1. ELECTROCHEMISTRY

2. References: Engg.Chemistry by Jain and Jain
Engg.Chemistry by Dr. R.V.Gadag and Dr. A.Nithyananda Shetty
Principles of Physical Chemistry by Puri and Sharma

3. Electrochemistry is a branch of chemistry which deals with the properties and behavior of electrolytes in solution and inter-conversion of chemical and electrical energies.

4. An electrochemical cell can be defined as a single arrangement of two electrodes in one or two electrolytes which converts chemical energy into electrical energy or electrical energy into chemical energy.
It can be classified into two types:
Galvanic Cells.
Electrolytic Cells.

5. Galvanic Cells:
A galvanic cell is an electrochemical cell that produces electricity as a result of the spontaneous reaction occurring inside it. Galvanic cell generally consists of two electrodes dipped in two electrolyte solutions which are separated by a porous diaphragm or connected through a salt bridge. To illustrate a typical galvanic cell, we can take the example of Daniel cell.

6. Daniel Cell.

7. At the anode: Zn → Zn 2+ + 2e-
At the cathode: Cu e- → Cu
Net reaction: Zn(s)+Cu 2+ (aq)→ Zn 2+ (aq)+ Cu(s)

8. ELECTROLYTIC CELL
An electrolytic cell is an electro –chemical cell in which a non- spontaneous reaction is driven by an external source of current although the cathode is still the site of reduction, it is now the negative electrode whereas the anode, the site of oxidation is positive.

10. Representation of galvanic cell.
Anode Representation:
Zn│Zn2+ or Zn ; Zn2+
Zn │ ZnSO4 (1M) or Zn ; ZnSO4 (1M)
Cathode Representation:
Cu2+/Cu or Cu2+ ;Cu
Cu2+ (1M) ; Cu or CuSO4(1M)/Cu
Cell Representation:
Zn │ ZnSO4 (1M)║ CuSO4(1M)/Cu

11. Liquid Junction Potential.
Difference between the electric potentials developed in the two solutions across their interface .
Ej = Ø soln, R - Ø soln,L
Eg: *Contact between two different electrolytes (ZnSO4/ CuSO4).
*Contact between same electrolyte of different concentrations(0.1M HCl / M HCl).

12. Salt Bridge.
The liquid junction potential can be reduced (to about 1 to 2 mV) by joining the electrolyte compartments through a salt bridge.

13. Function Of Salt Bridge.
It provides electrolytic contact between the two electrolyte solutions of a cell.
It avoids or at least reduces junction potential in galvanic cells containing two electrolyte solutions in contact.

14. Emf of a cell.
The difference of potential, which causes a current to flow from the electrode of higher potential to one of lower potential.
Ecell = Ecathode- Eanode
The E Cell depends on:
the nature of the electrodes.
temperature.
concentration of the electrolyte solutions.

15. Standard emf of a cell(Eo cell) is defined as the emf of a cell when the reactants & products of the cell reaction are at unit concentration or unit activity, at 298 K and at 1 atmospheric pressure.

16. The emf cannot be measured accurately using a voltmeter :
As a part of the cell current is drawn,thereby causing a change in the emf.
As a part of the emf is used to overcome the internal resistance of the cell.

18. The emf of the cell Ex is proportional to the length AD.
Ex α AD
The emf of the standard cell Es is proportional to the length AD1.
Es α AD1
Ex ═   AD
Es AD1
Ex = AD x Es
AD1

19. Standard Cell.
It is one which is capable of giving constant and reproducible emf.
It has a negligible temperature coefficient of the emf.
The cell reaction should be reversible.
It should have no liquid junction potential.
Eg: Weston Cadmium Cell. The emf of the cell is V at 293 K and V at 298 K.

20. Weston Cadmium Cell Sealed wax Cork Soturated solution of CdSO4.8/3H2O
crystals
Paste of Hg2SO4
Cd-Hg
12-14% Cd
Mercury, Hg

21. Cell representation:
Cd-Hg/Cd2+// Hg2SO4/Hg
At the anode:
Cd (s) → Cd2+ + 2e-
At the cathode:
Hg2SO4(s) + 2e- → 2 Hg (l)+ SO42-(aq)
Cell reaction:
Cd + Hg22+ → Cd2+ + 2Hg

22. Origin of single electrode potential.
Consider Zn(s)/ ZnSO4
Anodic process: Zn(s) → Zn2+(aq)
Cathodic process: Zn2+(aq) → Zn(s)
At equilibrium: Zn(s) ↔ Zn2+(aq)
Metal has net negative charge and solution has equal positive charge leading to the formation of an Helmholtz electrical layer.

23. Single electrode potential.
Electric layer on the metal has a potential Ø (M).
Electric layer on the solution has a potential Ø (aq)
Electric potential difference between the electric double layer existing across the electrode /electrolyte interface of a single electrode or half cell.

24. Helmholtz double layer
De-electronation
Electronation
Helmholtz double layer

25. MEASUREMENT OF ELECTRODE POTENTIAL.
It is not possible to determine experimentally the potential of a single electrode.
It is only the difference of potentials between two electrodes that we can measure by combining them to give a complete cell.
By arbitrarily fixing the potential of reversible hydrogen electrode as zero it is possible to assign numerical values to potentials of the various other electrodes.

26. Sign Of Electrode Potential.
The electrode potential of an electrode:
Is positive: If the electrode reaction is reduction when coupled with the standard hydrogen electrode
Is negative: If the electrode reaction is oxidation when coupled with standard hydrogen electrode. According to latest accepted conventions, all single electrode potential values represent reduction tendency of electrodes.

27. when copper electrode is combined with SHE, copper electrode acts as cathode and undergoes reduction hydrogen electrode acts as anode.
H2(g) → 2H+ +2e- (oxidation)
Cu2+ +2e- → Cu (reduction)
Hence electrode potential of copper is assigned a positive sign. Its standard electrode potential is 0.34 V.

28. When zinc is coupled with S. H. E
When zinc is coupled with S.H.E. zinc electrode acts as anode and hydrogen electrode acts as cathode.
Zn → Zn2+ +2e-
2H+ + 2e-→ H2.
Hence, electrode potential of zinc is negative. The standard electrode potential of zinc electrode is V.

29. Nernst Equation.
It is a quantitative relationship between electrode potential and concentration of the electrolyte species.
Consider a general redox reaction:
Mn+(aq) + ne- → M(s) ----(1)
We know that, ΔG =-nFE (2)
ΔGo=-nFEo-----(3)
ΔG =ΔGo +RT ln K

30. ΔG =ΔGo +RT ln K
ΔG =ΔGo +RT ln[M]/[Mn+]-----(4)
-nFE= -nFEo + RT ln [M]/[Mn+]----(5)
E= Eo – RT/nF ln 1/[Mn+]------(6)
E=Eo RT/nF log 1/[Mn+]---(7)
At 298K,
E= Eo /n log 1/[Mn+] (8)

31. problems
1. A galvanic cell consists of copper plate immersed in 10 M solution of CuSO4 and iron plate immersed in 1M FeSO4 at 298K. If E0cell=0.78 V, write the cell reaction and calculate E.M.F. of the cell.

32. Solution:
Cell reaction:
Fe + Cu ↔ Fe Cu
ECell= E0Cell /2 log [Fe2+ ]/[Cu2+]
ECell= log 10/1
=0.8096V

33. Calculate E. M. F. of the zinc – silver cell at 25˚C when [Zn2+ ] = 1
Calculate E.M.F. of the zinc – silver cell at 25˚C when [Zn2+ ] = 1.0 M and [Ag+] = 10 M (E0cell=1.56V at 25˚C). Write the cell representation and cell reaction

34. Solution:
Cell representation
Zn/ Zn2+((1M)//Ag+(10M) /Ag
Cell reaction:
Zn Ag ↔ Zn Ag
ECell= E0Cell /2 log [Zn2+ ]/[Ag+]2
ECell= log 10/1.0
= V

35. The emf of the cell
Mg│ Mg 2+ (0.01M)║ Cu 2+ /Cu is measured to be 2.78 V at 298K. The standard eletrode potential of magnesium electrode is V. Calculate the electrode potential of copper electrode

36. Cell reaction:
Mg + Cu ↔ Mg Cu
E= Eo /n log 1/[Mn+]
EMg= EoMg /2 log 1/[Mg2+]
= V
Ecell=ECu-EMg
2.78 = ECu-[-2.429]
ECu =
= V

37. The emf of the cell
Cu│ Cu 2+ (0.02M)║ Ag+ /Ag is measured to be 0.46 V at 298K. The standard eletrode potential of copper electrode is 0.34 V. Calculate the electrode potential of silver. electrode

38. Energetics of Cell Reactions.
Net electrical work performed by the cell reaction of a galvanic cell:
W= QE (1)
Charge on 1mol electrons is F(96,500)Coulombs.
When n electrons are involved in the cell reaction,
the charge on n mole of electrons = nF

39. Q = nF
Substituting for Q in eqn (1)
W = nFE (2)
The cell does net work at the expense of
ΔG accompanying. ΔG = -nFE
ΔG = nFE

40. From Gibbs – Helmholtz equation.
ΔG = ΔH + T [δ(ΔG)/ δT]P (2)
-nFE = ΔH – nFT (δ E/ δT)P
ΔH = nFT (δ E/ δ T)P – nFE
ΔH = nF[T(δ E/ δT)P –E]
We know that, [δ (∆G)/ δT]P = - ΔS
ΔS = nF (δE/ δT)P

41. Problem: Emf of Weston Cadmium cell is 1. 0183 V at 293 K and 1
Problem: Emf of Weston Cadmium cell is V at 293 K and 1.0l81 V at 298 K.
Calculate ∆G, ΔH and ΔS of the cell reaction at 298 K.
Solution:- ∆G: ∆G = - n FE
n = 2 for the cell reaction; F = 96,500 C E= V at 298 K
∆G = -2 x 96,500 x J = KJ

42. ∆H: ∆H = nF [ T (δE /δT)P – E]
= VK-1
T = 298 K
∆H = 2 x 96,500 { 298 x ( ) – )
= KJ
ΔS: ΔS = nF (δE / δT) P
= 2 x 96,500 x ( ) = -7.72JK-1

43. Classification of Electrodes.
Gas electrode ( Hydrogen electrode).
Metal-metal insoluble salt (Calomel electrode).
Ion selective electrode.(Glass electrode).

44. Gas electrode.
It consists of gas bubbling over an inert metal wire or foil immersed in a solution containing ions of the gas.
Standard hydrogen electrode is the primary reference electrode, whose electrode potential at all temperature is taken as zero arbitrarily.

45. Construction.

46. Representation: Pt,H2(g)/ H+
Electrode reaction: H+ + e- 1/2 H2(g)
The electrode reaction is reversible as it can undergo either oxidation or reduction depending on the other half cell.
If the concentration of the H+ ions is 1M, pressure of H2 is 1atm at 298K it is called as standard hydrogen electrode (SHE).

47. Applications.
To determine electrode potential of other unknown electrodes.
To determine the pH of a solution.
E=Eo RT/nF log [H2]1/2/[H+]
= log 1/[H+]
= pH.
Cell Scheme: Pt,H2,H+(x)// SHE

48. The emf of the cell is determined.
E (cell) = E (c) – E(A)
= 0 – ( pH)
E (cell) = pH
pH = E(cell)/

49. Limitations. Constuction and working is difficult.
Pt is susceptible for poisoning.
Cannot be used in the presence of oxidising agents.

50. Metal –metal salt ion electrode.
These electrodes consist of a metal and a sparingly soluble salt of the same metal dipping in a solution of a soluble salt having the same anion.
Eg: Calomel electrode.
Ag/AgCl electrode.

51. Construction.

52. Representation: Hg; Hg2Cl2 / KCl
It can act as anode or cathode depending on the nature of the other electrode.
As anode: 2Hg + 2Cl- → Hg2Cl2 + 2e-
As Cathode: Hg2Cl2 + 2e- → 2Hg + 2 Cl-

53. E = Eo – RT/2F log [Cl-)2
= Eo log [Cl-] at 298 K
Its electrode potential depends on the concentration of KCl.
Conc. of Cl Electrode potential
0.1M V
1.0 M V
Saturated V

54. Applications.
Since the electrode potential is a constant it can be used as a secondary reference electrode.
To determine electrode potential of other unknown electrodes.
To determine the pH of a solution.
Pt,H2/H+(X) // KCl,Hg2Cl2,Hg
pH = E(cell) – /

55. Ion Selective Electrode.
It is sensitive to a specific ion present in an electrolyte.
The potential of this depends upon the activity of this ion in the electrolyte.
Magnitude of potential of this electrode is an indicator of the activity of the specific ion in the electrolyte.
*This type of electrode is called indicator electrode.

56. Glass Electrode:`

57. a sensing part of electrode, a bulb made from a specific glass
Scheme of typical pH glass electrode
a sensing part of electrode,
a bulb made from a specific glass
sometimes electrode contain small amount
of AgCl precipitate inside the glass
electrode
3 internal solution, usually 0.1M HCl for pH
electrodes
4.internal electrode, usually silver chloride
electrode or calomel electrode
5.body of electrode, made from non-
conductive glass or plastics.
6.reference electrode, usually the same type
as 4
7.junction with studied solution, usually made from ceramics or capillary with asbestos or quartz fiber.

58. The hydration of a pH sensitive glass membrane involves an ion-exchange reaction between singly charged cations in the interstices of the glass lattice and protons from the solution.
H Na Na H+
Soln glass soln glass
Eg = Eog – pH

59. Electrode Potential of glass electrode.
The overall potential of the glass electrode has three components:
The boundary potential Eb,
Internal reference electrode potential Eref.
Asymetric potential Easy.- due to the difference in response of the inner and outer surface of the glass bulb to changes in [H+].
Eg = Eb + Eref. + Easy.

60. Eb = E1 – E2
= RT/nF ln C1 – RT/nF ln C2
= L + RT/nF ln C1
Eb depends upon [H+]
Eg = Eb + EAg/AgCl + Easy.
= L + RT/nF ln C1 + EAg/AgCl + Easy.
= Eog + RT/nF ln C1
= Eog log [H+]
Eg = Eog – pH.

61. Advantages:
It can be used without interference in solutions containing strong oxidants, strong reductants, proteins, viscos fluids and gases as the glass is chemically robust.
It can be used for solutions having pH values 2 to 10. With some special glass (by incorporation of Al2O3 or B2O3) measurements can be extended to pH values up to 12.
It is immune to poisoning and is simple to operate
The equilibrium is reached quickly & the response is rapid

62. 5. It can be used for very small quantities of the solutions
5. It can be used for very small quantities of the solutions. Small electrodes can be used for pH measurement in one drop of solution in a tooth cavity or in the sweat of the skin (micro determinations using microelectrodes)
6. If recently calibrated, the glass electrode gives an accurate response.
7. The glass electrode is much more convenient to handle than the inconvenient hydrogen gas electrode.

63. Disadvantages:
The bulb of this electrode is very fragile and has to be used with great care.
The alkaline error arises when a glass electrode is employed to measure the pH of solutions having pH values in the range or greater. In the presence of alkali ions, the glass surface becomes responsive to both hydrogen and alkali ions. Low pH values arise as a consequence and thus the glass pH electrode gives erroneous results in highly alkaline solutions.

64. The acid error results in highly acidic solutions (pH less than zero)Measured pH values are high.
Dehydration of the working surface may cause erratic electrode performance. It is crucial that the pH electrode be sufficiently hydrated before being used. When not in use, the electrode should be stored in an aqueous solution because once it is dehydrated, several hours are required to rehydrate it fully.

65. As the glass membrane has a very high electrical resistance (50 to 500 mΩ), the ordinary potentiometer cannot be used for measurement of the potential of the glass electrode. Thus special electronic potentiometers are used which require practically no current for their operation.

66. Standardization has to be carried out frequently because asymmetry potential changes gradually with time. Because of an asymmetry potential, not all glass electrodes in a particular assembly have the same value of EoG . For this reason, it is best to determine EoG for each electrode before use.
The commercial verson is moderately expensive

67. Limitations.
The bulb is very fragile and has to be used with great care.
In the presence of alkali ions, the glass surface becomes responsive to both hydrogen and alkali ions. Measured pH values are low.
In highly acidic solutions (pH less than zero) measured pH values are high.
When not in use, the electrode should be stored in an aqueous solution.

68. Applications. Determination of pH: Cell: SCE ║Test solution / GE
E cell = Eg – Ecal.
E cell = Eog – pH –
pH = Eog -Ecell – Ecal. /

69. Problems The cell SCE ΙΙ (0.1M) HCl Ι AgCl(s) /Ag
gave emf of 0.24 V and 0.26 V with buffer having pH value 2.8 and unknown pH value respectively. Calculate the pH value of unknown buffer solution. Given ESCE= V

70. Eog= pH +Ecell + Ecal.
= x
=0.648 V
pH = Eog -Ecell – Ecal. /
= /0.0592
= 2.46

71. CONCENTRATION CELLS.
Two electrodes of the same metal are in contact with solutions of different concentrations.
Emf arises due to the difference in concentrations.
Cell Representation:
M/ Mn+[C1] ║ Mn+/M[C2]

72. Construction.

73. At anode: Zn →Zn2+(C1) + 2e-
At cathode: Zn2+(C2) + 2e-→ Zn
Ecell = EC-EA
= E0 + (2.303RT/ nF)logC [E0+(2.303RT/nF)logC1]
Ecell = (0.0592/n) log C2/C1
Ecell is positive only if C2 > C1

74. Anode - electrode with lower electrolyte concentration.
Cathode – electrode with higher electrolyte concentration.
Higher the ratio [C2/C1] higher is the emf.
Emf becomes zero when [C1] = [C2].

75. Problems
Zn/ZnSO4(0.001M)||ZnSO4(x)/Zn is 0.09V at 25˚C. Find the concentration of the unknown solution.

76. Ecell = /n log C2/C1
0.09 =(0.0592/2) log ( x / 0.001)
x =1.097M

77. 2. Calculate the valency of mercurous ions
with the help of the following cell.
Hg/ Mercurous || Mercurous /Hg
nitrate (0.001N) nitrate (0.01N) when the emf observed at 18˚ C is V
Ecell=(2.303 RT/nF) log C2/C1

78. Ecell=(2.303 RT/nF) log C2/C1
0.029 = 2.303RT/n) log (0.01/0.001)
0.029 =0.057 x 1/ n
n = 0.057/ =̃ 2
Valency of mercurous ions is 2, Hg2 2+

79. Assignment Answer the following questions:
1.Distinguish between electrolytic and galvanic cells.
2.Explain the origin of electrode potential. What are the sign conventions for electrode potential?
3.Give reasons for the following.
i) The glass electrode changes its emf over a period
of time.
ii) KCl is preferred instead of NaCl as an electrolyte
in the preparation of salt bridge
4. What is meant by a standard cell? Give an example

80. 5. Quote any four limitations of glass electrode
6.Define liquid junction potential. How it can be eliminated or minimized?
7.Derive Nernst equation for the single electrode potential.
8.Describe potentiometric determination of emf of a cell.
9.Writ e construction and working of Calomel Electrode
10.What are concentration cells? Show that emf of concentration cell becomes zero at a certain point of its working.