Latimer diagram
A Latimer diagram of a chemical element is a summary of the standard electrode potential data of that element. This type of diagram is named after Wendell Mitchell Latimer (1893–1955), an American chemist.
Construction
[edit]In a Latimer diagram, because by convention redox reactions are shown in the direction of reduction (gain of electrons), the most highly oxidized form of the element is on the left side, with successively lower oxidation states to the right side. The species are connected by arrows, and the numerical value of the standard potential (in volts) for the reduction is written at each arrow. For example, for oxygen, the species would be in the order O2 (0), H2O2 (–1), H2O (-2):
The arrow between O2 and H2O2 has a value +0.68 V over it, it indicates that the standard electrode potential for the reaction:
- O2(g) + 2 H+ + 2 e− ⇄ H2O2(aq)
is 0.68 volts.
Application
[edit]Latimer diagrams can be used in the construction of Frost diagrams, as a concise summary of the standard electrode potentials relative to the element. Since ΔrGo = -nFEo, the electrode potential is a representation of the Gibbs energy change for the given reduction. The sum of the Gibbs energy changes for subsequent reductions (e.g. from O2 to H2O2, then from H2O2 to H2O) is the same as the Gibbs energy change for the overall reduction (i.e. from O2 to H2O), in accordance with Hess's law. This can be used to find the electrode potential for non-adjacent species, which gives all the information necessary for the Frost diagram.
It must be stressed that standard reduction potentials are not additive values. They cannot be directly summed up, or subtracted, from the values in volt indicated in a Latimer diagram. If needed, their calculation must be performed via the difference in Gibbs free energies. The easiest way to proceed is simply to use energies (nE) directly expressed in electron-volt (eV), because the Faraday constant F and the sign minus simplifies on both side of the equation. So, the values of E in volt must be simply multiplied by the number (n) of electron transferred in the considered half-reaction. Since the Faraday constant can disappear from the equation, no need to calculate ΔrGo expressed in joule.
A simple examination of a Latimer diagram can also indicate if a species will disproportionate in solution under the conditions for which the electrode potentials are given: if the potential to the right of the species is higher than the potential on the left, it will disproportionate. Therefore, hydrogen peroxide H2O2 is unstable and will disproportionate in O2 and H2O.
See also
[edit]References
[edit]- Atkins, Peter; Overton, Tina (2010). "5. Oxidation and Reduction: The diagrammatic presentation of potential data §5.12 Latimer diagrams". Shriver and Atkins' Inorganic Chemistry. OUP Oxford. pp. 162–163. ISBN 978-0-19-923617-6.
- Rieger, P.H. (1993). "§1.3 Some Uses of Standard Potentials: Latimer Diagrams". Electrochemistry. Springer. p. 18. ISBN 978-0-412-04391-8.
 | This electrochemistry-related article is a stub. You can help Wikipedia by expanding it. |