Electrochemical cell
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[[Image:ElectrochemCell.png|thumb|250px|right|A demonstration electrochemical cell setup resembling the [[Daniell cell]]. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.]]
An '''electrochemical cell''' is a device used for generating an [[electromotive force]] ([[voltage]]) and current from [[electrochemistry|chemical reactions]]. The current is caused by the reactions releasing and accepting [[electron]]s at the different ends of a conductor. A common example of an electrochemical cell is a standard 1.5-volt '''[[Battery (electricity)|battery]]'''.
==Overview==
An electrochemical cell consists of two half-cells. The two half-cells may use the same electrolyte, or they may use different electrolytes. Each ''half-cell'' consists of an electrode, and an electrolyte. One half-cell undergoes oxidation, and the other one undergoes reduction. The reactions may involve the electrolyte, the electrodes or an external substance (as in fuel cells which may use hydrogen gas as a reactant). In a full electrochemical cell, ions, atoms or molecules from one half-cell lose electrons (oxidation) to their [[electrode]] while ions, atoms or molecules from the other half-cell gain electrons (reduction) from their electrode. Finally, a ''[[salt bridge]]'' is often employed to provide electrical contact between two half-cells with very different electrolytes—to prevent the solutions from mixing. This can simply be a strip of [[filter paper]] soaked in saturated potassium nitrate (V) solution. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell below (between the electrodes).
[[image:bunsen_cell.jpg|thumb|left|200px|The [[Bunsen cell]], invented by [[Robert Bunsen]].]]
Each half-cell has a characteristic voltage. Different choices of substances for each half-cell give different potential differences. Each reaction is undergoing an [[Chemical equilibrium|equilibrium]] reaction between different oxidation states of the ions—when equilibrium is reached the cell cannot provide further voltage. In the half-cell which is undergoing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Similarly, in the reduction reaction, the further the equilibrium lies to the ion/atom with the more ''negative'' oxidation state the higher the potential.
The cell potential can be predicted through the use of [[electrode potential]]s (the voltages of each half-cell). (Tables of electrode potentials are available in chemistry reference books.) The difference in voltage between electrode potentials gives a prediction for the potential measured.
Cell potentials have a possible range of about zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts.
==Cell types==
* [[Secondary cell|Accumulator]]
* [[Concentration cell]]
* [[Electrolytic cell]]
* [[Galvanic cell]]
* [[Galvanic corrosion#Lasagna cell|Lasagna cell]]
* [[Lemon battery]]
==See also==
{{portal|energy}}
* [[Electrochemical potential]]
* [[Nicad|Nickel Cadmium battery]]
* [[Alkaline battery]]
* [[Cell notation]]
[[Category:Electric batteries]]
[[Category:Electrochemistry]]
[[es:Célula electrolítica]]
[[fr:Pile électrique]]
[[he:תא אלקטרוכימי]]
[[no:Galvanisk element]]
[[pt:Célula electroquímica]]
[[ru:Химические источники тока]]
[[sv:Elektrokemisk cell]]
[[zh:电化电池]]