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Electrolytic cell
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An electrolytic cell decomposes 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 hydroxide, and bauxite into aluminium and other chemicals.
lectrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode).
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Encyclopedia
An electrolytic cell decomposes 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 hydroxide, and bauxite into aluminium and other chemicals.
Components
An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. 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 battery or Galvanic 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-cells, 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 battery, such as a AA NiMH cell or a single cell of a lead-acid battery, acts as a galvanic 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 silver ions Ag+), to be reduced by reacting with (negatively-charged) electrons on the cathode. Likewise he defined the anode as the electrode to which anions flow (negatively charged ions, like chloride ions 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 battery, 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 chloride ions (Cl-) to chlorine gas, releasing electrons to the anode. Likewise the cathode reduces sodium ions (Na+), which accept electrons from the cathode and deposits on the cathode as sodium metal.
NaCl dissolved in water can also be electrolyzed. The anode oxidixes chloride ions (Cl-), and Cl2 gas is still produced. However, at the cathode, instead of sodium ions being reduced to sodium metal, water molecules are reduced to hydroxide ions (OH-) and hydrogen gas (H2). The overall result of the electrolysis is the production of chlorine gas and aqueous sodium hydroxide (NaOH) solution.
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
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