Electrochemical cells are contained systems in which oxidation–reduction reactions occur. There are three fundamental types of electrochemical cells: galvanic cells (also known as voltaic cells), electrolytic cells, and concentration cells. In addition, there are specific commercial cells such as Ni–Cd batteries. Galvanic cells and concentration cells house spontaneous reactions, whereas electrolytic cells contain nonspontaneous reactions. Remember that spontaneity is indicated by the change in Gibbs free energy, ΔG. All three types contain electrodes where oxidation and reduction take place. For all electrochemical cells, the electrode where oxidation occurs is called the anode, and the electrode where reduction occurs is called the cathode. Other descriptors of electrochemical cells include the electromotive force (emf), which corresponds to the voltage or electrical potential difference of the cell. If the emf is positive, the cell is able to release energy (ΔG < 0), which means it is spontaneous. If the emf is negative, the cell must absorb energy (ΔG > 0), which means it is nonspontaneous. For all electrochemical cells, the movement of electrons is from anode to cathode, and the current (I) runs from cathode to anode. This can be a point of confusion as in physics, it is typical to state that current is the direction of flow of a positive charge through a circuit. Modern chemists are interested in the flow of electrons, but may discuss the current (a theoretical flow of positive charge) as a proxy for the flow of electrons; the current and the flow of electrons are always of equal magnitude but in opposite directions.
Faraday’s laws state that the liberation of gas, and deposition of elements, on electrodes is directly proportional to the number of electrons being transferred during the oxidation–reduction reaction. Here, normality or gram equivalent weight is used. These observations are proxy measurements of the amount of current flowing in a circuit. According to this law, one mole of metal M (s) will logically be produced if n moles of electrons are supplied to one mole of Mn+. Additionally, the number of moles of electrons needed to produce a certain amount of M (s) can now be related to the measurable electrical property of charge. 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:
The Faraday constant is the name given to this value, and one faraday (F) is the same as the amount of charge contained in one mole of electrons (1 F = 96,485 C).
The non-rechargeable batteries in households products, such as flashlights, are galvanic cells, also called voltaic cells. Accordingly, because household batteries are used to supply energy to a flashlight or remote control, the reactions in these cells must be spontaneous. This means that the reaction’s free energy is decreasing (ΔG < 0) as the cell releases energy to the environment. By extension, if the free energy change is negative for these cells, their electromotive force (Ecell) must be positive; the free energy change and electromotive force always have opposite signs. In the galvanic (voltaic) cell, two electrodes of distinct chemical identity are placed in separate compartments, which are called half-cells. The two electrodes are connected to each other by a conductive material, such as a copper wire. Along the wire, there may be other various components of a circuit, such as resistors or capacitors, but for now, we’ll focus on the battery itself. Surrounding each of the electrodes is an aqueous electrolyte solution composed of cations and anions. As shown in the Galvanic cell (Daniell cell) below, the cations in the two half-cell solutions can be of the same element as the respective metal electrode. Connecting the two solutions is a structure called a salt bridge, which consists of an inert salt.4 When the electrodes are connected to each other by a conductive material, charge will begin to flow as the result of an oxidation–reduction reaction that is taking place between the two half-cells. The redox reaction in a galvanic cell is spontaneous, and therefore the change in Gibbs free energy for the reaction is negative (ΔG < 0). As the spontaneous reaction proceeds toward equilibrium, the movement of electrons results in a conversion of electrical potential energy into kinetic energy. By separating the reduction and oxidation half-reactions into two compartments, we are able to harness this energy and use it to do work by connecting various electrical devices into the circuit between the two electrodes.
In the Galvanic cell, a zinc electrode is placed in an aqueous ZnSO4 solution, and a copper electrode 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 net reaction is
Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)
However, if only a wire were provided for this electron flow, the reaction would soon stop because an excess positive charge would build up on the anode, and an excess negative charge would build up on the cathode.4 Eventually, the excessive charge accumulation would provide a countervoltage large enough to prevent the oxidation–reduction reaction from taking place, and the current would cease. 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, which contains ions that will not react with the electrodes or with the ions in solution. While the anions from the salt bridge (Cl–) diffuse into the solution on the anode side (ZnSO4) to balance out the charge of the newly created Zn2+ ions, the cations of the salt bridge (K+) flow into the solution on the cathode side (CuSO4) to balance out the charge of the sulfate ions left in solution when the Cu2+ ions are reduced to Cu and precipitate onto the electrode. This precipitation process onto the cathode itself can also be called plating or galvanization.
A concentration cell is a special type of galvanic cell. Like all galvanic cells, it contains two half-cells connected by a conductive material, allowing a spontaneous oxidation–reduction reaction to proceed, which generates a current and delivers energy. The distinguishing characteristic of a concentration cell is in its design: the electrodes are chemically identical4 For example, if both electrodes are copper metal, they have the same reduction potential. Therefore, current is generated as a function of a concentration gradient established between the two solutions surrounding the electrodes. The concentration gradient results in a potential difference between the two compartments and drives the movement of electrons in the direction that results in equilibration of the ion gradient. The current will stop when the concentrations of ionic species in the half-cells are equal. This implies that the voltage (V) or electromotive force of a concentration cell is zero when the concentrations are equal; the voltage, as a function of concentrations, can be calculated using the Nernst equation. In a biological system, a concentration cell is best represented by the cell membrane of a neuron. Sodium and potassium cations, and chlorine anions, are exchanged as needed to produce an electrical potential. The actual value depends on both the concentrations and charges of the ions. In this way, a resting membrane potential (Vm) can be maintained. Disturbances of the resting membrane potential, if sufficiently large, may stimulate the firing of an action potential.
A lead–acid battery, also known as a lead storage battery, is a specific type of rechargeable battery. As a voltaic cell, when fully charged, it consists of two half-cells, a Pb anode and a porous PbO2 cathode, connected by a conductive material. When fully discharged, it consists of two PbSO4 electroplated lead electrodes with a dilute concentration of H2SO4. Nickel–cadmium batteries are also rechargeable cells. They consist of two half-cells made of solid cadmium (the anode) and nickel (III) oxide-hydroxide (the cathode) connected by a conductive material, typically potassium hydroxide (KOH). Most of us are familiar with AA and AAA cells made of Ni–Cd materials, inside of which the electrodes are layered and wrapped around in a cylinder.
Cell Membrane Potential
Useful tip: SALTY BANANA
High NaCl on outside of cell and high K+ on inside of cell.
1) Matthew Sadiku (2009). Elements of electromagnetics. p. 104.
2) Farago, P.S. (1961) An Introduction to Linear Network Analysis, pp. 18–21, The English Universities Press Ltd.
3) Keithley, Joseph F (1999). Daniell Cell. John Wiley and Sons. pp. 49–51.
4) Verschuur, Gerrit L. (1993). Hidden Attraction: The History and Mystery of Magnetism. New York: Oxford University Press. p. 76
5) Dorland’s (2012). Dorland’s Illustrated Medical Dictionary (32nd ed.). Elsevier Saunders.
6) Black, J.A., Sontheimer, H., Oh, Y., and Waxman, S.G. (1995). In The Axon, S. Waxman, J. Kocsis, and P. Stys, eds. Oxford University Press, New York, pp. 116–143.