UNIT:3:- E L E C T R O C H E M I S T R Y
CONDUCTIVITY
1# With the help of graph , explain why is it not easy to determine Λ m0 for a weak electrolyte by extrapolating the conc.-Molar
conductance curve as for strong electrolyte
OR,
How does Molar conductance varies with concentration in case of strong and weak electrolyte? Explain with the graph .
OR,
How would you explain the sharp increase in the molar conductivity of a weak electrolyte on dilution ? (2)
2# State and explain Kohlrausch’s law .Write the equation of Λm0 of CaCl2 and give its three applications . (2)
3# Define the following terms with their units (i) Cell constant (ii) Conductivity (iii) Molar conductance. ( 1+ 1 + 1)
4# (a) What is the relationship between conductivity and molar conductivity ? (1) ( AI-2008)
(b) How conductivity and molar conductivity for the solution of an electrolyte vary with concentration. (1)
(c) Why does the conductivity of a solution decrease with dilution? (1)
(d) Suggest a way to determine the Λm value of water. (1)
(e) State the factors that determine the magnitude of conductivity of an electrolyte (1)
(f) What is the effect of temperature on molar conductance ? (1)
(g) How is the molar conductivity related to the degree of dissociation of a sparingly soluble salt ? (1)
CORROSION:
5#What is meant by ―corrosion of metals‖ ? Give its mechanism ? Describe three methods for preventions of corrosion ? (3)
6# Explain how rusting of iron is envisaged as setting up of an electrochemical cell. (2)
7# What is the role of H+ during the rusting of iron . Also mention its source. (1)
8 # Write the one application of bisphenol. (1)
9 # Rusting of iron is quicker in saline water than in ordinary water .Expalin ? (1) (2006-AI)
10# Why does an alkaline medium inhibit the rusting of iron ? (1) (2006-D)
11# What is meant by cathodic protection of Fe ? Name two metals used for it (2M)
12#What is sacrificial protection from rusting? Which metal is generally used for this purpose. (2)
13 # How is cathodic protection of iron is different from galvanization . (1)
14 # ‖Iron does not rust when coating is broken in a galvanized iron pipe but rusting occurs if coating of tin over iron is broken ―Explain (1)
15# CO2 is always present in natural water . Explain its effects (increases , stops or no effect) on rusting of iron ? (1)
16# Discuss electrical protection for preventing rusting of iron pipes in underground water . (1)
CELLS AND BATTERIES:
17# Write the chemistry of recharging the lead storage battery, highlighting all the materials that are involved during recharging.(2)
18# Suggest two materials other than hydrogen that can be used as fuels in fuel cells. (1)
19# Write the name of the electrodes and electrolyte used in (i) fuel cell (ii) mercury cell (iii) in dry cell (iv) lead-storage cell
20#What is a fuel cell ? What is the basis of obtaining electrical energy in fuel cells? Give the electrode reactions H 2-O2 fuel cell .
Write two advantages of fuel cell ? How do they resemble and differ from galvanic cells?
(3M)
21# Why does a mercury cell gives a constant voltage throughout its life ?
(1M)
22# Why does dry cell become dead after a long time , even if it has not been used ?
(1M)
23# Write a cell reactions which occur in lead storage battery (i) when the battery is in use and (ii) when the battery is on charging . (1)
24# Describe a Leclanche cell with special reference to (i) the electrodes used and (ii) the reactions occurring at the electrodes in the cell. (2)
25# Describe the composition of anode and cathode in a Mercury Cell . Write the electrode reaction for this cell. (2)
26# Distinguish between primary cells and secondary cells with suitable examples ? (2)
GALVANIC CELL:
27 # Mention two important functions of salt bridge ? (1)
28# How would you determine the standard electrode potential of the system Mg 2+|Mg ?
29# Can you store copper sulphate solutions in a zinc pot?
30# Consult the table of standard electrode potentials and suggest three substances that can oxidize ferrous ions under suitable conditions.
31# Arrange the following metals in the order in which they displace each other from the solution of their salts. Al, Cu, Fe, Mg and Zn.
32# Given the standard electrode potentials,
K+/K = –2.93V, Ag+/Ag = 0.80V, Hg2+/Hg = 0.79V Mg2+/Mg = –2.37 V , Cr3+/Cr = – 0.74V
Arrange these metals in their increasing order of reducing power.
33# Using the standard electrode potentials, predict if the reaction between the following is feasible:
(i) Fe3+(aq) and I–(aq) (ii) Ag+ (aq) and Cu(s) (iii) Fe3+ (aq) and Br– (aq) (iv) Ag(s) and Fe 3+ (aq (v) Br2 (aq) and Fe2+ (aq).
[ GIVEN : E0 for Ag+/Ag = 0.80V ;
Fe+2 / Fe = -0.44V ;
I2 / I ˉ = 0.54 V ;
Br2 /Brˉ= 1.09 V ;
+3
+2
+2
Fe / Fe = 0.77 V;
Cu /Cu =0.34V ]
34# How does electrochemical series help in predicting whether a metal can liberate hydrogen from acid or not?
35# ―E‖ is an intensive parameter but ∆ rG is an extensive property and the value depends on ―n‖.—Explain?
36#.Depict the galvanic cell in which the reaction Zn(s) + 2Ag + (aq)  Zn+2 (aq) + 2Ag(s) takes place , further show :
( AI-2008)
(i) which of the electrode is negatively charged, (ii) The carriers of the current in the cell, and (iii) individual reactions at each electrode.
37 # Write the relation between cell potential and equilibrium constant
38# What is the relationship between the standard EMF of the cell and the equilibrium constant of the cell reaction at 298K ?
39# Describe the working and construction of Standard Hydrogen Electrode(SHE) with a labeled diagram.
40# What is the effect of an increase in concentration of Zn ions on the electrode potential of Zn-electrode for which E0 = -0.76 V .
( Page-2)
41# Why is not possible to measure the absolute value of the electrode potential –Explain?
42# Under what condition will a galvanic cell send no current into outer circuit ?
43#.On the basis of the E0 values stated for acid solution, predict whether Ti+4 species may be used to oxidize Fe++ to Fe+3
(a) Ti+4 + e  Ti+3 : E0 = 0.01 V
(b) Fe+3 +e  Fe++ : E0 = +0.77 V
(AI-2007)
44#How can we measure the single electrode potential ?Explain with one example ?
45#Write four applications of Electrochemical cell
46# What does the negative value of E0cell indicate ?
47# What is CONCENTRATION CELL?Describe the working of concentration cell with suitable example Give the expression of the cell
ELECTROLYSIS:48 # State and explain Faraday’s law of electrolysis .What is electrochemical equivalent ? (3)
What is the relation between chemical equivalent and electrochemical equivalent ?
49# Predict the products of electrolysis in each of the following : (1 mark each)
(a) An aqueous solution of NaCl with platinum electrodes.
(b) Molten NaCl with platinum electrodes.
(c) A dilute solution of H2SO4 with platinum electrodes
(d) A conc.solution of H2SO4 with platinum electrodes
(e) An aqueous solution of AgNO3 with silver electrodes.
(f) An aqueous solution of AgNO3 with platinum electrodes.
(f) An aqueous solution of CuCl2 with platinum electrodes. (g) An aqueous solution of CuCl2 with copper electrodes. .
(i) Molten PbBr2 with platinum electrodes
50# Why is that aluminum metal can’t be obtained by electrolysis of an aqueous solution of a salt of aluminum ? (1)
51#.Explain why electrolysis of aq.solution of NaCl gives H 2 at cathode and Cl2 at anode .Write overall reaction (2)
( E0Na+/Na = -2.71 V ; E0(H2O/H2)= -0.83 V ; E0 Cl2 /2Cl- = +1.36 V ; E0H2O/O2 = + 1.23 V )
52# (a) Explain --Cu does not dissolve in HCl but in HNO3 while both the acid contains H+ . OR , Can H+ oxidize Cu ? (1)
(c) Explain-- Can H2 gas reduce Cu++ ? (1)
(d) Explain – fluorine gas is the strongest oxidizing agent and Fluoride ion is the weakest reducing agent? (1)
(e) Suggest a list of metals that are extracted electrolytically. (1)
(f) Name the anode and cathode during electrorefining of copper. (1)
(g) Why Na,Mg, Al are normally extracted through electrometallurgy.What are the starting materials taken while extracting these metals
53# (a) What is Overvoltage ? (1)
(b) What is the function of platinised platinum in SHE (1)
(c) Write Debye- Huckel –Onsager equation . what do different symbol signify ? In which type of electrolyte it is valid , strong or weak .
ADITIONAL QUESTIONS FOR PRACTICE ********
1# Reduction potential of Cu and Zn are +0.34 and -0.76 V respectively.Which of them is a stronger reducing agent ?
2# Can Ni spatula be used to stir a solution of CuSO4 ?Support your answer with reason ?
(Reduction potential of Ni and Cu are -0.25 and +0.34 V respectively)
3# The compound AgF2 is unstable.However,if formed, the compounds acts as a very strong oxidizing agent .Why?
4# How does the molar conductivity of KCl solution vary with increasing concentration ?
5# What is the sign of ∆G for (i) electrolytic cell and (ii) electrochemical cell
6# What is the charge on one mole of electrons.
7# Differentiate inert and reactive electrode with suitable examples.
8# Mention the two problems faced during the measurement of resistance of an ionic solutions and how these difficulties were resolved
9# What is the conductivity of pure water ?
10# Write the symbolic notations for SHE and its potentials . ( both anodic and cathodic half cell reactions)
11# Under what condition will a galvanic cell send no current into outer circuit ?
12# Write the expression to relate the molar conductivity of electrolyte in terms of degree of dissociation.
13# Write the equation showing the relationship between std.free energy and std.cell potential .
14# Write the representation of a Daniel cell ?
15# How cell potential and emf differ?
16# Differentiate between (a) electrolytic and galvanic cell (b) electrolytic and metallic conduction
17# What percentage of sulphuric acid is used in the lead storage battery .
18# What is the EMF of the cell when the cell reaction attains equilibrium ?
19# What flows in the internal circuit of a Galvanic cell ?
20# Out of Zn and Sn which one protects iron better even after cracks and why ?
21# How can you increase the reduction potential of an electrode ?
22# Why a cell stops working after some time ?
23# Which type of metal can be used in cathodic protection of Iron against rusting ?
NUMERICAL QUESTIONS --- E L E C T R O C H E M I S T R Y
G a l v a n I c C e l l :Q.1# Represent the cell in which the following reaction takes place
Mg(s) + 2Ag+(0.0001M) → Mg2+(0.130M) + 2Ag(s)
Calculate E0 cell , E cell , ΔrG0 and equilibrium constant of the reactions. Given : E0 Mg++/Mg = ─2.36 V and E0 Ag+/Ag = 0.80 V
Q.2# Calculate the standard cell potentials of galvanic cell in which the following reactions take place:
(i) 2Cr(s) + 3Cd2+(0.001M) 2Cr3+(0.00001M) + 3Cd . Calculate the ECell
(ii) Fe2+(aq) + Ag +(aq) Fe3+(aq) + Ag(s)
Calculate the ΔrG0 and equilibrium constant of the reactions. Given : E0 Cd++/Cd = 0.40 V and E0 Cr+++/Cr = ─ 0.74 V
Q.3# Calculate the potential of hydrogen electrode in contact with a solution whose pH is 10.
Q.4 # Write the Nernst equation and emf of the following cells at 298 K:
(i) Mg(s)|Mg2+(0.001M)||Cu2+(0.0001 M)|Cu(s)
(ii) Fe(s)|Fe2+(0.001M)||H+(1M)|H2(g)(1bar)| Pt(s)
(iii) Sn(s)|Sn2+(0.050 M)||H+(0.020 M)|H2(g) (1 bar)|Pt(s)
(iv) Pt(s)|Br2(l)|Br–(0.010 M)||H+(0.030 M)| H2(g) (1 bar)|Pt(s).
Q.5 # In the button cells widely used in watches and other devices the following reaction takes place:
Zn(s) + Ag2O(s) + H2O(l) Zn2+(aq) + 2Ag(s) + 2OH–(aq) ,
( 2005-D)
Determine ΔrG0 and E0 for the reaction. E0 Ag2O/Ag = 0.344 V and E0 Zn++/Zn = ─ 0.76V
Q.6#Calculate the equilibrium constant of the reaction:
Cu(s) + 2Ag+(aq) →Cu2+(aq) + 2Ag(s) E0cell = 0.46 V
Q.7# The standard electrode potential for Daniell cell is 1.1V.
Calculate the standard Gibbs energy for the reaction: Zn(s) + Cu2+(aq) →Zn2+(aq) + Cu(s)
Q.8# Calculate the emf of the cell in which the following reaction takes place
Ni(s) + 2Ag+ (0.002 M) →Ni2+ (0.160 M) + 2Ag(s) , Given that E0(cell) = 1.05 V
Q.9 # The cell in which the following reaction occurs: 2Fe3+(aq) + 2Iˉ( aq )  2Fe2+ (aq ) + I2(s)has
E0 cell = 0.236 V at 298 K. Calculate the standard Gibbs energy and the equilibrium constant of the cell reaction.
Q.10## A cell consists of two hydrogen electrodes . The negative electrode is in contact with a solution having 10-6 M H+
ion concentration . Calculate the concentration of H+ ions at the positive electrode , if the EMF of the cell is found to be
0.118 V at 298K . ( Ans =10-4M )
Q.11# # At what PH of HCl solution will hydrogen gas electrode show electrode potential of – 0.118 V ? H2 gas is bubbled at
298 K and 1 atm pressure . ( Ans  PH = 2)
Q.12 # Calculate the E0 Cu+/Cu , when E0 Cu++/Cu = 0.337 V and E0 Cu++/Cu+ = 0.153 V .
E L E C T R O L Y S I S :Q.13# How much charge is required for the following reductions:
(i) 1 mol of Al3+ to Al. (ii) 1 mol of Cu2+ to Cu. (iii) 1 mol of MnO4 – to Mn2+ (iv) 1 mol of Cr2O7 2–
Q.14# How much electricity in terms of Faraday is required to produce
(i) 20.0 g of Ca from molten CaCl2.
(ii) 40.0 g of Al from molten Al2O3.
Q.15# How much electricity is required in coulomb for the oxidation of
(i) 1 mol of H2O to O2.
(ii) 1 mol of FeO to Fe2O3.
Q.16 # Consider the reaction: Cr2O7 2– + 14H+ + 6e– 2Cr3+ + 8H2O
What is the quantity of electricity in coulombs needed to reduce 1 mol of Cr2O72–?
Q.17# If a current of 0.5 ampere flows through a metallic wire for 2 hours, then how many electrons would flow through
the wire?
Q.18 # A solution of Ni(NO3)2 is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes.
What mass of Ni is deposited at the cathode?
Q.19 # A solution of CuSO4 is electrolysed for 10 minutes with a current of 1.5 amperes. What is the mass of copper
deposited at the cathode?
Q.20 # Three electrolytic cells A,B,C containing solutions of ZnSO4, AgNO3 and CuSO4, respectively are connected in
series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B.
How long did the current flow? What mass of copper and zinc were deposited? (AI-2008) (5M)
Q.21 # Silver is electrodeposited on a metallic vessel of surface area 800 cm2 by passing current of 0.2 amp for 3 hrs.
Calculate the thickness of silver deposited .(density of Ag = 10.47 g/cm3 . At. Mass of Ag = 107.92 amu )
Q.22 # In an industrial plant ,aluminium is produced by the electrolysis of alumina dissolved in cryolite.This takes a
current of 20000 amp.If the current efficiency is 90 % ,how much Al will be produced per day . (Ans=145.1 kg)
ConductIvIty
Q.23 # The resistance of a conductivity cell containing 0.001M KCl solution at 298 K is 1500 Ώ. What is the cell constant if
conductivity of 0.001M KCl solution at 298 K is 0.146 × 10–3 S cm–1. (AI-2008)
Q.24# The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 S cm–1. Calculate its molar conductivity.
Q.25# Conductivity of 0.00241 M acetic acid is 7.896 × 10–5 S cm–1. Calculate its molar conductivity and if Λmfor acetic
acid is 390.5 S cm2 mol–1, what is its dissociation constant? (AI-2008) (5M)
Q.26# Resistance of a conductivity cell filled with 0.1 mol L–1 KCl solution is 100 Ω. If the resistance of the same cell when
filled with 0.02 mol L–1 KCl solution is 520 Ω , calculate the conductivity and molar conductivity of 0.02 mol L–1 KCl
solution. The conductivity of 0.1 mol L–1 KCl solution is
1.29 S/m. ( AI-2007) ( 3M)
Q.27# The electrical resistance of a column of 0.05 mol L–1 NaOH solution of diameter 1 cm and length 50 cm is 5.55 × 103
ohm. Calculate its resistivity, conductivity and molar conductivity.
Q.28# Λm for NaCl, HCl and NaAc are 126.4, 425.9 and 91.0 S cm2 mol–1 respectively. Calculate Λm for HAc.
Q.29# The conductivity of 0.001028 mol L–1 acetic acid is 4.95 × 10–5 S cm–1. Calculate its dissociation constant if Λmfor
acetic acid is 390.5 S cm2 mol–1.
Q.30# The molar conductivity of 0.025 mol L–1 methanoic acid is 46.1 S cm2 mol–1.
Calculate its degree of dissociation and dissociation constant.
[ Given λ0(H+) = 349.6 S cm2 mol–1 and λ0 (HCOO–) = 54.6 S cm2 mol–1 ]
Extraaaaaaaaaaaaaaaaaa
Q.31# The molar conductivity of KCl solutions at different concentrations at 298 K are given below:
c/mol L–1
Λm/S cm2 mol–1
0.000198
148.61
0.000309
148.29
0.000521
147.81
0.000989
147.09
Show that a plot between Λm and c1/2 is a straight line. Determine the values of Λmand A for KCl.
Q. 32# The conductivity of sodium chloride at 298 K has been determined at different concentrations and the results are
given below:
Concentration/M
0.001
0.010 0.020
0.050 0.100
2
–1
10 × k/S m
1.237
11.85
23.15
55.53 106.74
Calculate Λm for all concentrations and draw a plot between Λm and √ c. Find the value of Λm
------------------------------------------------------------------------------------ ----------------------------------------------------------------------------------------
READY RECONER:-*Preferential Discharge Theory:If more than one type of ion is attracted towards a particular electrode, then ion discharged is one which requires least energy.
The decreasing order of discharge potential OR , the increasing order of deposition of cations and anions as follows:
CATIONS:- K+, Na+, Ca++, Mg++, Al+3, Zn+2, H+, Cu+2, Ag+,Au+3 : ANIONS:- SO42¯ ,NO3¯, OH¯(H2O),Cl¯, Br¯, I¯
*Faraday’s Law: (1) The amount of chemical reaction which occurs at any electrode during electrolysis by a current is proportional to
the quantity of electricity passed through the electrolyte (solution or melt).
m = z Q = z I t : m mass of subs deposited, Q quantity of electricity passed (in coulombs) , I is current in ampere
t is time in second . : Z= m / Q = the mass of the substance of deposited when 1 amp current is passed for 1 sec i.e. 1 coulomb.
ECE:- The mass of the substance deposited when current of one ampere is passed for one second.
( One Faraday deposits one-gm-equivalent of the substance) ECE=Z=electrochemical equivalent = Eq.Wt. of substance / 96500
(2) The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are
proportional to their chemical equivalent weights .
CHEMICAL EQUIVALENT(E) = Atomic Mass of Metal ÷ Number of electrons required to reduce the cation
m1 / m2 = Z1It / Z2It = E1 /E2 :
E1 /E2 = Z1 /Z2 (when same charge is passed through two cells)
W1 of Cu deposited / W1 of Ag deposited = Eq. mass of Cu / Eq. mass of Ag
OVERPOTENTIAL:- The difference between the potential of the electrode when gas evolution is actually observed and the theoretical (i.e
reversible) value for the same evolution
Moreover, some of the electrochemical processes although feasible, are so slow kinetically that at lower voltages these
don’t seem to take place and extra potential (called overpotential) has to be applied, which makes such process more difficult to occur. is
called overvoltage .
 Transference of electrons form water is kinetically slow process as it requires relatively large activation energy.The slowness of
electrode reaction creates some electrical resistance at the electrode surface .Hence , extra potential or voltage is required to provide
activation energy and to overcome resistance in case of oxidation of water of molecules .This extra voltage required for oxidation of
water is called overvoltage. Due to overvoltage ,the oxidation of Cl¯ ions occurs in preference to H 2O.
ELECTRODE POTENTIAL:
Absolute value of electrode potential of an electrode cannot be determined a half cell by itselfcannot cause movement of charge(flow of
electrons) .It is due to the fact that once equilibrium is reached between the electrode and the solution in a half cell ,no further displacement of
charges can occur unless and until it is connected to another half cell with different electrode potential .This difficulty is overcome by finding
the electrode potential of various electrode relative to a reference electrode whose electrode potential is arbitrarily fixed. Standard Hydrogen
Electrode(SHE) as reference electrode whose electrode potential is arbitrarily taken to be zero at all temp.
EFFECT OF DILUTION ON CONDUCTIVITY:
Increase in volume is more than increase in total no. of ions .Actually,total no. of ions / cm3 of volume is decreased and hence
net value of conductivity decreases.
(2) 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.
(3) Weak electrolytes like acetic acid have lower degree of dissociation at higher concentrations and hence for such electrolytes,
the change in m with dilution is due to increase in the degree of dissociation and consequently the number of ions in total volume
of solution that contains 1 mol of electrolyte. In such cases m increases steeply on dilution, especially near lower concentrations.
Therefore, m cannot be obtained by extrapolation of m to zero concentration. At infinite dilution (i.e., concentration c zero)
electrolyte dissociates completely (=1),but at such low concentration the conductivity of the solution is so low that it cannot be
measured accurately. Therefore, mfor weak electrolytes is obtained by using Kohlrausch law of independent migration of ions
(1)