PHYSICAL CHEMISTRY

ELECTROCHEMISTRY

" A SPECIALLY DESIGNED KIT FOR LEARNING."

CONTENTS
THE KEY  Basic principles of subjects. An outline of the topics to be
discussed in class lectures.
THE ATLAS  Basic layout of subject. A route map correlating different subtopics
in coherent manner.
EASY RIDE  Introductory problems to get first hand experience of problem
solving.
PROFICIENCY TEST  To check you newly acquired concepts.
THE MIDDLE GAME  A collection of good problems.
ZENITH  Enjoy the ultimate heights of problem solving.
SUHANA SAFAR  A collection of previous ten years JEE problems.
 KEY CONCEPTS

ELECTROCHEMICAL CELLS
An electrochemical cell consists of two electrodes (metallic conductors) in contact with an electrolyte
(an ionic conductor).
An electrode and its electrolyte comprise an Electrode Compartment.

Electrochemical Cells can be classified as:

(i) Electrolytic Cells in which a nonspontaneous reaction is driven by an external source of current.

(ii) Galvanic Cells which produce electricity as a result of a spontaneous cell reaction

Note: In a galvanic cell, cathode is positive with respect to anode.

In a electrolytic cell, anode is made positive with respect to cathode.

ELECTROLYSIS
The decomposition of electrolyte solution by passage of electric current, resulting into deposition of
metals or liberation of gases at electrodes is known as electrolysis.

ELECTROLYTIC CELL
This cell converts electrical energy into chemical energy.
The entire assembly except that of the external battery is
known as the electrolytic cell

ELECTRODES
The metal strip at which positive current enters is called anode; anode is positively charged in electrolytic
cell. On the other hand, the electrode at which current leaves is called cathode. Cathodes are negatively
charged.

Anode Positive Loss of electron Positive

or oxidation current
takes place enters

Cathode Negative Gain of electron Current

or reduction leaves
takes place

ELECTROLYSIS OF MOLTEN SODIUM CHLORIDE

NaCl(molten)  Na+ + Cl–
Reactions at anode (oxidation) : cathode (reduction)
–
2Cl  Cl2(g) + 2e : – 2Na+ + 2e–  2Na(l)
There are two types of electrodes used in the electroytic cell, namely attackable and non - attackable.
The attackable elecrodes participitate in the electrode reaction. They are made up of reactive metals like
Zn, Cu, Ag etc. In such electrodes, atom of the metal gets oxidised into the corresponding cation, which
is passed into the solution. Thus, such anodes get dissolved and their mass decreases. On the other
hand, non-attackable electrodes do not participiate in the electrode reaction as they made up of unreactive
elements like Pt, graphite etc. Such electrodes do not dissolve and their mass remain same.
FARADAY’S LAWS OF ELECTROLYSIS:
(i) First law of electrolysis :
Amount of substance deposited or liberated at an electrode is directly proportional to amount of charge
passed (utilized) through the solution.
wQ
W = weight liberated, Q = charge in coulomb
w = ZQ
Z = electrochemical equivalent
when Q = 1 coulomb, then w = Z
Thus, weight deposited by 1 coulomb charge is called electrochemical equivalent.
Let 1 ampere current is passed till ‘t’ seconds .
Then, Q = It  w = ZIt
1 Faraday = 96500 coulomb = Charge of one mole electrons
One faraday is the charge required to liberate or deposit one gm equivalent of a substance at
corresponding electrode.
Let ‘E’ is equivalent weight then ‘E’ gm will be liberated by 96500 coulomb.
E E
 1 Coulomb will liberate gm ; By definition, Z
96500 96500
ItE
 W
96500
When a gas is evolved at an electrode, then above formula changes as,
ItV e
V
96500
where V = volume of liberated gas, Ve = equivalent volume of gas.
Equivalent volume may be defined as:
The volume of gas liberated by 96500 coulomb at STP.
(ii) Second law of electrolysis :
When same amount of charge is passed through different electrolyte solutions connected in series then
weight of substances deposited or dissolved at anode or cathode are in ratio of their equivalent
weights. i.e. w1/w2 E1/E2

QUALITATIVE ASPECTS OF ELECTROLYSIS

In the electrolysis process we have discussed above, we have taken molten salt as electrolyte, which
contains only one cation and anion. Now, if the electrolyte taken contains more than one cation and
anion (for example, aqueous solution of the ionic electrolyte), then the cation and anion that will get
discharged depends on the ability of cation to get reduced and the ability of anion to get oxidised.
The ability of an ion to get oxidised or reduced depends upon the size, mass, positive charge, negative
charge etc. Thus, it is not possible to predict qualitatively that which ion would be discharged first, as one
factor might enhance the ability to discharge while the other factor may hamper it. This can only be
predicted on the basis of quantitative value assigned based on the cumulative effect of all the factors reponsible
for an ion's ability to discharge. The value is referred as standard potential, which is determined by keeping
the concentration of ion as 1 M, pressure of gas at 1 atm, and the measurement done at 25°C. For a cation,
the standard reduction potential (SRP) values are compared. The cation having higher standard reduction
potential value is discharged in preference to cation with lower SRP value provided the ions are at 1 M
concentration. For an anion, the standard oxidation potential (SOP) values are compared and anion having
higher SOP is preferentially discharged, if the concentration is 1 M for each of the ion. The SRP values at
25°C for some of the reduction half reactions are given in the table below.
S.NO. Reduction half cell reaction E° in volts at 25°C
1. F2 + 2e– 2F– + 2.65
2. S 2 O82  + 2e– 2SO 24 + 2.01
3. Co 3+ + e– Co 2+ + 1.82
4. PbO2 + 4H+ + SO 24 + 2e– PbSO4 + 2H2O + 1.65
 + – 2+
5. MnO 4 + 8H + 5e Mn + 4H2O + 1.52
3+ –
6. Au + 3e Au + 1.50
–
7. Cl2 + 2e 2Cl– + 1.36
8. Cr2 O 72  + 14H+ + 6e– 2Cr3+ + 7H2O + 1.33
9. O2 + 4H+ + 4e– 2H2O + 1.229
– –
10. Br2 + 2e 2Br + 1.07
 +
11. NO3 + 4H +3e NO + 2H2O + 0.96
12. 2Hg2+ + 2e– Hg22 + 0.92
13. Cu2+ + I– + e– CuI + 0.86
+ –
14. Ag + e Ag + 0.799
2 –
15. Hg 2 + 2e 2 Hg + 0.79
3+ – 2+
16. Fe + e Fe + 0.77
–
17. I2 + 2e 2I– + 0.535
+ –
18. Cu + e Cu + 0.53
2+ –
19. Cu + 2e Cu + 0.34
–
20. Hg2Cl2 + 2e 2Hg + 2Cl– + 0.27
–
21. AgCl + e Ag + Cl– + 0.222
22. Cu2+ + e– Cu+ + 0.15
4+ –
23. Sn + 2e Sn2+ + 0.13
24. 2H+ + 2e– H2 0.00
3+ –
25. Fe + 3e Fe – 0.036
26. Pb2+ + 2e– Pb – 0.126
2+ –
27. Sn + 2e Sn – 0.14
28. AgI + e– Ag + I– – 0.151
29. Ni2+ + 2e– Ni – 0.25
2+ –
30. Co + 2e Co – 0.28
31. Cd2+ + 2e– Cd – 0.403
32. Cr3+ + e– Cr2+ – 0.41
2+ –
33. Fe + 2e Fe – 0.44
3+ –
34. Cr + 3e Cr – 0.74
2+ –
35. Zn + 2e Zn – 0.762
36. 2H2O + 2e– H2 + 2OH– – 0.828
37. Mn2+ + 2e– Mn – 1.18
38. Al3+ + 3e– Al – 1.66
39. H2 + 2e– 2H– – 2.25
40. Mg2+ + 2e– Mg – 2.37
41. Na+ + e– Na – 2.71
42. Ca2+ + e– Ca – 2.87
43. Ba2+ + 2e– Ba – 2.90
+ –
44. Cs + e Cs – 2.92
+ –
45. K +e  – 2.93
46. Li+ + e– Li – 3.03
 When solution of an electroyte contains more than one type of cations and anions at concentrations
different than 1 M, the discharge of an ion does not depend solely on standard potentials but also
depends on the concentration of ion in the solution. This value is refered as potential, called as
reduction potential for cation and oxidation potential for anion. The relation between reduction
potential and standard reduction potential is given by Nernst equation, as
RT [concentration of product ]
ERP = E°RP – ln
nF [concentration of reac tan t ]
where ERP = Reduction potential of cation and E°RP = Standard reduction potential of cation.
Thus, it is possible that a cation (A+) with lower standard reduction potential getting discharged in
preference to cation (B+) having higher standard reduction potential because their concentration
might be such that the reduction potential of A+ is higher than that of B+.
When two metal ions in the solution have identical values of their reduction potentials, the simultaneous
deposition of both the metals will occur in the form of an alloy.

GALVANIC CELL
This cell converts chemical energy into electrical energy.

Galvanic cell is made up of two half cells i.e., anodic and cathodic. The cell reaction is of redox kind.
Oxidation takes place at anode and reduction at cathode. It is also known as voltaic cell. It may be
represented as shown in Fig. Zinc rod immersed in ZnSO4 behaves as anode and copper rod immersed
in CuSO4 behaves as cathode.
Oxidation takes place at anode:
Zn Zn2+ + 2e– (loss of electron : oxidation)
Reduction takes place at cathode:
Cu2+ + 2e– Cu (gain of electron ; reduction)
Over all process:
Zn(s) + Cu2+ Cu(s) + Zn2+

In galvanic cell like Daniell cell; electrons flow from anode (zinc rod) to the cathode (copper rod)
through external circuit; zinc dissolves as Zn2+ ; Cu2+ ion in the cathode cell picks up two electron and
become deposited at cathode.

SALT BRIDGE
Two electrolyte solutions in galvanic cells are seperated using salt bridge as represented in the Fig. salt
bridge is a device to minimize or eliminate the liquid junction potential. Saturated solution of salt like
KCI, KNO3, NH4Cl and NH4NO3 etc. in agar-agar gel is used in salt bridge. Salt bridge contains high
concentration of ions viz. K+ and NO3 at the junction with electrolyte solution. Thus, salt bridge carries
whole of the current across the boundary ; more over the K+and NO3 ions have same speed. Hence,
salt bridge with uniform and same mobility of cations and anions minimize the liquid junction potential &
completes the electrical circuit & permits the ions to migrate.
Representation of a cell (IUPAC conventions ): Let us illustrate the convention taking the example of Daniel
cell.
(i) Anodic half cell is written on left and cathodic half cell on right hand side.
Zn(s) | ZnSO4 (sol) || CuSO4 (sol) | Cu(s)
(ii) Two half cells are separated by double vertical lines: Double vertical lines indicate salt bridge or any type
of porous partition.
(iii) EMF (electromotive force) may be written on the right hand side of the cell.
(iv) Single vertical lines indicate the phase separation between electrode and electrolyte solution.
Zn | Zn2+ || Cu2+ | Cu
(Illustration of Phase boundary)
(v) Inert electrodes are reprsented in the bracket
Zn | ZnSO4 || H+ | H2, Pt

CONCEPT OF ELECTROMOTIVE FORCE (EMF) OFA CELL

Electron flows from anode to cathode in external circuit due to a pushing effect called or electromotic
force (e.m.f.). E.m.f. is some times called as cell potential. Unit of e.m.f. of cell is volt.
EMF of cell may be calculated as :
Ecell = reduction potential of cathode Reduction potential of anode
Similarly, standard e.m.f. of the cell (E°) may be calculated as
E°cell = Standard reduction potential of cathode  Standard reduction potential of anode

SIGN CONVENTION OF EMF

EMF of cell should be positive other wise it will not be feasible in the given direction .
Zn | ZnSO4 || CuSO4 | Cu E = +1.10 volt (Feasible)
Cu | CuSO4 || ZnSO4 | Zn E = 1.10 volt (Not Feasible)
NERNST EQUATION
Walter Nernst derived a relation between cell potential and concentration or Reaction quotient.

G = G° + RT ln Q ..(1)
where G and G° are free energy and standard free energy change; ‘Q’ is reaction quotient.
Let n, Faraday charge is taken out from a cell of e.m.f. (E) then electrical work done by the cell
may be calculated as,
Work done = Charge x Potential = nFE
From thermodynamics we know that decrease in Gibbs free energy of a system is a measure of
reversible or maximum obtainable work by the system if there is no work due to volume expansion
 G = nFE and G° = nFE°
Thus from Eq. (i),we get nFE = -nFE° + RT lnQ
0.0591
At 25°C, above equation may be written as E  E0  log Q
n
Where ‘n’ represents number of moles of electrons involved in process.
E, E° are e.m.f. and standard e.m.f. of the cell respectively.
In general , for a redox cell reaction involving the transference of n electrons
aA + bB  cC + dD, the EMF can be calculated as:
 0.0591 [C]c [D]d
ECell = E°Cell – log
n [A ]a [B]b
Prediction and feasibility of spontaniety of a cell reaction.
Work done by the cell = nFE;
It is equivalent to decrease in free energy G = –nFE
Under standard state G0 = –nFE0 (i)
(i) From thermodynamics we know,  G = negative for spontaneous process. Thus from eq.(i) it is
clear that the EMF should be +ve for a cell process to be feasible or spontaneous.
(ii) When G = positive, E = negative and the cell process will be non spontaneous.
(iii) When G = 0 , E = 0 and the cell will attain the equilibrium.
Reactions G E
Spontaneous (–) (+)
Non- spontaneous (+) (–)
Equilibrium 0 0
Standard free energy change of a cell may be calculated by electrode potential data.
Substituting the value of E0 (i.e., standard reduction potentialof cathode- standard reduction potential of
anode) in eq. (i) we may get G0.
Let us see whether the cell (Daniell) is feasible or not: i.e. whether Zinc will displace copper or not.
Zn | (s) | ZnSO4 (sol) || CuSO4 (sol) | Cu(s)
E 0Zn 2 / Zn  0.76volt ; E 0Cu 2 / Cu  0.34volt

E 0 cell E 0Cu 2 / Cu  E 0zn 2 / Zn

=0.34 –(–0.76) = +1.10 volt
Since E0 = +ve , hence the cell will be feasible and zinc will displace copper from its salt solution. In the
other words zinc will reduce copper.

THERMODYNAMIC TREATMENT OF NERNST EQUATION

Determination of equilibrium constant : We know, that

0.0591
E  E0  logQ ..(i)
n
At equilibrium, the cell potential is zero because cell reactions are balanced, i.e. E = 0
 From Eq. (i), we have

0.0591  nE 0 
0 K eq anti log
0E  logK eq or 
n  0.0591 

Heat of Reaction inside the cell: Let n Faraday charge flows out of a cell of e.m.f. E, then
G = nFE (i)
Gibbs Helmholtz equation (from thermodynamics ) may be given as,
 G 
G = H + T  T  (ii)
p

From Eqs. (i) and (ii), we have

   nFE    E 
 nFE HT  HnFT  
 T  p  T  p
  E 
 H nFE nFT  
 T  p
Entropy change inside the cell : We know that G = H - TS or G = H  TS ...(i)
where G = Free energy change ; H = Enthalpy change and S = entropy change.
According to Gibbs Helmoholtz equation,
 G 
GH  T
 T  p

 G 
G  H  T 
 T  p
From Eqs. (i) and (ii), we have
 G   G 
 TST or S   
 T  p  T  p

 E 
or SnF 
 T  p

 E 
where  T  is called temperature coefficient of cell e.m.f.
p

DIFFERENT TYPES OF HALF-CELLS AND THEIR REDUCTION POTENTIAL

(1) Gas-Ion Half Cell:

In such a half cell, an inert collector of electrons, platinum or graphite is in contact with gas and a solution
containing a specified ion. One of the most important gas-ion half cell is the hydrogen-gas-hydrogen ion
half cell. In this half cell, purified H2gas at a constant pressure is passed over a platinum electrode which
is in contact with an acid solution.
H+(aq) + e– 1/2 H2

E  E  
0.0591
0
log
pH 2  1/ 2

H /H H /H 1 H  
2 2
 

(2) Metal-Metal Ion Half Cell:

This type of cell consist of a metal M in contact with a solution containing Mn+ ions.
Mn+(aq) + ne– M(s)
0.0591 1
E n
E 0 n
 log
M /M M /M n M n  
 

(3) Metal-Insoluble Salt - Anion Half Cell:

In this half cell, a metal coated with its insoluble salt is in contact with a solution containing the anion of the
insoluble salt. eg. Silver-Silver Chloride Half Cell:
This half cell is represented as Cl–/AgCl/Ag. The equilibrium reaction that occurs at the electrode is
AgCl(s) + e– Ag(s) + Cl–(aq)
 0.0591   
E E 0  log Cl
Cl

/ AgCl / Ag Cl

/ AgCl / Ag 1  
potential of such cells depends upon the concentration of anions. Such cells can be used asReference Electrode.
(4) Oxidation-reduction Half Cell:
This type of half cell is made by using an inert metal collector, usually platinum, immersed in a solution
which contains two ions of the same element in different states of oxidation. eg. Fe2+ - Fe3+ half cell.
Fe3+(aq) + e– Fe2+(aq)
Fe 2  
0 0.0591  
E E  log
Fe
3
/ Fe
2
Fe
3
/ Fe
2 1 Fe 3  
 
CONCENTRATION CELL
The cells in which electrical current is produced due to transport of a substance from higher to lower
concentration. Concentration gradient may arise either in electrode material or in electrolyte. Thus there
are two types of concentration cell .
(i) Electrode concentration cell
(ii) Electrolyte concentration cell

Electrode Gas concentration cell :

Pt, H2(P1) | H+(C) | H2(P2), Pt
Here, hydrogen gas is bubbled at two different partial pressures at electrode dipped in the solution of
same electrolyte.
Cell process : 1 / 2H 2 ( p1 )H  (c)e  (Anode process)

 
H (c)e 1 / 2H 2 ( p2 )
1 / 2H 2 ( p1)  1 / 2H 2 ( p2 )

1/ 2
2.303RT  p 2 
 E  log  
F  p1 

 2.303RT   p 2  0.059  p1 
or E log
   , At 25 C ,
0 E  log  
 2F   p1  2F p
 2 
For spontanity of such cell reaction, p1>p2
Electrolyte concentration cells:
Zn(s) | ZnSO4 (C1) || ZnSO4 (C2) | Zn(s)
In such cells, concentration gradient arise in electrolyte solutions. Cell process may be givewn as,
Zn s Zn 2  C1  2e (Anodic process)

Zn 2 (C 2 )  2e  Zn (s)
(Over all process)
Zn 2 (C 2 ) Zn 2 (C1 )

 From Nernst equation, we have

 2.303RT  C1  2.303RT  C 2 
E0 log   or E log  
2F  C 2  2F  C1 
For spontanity of such cell reaction, C2> C1
CONDUCTANCE
Introduction:
Both metallic and electrolytic conductors obey Ohm's law
i.e. V = IR
where V = Potential difference in volt; I = Current in ampere ; R = resistance in Ohm
We know, resistance is directly proportional to length of conductor and inversely proportional to
cross sectional area of the conductor.
l l
R or R  (  = Specific resistance )
A A
Specific resistance is the resistance of a conductor having lengths of 1 cm and corss sectional
area of 1 cm2.
Unit of R is ohm and unit of specific resistance is ohm cm
Reciprocal of resistance is called as conductance and reciprocal of specific resistance is called as
specific conductance.
1 1A A
 or CK
R l l
where C = conductance ohm ; K = specific conductance ohm1cm1 .
1

Mho and siemens are other units of conductance

l
K C
A
Specific conductance= Cell constant x Conductance

SPECIFIC CONDUCTANCE IS CONDUCTANCE OF 1 CM3 OFAN ELECTROLYTE SOLUTION.

In case of electrolytic solution, the specific conductance is defined as the conductance of a solsution of
definite concentration enclosed in a cell having two electrodes sof unit area separated by 1 cm apart.
1. Equivalent Conductance
Equivalent conductance is the conductance of an electrolyte solution containing 1 gm equivalent
of electrolyte. It is densoted by  .
=Kx V
( = ohm cm1 x cm3 = ohm1 cm2)
 1

Usually concen ration of electrolyte solution is expressed as C gm equivalent per litre.

1000
Thus, V
C

Volumehaving1gmequivalentelectrolytein thesolutionThus, K  1000

C
2. Molar Conductance
Molar conductance may be defined as conductance of an electrolyte solution having 1 gm mole
electrolyte in a litre. It is denoted by  m .
m=K×V
Usually concentration of electrolyte solution is expressed as ‘M’ gm mole elecrtrolyte per litre.
 1000
Thus, V
M
1000
Hence,  m K x
M
Relation between  and  m :  m= n × 
DETERMINATION OF  0m OR  0
A plot of  m vs C as found experimentally is as shown below graphically.

The m vs C plot of strong electrolyte being linear it can be extrapolated to zero concentration.
Thus, m values of the solution of the test electrolyte are determined at various concentrations the
concentrations should be as low as good.
m values are then plotted against C when a straight line is obtained. This is the extrapolated to zero
concentration. The point where the straight line intersects m axis is  0m of the strong electrolyte.
However, the plot in the case weak electrolyte being non linear, shooting up suddenly at some low
concentration and assuming the shape of a straight line parallel to m axis. Hence extrapolation in this
case is not possible. Thus, 0 of a weak electrolyte cannot be determined experimentally. It can, however,
be done with the help of Kohlrausch's law to be discussed later.
Kohlrausch's Law of Independent Migration of Ions
Kohlrausch determined 0 values of pairs of some strong electrolytes containing same cation say KF
and KCl, NaF and NaCl etc., and found that the difference in 0 values in each case remains the same:
 0m (KCl) –  0m (KF) =  0m (NaCl) –  0m (NaF)
He also detemined 0 values of pairs of strong electrolytes containing same anion say KF and NaF, KCl
and NaCl etc.and found that the difference in 0 values in each case remains the same.
 0m (KF) –  0m (NaF) =  0m (KCl) –  0m (NaCl)
This experimental data led him to formulate the following law called Kohlrausch's law of independent
migration of ions.
At infinite dilution when dissociation is complete, every ion makes some definite contribution towards
molar conductance of the electrolyte irrespective of the nature of the other ion which with it is associated
and that the molar conductance at infinite dilution for any electrolyte is given by the sum of the contribution
of the two ions. Thus,
 0m = 0  0
 Where 0 is the contribution of the cation and 0 is the contribution of the anion towards the
molar conductance at infinite dilution. These contributions are called molar ionic conductances at
infinite dilution. Thus, 0 is the molar ionic conductance of cation and 0 is the molar ionic
conductnace of anion, at infinite dilution. The above equation is, however, correct only for binary
electrolyte like NaCl, MgSO4 etc.
Application of Kohlrausch's law :
(1) Determination of  0m of a weak electrolyte:
In order to calculate  0m of a weak electrolyte say CH3COOH, we determine experimentally  0m values
of the following three strong electrolytes:
(a) A strong electrolyte containing same cation as in the test electrolyte, say HCl
(b) A strong electrolyte containing same anion as in the test electrolyte, say CH3COONa
(c) A strong electrolyte containing same anion of (a) and cation of (b) i.e. NaCl.
 0m of CH3COOH is then given as:
 0m (CH3COOH) =  0m (HCl) +  0m (CH3COONa) –  0m (NaCl)
Proof:
0
 0m (HCl) =  H   Cl ......................(i)
0
 0m (CH3COONa) =  CH3COO    Na  ......................(ii)
0 0
 0m (NaCl) =  Na    Cl ......................(iii)
Adding equation (i) and equation (ii) and subtracting (iii) from them:
0 0 0
 0m (HCl) +  ( CH 3COONa ) –  ( NaCl)   (H  )   (CH3COO0 )   0( CH3COOH)
0

(2) Determination of degree of dissociation () :

No. of moleculesionised 
=  m
total numberof moleculesdissolved  0m
(3) Determination of solubility of sparingly soluble salt
The specific conductivity of a saturated solution of the test electrolyte (sparingly soluble) made in
conductivity water is determined by the method as described above. From this the specific conductivity
of conductivity water is deducted. The molar conductance of the saturated solution is taken to be equal
to  0m as the saturated solution of a sparingly soluble salt is extremely dilute. Hence from equation (4).
1000
 0m = ,
C
where C is the molarity of solution and hence the solubility.
ATLAS
 EASY RIDE

ELECTROLYTIC CELL :

Q.1 Calculate the no. of electrons lost or gained during electrolysis of

(a) 3.55 gm of Cl– ions (b) 1 gm Cu2+ ions (c) 2.7 gm of Al3+ ions
Q.2 How many faradays of electricity are involved in each of the case
(a) 0.25 mole Al3+ is converted to Al.
(b) 27.6 gm of SO3 is convered to SO 32
(c) The Cu2+ in 1100 ml of 0.5 M Cu2+ is converted to Cu.
Q.3 0.5 mole of electron is passed through two electrolytic cells in series. One contains silver ions, and the
other zinc ions. Assume that only cathode reaction in each cell is the reduction of the ion to the metal.
How many gm of each metals will be deposited.
Q.4 The electrosynthesis of MnO2 is carried out from a solution of MnSO4 in H2SO4 (aq). If a current of
25.5 ampere is used with a current efficiency of 85%, how long would it take to produce 1 kg of MnO2?
Q.5 A constant current of 30 A is passed through an aqueous solution of NaCl for a time of 1.0 hr. How many
grams of NaOH are produced? What is volume of Cl2 gas at S.T.P. produced?
Q.6 If 0.224 litre of H2 gas is formed at the cathode, how much O2 gas is formed at the anode under identical
conditions?
Q.7 If 0.224 litre of H2 gas is formed at the cathode of one cell at S.T.P., how much of Mg is formed at the
cathode of the other electrolytic cell.
Q.8 Assume 96500 C as one unit of electricity. If cost of electricity of producing x gm Al is Rs x, what is the
cost of electricity of producing x gm Mg?
Q.9 Chromium metal can be plated out from an acidic solution containing CrO3 according to following equation:
CrO3(aq) + 6H+ (aq) + 6e–  Cr(s) + 3H2O
Calculate :
(i) How many grams of chromium will be plated out by 24000 coulombs and
(ii) How long will it take to plate out 1.5 gm of chromium by using 12.5 ampere current
Q.10 Calculate the quantity of electricity that would be required to reduce 12.3 g of nitrobenzene to aniline, if
the current efficiency for the process is 50 percent. If the potential drop across the cell is 3.0 volts, how
much energy will be consumed?
Q.11 How long a current of 2A has to be passed through a solution of AgNO3 to coat a metal surface of
80cm2 with 5m thick layer? Density of silver = 10.8g/cm3.

Q.12 3A current was passed through an aqueous solution of an unknown salt of Pd for 1Hr. 2.977g of Pd+n
was deposited at cathode. Find n.
Q.13 50mL of 0.1M CuSO4 solution is electrolyzed with a current of 0.965 A for a period of
200 sec. The reactions at electrodes are:
Cathode : Cu2+ + 2e  Cu(s) Anode : 2H2O  O2 + 4H+ + 4e.
Assuming no change in volume during electrolysis, calculate the molar concentration of Cu2+, H+ and
SO42at the end of electrolysis.
Q.14 A metal is known to form fluoride MF2. When 10A of electricity is passed through a molten salt for 330
sec., 1.95g of metal is deposited. Find the atomic weight of M. What will be the quantity of electricity
required to deposit the same mass of Cu from CuSO4?
Q.15 10g fairly concentrated solution of CuSO4 is electrolyzed using 0.01F of electricity. Calculate:
(a)The weight of resulting solution (b)Equivalents of acid or alkali in the solution.

Q.16 An electric current is passed through electrolytic cells in series one containing Ag(NO3)(aq.) and other
H2SO4(aq). What volume of O2 measured at 250C and 750mm Hg pressure would be liberated from
H2SO4 if
(a) one mole of Ag+ is deposited from AgNO3 solution
(b) 8 x 1022 ions of Ag+ are deposited from AgNO3 solution.

Q.17 Cadmium amalgam is prepared by electrolysis of a solution of CdCl2 using a mercury cathode. How
long should a current of 5A be passed in order to prepare 12% CdHg amalgam on a cathode of 2gm
Hg (Cd=112.4)

Q.18 After electrolysis of NaCl solution with inert electrodes for a certain period of time. 600 mL of the
solution was left. Which was found to be 1N in NaOH. During the same time, 31.75 g of Cu was
deposited in the copper voltameter in series with the electrolytic cell. Calculate the percentage yield of
NaOH obtained.

Q.19 Three electrolytic cells A, B, C containing solution of ZnSO4, AgNO3 and CuSO4, respectively are
connected in series. A steady current of 2 ampere was passed through them until 1.08 g of silver deposited
at the cathode of cell B. How long did the current flow? What mass of copper and of zinc were deposited?

Q.20 Copper sulphate solution (250 mL) was electrolysed using a platinum anode and a copper cathode. A
constant current of 2 mA was passed for 16 minutes. It was found that after electrolysis the concentration
of the solution was reduced to 50% of its original value. Calculate the concentration of copper sulphate
in the original solution.

Q.21 A solution of Ni(NO3)2 is electrolysed between platinum electrodes using a current of 5 ampere for 20
mintue. What mass of Ni is deposited at the cathode?

Q.22 A current of 3.7A is passed for 6hrs. between Ni electrodes in 0.5L of 2M solution of Ni(NO3)2.
What will be the molarity of solution at the end of electrolysis?

GALVANIC CELL :
Representation of Cell diagrams, complete and half cell reactions :
Q.23 Make complete cell diagrams of the following cell reactions :
(a) Cd2+ (aq) + Zn (s) Zn2+ (aq) + Cd (s)
(b) +
2Ag (aq) + H2 (g) 2H+ (aq) + 2Ag (s)
(c) Hg2Cl2 (s) + Cu (s) Cu2+ (aq) + 2Cl– (aq) + 2Hg (l)
(d) Cr2O 72 (aq.) + 14H+ (aq) + 6Fe2+ (aq) 6Fe3+ (aq) + 2Cr3+ (aq) + 7H2O (l)

Q.24 Write cell reaction of the following cells :

(a) Ag | Ag+ (aq) | | Cu2+ (aq) | Cu
(b) Pt | Fe2+ , Fe3+ | | MnO 4 , Mn2+, H+ | Pt
(c) Pt,Cl2 | Cl– (aq) | | Ag+ (aq) | Ag
(d) Pt, H2 | H+ (aq) | | Cd2+ (aq) | Cd
Q.25 Write half cells of each cell with following cell reactions :
(a) Zn (s) + 2H+ (aq) Zn2+ (aq) + H2 (g)
(b) 2Fe3+ (aq) + Sn2+ (aq) 2Fe2+ (aq) + Sn4+ (aq)
(c) MnO 4 (aq) + 8H+ (aq) + 5Fe2+ (aq) 2Fe3+ (aq) + Mn2+ (aq) + 4H2O (l)
(d) Pb (s) + Br2 (l) Pb2+ (aq) + 2Br– (aq)

Electrode potential and standard electrode potential :

Q.26 For the cell reaction 2Ce4+ + Co 2Ce3+ + Co2+
E ocell is 1.89 V. If E o is – 0.28 V, what is the value of E o ?
Co 2  | Co Ce 4  | Ce3

Q.27 Determine the standard reduction potential for the half reaction :
Cl2 + 2e–  2Cl–
Given Pt2+ + 2Cl– Pt + Cl2, E oCell = – 0.15 V
Pt2+ + 2e– Pt E° = 1.20 V

o
Q.28 What is E Cell if
2Cr + 3H2O + 3OCl– 2Cr3+ + 3Cl– + 6OH–
3+
2Cr + 3e – Cr,, E° = – 0.74 V
OCl + H2O + 2e–
– Cl– + 2OH–, E° = 0.94 V
o
G°, E Cell and Keq :
Q.29 Is 1.0 M H+ solution under H2SO4 at 1.0 atm capable of oxidising silver metal in the presence of 1.0 M
Ag+ ion?
o
Eo = 0.80 V, E H |H = 0.0 V
Ag  |Ag 2 ( Pt )

Q.30 If for the half cell reactions Cu2+ + e– Cu+ E° = 0.15 V

2+
Cu + 2e – Cu E° = 0.34 V
Calculate E° of the half cell reaction
Cu+ + e– Cu
+
also predict whether Cu undergoes disproportionation or not.

o o o
Q.31 If E Fe2 | Fe = – 0.44 V, E Fe3 | Fe 2 = 0.77 V. Calculate E Fe3 | Fe .
o o o
Q.32 If E Cu  |Cu = 0.52 V, E Cu 2 |Cu = 0.34 V, what is E Cell of the cell reaction
Cu + Cu2+ 2Cu+?
is cell reaction spontaneous?

Q.33 Calculate the EMF of a Daniel cell when the concentration of ZnSO4 and CuSO4 are 0.001 M and
0.1M respectively. The standard potential of the cell is 1.1V.

Q.34 Calculate the equilibrium constant for the reaction Fe + CuSO4 FeSO4 + Cu at 250C.
Given E0 (Fe/Fe2+) = 0.44V, E0 (Cu/Cu2+) = 0.337V.
Q.35 For a cell Mg(s) | Mg2+(aq) || Ag+ (aq) | Ag, Calculate the equilibrium constant at 250C. Also find the
maximum work that can be obtained by operating the cell.
E0 (Mg2+/Mg) = 2.37V, E0 (Ag+/Ag) = 0.8 V.

Q.36 The standard reduction potential at 250C for the reduction of water
2H2O + 2e H2 + 2OH is 0.8277 volt. Calculate the equilibrium constant for the reaction
2H2O l H3O + OH at 250C.
+

Q.37 At 250C the value of K for the equilibrium Fe3+ + Ag Fe2+ + Ag+ is 0.531mol/litre. The standard
electrode potential for Ag + + e – Ag is 0.799V. What is the standard potential for
Fe3+ + e– Fe2+ ?

Q.38 The EMF of the cell M | Mn+ (0.02M) || H+ (1M) | H2(g) (1 atm), Pt at 250C is 0.81V. Calculate the
valency of the metal if the standard oxidation of the metal is 0.76V.

Q.39 Equinormal Solutions of two weak acids, HA (pKa = 3) and HB (pKa = 5) are each placed in contact
with standard hydrogen electrode at 250C. When a cell is constructed by interconnecting them through
a salt bridge, find the emf of the cell.

Q.40 In two vessels each containing 500ml water, 0.5m mol of aniline (Kb= 109) and 25mmol of HCl are
added separately. Two hydrogen electrodes are constructed using these solutions. Calculate the emf of
cell made by connecting them appropriately.

Q.41 Calculate E0 and E for the cell Sn | Sn2+ (1M) || Pb2+ | Pb(103M), E0 (Sn2+| Sn) = 0.14V,
E0 (Pb2+| Pb) = 0.13V. Is cell representation is correct?

Q.42 At what concentration of Cu2+ in a solution of CuSO4 will the electrode potential be zero at 250C?
Given : E0 (Cu | Cu2+) = 0.34 V.

Q.43 A zinc electrode is placed in a 0.1M solution at 250C. Assuming that the salt is 20% dissociated at this
dilutions calculate the electrode potential. E0 (Zn2+| Zn) = 0.76V.

Q.44 From the standard potentials shown in the following diagram, calculate the potentials E1 and E 2 .

Q.45 For the reaction, 4Al(s) + 3O2 (g) + 6H2 O + 40H– 4 [Al(OH)4– ] ; Ecell = 2.73 V. If
G f (OH  ) = –157 kJ mol–1 and G f ( H 2 O) = –237.2 kJ mol–1, determine G f [Al (OH)4–].
Concentration cells :

Q.46 Calculate the EMF of the following cell

Zn | Zn2+ (0.01M) || Zn2+ (0.1 M) | Zn
at 298 K.

Q.47 Calculate the EMF of the cell,

Zn – Hg(c1M) | Zn2+ (aq)| Hg – Zn(c2M)
at 25°C, if the concentrations of the zinc amalgam are: c1 = 10g per 100g of mercury and
c2 = 1g per100 g of mercury.

Q.48 Calculate pH using the following cell :

Pt (H2) | H+ (x M) | | H+ (1 M) | Pt (H2) if Ecell = 0.2364 V.
1 atm 1 atm

Q.49 Calculate the EMF of following cells at 250C.

(i) Fe | Fe2+ (a1 = 0.3) || Sn2+ (a2 = 0.1) | Sn E0 (Fe2+/Fe) = –0.44 V
(ii) Pt, H2 (2atm) | HCl |H2 (10 atm), Pt. E0 (Sn2+/Sn) = 0.14 V

Q.50 EMF of the cell Zn | ZnSO4 (a1= 0.2) || ZnSO4(a2) | Zn is 0.0088V at 250C. Calculate the value of a2.

CONDUCTANCE
Conductivities and cell constant:

Q.51 The resistance of a conductivity cell filled with 0.01N solution of NaCl is 210 ohm a t 1 8 o C .
Calculate the equivalent conductivity of the solution. The cell constant of the conductivity cell is
0.88 cm1.
Q.52 The molar conductivity of 0.1 M CH3COOH solution is 4.6 S cm2 mole1 . What is the specific
conductivity and resistivity of the solution ?

Q.53 The conductivity of pure water in a conductivity cell with electrodes of cross sectional area 4 cm2
and 2 cm apart is 8 x 107 S cm1.
(i) What is resistance of conductivity cell ?
(ii) What current would flow through the cell under an applied potential difference of 1 volt?

Q.54 Resistivity of 0.1M KCl solution is 213 ohm cm in a conductivity cell. Calculate the cell constant
if its resistance is 330 ohm.

Q.55 Resistance of a 0.1M KCl solution in a conductance cell is 300 ohm and specific conductance of
0.1M KCl is 1.29 x 10-2 ohm-1 cm-1. The resistance of 0.1M NaCl solution in the same cell is 380
ohm. Calculate the equivalent conductance of the 0.1M NaCl solution.

Q.56 For 0.01N KCl, the resistivity 709.22 mho cm. Calculate the conductivity and equivalent
conductance.

Q.57 A solution containing 2.08 g of anhydrous barium chloride is 400 CC of water has a specific
conductivity 0.0058 ohm–1cm–1. What are molar and equivalent conductivities of this solution.
Application of Kohlrausch's law:

Q.58 Equivalent conductance of 0.01 N Na2SO4 solution is 112.4 ohm–1 cm2 eq–1. The equivalent
conductance at infinite dilution is 129.9 ohm–1 cm2. What is the degree of dissociation in 0.01 N
Na2SO4 solution?

Q.59 Specific conductance of a saturated solution of AgBr is 8.486×10–7 ohm–1cm–1 at 250C. Specific
conductance of pure water at 250C is 0.75×10–6 ohm–1 cm–2. m for KBr , AgNO3 and KNO3
are 137.4 , 133 , 131 ( S cm2 mol–1) respectively. Calculate the solubility of AgBr in gm/litre.

Q.60 Saturated solution of AgCl at 250C has specific conductance of 1.12×10–6 ohm–1 cm–1. The
  Ag+ and Cl– are 54.3 and 65.5 ohm–1 cm2 / equi. respectively. Calculate the solubility product
of AgCl at 250C.

Q.61 Hydrofluoric acid is weak acid. At 25 0 C, the molar conductivity of 0.002M HF is

176.2 ohm–1 cm2 mole–1. If its m = 405.2 ohm–1 cm2 mole–1, calculate its degree of dissociation
and equilibrium constant at the given concentration.

Q.62 The value of m for HCl, NaCl and CH3CO2Na are 426.1, 126.5 and 91 S cm2 mol–1 respectively..
Calculate the value of m for acetic acid. If the equivalent conductivity of the given acetic acid is 48.15
at 25° C, calculate its degree of dissociation.

Q.63 Calculate the specific conductance of a 0.1 M aqueous solution of NaCl at room temperature,
given that the mobilities of Na + and Cl– ions at this temperature are 4.26×10 –8 and
6.80×10–8 m2 v–1 s–1, respectively.

Q.64 For the strong electroytes NaOH, NaCl and BaCl2 the molar ionic conductivities at infinite dilution
are 248.1×10–4, 126.5 ×10–4 and 280.0 ×10-4 mho cm2 mol–1 respectively. Calculate the molar
conductivity of Ba(OH)2 at infinite dilution.

Q.65 At 25°C,  (H+) = 3.4982 ×10 –2 S m2 mol–1 and (OH–) = 1.98 ×10–2 S m2mol–1.
Given: Sp. conductnace = 5.1 ×10–6 S m–1 for H2O, determine pH and Kw.
 PROFICIENCY TEST

1. In highly alkaline medium, the anodic process during the electrolytic process is
4OH–  O2 + 2H2O + 4e–.

2. Compounds of active metals (Zn, Na, Mg) are reducible by H2 whereas those of noble metals (Cu, Ag,
Au) are not reducible.

EIt
3. The mass of a substance deposited on the cathode or anode during electrolysis is given by w = .
F

4. Faraday’s second law of electrolysis is related to the equivalent, mass of the electrolyte.

5. Equivalent conductance at infinite dilution of salt AB is equal to the sum of equivalent conductances of
ions, A+ and B– at infinite dilution.

6. The standard reduction potential of Cl– | AgCl | Ag half-cell is related to that of Ag+ | Ag half-cell through
the expression E    E   RT In K (AgCl).
Ag |Ag Cl |AgCl |Ag
 SP
F

7. The cell potential is given by Ecell = ERP(cathode)– ERP(anode).

8. A half-cell reaction is A (x+n) + ne–  Ax+. It is possible to determine the value of n from the
measurements of cell potential.

9. In a galvanic cell, the half-cell with higher reduction potential acts as a reducing agent.

10. In an electrode concentration cell, the cel reaction Zn(c1)  Zn(c2) will be spontaneous if c1 > c2.

11. The absolute value of electrode potential cannot be determined.

12. All chemical reactions used in galvanic cells are redox reactions.

13. The amount of the product formed by the passage of 1 coulomb of electricity through electrolyte is called
electrochemical equivalent of the substance.

14. The redox reaction involved in galvanic cell is a non- spontaneous process.

15. In galvanic cell, the cathode is a – ve terminal of the cell.

16. The specific conductance of a 0.1 N KCl solution at 23°C is 0.012 ohm–1 cm–1. The resistance of the
cell containing the solution at the same temperature was found to be 55 ohms. The cell constant is ____.

17. Dilute sulphuric acid on electrolysis liberates___________ at the anode.

18. The electrical conductivity of a solution of acetic acid will _______ if a solution of sodium hydroxide is
added.

19. A cation having a ________reduction potential is preferentially reduced at the cathode.

20. When an aqueous solution of sodium sulphate is electrolysed, the gases liberated at the anode &
cathode are ________ and __________, respectively.

21. A cell in which two electrodes of the same metal are dipped in solutions of metal ion of different
concentrations in called___________.

22. The half-cell involving the reaction,

Cr2O 72 (aq.) +14H+(aq.) + 6e– 2Cr3+(aq.) + 7H2O
is represented as _____________.

23. During discharge of lead storage battery, the overall reaction is___________.

24. In the calomel half-cell, the reduction reaction is ___________.

25. Temperature coefficient and change in enthalpy are related by the expression__________.
26. In salt bridge, the electrolyte used should be _________.

27. In electrochemical cell, the electrical neutrality in two half cells is maintained by _________.

28. The E° value for H2  2H+ + 2e– is ____________.

29. E°cell of E°oxi.(anode) + __________.

30. Coulomb refers to _______ of electricity while ampere refers to ________ at which it flows.

31. The cathodic reactions always involve__________.

32. During electrolysis of aqueous solution of CuSO4 using Pt electrodes the product at anode is ______.
33. The quantity of electricity required for complete reduction of 0.5 mole MnO 4 to Mn2+ is ______C.

34. During electrolysis process ________energy is converted into _______.

35.  eq × normality = _________.

 MIDDLE GAME

Q.1 Same quantity of electricity is being used to liberate iodine (at anode) and a metal x (at cathode). The
mass of x deposited is 0.617g and the iodine is completely reduced by 46.3 cc of 0.124M sodium
thiosulphate. Find the equivalent mass of x.
Q.2 The standard reduction potential values, E0(Bi3+|Bi) and E0(Cu2+|Cu) are 0.226V and 0.344V
respectively. A mixture of salts of bismuth and copper at unit concentration each is electrolysed at 250C.
to what value can [Cu2+] be brought down before bismuth starts to deposit, in electrolysis.
Q.3 In a fuel cell, H2 & O2 react to produce electricity. In the process, H2 gas is oxidized at the anode & O2
at the cathode . If 67.2 litre of H2 at STP react in 15 minutes, what is the average current produced ? If
the entire current is used for electrode deposition of Cu from Cu (II) solution, how many grams of Cu
will be deposited?
Anode : H2 + 2OH  2H2O + 2 e– Cathode : O2 + 2 H2O + 4e  4 OH–
Q.4 One of the methods of preparation of per disulphuric acid, H2S2O8, involve electrolytic oxidation of
H 2 SO 4 at anode (2H 2 SO 4  H 2 S 2 O 8 + 2H + + 2e  ) with oxygen and hydrogen as
byproducts. In such an electrolysis, 9.722 L of H2 and 2.35 L of O2 were generated at STP. What is
the weight of H2S2O8 formed?
Q.5 During the discharge of a lead storage battery the density of sulphuric acid fell from 1.294 to
1.139 g.ml1. H2SO4 of density 1.294 g mL1 is 39% and that of density 1.39 g mL1 is 20% by
weight. The battery holds 3.5L of acid and the volume practically remains constant during the discharge.
Calculate the number of ampere hours for which the battery must have been used. The discharging
reactions are:
Pb + SO42 PbSO4 + 2e (anode)
+
PbO2 + 4H + SO4 + 2e 2  PbSO4 + 2H2O (cathode)
Q.6 The emf of the cells obtained by combining Zn and Cu electrode of a Daniel cell with N calomel electrode
in two different arrangements are 1.083V and 0.018V respectively at 250C. If the standard reduction
potential of N calomel electrode is 0.28V and that of Zn is 0.76 V, find the emf of Daniel cell.
Q.7 Given t he standard reduction pot entials Tl+ + e  TI, E 0 = 0.34V and
TI3+ + 2e TI+, E0 = 1.25V. Examine the spontaneity of the reaction, 3TI+ 2TI + TI3+. Also
0
find E for this disproportionation.
Q.8 The emf of the cell Ag|AgI|KI(0.05M) || AgNO3(0.05M) |Ag is 0.788V. Calculate the solubility product
of AgI.
Q.9 The cell Pt, H2(1 atm) | H+(pH=x) || Normal calomel Electrode has an EMF of 0.67V at 250C. Calculate
the pH of the solution. The oxidation potential of the calomel electrode on hydrogen scale is 0.28 V.
Q.10 Estimate the cell potential of a Daniel cell having 1 M Zn++ & originally having 1 M Cu++ after sufficient
NH3 has been added to the cathode compartment to make NH3 concentration 2 M.
Kf for [Cu(NH3)4]2+ = 1 x 1012, E0 for the reaction,
Zn + Cu2+ Zn2+ + Cu is 1.1 V..
Q.11 Consider the cell Ag|AgBr(s)|Br ||AgCl(s), Ag | Cl at 25º C . The solubility product constants of AgBr
& AgCl are respectively 5 x 1013 & 1 x 1010 . For what ratio of the concentrations of Br & Cl ions
would the emf of the cell be zero ?
Q.12 The pKsp of Agl is 16.07 . If the Eº value for Ag+Ag is 0.7991 V . Find the Eº for the half cell reaction
AgI (s) + e Ag + I.
Q.13 Voltage of the cell Pt, H2 (1 atm)|HOCN (1.3 x 103 M)||Ag+ (0.8 M)|Ag(s) is 0.982 V . Calculate the
Ka for HOCN . Neglect [H+] because of oxidation of H2(g) .
Ag+ + e Ag(s) = 0.8 V..
Q.14 The normal potential of Zn referred to SHE is 0.76V and that of Cu is 0.34V at 250C. When excess of
Zn is added to CuSO4, Zn displaces Cu2+ till equilibrium is reached. What is the ratio of Zn2+ to Cu2+
ions at equilibrium?
Q.15 Calculate the potential of an indicator electrode versus the standard hydrogen electrode, which originally
contained 0.1M MnO4 and 0.8M H+ and which was treated with 90% of the Fe2+ necessary to reduce
all the MnO4 to Mn+2.
MnO4 + 8H+ + 5e Mn2+ + 4H2O, E0 = 1.51V

Q.16 Kd for complete dissociation of [Ag(NH3)2]+ into Ag+ and 2NH3 is 6 x 108. Calculate E0 for the
following half reaction; Ag(NH3)2+ + e Ag + 2NH3
+
Ag + e  Ag, 0
E = 0.799 V
Q.17 The overall formation constant for the reaction of 6 mol of CN with cobalt (II) is
1 x 1019. The standard reduction potential for the reaction
[Co(CN)6]3 + e Co(CN)64 is 0.83 V. Calculate the formation constant of [Co(CN)6]3
Given Co3+ + e Co2+ ; E0 = 1.82 V
Q.18 Calculate the emf of the cell
Pt, H2(1.0 atm) | CH3COOH (0.1M) || NH3(aq, 0.01M) | H2 (1.0 atm),
Pt Ka(CH3COOH) = 1.8 x 105, Kb (NH3) = 1.8 x 105.
Q.19 A current of 3 amp was passed for 2 hour through a solution of CuSO4 ,3 g of Cu2+ ions were deposited
as Cu at cathode. Calculate percentage current efficiency of the process.
Q.20 The Edison storage cell is represented as Fe(s) | FeO(s) | KOH(aq) | Ni2O3 (s) | Ni(s) The halfcell
reaction are
Ni2O3(s) + H2O(i) + 2e 2NiO(s) + 2OH E0 = + 0.40V
FeO(s) + H2O(l) + 2e Fe(s) + 2OH E0 =  0.87V
(i) What is the cell reaction?
(ii) What is the cell e.m.f.? How does it depend on the concentration of KOH?
(iii) What is the maximum amount of electrical energy that can be obtained from one mole of Ni2O3?
Q.21 For the galvanic cell : Ag|AgCl(s)| KCl (0.2M) || K Br (0.001 M)| AgBr(s) | Ag,
Calculate the EMF generated and assign correct polarity to each electrode for a spontaneous process
after taking into account the cell reaction at 250C.
[K sp( AgCl )  2.8x1010 ;K sp (AgBr )  3.3x10 13 ]

Q.22 An aqueous solution of NaCl on electrolysis gives H2(g), Cl2(g) and NaOH according to the reaction:
2Cl(aq) + 2H2O 2OH(aq) + H2(g) + Cl2(g)
A direct current of 25 amperes with a current efficiency of 62% is passed through 20 liters of NaCl
solution (20% by weight). Write down the reactions taking place at the anode and the cathode. How
long will it take to produce 1Kg of Cl2? What will be the molarity of the solution with respect to hydroxide
ion? (Assume no loss due to evaporation).

Q.23 An acidic solution of Cu2+ salt containing 0.4 g of Cu2+ is electrolyzed until all the copper is deposited.
The electrolysis is continued for seven more minutes with the volume of solution kept at 100 ml and the
current at 1.2 amp. Calculate the volume of gases evolved at NTP during the entire electrolysis.
Q.24 In the refining of silver by electrolytic method what will be the weight of 100 gm Ag anode if
5 ampere current is passed for 2 hours? Purity of silver is 95% by weight.
Q.25 Hydrogen peroxide can be prepared by successive reactions:
2NH4HSO4  H2 + (NH4)2S2O8
(NH4)2S2O8 + 2H2O  2NH4HSO4 + H2O2
The first reaction is an electrolytic reaction the second is steam distillation. What amount of current
would have to be used in first reaction to produce enough intermediate to yield 100 gm pure H2O2 per
hour? Assume 50% anode current efficiency.
Q.26 Dal lake has water 8.2×1012 litre approximately. A power reactor produces electricity at the rate of
1.5×106coulomb per second at an appropriate voltage.How many years would it take to electrolyse the
lake?
Q.27 Calculate the potential at 25°C for the cell
Cd | Cd2+ (2.00 M) || Pb2 (0.0010 M) | Pb
Given E°cell = 0.277 V.
Q.28 Calculate E° for the following reactions at 298 K,
Ag ( NH 3 ) 2 + e– Ag + 2NH3

Ag (CN ) 2 + e– Ag + 2CN–

Given: E   0.7991 V, K Ins [Ag ( NH 3 ) 2 ] = 6.02 × 10–8 and K Ins [Ag(CN ) 2 ] = 1.995 ×10–19
Ag |Ag

Q.29 Determine the degree of hydrolysis and hydrolysis constant of aniline hydrochloride in M/32 solution of
salt at 298 K from the following cell data at 298 K.
Pt | H2 (1 atm) | H+(1M) || M/32 C6H5NH3Cl | H2 (1 atm) | Pt ; Ecell= – 0.188 V.
Q.30 The emf of the cell, Pt | H2 (1 atm), | H+ (0.1 M, 30 ml) || Ag+ (0.8 M) | Ag is 0.9 V. Calculate the emf
when 40 ml of 0.05 M NaOH is added.
Q.31 Given, E° = –0.268 V for the Cl– | PbCl2 | Pb couple and – 0.126 V for the Pb2+ | Pb couple, determine
Ksp for PbCl2 at 25°C?
Q.32 The equivalent conductance of 0.10 N solution of MgCl2 is 97.1 mho cm2 equi–1 at 250C. a cell
with electrode that are 1.5 cm2 in surface area and 0.5 cm apart is filled with 0.1 N MgCl2 solution.
How much current will flow when potential difference between the electrodes is 5 volt.
Q.33 A dilute aqueous solution of KCl was placed between two electrodes 10 cm apart, across which a
potential of 6 volt was applied. How far would the K+ ion move in 2 hours at 250C? Ionic
conductance of K+ ion at infinite dilution at 250C is 73.52 ohm–1 cm2 mole–1?
Q.34 When a solution of specific conductance 1.342 ohm–1 metre–1 was placed in a conductivity cell with
parallel electrodes, the resistance was found to be 170.5 ohm. Area of electrodes is 1.86×10–4 m2.
Calculate separation of electrodes.
Q.35 The specific conductance at 250C of a saturated solution of SrSO4 is 1.482×10–4 ohm–1 cm–1while
that of water used is 1.5×10–6 mho cm–1. Determine at 250C the solubility in gm per litre of SrSO4
in water. Molar ionic conductance of Sr2+ and SO42– ions at infinite dilution are 59.46 and
79.8 ohm–1 cm2 mole–1 respectively. [ Sr = 87.6 , S = 32 , O = 16 ]
 ZENITH

Q.1 A lead storage cell is discharged which causes the H2SO4 electrolyte to change from a concentration of
34.6 % by weight (density 1.261g ml–1 at 25°C) to 27 % by weight. The original volume of electrolyte
is one litre. Calculate the total charge released at anode of the battery. Note that the water is produced
by the cell reaction as H2SO4 is used up. Over all reaction is
Pb(s) + PbO2(s) + 2H2SO4(l)  2PbSO4(s) + 2H2O(l)

Q.2 Assume that impure copper contains only iron, silver and a gold as impurities. After passage of 140 A, for
482.5s of the mass of the anode decreased by 22.260g and the cathode increased in mass by 22.011 g.
Estimate the % iron and % copper originally present.

Q.3 100ml CuSO4(aq) was electrolyzed using inert electrodes by passing 0.965 A till the pH of the resulting
solution was 1. The solution after electrolysis was neutralized, treated with excess KI and titrated with
0.04M Na2S2O3. Volume of Na2S2O3 required was 35 ml. Assuming no volume change during
electrolysis, calculate:
(a) duration of electrolysis if current efficiency is 80% (b) initial concentration (M) of CuSO4.

Q.4 Calculate the equilibrium constant for the reaction:

3Sn(s) + 2Cr2O72– + 28H+  3Sn4+ + 4Cr3+ + 14H2O
E0 for Sn/Sn2+ =0.136 V E0 for Sn2+/Sn4+ = – 0.154 V
E0 for Cr2O72–/Cr3+ = 1.33 V

Q.5 Calculate the equlibrium concentrations of all ions in an ideal solution prepared by mixing 25.00 mL of
0.100M Tl+ with 25.00mL of 0.200M Co3+.
E0 ( Tl+ /Tl3+ )= –1.25 V ; E0 (Co3+/Co2+ ) = 1.84 V

Q.6 Calculate the voltage, , of the cell at 250 C

Mn(s) | Mn(OH)2(s) | Mn2+(x M), OH –(1.00 x 10–4M) || Cu2+(0.675M) | Cu(s)
given that Ksp = 1.9 x 10–13 for Mn(OH)2(s) E0 (Mn2+/Mn) = –1.18 V

Q.7 Calculate the voltage, E, of the cell

Ag(s) | AgIO3(s) | Ag+(x M), HIO3 (0.300M) || Zn2+ (0.175M) | Zn(s)
if Ksp = 3.02 x 10–8 for AgIO3(s) and Ka = 0.162 for HIO3.
Q.8 The voltage of the cell
Pb(s) | PbSO4(s) | NaHSO4(0.600M) || Pb2+(2.50 x 10–5M) | Pb(s)
is E = +0.061 V. Calculate K2 = [H+] [SO42–] / [HSO4–], the dissociation constant for HSO 4 .
Given : Pb (s) + SO42–(aq) PbSO4 (s) + 2e– (E0 = 0.356) E0(Pb2+/Pb) = –0.126 V
Q.9 The voltage of the cell
Zn(s) | Zn(CN)42–(0.450M), CN–(2.65 x 10–3M) || Zn2+(3.84 x 10–4M) | Zn(s)
is E = +0.099 V. Calculate the constant Kf for Zn2+ + 4CN– Zn(CN)42–, the only
2+ –
Zn + CN complexation reaction of importance.
Q.10 An external current source giving a current of 5.0 A was joined with Daniel cell and removed after 10
hrs. Before passing the current the LHE and RHE contained 1L each of 1M Zn2+ and Cu2+ respectively.
Find the EMF supplied by the Daniel cell after removal of the external current source. E0 of Zn2+/Zn and
Cu2+/Cu at 25°C is 0.76 and +0.34V respectively.
Q.11 Determine at 298 for cell
Pt | Q, QH2, H+ || 1M KCl | Hg2Cl2(s) | Hg(l) | Pt
(a) it's emf when pH = 5.0
(b) the pH when Ecell = 0
(c) the positive electrode when pH = 7.5
given E0RP(RHS) = 0.28, E0RP(LHS) = 0.699

Q.12 At 25°C, Hf (H2O,l) = –56700 cal / mol and energy of ionization of H2O (l) = 19050 cal/mol. What
will be the reversible EMF at 25°C of the cell,
Pt | H2(g) (1 atm) | H+ || OH– | O2(g) (1 atm) | Pt, if at 26°C the emf increas by 0.001158 V.

Q.13 Calculate the cell potential of a cell having reaction: Ag2S + 2e– 2Ag + S2– in a solution buffered at
pH = 3 and which is also saturated with 0.1 M H2S.
For H2S : K1 = 10–8 and K2 = 1.1 × 10–13, Ksp(Ag2S) = 2 × 10–49, EAg  / Ag  0.8.

Q.14 Calculate the solubility and solubility product of Co2 [Fe(CN)6] in water at 250C from the following
data:
Conductivity of a saturated solution of Co2[Fe(CN)6] is 2.06 × 10–6 –1 cm–1 and that of water used
4.1 × 10–7–1 cm–1 . The ionic molar conductivities of Co2+ and Fe(CN)64– are 86.0 –1 cm2 mol–1
and 444.0  –1 cm–1mol–1.

Q.15 A sample of water from a large swimming pool has a resistance of 9200  at 25°C when placed in a
certain conductance cell. When filled with 0.02 M KCl solution, the cell has a resistance of 85  at
25°C. 500 gm of NaCl were dissolved in the pool, which was throughly stirred. A sample of this solution
gave a resistance of 7600 . Calculate the volume of water in the pool.
Given : Molar conductance of NaCl at that concentration is 126.5 –1 cm2 mol–1 and molar conductivity
of KCl at 0.02 M is 138  –1 cm2 mol–1.
 SUHANA SAFAR
OBJECTIVE
Q.1 A dilute aqueous solution at Na2SO4 is electrolysed using platinum electrodes.The products at the anode
and cathode are
(A) O2, H2 (B) S2O82 , Na (C) O2, Na (D) S 2 O 82 , H2

Q.2 The standard reduction potentials of Cu2+/ Cu and Cu2+ / Cu+ are 0.337 and 0.153 V respectively. The
standard electrode potential of Cu+ / Cu half cell is
(A) 0.184 V (B) 0.827 V (C) 0.521 V (D) 0.490 V

Q.3 A standard hydrogen electrons has zero electrode potential because

(A) hydrogen is easier to oxidise
(B) this electrode potential is assumed to be zero
(C) hydrogen atom has only one electron
(D) hydrogen is the lighest element.

Q.4 The standard reduction potential values of the three metallic cations X, Y, Z are 0.52, –3.03, and
–1.18 V respectively. The order of reducing power of the corresponding metals is
(A) Y > Z > X (B) X > Y > Z (C) Z > Y > X (D) Z > X > Y

Q.5 A gas X at 1 atm is bubbled through a solution containing a mixture of 1 M Y– and 1 M Z– at 25°C. If the
reduction potential of Z > Y > X, then
(A) Y will oxidise X and not Z (B) Y will oxidise Z and X
(C) Y will oxidise both X and Z (D) Y will reduce both X and Z.

Q.6 For the electrochemical cell, M | M+ || X– | X, E° (M+/M) = 0.44 V and E° (X/X–) = 0.33V. From this
data , one can deduce that
(A) M + X  M+ + X– is the spontaneous reaction
(B) M+ + X–  M + X is the spontaneous reaction
(C) Ecell= 0.77 V
(D) Ecell= –0.77 V

Q.7 The reaction,

3ClO–(aq)  ClO 3 (aq) + 2Cl–(aq)
is an example of
(A) Oxidation reaction (B) Reduction reaction
(C) Disproportionation reaction (D) Decomposition reaction

Q.8 The correct order of equivalent conductance at infinite dilution of LiCl, NaCl and KCl is
(A) LiCl > NaCl > KCl (B) KCl > NaCl > LiCl
(C) NaCl > KCl > LiCl (D) LiCl > KCl > NaCl
Q.9 Saturated solution of KNO3 is used to make salt bridge because
(A) velocity of K+ is greater than that of NO3

(B) velocity of NO3 is greater than that of K+

(C) velocities of both K+ and NO3 are nearly the same

(D) KNO3 is highly soluble in water

Q.10 Standard electrode potential data are useful for understanding the suitablilty of an oxidant in a redox
 titration. Some half cell reactions and their standard potentials are given below:
MnO 4 (aq) + 8H+(aq) + 5e– Mn2+ (aq) + 4H2O (l); E° = 1.51 V

Cr2O 72 (aq) + 14 H+ (aq) + 6e– 2Cr3+ (aq) +7H2O (l); E° = 1.38 V
Fe3+ (aq) + e– Fe2+ (aq); E° = 0.77 V
Cl2 (g) + 2e – 2Cl– (aq); E° = 1.40 V
Identify the only incorrect statement regarding quantitative estimation of aqueous Fe(NO3)2
(A) MnO 4 can be used in aqueous HCl

(B) Cr2O 72 can be used in aqueous HCl

(C) MnO 4 can be used in aqueous H2SO4

(D) Cr2O 72 can be used in aqueous H2SO4

Q.11 In the electrolytic cell, flow of electrons is from:
(A) Cathode to anode in solution (B) Cathode to anode through external supply
(C) Cathode to anode through internal supply (D) Anode to cathode through internal supply.

Q.12 Zn | Zn2+ (a = 0.1M) || Fe2+ (a = 0.01M)|Fe. The emf of the above cell is 0.2905 V. Equilibrium
constant for the cell reaction is
(A) 100.32/0.0591 (B) 100.32/0.0295
(C) 100.26/0.0295 (D) e0.32/0.295

Q.13 The half cell reactions for rusting of iron are:

1
2H+ + O + 2e– H2O; E0 = + 1.23 V, Fe2+ + 2e– Fe; E0 = –0.44 V
2 2
G0 (in kJ) for the reaction is:
(A) –76 (B) –322
(C) –122 (D) –176

SUBJECTIVE

Q.14 The standard reduction potential for Cu2+ / Cu is 0.34 V. Calculate the reduction potential at
pH = 14 for the above couple. Ksp of Cu(OH)2 is 1 x 1019.

Q.15 Electrolysis of a solution of MnSO4 in aqueous sulphuric acid is a method for the preparation of MnO2
as per the reaction, Mn2+aq + 2H2O  MnO2(s) + 2H+aq + H2(g)
Passing a current of 27A for 24 hours gives one kg of MnO2. What is the value of current efficiency?
Write the reaction taking place at the cathode and at the anode.

Q.16 How many grams of silver could be plated out on a serving tray by electrolysis of a solution containing
silver in +1 oxidation state for a period of 8.0 hours at a current of 8.46 Amperes? What is the area of
the tray if the thickness of the silver plating is 0.00254 cm? Density of silver is 10.5 g/cm3.

Q.17 Calculate the equilibrium constant for the reaction

Fe2+ +Ce4+ Fe3+ + Ce3+ , [given : E 0 Ce 4  / Ce3 1.44V;E 0 Fe3 / Fe2  0.68V ]
Q.18 Calculate the equilibrium constant for the reaction, 2Fe3+ + 3I 2Fe2+ + I3. The standard
 reduction potentials in acidic conditions are 0.77 and 0.54 V respectively for Fe3+ / Fe2+ and I3 / I
couples.
Q.19 Find the solubility product of a saturated solution of Ag2CrO4 in water at 298 K if the emf of the cell
Ag|Ag+ (satd.Ag2 CrO4 soln.) || Ag +(0.1 M) | Ag is 0.164 V at 298K.
Q.20 Copper sulphate solution (250 mL) was electrolysed using a platinum anode and a copper cathode. A
constant current of 2 mA was passed for 16 mintue. It was founf that agfter electrolysis, the absorbance
of the solution was reduced to 50% of its original value. Calculate the concentration of copper sulphate
in the solution to begin with.
Q.21 The following electrochemical cell has been set up
Pt(I) | Fe3+, Fe2+(a =1) || Ce4+ , Ce3+ (a = 1) | Pt(II)

E Fe 3 / Fe 2  = 0.77 V and E Ce 4  / Ce 3 = 1.61 V
If an ammetter is connected between the two platinum electrodes. predict the direction of flow of current.
Will the current increase or decrease with time?
Q.22 The standard potential of the following cell is 0.23 V at 150 C & 0.21 V at 350 C
Pt | H2(g) | HCl (aq) | AgCl(s) | Ag(s)
(i) Write the cell reaction.
(ii) Calculate H0 ,S0 for the cell reaction by assuming that these quantities remain unchanged in the range
150C to 350C.
(iii) Calculate the solubility of AgCl in water at 250C. Given standard reduction potential of the
Ag+/Ag couple is 0.80 V at 250C.
Q.23 Two students use same stock solution of ZnSO4 and a solution of CuSO4. The e.m.f of one cell is 0.3 V
higher than the other. The conc. of CuSO4 in the cell with higher e.m.f value is 0.5 M. Find out the conc.
 2.303 RT 
of CuSO4 in the other cell   0.06  .
 F 
Q.24 Find the equilibrium constant at 298 K for the reaction,
Cu2+(aq) + In2+(aq) Cu+(aq) + In3+(aq)
  
Given that E Cu 2  |Cu   0.15V , E In3 |In   0.42V , E In 2  |In   0.40V
0
Q.25(a)Calculate G f of the following reaction

Ag  (aq ) + Cl  (aq ) AgCl(s)

0 0 0
Given : G f (AgCl) = –109 kJ/mole, G f (Cl–) = –129 kJ/mole, G f (Ag+) = 77 kJ/mole
Represent the above reaction in form of a cell
Calculate E0 of the cell. Find log10KSP of AgCl

(b) 6.593 × 10–2 g of metallic Zn (amu = 65.39) was added to 100 ml of saturated solution of AgCl.
Calculate log10 Zn 
2
, given that

 
Ag  2

Ag+ + e– Ag E0 = 0.80 V

2+
Zn + 2e – Zn E0 = –0.76V
Also find how many moles of Ag will be formed?
 ANSWER KEY
EASY RIDE

Q.1 (a) 6.02 × 1022 electrons lost, (b) 1.89 × 1022 electrons gained, (c) (b) 1.80 × 1023 electrons gained
Q.2 (a) 0.75 F, (b) 0.69 F, (c)1.1 F Q.3 (i) 54 gm, (ii) 16.35 gm
Q.4 5
1.023 × 10 sec Q.5 1.12 mol, 12.535 litre
Q.6 0.112 litre Q.7 0.24 gms
Q.8 Rs. 0.75x Q.9 (i) 2.1554 gm ; (ii) 1336. 15 sec
Q.10 115800C, 347.4 kJ Q.11 t = 193 sec
Q.12 n=4 Q.13 Cu2+ = 0.08M, H+ = 0.04M, SO24 = 0.1M
Q.14 A = 114, Q = 5926.8C Q.15 Final weight = 9.6g,0.01Eq of acid
Q.16 (a) V(O2)=6.2L, (b)V(O2) = 0.824L Q.17 t = 93.65 sec.
Q.18 60 % Q.19 (i) 482.5 sec (ii) 0.3175 gm (iii) 0.327 gm
Q.20 7.958 ×10 M –5 Q.21 1.825 g Q.22 2M
Q.23 2+ 2+ + + 2+ –
(a) Zn | Zn | | Cd | Cd, (b) Pt, H2 | H | | Ag | Ag , (c) Cu | Cu | | Cl | Hg2Cl2 | Hg
(d) Pt | Fe2+, Fe3+ | | Cr2O 72 , Cr3+ | Pt

Q.24 (a) 2Ag + Cu2+  2Ag+ + Cu, (b) MnO 4 + 5Fe2+ + 8H+  Mn2+ + 5Fe3+ + 4H2O
(c) 2Cl– + 2Ag+  2Ag + Cl2, (d) H2 + Cd2+  Cd + 2H+
Q.25 Anode Cathode
(a) Zn | Zn2+ H+, H2 | Pt
(b) Pt | Sn2+, Sn4+ Fe3+, Fe2+ | Pt
(c) Pt | Fe2+, Fe3+ MnO 4 , Mn2+ | Pt
(d) Pb | Pb2+ Br2, Br– | Pt ]
Q.26 1.61 V Q.27 1.35 V Q.28 1.68 V
Q.29 – 0.80 V, NO Q.30 0.53 V, disproportionation Q.31 – 0.0367 V
o
Q.32 E Cell = – 0.36 V, not spontaneous Q.33 E =1.159V
Q.34 Kc = 2.18 x 10 26 107 0
Q.35 Kc = 1.864 x 10 , G = – 611.8 kJ
Q.36 Kw 10 14 Q.37 E0 = 0.7826 V Q.38 n = 2
Q.39 E = 0.059 Q.40 E = 0.395 V
Q.41 E0cell = +0.01V, Ecell = 0.0785V, correct representation is Pb|Pb2+ (103M)||Sn2+(1M)|Sn
Q.42 [Cu2+] = 2.97 x 1012M for E =0 Q.43 E = 0.81eV
Q.44 0.52 V, 0.61 V 3
Q.45 –1.30 ×10 kJ mol –1

Q.46 0.0295 V Q.47 0.0295 V Q.48 pH = 4

Q.49 (i) E = 0.286V; (ii) E = 0.0206V Q.50 a2 = 0.1006 M
Q.51 2
419 S cm equivalent –1
Q.52 0.00046 S cm1 ; 2174 ohm cm
5 6
Q.53 (i) 6.25 x 10 ohm, (ii) 1.6 x 10 amp Q.54 1.549 cm–1
Q.55 101.8 ohm-1 cm2 / gm equivalent
Q.56 0.0141 mho g equiv-–1 m2, 0.141 mho m-1
Q.57 (i) 232 Mho cm2 mol–1 , (ii) 116 Mho cm2 equivalent –1
Q.58 0.865 Q.59 1.33 ×10–4 gm/litre Q.60 8.74 × 10–11 mole2 /litre2
Q.61  = 0.435 , k = 6.7×10–4
Q.62 (i) 390.6 S cm2 mol–1 (ii) 12.32%
Q.63 1.067 S m–1
Q.64 523.2 ×10–4 mho cm2 mol–1
Q.65 (i) 6.98 (ii) 1.08 × 10–14
 PROFICIENCY TEST

1. T 2. F 3. T 4. T 5. T
6. F 7. T 8. T 9. F 10. T
11. T 12. T 13. T 14. F 15. F
16. 0.66 cm–1 17. O2 18. increase 19. higher
20. O2 & H2 21. Electrolyte concentration cell
22. Cr2O 72 (aq.), Cr3+(aq.), H+ | Pt
23. Pb(s) + PbO2(s) + 2H2SO4  2PbSO4(s) + 2H2O (l)
24. Hg2Cl2(s) + 2e–  2Hg(l) + 2Cl–(aq.)
  dE  
25. H = nF T   E  26. inert, i.e., should not interfere with net cell reaction
  dT  
27. Salt bridge or porous partition 28. zero 29. E°red cathode
30. Amount, rate 31. reduction process 32. Oxygen
33. 2.5 × 96500 C

MIDDLE GAME

Q.1 Eq. wt. = 107.3 Q.2 [Cu2+] = 104M Q.3 643.33amp,190.5g

Q.4 43.456g Q.5 265 Amp. hr. Q.6 E = 1.1 V
Q.7 E° = 1.59V, non-spontaneous Q.8 Ksp = 1.1 x 1016
Q.9 pH = 6.61 0
Q.10 E = 0.71V Q.11 [Br ] : [Cl] = 1 : 200
Q.12 0 
E = 0.1511V Q.13 Ka = 6.74 x 10 4 Q.14 [Zn2+]/[Cu2+] = 1.941 x 1037
Q.15 1.39V Q.16 0.373V Q.17 Kf = 8.227 x 1063
Q.18 –0.46 V Q.19 42.2 % Q.20 (ii). 1.2V, (iii) 245.1 kJ
Q.21 –0.037 V 
Q.22 A48.71 hour, [OH ] = 1.41 M
Q.23 V(O2) = 99.68 mL, V(H2) = 58.46 mL, Total vol. = 158.1 mL
Q.24 57.5894 gm Q.25 315.36 A Q.26 1.9 million year
Q.27 0.179 V
Q.28 0.372 V , – 0.307 V Q.29 h = 2.12 ×10–2, Kh= 1.43 × 10–5 M
Q.30 0.95 V Q.31 1.536 ×10–5 M3 Q.32 0.1456 ampere
–2
Q.33 3.29 cm Q.34 4.25×10 metre Q.35 0.1934 gm/litre

ZENITH

Q.1 1.21 × 105 coulomb Q.2 Cu = 98.88%, Fe = 0.85% Q.3 1250 s, 0.064 M
Q.4 K = 10268 Q.5 Tl+ = 10–8; Co3+ = 2 × 10–8 Q.6 1.66V
Q.7 –1.188V Q.8 10–2 Q.9 5.24 x 1016
Q.10 1.143V Q.11 (a) –0.124 V, (b) 7.1, (c) calomel electrode
Q.12 0.4414 V Q.13 – 0.167 V Q.14 KSP = 7.682 × 10–17
Q.15 2 × 105 dm3
 SUHANA SAFAR

OBJECTIVE

Q.1 A Q.2 C Q.3 B Q.4 A Q.5 A

Q.6 B Q.7 C Q.8 B Q.9 C Q.10 A

Q.11 C Q12 B Q.13 B

SUBJECTIVE

Q.14 E0 = 0.22N

Q.15  = 94.8%; Cathode : 2H+ + 2e  H2, Anode : Mn2+  Mn4+ + 2e

Q.16 WAg = 272.2g, area =1.02 x104 cm2 Q.17 Kc = 7.6 x 1012 Q.18 KC = 6.26 x 107

Q.19 Ksp = 2.287 x 1012 M3 Q.20 7.95 × 10–5M

Q.21 decrease with time

Q.22 H0 = –49987 Jmol–1 , S0 = –96.5 J mol–1 K–1 , s = 1.47x10–5 M

Q.23 0.05 M Q.24 KC = 1010

Q.25 (a) E0 = 0.59 V, log10KSP = –10
(b) 52.8, 10–6 moles