Modern inorganic chemistry

100 ACIDS AND BASES: OXIDATION AND REDUCTION

Table 4.3

REDOX POTENTIALS FOR ION-ION SYSTEMS (ACID SOLUTIONS)

£(V)

Sn^4 + (aq) +

Increasing
oxidising
power

yI 2 (s) -f e~

Fe^3 + (aq) +

iBr 2 (l) + e~

IOj(aq) -f (

02 (g) + 4H

Cr 2 O?"(aq)

iCl 2 (g) + e'

MnO^Caq)

iF 2 (g) 4- e~

2e" -^Sn^2 ^(aq)

-> I~(aq)

e~ -> Fe^2 + (aq)

-> Br"(aq)

jH 3 O^ ^ 5e"" -+

JO + 4e" -^6H

-f 14H 3 O+ -f 6

• -* CT(aq)
  f 8H 3 O+ -h 5^~
  
  
  • F-(aq)
  
  
  
  

iI 2 (s) + 9H 2 O

20

r -*2Cr3+(aq)

H-21H 2 O

-M2+(aq)

+ 12H 2 O

+ 0.15

+ 0.54

+ 0.76

Increasing -(-1.07

reducing 4-1.19

power +1.23

+ 1.33

+ 1.36

+ 1-52

+ 2.80

THE EFFECT OF CONCENTRATION AND

TEMPERATURE ON POTENTIALS

Changes in ion concentration and temperature influence redox
potentials by affecting the equilibrium

M(s) ^ M*+(aq) + ne~

The change in the redox potential is given quantitatively by the
Nernst equation :

RT

where £ is the actual electrode potential, E^ is the standard electrode
potential, R the gas constant, Tthe temperature in K, F the Faraday
constant and n the number of electrons.
Substituting for R and F and for a temperature of 298 K this
equation approximates to :

The redox (electrode) potential for ion-ion redox systems at any
concentration and temperature is given by the Nernst equation in
the form

^ RT rOxidised state"!

+ ~nF ge [TedSoedlt ate]