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(2011年08月04日) Oxidation Number
A traditional fiction which can be most useful.

An increase in the oxidation number is an oxidation. Conversely, a decrease is a reduction.

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Wikipedia: Oxidation number
Oxidation Numbers & Redox Reactions by Gwen Sibert.

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(2011年08月04日) Salt Bridge
Often made with an inert electrolyte gelified with agar, in a glass tube.

Instead of gelification, porous plugs may be used at both ends of the tube (the idea is to allow electrical contact but prevent a transfer of ions).

Alternately, the tube can simply be replaced by strips of filter paper soaked with the inert electrolyte (this setup may have a substantial ohmic resistance but it's adequate for voltage measurements with negligible electric currents).

Traditional electrolytes for a salt bridge include potassium or sodium chloride (KCl or NaCl). Nitrates are also used (KNO₃ or NaNO₃ ).

For example, a salt bridge can be used to connect two solutions of the same salt at different concentrations surrounding electrodes of the same metal A voltage is then observed between the two electrodes (as explained next) which tends to cause a cureent in the direction that would reduce the difference between the concentrations.

Wikipedia: Salt bridge

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(2011年08月04日) Concentration Cells & Nernst Equation
The voltage difference caused by a difference in concentrations.

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Concentration cell | Nernst equation | Walther Nernst (1864-1941)

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(2003年10月11日) Redox Reactions
An oxidizer gains the electrons which a reductant loses.
(The reductant is oxidized, the oxidizer is reduced.)

Oxidation is loss (of electrons) reduction is gain (OIL RIG mnemonic).

A redox reaction transfers electrons from a reducer (reductant, or reducing agent) to an oxidizer (oxidant, or oxidizing agent). Said reducer is oxidized by losing electrons. The oxidizer is reduced by gaining them.

Every half-reaction above is written as the reduction of an oxidizer, but the reverse direction (the oxidation of the reducer on the right-hand side) is more common for the reactions with a low redox potential (listed in volts V): In a complete redox reaction, a reduction occurs as written above only if a balancing oxidation with a lower redox potential occurs in the reverse direction. For example, the nitrate ion has a higher potential than the cupric ion and nitric acid may thus oxidize copper metal. (The opposite is true between hydrogen and cupric ions, so an ordinary acid can't oxidize copper.)

2 NO₃^- + 4 H ⁺ + Cu ® 2 NO₂ + 2 H₂O + Cu⁺⁺

In a balanced redox reaction, the difference Df between the potentials of both half-reactions is simply the change in free enthalpy DG (G = H-TS) per unit of electric charge transferred. If n moles of electrons are involved, this translates into n moles of electronvolts in DG for each volt in Df. Therefore:

DG = -n F Df = -n Df (96485 J/V) = -n Df (23.06 kcal/V)

A joule per volt (J/V) is a coulomb (C). The bracketed factor corresponds to a mole of electrons (Faraday's constant, F ) in two different units.

Only the Df (or DG) of an actual redox reaction has a physical meaning, while all the half-reactions are convenient fictions whose redox potentials are defined within an additive constant, which is conventionally set to 0 V for hydrogen. [Another convention is used for the related "Oxydo-Reduction Potential" (ORP) measured directly for aqueous solutions, which lets 1 V be the ORP of chlorine.]

The standard redox potential (Df^o ) tabulated for a normal pressure of 1 atm (101325 Pa) at 25°C (77°F) is understood for unit (1M) concentrations of both reactants and products, otherwise the Nernst equation is used:

DG = DG^o + n RT ln ( [products] )

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[reactants]

Df = Df^o - RT ln ( [products] )

vinculum vinculum

F [reactants]

Therefore, even if the comparison of standard redox potentials seems to imply that a reaction does not occur, what actually evolves is an equilibrium where the concentration of "products" is small, or even utterly negligible...

Note that RT / F = kT / e is equal to 25.6926 mV (at 25°C = 298.15 K). This is precisely the thermal voltage which appears in Shockley's Ideal Diode Equation and elsewhere...

Wikipedia: Standard electrode potential (table)
Standard Reduction Potentials at 25°C by Ken Costello (Chemistryland)
Standard Reduction Potential (UC Davis ChemWiki)
Standard Reduction Potentials (Reference Tables for Chemistry) by Harry Clark
Standard Reduction Potentials & Temperature Coefficients in Water at 298.15 K by Sreven G. Brascht (1988)
Calculating the standard half-cell reduction potential (Chemical Forums) by Kimi85 & Borek