the solution surrounding the anode. This charge gradient is dissipated by the presence of a salt bridge, which permits the exchange of cations and anions. The salt bridge contains an inert electrolyte, usually KCl or NH4NO3, whose ions will not react with the electrodes or with the ions in solution. At the same time the anions from the salt bridge (such as Cl−) diffuse from the salt bridge of the Daniell cell into the ZnSO4 solution to balance out the charge of the newly created Zn2+ ions, the cations of the salt bridge (such as K+) flow into the CuSO4 solution to balance out the charge of the SO42− ions left in solution when the Cu2+ ions deposit as copper metal.
During the course of the reaction, electrons flow from the zinc bar (anode) through the wire and the ammeter, toward the copper bar (cathode). The anions (Cl−) flow externally (via the salt bridge) into the ZnSO4, and the cations (K+) flow into the CuSO4. This flow depletes the salt bridge and, along with the finite quantity of Cu2+ in the solution, accounts for the relatively short lifetime of the cell.
Instead of an ammeter that simply measures the current, one can place a device that is powered by electric current so as to extract the potential energy of the electrons. That is, after all, why galvanic cells are useful. The common dry cell battery and the lead-acid storage battery found in cars are examples of galvanic cells.
A redox reaction occurring in an electrolytic cell has a positive ∆G and is therefore nonspontaneous. In electrolysis, electrical energy is required to induce reaction; i.e., instead of extracting work from a spontaneous redox reaction, we supply energy to force a nonspontaneous redox reaction to occur. The oxidation and reduction half-reactions are usually placed in one container. Where the ammeter or electrical device used to be for the galvanic cell, we need to place a source of electrical power, like a battery, instead (see figure on the next page).
Michael Faraday was the first to define certain quantitative principles governing the behavior of electrolytic cells. He theorized that the amount of chemical change induced in an electrolytic cell is directly proportional to the number of moles of electrons that are exchanged during a redox reaction. The number of moles exchanged can be determined from the balanced half-reaction. In general, for a reaction that involves the transfer of n electrons per atom:
Mn+ + ne− → M(s)
one mole of M(s) will be produced if n moles of electrons are supplied.
The number of moles of electrons needed to produce a certain amount of M(s) can now be related to a measurable electrical property. One electron carries a charge of 1.6 × 10−19 coulombs (C). The charge carried by one mole of electrons can be calculated by multiplying this number by Avogadro’s number, as follows:
(1.6 × 10−19)(6.022 × 1023) = 96,487 C/mol e−
This number is called Faraday’s constant, and one Faraday (F) is equivalent to the amount of charge contained in one mole of electrons (1 F = 96,487 coulombs, or J/V).
An example of an electrolytic cell, in which molten NaCl is electrolyzed to form Cl2(g) and Na(l), is given below:
In this cell, Na+ ions migrate toward the cathode, where they are reduced to Na (l). Similarly, Cl− ions migrate toward the anode, where they are oxidized to Cl2 (g).
The anode of an electrolytic cell is considered positive, since it is attached to the positive pole of the battery and so attracts anions from the solution. The anode of a galvanic cell, on the other hand, is considered negative because the spontaneous oxidation reaction that takes place at the galvanic cell’s anode is the original source of that cell’s negative charge, i.e., is the source of electrons. In spite of this difference in designating charge, oxidation takes place at the anode in both types of cells, and electrons always flow through the wire from the anode to the cathode.