In electrochemistry, electrode potential (also called reversible potential difference or reversible potential difference of an electrode, and abbreviated E) is the potential difference of a half-reaction which occurs across a reversible cell made up of any electrode and a standard hydrogen electrode.

Different half-reactions attached through electrochemical cells give different potential readings. For example, on connection, the below half-reactions give a reading of 0.59V:

(1) Fe³⁺ + e^- -> Fe²⁺

(2) Cl₂ + 2e^- -> 2Cl^-

Whereas connecting half-reaction 1 (Fe³⁺/Fe²⁺) to I₂/2I^- results in a reading of 0.23V:

(1) Fe³⁺ + e^- -> Fe²⁺

(3) I₂ + 2e^- -> 2I^-

It is clear, then, that both half-reactions contribute to the potential measured. In order to isolate the potentials of individual half-reactions, measurements against hydrogen as the second half-cell are made:

(1) Fe³⁺ + e^- -> Fe²⁺

(4) 2H⁺ + 2e^- -> H₂

When a reading is taken with hydrogen as the second half-cell, the potential measured is termed a half-reaction's standard electrode potential. The above cell gives the standard electrode potential of the Fe³⁺/Fe²⁺ half-reaction as +0.77V. In this case a positive potential indicates that the equilibrium of the Fe³⁺/Fe²⁺ reaction lies towards Fe²⁺, meaning that it accepts electrons from the hydrogen half-reaction. A negative potential reading would indicate that the half-reaction donates electrons to hydrogen.

Once the standard electrode potentials of two substances are known, it is possible to predict the direction electrons will travel from one substance to another, as well as the potential measured between them. The more positive half-reaction will always accept electrons, whilst the less positive (not necessarily negative) reaction will donate electrons. The voltage measured between the half-reactions is equal to the difference in standard electrode potentials between the substances. For example given our original reaction:

(1) Fe³⁺ + e^- -> Fe²⁺ EP = 0.77 (See Table of standard electrode potentials)

(2) Cl₂ + 2e^- -> 2Cl^- EP = 1.36

Cl₂/Cl^- is clearly the most positive, so it will accept electrons, whilst Fe³⁺/Fe²⁺ donates electrons:

(1) Fe²⁺ -> Fe³⁺ + e^-

(2) Cl₂ + 2e^- -> 2Cl^-

Combining the two half-reactions brings the chemical formula:

(5) Cl₂ + 2e^- + 2Fe²⁺ -> 2Cl^- + 2Fe³⁺ + 2e^-

(In order to balance the number of electrons in both reactions, the Fe³⁺/Fe²⁺ half-reaction was doubled.) Simplifying the reaction by removing the electrons which appear on both sides gives:

(6) Cl₂ + 2Fe²⁺ -> 2Cl^- + 2Fe³⁺

The voltage that will be measured is the difference between the potentials i.e., 1.36 - 0.77, which equals 0.59V.

The use of standard electrode potentials is not restricted to electrochemical cells - the same predictions we made are equally true of reactions occurring naturally.

See also

• Nernst equation
• Electrochemical cell
• Electrochemical potential