Electrochemistry deals with the links between chemical reactions and electricity. This includes the study of chemical changes caused by the passage of an electric current across a medium, as well as the production of electric energy by chemical reactions. Electrochemistry also embraces the study of electrolyte solutions and the chemical equilibria that occur in them.
Many chemical reactions require the input of energy. Such reactions can be carried out at the surfaces of electrodes in cells connected to external power supplies. These reactions provide information about the nature and properties of the chemical species contained in the cells, and can also be used to synthesize new chemicals. The production of chlorine and aluminum and the electroplating and electrowinning of metals are examples of industrial electrochemical processes. Electrochemical cells that produce electric energy from chemical energy are the basis of primary and secondary (storage) batteries and fuel cells. Other electrical phenomena of interest in chemical systems include the behavior of ionic solutions and the conduction of current through these solutions, the separation of ions by an electric field (electrophoresis ), the corrosion and passivation of metals, electrical effects in biological systems (bioelectrochemistry), and the effect of light on electrochemical cells (photoelectrochemistry).
An electrochemical cell generally consists of two half-cells, each containing an electrode in contact with an electrolyte. The electrode is an electronic conductor (such as a metal or carbon) or a semiconductor. Current flows through the electrodes via the movement of electrons. An electrolyte is a phase in which charge is carried by ions. For example, a solution of table salt (sodium chloride, NaCl) in water is an electrolyte containing sodium cations (Na+) and chloride anions (Cl−). When an electric field is applied across this solution, the ions move: Na+ toward the negative side of the field and Cl− toward the positive side.
The half-cells are connected by a cell separator that allows ions to move between the half-cells but prevents mixing of the electrolytes. The separator can consist of a salt bridge, or tube of aqueous solution plugged at both ends with glass wool, or it can be an ion exchange membrane or a sintered-glass disk. In some cases both half-cells use the same electrolyte, so that the electrochemical cell consists of two electrodes in contact with a single electrolyte. Electrochemical cells are usually classified as either galvanic or electrolytic. In galvanic cells, reactions occur spontaneously at the electrode–electrolyte interfaces when the two electrodes are connected by a conductor such as a metal wire. Galvanic cells convert chemical energy to electric energy and are the components of batteries, which usually contain several cells connected in series. In electrolytic cells, reactions are forced to occur at the electrode–electrolyte interfaces by way of an external source of power connected to both electrodes. Electric energy from the external source is converted to chemical energy in the form of the products of the electrode reactions.
The galvanic cell shown in Figure 1 is known as the Daniell cell and was used as an early source of energy. It consists of a zinc (Zn) electrode in contact with an aqueous zinc sulfate solution and a copper (Cu) electrode in contact with an aqueous copper sulfate solution. When the external switch is closed, an atom of zinc on the zinc electrode is oxidized to zinc ion, liberating two electrons.
Zn → Zn2+ + 2e − (1)
The electrons pass through the external wire and reduce a copper ion to an atom of copper metal on the surface of the copper electrode.
Cu2+ + 2e − → Cu (2)
The electron flow in the external circuit represents an electric current produced by the cell. Ions flow within the electrolytes and across the salt bridge, as shown in the figure, to prevent an imbalance of ionic charge in the solutions that could result from the occurrence of these two electrode reactions. The overall cell reaction is the reduction of copper ion by zinc.
Cu2+ + Zn → Cu + Zn2+ (3)
The electrolytic cell shown in Figure 2 is the industrial chloralkali cell in which brine (an aqueous sodium chloride solution) is electrolytically converted to chlorine and caustic soda (sodium hydroxide, NaOH). The external power source supplies electric energy to drive the overall reaction.
2Cl− + 2H2O → Cl2 + H2 + 2OH− (4)
Chloride ion is oxidized to chlorine gas at the carbon electrode, and water is reduced to hydrogen gas (H2) and hydroxide ion (OH−) at the iron electrode. The electrolytes are maintained as electrically neutral by a flow of sodium ions through the separator (such as an ion exchange membrane).
The electrode where oxidation occurs, the zinc electrode in Figure 1 and the carbon electrode in Figure 2, is called the anode, while the electrode where reduction occurs is called the cathode. Reactions (1) and (2) are known as half-reactions, whereas reactions (3) and (4) are called oxidation-reduction (redox) reactions.
Current and potential (or voltage ) are the two electrical variables of greatest interest in electrochemical cells. Current is related to the rate of the electrode
reactions, and the potential, to the cell energetics. Current is measured in amperes (A), or the amount of electricity in coulombs (C) that passes across a medium per second(s). Potential between the two electrodes is measured in volts (V) with a voltmeter. Potential (V) has units of energy or work (joules, J) per amount of electric charge (C). That is, 1 V = 1 J/C, so that the cell potential is a measure of the energy of the cell reaction. The cell is said to be at open circuit when no current flows; that is, when there are no external connections to the electrodes. Under these conditions, no electrode reactions occur.
Measurements of the potentials of galvanic cells at open circuit give information about the thermodynamics of cells and cell reactions. For example, the potential of the cell in Figure 1, when the solution concentrations are 1 molar (1 M) at 25°C, is 1.10 V. This is called the standard potential of the cell and is represented by E°. The available energy (the Gibb's free energy ΔG °) of the cell reaction given in equation (3) is related to E° by
ΔG ° = −nF E° (5)
where n is the number of electrons transferred in the reaction (in this case two) and F is a proportionality constant, called the Faraday (96,485 coulombs/equivalent). The cell potential is the difference in potential of the two half-cells. Tables of standard electrode potentials of half-reactions have been compiled; representative values are given in Table 1. These are frequently tabulated with respect to the standard or normal hydrogen electrode (SHE or NHE), which is arbitrarily assigned a half-cell potential of zero. Thus the value, +0.34 V, is assigned to the half-reaction Cu2+ + 2e − → Cu and
E° = +0.34 V vs NHE (6)
Similarly, the standard potential for the Zn/Zn2+ cell yields Zn2+ + 2e − → Zn and
|REPRESENTATIVE STANDARD POTENTIALS|
|Half Reaction||Eo vs NHE|
|SOURCE: Bard, Allen J. et al., eds. (1985). Standard Potentials in Aqueous Solutions. New York: Marcel Dekker.|
|Li+ + e− → Li||−3.045|
|Mg2+ + 2e− → Mg||−2.356|
|Al3+ + 3e− → Al||−1.67|
|Zn2+ + 2e− → Zn||−0.7626|
|Cr3+ + e− → Cr2+||−0.424|
|2H+ + 2e− → H2||0.000|
|Cu2+ + 2e− → Cu||0.340|
|O2 + 4H+ + 4e− → 2H2O||1.229|
|F2 + 2e− → 2F−||2.87|
E° = −0.76 V vs NHE (7)
The difference between these two half-cell potentials yields the standard potential of the Zn-Cu cell.
The standard potential applies to a half-cell when all the reactants are present at unit activity; that is, when the solution species are near a concentration of 1 molar. The actual half-cell potential E is a function of the solution concentrations and is related to these and to the standard potential E° by the Nernst equation. The Nernst equation for the generalized half-reaction
Oxidized Species + ne − → Reduced Species (8)
where R is the universal gas constant, T is the temperature on the Kelvin (absolute) scale, and the terms [Oxidized Species] and [Reduced Species] denote the activities of the species in the half-cells. The activities of pure solids and liquids are taken as unity. At 25°C, (RT/F ) = 0.025 V.
In most flashlights, toys, and remote controllers for televisions, primary batteries are used. The cell reactions in primary batteries are irreversible. During use, reactants are converted to products, and when the reactants are used up, the battery is "dead." The inexpensive flashlight batteries sold in retail stores use a design called a Leclanche dry cell. The body of the battery is made of zinc, which acts as the anode. A carbon rod in the center of the cell serves as the cathode. It is surrounded by a moist paste of graphite powder (carbon), manganese dioxide (MnO2), and ammonium chloride (NH4Cl). The anode reaction is the oxidation of the zinc cylinder to zinc ions. The cathode reaction involves the reduction of manganese dioxide. A simplified version of the overall reaction is
Zn + 2MnO2 + H2O → Zn2+ + Mn2O3 + 2OH− (10)
Alkaline cells are similar, except that the zinc case is porous and the paste around the carbon cathode is moist manganese dioxide and potassium hydroxide. These are more expensive than ordinary zinc-carbon cells, but they maintain a high voltage longer.
The lead-acid storage battery used in automobiles is a secondary battery; it is rechargeable. That is, the automobile battery operates as a galvanic cell when used to start the engine (when discharging), and as an electrolytic cell when it is charged by the alternator or by an external battery charger. The anode consists of porous lead plates in contact with a sulfuric acid (H2SO4) solution. The cathode consists of lead dioxide (PbO2) plates, also in sulfuric acid. Electrons flow from the lead plates to the lead oxide plates. As lead (Pb) loses electrons, it forms lead ions (Pb2+) that react with sulfate ions (SO42−) in solution to form insoluble lead sulfate (PbSO4). When PbO2 gains electrons, it too reacts with SO42− ions in solution to form solid PbSO4. The cell reaction is
Pb + PbO2 + 4H+ + 2 SO42− → 2PbSO4 + 2H2O (11)
and proceeds from left to right when the battery is discharging and from right to left when charging.
The rechargeable nickel-cadmium (Ni-Cad) batteries are used in a variety of cordless appliances such as telephones, battery operated tools, and portable computers. During discharge, cadmium metal (Cd) acts as the anode, and nickel dioxide (NiO2) as the cathode. Both metals form insoluble hydroxides due to the presence of the potassium hydroxide electrolyte. The cell reaction during discharge is
NiO2 + Cd + 2H2O → Ni(OH)2 + Cd(OH)2 (12)
The reaction is reversed during charging.
see also Aluminum; Faraday, Michael.
Cynthia G. Zoski
Bard, Allen J.; Parsons, Roger; and Jordan, Joseph, eds. (1985). Standard Potentials in Aqueous Solutions. New York: Marcel Dekker.
Bard, Allen J., and Faulkner, Larry R. (2000). Electrochemical Methods: Fundamentals and Applications. New York: Wiley.
Bockris, J. M., and Reddy, A. K. N. (1998). Modern Electrochemistry. New York: Plenum.
Oldham, Keith B., and Myland, Janice C. (1994). Fundamentals of Electrochemical Science. New York: Academic Press.
Rieger, Philip H. (1994). Electrochemistry, 2nd edition. New York: Chapman and Hall.
"Electrochemistry." Chemistry: Foundations and Applications. . Encyclopedia.com. (December 13, 2018). https://www.encyclopedia.com/science/news-wires-white-papers-and-books/electrochemistry
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electrochemistry, science dealing with the relationship between electricity and chemical changes. Of principal interest are the reactions that take place between electrodes and the electrolytes in electric and electrolytic cells (see electrolysis), as well as the reactions that take place in an electrolyte as electricity passes through it. The principles of electrochemistry are applied in a variety of ways, e.g., in electroplating and in the generation of electricity by magnetohydrodynamics. See battery; voltaic cell.
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"electrochemistry." The Columbia Encyclopedia, 6th ed.. . Retrieved December 13, 2018 from Encyclopedia.com: https://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/electrochemistry
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