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© 2000-2015 Gérard P. Michon, Ph.D.    

Electrochemistry    

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 International Year 
 of Chemistry - 2011

Ionization,
Oxidation-Reduction
and Electrochemistry

 Coat-of-arns of 
 Joseph Priestley (1734-1804)  Coat-of-arms of 
 Luigi Galvani (1737-1798)  Coat-of-arms of Count 
 Alessandro Volta (1745-1827)

(2011-08-04)     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.

 Come back later, we're
 still working on this one...

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


(2011-08-04)     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 NaNO).

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


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

 Come back later, we're
 still working on this one...

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


(2003-10-11)     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 ]

Redox reactions are best described as transfers of electrons between chemical species:  An oxidizer (oxidant, or oxidizing agent) is "reduced" by gaining electrons; a reducer (reductant, or reducing agent) is "oxidized" by losing them.

Some Redox Half-Reactions

Potential  Df
(25°C, 1 atm)
Hydrofluoric acidF2  +  2 H+  +  2 e- ® 2 HF(aq) (+3.05 V)
Fluorine½ F2  +  e- ® F - (+2.866 V)
SulfateS2O8- -  +  2 e- ® 2 SO4- - (+2.010 V)
PeroxideH2O2  +  2 H+  +  2 e- ® 2 H2O (+1.77 V)
Gold (aurous)Au+  +  e- ® Au (+1.692 V)
PermanganateMnO4-  +  4 H +  +  3 e- ® MnO2  +  2 H2O (+1.679 V)
PermanganateMnO4-  +  8 H +  +  5 e- ® Mn++  +  4 H2O (+1.507 V)
Gold (auric)Au+++  +  3 e- ® Au (+1.498 V)
HypochloriteHClO  +  H +  +  2 e- ® Cl -  +  H2O (+1.490 V)
Chlorine½ Cl2  +  e-   ®   Cl- (+1.35827 V)
Oxygen ½ O2  +  2 H +  +  2 e- ® H2O (+1.229 V)
PlatiniumPt++  +  2 e- ® Pt (+1.188 V)
Bromine Br2(aq)  +  2 e- ® 2 Br- (+1.0873 V)
Bromine Br2(l)  +  2 e- ® 2 Br- (+1.066 V)
NitrateNO3-  +  4 H +  +  3 e- ® NO  +  2 H2O (+ 0.96 V)
NitrateNO3-  +  2 H +  +  e- ® NO2  +  H2O (+ 0.80 V)
Silver Ag+  +  e- ® Ag (+0.7996 V)
Mercury Hg2++  +  2 e- ® 2 Hg (+0.796 V)
Ferrous Ion Fe+++  +  e- ® Fe++ (+0.769 V)
PeroxideO2  +  2 H+  +  2 e- ® H2O2 (+0.68 V)
PermanganateMnO4-  +  2 H2O  +  3 e- ® MnO2  +  4 OH - (+0.59 V)
Iodine I  +  e- ® I - (+0.534 V)
Copper (cuprous Cu+  +  e- ® Cu (+0.518 V)
HydroxideO2  +  2 H2O  +  4 e- ® 4 OH - (+0.40 V)
Copper (cupric) Cu++  +  2 e- ® Cu (+0.3419 V)
Silver Chloride AgCl  +  e- ® Ag  +  Cl - (+0.2198 V)
Thiosulfate (S2O3)22-  +  2 e- ® 2 S2O32- (+0.08 V)
Silver Bromide AgBr  +  e- ® Ag  +  Br- (+0.07133 V)
Hydrogen H +  +  e- ® ½ H2 ( 0 V )
Iron (ferric) Fe+++  +  3 e- ® Fe (-0.04 V)
Methanoate
(or formate)
CO2  +  2 H +  +  2 e- ® HCOOH (-0.11 V)
Silver Iodide AgI  +  e- ® Ag  +  I - (-0.15224 V)
Ethanedioate
(or oxalate)
2 H2CO3 + 2 H + + 2 e- ® H2C2O4 + 2 H2O (-0.386 V)
Ethanedioate
(or oxalate)
2 CO2  +  2 H +  +  2 e- ® H2C2O4 (-0.43 V)
Iron (ferrous) Fe++  +  2 e- ® Fe (-0.44 V)
Thiosulfate 2 SO32- +  3 H2O  +  4 e- ® S2O32- +  6 OH- (-0.571 V)
Zinc Zn++  +  2 e- ® Zn (-0.7618 V)
Hydroxide H2O  +  e- ® ½ H2  +  OH- (-0.8277 V)
Aluminum Al+++  +  3 e- ® Al (-1.677 V)
Magnesium Mg++  +  2 e- ® Mg (-2.372 V)
Sodium Na+  +  e- ® Na (-2.7143 V)
Calcium Ca++  +  2 e- ® Ca (-2.868 V)
Strontium Sr++  +  2 e- ® Sr (-2.899 V)
Barium Ba++  +  2 e- ® Ba (-2.912 V)
Potassium K+  +  e- ® K (-2.931 V)
Lithium Li+  +  e- ® Li (-3.0401 V)

Each of the above half-reactions 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 relation holds 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) and the above bracketed constant is the Faraday constantF, the charge of a mole of electrons) 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)  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   =   DGo  +  n  RT    ln  [products]   )
vinculum
[reactants]
 
Df   =   Dfo  -  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

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