1. ELECTRO CHEMISTRY
UNIT -I

2. Contents: Conductance-Electrolyte in solution
Conductance-Specific, Equivalent and Molar conductance
Ionic mobilities
Kolrausch’s Law
Application of conductance

3. EMF
Galvanic Cells
Types of Electrodes
Nernst equation
Concentration Cells
Galvanic series
Potentiometric titrations
Numerical problems.
Batteries
Fuel cells

4. Conductance
The reciprocal of the resistance is called conductance. It is denoted by C.
C = 1 / R
Conductors allows electric current to pass through them. Examples are metals, aqueous solution of acids, bases and salts etc.
Unit of conductance is ohm-1 or mho or Siemen (S)

5. Classification of Conductors
These may be divided into three main categories; they are:
I) Gaseous II) Metallic or electronic III) Electrolytic.
Gases conduct electricity with difficulty and only under the influence of high potentials or if exposed to the action of certain radiations.

6. Metallic or electronic conductors :
Conductors which transfer electric current by transfer of electrons, without transfer of any matter, are known as metallic or electronic conductors. Metals such as copper, silver, aluminum, etc., non-metals like carbon (graphite - an allotropic form of carbon) and various alloys belong to this class.

7. Classification of Conductors….
Electrolytic conductors : (a) Conductors like aqueous solutions of acids, bases and salts in which the flow of electric current is accompanied by chemical decomposition are known as electrolytic conductors.
b)The substances whose aqueous solutions do not conduct electric current are called non-electrolytes. Solutions of cane sugar, glycerine, alcohol, etc., are examples of non-electrolytes.

8. Electrolytes: Substances whose solution in water conducts electric
current. Conduction takes place by the movement of ions.
Examples are salts, acids and bases.
Substances whose aqueous solution does not conduct electricity are called non electrolytes.
Examples are solutions of cane sugar, glucose, urea
etc.

9. Types of Electrolytes
Strong electrolyte are highly ionized in the solution.
Examples are HCl, H2SO4, NaOH, KOH etc
Weak electrolytes are only feebly ionized in the solution.
Examples are H2CO3, CH3COOH, NH4OH etc

10. Difference between electronic & electrolytic conductors
(3) Conduction increases with increase in temperature
(3) Conduction decreases with increase in temperature
(2) Flow of electricity is due to the movement of ions
(2) Conduction is due to the flow of electron
(1)Flow of electricity takes place by the decomposition of the substance.
(1) Flow of electricity take place without the decomposition of substance.
Electrolytic conductors
Electronic conductors

11. Specific Resistance:
Resistance refers to the opposition to the flow of current.
For a conductor of uniform cross section (a) and length (l); Resistance R,
Where ‘ ρ’ is called resistivity or specific resistance.

12. Specific conductance or conductivity
Conductance of unit volume of cell is specific conductance.
   If l = 1 cm and a = 1 cm2, then
R = ρ                            
Κ= 1/ρ, Κ = kappa - the specific conductance          
ρ = a/l. R or 1/ρ = 1/a.1/R
K = 1/a×C (1/z = cell constant)
Specific conductance = cell constant x Conductance
The unit of specific conductance is ohm-1 cm-1.

13. Specific conductance or conductivity
Representation of specific conductance
Specific conductance depend on the nuber of ions present in unit volume (1 ml ) solution

14. Equivalent Conductance(λ / /\):
It is the conductance of one gram equivalent of the electrolyte dissolved in V cc of the solution.
/\ = KV
In case, if the concentration of the solution is c g equivalent per litre, then the volume containing 1 g equivalent of the electrolyte will be 1000/C.
So equivalent conductance
/\ = K / c
/\ = k × 1000/N
         Where N = normality
                The unit of equivalent conductance is ohm-1 cm-2 equi-1.

15. Molar conductance
The molar conductance is defined as the conductance of all the ions produced by ionization of 1 g mole of an electrolyte when present in V mL of solution. It is denoted by.
Molar conductance     Λ m = k ×V  
                              
Where V is the volume in mL containing 1 g mole of the electrolyte. If c is the concentration of the solution in g mole per litre, then
Λ m = k × 1000/c
It units are ohm-1 cm2 mol-1.
              
Equivalent conductance =  (Molar conductance)/n
               Where            n = (Molecular mass) / (Equivalent mass)

16. Effect of Dilution on Conductivity
Specific conductivity decreases on dilution.
Equivalent and molar conductance both increase with dilution and reaches a maximum value.
The conductance of all electrolytes increases with temperature.

17. Illustrative Example
The resistance of 0.01N NaCl solution at 250C is 200 ohm. Cell constant of conductivity cell is unity. Calculate the equivalent conductance and molar conductance of the solution.
Solution:
Conductance of the cell=1/resistance
=1/200
=0.005 S.
Specific conductance=conductance x cell constant
=0.005 x 1
=0.005 S cm-1

18. Equivalent Conductance = Specific conductance x (1000/N)
Solution Cont.
Equivalent Conductance = Specific conductance x (1000/N)
= x 1000/0.01
= 500 ohm-1cm2eq-1
Molar Conductivity = Equivalent conductivity x n-factor
= 500 x 1
= 500 ohm-1mol-1cm2

19. Kohlrausch’s Law “Limiting molar conductivity of an electrolyte can be
represented as the sum of the individual contributions of
the anion and cation of the electrolyte.”
Where are known as ionic conductance of anion and cation at infinite dilution respectively.

20. Application of Kohlrausch’s law
1. It is used for determination of degree of dissociation of a weak electrolyte.
Where,
represents equivalent conductivity at infinite dilution.
represents equivalent conductivity at dilution v.
2. For obtaining the equivalent conductivities of weak electrolytes at infinite dilution.
3. Determination of the solubility of a sparingly soluble salt

21. Thus conductance is given by product of charge and velocity.
Ionic Mobilities:
Equivalent conductivity of an electrolyte at any dilution is directly proportional to the charged carried by the ions and their velocities.
Thus conductance is given by product of charge and velocity.
At infinite dilution charge is constant
( because of complete ionisation) therefore at infinite dilution equivalent conductivity depends only on ionic velocities.

22. where u = ionic velocity , u – and u + are the velocities of ions
Ionic mobilities cont..
Λ∞ / λ ∞ α u
λ ∞ = K u
Λ+ α u + or Λ- α u –
λ + = K u+ or Λ = K u –
where u = ionic velocity , u – and u + are the velocities of ions
λ ∞ = equivalent conductivity at infinite dilution
K is proportionality constant

23. Electromotive Force
The maximum potential difference between the electrodes of a voltaic cell is referred to as the electromotive force (emf) of the cell, or Ecell.
It can be measured by an electronic digital voltmeter which draws negligible current. 23 2

24. Electrons are driven (“pushed”) through conducting wire in the direction of anode  cathode by cell force
Origin of cell force is maximum electric potential difference between electrodes or electromotive force (Ecell) or cell potential
Potential difference - difference in electrical potential (electrical pressure) between two electrodes; standard unit of cell potential difference is the Volt 24

25. "9"
Calomel Electrode
Hg0l , Hg2Cl2 (sat’d) , KCl (X M = 0.09M)

26. Saturated Calomel Electrode
H2 (1 bar)
potentiometer, eo = V
KCl / agar salt bridge
inert Pt (s) electrode
2 Cl-
2 H+
2 e-
KCl (sat aq)
KCl (s)
Hg2Cl2 (s)
Hg (l)
Pt wire

27. (½ Hg2Cl2 (s) + e – ----------> Hg (l) + Cl – aq)
In standard cell notation this cell would be described as
Pt / H2 (g) / soln with a H+ // KCl (sat) / Hg2Cl2 (s) / Hg (l) / Pt
The anode half-cell reaction is
½ H2 (g) > H + (aq) + e –
The cathode half-cell reaction is
(½ Hg2Cl2 (s) + e – > Hg (l) + Cl – aq)
The net cell reaction is
½ H2 (g) + ½ Hg2Cl2 (s) > Hg (l) + Cl – (aq) + H + (aq)

28. E left = E0CE + 0.0591 / n log [Hg2Cl2] / [Hg0]2[Cl-]2
Anodic Side
Hg2Cl ↔ Hg Cl-
Hg e- ↔ 2Hg0
Hg2Cl2+ 2e- ↔ 2Hg0 + 2Cl-
E left = E0CE / n log [Hg2Cl2] / [Hg0]2[Cl-]2
= ( / 2) log [1/1*(0.09)2]
= volt.

29. Ion Selective Electrode: Ex: Glass Electrode

30. The glass electrode is hydrogen ion responsive electrode.
A very thin membrane of the special glass is sealed on to the end of a glass tube. Which is smooth and square.
The membrane potential can be developed by exchange of ions between glass membrane and the solution.
Cell representation
Ag(s)AgCl(s) Cl- //glass membraneelectrolyte
E cell = E GE – E SCE

31. Quinhydrone Electrode
It is a secondary electrode
Quinhydrone is a combination of quinone and hydroquinone
Q + 2H + + 2e-  H2Q
Pt (s) quinhydrone (test solution) // KCl (sat) / Hg2Cl2 (s) / Hg (l)

32. E QE = E0QE + (0.0591/n) log [Q]*[H+]2 / [H2Q]
= /2 log [H+]2
= log [H+]
= pH

33. -nFEcell = -nFE°cell + RTln(Q)
Nernst Equation
Recall, in general:
DG = DG° + RTln(Q)
However:
DG = -nFEcell
-nFEcell = -nFE°cell + RTln(Q)
Ecell = E°cell - (RT/nF)ln(Q)
Ecell = E°cell - (0.0591/n)log(Q)
The Nernst Equation

34. Concentration Cells Consider the cell presented on the left.
The 1/2 cell reactions are the same,
it is just the concentrations that differ.

35. Concentration Cells (cont.)
Ag+ + e- Ag E°1/2 = 0.80 V
• What if both sides had 1 M
concentrations of Ag+?
• E°1/2 would be the same;
therefore, E°cell = 0.

36. Concentration Cells (cont.)
Anode:
Ag Ag+ + e- E1/2 = ? V
Cathode:
Ag+ + e Ag E1/2 = 0.80 V
Ecell = E°cell - (0.0591/n)log(Q)
0 V 1
Ecell = - (0.0591)log(0.1) = V

37. Galvanic Cell Types of electrochemical systems
electrolytic - chemical reaction which occurs when electrical current is passed through solution
voltaic/galvanic - spontaneous reactions able to generate a supply of electricity (e.g., batteries) / which converts chemical energy to electrical energy. 37

38. Contd..
An electrochemical cell is a system consisting of electrodes that dip into an electrolyte in which a chemical reaction either uses or generates an electric current.
A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current.
An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise non spontaneous reaction. 38

39. Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
A complete redox reaction takes place in a galvanic cell
Overall reaction separated into half-reactions which take place at the anode and cathode
Given the following reaction:
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
the two half-reactions are:
oxidation half-reaction: Zn(s)  Zn2+(aq) + 2e- (anode)
reduction half-reaction: Cu2+(aq) + 2e-  Cu(s) (cathode)
Electrode reactions:
Anode: site of oxidation; electrons originate there; neg. pole of cell (anions migrate toward)
Cathode: site of reduction; electrons consumed there; pos. pole of cell (cations migrate toward) 39

40. Figure : Atomic view of voltaic cell
Figure : Atomic view of voltaic cell. In a voltaic cell, two half-cells are connected in such a way that electrons flow from one metal electrode to the other through an external circuit.
The electrons flow through the external circuit to the copper electrode where copper ions gain the electrons to become copper metal. 40

41. Figure Two electrodes are connected by an external circuit
Daniel Cell Design 41

42. The Electrochemical Cell

43. Zinc is (much) more easily oxidized than Copper
Maintain equilibrium electron densities
Add copper ions in solution to Half Cell II
Salt bridge only carries negative ions
This is the limiting factor for current flow
Pick a low-resistance bridge

44. The Electrochemical Series
Most wants to reduce (gain electrons)
Gold
Mercury
Silver
Copper
Lead
Nickel
Cadmium
But, there’s a reason it’s a sodium drop
Iron
Zinc
Aluminum
Magnesium
Sodium
Potassium
Lithium
Most wants to oxidize (lose electrons)

45. Reactions at electrodes
Left: Galvanic cell. Electrons are deposited on the anode (so it is neg) and collected from the cathode (so it is positive)
Right: Electro-lytic cell. Elec-trons are forced out of the anode (positive) and into the cathode (negative)

46. Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
Cell Notation
Cell notation is used to describe structure of galvanic cell
For the Zn/Cu cell, the galvanic cell notation is:
Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
= phase boundary
= salt bridge
anode reaction: to the left of the salt bridge
cathode reaction: to the right of the salt bridge
both half-cell reactions in order of spontaneous reaction
Zinc solid reacts to form zinc(II) ion at the anode
Copper(II) ion reacts to form copper metal at the cathode 46

47. Potentiometric titrations
Principle
It measures the change in potential , can be used for all kinds of titration :
1- acid base
2-redox
3-complexometry.

48. Eright= E0+0.0591/n log OX/RED
Potentiometry
Ecell=ERt-ELf
Eleft= E /n log OX/RED
Eright= E /n log OX/RED

49. When it is used
It is used when the endpoints are very difficult to determine , either when:
1- very diluted solution.
2-coloured and turbid solution
3-absence of a suitable indicator

50. Instrument Combined glass electrode ( double function electrode (
Potentiometer / PH meter
Magnetic stirrer

51. AIM: Titration of a Strong acid ( Hydrochloric acid ) against a strong base ( NaOH) REAGENTS REQUIRED: Standard NaOH solution Hydrochloric acid

52. Procedure:
10.0 ml of the given hydrochloric acid solution is pipetted out into a 100ml beaker and about 40 ml of distilled water is added so that the tips of combined glass electrode is completely immersed in solution. 
The potential / pH of the solution is noted before adding the alkali. 
Then standardized sodium hydroxide is added from a burette with 1ml increment and the readings (pH) are noted while shaking thoroughly the contents of the beaker during the addition. 
Near the equivalence point sodium hydroxide is added drop wise.  To get a neat curve titration is continued with a few more increments.   

53. Potentiometric titration curve of 0. 1 N NaOH against 0
Potentiometric titration curve of 0.1 N NaOH against 0.1N Hydrochloric acid acid

54. Calculation The normality of the hydrochloric acid from the formula
V1 N1 =  V2 N2. 
The concentration of hydrochloric acid (in gm/l) as follow
C = Eq. wt x normality of hydrochloric acid

55. Batteries

56. Battery
A device that stores energy for later
release as electricity.

57. Battery (Ancient) History
1800 Voltaic pile: silver zinc
1836 Daniell cell: copper zinc
1859 Planté: rechargeable lead-acid cell
1868 Leclanché: carbon zinc wet cell
1888 Gassner: carbon zinc dry cell
1898 Commercial flashlight, D cell
1899 Junger: nickel cadmium cell
Daniell cell was first reliable source of energy
Solved corrosion problems in Voltaic cell

58. Battery History 1946 Neumann: sealed NiCd
1960s Alkaline, rechargeable NiCd
1970s Lithium, sealed lead acid
1990 Nickel metal hydride (NiMH)
1991 Lithium ion
1992 Rechargeable alkaline
1999 Lithium ion polymer
v battery introduced
1991 Early lithium battery recall (mobile phone battery burst into flame)

59. Battery Nomenclature Duracell batteries 9v battery 6v dry cell
More precisely
Two cells
A real battery
Another battery
A battery contains two or more cells
All products here are dry cells (car batteries are wet cells)

60. Classification
Primary Batteries
Secondary Batteries

61. Non rechargeable and are meant for
Primary Batteries
Non rechargeable and are meant for
single use.
They are discarded after use

62. Primary (Disposable) Batteries
Zinc carbon (flashlights, toys)
Lithium (photoflash)
Heavy duty zinc chloride (radios, recorders)
Alkaline (all of the above)
Silver, mercury oxide (hearing aid, watches)

63. Standard Zinc Carbon Batteries
Chemistry
Zinc (-), manganese dioxide (+)
Zinc, ammonium chloride aqueous electrolyte
Features
Inexpensive, widely available
Inefficient at high current drain
Poor discharge curve (sloping)
Poor performance at low temperatures

64. Lithium Cells Liquid Cathode Lithium Cell
Solid Cathode Lithium Cells ( Li - MnO2)
Solid Electrolyte Lithium Cell

65. Lithium Manganese Dioxide
Chemistry
Lithium (-), manganese dioxide (+)
Alkali metal salt in organic solvent electrolyte
Features
High energy density
Long shelf life (20 years at 70°C)
Capable of high rate discharge
Expensive

66. IBM ThinkPad Backup Battery
Panasonic CR2032 coin-type lithium-magnesium dioxide primary battery
Application: CMOS memory backup
Constant discharge, ~0.1 mA
Weight: 3.1g
220 mA-h capacity

67. Heavy Duty Zinc Chloride Batteries
Chemistry
Zinc (-), manganese dioxide (+)
Zinc chloride aqueous electrolyte
Features (compared to zinc carbon)
Better resistance to leakage
Better at high current drain
Better performance at low temperature

68. Battery Nomenclature Duracell batteries 9v battery 6v dry cell
More precisely
Two cells
A real battery
Another battery
A battery contains two or more cells
All products here are dry cells (car batteries are wet cells)

69. Standard Alkaline Batteries
Chemistry
Zinc (-), manganese dioxide (+)
Potassium hydroxide aqueous electrolyte
Features
50-100% more energy than carbon zinc
Low self-discharge (10 year shelf life)
Good for low current (< 400mA), long-life use
Poor discharge curve

70. Alkaline-Manganese Batteries (2)

71. Secondary Batteries cycle use.
These are rechargeable and are meant for multi
cycle use.
After every use the electrochemical reaction could
be reversed by external application till the
capacity fades or lost due to leakage or internal
short circuit.

72. Secondary (Rechargeable) Batteries
Nickel cadmium
Lead acid
Nickel metal hydride
Alkaline
Lithium ion
Lithium ion polymer

73. Nickel Cadmium Batteries
Chemistry
Cadmium (-), nickel hydroxide (+)
Potassium hydroxide aqueous electrolyte
Features
Rugged, long life, economical
Good high discharge rate (for power tools)
Relatively low energy density
Toxic
Used in power tools
Use restricted in some countries

74. Ni Cd Recharging Over 1000 cycles (if properly maintained)
Fast, simple charge (even after long storage)
C/3 to 4C with temperature monitoring
Self discharge
10% in first day, then 10%/mo
Trickle charge (C/16) will maintain charge
Memory effect
Overcome by 60% discharges to 1.1V

75. NiCd Memory Effect

76. Lead Acid Batteries Chemistry Features Lead Sulfuric acid electrolyte
Least expensive
Durable
Low energy density
Toxic

77. Lead Acid Recharging Low self-discharge No memory
40% in one year (three months for NiCd)
No memory
Cannot be stored when discharged
Limited number of full discharges
Danger of overheating during charging

78. Lead Acid Batteries Ratings Deep discharge batteries
CCA: cold cranking amps (0F for 30 sec)
RC: reserve capacity (minutes at 10.5v, 25amp)
Deep discharge batteries
Used in golf carts, solar power systems
2-3x RC, CCA of car batteries
Several hundred cycles

79. Lithium Ion Batteries Chemistry Features
Graphite (-), cobalt or manganese (+)
Nonaqueous electrolyte
Features
40% more capacity than NiCd
Flat discharge (like NiCd)
Self-discharge 50% less than NiCd
Expensive

80. Lithium Ion Recharging
300 cycles
50% capacity at 500 cycles

81. Battery Capacity Type Capacity (mAh) Density (Wh/kg) Alkaline AA 2850
124
Rechargeable
1600 80
Ni Cd AA
750 41
Ni MH AA
1100 51
Lithium ion
1200
100
Lead acid
2000 30
Sources differ as to exact numbers
Density should take packaging into account, not just chemistry
Consult product specification sheets

82. Fuel Cells
It is an electro chemical cell which converts chemical energy into electrical energy

84. Working of a Fuel Cell
It operates similarly to a battery, but it does not run down nor does it require recharging
As long as fuel is supplied, a Fuel Cell will produce both energy and heat

85. A Fuel Cell consists of two catalyst coated electrodes surrounding an electrolyte
One electrode is an anode and the other is a cathode
The process begins when Hydrogen molecules enter the anode
The catalyst coating separates hydrogen’s negatively charged electrons from the positively charged protons

86. Working of a Fuel Cell
The electrolyte allows the protons to pass through to the cathode, but not the electrons
Instead the electrons are directed through an external circuit which creates electrical current

88. While the electrons pass through the external circuit, oxygen molecules pass through the cathode
There the oxygen and the protons combine with the electrons after they have passed through the external circuit
When the oxygen and the protons combine with the electrons it produces water and heat

89. Chemistry of a Fuel Cell
Anode side:
2H H e-
Cathode side:
O2 + 4H+ + 4e H2O
Net reaction:
2H2 + O H2O

91. Individual fuel cells can then be placed in a series to form a fuel cell stack
The stack can be used in a system to power a vehicle or to provide stationary power to a building

92. How can Fuel Cell technology be used?
Transportation
All major automakers are working to commercialize a fuel cell car
Automakers and experts speculate that a fuel cell vehicle will be commercialized by 2010
50 fuel cell buses are currently in use in North and South America, Europe, Asia and Australia
Trains, planes, boats, scooters, forklifts and even bicycles are utilizing fuel cell technology as well

93. How can Fuel Cell technology be used?
Stationary Power Stations
Over 2,500 fuel cell systems have been installed all over the world in hospitals, nursing homes, hotels, office buildings, schools and utility power plants
Most of these systems are either connected to the electric grid to provide supplemental power and backup assurance or as a grid-independent generator for locations that are inaccessible by power lines

94. THANK U