Electrocehmistry | Research & Encyclopedia Articles

Electrocehmistry

Electrochemistry is the study of reactions caused by electric current and reactions which are capable of generating electric current. It may also be viewed as the study of the interconversion of chemical and electrical energy. Electrochemistry involves oxidation-reduction reactions. These reactions involve the transfer of electrons and consequent changes in the oxidation states of atoms during chemical reactions. Oxidation-reduction (or redox) reactions occur without passage of electric current when they take place between reagents that were in direct contact: ions or molecules in solution, gases mixed in the vapor phase, or liquids contacting solids. In electrochemistry, however, the reagents undergoing oxidation and reduction are actually physically separated from one another and contained in the compartments of a device called a cell. The reagents that are coupled in these oxidation-reduction reactions use some conduit to transfer charge from one compartment to another--in many cases a wire, in others a gel or solid that allows the passage of electrons from one species to another.

Redox reaction are central to electrochemistry because they are chemical processes that involve the flow of electrons from one chemical entity to another. They can be either a source of electric current, as they are in a battery, or can be driven by electric current, as they are in electroplating processes.

Electrons flowing through a wire or ions flowing through an aqueous solution are both examples of electric current. If the electrons gained and lost by the components in a spontaneous redox reaction can be channeled through a wire or some other conducting material, then the conversion of chemical energy to electrical energy can be measured and the energy released by the spontaneous reaction can be used to do work. By the same token, if electrical current can be added to a system, a nonspontaneous redox reaction can be made to occur by converting electrical energy back into chemical energy to force a reaction to take place.

A voltaic cell generates electrical energy from a spontaneous redox reaction. Common batteries in watches, calculators and other devices are voltaic cells. In the cell, two electrodes, conductors by which electrons enter or leave a conducting medium, are connected by a material that carries the electrons between separate chambers containing the species involved in the electrochemical reaction. Both electrodes are immersed in solutions or other conducting materials, such as a moist paste, containing ions called electrolytic solutions or electrolytes. As the electrochemical reaction proceeds in a cell, ions are produced on one side and consumed on the other. To maintain a charge balance in an electrochemical reaction taking place in two different containers, a salt bridge can be used. The salt bridge contains a high concentration of an electrochemically unreactive salt, such as KCl, supported in an aqueous gel. Ions flow out of the salt bridge into the cell to maintain electrical neutrality in the solutions.

When the circuit is complete, electrons spontaneously flow from the electrode in contact with the more easily oxidized substance (in some case, the electrode material itself) to the electrode in contact with the more readily reduced material (again, this can be the electrode material). The reactions resulting from this flow of electric current occur at the electrode surfaces within each chamber in the cell. Oxidation occurs at the electrode from which electrons flow and reduction takes place at the electrode that receives the electrons. Oxidation and reduction occur simultaneously and to the same extent.

Each electrode and its surrounding electrolyte make up a half-cell. By definition, reduction takes place in one half-cell at the electrode called the cathode, and oxidation takes place in the other at the electrode called the anode. The sum of the half-reactions taking place in both cells is the overall redox reaction in the electrochemical cell. The term anode comes from the Greek anodes, or "up and out," for the electrode from which electrons leave a cell and oxidation occurs. Cathode comes from the Greek kathodos, or "down and in," for the electrode by which the electrons enter a cell and reduction occurs.

Because electrons originate at the anode in a voltaic cell, the anode has a (-) charge; electrons enter the cathode, which has a (+) charge. This polarity of the terminals and the role of the electrodes are accurate for any cell that is operating spontaneously and producing electricity. The electric current produced by this type of cell can be used to perform work.

Voltaic cells are often called galvanic cells in honor of Luigi Galvani, who is credited with discovering the phenomenon of electricity. The volt and, hence, voltaic cell are named in honor of Alessandro Volta, who developed some of the first cells that could generate a voltage.

A short-hand notation has been developed to depict the physical arrangement and define the chemical reactions taking place in an electrochemical cell:

anodic half-cell cathodic half-cell

(electrode)│(ionn+)aq∥(ionn+)aq│(electrode)

For example, the reaction: Zno + Cu2+ ⟹Zn2+ + Cuo, would be represented by:

Anodic half-cell Cathodic half-cell

Zno│Zn2+(aq)∥Cu2+(aq)│Cuo

The single vertical line "│" indicates the boundary between a solid electrode and the solution in which it is immersed. The double line "∥" symbolizes the salt bridge.

Half-cells are characterized by their standard reduction potential, which is a measure of how readily a particular species is reduced relative to the hydronium ion. The more readily reduced a species is, the more positive its standard reduction potential. The better an oxidizing agent a species is the more readily it will itself be reduced. Therefore, the standard reduction potential is also a measure of the strength of an oxidizing agent.

For example, the silver ion is a stronger oxidizing agent than the hydronium ion, and so the silver ion is spontaneously reduced and the hydrogen spontaneously oxidized when a silver metal/silver ion half cell is connected to a hydrogen gas/hydronium ion standard half-cell and the standard reduction potential for the silver ion is positive. in this cell. The Ag+(aq) ion is also a stronger oxidizing agent than the Cu2+(aq) ion and has a more positive standard reduction potential (+0.80V vs. +0.34V). If the standard half-cell for copper is connected to the standard half-cell for silver, a spontaneous reaction occurs in which the silver half-cell with its higher standard reduction potential undergoes reduction. Therefore, it functions as the cathode and the copper half-cell functions as the anode.

Electrical current is also used to cause non-spontaneous reactions to occur. An important industrial use of electrochemistry is the production of metals and halogen gases from salts. When a salt such as sodium chloride is melted and electrodes are immersed in it, application of a voltage difference across electrodes leads to a reduction of the sodium cation to metallic sodium and oxidation of the chloride ion to chlorine gas. This decomposition of an electrolyte caused by passing an electric current through it is called electrolysis. The apparatus in which an oxidation-reduction reaction occurs as a result of an applied voltage is called an electrolytic cell.

In an electrolytic cell, energy must be supplied in the form of electric current from an external source. The negative terminal of a power source pumps electrons into the electrode to which it is attached; that electrode is the cathode at which the reduction of sodium occurs. The positive terminal of the power source is connected to the electrode that functions as the anode; the oxidation of chloride ion to chlorine gas occurs at this site. In an electrolytic cell the reduction takes place at the cathode, just as it does in the voltaic cell. But because electrons must be supplied to the cathode by an external power source to drive this process, the cathode in an electrolytic cell has a negative charge. Its polarity is reversed in comparison to cathode in the voltaic cell, but its function--being the site of reduction--is the same.

Because electrolytic reactions require the input of electrical energy to cause a non-spontaneous reaction to take place, the products of electrolytic reactions typically must be physically separated from each other to prevent them from reacting with each other in the spontaneous (reverse) reaction direction. Electrolysis experiments established the composition of water--the fact that it is made up of hydrogen and oxygen and that the ratio of these components in the water molecule is 2:1. To electrolyze water or other non- ionic substances, an ionic substance, an electrolyte, is added to the water so that current will flow and the reaction will proceed at a relatively low voltage. The ions are said to "carry" the current by "migration" to the electrodes.

Despite the explanation used above, however, individual ions do not "migrate" toward electrodes. They are not "attracted" at long distances through water across cells. The insulating power of water is so great that the effect of a positively charged ion on a negatively charged ion falls rapidly to insignificance in a distance equivalent to only a few molecular diameters. The attraction of an ion toward an electrode is certainly not important over distances of millimeters or centimeters within cells. Ions change position over time as an electrochemical reaction proceeds, but the change is caused by normal processes of mixing and charge neutralization, not by the attraction of individual ions by the charge at or near an electrode.

As an electrochemical reaction such as the electrolysis of water proceeds, charge is reduced on species at the cathode and charge is increased on species in the vicinity of the anode. If there were no way to neutralize this change in charge, the electrochemical reaction would soon cease because too much energy would be required to maintain the charge separation. The region of excess positive charge would no longer be able to give up electrons and the region of excess negative charge would no longer accept additional electrons. However, other ions in the vicinity of the electrodes are buffeted about by the motion of the solvent molecules and, once near an ion of the opposite charge, tend to stay nearby. Eventually, the concentrations of ions throughout the solution provide for a uniform neutrality.

The Chlor-alkali Process for the production of sodium hydroxide chlorine gas is carried out on a very large scale; it consumes 1% of the total electricity produced in the United States annually in 1988. The applications of chlorine and sodium hydroxide are so varied that hardly a consumer product exists that does not depend on one or both of these compounds. The chlor-alkali process involves the electrolysis of brines, which are saturated aqueous solutions of sodium chloride. The decorative electroplating of jewelry using gold and silver was the first electrochemical process to be patented and is one of the oldest uses of electrochemistry. Electroplating consists of applying a metallic coating to an article by passing an electric current through a solution of electrolyte in contact with the article. Items ranging from jewelry to tableware ("silverware") to automobile bumpers ("chrome") are covered with precious metals or protective coatings by electroplating.

In this process, the less expensive metal is pre-formed into the desired shape and immersed in a solution of a salt of the plating metal. In the cell, the object to be plated is the cathode, the electrolyte contains the ionic form of the metal to be deposited on the cathodic surface, and an anode is immersed in the solution to complete the circuit. A power source is used to drive the electrochemical reaction. Its positive terminal is connected to the anode and its negative terminal to the cathode where the reduction of the metal ion takes place. When the circuit is complete, the cationic form of the metal to be coated on the object is reduced to the metal on the cathode. The reduced metal typically forms a metallic alloy on the cathode .

The quantitative relationships that govern the amount of material that is produced in electrochemical processes are called Faraday's Laws, afterMichael Faraday, a British scientist of the nineteenth century (1791-1867).

The amount of energy that a voltaic cell can generate or the energy that is required for an electrolytic cell to produce a certain amount of material can be calculated from the total number of mols of electrons that flow during the process and the potential (also called the electromotive force) of the cell. The calculated energy can, in turn, be used to determine the equilibrium constant for the overall reaction that takes place in the cell.

Electrochemical potentials also arise when half-cells containing the same materials but at different concentrations are connected to one another. The potential difference between the two half-cells is greatest when the concentration difference is large. If there is a means for ions to migrate to bring the concentrations of particular ions closer to the same concentration in each half-cell the potential difference decreases. The Nernst equation defines the relationship of the potential of a cell as a function of the concentration of its components and the temperature. Many processes that result in concentration changes in electrochemical cells can be analyzed using the Nernst equation. The concept of concentration cells and the dependence of cell potential on ion concentration is central to our understanding of nerve impulse transmission.

Another important area in which the application of the concepts of electrochemistry is useful is metallic corrosion. Many metals are readily oxidized by oxygen in the atmosphere. Typically the way in which atmospheric corrosion occurs is through the action of oxygen dissolved in water in contact with the metal. The corrosive action is best understood as involving neutral metal atoms, oxidized metal ions, hydronium ions and hydroxide ions in the water and dissolved oxygen gas.

A very common application of electrochemistry is the use of batteries to store electrical energy for use on demand. Although the actual reactions taking place within most batteries are very complex, they can be understood using simple considerations of oxidation and reduction reactions. Dry cells, NiCad batteries and fuel cells are common devices used to supply electrical energy on demand. Research continues on ways to harness electrochemical energy that use lighter, cheaper materials to meet the personal and industrial demands of our increasingly energy intensive world.

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