1. Instrumental Methods II
SATISH PRADHAN DNYANASADHANA COLLEHE, THANE
Paper III
Unit II
Instrumental Methods II
Instruments Based on Electrochemical Properties of the Analytes
Potentiometry By
Dr. Bhushan Langi

2. Introduction: Electrochemistry
Electrochemistry is the study of the relationship between chemical transformation and electrical energy.
Electrochemical cell
It is the cell in which conversion of chemical energy to electrical energy can occur in either direction.

3. Galvanic or Voltaic Cell
Electrochemical cell is of two types-
Electrochemical Cell
Electrolytic Cell
Galvanic or Voltaic Cell
Converts electrical energy into chemical energy with non-spontaneous redox reaction.
Converts chemical energy into electrical energy with spontaneous redox reaction.

4. Electrochemical cell consists of two electrodes viz
Electrochemical cell consists of two electrodes viz. anode (oxidation reaction) and cathode (reduction reaction)
The electrode at which an oxidation reaction takes place is called as anode whereas the electrode at which a reduction reaction is takes place is called as cathode.
Electrode Charges
Cell
Anode
Cathode
Electrolytic Cell
Positive
Negative
Galvanic or Voltaic Cell

5. Redox Reaction:
Oxidation: Loss of electrons causes increase in oxidation state.
Reduction: Gain of electrons causes decrease in oxidation state.

6. Difference between Electrolytic and Galvanic or Voltaic Cells

7. Cell Notation of Electrochemical Cell:
For drawing a cell diagram, we follow the following conventions.
The anode is always placed on the left side, and the cathode is placed on the right side.
The salt bridge is represented by double vertical lines (||). 
The difference in the phase of an element is represented by a single vertical line (|), while changes in oxidation states are represented by commas (,).
Concentration of aqueous solutions are written in parentheses after the symbol for ion or molecule.

8. Consider the following reaction:
Two half Reactions
At anode (Oxidation):
At cathode (Reduction):
Half-Cell:  
When electrodes are immersed in a solution containing ions of the same metal, it is called a half-cell.

9. Salt Bridge:
U-shaped tube containing an electrolyte (typically in the form of a gel), providing electrical contact between two solutions.
The electrolyte used to prepare salt bridge is generally KNO3, KCl or NH4NO3

10. Cell Potential: (Ecell)
Cell potential is the potential difference between two electrodes.
Ecell = Ecathode – Eanode
In Standard state,
Eocell = Eocathode – Eoanode

11. Reference Electrode: Indicator Electrode:
A reference electrode is that electrode whose potential is known and remain constant.
e.g. Saturated calomel electrode (ESCE = 0.242)
Indicator Electrode:
An indicator electrode is that electrode whose potential depends on the activity of ions being titrated or estimated.
e.g. To carry out acid-base potentiometric titration Hydrogen gas. Quinhydrone electrode and glass electrodes are used as indicator electrode.

12. Potentiometric Titrations:
Principle:
Determination of equivalence point of titration on the basis of potential measurement using a suitable set up of galvanic cell is called as potentiometric titration.
Galvanic cell consists of two half cells. One half cell is called as indicator half cell and second half is called reference half cell.
During the potentiometric titration, the potential of indicator half cell changes whereas the potential of the reference half cell remains constant.

13. E = Eo - 2.303 RT nF log10 𝑎(𝑅𝑒𝑑𝑢𝑐𝑒𝑑 𝑆𝑡𝑎𝑡𝑒) 𝑎 (𝑂𝑥𝑖𝑑𝑖𝑠𝑒𝑑 𝑆𝑡𝑎𝑡𝑒)
The electrode potential on the activity or concentration of the ion with which the electrode is reversible is given by the Nernst equation.
E = Eo RT nF log10 𝑎(𝑅𝑒𝑑𝑢𝑐𝑒𝑑 𝑆𝑡𝑎𝑡𝑒) 𝑎 (𝑂𝑥𝑖𝑑𝑖𝑠𝑒𝑑 𝑆𝑡𝑎𝑡𝑒)
Where,
E = Electrode potential of the indicator electrode
Eo = Standard electrode potential of the indicator electrode
n = Number of electrons involved in electrode reaction
R = Gas Constant
F = Faraday Constant (96500 Coulomb)
a(Reduced State) = Product of activities of species involved in reduced state
a(Oxidation State) = Product of activities of species involved in reduced state

14. During the potentiometric titration it is observed that,
Initially the emf changes gradually
Towards the equivalence point it is quite rapidly
After the equivalence point it is changes gradually
The equivalence point can therefore determined by finding the quantity of titrant added at the point at which the rate of change of emf or potential is maximum.

15. ------------------------------------
The emf of the cell after each addition of titrant is measured.
The graph of
E vs V
or ΔE/ΔV vs V
or ΔE2/ΔV2 vs V is plotted from which equivalence point can be determined.
EMF
ΔE/ΔV
Δ2E/Δ2V
Volume of Titrant in cm3
Volume of Titrant in cm3
Volume of Titrant in cm3

16. Role of Reference electrode and Indicator Electrode:
A reference electrode is that electrode whose potential is known and remain constant.
e.g. Saturated calomel electrode (ESCE = 0.242)
Indicator Electrode:
An indicator electrode is that electrode whose potential depends on the activity of ions being titrated or estimated.
e.g. To carry out acid-base potentiometric titration Hydrogen gas. Quinhydrone electrode and glass electrodes are used as indicator electrode.
Indicator electrodes for acid-base titrations
In acid-base titration hydrogen ion is the ion being titrated, its concentration changes during the titration. Hence, in acid-base titration indicator electrode should be an electrode reversible to hydrogen ion.

17. Some examples of Indicator electrode used in acid-base titration.
Quinhydrone electrode:
Representation:
Electrode Reaction:
Pt H2Q, Q, H+(aq)
Expression for the Reaction Q =
aH2Q
aQ . a2H+

18. Expression for the electrode potential:
EQuin =
E0 -
2.303 RT
log10
aH2Q 2F
aQ . a2H+
EQuin =
E0 -
2.303 RT
log10
aH2Q +
log10 a2H+ 2F aQ
EQuin =
E0 +
2.303 RT
log10 a2H+ -
log10
aH2Q 2F aQ
Quinhyrdone is an equimolar mixture of quinone and hydroquinone.
Solubility of both in pure water is very small.
Thus, in an equimolar mixture of the two, the amounts that dissolved can be assumed to be equal.
Therefore, the activity ratio aH2Q/aQ can be assumed to be equal.

19. EQuin =
E0 +
2.303 RT
log10 a2H+ 2F
EQuin =
E0 +
2.303 RT
log10 aH+ F
At 298 K,
EQuin =
E0Quin - pH

20. Construction:
The electrode is set up by adding a small quantity of commercial hydroquinone to the experimental solution so as to saturate it.
This solution is stirred and then an electrode of platinum is introduced for electrical contact. The platinum strip used should be clean.
The electrode can be combined with saturated calomel electrode to form cell, the potential of which can be determined.

21. Hg(l) – HgCl2(s) H+(aq), H2Q, Q Pt
In acidic medium:
The cell set up: Θ Ꚛ
Hg(l) – HgCl2(s) H+(aq), H2Q, Q Pt
Calomel
KCl salt Bridge
Quinhydrone
Ecell = Ecathode – Eanode
Ecell = EQuin – ECalomel
EQuin =
E0Quin pH
ECell =
E0Quin pH - ECalomel

22. ECell = pH
ECell = pH
pH = 𝟎.𝟒𝟓𝟕 − 𝐄𝐜𝐞𝐥𝐥 𝟎.𝟎𝟓𝟗𝟏𝟔

23. Pt H2Q, Q, H+(aq), Hg(l) – HgCl2(s)
b. In alkaline medium:
The cell set up: Ꚛ Θ
Pt H2Q, Q, H+(aq), Hg(l) – HgCl2(s)
Quinhydrone
KCl salt Bridge
Calomel
Ecell = Ecathode – Eanode
Ecell = ECalomel – EQuin
EQuin =
E0Quin pH
ECell =
ECalomel - E0Quin pH
ECell =
0.242 – – pH

24. ECell = pH
pH = 𝐄𝐜𝐞𝐥𝐥 + 𝟎.𝟒𝟓𝟕 𝟎.𝟎𝟓𝟗𝟏𝟔

25. Merits and Demerits of the electrode:
Electrode is easy to set up.
It is also easy to handle.
It can be functioning satisfactorily also in highly acidic solution.
It is used to measure the pH of aqueous and non-aqueous solution.
Demerits:
This electrode is functioning only in the pH range of 1 to 8
With the solution of pH greater than 8, the activity ratio is no longer remain equal to 1.
It cannot be functioning in presence of oxidising and reducing agents that can react rapidly with either hydroquinone or quinone.

26. 2. Glass electrode: Representation:
The glass electrode belongs to category of ion selective electrodes.
It is characterized by presence of a glass membrane which shows selectivity for a particular ion.
The glass electrode used for the determination of pH shows a preferential response for hydrogen ions.
Representation:
Ag(s) – AgCl(s) HCl Test - solution
Glass
Membrane

27. The potential developed by the electrode is proportional to the difference in the hydrogen ion concentration on the both side of the glass membrane.
When the hydrogen ion concentration, on one side of the glass membrane is held constant, then the potential developed by the electrode will become proportional to the hydrogen ion concentration of the test solution.
The Ag(s)-AgCl(s) electrode provides the constant potential and the solution of 0.1N HCl provides the constant hydrogen ion concentration on one side of the glass membrane.

28. Electrode Potential: Cell Set-up: EG = E0Glass - 0.05916 pH
Glass Electrode
Salt Bridge
Calomel Electrode

29. Construction:
The glass electrode consists of a glass tube ended into a glass bulb containing a solution of constant pH and electrode of constant potential (Ag-AgCl).
Special electronic circuits have to be devised for the measurement of the pH of the solution.
Instruments designed for this purpose, that make use of a glass and calomel electrode assembly and provide directly the pH of the solution are known as pH meter.

30. Diagram:

31. Mechanism:
This electrode involves the exchange of hydrogen ions of solution with silver ion of the glass electrode.
The extent of the exchange is depending on the concentration of H+ ion concentration in the respective solutions.
Due to this there is development of potential that is similar to the junction potential at each surface.
As long as the activity of water on the both sides of the glass membrane are equal, the response of the electrode remains proper.
Whenever ions other than hydrogen ion participate in the exchange process, the response of the electrode changes.

32. Merits and Demerits of the electrode:
It provides a measure of pH in the pH range of 1 – 9.
Using a pH meter, pH of the solution can be directly read.
The electrode can be used in all aqueous solutions.
Electrode is not affected by oxidizing and reducing agents or by any organic compound.
pH can be determined even for small volume of solution.

33. Demerits:
The electrode cannot function in highly acidic or alkaline medium
It cannot produce proper response with pH > 9 or <0.5.
It cannot function in non-aqueous medium.
It needs standardisation every time before the use.

34. It is an electrode of constant known potential.
Reference Electrode:
It is an electrode of constant known potential.
This electrode is of two types
Primary Reference Electrode
Secondary Reference Electrode
Primary Reference Electrode:
It is a electrode whose potential under standard conditions is to be defined to be exactly zero.
A Standard Hydrogen Electrode (SHE) is referred as a primary standard electrode.
It consist of a Pt electrode, H2 gas and H+ ions in aqueous medium.
The cell is represented as:
Pt(s) H2(g), 1 atm H+ (aq.), 1M

35. The cell potential for the standard hydrogen electrode is defined to be exactly zero.
Limitations:
The Standard Hydrogen electrode is difficult to set up.
It is also difficult to handle.
It cannot be used in the solution containing oxidising and reducing agents.
It cannot be used also with unsaturated organic compounds and compounds containing sulphur and arsenic.

36. 2. Secondary Reference Electrode:
It is a electrode whose potential is measured against standard hydrogen electrode and can be used to measure potential of other electrodes.
e.g. Calomel electrode, Silver-silver chloride and mercury-mercurous sulphate electrode.
The calomel electrode is easy to work with and can be used in any solution.
The Saturated Calomel Electrode (SCE) is represented as:
Hg-Hg2Cl2(s) KClsat.

37. Calomel Electrode:
It consists of a thin platinum wire sealed through a glass tube which is in contact with mercury and mercurous chloride paste (calomel).
This tube is enclosed in a wider glass tube filled with KCl solution and provided with a porous plug at the base.
The concentration of the KCl solution may be 1M, 0.1M or saturated.
The potential of the Calomel electrode is varies with the concentration of KCl solution.

38. Experimental Set Up:
Diagram:
Constructions:
It consists of pair of electrodes.
One is acting as indicator electrode reversible to ion which is being estimated and second is reference electrode like SCE whose potential is known and remain constant.
Both the electrodes are dipped in the experimental solution.
A stirrer is added to stir the solution after each addition of titrant.
Electrodes are connected to the appropriate terminals of potentiometer.

39. Procedure:
A known volume of the sample solution containing the ions to be titrated is placed in a beaker and indicator electrode reversible to the ions to be titrated is dipped in it along with SCE.
The electrodes then connected to the appropriate terminals of the potentiometer.
The emf of the cell thus set up is noted before adding any amount of the titrant. This Ecell corresponds to the zero volume of the titrant added.
The titrant is added from the burette in small increments and after addition of each increments the solution is stirred and emf is noted.
Initially, large increments of the titrant (1 to 2 cm3) may be added, but near the equivalence point where potential changes suddenly small increments of the titrant (0.1 to 0.2 cm3) should be added at a time.

40. ------------------------------------
Since the potential changes are sharp near to the equivalence point, it is necessary to take large number of observations with small increments of the titrant before and after the equivalence point.
The equivalence point of the titration can be obtained by plotting any one of the graphs viz. E vs V or ΔE/ΔV vs V or ΔE2/ΔV2 vs V.
EMF
ΔE/ΔV
Δ2E/Δ2V
Volume of Titrant in cm3
Volume of Titrant in cm3
Volume of Titrant in cm3

41. ER(cathode) – EL(anode)
Application of Neutralization Reactions with reference to titration of a strong acid against strong base: (Using Quinhydrone Electrode)
Cell Representation: Or Θ Ꚛ
Hg(l) – HgCl2(s) H+(aq), H2Q, Q Pt Θ Ꚛ
Hg(l) – HgCl2(s) Definite vol. of acid solution, H2Q, Q Pt
SCE
Quinhydrone
The emf of the cell is given by
Ecell =
ER(cathode) – EL(anode)
Ecell =
EQuin – ESCE

42. Ecell = EQuin – ESCE Where, ESCE = 0.242 For EQuin Electrode Reaction
H2Q
By applying Nernst Equation
EQuin =
E0 -
2.303 RT
log10
aH2Q 2F
aQ . a2H+
EQuin =
E0 -
2.303 RT
log10
aH2Q +
log10 a2H+ 2F aQ

43. EQuin =
E0 +
2.303 RT
log10 a2H+ -
log10
aH2Q 2F aQ
Quinhyrdone is an equimolar mixture of quinone and hydroquinone.
Solubility of both in pure water is very small.
Thus, in an equimolar mixture of the two, the amounts that dissolved can be assumed to be equal.
Therefore, the activity ratio aH2Q/aQ can be assumed to be equal.
EQuin =
E0 +
2.303 RT
log10 a2H+ 2F
EQuin =
E0 +
2.303 RT
log10 aH+ F
At 298 K,
EQuin =
E0Quin - pH

44. Ecell =
EQuin – ESCE
Ecell =
E0Quin – pH - ESCE
E0Quin
= 0.7
ESCE =
Ecell =
0.7 – pH – 0.242
Ecell =
0.458 – pH
pH = − Ecell

45. Graphical Methods of determining the equivalence point in potentiometric titration of strong acid and strong base:
Consider the titration of HCl (Strong acid) against NaOH (Strong base). In potentiometric titration, emf is measured for different volumes of titrant added. The equivalance point then can be determined graphically by any one of the following method
1. From the graph of Ecell vs Volume of the titrant (NaOH): Normal Curve Method
In this method the emf of the cell is plotted against the volume of titrant added.
The graph of Ecell vs volume of the titrant is always S shaped or sigmoid curve.
If the curve is vertical near the equivalence point, the equivalence point can be determined by bisecting the straight portion of the curve.
EMF
Volume of Titrant in cm3

46. When the curve is symmetrical, the midpoint of the steep portion of the curve is the equivalence point.
However, if the curve is not symmetrical, then the above method will not give correct results. In this case a method of parallel tangents or the circle fitting method is adopted to determine equivalence point.
2. From the graph of 𝚫𝐄 𝚫𝐕 vs Volume of the titrant (NaOH): First Derivative Curve
Plot of 𝚫𝐄 𝚫𝐕 vs volume of titrant added shows two portions the ascending portion representing the increase in 𝚫𝐄 𝚫𝐕 and descending portion representing decrease in 𝚫𝐄 𝚫𝐕 . These two portions are then extra plotted which will intersect at the corresponding to the maximum value of 𝚫𝐄 𝚫𝐕 and this will be the equivalence point.
ΔE/ΔV
Volume of Titrant in cm3

47. ------------------------------------
3. From the graph of 𝚫𝟐𝐄 𝚫𝟐𝐕 vs Volume of the titrant (NaOH): Second Derivative Curve
This method is very accurate method for locating equivalence point of titration.
The method is based on the fact that the point at which the first derivatives is maximum.
At the same point, the second derivative is zero. Therefore, at the equivalence point 𝚫𝟐𝐄 𝚫𝟐𝐕 = 0.
The graph consists of two sets of 𝚫𝟐𝐄 𝚫𝟐𝐕 values.
Δ2E/Δ2V
Volume of Titrant in cm3
One with positive values of 𝚫𝟐𝐄 𝚫𝟐𝐕 and second is with negative values of 𝚫𝟐𝐄 𝚫𝟐𝐕 .
The maximum positive values of 𝚫𝟐𝐄 𝚫𝟐𝐕 is joined with maximum negative values of 𝚫𝟐𝐄 𝚫𝟐𝐕 .
The line cut the volume axis at which 𝚫𝟐𝐄 𝚫𝟐𝐕 is zero which gives Veq.

48. Advantages of Potentiometric Titrations:
All the categories of acid-base titrations can be carried out by potentiometrically.
Acid-base titration in non-aqueous medium can also be carried out.
Precipitation titration can be carried out by using metal-metal ion electrode or metal-metal insoluble salt type electrode.
Complexometric titration can be carried out by using mercury as indicator electrode.
Many redox system can be carried out potentiometrically.
Disadvantages of Potentiometric Titrations:
This method is time consuming.
It requires a large number of observations.
In this titration end point is obtained graphically.

49. Applications of Potentiometric Titrations:
To determine alkalinity and carbonate content in sea water
To determine the phosphoric acid content in aerated drinks.
To determine acetic acid content in commercial vinegar sample.
To determine dissociation constant of dibasic and tribasic acid.
To estimate small charges of CO2 or O2 content in natural waters.