Potentiometry: Electrical potential, electrochemical cell, reference electrodes, indicator electrodes, measurement of potential and Ph, construction and working of electrodes, Potentiometric titrations, methods of detecting end point, Karl Fischer titration.

1. DISCOVER . LEARN . EMPOWER
Potentiometry
UNIVERSITY INSTITUTE OF PHARMA SCIENCES
Pharm D (3rd Year)
Pharmaceutical Analysis
Mr. Yunes Alsayadi
Assistant Professor of
Pharmaceutical Analysis

2. Potentiometry
• Potentiometry is one of the
methods of electroanalytical
chemistry. It is usually employed
to find the concentration of a
solute in solution.
• In potentiometric
measurements, the potential
between two electrodes is
measured using a high
impedance voltmeter.

3. History
• In potentiometry we measure the potential of an electrochemical cell
under static conditions. Because no current—or only a negligible
current—flows through the electrochemical cell, its composition
remains unchanged.
• For this reason, potentiometry is a useful quantitative method of
analysis.
• The first quantitative potentiometric applications appeared soon after
the formulation, in 1889, of the Nernst equation, which relates an
electrochemical cell’s potential to the concentration of electroactive
species in the cell,

4. Potentiometry Principles
• The potential difference between the two electrodes used forms the
basis of the potentiometry principle. The addition of a titrant leads to
a change in the ionic concentration, causing changes in the potential
difference.
• The indicator electrode measures this potential difference. The
reference electrode has a potential value and remains stable when
dipped into the sample solution.
• The salt bridge is a divide used during potentiometric titration to
avoid the intervention of the analyte solution with the reference
solution. Analyte solution is the solution whose potential we must
determine.

5. • We can calculate the total electric potential or the potential
difference as:
• Ecell = Eind – Eref + Ej
• Here,
• Ecell = potential of the whole cell
• Eind = potential of the indicator
• Eref = potential or electromotive force of the reference electrode
• Ej = potential at the junction across the salt bridge
The measured electrical potential depends on the concentration of the ions in
contact with the indicator electrode.

6. Electrical Potential
• Potentiometry is based on the
measurement of the potential of
an electrode system (e.g.
electrochemical cell).
Potentiometric measurement
system consists of two electrodes
called reference and indicator
electrode, potentiometer and a
solution of analyte.
• Potential of an indicator electrode
depends mainly on the
concentration of the analyte ions.
Potentiometric measurement system
/ Potentiometer (for pH
measurement)

7. Electrochemical cell
• Electrochemical cell consists of two
solutions connected by a salt bridge and
electrodes to form electrical circuit.
• Sample cell consists of solutions of ZnSO4
and CuSO4.
• Metallic Zn and Cu electrodes are
immersed in respective solutions.
Electrodes have contacts firstly through
wires connected to the voltmeter and
secondly through solutions and a salt
bridge, forming an electric circuit.
• Salt bridge consists of a tube filled with
saturated salt solution (e.g. KCl solution).
• The ends of the tube are capped with
porous frits that prevent solutions from
mixing, but permit movement of ions.
A galvanic electrochemical
cell

8. • Three distinct charge transfer processes are described for the system
in Fig :
1. Electrons move in electrodes and wires from zinc electrode to
copper electrode.
2. Ions move in solutions:
a. In solution on the left, zinc ions move away from the electrode and sulfate
ions move towards it.
b. In solution on the right, copper ions move towards the electrode and
negatively charged ions (sulfate) away from it.
c. In salt bridge positive ions move right and negative ions left.
3. On the surfaces of electrodes electrons are transferred to ions or
vice versa:
a. Zinc electrode dissolves: Zn → Zn2+ + 2e-
b. Metallic copper is deposited on the electrode surface: Cu2+ + 2e- → Cu ↓

9. Nernst equation
• Potential on an electrode depends
on the ions present in the solution
and their concentration. This way
electrochemical cells can be used
to determine ions and their
concentration in solution. The
dependence of potential between
electrodes from concentration of
ions is expressed by Nernst
equation:
E – electrode potential,
E0 – standard potential of the
electrode,
R – universal gas constant (8.314
J/(K•mol)),
F – Faraday constant (96485 C/mol),
T – temperature in kelvins,
n – charge of the ion or number of
electrons participating in the
reaction,
a – activity of the ions.

10. Reference electrodes
• The electrode whose half-cell
potential is known and is constant
and completely insensitive to the
composition of the solution is
called a reference electrode.
• The reference electrode can act as
both anode or cathode depending
upon the nature of other
electrodes.
• A reference electrode has a stable
and well defined electrochemical
potential (at constant
temperature), against which the
applied or measured potentials in
an electrochemical cell are
referred.
• The reference electrodes are
classified into two types:
1.Primary reference electrode: The
standard hydrogen electrode is
called a primary reference
electrode.
2.Secondary reference
electrode: The electrode whose
potential is determined by
connecting to the standard
hydrogen electrode is called a
secondary reference electrode.
Example: Calomel electrode

11. Standard Hydrogen Electrode (SHE)
• Standard hydrogen electrode scheme:
1) Platinized platinum electrode,
2) Hydrogen gas,
3) Acid solution with an activity of H+ = 1 mol/L,
4) Hydroseal for prevention of oxygen
interference,
5) Reservoir via which the second half-element of
the galvanic cell should be attached.
• The connection can be direct, through a narrow
tube to reduce mixing, or through a salt bridge,
depending on the other electrode and solution.
• This creates an ionically conductive path to the
working electrode of interest.

12. Standard Hydrogen Electrode (SHE)
• It is rarely used for routine analytical work, but is important because it is
the reference electrode used to establish standard-state potentials for
other half-reactions.
• A conventional salt bridge connects the SHE to the indicator half-cell.
• The shorthand notation for the standard hydrogen electrode is
• Pt(s), H2 (g, 1 atm) II H+ (aq)
• 2H+(aq) + 2e  H2(g); Eo = 0 V

13. Saturated Calomel Electrode (SCE)
• The saturated calomel electrode (SCE) is a reference electrode based on the
reaction between elemental mercury and mercury(I) chloride.
• It has been widely replaced by the silver chloride electrode, however the calomel
electrode has a reputation of being more robust.
• The aqueous phase in contact with the mercury and the mercury(I) chloride
(Hg2Cl2, "calomel") is a saturated solution of potassium chloride in water.
• The electrode is normally linked via a porous frit to the solution in which the
other electrode is immersed. This porous frit is a salt bridge.
• The calomel electrode is very practical and very robust, and is one of the most
common electrodes used in corrosion studies.
• The SCE is used in pH measurement, cyclic voltammetry and general aqueous
electrochemistry. The SCE met the criteria of ideal reference electrodes;
therefore, SCE is the most popular reference electrode for use in aqueous
solutions.

15. Working
• The electrode is represented as :
KCl(satu.)Hg2Cl2(s)Hg(l)
• Oxidation : If the electrode serves as
anode, then half reaction that occurs on it
will be oxidation. Mercury is first oxidised
to mercuric ions.
2Hg(l)⟶Hg2
+2 +2e−
The chloride ions supplied by KCl solution
combine with mercuric ions [Hg2
+2] to
form insoluble mercurous chloride. Thus,
Hg2
2+ +2Cl−⟶Hg2Cl2(s)
Overall reaction is,
2Hg(l)+2Cl−
(aq)⟶Hg2Cl2(s)+2e−
• Reduction : If the electrode is cathode in
the galvanic cell, the half reaction that
occurs on it will be reduction :
Hg2Cl2(s)+2e−⟶2Hg(l)+2Cl−
(aq)
• The potential of calomel electrode
decreases with increase in the
concentration of chloride ions at a given
temperature. Thus, electrode is
reversible with respect to concentration
of chloride ions.

16. • The shorthand notation for the calomel electrode half-cell is:
Hg(l) | Hg2Cl2 (sat'd), KCl (aq, saturated)
• The SCE has the advantage that the concentration of Cl-, and, therefore,
the potential of the electrode, remains constant even if the KCl solution
partially evaporates.
• A significant disadvantage of the SCE is that the solubility of KCl is
sensitive to a change in temperature. At higher temperatures the
concentration of Cl- increases, and the electrode's potential decreases.
• Electrodes containing unsaturated solutions of KCl have potentials that
are less temperature-dependent, but experience a change in potential if
the concentration of KCl increases due to evaporation.
• Another disadvantage to calomel electrodes is that they cannot be used
at temperatures above 80 °C.

17. INDICATOR ELECTRODE
• The electrode is used to measure the potential of a solution to which
it is dipped inside along with reference electrode.
• Indicator Electrode: electrode that responds to analyte and
donates/accepts electrons while Reference Electrode: second ½ cell at
a constant potential
• Cell voltage is difference between the indicator and reference
electrode.

18. Indicator Electrodes
1.) Two Broad Classes of Indicator Electrodes
 Metal Electrodes
- Develop an electric potential in response to a redox reaction at the metal surface
 Ion-selective Electrodes
- Selectively bind one type of ion to a membrane to generate an electric potential
Remember an electric potential is generated by a separation of charge

19. METALLIC INDICATOR ELECTRODES

20. Electrodes of the First Kind
• Pure metal electrode in direct equilibrium with its cation
• Metal is in contact with a solution containing its cation.
M+n(aq) + ne-  M(s)

21. Disadvantages of First Kind Electrodes
•Not very selective
• Ag+ interferes with Cu+2
•May be pH dependent
• Zn and Cd dissolve in acidic solutions
•Easily oxidized (deaeration required)
•Non-reproducible response

22. Electrodes of the Second Kind
• Respond to anions by forming precipitates or stable complex
• Examples:
1. Ag electrode for Cl- determination
2. Hg electrode for EDTA determination

23. Inert Metallic (Redox) Electrodes
• Inert conductors that respond to redox systems
• Electron source or sink
• An inert metal in contact with a solution containing the soluble
oxidized and reduced forms of the redox half-reaction.
• May not be reversible
• Examples:
• Pt, Au, Pd, C

24. MEMBRANE ELECTRODES
• Consist of a thin membrane separating 2 solutions of different ion concentrations
• Most common: pH Glass electrode

25. Glass pH Electrode

26. Properties of Glass pH electrode
• Potential not affected by the presence of oxidizing or reducing
agents
• Operates over a wide pH range
• Fast response
• Functions well in physiological systems
• Very selective
• Long lifespan

27. Liquid Membrane Electrodes
• Potential develops across the interface between the analyte solution and a
liquid ion exchanger (that bonds with analyte)
• Similar to a pH electrode except that the membrane is an organic polymer
saturated with a liquid ion exchanger
• Used for polyvalent ions as well as some anions
• Example:
• Calcium dialkyl phosphate insoluble in water, but binds Ca2+ strongly

28. Responsive to Ca2+
0.1 M CaCl2

30. Crystalline-Membrane Electrodes
•Solid state electrodes
•Usually ionic compound
•Crushed powder, melted and formed
•Sometimes doped to increase conductivity
•Operation similar to glass membrane

31. Crystalline-Membrane Electrodes
• AgX membrane: Determination of X-
• Ag2S membrane: Determination of S-2
• LaF3 membrane: Determination of F-

32. Gas Sensing Probes
• A galvanic cell whose potential is related to the concentration of a
gas in solution
• Consist of RE, ISE and electrolyte solution
• A thin gas-permeable membrane (PTFE) serves as a barrier
between internal and analyte solutions
• Allows small gas molecules to pass and dissolve into internal
solution
• O2, NH3/NH4
+, and CO2/HCO3
-/CO3
2-

33. Gas Sensing
Probe

34. 34
Potentiometric Titration Method
Potentiometric Titration is done via the usage of two electrodes – an indicator electrode and a reference
electrode (generally a hydrogen electrode or a silver chloride electrode). One half-cell is formed with the
indicator electrode and the ions of the analyte, which is generally an electrolyte solution. The other half-cell is
formed by the reference electrode.
The overall cell potential can be calculated using the formula given below.
Ecell=Eind–Eref+Esol
Where the potential drop between the indicator and reference electrodes over the electrolyte solution is
given by Esol.
The overall cell potential, Ecell is calculated in every interval where the titrant is measured and added. Now, a
graph is plotted with the Potential difference on the Y-axis and the volume on the X-axis as shown below.
It can be observed from the graph that the electric potential of the cell is dependent on the concentration of
ions which are in contact with the indicator electrode. Therefore, the Ecell is measured with each addition of
the titrant.

35. 35
Types of Potentiometric Titration
There are four types of titration that fall under the category of potentiometric titration, namely acid-base
titration, redox titration, complexometric titration, and precipitation titration. A brief description of each of
these types of titration is given below.
1. Acid-Base Titration: This type of potentiometric titration is used to determine the concentration of a given
acid/base by neutralizing it exactly using a standard solution of base/acid whose concentration is known.
2. Redox Titration: This type of potentiometric titration involves an analyte and titrant that undergo a redox
reaction. An example of this type of titration would be the treatment of an iodine solution with a reducing
agent which produces iodide ion (a starch indicator is used to get the endpoint).
3. Complexometric Titration: This type of titration can also be referred to as chelatometry. In this method, a
coloured complex is formed, indicating the end point of the titration. This method is used to determine a
mixture of metal ions in a given solution.
4. Precipitation Titration: This type of titration involves a reaction between the given analyte and the titrant
wherein an insoluble precipitate is formed. The end-point of this titration is noted when the addition of the
titrant no longer forms a precipitate.

36. POTENTIOMETRIC TITRATION
• Involves measurement of the potential of a suitable
indicator electrode as a function of titrant volume
• Provides MORE RELIABLE data than the usual titration
method
• Useful with colored/turbid solutions
• May be automated
• More time consuming

37. End point
• The endpoint of the titration is the point at which the colour of the solution
changes completely due to the formation of product due to the addition of
indicator.
• It is to be noted that weak acids only show one endpoint in the titration.
• Example: In acid base titration, due to addition of Phenolphthalein, the solution
changes its colour.
• The main difference between equivalence and endpoint is that the equivalence
point is a point where the chemical reaction comes to an end, while the endpoint
is the point where the colour change occurs in a system.
• since there is only a slight difference between an equivalent point and an
endpoint, it can be considered the same for laboratory purposes.

38. Methods of detecting end point
• Visual Inspection: In some cases, the endpoint can be determined
visually by observing a sudden change in the color of an indicator
solution. For example, phenolphthalein changes from colorless to
pink in acidic solutions, indicating the endpoint in acid-base titrations.
• pH Measurement: pH meters can be used to monitor the pH of the
solution throughout the titration. A significant change in pH often
corresponds to the endpoint, especially in acid-base titrations.

39. Considerations for Specific Titrations
• Different types of potentiometric titrations require specific
considerations for endpoint determination:
• Acid-Base Titrations: In acid-base titrations, the endpoint is often
detected using pH indicators or pH meters. The pH at the equivalence
point is typically close to 7 for neutralization reactions.
• Redox Titrations: In redox titrations, the endpoint is determined by
monitoring changes in the voltage due to the transfer of electrons
between the analyte and titrant.
• Complexometric Titrations: In complexometric titrations, the
endpoint is typically detected using indicators or electrodes sensitive
to metal ions. Ethylenediaminetetraacetic Acid (EDTA) is the example

40. Karl Fischer titration
• Karl Fischer (March 24, 1901 – April 16, 1958) was a German chemist
Published a method in 1935 to determine trace amounts of water in
samples.
• This method is now called Karl Fischer titration. Abbreviations: KF or KFT
• It remains the primary method of water content determination used
worldwide by:
Government – Food Science
 Academia – Research –
Industry – Quality Control
• Karl Fischer titration is a classic titration method that
uses coulometric or volumetric titration to determine trace amounts
of water in a sample.

41. Chemical principle
• The principle of Karl Fischer titration is based on the oxidation reaction between
iodine and sulphur dioxide.
• Water reacts with iodine and sulphur dioxide to form sulphur trioxide and hydrogen iodide.
• An endpoint is reached when all the water is consumed.
• The chemical equation for the reaction between sulphur dioxide, iodine, and water (which is employed during
Karl Fischer titration) is provided below.
• I2 + SO2 + H2O → 2HI + SO3
• This elementary reaction consumes exactly one molar equivalent of water vs. iodine. Iodine is added to the
solution until it is present in excess, marking the end point of the titration, which can be detected by
potentiometry. The reaction is run in an alcohol solution containing a base, which consumes the sulfur
trioxide and hydroiodic acid produced.

42. Karl Fischer Titration Equipment
• Drying tube, sample injection
cap, electrode analysis, Drain
cook, a cathode chamber,
detection electrode, rotor,
anode chamber, KF reagent.
• Ingredients of KF reagent:
Iodine, Buffer (Imidazole),
sulphur dioxide, solvent
(methanol).

43. Karl Fischer Titration Procedure
• The Karl Fischer titration experiment can be performed in two
different methods. They are:
• Volumetric determination – This technique is suitable to determine
water content down to 1% of water. The sample is dissolved in
KF methanol and the iodine is added to KF Reagent. The endpoint is
detected potentiometrically.
• Coulometric determination – The endpoint is detected in this
experiment electrochemically. Iodine required for KF reaction is
obtained by anodic oxidation of iodide from solution.

44. Advantages of Karl Fischer Titration
• It is fitted for determining water in
gases, liquids and solids.
• The coulometric titrator helps in
detecting free water, dissolved
water, and emulsified water.
• It is a swift process which demands
a minimal amount of sample
preparation.
• Extremely accurate method.
Limitations of Karl Fischer Titration
• It is a destructive technique.
• The solvent consumption is high as
the manual volumetric titration
demands reloading during each
determination.
• Coulometric titration is fitted only
for samples that contain a small
amount of water.
• Coulometric titration takes
extremely long periods to
determine.

45. • A.I. Vogel, Text Book of Quantitative Inorganic Analysis; 5th Edition, John Wiley &
Sons, Inc., New York.
• A.H. Beckett and J.B. Stenlake's, Practical Pharmaceutical Chemistry Vol I and II,
Stahlone Press of University of London.
• P. Gundu Rao, Inorganic Pharmaceutical Chemistry.
• Bentley and Driver's Textbook of Pharmaceutical Chemistry.
• http://www.anachem.umu.se/jumpstation.htm
• http://userwww.service.emory.edu/~kmurray/mslist.html
• http://www.anachem.umu.se/jumpstation.htm
References

46. THANK YOU
For queries
Email: yunes20171@gmail.com