Talk:Standard electrode potential/Temp

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The standard electrode potential, abbreviated E^o is the potential, under standard conditions of an electrochemical cell in a cell in which the anode is the standard hydrogen electrode and the cathode is the half cell of interest.

The standard conditions for measurement are:

all components in solution present at an activity of 1 mol/L (this implies a pH of zero if hydrogen ions are involved in the reaction);
all gaseous components present at a partial pressure of 1 bar;
a temperature of 25 °C.

The potentials of electrodes under non-standard conditions cane be calculated using the Nernst equation (see also below).

Standard electrode potentials may be either measured directly or calculated from other thermodynamic data. A selection of values is given on the accompanying data page: by convention, standard electrode potentials are listed for the reduction reaction, with the most negative potentials appearing first.

1 Relationship to cell potentials
2 Effect of concentration
3 Measurement of standard electrode potentials
4 Relationship to other thermodynamic quantities
5 External link

[edit] Relationship to cell potentials

The potential of an isolated electrode is undefined: only the potential difference relative to another electrode can be measured. Where this second electrode is a standard hydrogen electrode, the potential difference between the electrodes will be equal to the standard electrode potential. However, the standard hydrogen electrode is not particularly practical to use, even for laboratory measurements, and all but irrelevant in wider applications. To illustrate the application of standard electrode potentials, the case of a Galvanic cell will be examined.

A Galvanic cell consists of a copper electrode immersed in a solution of copper ions and a zinc electrode immesed in a solution of zinc ions, with a salt bridge connecting the two solutions so that they are in electrical contact without being able to mix. The two half-reactions and the corresponding standard electrode potentials are:

Zn²⁺ + 2e⁻ ⇌ Zn E^o = −0.76 V

Cu²⁺ + 2e⁻ ⇌ Cu E^o = +0.34 V

The difference between the two potentials is 1.10 V, and this is, in effect, the cell potential assuming that the two soutions have an activity of 1 mol/L: this quantity is also known as the standard electromotive force of the cell.

The zinc electrode, which has the more negative potential, is the negative pole of the cell, that is to say the anode, the electrode at which an oxidation reaction takes place. The copper electrode is the cathode, the positive pole of the cell at which a reduction reaction takes place. The overall reaction when the cell delivers an electric current is:

Zn + Cu²⁺ → Zn²⁺ + Cu

Zinc is oxidized to Zn²⁺ ions, and Cu²⁺ ions are reduced to copper metal.

[edit] Effect of concentration

The potential of an electrode depends on the concentration of the componants in solution. The dependence is given by the Nernst equation:

$E = E^\circ - {{RT}\over{nF}}\ln{{a_{\rm{red}}}\over{a_{\rm{ox}}}}$

where R is the universal gas constant, R =N_Ak_B = 8.3145 J.K⁻¹.mol⁻¹, T is the thermodynamic temperature, n is the number of electrons transferred during the reaction, F is the Faraday constant, F = N_Ae = 96.485 kC/mol, a_red is the activity of the reduced species and a_ox is the activity of the oxidized species. Activities will be replaced by concentrations for the examples.

Hence a copper electrode immersed in a solution of 0.01 mol/L copper(II) ions has a potential, relative to the standard hydrogen electrode of

$E = E^\circ - {{RT}\over{nF}}\ln{{a_{\rm{red}}}\over{a_{\rm{ox}}}} = +0.34 - 0.0128\ln{1\over{0.01}} = +0.28 \rm{V}$

Many half-reactions involve hydrogen ions: this is particularly the case for reactions of oxoanions. The potentials of these half cells will depend on the pH. The following version of the Nernst equation is often used in such cases:

$E = E^\circ - {{0.0591\nu}\over{n}}\rm{pH}$

where ν is the stoichiometric coefficient of the hydrogen ion in the half-reaction.

[edit] Measurement of standard electrode potentials

[edit] Relationship to other thermodynamic quantities

Not all electrode potentials can be measured directly, even after correcting for non-standard concentrations. The potential of the sodium electrode, Na⁺ + e⁻ ⇌ Na, is listed as E^o = −2.71 V, but an electrode constructed of sodium metal would not survive immersion in an aqueous solution of sodium ions whatever the concentration.

At the same time, the knowledge of electrode potentials allows the prediction of the direction of a cell reaction, in the same way that the knowledge of the Gibbs free energy change allows the prediction of the direction of a spontaneous chemical process. The two quantities are indeed related, by the equation:

Δ_rG = −nFΔE

where n is the number of electrons transferred during the reaction, F is the Faraday constant, F = N_Ae = 96.485 kC/mol, and ΔE is the cell potential.

Hence, the standard Gibbs free energy change of reaction for the reaction of the Galvanic cell can be calculated from the standard electrode potentials: in this case, n = 2:

Δ_rG^o = −2×96.485×1.10 = −212 kJ/mol