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ELECTROCHEMISTRY. 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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slide2
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
slide3
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.
slide4
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.
slide5

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.

slide7

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)

slide8
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.

representation of galvanic cell
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

liquid junction potential
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).

salt bridge
Salt Bridge.
  • The liquid junction potential can be reduced (to about 1 to 2 mV) by joining the electrolyte compartments through a salt bridge.
function of salt bridge
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.
emf of a cell
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.
slide15
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.
slide16
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.
slide18
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

standard cell
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.

slide20
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

slide21
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

origin of single electrode potential
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.

single electrode potential
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.
slide24

Helmholtz double layer

De-electronation

Electronation

measurement of electrode potential
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.
sign of electrode potential
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.

slide27
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.

slide28
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.

nernst equation
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

slide30
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- 2.303 RT/nF log 1/[Mn+]---(7)

At 298K,

E= Eo-0.0592/n log 1/[Mn+]-------(8)

slide31

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.
slide32
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

slide33
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
slide34
Solution:
  • Cell representation
  • Zn/ Zn2+((1M)//Ag+(10M) /Ag
  • Cell reaction:

Zn + 2Ag+↔ Zn2+ + 2Ag

ECell= E0Cell-0.0592/2 log [Zn2+ ]/[Ag+]2

ECell= 1.56 + 0.0592 log 10/1.0

=1.6192 V

slide35
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

slide36
Cell reaction:

Mg + Cu2+↔ Mg2+ + Cu

E= Eo-0.0592/n log 1/[Mn+]

EMg= EoMg-0.0592/2 log 1/[Mg2+]

=-2.4291V

Ecell=ECu-EMg

2.78 = ECu-[-2.429]

ECu=2.78-2.429

=0.3509 V

slide37
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

energetics of cell reactions
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

slide39
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
slide40
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

slide41
Problem: Emf of Weston Cadmium cell is 1.0183 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= 1.0181 V at 298 K

∆G = -2 x 96,500 x 1.0181 J = -196.5 KJ

slide42
∆H: ∆H = nF [ T (δE /δT)P – E]
  • (δE/δT)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
  • ΔS: ΔS = nF (δE / δT) P
  • = 2 x 96,500 x (0.00004) = -7.72JK-1
classification of electrodes
Classification of Electrodes.
  • Gas electrode ( Hydrogen electrode).
  • Metal-metal insoluble salt (Calomel electrode).
  • Ion selective electrode.(Glass electrode).
gas electrode
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.
slide46
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).
applications
Applications.
  • To determine electrode potential of other unknown electrodes.
  • To determine the pH of a solution.

E=Eo- 2.303 RT/nF log [H2]1/2/[H+]

= 0 -0.0591 log 1/[H+]

= -0.0591pH.

Cell Scheme: Pt,H2,H+(x)// SHE

slide48
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

limitations
Limitations.
  • Constuction and working is difficult.
  • Pt is susceptible for poisoning.
  • Cannot be used in the presence of oxidising agents.
metal metal salt ion electrode
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.

slide52
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-
slide53
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

applications1
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

ion selective electrode
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.

slide57

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.
slide58
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

electrode potential of glass electrode
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.

slide60
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.

slide61
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
slide62
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.

slide63
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.
slide64
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.
slide65
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.
slide66
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
limitations1
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.
applications2
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

problems
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= 0.2422 V

slide70
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

concentration cells
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]

slide73
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 > C1

slide74
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].
problems1
Problems
  • Zn/ZnSO4(0.001M)||ZnSO4(x)/Zn is 0.09V at 25˚C. Find the concentration of the unknown solution.
slide76
Ecell = 0.0592/n log C2/C1

0.09 =(0.0592/2) log ( x / 0.001)

x =1.097M

slide77
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

slide78
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+

slide79
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

slide80
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.