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The data below tabulates standard electrode potentials (E°), in volts relative to the standard hydrogen electrode (SHE), at: Temperature 298.15 K (25.00 °C; 77.00 °F); Effective concentration (activity) 1 mol/L for each aqueous or amalgamated (mercury-alloyed) species; Unit activity for each solvent and pure solid or liquid species; and
Where is the standard reduction potential of the half-reaction expressed versus the standard reduction potential of hydrogen. For standard conditions in electrochemistry (T = 25 °C, P = 1 atm and all concentrations being fixed at 1 mol/L, or 1 M) the standard reduction potential of hydrogen is fixed at zero by convention as it serves of reference.
The electric potential also varies with temperature, concentration and pressure. Since the oxidation potential of a half-reaction is the negative of the reduction potential in a redox reaction, it is sufficient to calculate either one of the potentials. Therefore, standard electrode potential is commonly written as standard reduction potential.
Scheme of the standard hydrogen electrode. The scheme of the standard hydrogen electrode: platinized platinum electrode; hydrogen gas; solution of the acid with activity of H + = 1 mol dm −3; hydroseal for preventing oxygen interference; reservoir through which the second half-element of the galvanic cell should be attached.
The reference half cell consists of a redox standard of known potential. The standard hydrogen electrode (SHE) is the reference from which all standard redox potentials are determined, and has been assigned an arbitrary half cell potential of 0.0 V. However, it is fragile and impractical for routine laboratory use.
The electrochemical series, which consists of standard electrode potentials and is closely related to the reactivity series, was generated by measuring the difference in potential between the metal half-cell in a circuit with a standard hydrogen half-cell, connected by a salt bridge. The standard hydrogen half-cell: 2H + (aq) + 2e − → H 2 (g)
To avoid possible ambiguities, the electrode potential thus defined can also be referred to as Gibbs–Stockholm electrode potential. In both conventions, the standard hydrogen electrode is defined to have a potential of 0 V. Both conventions also agree on the sign of E for a half-cell reaction when it is written as a reduction.
For example, the absolute half-cell entropy has been defined as the entropy of the absolute half-cell process defined above. [4] An alternative definition of the absolute half-cell entropy has recently been published by Fang et al. [5] who define it as the entropy of the following reaction (using the hydrogen electrode as an example):