Latimer diagram

Latimer diagrams , also called potential diagrams , are a representation of the reduction potentials of half-reactions , each for the different oxidation states of an element. Usually the highest oxidation number of the element is on the far left, while the oxidation number decreases to the right. The individual stages are connected with arrows, above which the reduction potential of the half-reaction is shown. These can refer to standard conditions (25  ° C , pH = 0, c = 1 mol·l ⁻¹ ) or to any other specified conditions (e.g. pH = 14).

The diagram is named after the American chemist Wendell Mitchell Latimer (1893–1955).

Examples

Diagram for chlorine in acidic solution (pH = 0)

${\displaystyle \mathrm {ClO_{4}^{-}{\xrightarrow {+1.20V}}\ ClO_{3}^{-}{\xrightarrow {+1.18V}}\ ClO_{2}^{-}{\xrightarrow {+1.65V}}\ HClO{\xrightarrow {+1.63V}}\ Cl_{2}{\xrightarrow {+1.36V}}\ Cl^{-}} }$

If the detailed reduction reaction equation for the first step is to be written down, the half-reaction has to be balanced with regard to the charge and also stoichiometrically :

${\displaystyle \mathrm {ClO_{4}^{-}+2H^{+}+2e^{-}\rightarrow ClO_{3}^{-}+H_{2}O\ mit\ E^{0}=+1.20V} }$

Diagram for selenium under standard conditions (pH = 0.25 ° C, 10 ⁵ Pa)

${\displaystyle \mathrm {SeO_{4}^{2-}{\xrightarrow {1,15V}}\ H_{2}SeO_{3}{\xrightarrow {0,74V}}\ Se{\xrightarrow {-0,08V}}\ H_{2}Se} }$

Latimer diagram for oxygen

The potential diagram for oxygen , hydrogen peroxide and water is shown on the right; the oxidation numbers of the oxygen atoms (0, −1 and −2) are given.

The first potential, 0.68 V, applies to the reduction of oxygen to hydrogen peroxide, i.e. for the reaction

${\displaystyle \mathrm {O_{2}(g)+2\ H^{+}+2\ e^{-}\rightleftharpoons H_{2}O_{2}(aq)\qquad mit\ } E^{0}=+0,68\mathrm {\ V\ bei\ pH=0} }$.

The second equilibrium reaction is the reduction of hydrogen peroxide to water at a potential of 1.78 V: . ${\displaystyle \mathrm {H_{2}O_{2}(aq)\ +2H^{+}+2\ e^{-}\rightleftharpoons 2\ H_{2}O\qquad mit\ } E^{0}=+1,78\mathrm {\ V\ bei\ pH=0} }$

The potential indicated on the lower arrow applies to the reduction of oxygen to water, i.e. for the reaction

${\displaystyle \mathrm {O_{2}\ +4H^{+}+4\ e^{-}\rightleftharpoons 2\ H_{2}O\qquad mit\ } E^{0}=+1,23\mathrm {\ V\ bei\ pH=0} }$.

The potential of this reaction, 1.23 V, also corresponds to the decomposition voltage of water and is the mean value of the two individual potentials 0.68 V and 1.78 V, since the same number of electrons is exchanged in these reactions. Hydrogen peroxide tends to break down into water and oxygen, it disproportionates . This can be seen in the potential diagram by the fact that the right potential, 1.78 V, is higher than the left.

application

Using a Latimer diagram, it is easy to see whether an oxidation state in solution is unstable and tends to disproportionate : If - as usual - the highest oxidation state is on the left, a species will disproportionate if the potential to the left of it is smaller than that to the right her. Accordingly, hydrogen peroxide is unstable.

If you want to determine the potential of a reaction in which the oxidation number changes by several units, you must not add the redox potentials of the steps in between. Instead, an average must be formed whose contributions are weighted with the number of electrons exchanged in each case: ${\displaystyle {\bar {E}}}$ ${\displaystyle E_{i}}$ ${\displaystyle n_{i}}$

${\displaystyle {\bar {E}}={\frac {\sum n_{i}\cdot E_{i}}{\sum n_{i}}}}$

See also

• Frost diagram , a graphic representation of the numerical values ​​of the Latimer diagram
• Pourbaix diagram showing the potential values ​​for different pH values

literature

• Wendell Mitchell Latimer : The Oxidation States of the Elements and their Potentials in Aqueous Solutions. Prentice Hall, New York 1938.

Individual evidence

1. ↑ Latimer Diagrams

Web links

Exercise on the Latimer diagram