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Standard electrode potential

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Electromotive force of a half reaction cell versus standard hydrogen electrode

Inelectrochemistry,standard electrode potential $E^{\ominus }$ , or $E_{red}^{\ominus }$ , is theelectrode potential (a measure of the reducing power of any element or compound) which the IUPAC "Gold Book" defines as"the value of the standardemf (electromotive force) of a cell in which molecular hydrogen understandard pressure is oxidized to solvated protons at the left-hand electrode".^[1]

Background

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The basis for anelectrochemical cell, such as thegalvanic cell, is always aredox reaction which can be broken down into twohalf-reactions:oxidation at anode (loss of electron) andreduction at cathode (gain of electron).Electricity is produced due to the difference ofelectric potential between the individual potentials of the two metalelectrodes with respect to theelectrolyte.

Although the overall potential of a cell can be measured, there is no simple way to accurately measure theelectrode/electrolyte potentials in isolation. 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.

Calculation

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Thegalvanic cell potential results from the voltage difference of apair of electrodes. It is not possible to measure an absolute value for each electrode separately. However, the potential of a reference electrode,standard hydrogen electrode (SHE), is defined as to 0.00 V. An electrode with unknown electrode potential can be paired with either the standard hydrogen electrode, or another electrode whose potential has already been measured, to determine its "absolute" potential.

Since the electrode potentials are conventionally defined as reduction potentials, the sign of the potential for the metal electrode being oxidized must be reversed when calculating the overall cell potential. The electrode potentials are independent of the number of electrons transferred —they are expressed in volts, which measure energy per electron transferred—and so the two electrode potentials can be simply combined to give the overallcell potential even if different numbers of electrons are involved in the two electrode reactions.

For practical measurements, the electrode in question is connected to the positive terminal of theelectrometer, while the standard hydrogen electrode is connected to the negative terminal.^[2]

Reversible electrode

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Standard reduction potential table

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Main article:Standard electrode potential (data page)

The larger the value of the standard reduction potential, the easier it is for the element to be reduced (gainelectrons); in other words, they are betteroxidizing agents.

For example, F₂ has a standard reduction potential of +2.87 V and Li⁺ has −3.05 V:

F
₂(g) + 2 e⁻⇌ 2 F⁻
= +2.87 V

Li⁺
+ e⁻⇌ Li(s) = −3.05 V

The highly positive standard reduction potential of F₂ means it is reduced easily and is therefore a good oxidizing agent. In contrast, the greatly negative standard reduction potential of Li⁺ indicates that it is not easily reduced. Instead, Li_(s) would rather undergo oxidation (hence it is a goodreducing agent).

Zn²⁺ has a standard reduction potential of −0.76 V and thus can be oxidized by any other electrode whose standard reduction potential is greater than −0.76 V (e.g., H⁺ (0 V), Cu²⁺ (0.34 V), F₂ (2.87 V)) and can bereduced by any electrode with standard reduction potential less than −0.76 V (e.g. H₂ (−2.23 V), Na⁺ (−2.71 V), Li⁺ (−3.05 V)).

In a galvanic cell, where aspontaneous redox reaction drives the cell to produce an electric potential,Gibbs free energy $\Delta G^{\ominus }$ must be negative, in accordance with the following equation:

\Delta G_{cell}^{\ominus }=-nFE_{cell}^{\ominus }

(unit: Joule = Coulomb × Volt)

wheren is number ofmoles of electrons per mole of products andF is theFaraday constant,~ 96 485 C/mol.

As such, the following rules apply:

E_{cell}^{\ominus }

> 0, then the process is spontaneous (galvanic cell):

\Delta G_{cell}^{\ominus }

< 0, and energy is liberated.

E_{cell}^{\ominus }

< 0, then the process is non-spontaneous (electrolytic cell):

\Delta G_{cell}^{\ominus }

> 0, and energy is consumed.

Thus in order to have a spontaneous reaction ( $\Delta G_{cell}^{\ominus }$ < 0), $E_{cell}^{\ominus }$ must be positive, where:

E_{cell}^{\ominus }=E_{cathode}^{\ominus }-E_{anode}^{\ominus }

where $E_{cathode}^{\ominus }$ is the standard potential at the cathode (called as standard cathodic potential or standard reduction potential and $E_{anode}^{\ominus }$ is the standard potential at the anode (called as standard anodic potential or standard oxidation potential) as given in thetable of standard electrode potential.