Electrolytic cell

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Nineteenth-century electrolytic cell for producing oxyhydrogen

An electrolytic cell is an electrochemical cell that utilizes an external source of electrical energy (voltage applied between two electrodes) to drive a chemical reaction that would not otherwise occur. This is in contrast to a galvanic cell, which itself is a source of electrical energy and the foundation of a battery. The net reaction taking place in a galvanic cell is a spontaneous reaction, i.e, the Gibbs free energy remains negative, while the net reaction taking place in an electrolytic cell is the reverse of this spontaneous reaction, i.e, the Gibbs free energy is positive.[1][2]

Principles

In an electrolytic cell, a current passes through the cell by an external voltage, causing a non-spontaneous chemical reaction to proceed. In a galvanic cell, the progress of a spontaneous chemical reaction causes an electric current to flow. An equilibrium electrochemical cell exists at the state between an electrolytic cell and a galvanic cell. The tendency of a spontaneous reaction to push a current through the external circuit is exactly balanced by a counter-electromotive force so that no current flows. If this counter-electromotive force is increased, the cell becomes an electrolytic cell, and if it is decreased, the cell becomes a galvanic cell.[3]

An electrolytic cell has three components: 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 can also function as electrolytes. When driven by an external voltage applied to the electrodes, the ions in the electrolyte are attracted to an electrode with the opposite charge, where charge-transferring (also called faradaic or redox) reactions can take place. Only with an external electrical potential (i.e., voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided can produce a chemical reaction which would not otherwise occur spontaneously.

Michael Faraday defined the cathode of a cell as the electrode to which cations (positively charged ions, like silver ions Ag+
) flow within the cell, to be reduced by reacting with electrons (negatively charged) from that electrode. Likewise, he defined the anode as the electrode to which anions (negatively charged ions, like chloride ions Cl−
) flow within the cell, to be oxidized by depositing electrons on the electrode. To an external wire connected to the electrodes of a galvanic cell (or battery), forming an electric circuit, the cathode is positive and the anode is negative. Thus positive electric current flows from the cathode to the anode through the external circuit in the case of a galvanic cell.

Applications

A video describing the process of electrolytic reduction as used on Captain Kidd's Cannon at The Children's Museum of Indianapolis

Electrolytic cells are often used to decompose chemical compounds, in a process called electrolysis—the Greek word lysis means to break up. Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminium and other chemicals. Electroplating (e.g., of copper, silver, nickel or chromium) is done using an electrolytic cell. Electrolysis is a technique that uses a direct electric current (DC).

Commercially, electrolytic cells are used in electrorefining and electrowinning of several non-ferrous metals. Almost all high-purity aluminium, copper, zinc and lead are produced industrially in electrolytic cells.

As already noted, water, particularly when ions are added (saltwater or acidic water), can be electrolyzed (subjected 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 oxidizes chloride ions (Cl−
), and Cl2 gas is 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, hydrogen and aqueous sodium hydroxide (NaOH) solution.

See also

References

  1. ^ Harris, Daniel C. (2010). Quantitative chemical analysis. (8th ed.). New York: W. H. Freeman and Company. p. GL7. ISBN 978-1429263092
  2. ^ Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2014). Fundamentals of Analytical Chemistry. (9th ed.). Belmont: Brooks/Cole, Cengage Learning. p. 446-449. ISBN 978-0-495-55828-6
  3. ^ Mortimer, Robert (2008). Physical Chemistry (3rd ed.). Elsevier Academic Press. p. 354. ISBN 978-0-12-370617-1
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