Part I. Electroanalytical Methods
1. Introduction
 Electroanalytical chemistry is deals with the relationship between
electricity and chemistry
 Analytical calculations are based on the measurement of
electrical quantities (current, potential, charge, or resistance)
and their relationship to chemical parameters
Advantages
 Measurements are easy to automate as they are electrical
signals
 Low concentrations of analytes are determined without
difficulty (high sensitivity)
 Far less expensive equipment than spectroscopy instruments
Fundamental concepts
Charge
one mole of electrons
= (1.602 x 10-19 C)(6.022 x 1023/mol) = 96,485 C/mol
= Faraday constant (F)
 The charge (q) transferred in a redox reaction is given by
q=nxF
Current (i)
 The quantity of charge flowing past a point in an
electric circuit per second
i = q/t
Ampere (A) = coulomb per second (C/s)
Potential
Work done by or on electrons when they move from one
point to another
w=Exq
or
E = w/q
Units: volts (V or J/C)
Ohm’s Law
i = E/R
R = resistance, ohm(Ω) or V/A
Electrochemical cell
 Made up of the electrodes and the contacting sample solution
 Electrical conductor is immersed in a solution of its own ions
 A potential difference (voltage) is created between the conductor
and the solution
 The system is a half-cell
 The metal conductor is an electrode and the solution
is an electrolyte
Galvanic cell
For the overall reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Cell Notation: Zn/Zn2+ // Cu2+/Cu
Nernst equation
For the half reaction
aO + ne-  bR
The half-cell potential (at 25 oC), E, is given by
E  EO 
RT  R 
ln
nF  O a
b



b




2.3RT
R

E  EO 
log
a 
nF
 O 
0.059  R  b 

EE 
log
a 
n
 O 
O
 E = Eo when [O] = [R] = 1M
 Concentration for gases are
expressed as pressures in bars or atm
Concentrations for pure solids,
liquids, and solvents are omitted
(activity = 1)
 Reduction is more favorable on the
negative side of Eo
 When a half reaction is multiplied
by a factor Eo remains the same
Standard Hydrogen Electrode (SHE)
 Reference electrode half-cell
 Used to measure Eo for half-reactions (half-cells)
 Connected to negative terminal (anode)
 Assigned Eo = 0.00 V under standard state conditions (T = 25
oC, concentration = 1M, pressure = 1 atm, pure solid or liquid)
 Platinized Pt electrode immersed in a solution of 1M HCl
 H2 gas (1 atm) is bubbled over the Pt electrode
2H+(aq, 1 M) + 2e-  2H2 (g, 1 bar)
Example: Calculate the half-cell and cell potentials for the
following cell.
Cu/CuSO4 (0.02M)//AgNO3(0.01 M)/Ag
Common reference electrodes
Saturated Calomel Electrode (SCE)
 Composed of metallic mercury in contact with saturated solution
of mercurous chloride (calomel, Hg2Cl2) in contact with saturated
KCl solution
 Pt wire is in contact with the metallic mercury
 E = +0.267 V at 25 oC
Silver/Silver Chloride Reference Electrode (Ag/AgCl)
 Consists of silver metal coated with silver chloride paste
 Immersed in saturated KCl and AgCl solution
 E = +0.222 V at 25 oC
Electrochemical process
Electrode (Example: Pt, C(glassy or graphite), Au, Ag)
 Conducts electrons into or out of a redox reaction system
 The electrode surface serves as a junction between an ionic
conductor and an electronic conductor
 From electronic conductor
When potential is applied, electrons are accumulated on the
electrode surface and causes capacitive current
If electrons cross the interface by redox reaction there is a
Faradaic current
From ionic conductor
There are different process when electroactive species moves to
the electrode surface
 Apart from redox reaction at the electrode surface, there will be
mass transport: conviction, migration and diffusion
2. Potentiometry
 Based on static (zero-current) measurements
 Involves measurement of potential (voltage) of an
electrochemical cell
 Used to obtain information on the composition of an analyte
Reference electrode //indicator electrode
 Indicator (sensing) electrode responds to the concentration of
the analyte species using Nernst equation
 Indicator electrode (Eind) is connected
to a reference electrode
(Eref, such as SCE, Ag/AgCl, to form a complete cell)
Ecell = Eind – Eref
Reference electrode is connected to the negative terminal of
the readout device (potentiometer)
Indicator electrode
 Electrode that responds to change in analyte activity, showing
high selectivity
 Ion-selective electrodes (ISE) –is one type of indicator electrode
which respond directly to the analyte
 Used for direct potentiometric measurements
 Selectively binds and measures the activity of one ion (no redox
chemistry)
Examples
 pH electrode
 Calcium (Ca2+) electrode, CaSE
 Chloride (Cl-) electrode, ClSE
ISE
 Made from a permselective ion-conducting membrane (ionexchange material that allows ions of one electrical sign to pass
through)
 Reference electrode is inbuilt
 Internal solution (solution inside electrode) contains ion of interest
with constant activity
 Ion of interest is also mixed with membrane
 Membrane is nonporous and water insoluble
ISE
 If C+ is the preferential ion that cross the membrane
 [C+] inside the electrode isn’t equal to [C+] outside the electrode
which results in a potential difference across the membrane
RT  [C  ] outer
E 
ln  
z i F  [C ] inner



 [C  ]outer
0.059
At 25 C, E 
log  
n
 [C ]inner
o

0.059
  cons 
log[ C  ]outer
n

Generally (at 25 oC)
 10-fold change in activity implies 59/n mV change in E
 n is the charge on the selective ion (negative for anions)
 n = +1 for K+, n = +2 for Ca2+, n = -2 for CO32-
Selectivity Coefficient (k)
 A measure of the ability of ISE to discriminate against an interfering
ion. It is assumed that ISEs respond only to ion of interest
0.059
E  cons 
log( ai  k j a j )
n
 In practice, no electrode responds to only one specific ion
 The lower the value of kj (selectivity coefficient) the more selective
is the electrode, k = 0 for an ideal electrode (implies no interference)
For k > 1 :- ISE responds better to the interfering ion than to the target
ion
For k = 1:- ISE responds similarly to both ions
For k < 1:- ISE responds more selectively to ion of interest
Empirical Calibration Plot
Potential (mV)
Slope = 59/n mV
Called Nernstian slope
p[C+]
 Used to determine the unknown concentration of analytes
 Departure from linearity is observed at low concentrations
pH glass electrode
 The most widely used for pH measurements (selective sense H+)
whose response is fast, stable, and has broad range
 Potential difference is 0.059 V when [H+] changes by a factor of
10 fold change
Thin glass membrane (bulb) consists of SiO4
 Most common composition is SiO2, Na2O, and CaO
Glass membrane contains
 Dilute HCl solution saturated in AgCl
 Inbuilt reference electrode (Ag wire coated with AgCl)
Glass pH Electrode
Fig. 1 Typical pH electrode: Left: Glass electrode (indicator) and SCE
(reference immersed in a solution of unknown pH; Right: Combined
probe consisting of both a glass electrode and Ag/AgCl reference
pH glass electrode
 Glass electrode response at 25 oC (potential across membrane
with respect to H+)
 Equilibrium establishes across the glass membrane with respect
to H+ in inner and outer solutions which produces the potential, E
E  cons - 0.059 pH
 Linearity between pH and potential
 Calibration plot yields slope = 59 mV/pH units
 Electrode is prevented from drying out by storing in aqueous
solution when not in use
Error in pH glassy electrode
 Glass electrodes respond to the
concentration of both hydrogen ion
and alkali metal ions in basic solution,
which causes alkaline error (four
different glass membranes is shown in
Fig. 2, curve C to F)
Other glass electrodes
Fig. 2 Alkaline and acidic errors
Glass electrodes for other cations
K+ -, NH4+-, Na+-selective electrodes
 Employs aluminosilicate glasses (Na2O, Al2O3, SiO2)
 Minimizes interference from H+ when solution pH > 5
Solid-state electrodes
 Solid membranes that are selective primarily to anions
Solid-state membrane may be
 single crystals (most common)
 polycrystalline pellets or mixed crystals
Examples
 Most common is fluoride-ion-selective electrode (limited pH range
of 0-8.5)(OH- is the only interfering ion due to similar size and charge)
 Iodide electrode (high selectivity over Br- and Cl-)
 Chloride electrode (suffers interference from Br- and I-)
Ion-Exchange Electrodes
 The basis is the ability of phosphate ions to form stable
complexes with calcium ions
 Selective towards calcium
 Employs cation-exchanger that has high affinity for calcium ions
(diester of phosphoric acid)
 Inner solution is a saturated solution of calcium chloride
 Cell potential is given by
0.059
E  cons 
log(aCa )
2
Applications of potentiometry
 Used as detectors for automated flow analyzers
(flow injection systems)
 High-speed determination of blood electrolytes in hospitals
(H+, K+, Cl-, Ca2+, Na+)
 For measuring soil samples (NO3-, Cl-, Li+, Ca2+, Mg2+)
 Coupling ion chromatography with potentiometric detection
 Column detectors for capillary-zone electrophoresis
3. Voltammetry
 Current is measured as function of applied potential
 Solid working electrodes are used and oxidation-reduction takes
place at or near the surface of the working electrode
 Graph of current versus potential is obtained
 Peak current is proportional to concentration of analyte
 Current versus potential plot is known as voltammogram
3. Voltammetry
Potentiostat
-Instrument that controls the
potential at a working electrode
- Connects the three electrodes
Electrochemical Cell
- Contains the three electrodes
immersed in the sample solution
- Electrodes are inserted through
holes in the cell cover
Working Electrode (WE)
Electrode at which the reaction of
interest occurs (Pt, Au, Ag, C)
Voltammetry
 RE is placed as close as
possible to WE to minimize
potential drop caused by the
cell resistance (iR)
 Current flow cannot occur
through RE hence the need
for CE to complete the current
path
 Current flows through
solution between WE and CE
 Voltage is measured
between WE and RE
Electrochemical Cell
Supporting electrolyte
 Inert or don’t involve in redox reaction and its main purpose is
decreases the resistance of the solution
 Eliminates migration effects
 Maintains a constant ionic strength
 Concentration range in usually 0.1 M – 1.0 M
 Should be in large excess of analyte concentration
Mass transport
Diffusion
 Spontaneous movement as a result of concentration gradient
 Movement from regions of high concentration to regions of
low concentration
Convection
 Transport to the electrode by gross physical movement
Driving force is an external mechanical energy
 Solution stirring or flowing or electrode rotation or vibration
 Physical movement as a result of density gradient
Migration
 Movement of charged particles along an electric field
 Charge is carried through the solution as a result of movement of
ions
Cyclic voltammetry
Involves linear scanning of potential of a stationary electrode
using a triangular waveform in unstirred solution

 Assume only O is present initially O + ne-  R
 A negative potential sweep results in the reduction of O to R
(starting from a value where no reduction of O initially occurs)
 As potential approaches Eo for the redox process, a cathodic current
is observed until a peak is reached
 The direction of potential sweep is reversed after going beyond the
region where reduction is observed
Cyclic voltammetry
 Peak current for a reversible couple is given by the Randles-Sevcik
equation (at 25 oC)
i p  2.69 x 10 n ACD ν
5
3/2
Where
n = number of electrons
A = electrode area (cm2)
C = concentration (mol/cm3)
D = diffusion coefficient (cm2/s)
ν = potential scan rate (V/s)
1/2 1/2
Reversible Systems
 ip is proportional to C
 ip is proportional to ν1/2 which implies electrode reaction is
controlled by diffusion
 ia/ic ≈ 1 for simple reversible couple
 For a redox couple
E 
o
E pa  E pc
2
 The separation between peak potentials
0.059
ΔEp  E pa  E pc 
V
n
Used to determine the number of electrons transferred and for a
fast one electron transfer ∆Ep = 59 mV
 Epa and Epc are independent of the scan rate
Irreversible Systems
 Systems with sluggish electron transfer
 Individual peaks are reduced in size and are widely separated
 Characterized by shift of the peak potential with scan rate
Quasi-reversible Systems
 Current is controlled by both charge transfer and mass transport
 Voltammograms are more drawn out
 Exhibit larger separation in peak potentials compared
to reversible systems
Polarography
 Voltammetry in which the working electrode is dropping mercury
 Makes use of potential ramp and conventional DC
 Wide cathodic potential range and a renewable surface
 Hence widely used for the determination of many reducible
species
 Cathodic potential scan is applied and current is measured
 Reduction begins at sufficiently negative potential [concentration
gradient increases and current rises rapidly to its limiting value (iL)]
Polarography
 Diffusion current is obtained by subtracting response due to
supporting electrolyte (background current)
 Analyte species entering region close to the electrode surface
undergo instantaneous electron transfer reaction
 Maximum rate of diffusion is achieved
 Current-potential plot provides polarographic wave (polarogram)
The Ilkovic Equation
i L  708nD m t C
1/2
D = cm2/s
C = mol/cm3
2/3 1/6
m = g/s
t=s
 iL is current at the end of drop life (the limiting current)
Half wave potential (E1/2)
 Potential at which the current is one-half its limiting value
 E1/2 is independent of concentration of species
E 1/2
Where
 DR 
RT

E 
log 
nF
 DO 
1/2
o
DR = diffusion coefficient of reduced species
DO = diffusion coefficient of oxidized species
 Experimental E1/2 is compared to literature values to identify
unknown analyte
Half Wave Potential (E1/2)
At 25 oC
E  E1/2
0.059
 iL  i 

log 

n
 i 
 A graph of E versus log[(iL-i)/i] is linear if reaction is reversible
(Nernstian behavior)
 Slope = 0.059/n and intercept = E1/2
 E = E1/2 when [Ox] = [Red]
Pulse voltammetry
 The basis of pulse techniques
is the difference in the rate of
the decay of the charging and
the faradaic currents following
a potential pulse.
 The charging current decays
exponentially, whereas the
faradaic current as a function
of 1/t1/2
 That is, the rate of decay of the
charging current is
considerably faster than that of
the faradaic current
36
Normal-pulse voltammetry
 Pulse amplitude increases linearly
with each drop
 Current is measured about 40 ms
after each pulse is applied
(at which time charging current is
negligible)
 Diffusion layer is thinner than
that of DC polarography
due to short pulse duration
 Higher faradaic current than DC
polarography
E
Eb

’
t
37
Normal-pulse voltammetry
 Voltammogram has a sigmoidal shape
 Limiting current (il) is given by
nFACD 1/2
il 
π( '- )
 tm (’-)= time after application of pulse
when the current is measured
 normal pulse is about 5-10 times more sensitive
i l, NP
i l,DC
1/2
 3t d 

 
 7t m 
38
Differential-pulse voltammetry
 Small pulses of constant
amplitude are superimposed on a
linear potential ramp applied to the
working electrode
 Potentials are applied just before
the end of each drop
 Current is sampled twice
Just before the pulse application (i1)
and late in the pulse life
(after ~ 40 ms) when the charging
current has decayed (i2)
39
Differential-pulse voltammetry
 ∆i (= i2 – i1) is plotted
against the applied potential
and displayed
(instrument does these)
 The charging current
contribution to the differential
current is negligible
 Detection limit is as low as
10-8 M (~ 1 μg/L)
The Voltammogram: Consists of
current peaks
 The height of peaks is directly
proportional the concentration of
analyte
40
Differential-pulse voltammetry
 The peak shaped response exhibits higher resolution than DC
polarography
 The peak potential (Ep) occurs near the polarographic half-wave
potential and can be used to identify the species
E p  E 1/2
ΔE

2
where ∆E = pulse amplitude
 The width at half-height of the peak (W1/2)
W1/2
3.52RT

nF
If n = 1, W1/2 ≈ 90.4 mV at 25 oC
 Irreversible redox systems produce lower and broader peaks
than reversible systems
41
Problem
The concentration of As(III) in water can be determined by
differential pulse voltammetry in 1 M HCl. The initial potential is set
to -0.1 V vs SCE and is scanned toward more negative potentials at a
rate of 5 mV/s. Reduction of As(III) to As metal occurs at a potential
of about -0.44 V vs SCE. The calibration curve for standard solution is
shown below. What is the concentration of As(III) in a sample of
water if its peak current is 1.10 A? Calculate the concentration of
As(III).
3.0
2.5
ipc/
2.0
1.5
Slope = 0.341
intercept = 0.021
R2 = 0.9998
1.0
0.5
0.0
0
2
4
6
Conce/M
8
10
Square-wave voltammetry
 Large amplitude differential
technique
 The wave form applied to the
working electrode is a symmetric
square wave superimposed on a
base staircase potential
 Current is sampled twice during
each square-wave cycle
 One at the end of the forward
pulse (i1) and one at the end of the
reverse pulse (i2)
43
Square-Wave Voltammetry
 Modulation amplitude is very
large
 Reverse pulses cause the
reverse reaction of any product
formed from the forward pulse
 The net current (i1 – i2) is then
plotted versus the base
staircase potential
 This gives the peak-shaped
voltammogram
44
Square-wave voltammetry
 The peak-shaped voltammogram is
symmetric about the half-wave
potential
 Peak current is proportional to the
concentration
 The difference current
voltammogram reaches a peak at
and has a dimensionless peak current,
p, that depends on n, Ep and Es
45
Stripping analysis
Two step technique
 Deposition step (preconcentration step)
Involves preconcentration of analyte species by reduction
(anodic stripping) or oxidation (cathodic stripping)
into a mercury electrode
 Stripping step
 Measurement step
 Rapid oxidation or reduction to strip the products back into the
electrolyte
Stripping analysis
 Very sensitive for trace analysis of heavy metals
 Favorable signal to background ratio
 About four to six metals can be measured simultaneously
at levels as low as 10-10 M
 Low cost instrumentation
 There are different versions of stripping analysis depending
on the nature of the deposition and stripping steps
Anodic Stripping Voltammetry (ASV)
 The most widely used stripping analysis
 Preconcentration is done by cathodic deposition at
controlled potential and time
 Metals are preconcentrated by electrodeposition into a
small-volume Hg electrode
 Deposition potential is usually 0.3 – 0.5 V more negative than
Eo for the analyte metal ion
Anodic Stripping Voltammetry (ASV)
 Metal ions reach the Hg electrode surface by diffusion
and convection
Electrode rotation or solution stirring is employed to
achieve convection
 Metal ions are reduced and concentrated as amalgams
Mn+ + ne- + Hg → M(Hg)
 Hg film electrodes or Hg drop electrodes may be used
Anodic Stripping Voltammetry (ASV)
Following preselected deposition period:
 Forced convection is stopped
 Anodic potential scan is employed (may be linear or pulse)
 Amalgamated metals are reoxidized (stripped off electrode)
 An oxidation (stripping) current then flows
M(Hg) → Mn+ + ne- + Hg
Cathodic Stripping Voltammetry (CSV)
 Mirror image of ASV
 Involves anodic deposition of analyte and subsequent stripping
by a potential scan in the negative direction
Pre-concentration
An- + Hg  HgA + neStripping step
HgA + ne-  An- + Hg
 Useful for measuring organic and inorganic compounds that
form insoluble salts with Hg (thiols, penicillin, halides, cyanides)
4. Coulometry
 Method in which charge (quantity of electricity)
 is measured
 Species being measured is converted quantitatively to a new species
The Methods Based on Electrolysis
 Electrogravimetry
 Constant-potential coulometry
 Constant-current coulometry (coulometric titrimetry)
Electrolysis
 A process causing a thermodynamically non-spontaneous
oxidation or reduction reaction to occur by application
of potential or current
Quantity of electricity
 Faraday’s law relates the number of moles of
analyte, nA to the charge
nA = Q/nF
where n is the number of moles of electrons in
the analyte half-reaction
 For a constant current of I amperes for t
seconds, the charge in coulombs Q is given by
Q = It
 For a variable current I, the charge is given by
Q = idt
Example
A constant current of 0.80 A was used to deposit copper at the
cathode and oxygen at the anode of an electrolytic cell. Calculate the
mass of each product that was formed in 15.2 min, assuming that no
other redox reactions occur.
The two half-reactions
Cu2+ + 2e-  Cu(s)
2H2O 4e- + O2 + 4H+
Controlled Potential Coulometry
 Three electrode system that permits applied potential pulse or
ramp at the working electrode
 Metal elements are deposited as potential is increased which
increases charge passing through cell
 The instrument is the coulometer which measures q
Applications
 Used to eliminate interferences from other reactions that take
place at different potentials
 Used to determine the number of electrons involved in a reaction
 Used for coulometric titrations
Instrumentation
Controlled potential Coulometry
Apparatus comprises of
 Potentiostat with DC output E
 Inert cathode and anode
 Stirring rod set-up
 Solution may be heated
 Working electrode can be either anode or cathode
 Controlled potential conditions
5. Conductometry
 Conductance: is the ability of the medium to carry the electric
current.
 There is migration of positively charged species towards the
cathode and negatively charged ones through the anode
1
G
R
From Ohm low
E
R
i
ohm-1(seimen)
E is the potential difference,
i is the current intensity.
57
The conductance of the solution depends on:
 Temperature: It is increased by increase of temperature.
 The concentration of ions: As the number of ions increases
the conductance of the solution increases.
 The size of the electrodes
A
L/A is cell constant
G  K
L
K  G
L
A
K is the specific conductance or conductivity
(ohm-1cm-1 or seimen/cm)
 Nature of ions: size, molecular weight, number of charges the
ion carries and other factors 58
Molar conductance 
 It is defined as the conductance of one mole of solute contained
between electrodes spaced one centimeter apart.
o = +o + -o
59
Instrumentments
60
DC verse AC
 The use of direct current is not good for practical work, since
the electrodes would quickly become polarized, which causes
a redox reaction
 Polarization can be prevented by:
 Using alternating current
 Employing platinum covered with platinum black, having
61
Conductivity
 Conductivity of a solution depends on
 Directly on the surface area of the
electrodes
 Inversely on the distance between
the electrodes
 Directly on the concentration of
ions in solutions
 Directly on the mobility of ions
and temperature
62
Application of conductivity
1. Direct measurements
 Checking purity of distilled water or other chemicals
 Determination of physical constants such as
ionization constant.
2. Conductometric titrations
 Very dilute solutions.
 Turbid and highly colored solutions.
 Reaction which is not complete and where there is no
suitable indicator, e.g. reaction between weak acid and
weak base.
63
Titration of strong acid with strong base
The reaction is represented by the following equations
e.g. H+ + Cl- + Na+ + OH-  H2O + Na+ + Cl-
64
Titration of weak acid with strong base
Very weak acid with strong base or a very weak base with
stronger acids
65
Titration of mixture of acids versus strong base
 Determination of mixture of hydrochloric acid (strong
acid) and acetic acid (weak acid) with sodium hydroxide
(strong base)
66