Standard Reduction Potentials, 25°C,
H+ = acidic solution, OH– = basic solution
E°red (Volts)
Reduction half reaction
E°red (V)
Reduction half reaction
E°red (V)
Li+ + e– —> Li(s)**
Rb+ + e– <–––> Rb(s)
K+ + e– —> K(s)
–3.045
–2.98
–2.925
MnO4– + e– —> MnO4–2
0.564
MnO4– + 2 H2O(l) + 3 e– —>
! MnO2(s) + 4 OH–
0.588
Ca+2 + 2 e– —> Ca(s)
–2.866
BrO3– + 3 H2O(l)+ 6 e– –> Br– + 6 OH– 0.61
Na+ + e– —> Na(s)
–2.714
ClO2– + H2O + 2 e– —> ClO– + OH–
0.66
Mg+2 + 2 e– —> Mg(s)
–2.363
O2(g) + 2 H+ + 2 e– —> H2O2(aq)
0.682
Al+3 + 3 e– —> Al(s)
–1.662
Fe+3 + e– —> Fe+2
0.771
Mn+2 + 2 e– —> Mn(s)
–1.185
Ag+ + e– —> Ag(s)
0.799
SO4–2 + H2O(l) + 2 e– —>
!
SO3–2 + 2 OH–
–0.93
O2(g) + 2 H2O(l) + 2 e– —>
! H2O2(aq) + 2 OH–
0.88
2 H2O(l)+ 2 e– –> H2(g) + 2 OH– –0.828
ClO– + H2O(l) + 2 e– –> Cl– + 2 OH–
0.89
Zn+2 + 2 e– —> Zn(s)
–0.763
0.96
Cr+3 + 3 e– —> Cr(s)
–0.744
NO3– + 4 H+ + 3 e– —>
! NO(g) + 2 H2O(l)
Fe+2 + 2 e– —> Fe(s)
–0.440
Br2(aq) + 2 e– —> 2 Br–
1.087
Cr+3 + e– —> Cr+2
–0.408
Co+2 + 2 e– –––> Co(s)
–0.28
Ni+2 + 2 e– —> Ni(s)
–0.250
Sn+2 + 2 e– —> Sn(s)
–0.136
CrO4–2 + 4 H2O(l) + 3 e– —>
! Cr(OH)3(s) + 5 OH–
ClO4– + 2 H+ + 2 e– -> ClO3– + H2O(l) 1.19
2 IO3– + 12 H+ + 10 e– —>
! I2(s) + 6 H2O(l)
1.195
O2(g) + 4 H+ + 4 e– –> 2H2O(l)
1.229
–0.13
MnO2(s) + 4 H+ + 2 e– —>
! Mn+2 + 2 H2O(l)
1.23
Pb+2 + 2 e– —> Pb(s)
–0.126
1.33
MnO2(s) + 2H2O(l) + 2 e– —>
! Mn(OH)2(s) + 2OH–
–0.05
Cr2O7–2 + 14 H+ + 6 e– —>
! 2 Cr+3 + 7 H2O(l)
Cl2(g) + 2 e– —> 2 Cl–
1.360
2H+ + 2 e– —> H2(g)
0.000
Au+3 + 3 e– —> Au(s)
1.498
NO3– + H2O(l) + 2 e– —>
! NO2– + 2 OH–
0.01
MnO4– + 8 H+ + 5 e– —>
! Mn+2 + 4H2O(l)
1.51
Sn+4 + 2 e– —> Sn+2
0.15
1.63
Cu+2 + e– —> Cu+
0.153
2 HClO + 2 H++ 2 e– —>
! Cl2(g) + 2 H2O(l)
SO4–2 + 4 H+ + 2 e– —>
! H2SO3(aq) + H2O(l)
0.172
H2O2(aq) + 2 H+ + 2 e– —>2 H2O(l)
1.776
••F2(g) + 2 e– —> 2 F–
2.87
ClO3– + H2O(l) + 2 e– —>
! ClO2– + 2 OH–
0.33
ClO4– + H2O(l) + 2 e– —>
! ClO3– + 2 OH–
0.36
Cu+2 + 2 e– —> Cu(s)
0.337
O2(g) + 2 H2O(l) + 4 e– —> 4 OH–
0.401
I2(s) + 2 e– —> 2 I–
0.536
Products are Reducing agents.
**Products close to Li (the top) are easily Oxidized,
! so are good Reducing agents.
! (easily react in opposite direction)
! Reactants are Oxidizing agents.
•• Reactants close to F2 (bottom) are easily Reduced,
! so are good Oxidizing agents.
! (easily react as written )
Summary: Electrochemistry
Voltaic Cells (Spontaneous)
Check the relative position of the two reduction half
reactions on the E°red chart.
The reaction above is the Oxidation and would be
written in the Opposite direction.
!
E°ox = – E°red on chart (switch sign)
The reaction below is the Reduction and would be
written the right way.
!
E°red = as written on chart
Oxidation = anode (both vowels)
= the – (negative) electrode because negative
electrons are produced there
Reduction = cathode (both consonants)
= the + (positive) electrode because it seems to
attract the negative electrons
Electrons flow in wire from the oxidation (anode, –)
to the reduction (cathode, +)
Cell potential, (or voltage, emf, E°cell)
!
!
!
E°cell = E°red + E°ox
(Remember to switch sign for E°ox)
Measured in Volts = electrical “force” or
! “pressure”
!
1 Volt = 1 Joule
!
Coulomb
A Coulomb is a unit of electric charge:
!
1 Coulomb (C) = amp x second
!
or!
1 amp = 1 C !
!
sec
Amp = unit of electric current
Determining Spontaneous reactions
A redox reaction occurs spontaneously as
written if E°cell (E°red + E°ox) is positive.
If E°cell is negative, the reaction would not
occur spontaneously. The reverse reaction
would occur spontaneously instead.
The following calculations are almost always
done at 25°C, 298 K.
emf and Free energy (∆G again!)
Standard conditions (1 M solutions, 1 atm)
!
∆G° = – n F E°cell
Any concentrations
!
∆G = – n F Ecell
Where 1 F (Faraday) = 96500 C (coulombs)
!
= the charge of 1 mole of electrons
!
n = number of electrons transferred in
!
!
the balanced reaction
emf and Equilibrium constant
!
(Voltage for 1 M, 1 atm is related to K)
!
!
E°cell =
0.0591 log K !
n
emf and Non-standard conditions
*The Nernst Equation:
(Voltage for any concentrations, not 1 M, 1 atm)
!
!
Ecell = E°cell – 0.0591 log Q !
n
Electrolytic Cells: (Non-spontaneous) Electricity is added so the reaction goes in the
! “wrong” direction. Electrons are “pushed” from the anode to the cathode.
Anode = Oxidation = + electrode (because electrons are removed from it),
! (attracts anions, negative ions)
Cathode = Reduction = – electrode ( because electrons are being pushed into it)!
! (attracts cations, positive ions)
In aqueous solutions, the combination that gives the smallest negative voltage occurs.
Electroplating: First Calculate # of Coulombs = amps x time in seconds
Then Grams = # Coulombs (1 mol e– ) (1 mol Metal ) (F W of Metal, g)
! !
96500 C!
n mol e–!
1 mol Metal