ELECTROCHEMISTRY

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ELECTROCHEMISTRY
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• 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

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• 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.

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• 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.

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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.
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Daniel Cell.

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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)
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• 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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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

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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).

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Salt Bridge.

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

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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.

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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.

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• 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.

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• 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

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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.

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• 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
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• 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

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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.

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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.

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De-electronation
Electronation
Helmholtz double layer
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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.

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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.

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• 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.

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• 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.

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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

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• ?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)

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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.

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• 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

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• 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

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• 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

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• 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

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• 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

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• 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

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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

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• 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

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• 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

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• 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

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• ?H ?H nF T (dE /dT)P E
• (dE/dT)p 1.0181 1.0183 / 298-293 -0.0002 /
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• -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

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Classification of Electrodes.

• Gas electrode ( Hydrogen electrode).
• Metal-metal insoluble salt (Calomel electrode).
• Ion selective electrode.(Glass electrode).

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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.

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Construction.
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• 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).

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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

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• 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

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Limitations.

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

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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.

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Construction.

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• 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-

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• 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

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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

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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.

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• Glass Electrode

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• 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.

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• 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

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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.

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• 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.

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• 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

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• 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.

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• 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.

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• 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.

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• 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.

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• 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

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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.

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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

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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

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• 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

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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

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Construction.

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• 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

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• 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.

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Problems

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

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• Ecell 0.0592/n log C2/C1
• 0.09 (0.0592/2) log ( x / 0.001)
• x 1.097M

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• 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

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• 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

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• 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

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• 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.