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1. Dr. Mohammed Hamidaddin
College of Pharmacy
Sana’a University
Department of Pharmaceutical Chemistry
427 PHC : Pharmaceutical Instrumental Analysis
ELECTROCHEMICAL METHODS

2. This part of the course is devoted to the principles underlying electrochemical
methods of analysis and its application in analysis of pharmaceuticals.
Emphasis will be concentrated on potentiometry and polarography and their
application to analysis of some drugs and dosage forms.
Course Description
Reference
Instrumental Analysis.
Henery H. Bauer
Gary D. Christian
James E. O’Reilly

3. 
Electrochemistry is a scientific discipline with a well developed system of
theories and quantitative relationships.

It has many applications and used in both fundamental and applied areas
of chemistry:

Study of the mechanisms and kinetics of electrochemical reactions.

Tool for the electrosynthesis of organic and inorganic compounds.

Solution of quantitative analytical problems (electrochemical methods).
Electrochemical methods are physicochemical methods which measure
certain electron quantity or property of a solution which has a quantitative
relationship with the concentration of the analyte.
Introduction

4. Generalities of Electrochemical Methods
Why Electrochemical Methods are not as Popular as Optical or Chromatographic Methods ?
1.Electrochemistry and electrochemical methods are not emphasized in typical
college curricula.
Nearly disappearance of fundamental electrochemistry from beginning general and
physical-chemistry courses, whereas the interaction of electromagnetic radiation with
matter and the energy level concerned is covered in first-year courses.
Electrochemical theory is really no more complex or abstruse, but probably not so well
unified as spectrochemical theory.
Electrochemical methods are not as widely used as spectrochemical or
chromatographic methods for quantitative analytical applications
The exception is the universal use of the potentiometric pH-meter

5. Generalities of Electrochemical Methods
2.Spectrochemical methods are somewhat more amenable to automation than
electrochemical methods.
Examples : are seen in clinical analysis laboratories: one particular hospital performed
a half-million chemical tests, 91% of which were done with spectrochemical methods
and instruments. Reason: this, of course, was due to its use for automated clinical
analyzers, which are primarily optical in approach.
Electrochemical methods are alternative:
When electrochemical methods can provide the same information as other methods.
Electrochemical methods are must:
When only electrochemistry-based methods will provide the answer or will provide the
best answer to the analytical problem at hand.

6. Advantages of Electrochemical Methods
It is very difficult to consider electrochemical methods in general versus other
methods in general, electrochemical methods do have certain advantages
1. Electrochemical instrumentation is comparatively inexpensive.
2.Electrochemical elemental analysis is generally specific for a particular chemical form
of an element:
- mixture of Fe
2+
and Fe
3+
- Inorganic mercury [Hg
0
, Hg
2+
, Hg
2
Cl
2
, ...] or organic [CH
3
Hg
+
, (CH
3
)
2
Hg, ...]
3. Electrochemical methods respond to the activity of a chemical species rather than to
the concentration (advantage or disadvantage ?)
- Calcium level in serum:
Free (physiologically effective) or protein-bound ?
Ion-selective electrode potentiometry or flame photometry.
- Lead in soils:
Lead in clay soils is less available for absorption by plants than from sandy soil.

7. Classification of Electrochemical Methods
Electrochemical method is defined as a method in which the electrical response
of a chemical system or sample is measured
Electrolyte: chemical system capable of conducting current.
Measuring system: used to apply and to measure electrical signals (currents or voltages).
Electrodes (conductors): serve as contacts between the electrolytes and measuring system.
Experimental system is divided into:
Anodes: oxidation occurs, electrons are abstracted from the electrolyte and pass
into the measuring system (external circuit).
Cathodes: reduction occurs, electrons flow from the cathode into the electrolyte.
Working (indicator) electrode: at which the reaction being studied takes place.
Reference electrode: which maintain a constant potential irrespective of changes in
current.
Counter electrode: which serve to allow current to flow through the electrolyte.

8. Classification of Electrochemical Methods
They can be classified into THREE main groups:
1. Methods depending on the movement of an electric field without occurrence of
electron transfer (conductometry).
2. Methods depending on measuring of current or voltage between two electrodes
where electron transfer occur (potentiometry, coulometry and voltammetry).
3. Methods in which the electrolytes are converted to weighable form by means of an
electric current (when electron transfer occur (electrogravimetry).
Electrochemical Methods
PotentiometryConductometryCoulometry VoltammetryElectrogravimetry
PolarographyStrippingCyclicAmperometry

9. Classification of Electrochemical Methods

10. Definitions and Terminology of Electrochemistry
Faradaic Processes
Charge (e.g. electrons) is transferred across the electrode-solution interface. In these
Processes, actual oxidation or reduction occurs; they are governed by Faraday’s law.
Nonfaradaic Processes
Transitory changes in current and/or potential at the
electrode-solution interface (by adsorption)
without occurrence of actual oxidation-reduction reaction.
Charging Current. Charging of an electrode is an
example of nonfaradaic processes.
Ideally polarized electrode.
Electrode at which no charge-transfer (redox) reactions
across the electrode-solution interface, regardless of
the potential imposed from an outside source of
voltage; Nonfaradaic process.
No real electrode can behave in this manner at all potentials
(e.g. Hg electrode in contact with deoxygenated NaCl solution)

11. Modes of Electrochemical Mass Transfer
Migration
It is the movement of charged substances in an electrical gradient, a result of the force
exerted on charged particles by an electric field; a result of simple attraction of positively
charged substances into a negatively charged electrode surfaces or repulsion from a
positively charged electrode.
In almost all electrochemical methods of analysis, migration effects serves no useful
purpose
It is usually swamped out by the addition of a relatively large amount (~ 0.1 or 1 M) of
background (inert or indifferent) electrolyte such as KCl or HNO
3
.
Current can then flow as a result of migration of K
+
or Cl

ions, with negligible migration of
the electroactive species, which then moves as a result of concentration difference
(diffusion)

12. Convection
It is the movement of electroactive material to the electrode by gross physical movement of
the solution. Generally, fluid flow occurs because of natural convection (caused by density
gradient) or forced convection (usually caused by stirring).
Modes of Electrochemical Mass Transfer
Diffusion
It is the movement of electroactive material to the electrode by the influence of
a gradient of chemical potential , that is due to the concentration gradient;
the substances moves from regions of high concentration to regions of low concentrations
in order to minimize or eliminate concentration differences.

13. It is an electrochemical method of analysis concerned with the
determination of an ion concentration via measurement of a potential of an
analyte-sensitive electrode (indicator electrode) dipped in a test solution.
POTENTIOMETRY

14. When metal [M
0
] is immersed in solution of its ions [M
n+
],there is two forces
(pressures).
1- Solution
pressure
Tendency of M
0
to go to solution
M
0
M
n+
+ ne (oxidation)
2- Ionic
pressure
Tendency of M
n+
ion to deposit on M
0

rode
M
n+
M
0
(reduction)
ne
Electrode: Presentation and Potential
Electrode Presentation: M
o
/M
n+
(conc., M ) Zn
o
/Zn
2+
(0.1 M )
M
0
M
n+
Ionic
pressure
Solution
pressure

15. According to the nature of metal, it may has:
1. If element has tendency to loose electrons and converted to its ions.
In this case the metal has high solution pressure, e.g. Zn°, Fe°, Co°, Ni°.
2. If element its metal ions have tendency to accept electrons and converted to
metal, i.e the metal has high ionic pressure, e.g Cu°, Hg°, Ag°.
Zn° Zn
2+
+ 2e
EMF produced (electrode potential) has a negative (-ve) sign.
Cu
o
Cu
2+
+ 2e
EMF produced (electrode potential) has a positive (+ve) sign.
i.e. If solution pressure > ionic pressure so metal rode will carry –ve charge so –ve
electrode potential e.g: Zn
0
/Zn
2+
i.e If solution pressure < ionic pressure so metal rode will carry +ve charge so +ve
electrode potential e.g: Cu
0
/Cu
2+
Electrode: Presentation and Potential

16. Calculation of Electrode Potential
Nernst equation:
Where:
E
25°C = Electrode potential at
25°C
E
o = Standard electrode
potential
n = Number of electrons
gained or lost
[M
n+
] = Molar concentration of
metal ion.
From Nernest equation, E
25°C
is a function of () ionic concentration
E
25°C
= E
o
+ Log [M
n+
]
n
0.059
It is the EMF produced when metal is immersed in 1 M solution of its ions
N.B.
1- The sign of the potential is similar to the charge on the metal electrode.
2- The potential of single electrode can't be measured directly, but measured against reference electrode
(standard electrode which has known and fixed potential) through Electrochemical Cell.
ionic concentration is 1 molar,
E
25°C
= E
o
which is called Standard Electrode Potential

17. Electrochemical Cell
Definition: Electrochemical cell is defined as two conductors or electrodes, usually
metallic immersed in the same electrolyte solution, or in two different electrolyte
solutions which are in electrical contact.
Classification:
Galvanic (Voltaic) cell : in which the electrochemical reactions occur spontaneously
when the two electrodes are connected by a conductor.
These cells are often employed to convert chemical energy to electrical energy
(e.g. Lead-acid battery, flashlight battery, fuel cells, etc.).
Electrolytic cell : in which the chemical reactions are caused to occur by the imposition
of an external voltage greater than the reversible (galvanic) voltage of the cell
These cells are used to carry out chemical reactions at the expense of electrical energy
(e.g. Cells involving synthesizing processes, and electroplating procedures, etc.).

18. Composition: It consists of 2 electrodes (each electrode is considered as half cell), one of them
is zinc electrode (Zn°/ Zn
2+
) and the other is copper electrode (Cu°/ Cu
2+
)
The two electrodes are joined together by liquid junction, known as salt bridge where it permit
transfer (passage) of electric current between the solutions present in the electrodes
Anodic reaction

Cathodic reaction
Galvanic (Voltaic) Cell (e.g. Danniell cell)
Zn° Zn
2+
+ 2eCu
o
Cu
2+
+ 2e
E
Cu°
= E
o
+ Log [Cu
2+
]
2
0.059
E
Zn°
= E
o
+ Log [Zn
2+
]
2
0.059

19. Schematic Presentation:
Electrochemical cell consisting of 2 electrodes :
zinc electrode (Zn°/ Zn
2+
) and the other is copper electrode (Cu°/ Cu
2+
)
Galvanic (Voltaic) Cell (e.g. Danniell cell)
- Zn
o
/Zn
2+
(c
Zn
2+) // Cu
2+
(c
Cu
2+) / Cu° +
Anode, -ve E
Phase boundary
Liquid junction
Cathode, +ve E
E
cell = E
cathode - E
anode
Cell Potential:
The EMF of electrochemical cell is the algebraic sum of the potential developed across the two
electrode-solution interfaces

20. In this type, external EMF is applied, which is transformed to chemical energy
Voltaic cell can be converted to electrolytic cell if we apply sufficiently large potential from external source that
its output opposing that of galvanic cell, where electrode reactions are reversed.
Copper electrode has become the anode and
zinc electrode has become the cathode.
Recharged battery
Electrolytic Cell
Cu
o
Cu
2+
+ 2eZn
2+
+ 2e Zn
o
Platting

21. It is a liquid junction connect between the two half cell without
mixing permit passage of electrical current.
It may be in the form of bend tube or inverted U, shape tube, filled
with agar gel prepared in saturated KCl or KNO
3
solution.
Ions in salt bridge must not pass to the two half cell, this can be achieved
by blocking the two ends of salt bridge with cotton wool or gelatin or agar.
Salt Bridge and Liquid Junction Potential
Liquid junction potential: Potential developed at boundaries of salt bridge.
Liquid junction potential is produced due to the difference in the rates of
migration of both cations and anions of the salt bridge, which leads to unequal
distribution of charges at the ends of salt bridge, thus producing a potential.
Reduction of liquid junction potential:
Choose the electrolyte of salt bridge, that its cations and anions have nearly the same mobility, so that, they move
by the same rate, leading to equal distribution of charges (e.g KCl or KNO
3
: K
+
= 73.5 , Cl

= 76.3 , NO
3

=71.5).
Use high concentration of electrolyte in salt bridge to reduce difference in rates of migration of ions.
NaCl can not be used in salt bridge, why?

22. Electrodes Used in Potentiometric Measurements
Gas-Ions (Hydrogen gas electrode: Pt/H
2,H
+
Metal-Ions (Copper electrode: Cu
0
/Cu
2+
Metal-metal insoluble salt-precipitating ions:
Silver electrode: Ag
0
, AgCl / Cl


Calomel electrode: Hg
0
, Hg
2
Cl
2
/ Cl

Ions-Ions (Quinhydrone electrode: Quinone (Q) / Hydroquinone (H
2
Q)
Reference electrodes
Indicator electrodes
Composition Function
Electrodes

23. These electrodes (half-cell) have known and constant potentials
Used to measure the potential of indicator electrode, through galvanic cell
1- Hydrogen gas electrodes (Normal hydrogen electrode; NHE)
2- Calomel electrodes (Saturated calomel electrode; SCE)
3 -Silver/silver chloride electrodes (Ag/AgCl electrode)
Reference Electrodes

24. Reference Electrodes
E
25°C  [H
+
]
E
25°C = E
o = Zero
Pt, H
2 / H
+
(1 N )Presentation:
2H
+
+ 2e H
2
Potential:
E
25°C
= E
o
+ 0.059 log [H
+
]
in NHE, [H
+
] = 1 , log 1 = zero
Reaction:
E
25°C
= E
o
+ Log [H
+
]
2
2
0.059
Hydrogen Gas Electrodes
Normal Hydrogen Electrode (NHE)

25. Galvanic cell with salt bridge
Reference Electrodes
Hydrogen Gas Electrodes
Normal Hydrogen Electrode (NHE)

26. Advantages
It is a primary reference electrode, as its potential is considered to be zero.
Can be used as the ultimate standard for other electrode potentials and pH values.
Disadvantages
It is difficult to prepare properly (e.g. difficult to keep H
2
gas at one atmosphere during all
determinations) and inconvenient to use.
Its potential is sensitive to oxidants and reductants in solution-anything that oxidizes H
2 or reduce
H
+
can interfere with the equilibrium of the electrode.
The catalytic Pt surface is poisoned by a variety of substances (e.g. As, CN
-
, H
2
S, and Hg), and is
coated by high-molecular weight substances (e.g. proteins and other surface active compounds).
Normal Hydrogen Electrode (NHE)
Reference Electrodes

27. Reference Electrodes
Calomel Electrodes
Saturated Calomel Electrode (SCE)
Hg / Hg
2Cl
2 (satd.), Cl
-
(X M ) Presentation:
Hg
2Cl
2 + 2e 2Hg + 2Cl

Reaction:
Potential:
2
0.059
E
25°C
= E
o
+ Log [Hg
2
2+
]
E
o
= +0.2676 V
Disadvantages
Its potential varies strongly with temperature, owing to the change in
solubility of KCl; can be used only at temperatures less than 80
°
C.
Advantages
Easily made and maintained, and its potential is quite reproducible.
Typical commercial calomel electrode

28. Silver/Silver Chloride Electrodes
Ag/AgCl Electrode
Ag / AgCl (satd.), Cl
-
(X M ) Presentation:
Potential:
Reaction:AgCl
+ e Ag + Cl

0.059E
25°C
= E
o
+ Log [Ag
+
]E
o
= +0.2223 V
Disadvantages
Their apparent equilibrium potentials may differ by several millivolts from one
electrode to another.
More care is necessary for the preparation of highly stable and reproducible
electrodes.
Advantages
Easily made by anodizing a silver wire in chloride media.
It is sufficiently stable for use at temperatures up to about 275
°
C, making
it a useful alternative to SCE at elevated temperatures.
Reference Electrodes

29. Indicator electrode is sensitive to the concentration of analyte.
Its potential changes rapidly with the change of concentration of the analyte.
It must give rapid and reproducible response.
Indicator Electrodes
Metallic Electrodes
First-Order Electrodes (Cation-Selective)

Indicator Electrodes
Second-Order Electrodes (Anion-Selective)
Glass Electrode (pH-Sensitive)
Inert Metal (Pt or Au) Electrodes (Redox-Sensitive)

Liquid Membrane Electrodes

Membrane Electrodes

30. Indicator Electrodes: Metallic Electrodes
M / M
n+
(X M ) Presentation:
Potential:
Reaction:M
M
n+
+ ne
0.059E
25°C
= E
o
+ Log [M
n+
]
First-Order Electrodes (Cation-Selective)
The metal (M) is immersed into a solution of its own ions (M
n+
).
First-Order Electrodes: electrode potential is
directly dependent on the concentration of
its own ions (cations).

31. Indicator Electrodes: Metallic Electrodes
Ag /AgCl / Cl
-
(X M ) Presentation:
Reactions:
Potential: E
25°C
= E
o
+ 0.059 Log K
sp / [Cl

]
Second-Order Electrodes (Anion-Selective)
The metal is coated with metal-insoluble salt
immersed into a solution of the precipitating anions
(e.g. Silver wire coated with AgCl immersed in Cl
-
ions).
Ksp is constant; E
25°C
 1/ [Cl

]
Ag
+
+ Cl


AgCl 
Ag

+ e

Ag
+
K
sp = [Ag
+
][Cl

]
Second-Order Electrodes: electrode potential is inversely
dependent on the concentration of its own ions (anions).
E
25°C
 1/ [Cl

]
Ag-AgCl Reference
Electrode; How ?

32. Indicator Electrodes: Metallic Electrodes
Inert Metal Electrodes (Redox-Selective)
Inert metals (Pt or Au) electrodes are frequently used to detect changes in concentrations
where both the oxidized and reduced forms are soluble ions (Redox –involving reactions);
oxidation of Fe
2+
by Ce
4+
Reactions:
Potential: E
25°C (Fe
3+
/Fe
2+
)
= E
o (Fe
3+
/Fe
2+
)
+ 0.059 Log [Fe
4+
] / [Fe
2+
]
Half reaction (ox): Fe
2+

Fe
3+
+ e
Half reaction (red): Ce
4+
+ e
Ce
3+
Potential: E
25°C (Ce
4+
/Ce
3+
)
= E
o (Ce
4+
/Fe
3+
)
+ 0.059 Log [Ce
4+
] / [Ce
3+
]
Ref. Ind.

33. Indicator Electrodes: Membrane Electrodes
Glass Electrode (pH-Selective)
Presentation of Whole Cell:
Internal RE / Internal electrolyte / H+-responsive glass membrane / Test solution // External RE
Glass Electrode; pH-sensitive

34. Measurement of pH by glass electrode
Indicator Electrodes: Membrane Electrodes
Presentation of Whole Cell:
Internal RE / Internal electrolyte / H
+
-responsive glass membrane / Test solution // External RE
Glass Electrode; pH-sensitive

35. Indicator Electrodes: Membrane Electrodes
Glass Electrode (pH-Selective)
-Glass is irregular three-dimentional structure of
silicate tetrahedra in which each oxygen atom is
shared by two silicate groups; Na
+
and Ca
2+
are
located in this array.
-When the electrode is immersed in aqueous
solution, cations from the surface of the glass
membrane are leached out and replaced by
protons to form a hydrated silica-rich layer about
500 Å thick.
-The external part of this hydrated gel layer can act as a cation exchange membrane.
-Potential is developed due to the exchange ion exchange between H
+
and one of the component ions
of glass matrix (e.g Na
+
).

36. Indicator Electrodes: Membrane Electrodes
Glass Electrode (pH-Selective)
Disadvantages
Must be calibrated fairly often, preferably with buffers within a pH unit of the pH to be
measured.
They can not be used in acidic fluoride media.
Exhibit both “acidic” and “alkaline” errors.
Advantages
A major advantage of glass electrodes is that there is no formal electron transfer
involved in their functioning; thus they are completely uninfluenced by presence of
oxidizing and/or reducing agents in the test solution.
Can be used in presence of catalytic poisons.

37. Indicator Electrodes: Ion-Selective
Disadvantages
ISE are subjected to large number of interferences.
Require frequent calibration.
ISE are not ultra-trace level sensors (not usable below 10
6
M).
Advantages
Measure activity directly, not concentration.
They have logarithmic response, which results in a constant albeit rather
than large error over the concentration range where the Nernest relation holds.
The linear working range of many electrodes is quite large.
They function well in colored or turbid samples, where spectral methods
generally do not.
Electrode measurements are reasonably rapid-equilibrium being reached.
The equipment is simple, quite inexpensive, and can be portable for field operations.
The method is virtually nondestructive for sample (as far as it is in liquid state).
Can be used with very small samples ( 1 mL).

38. POTENTIOMETRY:
Applications

39. Potentiometric Titrations: Instrumentation

40. Potentiometric Titrations: Acid-Base Titrations
Electrochemical Galvanic Cell:
Glass electrode (as indicator) // Calomel electrode (as reference)
Instrument :
pH - meter, which is potentiometer, its scale is in mv.
Determination of End Point :

41. Potentiometric Titrations: Acid-Base Reactions
Potentiometry is used when Ka is smaller than 10
7
̶

42. This electrode system is used during
argentometric determination of halides,
sulphides, mercaptans, phosphates and
oxalates.
Example: Potentiometric titration of chloride with AgNO
3
Potentiometric Titrations: Precipitation Reactions
Electrochemical Galvanic Cell:
Metallic silver electrode (as indicator) // Calomel electrode (as reference)

43. This electrode system is used for determination
of metal ions (e.g. Cd
2+
, Ca
2+
, Bi
3+
)
Example: Potentiometric titration of Cd
2+
with by EDTA
Potentiometric Titrations: Complex-Forming Reactions
Electrochemical Galvanic Cell:
Ion-selective electrode (as indicator) // Calomel electrode (as reference)

44. Examples:
Potentiometric Titrations: Redox Reactions
Electrochemical Galvanic Cell:
Inert metal (Pt) electrode (as indicator) // Calomel electrode (as reference)
Titration of Fe
2+
with Ce
4+ Titration of Fe
2+
with KMnO
4
Volume of Ce
4+
solution added

45. POLAROGRAPHY

46. Polarography is the branch of voltammetry in which a dropping mercury electrode
(DME) is used as indicator electrode.
It is the electrochemical analytical technique that deals with the effect of an
externally applied potential of an electrode in an electrolysis cell on the current that
flows through it.
The determination of an analyte concentration is achieved via studying of an
applied potential of the DME (analyte-sensitive electrode) dipped in the test
solution.
Polarography: Principles

47. The procedures normally involve the use of cell containing:
The test solution.
Stable reference electrode (usually SCE).
Small-area working indicator electrode (DME electrode).
Auxiliary (counter) electrode; which together with the working electrode carries the
electrolysis current to the measuring circuit. Its presence must not affect the
electrolysis (oxidation or reduction) of the electroactive analyte substance.
Polarography: Procedures

48. Drops of mercury fall from the orifice of the DME capillary at
a constant rate (usually 5-30 drops/min) .
Each drop is the electrode while attached to the column
of mercury in the capillary.
At the slow rate of voltage scanning (50-200 mV/min), the change
of DME potential during a single drop life time
(usually 2-12 sec) is neglected; thus the current measured on each
drop is obtained under practically potentiostatic conditions
(i.e. at constant potential).
The current-voltage curves (polarograms) obtained with DME are very reproducible, since the
surface of every new mercury drop is fresh, clean, and practically unaffected by electrolysis at
earlier drops. With 20 ml sample solution, 100 polarograms can be recorded without noticeable
change.
The small size of DME permits the analysis of small sample volumes (~ 0.01 ml).
Mercury is chemically inert in most aqueous solutions, and hydrogen is evolved on it only at quite
negative potentials; consequently, the reduction of many chemical species can be studied at
mercury electrodes, but not at electrodes made of most other materials.
However, the anodic dissolution of mercury makes it impossible to study reactions at potentials
more positive than ~ +0.4 V versus SCE.
Polarography: Procedures

49. Polarography: Instrumentation
Different Versions of Commercial Polarographs

50. Polarography: Instrumentation
Polarographic Cell
+ Anode+ Anode
 CathodeCathode

51. Instrumentation: Dropping-Mercury Electrode
Composition and Construction of DME
The DME consists of a glass capillary
attached to a mercury reservoir.
Drops of mercury fall from the orifice of
the capillary at a constant rate (usually 5-
30 drops/min) .
Each drop is the electrode while
attached to the column of mercury in the
capillary.

52. 1.Reproducible results are obtained due to the continuously renewed smooth surface of
the mercury drop which allows reproducible rapid electron transfer.
2.Mercury is chemically inert in most aqueous solutions and it forms amalgum with many
metals that reduced at the electrode surface.
3.The diffusion current (analytical signal) assumes a reproducible and steady value
immediately after each change of applied potential.
4.It can be used in acidic solutions as Large overvoltage is needed for reduction of H
+
to
H
2 (usually more negative than – 1.8 V).
5.Large hydrogen over potential on mercury allows the electrolysis of substances those
difficult to be reduced (e.g. alkali metal ions).
Advantages of DME:
Instrumentation: Dropping-Mercury Electrode

53. 1.At potential above +0.4 V, Hg metal is oxidized to Hg
+
producing anodic polarographic
wave which can masks other waves (i.e. cause interferes with the analyte wave).
2.DME is limited to the easily reducible or oxidizable substances (usually those require
potentials at +0.4 to –1.8 V).
3.The capillary is very small and can be blocked by any impurities in mercury.
4.Oxygen, unless it removed prior to the analysis, interferes with some analyte
substances.
Disadvantages of DME:
Instrumentation: Dropping-Mercury Electrode

54. 1.Decrease the resistance of the test solution and to ensure that the electroactive
species moves by diffusion and not by electrical migration in the voltage field across
the polarographic cell.
2.Provide optimum conditions for the particular analysis: (e.g. buffering at a preferred
pH value and elimination of interferences by selective complexation of some species).
Reasons for Use of Supporting Electrolyte in Polarographic Analysis
Instrumentation: Supporting Electrolyte
The total concentration of supporting electrolyte is usually between 0.1 and 1.0 M

55. Instrumentation: Removal of Dissolved Oxygen
Reasons for Removing of Dissolved Oxygen from Sample Solution
Oxygen must be removed by bubbling nitrogen through the solution for five min
prior to the analysis
Oxygen is an electroreducible species, thus its presence in test solution produces a double wave
in the range of 0 to − 1.0 V due reduction reaction (at pH 7 versus SCE):
O
2 +2H
+
+ 2e H
2O
2 ( E
½ =
− 0.05 V)
H
2
O
2
+ 2H
+
+ 2e 2H
2
O ( E
½
=

− 1.3)
This wave (due to large oxygen reduction current) interferes with the analyte wave.
Oxygen or its reduction products can react with many test solution components.

56. A simplistic look at the effects of deareation of the Cd
2+
wave
Instrumentation: Removal of Dissolved Oxygen

57. Polarography: Description of Typical Polarogram
Residual Current
1.It is a small current flows at the most positive
potentials; it is also called charging or
condenser current.
2.It is a nonfaradaic current often practically identical
with the current obtained at the same potential
range with the supporting electrolyte alone.
3.It is observed even when very pure solutions
are used.
4.It is considered as a condenser current made by the continual charging of new mercury drops
to the applied potential.
5.Also, the electrolyte solutions may contain traces of impurities giving small currents those
superimposed upon the condenser current; both types of current are included in the residual
current and must be subtracted from the total observed current.

58. Polarographic wave or step
The part of the polarogram where the current
increases
steeply with the applied potential.
Decomposition potential
The potential at which electrolysis of the substance
occurs giving steep rise in current (polarographic wave).
Limiting current
The highest practically remaining constant current (plateau part of the polarographic wave),
and it is virtually parallel to the charging current.
- Oscillations in currents as the mercury drop grows and falls.
-As the current increases, the magnitude of the oscillations increases in direct proportion.
-The recorded current does not fall to zero at the instant the drop falls because of the slow
response of the current-recording device.
Polarography: Description of Typical Polarogram

59. 59
Diffusion current
The difference between the limiting and charging
current (wave height); it is usually proportional to
the analyte concentration (used as a basis for
quantitative analysis).
Polarography: Description of Typical Polarogram
Polarogram Calibration curve
A
E, V Concentration
Basis of quantitative polarographic analysis

60. Half-wave potential (E½)
The potential at which the current reaches half the
magnitude of the diffusion current.
-It is practically independent on the concentration of the
electroactive compound.
-It characterizes the oxidation-reduction properties of the
studied substance (used as a basis for qualitative
analysis).
-It can be related to the standard oxidation-reduction
potential (E
0
) of the elecrochemical reaction involved.
A Polarogram showing constancy of E½ of Cd
ion reduction at different concentrations of CdCl
2
Polarography: Typical Polarogram

61. Ilkovič Equation:
I
d
= 708 n C D
1/2
m
2/3
t
1/6
I
d Diffusion current (microamperes)
n Number of electron involved in the reaction
D Diffusion coefficient (cm
2
sec¯
1
)
C Concentration (mM )
m Rate of the mercury flow (mg/sec)
t Lifetime of a drop of mercury (sec; 2-7)
n, D, m and t are constants; i
d
= k C (basis for quantitative analysis)
Polarography: Factors Affecting Diffusion Current

62. Polarography: Quantitative Analysis
Polarogram Calibration curve
A
E, V Concentration
Basis of quantitative polarographic analysis
Direct Polarographic Analysis: when the analyte substance is electrochemically active.
Indirect Polarographic Analysis: when the analyte substance is electrochemically inactive.
It should be converted by a chemical reaction into electrochemmically active derivative.

63. In Absence of Interference in E
1/2
Mixture of Cu
+
, Pb
2+
, Cd
2+
, Ni
2+
, Zn
2+
,
Mn
2+
, Al
3+
, and Ba
2+
is determined
simultaneously in 0.5 M NH
4OH, O.5 M
NH
4Cl as each cation has its
characteristic E
1/2
and shows separate
wave.
Determination of Inorganic Cations
A Polarogram showing the reduction of seven different cations
Polarography: Applications in Quantitative Analysis

64. In Presence of Interference in E
1/2
A mixture of lead (Pb
2+
) and thallium (Tl
+
) gives a
wave that can be resolved into two individual
waves ; the E
1/2 differs by only ~ 60 mV.
Determination of Inorganic Cations
Polarography: Applications in Quantitative Analysis
A Polarogram of Pb
2+
and Tl
+
showing the shift of waves
with change in background electrolyte
Upon addition of excess NaOH :
•Pb
2+
form a hydroxo complex; thus its wave is shifted to more –ve potentials.
•Tl
+
does not form a hydroxo complex; thus its wave is resolved at less –ve potential.
Generally, when a metal cation forms a complex, its reduction potential is made
more –ve (i.e. it is more difficult to be reduced.

65. Bromate, iodate, dichromate and vanidate can be determined by polarography in
strongly buffered medium.
Determination of Inorganic Anions
Determination of Inorganic Molecules
Gases (e.g. oxygen), hydrogen peroxides, elemental sulfur, and sulfur oxides.
Oxygen (O
2) in blood and lymphatic fluid can be determined by polarography;
O
2 has two reduction waves (discussed in previous sections).
CS
2
and SO
2
in atmosphere in industrial area can be also determined by polarography.
Polarography: Applications in Quantitative Analysis
A great number and variety of organic compounds can be determined by reduction at
the DME (e.g. quinones, indophenols, p-aminophenol, thiazine, phenazines,
formaldehyde, benzaldehyde, vanillin, pyridoxal, cinamaldhyde, hetrocyclic ring-
containing alkaloids, benzodiazepines, local anthaestics and chloramphenicol.
Determination of Organic Molecules

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Electrochemical Methods