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ELECTROCHEMISTRY

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Title: ELECTROCHEMISTRY


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 2 2e- ? 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.
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10
Representation of galvanic cell.
  • Anode Representation
  • ZnZn2 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 / 1.0 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.

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  • The emf of the cell Ex is proportional to the
    length AD.
  • Ex a AD
  • The emf of the standard cell Es is proportional
    to the length AD1.
  • Es a 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
    1.0183 V at 293 K and 1.0181 V at 298 K.

20
  • Weston Cadmium Cell

Sealed wax
Cork
Soturated solution of CdSO4.8/3H2O
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
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. 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 -0.74 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, ?G -nFE ----- (2)
  • ?Go-nFEo-----(3)
  • ?G ?Go RT ln K

30
  • ?G ?Go RT ln K
  • ?G ?Go RT lnM/Mn-----(4)
  • -nFE -nFEo RT ln M/Mn----(5)
  • E Eo RT/nF ln 1/Mn------(6)
  • EEo- 2.303 RT/nF log 1/Mn---(7)
  • At 298K,
  • E Eo-0.0592/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 E0cell0.78 V,
    write the cell reaction and calculate E.M.F. of
    the cell.

32
  • Solution
  • Cell reaction
  • Fe Cu2 ? Fe2 Cu
  • ECell E0Cell-0.0592/2 log Fe2 /Cu2
  • ECell 0.78 0.0296 log 10/1
  • 0.8096V

33
  • Calculate E.M.F. of the zinc silver cell at
    25C when Zn2 1.0 M and Ag 10 M
    (E0cell1.56V at 25C). Write the cell
    representation and cell reaction

34
  • Solution
  • Cell representation
  • Zn/ Zn2((1M)//Ag(10M) /Ag
  • Cell reaction
  • Zn 2Ag ? Zn2 2Ag
  • ECell E0Cell-0.0592/2 log Zn2 /Ag2
  • ECell 1.56 0.0592 log 10/1.0
  • 1.6192 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 -2.37 V.
    Calculate the electrode potential of copper
    electrode

36
  • Cell reaction
  • Mg Cu2 ? Mg2 Cu
  • E Eo-0.0592/n log 1/Mn
  • EMg EoMg-0.0592/2 log 1/Mg2
  • -2.4291V
  • EcellECu-EMg
  • 2.78 ECu--2.429
  • ECu 2.78-2.429
  • 0.3509 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
  • ?G accompanying. ?G -nFE
  • ?G nFE

40
  • From Gibbs Helmholtz equation.
  • ?G ?H T d(?G)/ dTP ------- (2)
  • -nFE ?H nFT (d E/ dT)P
  • ?H nFT (d E/ d T)P nFE
  • ?H nFT(d E/ dT)P E
  • We know that, d (?G)/ dTP - ?S
  • ?S nF (dE/ dT)P

41
  • Problem Emf of Weston Cadmium cell is 1.0183 V
    at 293 K and 1.0l81 V at 298 K.
  • Calculate ?G, ?H and ?S of the cell reaction
    at 298 K.
  • Solution- ?G ?G - n FE
  • n 2 for the cell reaction F 96,500 C
    E 1.0181 V at 298 K
  • ?G -2 x 96,500 x 1.0181 J -196.5 KJ

42
  • ?H ?H nF T (dE /dT)P E
  • (dE/dT)p 1.0181 1.0183 / 298-293 -0.0002 /
    5
  • -0.00004VK-1
  • T 298 K
  • ?H 2 x 96,500 298 x (-0.00004) 1.0181)
  • -198. 8 KJ
  • ?S ?S nF (dE / dT) P
  • 2 x 96,500 x (0-00004) -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.
  • EEo- 2.303 RT/nF log H21/2/H
  • 0 -0.0591 log 1/H
  • -0.0591pH.
  • Cell Scheme Pt,H2,H(x)// SHE

48
  • The emf of the cell is determined.
  • E (cell) E (c) E(A)
  • 0 (- 0.0592 pH)
  • E (cell) 0.0592 pH
  • pH E(cell)/ 0.0592

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 2.303 RT/2F log Cl-)2
  • Eo -0.0591 log Cl- at 298 K
  • Its electrode potential depends on the
    concentration of KCl.
  • Conc. of Cl- Electrode potential
  • 0.1M
    0.3335 V
  • 1.0 M 0.2810 V
  • Saturated 0.2422 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) 0.2422/ 0.0592

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
  • 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 0.0592 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 0.0592 log H
  • Eg Eog 0.0592 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. 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 10-12 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 mO), 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 0.0592 pH 0.2422
  • pH Eog -Ecell Ecal. / 0.0592

69
Problems
  • The cell SCE ?? (0.1M) HCl ? 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 0.2422 V

70
  • Eog 0.0592pH Ecell Ecal.
  • 0.0592x2.8 0.24 0.2422
  • 0.648 V
  • pH Eog -Ecell Ecal. / 0.0592
  • 0.648 -0.26-0.2422/0.0592
  • 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/ MnC1 Mn/MC2

72
Construction.

73
  • At anode Zn ?Zn2(C1) 2e-
  • At cathode Zn2(C2) 2e-? Zn
  • Ecell EC-EA
  • E0 (2.303RT/ nF)logC2-
    E0(2.303RT/nF)logC1
  • Ecell (0.0592/n) log C2/C1
  • Ecell is positive only if C2 gt 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 25C.
    Find the concentration of the unknown solution.

76
  • Ecell 0.0592/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 0.029 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 0.057/0.029 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.
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