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Electrode potentials of successive elementary half-reactions cannot be directly added. However, the corresponding Gibbs free energy changes (∆G°) must satisfy ∆G° = – z FE°, where z electrons are transferred, and the Faraday constant F is the conversion factor describing Coulombs transferred per mole electrons. Those Gibbs free energy ...
The values below are standard apparent reduction potentials (E°') for electro-biochemical half-reactions measured at 25 °C, 1 atmosphere and a pH of 7 in aqueous solution. [ 1 ] [ 2 ] The actual physiological potential depends on the ratio of the reduced ( Red ) and oxidized ( Ox ) forms according to the Nernst equation and the thermal voltage .
There are two classes of redox reactions: Electron-transfer – Only one (usually) electron flows from the atom, ion, or molecule being oxidized to the atom, ion, or molecule that is reduced. This type of redox reaction is often discussed in terms of redox couples and electrode potentials.
In aqueous solutions, redox potential is a measure of the tendency of the solution to either gain or lose electrons in a reaction. A solution with a higher (more positive) reduction potential than some other molecule will have a tendency to gain electrons from this molecule (i.e. to be reduced by oxidizing this other molecule) and a solution with a lower (more negative) reduction potential ...
An atom (or ion) whose oxidation number increases in a redox reaction is said to be oxidized (and is called a reducing agent). It is accomplished by loss of one or more electrons. The atom whose oxidation number decreases gains (receives) one or more electrons and is said to be reduced. This relation can be remembered by the following mnemonics.
Going from the bottom to the top of the table the metals: increase in reactivity; lose electrons more readily to form positive ions; corrode or tarnish more readily; require more energy (and different methods) to be isolated from their compounds; become stronger reducing agents (electron donors).
Depiction of common redox reactions in the environment. Adapted from figures by Zhang [1] and Gorny. [2] Redox pairs are listed with the oxidizer (electron acceptor) in red and the reducer (electron donator) in black. Relative favorability of redox reactions in marine sediments based on energy.
In general, these redox balances (the one-line balance or each half-reaction) need to be checked for the ionic and electron charge sums on both sides of the equation being indeed equal. If they are not equal, suitable ions are added to balance the charges and the non-redox elemental balance.