The document discusses different types of electrochemical cells including primary cells that produce electricity from non-reversible chemical reactions and secondary cells that can be recharged by passing electricity in the opposite direction of the spontaneous reaction. Examples of primary cells discussed include Daniel, mercury, dry, and alkaline cells, while examples of secondary cells include lead-acid, nickel-cadmium, nickel-metal hydride, and lithium-ion batteries. The working and reactions of common cells like lead-acid, alkaline, and dry cells are also explained.

2. Dr. SURENDRAN PARAMBADATH
(M.Sc, M.Phil, M.Tech)
Formerly: Post Doctoral Research Associate,
Nano-Information Materials Research Laboratory,
Pusan National University, Busan-South Korea
Currently: Assistant Professor
Govt. Polytechnic College, Perinthalmanna

3. Electrochemistry is the study of
Inter-convention of
electrical energy
and
chemical energy
Electrical Energy Chemical Energy

4. Electrolytic Cell
This device can convert electrical
energy in to chemical energy
Electrochemical Cell
This device can convert chemical
energy in to electrical energy

5. Conductors and Insulators
This classification is based on their ability to allow electric current
to pass through them.
Conductors are those substances which allow
electric current to pass through them.
Examples: Metals, Alloys, Graphite, Ionic compounds in fused
or dissolved state.
Insulators are those substances which do not
allow electric current to pass through them.
Examples: Glass, Wood, Paper, Organic Compounds etc.

6. Types of Conductors
1. Metallic Conductors
Eg: Metals and Alloys
2. Electrolytic Conductors
Eg: NaCl, CuSO4, etc
3. Semi Conductors
Eg: Ge doped with Ar or P.
4. Super Conductors
Eg: Mercury at 4K

7. Metallic Conductors Electrolytic Conductors
1. Due to movement of electrons 1. Due to movement of ions
2. No Chemical Change takes place 2. Electrolysis takes place
3. No transfer of matters 3. Transfer of matter in the forms of
ions
4. Conductance of metals 4. Conductance of metals increases
decreases with increase of with increase of temperature
temperature

8. Electrolytes are substances which conduct
electric current through them either in the
molten state or in the dissolved state.
Eg: NaCl, H2SO4, KOH, HNO3 etc
Non electrolytes are substances which do not
conduct electricity in the fused state or in
dissolved state.
Eg: Sugar. Urea, alcohol etc.

9. Strong electrolytes: Electrolytes that dissociate almost
completely into ions even at moderate concentration
are called strong electrolytes.
They have high conductivity.
Eg: HCl, HNO3, H2SO4 etc. NaOH, KOH
etc, NaCl, CuSO4 etc.
Weak electrolytes: Electrolytes which dissociate into
ions partially at moderate concentrations.
They have low conductivity.
Eg: Acetic acid, Oxalic acid, NH4OH etc.

10. Electrolysis is the process of decomposition of an electrolyte
by passage of electric current.
- battery
+
Cl2 (g) escapes
Na (l)
NaCl (l)
Na+
Cl- Na+
Cl-
(-) (+)
electrode electrode
half-cell Cl- Na+ half-cell
Na+ + e-  Na 2Cl-  Cl2 + 2e-

11. Molten NaCl
Observe the reactions at the electrodes
- +
battery
Cl2 (g) escapes
Na (l)
Na+ NaCl (l)
Cl- Na+
Cl
(-) - (+)
electrode electrode
half-cell Cl Na+ half-cell
-
Na+ + e-  Na All rights reserved. 2Cl-  Cl2 + 2e-
http://academic.pgcc.edu/~ssinex/E_cells.ppt. 11

12. 1st Law………….
The mass of substance discharged at an electrode during
electrolysis is directly proportional to the quantity of electricity
passed through the electrolyte.
m Q ………………………m = Zit
m = mass in grams of substance discharged
Q= quantity of electricity in coulombs
t= time in seconds
Z= Electrochemical equivalent
ECE may be defined as the mass of the substance discharged by passing one
coulomb of electricity.

13. 2nd Law…….
The law states that when the same quantity of electricity is
passed through different electrolytes connected in series, the
amount of substance discharged at the electrodes are directly
proportional to their chemical equivalent.
E = Equivalent weight m1/m2 = E1/E2
e-
- battery
+
- + - + - +
e- e- e-

15. The process of depositing a superior metal on am inferior metal
by passing electric current is called electroplating.
The base metal object, which is to be plated is
made the cathode in the electrolytic cell.
The rod of pure metal to be deposited on the
object is made the anode.
The electrolyte is a solution of a soluble salt of the
superior metal.

16. Examples for Superior Metals
Cr, Ni, Ag & Au
Examples for Inferior Metals
Fe & Cu

17. 1. To protect the inferior metal object from corrosion.
2. To increase the resistance to chemical attack
3. To improve its physical appearance so as to make it
more attractive.
4. To modify hardness
5. To repair damaged part of the machinery.
6. To strengthen light weight non metallic like
wood, glass, leather, cloth etc.
7. To obtain conducting surfaces, eg copper plating on
wooden or plastic radar antenna masts.

19. Step: 1 Wash the object with an organic solvent
to remove any grease or oil on it. Then wash
with dilute sulphuric acid to remove oxide film
from the surface.
Finally wash with chromic acid or detergent to
clean the surface thoroughly.

20. Step: 2 The metal surface should be rough so
that the deposit sticks firmly and permanently.

21. Step 3: The electrolyte is nickel sulphate
solution containing nickel chloride or nickel
ammonium sulphate solution.

22. Step 4: pH of the electrolytic solution is
maintained between 4 to 5.

23. Step 5: The cleaned object to be plated is made
the cathode of the electrolytic cell, and pure
nickel plate or block, the anode.

24. Reaction:
At Anode,
Ni Ni2+ + 2e-
At Cathode,
Ni2+ + 2e- Ni

25. The Electrolytes used for electroplating should be
1. Highly soluble
2. Stable towards oxidation, reduction or hydrolysis
3. A good conductor

26. Anodizing is a process of coating a base metal like Aluminium, or
Magnesium with a thin uniform and protective oxide film.
Anode: Base metal
Electrolyte: Chromic acid, dil. H2SO4 or Phosphoric acid
Cathode: Graphite rod or lead sheet
The anode coating being thicker than the natural oxide
film, it has greater resistance to corrosion and
mechanical injury.
By addition of suitable dyes and pigment to the
electrolyte, brightly colored, lustrous surface
coating are obtained.

27.  “Cells” are containers of liquid with electrodes:
Source or use of
electricity
Cell Electrode
– + Molten or
– +
– + aqueous
chemicals
• In “electrolytic cells”, electricity is used to force chemicals
to undergo a redox reaction
• In “galavanic cells”, electricity is produced spontaneously
from a redox reaction

29.  An apparatus that allows a redox reaction to occur by
transferring electrons through an external connector.
 Product favored reaction
> voltaic or galvanic cell
--> electric current
Batteries are voltaic cells
 Reactant favored reaction
> electrolytic cell ---> electric current used to
cause chemical change.

30. The device in which chemical
energy is converted into
electrical energy is called
galvanic cell.
Working w.r.t: Oxidation-
Reduction Reaction.

31. 1. Two half cells, namely zinc half cell and copper half cell.
In the former is a zinc rod dipped in a ZnSO4 solution and the
latter is a copper rod dipped in a CuSO4 solution. The two
metallic rods are called electrodes.
2. The two half cells are connected externally by a metallic
wire to a galvanometer through a key and internally by a
salt bridge.
3. The salt bridge is an inverted U-tube containing an arouse
solution of an inert salt like KCl, KNO3, NH4NO3 to which
some agar-agar or gelatin has been added to convert it into
a semi solid, ie gel. The ends of the U-tube are plugged with
glass wool.

32. 1.Permits the passage of electric current internally,
2.Maintains the electrical neutrality of the solution,
3.Prevents intermixing of the solutions,
4.It does not take part in cell reaction.
Zn(s) / Zn2+(aq) // Cu2+(aq)/Cu(s)

34. Anode
Cathode
Zinc plate is eaten away and Copper deposits on the
copper plate. Electrons produced at the zinc anode flow
through the outer circuit to the copper cathode.
Electric Current is assumed to flow from copper to zinc,
ie, from positive terminal to negative terminal.

35. Cu(s) / Cu2+(aq) // Ag+(aq)/Ag(s)
Mg(s) / Mg2+(aq) // Ni2+(aq)/Ni(s)
Fe(s) / Fe2+(aq) // Au3+(aq)/Au(s)
Al(s) / Al3+(aq) // Sn2+(aq)/Sn(s)

36. e-
- battery +
Na+
Cl- Na+
Cl e-
(-) - (+)
(+) Cu Zn (-)
Cl- Na+
Cathode Anode Cathode Anode

37. galvanic electrolytic
produces need
electrical power
current two source
electrodes
anode (-) conductive anode (+)
cathode (+) medium cathode (-)
salt bridge vessel No salt bridge

38. Electrolytic Cell Galvanic Cell
1 Electrical Energy is converted Chemical Energy is converted into
into chemical energy electrical energy
2 Electrical energy brings about a Electrical energy is generated by a redox
redox reaction reaction
3 Anode is positive while cathode Anode is negative while cathode is
is negative positive
4 Redox reaction takes place in the Oxidation and reduction reactions are
same container carried out separately
5 No salt bridge is required Salt bridge is generally required
6 Ions are discharged at both the Ions are discharged at the cathode while
electrodes anode is consumed.

40. One of the main uses of
electrochemical cells is the generation
of portable electrical energy.
Two or more cells are connected in
series to form a battery which acts as
a source of electrical energy.

41. A com er ci al C l m
m el ust f ul f i l l t he
f ol l ow ng r equi r em s.
i ent
I t s houl d be
c o mp a c t a nd l i ght
a n d e a s y t o
t r a ns por t .
I t s v ol t a ge mu s t
n o t v a r y mu c h
d u r i ng us e .

42. There are two category of
energy producing galvanic
cells

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

44. 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
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

45. 1.Primary Cells (Disposable)
In primary cells the redox reaction occurs only once and the cell
becomes dead since the chemical reactions in these are not
reversible.
Daniel cell, mercury cell, Dry cell etc….
Daniel Cell in the commercial
form consists of a zinc electrode
dipping in zinc sulphate solution
contained in a porous pot. The pot
is placed ia a cylindrical copper
vessel containing copper sulphate
solution.

46. Zinc Rod
(Anode)
-
.... ....
+
.... .... Copper Vessel (Cathode)
.... ....
.... ....
.... .... ZnSO4 Solution
.... ....
.... ....
.... ....
.... ....
.... .... Porous Pot
.... ....
.... ....
.... ....
.... .... CuSO4 Solution
.... ....
.... …. ....
CuSO4 Crystals
Daniel Cell (Commercial Form)

47. The porous pot allows the passage of only
ions from one solution to another and serves
the purpose of salt bridge in the conventional
galvanic cell. When connections are made as
shown electrons flow from zinc to copper
and current is assumed to flow from copper
to zinc.
Zn(s) + CuSO4 (aq) ZnSO4 (aq) + Cu(s)
Zn(s) / Zn2+(aq) // Cu2+(aq)/Cu(s)
The e.m.f of the cell is 1.1 V

48. Dry Cell Battery NS
Anode (-)
Zn ---> Zn2+ + 2e-
Cathode (+)
2 NH4+ + 2e-  2 NH3 + H2
Give the net E0 of the complete
recation

49. Alkaline Battery NS
Nearly same reactions as
in common dry cell, but
under basic conditions.
Anode (-): Zn + 2 OH-  ZnO + H2O + 2e-
Cathode (+): 2 MnO2 + H2O + 2e-  Mn2O3 + 2 OH-

50. Mercury Battery NS
Anode:
Zn is reducing agent under basic conditions
Cathode:
HgO + H2O + 2e- ---> Hg + 2 OH-

51. 2. Secondary cells (Rechargeable)
 Nickel cadmium
 Nickel metal hydride
 Alkaline
 Lithium ion
 Lithium ion polymer
 Lead acid

52. In a secondary cell, the chemical
reactions taking place are reversible
and can be reversed by passing
electricity.
Since these cells can be
recharged, they can be used again and
again.
A battery consists of two or more

53. Lead Storage Battery
Anode (-) Pb + HSO4- ---> PbSO4 + H+ + 2e- Eo = +0.36 V
Cathode (+) PbO2 + HSO4- + 3 H+ + 2e- ---> PbSO4 + 2 H2O E0 = +1.68 V

54. Lead Storage Battery
The lead acid accumulator used in automobiles consists of 3 to 6 cells
to get a voltage of 6 to 12.
The cell has anode made of spongy lead presses in to grids and
cathode made of lead dioxide, PbO2 presses into grid made of lead.
A number of lead plates are connected in parallel and a number of lead
dioxide plates are also connected in parallel.
The plates are arranged alternately, separated by thin perforated
plastic or fibre glass sheets.
The whole arrangement is suspended in the electrolyte which is
dilutee sulphuric acid of density 1.31 gml-1, taken in a plastic or hard
rubber vessel.

55. Working of Lead Storage Cell
When discharging a lead storage cell,
At Anode
Lead loses electrons which flow through the wire to the cathode.
Pb Pb2+ + 2e-
The lead ions combine with the sulphate ions from sulphuric acid forming a
precipitate of lead sulphate.
Pb2+ + SO42- PbSO4
At Cathode
The electrons flowing from anode react with PbO2 of cathode and PbO2 is
reduced to Pb2+ in presence of H+ ions from H2SO4.
PbO2 + 4H+ + 2e- Pb2+ + 2 H2 O
The lead ions formed at the cathode react with sulphate ions forming a
precipitate of lead sulphate.
Pb2+ + SO42- PbSO4

56. Overall Reaction
Pb + PbO2 + 4H+ + 2SO42- 2PbSO4 + 2H2O + Energy
Total EMF = 2V

57. When you charge a battery, you are
forcing the electrons backwards (from
the + to the -). To do this, you will need
a higher voltage backwards than
forwards. This is why the ammeter in
your car often goes slightly higher while
your battery is charging, and then
returns to normal.
In your car, the battery charger is
called an alternator. If you have a
dead battery, it could be the
battery needs to be replaced OR
the alternator is not charging the
battery properly.

59. Charging of Lead Storage Cell
During discharging, both the electrodes get covered with PbSO4 and the
dilute sulphuric acid is consumed and its density falls from 1.31 to 1.2 g/ml.
When recharging an external e.m.f greater than 2 volts is passed from a generator to
recharge the cell. The positive pole of the generator is connected to positive pole of
the storage cell.
At Anode (+ve terminal)
PbSO4 + 2H2O PbO2 + 4H+ + SO42- + 2e-
At Cathode (-ve terminal)
PbSO4 + 2e- Pb + SO42-
Overall reaction
2PbSO4 + 2H2O + Energy Pb + PbO2 + 4H+ + 2SO42-

60. Ni-Cad Battery
Anode (-) : Cd + 2 OH- ---> Cd(OH)2 + 2e-
Cathode (+) : NiO2 + 2H2O + 2e- ---> Ni(OH)2 + 2OH-
Cell representation: Cd CdO KOH NiO2 Ni

62. 1.The metal atoms convert to metal ions.
M Mn+ + ne- (Oxidation)
-
-
-
-
-
+ - +
+ - +
+ - +
+ +

63. 1.The metal ions in solution gains electrons
from metal leaving a positive charge on the
metal.
Mn+ + ne- M (Reduction)
+
+
+
- + -
- + -
- + -
- + -
- -

64. Whatever may be the process, an electrical double layer
generates in between the metal and the solution. This electrical
double layer generates a potential difference.
The potential difference set up between the metal and its ions in
the solution is called electrode potential.
It is a measure of the tendency of an electrode to lose or gain
electrons when it is in contact with its own ions in solution.
i) If oxidation takes place at the electrode, the
potential is called oxidation potential.
ii) If reduction takes place at the electrode, the
potential is called reduction potential.

65. Electromotive Force (EMF)
When two half cells are connected, due to the
difference in potential an electric current flows
from the electrode of higher potential to the
electrode of lower potential.
The difference in potentials of two half cells of a
cell is known as electromotive force or emf of
the cell or cell potential.
EMF = Ecathode - Eanode

66. Electrochemical Series
It is an arrangement of
elements in the increasing
order of their standard
reduction potential.

67. Metal SRP, Eo
Lithium----------------- -3.05 V
Potassium
Calcium
Sodium
Decreasing Magnesium Increasing
tendency Aluminum order of std
to loose Zinc reduction
electrons Nickel potential
Tin
Hydrogen--------------- 0.00
Copper
Silver
Platinum
Gold---------------------- +1.15 V

68. Characteristics of ECS
1. Metals lying above hydrogen are easily rusted.
2. Iron and metals above it decomposes steam,
liberating hydrogen gas.
3. Oxides of iron and metals below it are decomposed
easily.
4. Oxides of mercury and metals below it are
decomposed on heating.

69. Applications of
Electrochemical Series

70. 1. It gives an idea regarding the
tendency of elements to lose or gain
electrons.
Elements with lower reduction potential have a tendency to
lose electrons, that is greater tendency to get oxidized.
So they are good reducing agents.
Elements with higher reduction potential have a tendency to
receive electrons, that is greater tendency to get reduced.
So they are good oxidizing agents.

71. 2. Displacement Reaction
An element above in the series can
displace an element below it.
In otherwords, an element with lower
reduction potential can displace an element
with higher reduction potential.
Eg: Zinc has lower reduction potential than
Copper. Hence zinc displaces copper from
CuSO4 solution.

72. 3. When a cell is constructed, anode
should be a metal higher in the series
and cathode a metal lower in the
series.
Eg: When a cell is constructed using zinc and
copper, Zn which is higher in the series will be the
anode and copper will be the cathode.
+ -
Zn Cu
Anode Cathode

73. 4. A metal above hydrogen in the
series can displace H2 gas from dilute
acid. But a metal below hydrogen
cannot liberate H2 gas from acid.
2Na + H2SO4 Na2SO4 + H2
2K + H2SO4 K2SO4 + H2
Ca + H2SO4 CaSO4 + H2
Mg + H2SO4 MgSO4 + H2

75. Fuel cells are galvanic cells in which
chemical energy from combustion of
fuel such as H2, CO, CH4 (gases)
alcohols (liquids) can be converted
into electrical energy.
About 75% of the chemical energy
can be converted into electrical
energy.

77. H2/O2 as a Fuel
Cars can use electricity generated by H2/O2 fuel
cells.
H2 carried in tanks or generated from
hydrocarbons (fuel)

79. The cell consists of two electrodes made
of porous graphite impregnated with a
catalyst Pt, Ag or CuO.
They are placed in aqueous concentrated
(35%) solution of NaOH or KOH.
H2 gas and O2 gas are continuously
bubbled through the porous electrodes at
the anode and cathode respectively at a
pressure of 50 atm.

80. The reaction at the electrodes are,
At anode
2H2 + 4OH- 4H2O + 4e-
At cathode
O2 + 2H2O + 4e- 4OH-
Overall reaction
2H2 + O2 2H2O + energy
The cell will produce an emf of about 1 volt. It is used in military
equipments, manned space crafts and submarines.

82. Fuel: 1 to 2 molar methanol in water.
To keep the concentration of methanol constant a
mixture of effluent and fresh methanol is recycled.

83. Anode: Oxidation of methanol
CH3-OH + H2O CO2 + 6H+ + 6e-
The protons H+ move from anode to the cathode via the electrolyte.
Cathode: Reduction of oxygen
3/2 O2 + 6H+ + 6e- 3H2O
The overall reaction
CH3-OH + 3/2 O2 CO2 + 2H2O
This type of cells are used as energy source in
(i) Space vehicles (ii) Submarines (iii) military vehicles
(iv) automobiles.

88. Advantages of Fuel cells
1.It converts energy of the fuel
directly to electrical energy.
2.Do not cause pollution problems.
3.Fuel cells are light and compact.
4.Efficiency is very high (60-75%)
5.Energy supply is continuous and
without any drop.