Electrolytic cell
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2008-07-05T08:51:28Z
SJP
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Reverted edits by [[Special:Contributions/123.237.190.198|123.237.190.198]] to last version by Bobo192 (using [[WP:HG|Huggle]])
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==Basics==
'''Electrolytic cells''' decompose chemical compounds by means of electrical energy, in a process called [[electrolysis]]; the Greek word [[lysis]] means ''to break up''. The result is that the chemical energy is increased. Important examples of electrolysis are the decomposition of [[water]] into hydrogen and oxygen gas, and [[bauxite]] into aluminium and other chemicals.
An electrolytic cell contains three elements: an [[electrolyte]] and two electrodes (a [[cathode]] and an [[anode]]). The [[electrolyte]] is usually a [[solution]] of [[water]] or other [[solvent]]s in which one dissolves the normally inert compound into its [[ion]]ic form. When driven by an external [[voltage]] applied to the electrodes, the electrolyte provides ions that flow to and from the electrodes, where charge-transferring, or faradaic, or [[redox]], reactions can take place. Only for an external [[electrical potential]] (i.e. voltage) of the correct polarity and large enough magnitude can an electrolytic cell decompose a normally stable, or [[inert]] chemical compound in the solution. The electrical energy provided undoes the effect of spontaneous chemical reactions.
==Galvanic Cells Compared to Electrolytic Cells==
In contrast, a [[galvanic cell]], also known as a [[voltaic cell]], converts chemical energy into electrical energy, by using spontaneous chemical reactions that take place at the electrodes. Each galvanic cell has its own characteristic voltage (defined as the energy release per electron transfer from one electrode to the other). A simple galvanic cell will consist only of an electrolyte and two different electrodes. (Galvanic cells can also be made by connecting two [[half-cell]]s, each with its own electrode and electrolyte, by an ion-transporting "bridge", usually a [[salt bridge]]; these cells are more complex.) The electrodes typically are two metals, which naturally have different reaction potentials relative to the electrolyte. This causes ions of one of the electrodes to preferentially enter the solution at one electrode, and another ion to leave the solution at the other electrode. This generates an electric current across the electrolyte, which will drive electric current through a wire that makes an exterior connection to each of the electrodes. A galvanic cell uses electrodes of different metals, whereas an electrolytic cell may use the same metal for cathode and anode.
A rechargeable voltaic cell, such as a AA [[NiMH]] cell or a single cell of a [[lead-acid battery]], is called a voltaic cell when discharging (converting chemical energy to electrical energy), and an electrolytic cell when being charged (converting electrical energy to chemical energy).
==Anode and Cathode Definitions Depend on Charge and Discharge==
[[Michael Faraday]] defined the [[cathode]] as the electrode to which cations flow (positively charged ions, like Cu+), to be "reduced" by reacting with (negatively-charged) electrons on the cathode. Likewise he defined the [[anode]] as the electrode to which flow anions (negatively charged ions, like Cl-), to be "oxidized" by depositing electrons on the anode. Thus positive electric current flows from the cathode to the anode. To an external wire connected to the electrodes of a voltaic cell, thus forming an electric circuit, the cathode is positive and the anode is negative.
Consider two voltaic cells, A and B, with the voltage of A greater than the voltage of B. Mark the positive and negative electrodes as anode and cathode. Place them in a circuit with anode near anode and cathode near cathode, so the cells will tend to drive current in opposite directions. The cell with the larger voltage discharges, making it a voltaic cell. Likewise the cell with the smaller voltage charges, making it an electrolytic cell. For the electrolytic cell, the external markings of anode and cathode are opposite the chemical definition. That is, the electrode marked as anode for discharge acts as the cathode while charging and the electrode marked as cathode acts as the anode while charging.
==Uses==
As already noted, water, particularly when ions are added (salt water or acidic water) can be ''electrolyzed'' (subject to electrolysis). When driven by an external source of voltage, H+ ions flow to the cathode to combine with electrons to produce hydrogen gas in a reduction reaction. Likewise, OH- ions flow to the anode to release electrons and an H+ ion to produce oxygen gas in an oxidation reaction.
In molten sodium chloride, when a current is passed through the salt the anode oxidizes chlorine ions (Cl-), which release electrons to the anode. Likewise the cathode reduces sodium metal (Na+), which accept electrons from the cathode and deposits on the cathode as sodium metal. [[valence (chemistry)|valence]] shell.
NaCl dissolved in water can also be electrolyzed. The anode oxidixes chlorine ions (Cl-), and Cl<sub>2</sub> gas is still produced. However, instead of sodium metal, aqueous sodium hydroxide (NaOH) is produced, which stays dissolved in the water. Some of the water will also be electrolyzed, producing H<sub>2</sub> gas. The oxygen is found in the hydroxide ion which combines with the sodium ions to make the sodium hydroxide.
The net equation is:
2 H<sub>2</sub>O (l) + 2 NaCl (aq) --> Cl<sub>2</sub>(g/aq) + 2 NaOH (aq) + H<sub>2</sub> (g)
Commercially, electrolytic cells are used in electrorefining and [[electrowinning]] of several non-ferrous metals. Almost all high-purity [[aluminium]], [[copper]], [[zinc]] and [[lead]] is produced industrially in electrolytic cells.
==Cell types==
* [[Concentration cell]]
* [[Electrochemical cell]]
* [[Galvanic cell]]
[[Category:Electrochemistry]]
[[Category:Electric batteries]]
[[es:CĂ©lula electrolĂtica]]