Electrochemistry involves redox systems.
Therefore species amenable for analysis primarily
are involved in red-ox systems.
Electro-Analytical Chemistry
Terminology:
Red-ox reaction – one species undergoes a loss
of electrons another gains electrons.
e-
Preliminaries
Fe(III) + V(II)  Fe(II) + V(III)
oxidant reductant
Further the reaction is carried out so that oxidation and
reduction occur at different locations - electrodes - in
electrochemical set-ups.
Fe(III) + V(II)  Fe(II) + V(III)
Two electrodes when coupled constitute an
electrochemical cell.
The electron movement; q = quantity of charge,
i = rate of movement of charge. The electrical
potentials of the electrodes, Eel and the potential
difference of the electrodes Ecell that are involved and
can be measured.
Two types of cells are studied:
a. Galvanic (Voltaic) cell: Galvanic cell uses
spontaneous red-ox chemical reactions to produce
electrical energy; that would result in a flow of electrons.
G < 0.
b. Electrolytic cell: An electrolytic cell decomposes
chemical compounds by red-ox processes using
electrical energy from an outside source - electrolysis.
This is an energy demanding process –
non-spontaneous. G > 0.
In any type of cell: Anode – oxidation occurs
Cathode – reduction occurs
Galvanic Cell Cd + 2 AgCl(s)  Cd+2 + 2Ag + 2Cl-
Analytical
Chemistry
Qualitative Analysis
property characteristic
to analyte
i  rate of reaction
Ecell  G of reaction
Eel
G = -nFEcell
Quantitative analysis
property related to
concentration of analyte
i, q
i
Oxidation
A
reduction
q = quantity of current
=it
=nF
n = #moles electrons
1 mol e- = 96485C = 1F
1
Cd + 2Ag+  Cd+2 + 2Ag
V
Half Cell Reactions:
Cd (s) + 2 AgCl (s)  Cd+2 (aq) + 2Ag (s) + 2Cl- (aq)
anode; oxidation
Cd(s)  Cd+2(aq) + 2e
2AgCl(s) + 2e  2Ag(s) + 2Cl-(aq) cathode; reduction
The above two equations (half reactions) involve a
physical transfer of electrons (Faradaic Process)
By convention: anode (oxidation half reaction) - left
Notation: Cd|Cd+2(aq)||Ag+1(aq)|Ag
phase boundary
i, ampere; A
R resistance, 
E, volts, V
i = q/t
t, sec F = 96485 coulombs/mole
96485 C/mol
q = charge; C
1 C/sec = 1 A (quite a large charge flow rate)
Work (J) = Free energy from reaction = q E
Coulombs volts
Construction of electrodes:
The electrode is made out of species involved in
the half reaction. If a metal is not involved, Pt
provides electrical connectivity.
Convention
Negative terminal
Left – Black
ANODE
Convention
Positive terminal
Right – Red
CATHODE
G = -nF Ecell
i = E/R
Standard Hydrogen Electrode: SHE
ESHE = 0.00V @ 250C
by definition.
Standard Hydrogen Electrode: SHE
In any cell, when electrons move (current flows) between
electrodes the potential difference drops to zero.
2
Ecell = Difference in electrical potentials of the two
electrodes.
Measured Ecell positive
if anode connected to
negative terminal, …
Positive terminal
Red
pH meter is very close
to an ideal voltmeter.
Impedance → 
Draws negligible current.
Negative terminal
- reference electrode
slot.
- BNC – outer connector
Negative terminal
COM - Black
Electrodes – (Equilibrium) Electrode Potentials:
M(s)
M+
M+ (aq)
e
if 


M(s)




M+(aq) + e
Metal acquires a negative
potential w.r.t. solution
Bayonet Neill Concelman connector
The equilibrium set up at the electrodes is not the
conventional chemical equilibrium, rather it is an
electrochemical equilibrium.
The potential difference developed across the
interface also controls the equilibrium position
of the half reaction.
This type of equilibrium is also referred to as
frustrated equilibrium.
M(s)
M+
Faradaic processes: an interfacial
phenomenon.
M+ (aq)
M(s)




M+(aq) + e
e
Depending on the position of equilibrium
the metal acquires a negative/positive
potential w.r.t. solution.
CM+ (aq) is one determinant of the
position of equilibrium.
3
How electrode potentials develops at an electrode:
Electrode Interface :
Ex: Cu/Cu+2(aq)
Electro-chemical potentials of the species Cu+2 ion is
not the same in the two phases, that has not attained
equilibrium. Natural tendency is to equalize the
electrochemical potentials of the two species.
In order to achieve such a state, ion concentration
must change in the two phases at the interface.
Cu(s)
Cu+2(aq) + 2e
If the electrochemical equilibrium shifts to right;
the excess e- would remain on the Cu metal making
it more negative w. r. t. solution and vice versa.
The higher the tendency for the oxidation process
to occur, the higher would be the electron density on
the metal.
For cases where the oxidation is dominant it’s
electrode potential is more negative.
The electrode potentials of red-ox systems are
tabulated, relative to the standard hydrogen
electrode (SHE).
IHL
OHL
Bulk
Standard Electrode Potentials, Eo
Electrode reaction
Li+ + e− → Li
K+ + e − → K
.
.
.
Eº /V
.
.
Electrode reaction
− 3.045
AgI + e− → Ag + I−
.
− 2.925
Sn2+ + 2e− → Sn
.
.
.
.
.
− 2.923
2H+ + 2e− → H2
Ba2+ + 2e− → Ba
.
.
.
.
− 2.92
AgBr + e− → Ag + Br−
.
.
Cs+ + e− → Cs
Al3+ + 3e− → Al
.
.
.
Eº /V
.
.
.
.
.
.
.
.
.
.
.
.
− 0.152 2
.
− 0.136
.
0 exactly
.
.
+ 0.071 1
.
− 1.67
I−3 + 2e− → 3I−
Zn2+ + 2e− → Zn
.
.
.
.
− 0.762 6
Fe3+ + e− → Fe2+
Ga3+ + 3e− → Ga
.
.
.
.
− 0.529
Hg22+ + 2e− → 2Hg
Fe2+ + 2e− → Fe
.
.
.
.
− 0.44
Ag+ + e− → Ag
.
.
.
.
.
.
+ 0.799 1
Cr3+ + e− → Cr2+
.
.
.
.
− 0.424
Hg2+ +2e− → Hg22+
.
.
.
.
.
+ 0.911 0
Cd2+ + 2e− → Cd
.
.
.
.
− 0.042 5
Pd2+ + 2e− → Pd
.
.
.
.
.
.
V3+ + e− → V2+
Ni2+ + e− → Ni
.
.
.
.
.
.
.
.
.
.
− 0.255
Cl2 + 2e− → 2Cl−
.
.
.
.
− 0.257
Au3+ + 3e− → Au
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
+ 0.536
.
.
+ 0.771
.
.
.
.
.
+ 0.796 0
+ 0.915
.
+ 1.358 3
+ 1.52
Equilibrium Electrode potential:
The magnitude of the electrode potential depends on
the excess charge that exists above that of the
metal alone.
If the ‘excess of (negative charge) electrons’ is
present in the metal of the electrode the potential of
the electrode is negative, with the energy of the
electrons in the electrode high and vice versa.
An external power supply is capable of forcing an
excess (or a depletion electrons) from the metal of
the electrode, could lead to a non-equilibrium
condition at the interface - electrolysis.
In electrolysis, there is a net reaction forced by the
applied power source – later topic.
4
Absolute individual electrode potentials cannot be
measured, only potential differences; i.e. only relative
values can be measured.
Electrode potentials are measured against a standard
electrode and tabulated ; electrode potential of the
standard hydrogen electrode (SHE) is defined as
0.00V at 250C.
Potentiometer
Ecell = ERHS – ELHS
measures the potential
difference between
Ecell = Etest – ELHS
= 0, definition
two electrodes.
Ecell = Etest – ESHE
Ecell = Etest for this set up.
Potentiometer
-
+
Ecell = Etest = Eel
Etest or Eel
SHE|| Test electrode
Pt|H2(g) (p=1atm)|H+(aq) (a=1)
Hg(EDTA)-2(aq, 0.005M)+2e→Hg(l)+EDTA-4 (aq,0.015M)
Potentiometer
pH2=1atm
SHE
Measured Ecell for this set up = Electrode potential
of test electrode
Standard Electrodes and Standard Electrode
Potential Eoel:
When the activity (~concentration) of all species
involved in the half reaction are unity – standard
electrode.
Potential of such electrodes are defined as it’s standard
electrode potential.
Potentiometer
aH=1
-
+
SHE|| (Test) Standard electrode
High Impedance Voltmeter
Electrode reaction
Li+ + e− → Li
K+ + e − → K
.
.
.
Eº /V
.
.
Electrode reaction
− 3.045
AgI + e− → Ag + I−
.
− 2.925
Sn2+ + 2e− → Sn
.
.
.
.
.
− 2.923
2H+ + 2e− → H2
Ba2+ + 2e− → Ba
.
.
.
.
− 2.92
AgBr + e− → Ag + Br−
.
.
Cs+ + e− → Cs
Al3+ + 3e− → Al
.
.
.
Eº /V
.
.
.
.
.
.
.
.
.
.
.
.
− 0.152 2
.
− 0.136
.
0 exactly
.
.
+ 0.071 1
.
− 1.67
I−3 + 2e− → 3I−
Zn2+ + 2e− → Zn
.
.
.
.
− 0.762 6
Fe3+ + e− → Fe2+
Ga3+ + 3e− → Ga
.
.
.
.
− 0.529
Hg22+ + 2e− → 2Hg
Fe2+ + 2e− → Fe
.
.
.
.
− 0.44
Ag+ + e− → Ag
.
.
.
.
.
.
+ 0.799 1
Cr3+ + e− → Cr2+
.
.
.
.
− 0.424
Hg2+ +2e− → Hg22+
.
.
.
.
.
+ 0.911 0
Cd2+ + 2e− → Cd
.
.
.
.
− 0.042 5
Pd2+ + 2e− → Pd
.
.
.
.
.
.
V3+ + e− → V2+
Ni2+ + e− → Ni
.
.
.
.
.
.
.
.
.
.
− 0.255
Cl2 + 2e− → 2Cl−
.
.
.
.
− 0.257
Au3+ + 3e− → Au
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
.
+ 0.536
.
.
+ 0.771
.
.
.
.
.
+ 0.796 0
+ 0.915
.
+ 1.358 3
+ 1.52
Note: Large and negative electrode potential means
less tendency for reduction, tendency is to oxidize.
Electrode potential, by convention is a
measure of the ability to undergo reduction.
5
E-
V>0
E+
Reduction ‘assumed’
To occur at ‘test electrode’
If so Eoel = positive
a=1
Hg(EDTA)-2(a=1),
EDTA(a=1)/Hg
Electrode potential is an interfacial phenomenon.
Separation of charges across metal/solution interface
brings about the potential difference between the
solution and metal.
The position of equilibrium governed by the
activities (concentrations) of species (by way of
reaction quotient Q) involved in the half reaction;
determines the electrode potential and the inherent
tendency to undergo reduction.
Calculation of Electrode Potential (Nernst equation):
Calculation of Electrode Potential:
By convention electrode potentials are expressed for
reduction reactions.
If potential is positive, all what it means is that the
reduction reaction of the ‘test electrode’ has a higher
propensity to happen compared to the reduction at
the standard hydrogen electrode.
Eel  Eel0 
RT
ln Qel
nF
@ 25o C; volts
R = 8.314 J/K mol
T = temperature, K
n= number of electrons involved in half reaction
F = 96485 C
Qel = reaction quotient in activity
pure liquids, solvents, solids; a=1
6
Calculation of Electrode Potential (Nernst equation):
Calculation of Electrode Potential:
Electrode potentials are calculated for the reduction
process as reduction potentials Write the half
reaction as a reduction reaction, balanced in mass
and charge. Write the expression for Q, determine
the # electrons involved;
0
Substitute in the Nernst Equation. Eel  Eel 
RT
ln Qel
nF
E.g. MnO4- +8H+ +5e = Mn+2 + 4H2O
Tabulation
aMn2
RT
0
ln
EMnO  / Mn2  EMnO

2 
4
4 / Mn
5F aMnO  aH8 
4
Calculation of Cell Potential:
0
Ecell  Ecell

RT
ln Qrxn
nF
a 4
RT aHg aEDTA4
RT
ln
 Eel0 
ln EDTA
2F
2 F aHgEDTA2
aHgEDTA2
Calculation of E0cell
MnO4- +8H+ +5 Fe+2 = Mn+2 + 5Fe+3 + 4H2O
Tabulated
Elect Pot
MnO4  8 H   5e  Mn 2  4 H 2O
0
MnO4 / Mn 2
Fe 2  Fe 3  e
0
Fe3 / Fe2
Electrode
Reaction Pot
E
E -
Substitute in the Nernst Equation.
Eel  Eel0 
E E
Cell potentials are calculated for redox reaction.
Write the reaction, balanced.
Write the expression for Q of the reaction, recognize
the # electrons involved; for the reaction.
Eg. Hg(EDTA)2-(aq)+ 2e = Hg(l) + EDTA4- (aq)
0
MnO4 / Mn 2
0
Fe3 / Fe2
E0cell = E0red – E0oxd
E0cell = E0+ – E0E0cell = E0MnO4-/Mn+2 – E0Fe+3/Fe+2
Calculation of Cell Potential:
Eg. MnO4- +8H+ +5e = Mn+2 + 4H2O
Fe+2 = Fe+3 + e
Overall (all aqueous species)
MnO4- +8H+ +5 Fe+2 = Mn+2 + 5Fe+3 + 4H2O
0
Ecell  Ecell

0
Ecell  Ecell
Ecell can be calculated by calculation each electrode
potential (reduction) separately and subtracting the
potential at oxidation half from the reduction half.
Ecell = Ered process – Eoxd process
Ecell = E+ – E-
RT
ln Qrxn
nF
5
aMn2 aFe
3
RT

ln
8
5
5 F aMnO  aH  aFe
2
4
0
Ecell  Ecell

0
Ecell  Ecell

RT
ln Qrxn
nF
5
aMn2 aFe
3
RT
ln
8
5
5 F aMnO  aH  aFe
2
4
7
Alternate view of ‘some’ electrodes:
Electrons move from less positive electrode potential
electrode to more positive electrode potential
electrode, when connected by a conductor.
Ag/AgCl(s), HCl(aq)
AgCl(s) + e = Ag(s) + Cl-(aq)
0
Eel  E AgCl
/ Ag 
0.05916
log aCl 
1
Ag+(aq) + e = Ag(s)
0
Eel  E Ag
 / Ag 
a
0.05916
1
0.05916
0
log
log Cl 
 E Ag
 / Ag 
1
1
a Ag 
K sp , AgCl
Alternate view of ‘some’ electrodes:
Use of electrodes in Analytical Chemistry
Pb/PbF2(s), HF(aq)
The fact that the electrode potential (and the cell
potential) is dependent on the concentration of
species allows the use of electrodes as chemical
probes.
PbF2(s) + 2e = Pb(s) + 2F-(aq)
0
Eel  EPbF
2 / Pb 
0.05916
log aF2 
2
Pb+2(aq) + 2e = Pb(s)
0
Eel  EPb
 2 / Pb 
K
0.05916
1
0.05916
0
log
log sp ,2PbF 2
 EPb
 2 / Pb 
2
2
aPb 
aF 
The electrode potential at an electrode measures
the propensity of a reduction reaction to occur;
at the concentrations (activities) of the species that
has attained an electrochemical equilibrium at
the interface.
Not a viable
cell configuration.
Zn
Cu
ZnCl2 (aq)
Cu (NO3)2 (aq)
8
Electrode Interface :
Study of charge transfer reactions is the goal of most
techniques electrochemistry. The interfacial system of
is very complex and even in the absence of electron
transfer and processes other than electron transfer do
occur. Such processes can affect the electrical double
layer and therefore the electrode behavior.
The SHE is cumbersome to construct. Other half-cells
are being used as secondary standards.
Such processes include phenomena such as adsorption,
desorption, and charging of the interface as a result of
changing electrode potential.
These are called non-faradaic processes.
Silver/Silver chloride/KCl reference
AgCl(s) +e- = Ag(s) + Cl- (aq)
Silver/silver chloride:
AgCl(s) +e- = Ag(s) + Cl- (aq)
Potential @ 25°
vs. SHE vs. SCE
Ag/AgCl, KCl (0.1M)
0.288
0.047
Ag/AgCl, KCl (3M)
Ag/AgCl, KCl (3.5M)
0.210
0.205
-0.032
-0.039
Ag/AgCl, KCl (sat'd)
0.199
-0.045
Ag/AgCl, NaCl (3M)
Ag/AgCl, NaCl (sat'd)
0.209
0.197
-0.035
-0.047
http://glossary.periodni.com/glossary.php?en=silver%2Fsilverchloride+electrode
9
silver/silver-chloride
electrode
Calomel:
Hg2Cl2(s) + 2e- = 2Hg (l) + 2Cl-(aq)
Potential vs. SHE / V
t / °C
3.5 mol dm-3
sat. solution
15
20
0.212
0.208
0.209
0.204
25
0.205
0.199
30
35
0.201
0.197
0.194
0.189
Hg
Hg2Cl2(s)
KCl, Hg2Cl2 (aq,sat),
KCl(s)
frit
Calomel:
Hg2Cl2(s) + 2e- = 2Hg (l) + 2Cl-(aq)
Mercury/mercurous sulfate:
Hg2SO4 (s) + 2e- = 2Hg (l) + SO4-2 (aq)
Potential @ 25°
vs. SHE
0.334
vs. SCE
0.0925
Hg/Hg2Cl2, KCl (1M)
NCE (Normal Calomel)
0.280
0.0389
Hg/Hg2Cl2, KCl (3.5M)
0.250
0.006
Hg/Hg2Cl2, KCl (sat'd)
SCE (Sat'd Calomel)
0.241
0
Hg/Hg2Cl2, NaCl (sat'd)
SSCE
0.2360
-0.0052
Hg/Hg2Cl2, KCl (0.1M)
Potential @ 25°
vs. SHE
vs. SCE
Hg/Hg2SO4, H2SO4 (0.5M)
0.682
0.441
Hg/Hg2SO4, H2SO4 (1M)
0.674
0.430
Hg/Hg2SO4, K2SO4(sat'd)
0.64
0.40
10