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Electrochemistry is a branch of chemistry dealing with chemical reactions that involve electrical currents and potentials. Some chemical reactions that proceed spontaneously can generate electrical current, which can be used to do useful work; while other chemical reaction can be forced to proceed by using electrical current. While all this may sound rather esoteric, many practical devices based on these reactions, and many products made by these reactions are well-known, everyday household items.
Some of these batteries are used in today’s “hybrid automobiles” in conjunction with gasoline engines. They are also candidates for power source in a fully electric car, once the weight of the batteries and their price could be further reduced.
In another version of the electric car, the power would be provided by fuel cells. Fuel cells are somewhat like batteries except that, rather than having a limited amount of chemicals built in, they can be continuously fed by a fuel. A fuel-cell car would be occasionally filled-up by a fuel, much like today’s gasoline-engine cars. The most likely fuel will be hydrogen, but research is being done to use gasoline-like fuels. The most likely earliest, practical commercial application of fuel cells will be power sources for portable electronic devices, lap-top computers, cell-phones, etc. These fuel cells would use alcohol as a fuel; so, rather than occasionally recharging the device from an electrical source, one only will have to insert a new fuel cartridge. The main difficulty with fuel cells is their price. Otherwise, technically reliable and well-working fuel cells have been available for decades: the U.S. space program has been using them routinely, as early as the Apollo program. Of course, in a space capsule price is not the primary consideration, but reliability is.
Many everyday chemical products in a household are connected to electrochemistry. “Bleach” is made from the products of brine electrolysis (“chlorine” and “caustic soda”), or it can be made directly with an electrochemical cell. If you have a pool, you may have a small device to produce locally small amounts of chlorine or bleach, which you need to treat your pool water. Your drinking water is probably treated with chlorine, possibly produced on-site by small electrochemical cells or purchased from companies producing it by brine electrolysis. Chlorine is also a basic ingredient in many plastics, like PVC, used in many homes for piping.
Many cleaning agents, detergents, soaps are made with the “caustic soda” that is also generated in the brine-electrolysis cells. All the paper you use was produced using large amounts of caustic soda, and it was possibly treated by chlorine to bleach it. Caustic soda is also used in the preparation of many food products.
Everything around you that is made of aluminum was made by an electrochemical process, which is the only practical and economical way to produce aluminum metal from its ore. Many other commonly used metals can be produced from their ores or refined (purified) by electrochemical processes. Some of these are: copper, zinc, silver, lead, and many more. Metal articles are often protected from corrosion by coating them with a more corrosion resistant metal. More often than not, this is carried out by electroplating, chrome plating is a good example. Decorative coatings are also applied electrochemically, utensils and jewelry are often silver or gold plated to improve appearance. Sometimes the whole jewelry piece is made by an electrodeposition process called electroforming.
Another everyday use of electrochemical devices is in chemical analysis. In the laboratory, electrochemical techniques are widely used. But even in everyday life there are electrochemical applications. For example, the glucose sensor used by diabetics is often an electrochemical device.
The above is a brief and cursory overview of the many aspects of everyday life connected to electrochemistry. But how is this electrochemistry different from chemistry in general? What is the connection between electricity and chemistry?
Atoms and molecules consist of a heavy nucleus with electrons swirling around it. The general definition of “oxidation/reduction” is the donation of electrons by one molecule or atom to another. The species that donates the electron is said to be “oxidized”, and the species receiving the electron is said to be “reduced”. Oxidation and reduction always occur together, oxidation is always accompanied by something being reduced, and reduction is always accompanied by something being oxidized. The electron has to go from somewhere to somewhere. These electrons are the same electrons that carry electrical currents in the household circuit or any electronic device. Under normal conditions this electron movement between atoms or molecules cannot be used to carry out useful work, the electron movement happens within an extremely short distance; it can be considered an “internal current”. But, if the two parts of the reaction, the oxidation and the reduction, are separated by a large (atomically speaking) distance, and the two reaction sites are connected by a wire for the electrons to travel, now this “external current” can be used to carry out useful work, light a light bulb or run a motor. This is electrochemistry in practice!
One of the simplest examples is the oxidation of hydrogen by oxygen to produce water. It is well known that hydrogen gas is explosive if mixed with air, the mixture is unstable; all it needs is a spark or a small flame to start the reaction, which than proceeds spontaneously (and explosively) to produce water vapor. In this reaction, electrons move from the hydrogen to the oxygen, but this “electrical current” is useless as far as we are concerned. But this is the same reaction that in a “hydrogen-oxygen” fuel cell can produce useful current (say to drive an automobile in the future).
|Fig. 1. Operating principle of a fuel cell. (Copied from an article on solid oxide fuel cells.)|
There is another type of electrochemical reaction: one that will not proceed unless forced to do so. The reverse reaction of the hydrogen-oxygen fuel cell reaction is a good example. Water, which is composed of hydrogen and oxygen, can be decomposed to yield its pure elements. But this is not a spontaneous reaction (at least not under “normal” conditions, that is: at room temperature and atmospheric pressure), it has to be forced by an input of energy. One way to force it is by an electrical current between two electrodes immersed in some water solution. The reverse reactions of those shown in Figure 1 will occur: hydrogen and oxygen gas will be produced, and they can be collected separately. This is an old technology with many past and present practical uses, for example, the production of oxygen in submarines to refresh the sub’s atmosphere. In this type of electrochemical cells, electrical energy is used to produce a chemical change. See the Appendix for some details about the fuel cell and the water-electrolysis reactions.
Very briefly, this is the essential story of electrochemistry. All the practical examples given in the previous section are only variations on the same theme. Of course, electrochemists worry about a number of additional details involved in these systems. What are the electrical potential differences between the electrodes? How will these potentials change when current is flowing? How will the ions carry the current in the electrolyte between the electrodes? What are the best chemicals (to be oxidized/reduced) to be used? What are the best electrode materials to be used? What is the best electrolyte to be used? But, for the purposes of this introduction, these can just be ignored.
There may be some details worth mentioning. The electrodes in the electrochemical cells are usually made of metals or alloys, sometimes of carbon. The electrode where oxidation occurs (that is, where electrons are donated to the electrode) is called the “anode”, and the electrode where reduction occurs (that is, where electrons are accepted from the electrode) is called the “cathode”. The electrolyte is typically a solution in water, sometimes in other solvents, occasionally it is a molten salt or even an ionically conducting solid.
It may also be worth mentioning that the above described two examples of electrochemical cells, the hydrogen-oxygen fuel cell and the water-electrolysis cell, would be important components of the proposed “hydrogen economy”. It is envisioned that solar electricity could be used to electrolyze water and produce hydrogen fuel. Or hydrogen could be directly produced from water using the energy of the sunlight in photoelectrochemical cells. The hydrogen would be piped to the point of use, either central stations, or individual homes, where fuel cells would turn it back into electricity. The hydrogen could also be used in transportation system (electric cars, etc.) to provide electric power through fuel cells.
In batteries, two chemicals, which are eager to react, are placed into the battery during its production in such a way that they cannot react directly (being spatially separated), but they can donate/receive electrons (getting oxidized/reduced) at electrodes inserted into them, and thereby produce useful electrical current. Once all the material reacted away, the battery has to be replaced. With rechargeable batteries, an electrical current can be forced through the used battery (charging the battery) that will reverse the reaction and restore the original chemicals, so the battery can be used again to produce electrical current. The chemicals and the reactions involved in the batteries are generally more complex (sometimes very much so) than the simple hydrogen-oxygen example discussed above, but this is a detail, the overall scheme remains the same.
Many everyday metals are produced (or purified) in electrochemical cells where electrical energy is used to produce a chemical change. These are reactions that have to be forced to proceed, and often the use of electrical current is the best choice to do so; in some cases this is the only practical choice (as for aluminum production). The metal ores are dissolved in some solvent, and the resulting solution is decomposed in the cell. The metal ions are reduced at the cathode to produce a solid (in some cases a liquid) metal deposit of the pure metal. Under special conditions, these processes can be used to produce loose metal powders, with controlled shape and size, rather than solid metals. The same reaction occurs in the electroplating cells, where a corrosion resistant or decorative metal coating is applied to a metal object that is made the cathode of the cell. In these cells, the inevitably occurring coupled oxidation reaction at the anode is more often than not oxygen evolution from the water, this oxygen is considered a useless byproduct and is vented into the air.
Metal objects can also be formed electrochemically by using either the reduction or the oxidation processes that occur in the cells. Electroforming is a special application of electroplating where a prefabricated form is completely filled with a thick metal deposit (by reduction of the metal ions from the solution) to produce a solid metal object in the shape of the form; no further machining is needed to obtain the final object. This process is often used to form precious metal jewelry. The opposite reaction occurs in electrochemical machining. Here the metal to be formed is made the anode in the electrochemical cell where the metal will be dissolved (oxidized into a soluble metal compound) following the form of a very closely spaced cathode (in this case, the coupled hydrogen evolution on the cathode is considered a useless byproduct). This process is often used to produce metal object from alloys that are very hard and difficult to machine into complex shapes (for example, turbine blades). In all these processes the chemicals and the reactions involved can be rather complex (sometimes very much so) but this is a detail, the overall spatially-separated oxidation/reduction scheme remains the same.
Many other chemicals (organic or inorganic) are also produced in electrochemical cells. As already mentioned above, there are always two products generated; both oxidation and reduction must occur at the same time. Unfortunately, often only one of these products is of interest, and the other is a useless byproduct. Nevertheless, these processes are often still provide the most economical choice. There is one notable exemption: the electrolysis of brine (solution of common table salt). The oxidation reaction produces chlorine gas, while the reduction produces hydrogen gas and with it a basic solution (see Eq.  in the Appendix), “caustic soda”. Both the chlorine and the caustic soda are useful products. For decades, the hydrogen was generally considered a useless byproduct and was just burned; recently, in some installations the hydrogen is fed to hydrogen-oxygen fuel cells to regain some of the electrical power used in the process; this results in three useful products from one process. Both the chlorine and the caustic are very essential, large scale chemicals, they are used in many chemical processes and products, many of which are for household use. It indicates the scale of the process that it uses about 1.5% of all electrical power generated in the U.S.; among the electrochemical industries this is second only to the aluminum metal production, which uses about twice as much electrical power.
There is one useful electrochemical device that does not follow the central theme. “Electrochemical capacitors” are considered in conjunction with batteries and fuel cells for hybrid power sources in electric cars. There is no reaction (oxidation/reduction) occurring in these devices, they are simply a different (and in some respect better) versions of the classical electrical capacitors.
There are different ways to write the reactions for hydrogen-oxygen fuel cells, depending on what the electrolyte is and what ions predominate, Considering a cell with a solution formed in water as the solvent, the reactions occurring in a hydrogen-oxygen fuel cell would depend whether the solution is acidic or basic. The water is in equilibrium with its ions: “ H2O <==> H+ + OH- ” but in acidic solution the hydrogen ions predominate, while in basic solutions the hydroxyl ions predominate.
The hydrogen-oxygen fuel cell reactions in an acidic electrolyte are:
|||H2 ==> 2H+ + 2e-||(oxidation reaction)|
|||½O2 + 2H+ + 2e- ==> H2O||(reduction reaction)||||H2 + ½O2 ==> H2O||(overall reaction)|
The hydrogen-oxygen fuel cell reactions in a basic electrolyte are:
|||H2 + 2OH- ==> 2H2O + 2e-||(oxidation reaction)|
|||½O2 + H2O + 2e- ==> 2OH-||(reduction reaction)||||H2 + ½O2 ==> H2O||(overall reaction)|
In either case, the reactions involve the oxidation of hydrogen and the reduction of oxygen to produce water. This is a spontaneous reaction, which can be used to produce electrical current.
For the case of the reverse overall reaction, the decomposition of water into hydrogen and oxygen, the reactions for an acidic electrolyte are:
|||H2O ==> ½O2 + 2H+ + 2e-||(oxidation reaction)|
|||2H++ 2e- ==> H2||(reduction reaction)||||H2O ==> H2 + ½O2||(overall reaction)|
For the decomposition of water into hydrogen and oxygen, the reactions for a basic electrolyte are:
|||2OH- ==> ½O2 + H2O + 2e-||(oxidation reaction)|
|||2H2O + 2e- ==> 2OH- + H2||(reduction reaction)||||H2O ==> H2 + ½O2||(overall reaction)|
In either case, the reactions involve the decomposition of water into hydrogen and oxygen, essentially the oxidation of water to produce oxygen and, at the same time, the reduction of water to produce hydrogen. This is not a spontaneous reaction; it has to be forced to proceed using an electrical current.
Listings of electrochemistry books, review chapters, proceedings volumes, and full text of some historical publications are also available in the Electrochemistry Science and Technology Information Resource (ESTIR). (http://electrochem.cwru.edu/estir/)
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