This document provides information on electrochemistry and electrochemical cells. It defines electrochemistry as the study of electricity production from spontaneous chemical reactions and use of electrical energy for non-spontaneous reactions. It describes different types of electrochemical cells including galvanic cells that convert chemical to electrical energy and electrolytic cells that do the opposite. Key concepts discussed include electrode potentials, standard hydrogen electrode, Nernst equation, and factors affecting cell potential. Common electrochemical devices like batteries and the corrosion process are also summarized.

1. SUB CHEMSITRY
CLASS 12
ELECTROCHEMSITRY

2. ELECTROCHEMISTRY
Electrochemistry is that branch of chemistry which deals with the study of
production of electricity from energy released during spontaneous chemical
reactions and the use of electrical energy to bring about non-spontaneous
chemical transformations.
Importance of Electrochemistry
1. Production of metals like Na, Mg. Ca and Al.
2. Electroplating.
3. Purification of metals.
4. Batteries and cells used in various instruments.

3. Conductors
Substances that allow electric current to pass through them are known as
conductors.
Metallic Conductors or Electronic Conductors
Substances which allow the electric current to pass through them by the
movement of electrons are called metallic conductors, e.g.. metals.
Electrolytic Conductors or Electrolytes
Substances which allow the passage of electricity through their fused
state or aqueous solution and undergo chemical decomposition are
called electrolytic conductors, e.g., aqueous solution of acids. bases
and salts.
Electrolytes are of two types:
1. Strong electrolytes
The electrolytes that completely dissociate or ionise into ions are called
strong electrolytes. e.g., HCl, NaOH, K2SO4
2. Weak electrolytes
The electrolytes that dissociate partially are called weak electrolytes,
e.g., CH3COOH, H2CO3, NH4OH,H2S, etc.

4. ELECTROCHEMICAL CELL
 The cell which converts chemical energy to
electrical energy and vice versa is called an
electrochemical cell.
 TYPES
 Galvanic Cells -Converts chemical energy into
electrical energy.
 Galvanic cell is also called voltaic cell.
 Electrolytic Cells -Converts electrical energy into
chemical energy.

6. General Representation of an Electrochemical Cell
CELL DIAGRAM OR
REPRESENTATION OF A CELL

7. EXAMPLE
The Daniel cell is represented as follows :
Zn(s) | Zn2+ (C1 ) || Cu2+ (C2 ) | Cu (s)
Salt bridge
When the oxidation takes place then it is anode when there is
reduction takes place then thereis cathode.

8. GALVANIC CELL
 DANIEL CELL
 It is a Galvanic Cell in which Zinc and Copper are used for
the redox reaction to take place.
Zn (s) + Cu2+ (aq) Zn2+ (aq) + Cu(s)
Oxidation Half : Zn (s) Zn2+ (aq) + 2e–
Reduction Half : Cu2+(aq) + 2e– Cu(s)
NOTES:-
 Anode is assigned negative polarity
 cathode is assigned positive polarity.
 In Daniel Cell, Zn acts as the anode and Cu acts as the
cathode
 Ques…. What is the difference between electrolytic cell
and galvanic cell.

9. GALVANIC CELL

10. FUNCTION OF SALT BRIDGE
1. It completes the circuit and allows the flow of
current.
2. It maintains the electrical neutrality on both sides.
3. Salt-bridge generally contains solution of strong
electrolyte such as KNO3, KCL etc. KCl is
preferred because the transport numbers of
K+ and Cl-are almost same.

11. SOME TERMS
 CELL POTENTIAL
It is the difference between electrode potentials (of
both electrodes i.e anode and cathode) of the given cell. It is denoted by E cell.
 Electrode Potential
When an electrode is in contact with the solution of its ions
in a half-cell, it has a tendency to lose or gain electrons
which is known as electrode potential. It is expressed in
Volts.
OR
Potential difference between the electrode and electrolyte.
 Oxidation potential
The tendency to lose electrons in the above case is known as oxidation
potential. Oxidation potential of a half-cell is inversely proportional to the
concentration of ions in the solution.
 Reduction potential
The tendency to gain electrons in the above case is known as reduction
potential.

12. E cell is written as :
E cell = E right – Eleft
EXAMPLE:-
Cell reaction:
Cu(s) + 2Ag+ (aq) → Cu2+(aq) + 2 Ag(s)
Half-cell reactions:
Cathode (reduction): 2Ag+ (aq) + 2e– → 2Ag(s)
Anode (oxidation): Cu(s) → Cu2+(aq) + 2e
The cell can be represented as:
Cu(s)|Cu2+(aq)||Ag+ (aq)|Ag(s)
Ecell = Eright – Eleft
=E Ag+/ Ag – E Cu2+ /Cu
NOTES:-It is not possible to determine the absolute value
of electrode potential. For this a reference electrode [NHE
or SHE] is required.

13. STANDARD HYDROGEN
ELCTRODE(SHE)
 It is related to standard conditions.
 The standard conditions taken are :
(i) 1M concentration of each ion in the solution.
(ii) A temperature of 298 K.
(iii) 1 bar pressure for each gas

14. EMF(ELECTROMOTIVE FORCE)
 It is the difference between the electrode potentials of
two half-cells and cause flow of current from
electrode at higher potential to electrode at lower
potential. It is also the measure of free energy
change. Standard emf of a cell. It is denoted by
E0(prnounced as e not).

16. ELECTROCHEMICAL SERIES
 The half cell potential values are standard values and are
represented as the standard reduction potential values as
shown in the table at the end which is also called
Electrochemical Series.

18. APPLICATIONS OF ELECTROCHEMICAL
SERIES
 The lower the value of E°, the greater the tendency
to form cation.
M → Mn+ + ne
 Reducing character increases down the series.
 Reactivity increases down the series.
 Determination of emf; emf is the difference of
reduction potentials of two half-cells.
•Eemf = ERHS – ELHS
 If the value of emf is positive. then reaction take
place spontaneously, otherwise not.

19. NERNST EQUATION
 It relates electrode potential with the concentration
of ions.
.

20. NERNST EQUATION
Concentration of pure solids and liquids is taken as unity.

21. APPLICATIONS OF NERNST EQUATION

•Relationship between free energy change and equilibrium constant
ΔG° = – 2.303RT log Kc

22. CONDITIONS FOR SPONTANEOUS
REACTION(IMPORTANT)

23.  CONDUCTANCE: It is the reciprocal of resistance.
 The ease with which the electric current flows
through a conductor.It is denoted by G.
G = (1/R), units ohm-1 mhos or Ω-1
 SPECIFIC CONDUCTIVITY -It is the reciprocal of
resistivity .It is denoted by kappa(k).

24. NOTE:-Specific conductivity decreases on dilution. This is
because concentration of ions per cc decreases upon dilution.
 Molar Conductivity (Λm)
The conductivity of all the ions produced when 1 mole of an
electrolyte is dissolved in V mL of solution is known as molar
conductivity.It is related to specific conductance as
Λm = (k x 1000/M) where. M = molarity.
It units are Ω-1 cm2 mol-1 or S cm2 mol-1.
 Equivalent conductivity (Λe)
The conducting power of all the ions produced when 1 g-
equivalent of an electrolyte is dissolved in V mL of solution, is
called equivalent conductivity. It is related to specific
conductance as
Λe = (k x 1000/N) where. N = normality.
Its units are ohm-1 cm2 (equiv-1) or mho cm2 (equiv-1) or S
cm2 (g-equiv-1).

25. MOLAR CONDUCTIVITY & EQUIVALENT
CONDUCTIVITY

26. FACTORS AFFECTING ELECTROLYTIC
CONDUCTANCE
 Electrolyte -is a substance that dissociates in
solution to produce ions and hence conducts
electricity in dissolved or molten state.
 Examples : HCl, NaOH, KCl (Strong electrolytes).
CH3COOH, NH4OH (Weak electrolytes).
 Nature of electrolyte or interionic attractions
 Temperature

28. VARIATION OF CONDUCTIVITY AND MOLAR
CONDUCTIVITY WITH DILUTION
 Conductivity always decreases with decrease in
concentration both, for weak and strong
electrolytes. This can be explained by the fact that
the number of ions per unit volume that carry the
current in a solution decreases on dilution.
 Molar conductivity increases with decrease in
concentration. This is because the total volume, V,
of solution containing one mole of electrolyte also
increases

30. LIMITING MOLAR CONDUCTIVITY (Λ0M )
 The value of molar conductivity when the
concentration approaches zero is known as limiting
molar conductivity or molar conductivity at infinite
dilution.
 It is possible to determine the molar conductivity at
infinite dilution λ0m.
 In case of strong electrolyte by extrapolation of
curve of λ0mvs c.
 In caseof weak electrolyte at infinite dilution cannot
be determined by extapolation of the curve as the
curve becomes almost parallel to y-axis when
concentration approaches to zero.

31. KOHLRAUSCH’S LAW
 It states that the limiting molar conductivity of an
electrolyte can be represented as the sum of the
individual contributions of the anion and cation of
the electrolyte.
 In general, if an electrolyte on dissociation gives v+
cations and v– anions then its limiting molar
conductivity is given by:-

32. APPLICATIONS OF LAW

33. ELECTROLYSIS
 It is the process of decomposition of an electrolyte when
electric current is passed through either its aqueous
solution or molten state.
 PREDICT THE PRODUCT OF ELECTROLYSIS:-
 when an aqueous solution of an electrolyte is
electrolysed, if the cation has higher reduction potential
than water (-0.83 V), cation is liberated at the cathode
(e.g.. in the electrolysis of copper and silver salts)
otherwise H2 gas is liberated due to reduction of water
(e.g., in the electrolysis of K, Na, Ca salts, etc.)
 Similarly if anion has higher oxidation potential than
water (- 1.23 V), anion is liberated (e.g., Br-), otherwise
O2 gas is liberated due to oxidation of water (e.g., in
caseof F-, aqueous solution of Na2SO4 as oxidation
potential of SO2-
4 is – 0.2 V).

37. BATTERIES
 Cells and batteries are available in a wide variety of
types.
 TYPES
 A primary cell cannot be recharged and cannot be
used.
 A secondary cell, or storage cell, can be recharged
and can be use.

38. PRIMARY CELL
 DRY CELL:-
Anode : Zn container
Cathode : Carbon (graphite) rod surrounded by powdered MnO2
and carbon.
Electrolyte : NH4 Cl and ZnCl2
Reaction :
Anode : Zn → Zn2+ + 2e–
Cathode : MnO2 + NH4
+ + e – → MnO (OH) + NH3
The standard potential of this cell is 1.5 V and it falls as the
cell gets discharged continuously and once used it cannot be
recharged.

40.  MERCURY CELL:
 It offers a rather more stable voltage.
 The emf of the Mercury Cell is 1.35 V.
 Usually, the mercury cell is costlier. This is the reason, why they
are used only in sophisticated instruments such as camera,
hearing aids, and watches etc.

41.  The reactions during discharge are,
At anode: Zn(Hg) + 2OH– →Zn (OH)2 + 2e–
At cathode: HgO + H2O + 2e– →Hg + 2OH–
 Overall reaction:
Zn(Hg) + HgO(s) →Zn(OH)2+ Hg(l)

42. SECONDARY CELL
 Secondary cells are those which can be recharged
again and again for multiple uses. e.g. lead storage
battery and Ni – Cd battery.

43. LEAD STORAGE
 Anode : Lead (Pb)
 Cathode : Grid of lead packed with lead oxide (PbO2 )
 Electrolyte : 38% solution of H2 SO4
Discharging Reactions
 Anode:
Pb(s) + SO4
2–(aq) PbSO4 (s) + 2e–
 Cathode:
PbO2 (s) + 4H+ (aq) + SO4 2–(aq) + 2e– PbSO4 (s) +
2H2O(l)
 Overall Reaction :
Pb(s) + PbO2 (s) + 2H2 SO4 (aq) 2PbSO4 (s) + 2H2O(l)

44. CORROSION
 It involves a redox reaction and formation of an
electrochemical cell on the surface of iron or any other metal.
Anode : 2Fe (s) → 2 Fe2+ + 4e– Eº = + 0.44 V
Cathode : O2 (g) + 4H+ + 4e– → 2H2 O(l) Eº = 1.23 V
Overall reaction:
2Fe (s) + O2 (q) + 4H+ → 2Fe2+ + 2H2O

46. IMPORTANT QUESTIONS
a) Following reactions occur at cathode during the electrolysis of aqueous
silver chloride solution :
Ag+(aq) + e– → Ag(s) E° = +0.80 V
H+(aq) + e– →1/2 H2 (g) E° = 0.00 V
On the basis of their standard reduction electrode potential (E°) values,
which reaction is feasible at the cathode and why ?
(b) Define limiting molar conductivity. Why conductivity of an electrolyte solution
decreases with the decrease in concentration ?
(c) Calculate emf of the following cell at 25 °C :
Fe | Fe2+(0.001 M) || H+(0.01 M) | H2(g) (1 bar) | Pt(s)
E°(Fe2+ | Fe) = –0.44 V E°(H+ | H2) = 0.00 V
(d) From the given cells : Lead storage cell, Mercury cell, Fuel cell and Dry cell .
Answer the following :
(i) Which cell is used in hearing aids ?
(ii) Which cell was used in Apollo Space Programme ?
(iii) Which cell is used in automobiles and inverters ?
(iv) Which cell does not have long life ?

47. (e)Calculate the degree of dissociation (α) of acetic acid if its molar
conductivity (∧m) is 39.05 S cm2mol–1. Given λo (H+) = 349.6 S cm2
mol–1 and λo(CH3COO–) = 40.9 S cm2mol–1.
(f) Calculate the mass of Ag deposited at cathode when a current of 2
amperes was passed through a solution of AgNO3 for 15 minutes. (Given
: Molar mass of Ag = 108 g mol–1 1F = 96500 C mol–1)
(g) Define fuel cell.
(h) Write the cell reaction and calculate the e.m.f. of the following cell at 298
K :
Sn (s) | Sn2+ (0·004 M) || H+(0·020 M) | H2(g) (1 bar) | Pt (s) (Given :E0
Sn2+/Sn = – 0·14 V)
(i) Give reasons :
 On the basis of Eo values, O2 gas should be liberated at anode but it is
Cl2 gas which is liberated in the electrolysis of aqueous NaCl.
 Conductivity of CH3COOH decreases on dilution.
(j) For the reaction 2AgCl (s) + H2 (g) (1 atm) 2Ag (s) + 2H+ (0·1 M)
+ 2Cl–(0·1 M),
∆Go= – 43600 J at 25 C. Calculate the e.m.f. of the cell. [log 10–n= – n]
(k) Define fuel cell and write its two advantages.

48. (l) E°cell for the given redox reaction is 2.71 V
Mg+ Cu2+(0.01v) Mg2+(0.001V) + Cu (s)
Calculate Ecell for the reaction. Write the direction of flow of
current when an external opposite potential applied is i) less
than 2.71 V and (ii) greater than 2.71 V
(m) A steady current of 2 amperes was passed through two
electrolytic cells X and Y connected in series containing
electrolytes FeSO4 and ZnSO4 until 2.8 g of Fe deposited at
the cathode of cell X. How long did the current flow? Calculate
the mass of Zn deposited at the cathode of cell Y. (Molar mass
: Fe = 56 g mol–1 Zn = 65.3 g mol–1, 1F = 96500 C mol–1)

49. THANK U