Electrochemistry
Thermodynamics at the electrode
Learning objectives
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You will be able to:
Identify main components of an electrochemical cell
Write shorthand description of electrochemical cell
Calculate cell voltage using standard reduction potentials
Apply Nernst equation to determine free energy change
Apply Nernst equation to determine pH
Calculate K from electrode potentials
Calculate amount of material deposited in electrolysis
Energy in or energy out
 Galvanic (or voltaic) cell relies on
spontaneous process to generate a potential
capable of performing work – energy out
 Electrolytic cell performs chemical reactions
through application of a potential – energy in
Redox Review
 Oxidation is...
 Loss of electrons
 Reduction is...
 Gain of electrons
 Oxidizing agents oxidize and are reduced
 Reducing agents reduce and are oxidized
Redox at the heart of the matter
 Zn displaces Cu from CuSO4(aq)
 In direct contact the enthalpy of reaction is
dispersed as heat, and no useful work is done
 Redox process:
 Zn is the reducing agent
 Cu2+ is the oxidizing agent
2
Zn( s)  Zn (aq)  2e
2
Cu (aq)  2e  Cu( s)
Separating the combatants
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Each metal in touch with a solution of its own ions
External circuit carries electrons transferred during the redox process
A “salt bridge” containing neutral ions completes the internal circuit.
With no current flowing, a potential develops – the potential for work
Unlike the reaction in the beaker, the energy released by the reaction in
the cell can perform useful work – like lighting a bulb
Labelling the parts
Odes to a galvanic cell
 Cathode
 Where reduction occurs
 Where electrons are
consumed
 Where positive ions
migrate to
 Has positive sign
 Anode
 Where oxidation occurs
 Where electrons are
generated
 Where negative ions
migrate to
 Has negative sign
The role of inert electrodes
 Not all cells start with elements as the redox
agents
 Consider the cell
Fe( s)  2 Fe3 (aq)  3Fe2 (aq)
 Fe can be the anode but Fe3+ cannot be the
cathode.
 Use the Fe3+ ions in solution as the
“cathode” with an inert metal such as Pt
Anode
Catho
de
Oxidati
on
Reduct
ion
Cell notation
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Anode on left, cathode on right
Electrons flow from left to right
Oxidation on left, reduction on right
Single vertical = electrode/electrolyte boundary
Double vertical = salt bridge
Anode:
Zn →Zn2+
+ 2e
Cathode:
Cu2+ + 2e
→Cu
Vertical │denotes different phase
 Fe(s)│Fe2+(aq)║Fe3+(aq),Fe2+(aq)│Pt(s)
 Cu(s)│Cu2+(aq)║Cl2(g)│Cl-(aq)│C(s)
Connections: cell potential and free
energy
 The cell in open circuit generates an
electromotive force (emf) or potential or
voltage. This is the potential to perform
work
 Energy is charge moving under applied
voltage
1J  1C 1V
Relating free energy and cell
potential
 The Faraday:
F = 96 485 C/mol e
G  nFE
Standard conditions (1 M, 1 atm, 25°C)
G  nFE
Standard Reduction Potentials
 The total cell potential is the sum of the potentials
for the two half reactions at each electrode
 Ecell = Ecath + Ean
 From the cell voltage we cannot determine the
values of either – we must know one to get the
other
 Enter the standard hydrogen electrode (SHE)
 All potentials are referenced to the SHE (=0 V)
Unpacking the SHE
 The SHE consists of a Pt electrode in contact with
H2(g) at 1 atm in a solution of 1 M H+(aq).
 The voltage of this half-cell is defined to be 0 V
 An experimental cell containing the SHE half-cell
with other half-cell gives voltages which are the
standard potentials for those half-cells
Ecell = 0 + Ehalf-cell
Zinc half-cell with SHE
 Cell measures 0.76 V
 Standard potential for Zn(s) = Zn2+(aq) + 2e = 0.76
V
Where there is no SHE
 In this cell there is no SHE and the
measured voltage is 1.10 V
2
2
Zn Zn (aq) C u (aq) Cu
2
2
Zn( s)  Cu (aq)  Zn (aq)  Cu( s)
2
Zn( s)  Zn (aq)  2e, E  0.76V
2
o
Cu (aq)  2e  Cu( s), E  0.34V
o
Standard reduction potentials
 Any half reaction can be written in two ways:
 Oxidation:
M = M+ + e (+V)
 Reduction:
M+ + e = M (-V)
 Listed potentials are standard reduction
potentials
Applying standard reduction
potentials
 Consider the reaction
Zn( s)  2 Ag  (aq)  Zn 2 (aq)  2 Ag ( s)
 What is the cell potential?
 The half reactions are:
2

Zn
(
s
)

Zn
(aq)  2e
Ag (aq)  e  Ag ( s)
 E° = 0.80 V – (-0.76 V) = 1.56 V
 NOTE: Although there are 2 moles of Ag
reduced for each mole of Zn oxidized, we do not
multiply the potential by 2.
Extensive v intensive
 Free energy is extensive property so need to
multiply by no of moles involved
G  nFE
 But to convert to E we need to divide by no of
electrons involved
E   G

 E is an intensive property

nF
The Nernst equation
 Working in nonstandard conditions

G  G  RT ln Q

 nFE  nFE  RT ln Q

E  E  RT

nF
ln Q
E  E  0.0592 log Q
n
Electrode potentials and pH
 For the cell reaction

H 2 ( g )  2H (aq)  2e
 The Nernst equation
EH
2 2 H
EH

E
2 2 H


H 2 2 H 
0.06V

n
 
 2 

H
 log


pH 2 

  
0.06V

log H 
n
2
 Half-cell potential is proportional to pH
The pH meter is an electrochemical cell
 Overall cell potential is proportional to pH
Ecell  0.06V  pH   Eref
pH 
Ecell  Eref
0.06V
 In practice, a hydrogen electrode is
impractical
Calomel reference electrodes
 The potential of the calomel electrode is known vs
the SHE. This is used as the reference electrode
in the measurement of pH
Hg 2Cl2 ( s)  2e  2Hg (l )  2Cl

 The other electrode in a pH probe is a glass
electrode which has a Ag wire coated with AgCl
dipped in HCl(aq). A thin membrane separates the
HCl from the test solution
Cell potentials and equilibrium
G  nFE
 Lest we forget…
 So then

G   RT ln K

nFE   RT ln K
 and E   RT
2.303RT
ln K 
log 10 K
nF
nF
Cell potential a convenient way to
measure K
Many pathways to one ending
 Measurement of K from different
experiments
c
d
 Concentration data
C  D
a
b
A B
 Thermochemical data
 Electrochemical data
G    RT ln K

nFE   RT ln K
Batteries
 The most important application of galvanic
cells
 Several factors influence the choice of
materials
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Voltage
Weight
Capacity
Current density
Rechargeability
Running in reverse
 Recharging a battery requires to run the
process in reverse by applying a voltage
 In principle any reaction can be reversed
 In practice it will depend upon many factors
 Reversibility depends on kinetics and not
thermodynamics
 Cell reactions that involve minimal structural
rearrangement will be the easiest to reverse
Lithium batteries
 Lightweight (Molar mass Li = 6.94 g)
 High voltage
 Reversible process
Fuel cells – a battery with a
difference
 Reactants are not contained within a sealed
container but are supplied from outside
sources
anode : 2 H 2 ( g )  4OH  (aq)  4H 2O(l )  4e
cathode : O2 ( g )  2 H 2O(l )  4e  4OH  (aq)
overall : 2H 2 ( g )  O2 ( g )  2H 2O(l )
Store up not treasures on earth
where moth and rust…
 An electrochemical mechanism for corrosion of iron. The metal and a
surface water droplet constitute a tiny galvanic cell in which iron is
oxidized to Fe2+ in a region of the surface (anode region) remote from
atmospheric O2, and O2 is reduced near the edge of the droplet at
another region of the surface (cathode region). Electrons flow from
anode to cathode through the metal, while ions flow through the water
droplet. Dissolved O2 oxidizes Fe2+ further to Fe3+ before it is deposited
as rust (Fe2O3·H2O).
Mechanisms
 Why does salt enhance rusting?
 Improves conductivity of electrolyte
 Standard reduction potentials indicate which
metals will “rust”
 Aluminium should corrode readily. It
doesn’t. Is thermodynamics wrong?
 No, the Al2O3 provides an impenetrable barrier
No greater gift than to give up your
life for your friend
 A layer of zinc protects iron from oxidation, even when the
zinc layer becomes scratched. The zinc (anode), iron
(cathode), and water droplet (electrolyte) constitute a tiny
galvanic cell. Oxygen is reduced at the cathode, and zinc is
oxidized at the anode, thus protecting the iron from
oxidation.
Electrolysis
 Electrolysis of a molten salt using inert electrodes
 Signs of electrodes:
 In electrolysis, anode is positive because electrons are removed
from it by the battery
 In a galvanic cell, the anode is negative because is supplies
electrons to the external circuit
Anode : 2Cl  (l )  Cl2 ( g )  2e
Cathode : 2 Na  (l )  2e  2 Na(l )


Overall : 2 Na (l )  2Cl (l )  2 Na(l )  Cl2 ( g )
Electrolysis in aqueous solutions – a
choice of process
 There are (potentially)
competing processes
in the electrolysis of an
aqueous solution
 Cathode Cathode : 2 Na  (l )  2e  2 Na(l )...E  2.71V
Cathode : 2H 2O(l )  2e  H 2 ( g )  2OH  (aq)...E  0.83V
 Anode
Anode : 2Cl  (l )  Cl2 ( g )  2e...E  1.36V
Anode : 2 H 2O(l )  O2 ( g )  4 H   4e...E  1.23V
Thermodynamics or kinetics?
 On the basis of thermodynamics we choose
the processes which are favoured
energetically
Anode : 2 H 2O(l )  O2 ( g )  4 H   4e...E  1.23V
Cathode : 2H 2O(l )  2e  H 2 ( g )  2OH  (aq)...E  0.83V
 But…chlorine is evolved at the anode
The role of overpotentials
 Thermodynamic quantities prevail only at
equilibrium – no current flowing
 When current flows, kinetic considerations
come into play
 Overpotential represents the additional
voltage that must be applied to drive the
process
 In the NaCl(aq) solution the overpotential for
evolution of oxygen is greater than that for
chlorine, and so chlorine is evolved
preferentially
 Overpotential will depend on the electrolyte
and electrode. By suitable choices,
overpotentials can be minimized but are never
eliminated
 The limiting process in electrolysis is usually
diffusion of the ions in the electrolyte (but not
always)
 Driving the cell at the least current will give
rise to the smallest overpotential
Electrolysis of water
 In aqueous solutions of
most salts or acids or
bases the products will
be O2 and H2
Cathode : 2H 2O(l )  2e  H 2 ( g )  2OH  (aq)...E  0.83V

Anode : 2 H 2O(l )  O2 ( g )  4 H  4e...E  1.23V
Quantitative aspects of electrolysis
 Quantitative analysis
uses the current
flowing as a measure
of the amount of
material
 Charge = current x
time
 Moles =
charge/Faraday