CE – 1 ELECTROCHEMISTRYeinsteinclasses.com/Electrochemistry.pdf · It is defined as the conducting power of all the ions present in one mol of an ... thickness of 10–2 cm

CE – 1

Einstein Classes, Unit No. 102, 103, Vardhman Ring Road Plaza, Vikas Puri Extn., Outer Ring Road

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ELECTROCHEMISTRY

C1 Conductance :

The resistance of any conductor varies directly as its length (l) and inversely as its cross-sectional area (a).

aR

l

aR

l .

where is a constant depending upon the nature of the material and is called specific resistance orresistivity of the material.

Specific resistance is the resistance of one centrimetre cube of a material.

The reciprocal of the specific resistance,

1 is called specific conductance or conductivity..

Conductivity is the conductance of one cm cube of a material.

taking K1

(specific conductance), C

R

1 (conductance) K = C ×

a

l....

a

l is called cell constant, l = distance between two electrodes , a = cross-sectional area of electrodes

C2A Equivalent Conductance :

It is defined as the conducting power of all the ions present in one gram equivalent of an electrolyte in agiven solution. At concentration N (in gm equivalent L–1) equivalent conductance is denoted by

eq

(equivalent conductance at concentration c) c =

N

1000k ; units: ohm–1 cm2, eq–1

Equivalent conductance increases with dilution. When the solution is infinitely diluted the equivalentconductance is denoted as

. The

can be determined by extrapolation method, in which graph between

c and C is extended to zero concentration.

C2B Molar Conductance :

It is defined as the conducting power of all the ions present in one mol of an electrolyte in a given solution.

M

1000km

unit : ohm–1 cm2 mol–1

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eq

and m both increases with decrease in concentration.

C2C Kohlrausch’s Law :

At infinite dilution, when ionisation is complete, each ion makes its contribution towards equivalentconductance of the electrolyte and at infinite dilution equivalent conductance is given by the sum of theequivalent conductances of the contributing ions.

Thus for AxB

y,

m = x

a + y

b

where, a is the molar conductance of A(say cation) and

b that of B(say anion) at infinite dilution.

(degree of dissociation of weak electrolyte)

m

cm

Practice Problems :

1. Molar conductance at infinite dilution of BaCl2, H

2SO

4 and HCl aq. solution are x

1, x

2 and x

3

respectively. Molar conductance of BaSO4 solution is :

(a) x1 + x

2 – x

3(b) x

1 – x

2 – x

3(c) x

1 + x

2 – 2x

3(d) x

1 – 2x

2 + x

3

[Answers : (1) c]

C3A Electrical Conductors :

All substances which can allow the flow of electricity are known as electrical conductors. Mainly we havetwo types of electrical conductors which are given below :

1. Electronic Conductor : They are those conductors in which the flow of electricity is due to the movementof loosely bonded electrons in their own standard state. In this case, the movement of matter does not takeplace during the flow of electricity. For example :

(a) all metals in their elemental state (b) graphite and (c) alloys

2. Electrolytical Conductors :

In this case of electrolytic conductors, the flow of electricity is due to the movement of ions i.e., here actualtransport of matter takes place.

C3B Electrolytes :

Electrolytes may be pure substances (e.g., salts, acids or bases) in their fused states or more commonly theyare aqueous solutions of these compounds or they are sometimes pure liquid.

There are two types of electrolytes :

(a) Strong Electrolyte : The compounds which are 100% ionised at any dilution, are treated asstrong electrolytes, for e.g., HClO

4, HI, HBr, HCl, H

2SO

4, HNO

3, NaOH, KOH, NaCl etc.

(b) Weak Electrolyte : The compounds which are less or feebly ionised (at lower dilution) aretreated as weak electrolytes.

All weak acids, weak bases or the salts having less ionic character are treated as weak

electrolytes e.g. CH3COOH, , H

2CO

3, Mg(OH)

2, Zn(OH)

2 etc.

C4A Electrolysis :

Electrolysis is a process which involves a chemical change at the electrodes when electricity is passesthrough an electrolyte.

“Electrolysis is a process which involves a chemical change at the electrodes when electricity is passedthrough an electrolyte”.

It was found experimentally that the positive ions from the electrolyte are attracted on the negativeelectrode (cathode) and get reduced. Similarly, the negative ions are attracted on the positive electrode(anode) and are oxidised.

Hence due to electrolysis, the species under consideration gets decomposed at cathode and anode.

Faraday has established a relationship between the amount of electricity passes through the electrolyte andthe amount of chemical change occurring at an electrode. This relationship is known as Faraday’s law ofelectrolysis.

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C4B Faraday’s Laws of Electrolysis :

First Law of Electrolysis :

The amount of chemical change produced is proportional to the quantity of electric charge passing throughan electrolysis cell.

Suppose after passing Q coulombs of electricity (amount of electricity), W amount of a substance hasappeared (or disappeared).

Thus W Q

W = ZQ (Where Z is a constant, which is known as electrochemical equivalent)

W = Z i t..... (Q = i t, where i is current passed for t seconds)

* Also Q = nF..... (F is a faraday constant i.e., it is charged carried by one mole of electron,where n is no.of mole of electrons transfer takes place).

Second Law of Electrolysis :

When same amount of amount of electricity is passed through different electrolytes, the weight of thesubstance appeared or disappeared by any electrode is always directly proportional to the equivalent weightof the substance appeared or disappeated.

According to the second law of electrolysis W E, in general, 2

2

1

1

E

W

E

W

C4C Relation between Electrochemical Equivalent and Equivalent Weight :

By the use of above two laws E Z

Thus, E = FZ.... (where F is Faraday’s constant, E is equivalentweight and Z is electrochemical equivalent).

If we combine first and second law of electrolysis then following expressions takes place :

F

It

E

Wfactor.nmolor

F

Q

E

Wneq , (n

eq)

oxidised = (n

eq)

reduced

Practice Problems :

1. The density of Cu is 8.94 g cm–3. The quantity of electricity needed to plate an area 10 cm 10 cm to athickness of 10–2 cm using CuSO

4 solution is

(a) 13586 C (b) 27172 C (c) 40758 C (d) 20348 C

2. The same amount of electricity was passed through two cells containing molten Al2O

3 and molten

NaCl. If 1.8 g of Al were liberated in one cell, the amount of Na liberated in the other cell is

(a) 4.6 g (b) 2.3 g (c) 6.4 g (d) 3.2 g

3. Silver is removed electrolytically from 200 mL of a 0.1 N solution of AgNO3 by a current of 0.1

ampere. How long will it take to remove half of the silver from the solution (At. wt. of Ag = 108 g)

(a) 10 sec (b) 16 sec (c) 100 sec (d) 9650 sec

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4. Consider the following electrolysis

(1) CuSO4

(2) Fe2(SO

4)

3(3) AlCl

3(4) AgNO

3

The quantity of electricity needed to electrolyse completely 1 M solutions of these electrolytes will be

(a) 2F, 6F, 3F and 1F respectively (b) 6F, 2F, 3F and 1F respectively

(c) 2F, 6F, 1F and 3F respectively (d) 6F, 2F, 1F and 3F respectively

5. Assume that during electrolysis of AgNO3, only H

2O is electrolysed and O

2 is formed as :

2H2O 4H+ + O

2 + 4e—

O2 formed at N.T.P. due to passage of 2 amperes of current for 965 s :

(a) 0.112 L (b) 0.224 L (c) 11.2 L (d) 22.4 L

[Answers : (1) b (2) a (3) d (4) a (5) a]

C5A Electrochemical Cell :

A devide that converts chemical energy into electrical energy. The cell is based on the principle of indirectredox reactions, i.e, the oxidation and reduction reactions takes place in different container.

The electrochemical cells (galvanic or voltaic) consists of two half cells connected with each other bymeans of an electric wire to allow an indirect redox reaction. Although the solutions of two half cells are indifferent containers, but for the continuous flow of electricity, transfer of ions from one solution to anothersolution are always allowed. Oxidation at anode and is a –ve electrode, reduction at cathode and is a +veelectrode. For e.g., Zn-Cu galvanic cell is represented as follows :

C5B Electrode Potential :

It is the potential difference established between the electrode and its electrolyte. It is of two types i.e.,Reduction Potential and Oxidation Potential.

Reduction Potential is tandency of an electrode to receive electrons for the deposition of the solvatedpositive ions from its own solution. Whereas oxidation potential is the tandency of an electron to getoxidised in its own solution.

C5C E.M.F. (Electromotive Force) or Cell Potential of a Cell :

The difference between the reduction potentials of two half cells constituting the Galvanic cell is known asits cell potential (or e.m.f.) i.e.,

Ecell

(e.m.f.) = (Reduction Potential)Cathode – (Reduction Potential)Anode

or, ECell

= ECathode

– EAnode

For the spontaneous flow of electricity from the Galvanic cell, the e.m.f. of the cell must be positive.

C5D Concentration Effect in Voltaic Cell – Nernst Equation

Nerst equation gives a quantitative relationship between the concentration of ions and electrode potential.For a general electrode reaction : Mn+ + L(s) Ln+ + M(s).

Emf(E) = E0 – 2.303 ]M[

]L[log

nF

RTn

n

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Practice Problems :

1. By how much will the potential of half-cell Cu2+ | Cu change if the solution is diluted to 100 times at298 K ?

(a) Increase by 59 mV (b) Decreases by 59 mV

(c) Increase by 29.5 mV (d) Decreases by 29.5 mV

2. Which of the following changes will increase the EMF of the cell :

Co(s) | CoCl2(M

1) || HCl (M

2) | Pt (H

2, g)

(a) increase in the volume of CoCl2 solution from 100 mL to 200 mL

(b) increase M2 from 0.01 M to 0.50 M

(c) increase the pressure of the H2(g) from 1.00 to 2.00 atm

(d) increase M1 from 0.01 M to 0.50 M

3. The measured voltage of the following cell is 0.9 V at 250C. Pt, H2(1 atm) | H+(aq) || Ag+ (1.0 M) | Ag

If V80.0E0

Ag/Ag . The pH of the aqueous solution of H+ ion is

(a) 1.69 (b) 2.50 (c) 5.20 (d) 9.69

[Answers : (1) b (2) b (3) a]

C6 Application of Nernst Equation

1. Electrical work :

G = –nFE (in a given state), G0 = –nFE0 (in a standard state), G0 = –RT In K aq

G10 + G

20 = G

30 (when different no. of element are involved, if equal no of electrons are involved

E10 + E

20 = E

30.

2. At equilibrium E = 0 and Q = K (equilibrium constant)

Klogn

059.0Kln

nF

RTE 10

0

3. For standard hydrogen electrods (SHE)

E0(SHE)

= 0.00 V.

SHE Pt | H2(g)

(1 atm) | HCl aq (1M)

SHE, colomel-electrode and silver, silver-chloride electrodes are used as reference half cells.

HH/H

p059.0E2

4. Concentration cells

Zn | Zn2+ (C1) || Zn2+ (C

2) | Zn

1

2cell

C

Clog

n

059.0E ........[C

1 > C

2]

Pt (H2) (P

1) | HCl | Pt (H

2) (P

2)

Ecell

= 0.059 log 2

1

P

P......[P

1 > P

2]

Practice Problems :

1. E0 for the cell Zn | Zn2+(aq) || Cu2+ (aq) | Cu is 1.10 V at 250C. The equilibrium constant for the

reaction Zn + Cu2+ (aq) Cu + Zn2+ (aq) is of the order of

(a) 10–37 (b) 1037 (c) 10+18 (d) 1017

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2.21 p2

p2 )H(Pt|)M1(H|)H(Pt (where p

1 and p

2 are pressures) cell reaction will be spontaneous if :

(a) p1 = p

2(b) p

1 > p

2

(c) p2 > p

1(d) p

1 = p

2 = 1atm

3. The standard reduction potential of Cu2+ | Cu and Cu2+ | Cu+ are 0.337 and 0.153 V respectively. Thestandard electrode potential of Cu+ | Cu half cell is

(a) 0.184 V (b) 0.827 V (c) 0.521 V (d) 0.490 V

[Answers : (1) b (2) b (3) c]

C7A Batteries

Any battery (actually it may have one or more than one cell connected in series) or cell that we use as asource of electrical energy is basically a galvanic cell where the chemical energy of the redox reaction isconverted into electrical energy.

1. Primary Batteries : In primary batteries, the reaction occurs only once and battery then becomes deadafter use over a period of time and cannot be reused again, for e.g., :

(a) Dry cell : Which is used commonly in our transistors and clocks. The cell consists of azinc container that also acts as anode and the cathode is a carbon (graphite) rod surrounded bypowdered magnese dioxide and carbon. The space between the electrodes is filled by a moistpaste of NH

4Cl and ZnCl

2. The electrode reactions are complex, but they can be written

approximately as follows :

Anode : Zn(s) Zn2+ + 2e–

Cathode : MnO2 + NH

4+ + e– MnO(OH) + NH

3

Ammonia produced in the reaction forms complex with Zn2+ to give [Zn(NH3)

4]2+. The cell has

a potential of nearly 1.5 V.

(b) Mercury cell : Suitable for the low current devides like hearing aids and camera etc. consists ofzinc-mercury amalgam as anode and a paste of HgO and carbon as the cathode. The electrolyteis a paste of KOH and ZnO. The electrode reactions for the cell are given below :

Anode : Zn(Hg) + 2OH– ZnO(s) + H2O + 2e–

Cathode : HgO + H2O + 2e– Hg(l) + 2OH–

2. Secondary Batteries : A secondary cell after use can be rechanged by passing current through it inopposite direction so that it can be used again. A good secondary cell can undergo a large number ofdischarging and charging cycles, for e.g. :

(a) Lead storage battery : commonly used in automobiles and invertors. It consists of a lead anodeand a grid of lead packed with lead dioxide (PbO

2) as cathode. A 38% solution of sulphuric acid

is used as an electrolyte.

The cell reactions when the battery is in use are given below :

Anode : Pb(s) + SO42–(aq) PbSO

4(s) + 2e–

Cathode : PbO2(s) + SO

42–(aq) + 4H+(aq) + 2e– PbSO

4(s) + 2H

2O (l)

i.e., overall cell reaction consisting of cathode and anode reaction is :

Pb(s) + PbO2(s) + 2H

2SO

4(aq) 2PbSO

4(s) + 2H

2O(l)

On charging the battery the reaction is reversed and PbSO4(s) on anode and cathode is converted

into Pb and PbO2, respectively.

(b) Nickel-cadmium cell : Which has longer life than the lead storage cell but more expensive tomanufacture. We shall not go into detains of working of the cell and the electrode reactionsduring charging and discharging. The overall reaction during discharge is :

Cd(s) + 2Ni(OH)3(s) CdO (s) + 2Ni(OH)

2 (s) + H

2O(l)

3. Fuel Cells : Production of electricity by thermal plants is not a very efficient method and is a major sourceof pollution. In such plants, the chemical energy (heat of combustion) of fossil fuels (coal, gas or oil) is firstused for converting water into high pressure steam. This is then used to run a turbine to produce electricity.

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It is now possible to make such cells in which reactants are fed continuously to the electrodes and productsare removed continuously from the electrolyte compartment. Galvanic cell that are designed to convert theenergy of combustion of fuels like hydrogen, methane, methanol etc. directly into electrical energy arecalled fuel cells. One of the most successful fuel cells uses the reaction of hydrogen with oxygen to formwater. The cell was used for providing electrical power in the Apollo space programme. The water vapoursproduced during the reaction were condensed and added to the drinking water supply for the astronauts.

Catode : O2(g) + 2H

2O(l) + 4e– 4OH–(aq)

Anode : 2H2 + 4OH–(aq) 4H

2O(l) + 4e–

Overall reaction being : 2H2(g) + O

2(g) 2 H

2O(l)

C7B Corrosion

Corrosion slowly coats the surfaces of metallic objects with oxides or other salts of the metal. The rustingof iron, tarnishing of silver, development of green coating on copper and bronze are some of the examplesof corrosion.

In corrosion, a metal is oxidised by loss of electrons to oxygen and formation of oxides. Corrosion of iron(commonly known as rusting) occurs in presence of water and air. The chemistry of corrosion is quitecomplex but it may be considered essentially as an electrochemical phenomenon. At a particular spot of anobject made of iron, oxidation takes place and that spot behaves as an anode and we can write thereaction :

Anode : 2 Fe(s) 2 Fe2+ + 4e– E0(Fe2+, Fe) = –0.44 V

Cathode : O2(g) + 4H+(aq) + 4e– 2 H

2O (l) [E0 = 1.23 V]

The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in theform of hydrated ferric oxide (Fe

2O

3 . xH

2O) and with further production of hydrogen ions.

Prevention of corrosion : one of the simplest method of preventing corrosion is to prevent the surface ofthe metallic object to come incontact with atmosphere. This can be done by covering the surface by paint orby some chemicals (e.g. bisphenol). Other simple method is to cover the surface by other metals (Sn, Znetc.) that are inert or react to save the object. An electrochemical method is to provide a sacrificial electrodeof another metal (like Mg, Zn etc.) which corrodes itself but saves the object.

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SINGLE CORRECT CHOICE TYPE

1. A gas X at 1 atm is bubbled through a solutioncontaining a mixture of 1 M Y— and 1 M Z— at 250C.If the reduction potential of Z > Y > X, then

(a) Y will oxidize X and not Z

(b) Y will oxidize Z and not X

(c) Y will oxidize both X and Z

(d) Y will reduce both X and Z

2. The E0cell

in which the reaction :

MnO—4 + Fe2+ + H+ Mn2+ + Fe3+ + H

2O occurs is

0.59 V at 250C. The equilibrium constant for thereaction is approximately of the order of

(a) 50 (b) 10

(c) 1050 (d) 105

3. Using the standard potential values given below,decide which of the statements I, II, III, IV arecorrect. Choose the right answer from (a), (b), (c)and (d)

Fe2+ + 2e— = Fe, E0 = 0.44 V

Cu2+ + 2e— = Cu, E0 = +0.34 V

Ag+ + e— = Ag, E0 = +0.80 V

I. Copper can displace iron from FeSO4

solution

II. Iron can displace copper from CuSO4

solution

III. Silver can displace copper from CuSO4

solution

IV. Iron can displace silver from AgNO3

solution

(a) I and II (b) II and III

(c) II and IV (d) I and IV

4. For the cell Zn(s) | Zn2+ || Cu2+ | Cu(s), the standardcell voltage, E0cell is 1.10 V. When a cell using thesereagents was prepared in the lab, the measured cellvoltage was 0.98 V. One possible explanation forthe observed voltage is :

(a) There were 2.00 mol of Zn2+ but only 1.00mol of Cu2+.

(b) The Zn electrode had twice the surfaceof the Cu electrode.

(c) The [Zn2+] was larger than the [Cu2+]

(d) The volume of the Zn2+ solution waslarger than the volume of the Cu2+

solution.

5. Cost of electricity for the production of x L H2 at

NTP at cathode is Rs. x, then cost of electricity forthe production of x L O

2 gas at NTP at anode will

be (assume 1 mol of electrons as one unit ofelectricity) :

(a) 2x (b) 4x

(c) 16x (d) 32x

6. Salts of A (atomic weight 7), B (atomic weight 27)and C (atomic weight 48) were electrolysed underidentical conditions using the same quantity ofelectricity. It was found that when 2.1 g of A wasdeposited, the weights of B and C deposited were2.7 g and 7.2 g. The valencies of A, B and C arerespectively

(a) 3, 1 and 2 (b) 1, 3 and 2

(c) 3, 1 and 3 (d) 2, 3 and 2

7. The cell reaction for the given cell is spontaneousif :

PtCl2|Cl—(1M) || Cl— (1M) | PtCl

2

P1

P2

(a) P1 > P

2(b) P

1 < P

2

(c) P1 = P

2(d) P

2 = 1 atm

8. The charge in coulomb on 1 g ion of N–3 is

(a) 2.89 × 105 (b) 2.89

(c) 1.89 × 102 (d) 5.89 × 109

9. The molar conductance of KCl, KNO3 and AgNO

3

are 149.9, 145.0, and 133.4 ohm–1 cm2 mol–1. (all at250C). The molar conductance of AgCl at thistemperature is

(a) 144.4 (b) 120.8

(c) 138.3 (d) 178.2

10. Given the following half cell reaction andcorresponding reduction potentials :

(i) A + e— A— E0 = – 0.24 V

(ii) B— + e— B2— E0 = 1.25 V

(iii) C— + 2e— C3— E0 = 0.68 V

(iv) E + 4e— E4— E0 = 0.38 V

The largest potential resulted by the combinationof two half cells is

(a) E0 = 2.50 V (b) E0 = 1.50 V

(c) E0 = 0.50 V (d) E0 = 1.20 V

11. A solution of sodium sulphate in water iselectrolysed using inert electrodes. The products atthe cathode and anode are respectively

(a) H2, O

2(b) O

2, H

2

(c) O2, Na (d) O

2, SO

2

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12. The value of the standard potentials for reductionreactions of A+/A, B2+/B, C2+/C and D2+/D at 250 are0.80, 0.34, –0.76 and –1.66 volts respectively. Thecorrect sequence in which these metals will bedeposited on the cathode is

(a) A, B, C, D (b) D, C, B, A

(c) A, C, B, D (d) D, B, C, A

13. The standard reduction potential for half-reactionfor four different elements A, B, C and D are

A2 + 2e— 2A— –E0 = +2.85 V

B2 + 2e 2B— E0 = +1.36 V

C2 + 2e— 2C— E0 = +1.06 V

D2 + 2e— 2D— E0 = +0.53 V

The strongest oxidising and reducing agents amongthese

(a) would be A and D respectively

(b) would be D and A respectively

(c) would be B and C respectively

(d) cannot be ascertained from the given dataas the species being subjected tooxidation or reduction have not beenindicated

14. The reaction ½ H2(g) + AgBr(s) H+(aq) + Br—

(aq) + Ag(s) occurs in the galvanic cell

(a) Ag | AgBr(s) | KBr (aq) || AgNO3 (aq) | Ag

(b) Pt | H2(g) | HBr (aq) || AgNO

3 (aq) | Ag

(c) Pt | H2(g) | HBr (aq) | AgBr(s) | Ag

(d) Pt | H2(g) | KBr (aq) || AgBr (s) | Ag

15. In a electrolytic cell is which of the followingstatement is correct :

(a) oxidation occurs at cathode

(b) reduction occurs at anode

(c) anode acts as a negative terminal

(d) flow of electron takes place from anodeto cathode

16. The standard oxidation potentials of the electrodesAg|Ag+, Sn|Sn2+, Ca|Ca2+, Pb|Pb2+ are –0.8, 0.1.36,2.866 and 10.126 V respectively. The most power-ful oxidising agent among these metals is

(a) Pb (b) Ca

(c) Sn (d) Ag

17. Given that 0

Fe|Fe

0

Fe|Fe 23 EandE are –0.36 V and

– 0.439 V, respectively. The value of 0

Pt|Fe,Fe 23E

would be

(a) (–0.36 – 0.439)V

(b) [3(–0.36) + 2(–0.439)]V

(c) (–0.36 + 0.439)V

(b) [3(–0.36) – 2(–0.439)]V

18. For the half cell

+ 2H+ + 2e—, E0 = 1.30 VV

At pH = 2, electrodes potential is :

(a) 1.36 V (b) 1.30 V

(c) 1.42 V (d) 1.20 V

19. Ag | Ag+ (1M) || Ag+ (2M) | Ag

1 L solution 1 L solution

0.5 F of electricity in the LHS (anode) and 1F ofelectricity in the RHS (cathode) is first passedmaking them independent electrolytic cells at 298K. EMF of the cell after electrolysis will be :

(a) increased

(b) decreased

(c) no change

(d) time is also required

20. During electrolysis of acidified water, O2 gas is

formed at the anode. To produce O2 gas at the

anode at the rate of 0.224 c.c. per second at STP,current passed is

(a) 0.224 A (b) 2.24 A

(c) 9.65 A (d) 3.86 A

21. 100 mL of a buffer of 1M NH3(aq) and 1 M NH

4+

(aq) are placed in two voltaic cells separately. Acurrent of 1.5 A is passed through both cells for 20minutes. If electrolysis of water only takes place :

2H2O + O

2 + 4e— 4OH— (RHS)

2H2O 4H+ + O

2 + 4e— (LHS) then pH of the :

(a) LHS half-cell will increase

(b) RHS half-cell will increase

(c) both half-cells will increase

(d) both half-cells will decrease

22. A hydrogen electrode placed in a buffer solution ofCH

3COONa and acetic in the ratio’s x : y and y : x

has electrode potential values E1 volt and E

2 volt

respectively at 2500C. The pKa values of acetic acidis (E1 and E2 are oxidation potential)

(a)118.0

EE 21 (b) 0

118.0

EE 12

(c)118.0

EE 21 (d) 0118.0

EE 21

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23. A 100 watt. 110 volt incandescent lamp in series withan electrolytic cell containing cadmium sulphatesolution. The mass of cadmium will be depositedby the current flowing for 10 hours is(At. wt. Cd = 112.4).

(a) 16.02 (b) 19.06

(c) 20.22 (d) 25.22

24. The solubility product of AgI from the followingdata

V15.0EandV80.0E 0

Ag|AgI|I

0

Ag/Ag

(a) 8.9 × 10–17 (b) 6.9 × 10–14

(c) 2.3 × 10–10 (d) 5.2 × 10–9

25. Zn + Cu2+ (aq) Cu + Zn2+ (aq). Reaction

quotient is ]Cu[

]Zn[Q

2

2

. Variation of Ecell

with log

Q is of the type with OA = 1.10 V. Ecell

will be 1.1591V when :

(a) [Cu2+]/[Zn2+] = 0.01

(b) [Zn2+]/[Cu2+] = 0.01

(c) [Zn2+]/[Cu2+] = 0.1

(d) [Zn2+]/[Cu2+] = 1

26. How much will the reduction potential of ahydrogen electrode change when its solutioninitially at pH = 0 is neutralised to pH = 7

(a) Increase by 0.059 V

(b) Decrease by 0.059 V

(c) Increase by 0.41 V

(d) Decrease by 0.41 V

27. Chromium plating can involve the electrolysis ofan electrolyte of an acidified mixture of chromicacid and chromium sulphate. If during electrolysisthe article being plated increases in mass by 2.6 gand 0.6 dm3 of oxygen are evolved at an inertanode, the oxidation state of chromium ions beingdischarged must be

(assuming Cr = 52 and 1 mole of gas at roomtemperature and pressure occupies a volume of24 dm3)

(a) –1 (b) Zero

(c) + 1 (d) + 2

28. Equal quantities of electricity are passed throughthree voltmeters containing FeSO

4, Fe

2(SO

4)

3 and

Fe(NO3)

3 consider the following statements in this

regard.

1. The amount of iron deposited in FeSO4

and Fe2(SO

4)

3 is equal

2. The amount of iron deposited in Fe(NO3)

3

is two thirds of the amount of irondeposited in FeSO

4.

3. The amount of iron deposited in Fe2(SO

4)

3

and Fe(NO3)

3 is equal to these

statements

(a) 1 alone is correct

(b) 1 and 2 are correct

(c) 2 and 3 are correct

(d) 3 alone is correct

29. In an electrolysis of an aqueous solutioncontaining sodium ions, 2.4 L of oxygen at STP wasliberated at anode. The volume of hydrogen at STPliberated at cathode would be

(a) 1.2 L (b) 2.4 L

(c) 2.6 L (d) 4.8 L

30. One faraday of current was passed through theelectrolysis cells placed in series containingsolutions of Ag+, Ni++ and Cr+++ respectively. Theamount of Ag, Ni and Cr, atomic masses 108, 59,and 52 g mol–1 respectively; deposited will be

Ag Ni Cr

(a) 108 g 29.5 g 17.5 g

(b) 108 g 59.0 g 52.0 g

(c) 108 g 108 g 108 g

(d) 108 g 117.5 g 166 g

ANSWERS (SINGLE CORRECTCHOICE TYPE)

21. b

22. a

23. b

24. a

25. b

26. d

27. d

28. c

29. d

30. a

1. a

2. c

3. d

4. c

5. a

6. b

7. b

8. a

9. c

10. b

11. a

12. a

13. a

14. c

15. d

16. d

17. d

18. c

19. c

20. d

CE – 11


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EXCERCISE BASED ON NEW PATTERN

COMPREHENSION TYPE

Comprehension-1

Tollens reagent (ammonical solution of silvernitrate) is used to test aldehydes. The following dataare available.

Ag+ + e– Ag : E10 = 0.80 V

C6H

12O

7 + 2H+ + 2e–

C6H

12O

6 + H

2O; E

20 = 0.05 V

[Ag(NH3)

2]+ + e–

Ag(s) + 2NH3

E30 = 0.373 V

1. The value of log K0eq

for the reaction

C6H

12O

6 + 2Ag+ + H

2O = C

6H

12O

7 + 2H+ + Ag is

(a) 12.7 (b) 25.4

(c) 29.27 (d) 58.54

2. The use of NH3 makes the pH of solution equal to

11. This causes

(a) decrease in the value of E2

(b) increase in the value of E2

(c) increase in the value of E1

(d) increase in the value of E10

3. Ammonia is used in this reaction rather than anyother base. This is due to the fact that

(a) [Ag(NH3)

2]+ is a weaker oxidizing agent

than Ag+

(b) ammonia prevents the decomposition ofgluconic acid

(c) silver precipitates gluconic acid as itssilver salt

(d) the standard reduction potential of[Ag(NH

3)

2]+ is changes.

Comprehension-2

A saturated solution of silver bromide is made10–7 M in silver nitrate. Given :

Ksp

0(AgBr) = 3.0 × 10–13, m(Ag+) = 6 × 10–3 S m2

mol–1,

m(NO

3–) = 7 × 10–3 S m2 mol–1,

m(Br–) = 8 × 10–3

S m2 mol–1, and K(water) = 7.5 × 10–6 S m–1.

4. Calculate the specific conductance of AgNO3

(a) 5 × 10–7 (b) 13 × 10–7

(c) 14 × 10–3 (d) 14 × 10–7

5. The specific conductor of AgBr is

(a) 5 × 10–7 (b) 13 × 10–7

(c) 14 × 10–3 (d) 70 × 10–7

6. Specific conductance of solution is

(a) 158 × 10–7 (b) 13 × 10–7

(c) 14 × 10–3 (d) 70 × 10–7

Comprehension-3

For the reaction MnO4– + 8H+ + 5Fe2+ Mn2+ +

4H2O + 5Fe3+, it is given that E0(MnO

4–, Mn2+, H+|Pt)

= 1.51 V and E0(Fe3+, Fe2+|Pt) = 0.77 V.

7. The cell emf could be increased above the standardemf by

(a) increasing [Mn2+]

(b) increasing [Fe3+]

(c) decreasing [MnO4–]

(d) decreasing pH of the solution

8. Reducing [Fe3+] to 0.50 M keeping all otherconcentrations at unity, the emf of the cell will bechanged by

(a) –0.059 V (b) –0.0178 V

(c) 0.059 V (d) 0.0178 V

9. Reducing [MnO4

–] to 0.50 M keeping all otherconcentrations at unity, the change in emf of thecell will be changed by

(a) –0.018 V (b) 0.0036 V

(c) 0.018 V (d) –0.0036 V

Comprehension-4

The Edison storage cell is represented asFe(x) | FeO(s) | KOH (aq) | Ni

2O

3(s) | Ni(s)

The half-cell reaction are

Ni2O

3(s) + H

2O(l) + 2e— 2NiO(s) + 2OH—

E0 = +0.40 V

FeO(s) + H2O(l) + 2e— Fe(s) + 2OH—

E0 = –0.87 V

10. The cell reaction is

(a) Ni2O

3(s) + Fe(s) 2NiO(s) + FeO(s)

(b) 2NiO(s) + FeO(s) Ni2O

3(s) + Fe(s)

(c) Ni2O

3(s) + Fe(s) Ni(s) + Fe2+

(d) Ni2O

3(s) + Fe(s) NiO + Fe2+

11. What is the cell e.m.f. ? How does it depend on theconcentration of KOH ?

(a) Ecell

= 0.27 V independent ofconcentration of KOH

(b) Ecell

= 1.27 V independent ofconcentration of KOH

(c) Ecell

= –1.27 V dependent ofconcentration of KOH

(d) Ecell

= 1.27 V dependent ofconcentration of KOH

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12. The maximum amount of electrical energy thatcan be obtained from one mole of Ni

2O

3

(a) 24.52 kJ (b) 2.452 kJ

(c) 0.2452 kJ (d) 245.11 kJ

Comprehension-5

The standard reduction potential for the followinghalf-cell is 0.78 V

NO–3(aq) + 2H+(aq) + e— NO

2(g) + H

2O

13. The reduction potential in 8 M H+ is

(a) 0.887 V (b) –0.887 V

(c) 0.32 V (d) –0.78 V

14. The reduction potential of the half-cell in a neutralsolution. Assume all the other species to be at unitconcentration.

(a) 0 (b) 0.887

(c) 0.0474 V (d) –0.0474 V

15. How much will the reduction potential of the abovementioned half cell will change when its solutioninitially at pH = 0 is neutralized to pH = 7 whilekeeping the other ions concentration to be 1 molar.

(a) decreases by 0.41 V

(b) increases by 0.826 V

(c) decreases by 0.826 V

(d) increases by 0.41 V

MATRIX-MATCH TYPE

Matching-1

Column - A Column - B

(A) Galvanic cell (P) Ecell

= ER – E

L

(B) Cathode of an (Q) reduction

electrolytic cell potentials

(C) Electrode potential (R) Q/t

(D) Current, I (S) Symbol : F

(E) Faraday constant (T) negativelycharged

Matching-2

Column-A Column-B

(A) Faraday’s required (P) 1/3 mol of Al

to reduce Cr2O

7– to Cr3+ deposited

(B) By passing one Faraday (Q) 6 F

of electricity

(C) Zn | Zn2+ (C1 = 0.05) || (R) –0.42 V

Zn2+ (C2 = 0.5) | Zn

(D) E0OCl–/Cl– = .94 (S) 0.0295 V

and E0Cl–/Cl

2 = –1.36 v.

The E0OCl–/Cl2 will be

(T) 3 F

Matching-3

Column-A Column-B

(A) The equivalent (P) 189 ohm–1 cm2

conductivity of 1M mol–1

H2SO

4 solution whose

conductivity is 26 × 10–2

ohm–1 cm–1 is

(B) Given that : (Q) 130 ohm–1 cm2

m[Al

2(SO

4)

3] = eq–1

858 ohm–1 cm2 mol–1

m(SO

42–) = 160

ohm–1 cm2 mol–1

then m (Al3+) is

(C) Given that : (R) 1.66 cm–1

m

c(NH4OH) =

9.33 ohm–1 cm2

m(NH

4OH) =

238.3 ohm–1 cm2 mol–1

the degree of dissociation

of NH4OH is

(D) The conductivity of (S) 5 %

N/10 KCl solution at

200C is 0.0212 ohm–1

cm–1 and the resistance

to the cell containing this

solution at 200C is 55 ohm.

The cell constant is

(T) 3.92 %

MULTIPLE CORRECT CHOICE TYPE

1. Rusting of iron can be prevented

(a) by electroplating the metal with silver

(b) by electroplating the metal with Gold

(c) by electropting the metal with Zn

(d) by connecting the iron material with Mg

2. For the cell :

Tl|Tl+(0.0001M)||Cu2+(0.1 M)|Cu

Ecell

= 0.83 at 298 k the cell potential can be increasedby

(a) increasing [Cu2+]

(b) increasing (Tl+)

(c) decreasing [Cu2+]

(d) decreasing [Tl+]

3. When net cell reaction is spontaneous which of thefollowing are correct

(a) E0 cell is negative

(b) Ecell

> 0

(c) Ecell

= E0cell

(d) G < 0

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4. In which case Ecell

– E0cell

= zero

(a) Cu|Cu2+ (0.01M) || Ag+(0.1M) | Ag

(b) Pt(H2)|PH = 1 || Zn2+ (0.01M) | Zn

(c) Pt(H2)|PH = 1||Zu2+(1M)|Zn

(d) Pt(H2)|H+(0.01M) || Zn2+ (0.01M) | Zn

5. By passing IF of electricity

(a) 1 mol of Zn deposited

(b) 12 g of Mg deposited

(c) 1/3 mol of Al deposited

(d) 11.2 L of O2 at N.T.P. evolved

6. KI solution containing starch turns blue onadditionof Cl

2. Which of the following statements

helps to explain this :

(a) Reduction potential of Cl2 > I

2

(b) E0OX

of Cl2 > E0

OH of I

2

(c) the product of Cl2 and starch is blue in

colour

(d) the product of I2 and starch is blue in

colour

7. The chemical reactions taken place in the processof rusting of iron are listed below, pick up the cor-rect chemical reactions

(a) 2Fe(s) + 4H++ O2 2Fe2+aq + 2H

2O

(b) 4Fe2+aq + O2(g) + 4H

2O 2Fe

2O

3 +

8H+

(c) Fe2O

3 + x H

2O Fe

2O

3 xH

2O

(d) 3Fe + 4H2O Fe

2O

3 + FeO + 4H

2

8. Which of the following are not reference electrode

(a) Normal hydrogen electrode

(b) Calomel electrode

(c) Silver-Silver chloride electrode

(d) Platinium electrode

9. For a galvanic cell,

(a) anode is a negative terminal and cathodeis a positive terminal

(b) oxidation takes place at anode andreduction at cathode.

(c) electrons in the external wire move fromanode to cathode.

(d) Ecell

= ER – E

L

10. The emf of the cell

Pt|H2(g)||HCl(c

1)||HCl(c

2)|H

2(g)|Pt can be increased

by

(a) decreasing c2

(b) decreasing c1

(c) increasing c1

(d) increasing c2

(Answers) EXCERCISE BASED ON NEW PATTERN

COMPREHENSION TYPE

1. b 2. a 3. a 4. b 5. d 6. a

7. d 8. d 9. d 10. a 11. b 12. d

13. a 14. d 15. c

MATRIX-MATCH TYPE

1. [A-P; B-T; C-P; D-R; E-S] 2. [A-Q; B-P; C-S; D-R] 3. [A-Q, B-P, C-T, D-R]

MULTIPLE CORRECT CHOICE TYPE

1. c, d 2. a, d 3. b, d 4. a, b 5. b, c 6. a, d

7. a, b,c 8. a, c 9. a, b, c, d 10. b, d

CE – 14


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INITIAL STEP EXERCISE

(SUBJECTIVE)

1. A current of 1.50 A flows through a cell containingaqueous NiSO

4 for 40 min. What mass of Ni is

deposited on the cathode and what volume at (STP)of oxygen will be evolved at the anode ? (At. wt. ofNi = 58.9).

2. In an electrolysis experiment electric current waspassed for 5 hrs through two cells connected inseries. The first cell contains gold salt and secondcontain CuSO

4 solution. 9.83 g of gold was

deposited in the first cell. If the oxidation numberof Au is +3, find the amount of copper deposited inthe second cell. Also calculate the magnitude of thecurrent in amperes. (At wt. of Au = 197 and At. wt.of Cu = 63.5)

3. An electric current is passed through twoelectrolytic cells connected in series, onecontaining a solution of silver nitrate and the othersolution of sulphuric acid. What volume of oxygenmeasured at 250C and 750 mm of Hg would beliberated from H

2SO

4 if (i) 1 mole and (ii) 8 × 1022

ions of Ag+ are deposited from the silver nitratesolution ?

4. Calculate the e.m.f. of the following cell andpredict whether the given cell representation iscorrect or wrong. If wrong, write the correctrepresentation and correct cell reaction

Cu(s) | Cu2+ (aq) || Zn2+(aq) | Zn(s)

Given :

Cu/Cu0

2E = +0.34 V and V76.0E Zn/Zn0

2

5. Calculate emf of the cell

Pt(H2) | CH

3COOH (0.1 M) || NH

4OH (0.01 M) |

(H2) Pt

Ka for CH

3COOH = 1.8 × 10—5 and K

b for

NH4OH = 1.8 × 10–5

6. Consider the reaction 2Ag+ + Cd 2Ag + Cd2+.The standard electrode potentials for Ag+/Ag andCd2+/Cd couples 0.80 V and 0.40 V respectively. (i)What is the standard potential for this reaction ?(ii) For the electrochemical cell in which the abovereaction takes place which electrode is negativeelectrode ? (iii) Will the total emf of the givenreaction be more negative or positive, if theconcentration of Cd2+ ions is 0.1 M rather than1.0 M ?

7. A zinc rod is placed in 0.1 M solution of ZnSO4 at

250C. Assuming that the salt is dissociated to theextent of 95% at this dilution, calculate thepotential of the electrode at this temperature. Given

that V76.0E Zn/Zn0

2

8. The emf of the cell

Zn | ZnCl2(0.05 mol dm–3) | Ag+ | AgCl(s) | Ag

is 1.015 V at 298 K, the silver electrode beingpositive, while the temperature coefficient of its emfis –0.00492 VK–1. Write down the equation for thereaction occuring when the cell is allowed todischarge and calculate the changes in (a) freeenergy (G), and (b) heat content (H) and (c)entropy (S) accompanying this reaction, at 298 K.

9. The standard reduction potential of Cu2+/Cu andAg+/Ag electrodes are 0.337 and 0.799 voltrespectively. Construct a galvanic cell using theseelectrodes so that its standard emf is positive. Forwhat concentration of Ag+ will the emf of the cell at250C be zero, if the concentration of Cu2+ is0.10 M ?

10. The electrolysis of a metal salt solution was carriedout by passing a current of 4 amp for 45 min. Itresulted in deposition of 2.977 of a metal. If atomicmass of the metal is 106.4 g mol–1, calculate thecharge on the metal cation.

11. 40 ml of 0.126 M NiSO4 solution is electrolysed by a

current of 0.05 amp for 40 min.

(a) Write equation for the reactionsoccurring at each electrode

(b) How many coulombs of electricity arepassed through the electrolyte

(c) How many grams of product is depositedat cathode

(d) How long the same current will have tobe passed to remove completely themetal ions from the solution.

12. Silver is electrodeposited on a metallic vessel ofsurface area 800 cm2 by passing a current of 0.2ampere for 3 hours. Calculate the thickness of thesilver deposited. Given the density of silver as10.47 g/cc.

13. Calculate standard electrode potential of Ni+2/Nielectrode if the cell potential of a cell Ni | Ni+2 (0.01M) || Cu2+ (0.1 M) | Cu is 0.59 V. Given that

V34.0E0

Cu/Cu2

14. Calculate the maximum possible electrical workthat can be obtained from the following cell understandard conditions at 250C.

Zn | Zn+2(aq) || Cu2+ (aq) | Cu

At 250C,

V34.0EandV76.0E 0

Cu/Cu

0

Zn/Zn 22

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15. Calculate the E.M.F. of the cell

Mg | Mg2+ (0.2 M) || Ag+ (1 × 10–3 M) | Ag

V80.0E,V37.2E 0

Ag/Ag

0

Mg/Mg2

What will be the effect on EMF if concentrations ofMg2+ ion is decreased to 0.1 M

16. The measured voltage of the following cell is 0.9 Vat 250C.

Pt, H2(1 atm) | H+(aq) || Ag+ (1.0 M) | Ag

If V80.0E0

Ag/Ag . Calculate the pH of the

aqueous solution of H+.

17. How many grams of silver could be plated out on aserving tray by electrolysis of a solution containingsilver in + 1 oxidation state for a period of 8.0 hoursat a current of 8.46 Amperes ? What is the area ofthe tray if the thickness of the silver plating is0.00254 cm ? Density of silver is 10.5 g/cm3.

18. A 100 watt. 110 volt incandescent lamp in series withan electrolytic cell containing cadmium sulphatesolution. What mass of cadmium will be depositedby the current flowing for 10 hours ? (At. wt.Cd = 112.4).

19. A current of 10.0 A is passed through 1.0 L of 1.0 MHCl solution for 1.0 h. Calculate the pH of thesolution at the end of the experiment. What is thevolume of total gas evolved at STP ?

20. Consider the cell

Zn | Zn2+ (aq. 1.0 M) || Cu2+ (aq 1.0 M) | Cu

The standard reduction potentials are 0.350 V forCu2+ (aq) + 2e— Cu(s) and – 0.763 V for Zn2+ (aq)+ 2e— Zn(s)

(i) Write down the cell reaction

(ii) Calculate the emf of the cell

(iii) Is the cell reaction spontaneous or not ?

21. Given the following cell

Al | Al3+ (0.1 M) || Fe2+ (0.2 M) | Fe

V44.0EandV66.1E 0

Fe/Fe

0

Al/Al 23 .

Calculate the maximum work that can be obtainedby the cell.

22. Determine the equilibrium constant of thefollowing reaction at 298 K

2Fe3+ + Sn2+ 2Fe2+ + Sn4+

Also predict whether Sn2+ ions can reduce Fe3+ ionsto Fe2+ quantitatively or not

V771.0E0

Fe/Fe 22 , V150.0E0

Sn/Sn 24

23. Calculate pH of the following half cell, Pt H2 | H

2SO

4.

The oxidation electrode potential is + 0.3 V.

24. (a) Calculate the electrode potential at acopper electrode dipped in a 0.1 Msolution of copper sulphate at 298 K;assuming CuSO

4 to be completely

dissociated. The standard electrodepotential of Cu2+ | Cu system is + 0.34

volts at 298 K.

(b) At what concentration of copper ions willthis electrode have a potential of zeovolt ?

25. A current was passed through a series of cellscontaining AgNO

3, CuSO

4 and H

2SO

4 solutions for

a period of 25 minutes. If the weight of silverdeposited was 0.5394 g, what would be (i) the weightof copper and (ii) the volume of H

2 at N.T.P.

liberated by the current ? Also calculate themagnitude (strength) of current assuming that itremained constant. [At. wt. Ag = 108, Cu = 63.5].

26. A solution of a salt of a metal of atomic weight 112was electrolysed for 150 minutes with a current of0.15 amperes. The weight of metal deposited was0.783 mg. Find the equivalent weight and valencyof the metal in the salt.

27. A current deposits 10–2 kg of Cu in 3.96 × 103 minutefrom a solution of Cu++ ions. What is the strengthof current in amperes ? How many grams of Cuwill the same current deposit from a solution ofcuprous ions ?

28. A 1.5276 g sample of CdCl2 was converted to

metallic cadmium by an electrolytic process.0.9367 g of cadmium was obtained. What is theatomic mass of cadmium from this experiment ifthe atomic mass of chlorine is taken as35.453 g mol—1.

29. Calculate heat of reaction inside the cell

Zn(s) + 2AgCl(s) ZnCl2(0.555 M) + 2Ag(s)

Given that E = 1.015 V at 00 C and

14

P

VK1002.4T

E

30. (i) The chemical reaction :

Cl2(g) + SO

2(g) + 2H

2O(l) 2Cl—(aq) +

3H+(aq) + HSO4— (aq)

proceeds readily and rapidly in aqueousacid solution. Write the half cellreactions and construct the cell.

(ii) If the fully charged cell initially held1.0 M of Cl

2, for how many days could it

sustain a current of 0.05 A assuming thatthe cell becomes inoperative when 90%of the initial Cl

2 has been consumed ?

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31. Zinc granules are added in excess to a 500 ml of 1.0M nickel nitr ate solut ion at 250C until theequilibrium is reached. If the standard reductionpotential of Zn2+ | Zn and Ni2+ | Ni are –0.75 V and–0.24 V respectively, find out the concentration ofNi2+ in soluble at equilibrium.

32. (a) The standard reduction potential ofCu++ / Cu and Ag+ / Ag electrodes are0.337 and 0.799 volt respectively.Construct a galvanic cell using theseelectrodes so that its standard e.m.f. ifpositive. For what concentration of Ag+

will the e.m.f. of the cell, at 250C, be zeroif the concentration of Cu++ is 0.01 M ?

FINAL STEP EXERCISE

(SUBJECTIVE)

1. 19 g fused SnCl2 was electrolysed using inert

electrode 0.119 g Sn was deposited at cathode. Ifnothing was given out during electrolysis, calculatethe ratio of weight of SnCl

2 and SnCl

4 in fused state

after electrolysis. (At. wt. of Sn = 119)

2. For the following galvanic cell, calculate the e.m.f.at 250C. Assign the correct polarity for thespontaneous reaction to take place

Ag | AgCl(s), KCl (0.2 M) || KBr (0.001 M), AgBr(s)| Ag

The solubility product of AgCl and AgBr are2.8 × 1010 and 3.3 × 1013 respectively.

3. Calculate the e.m.f. of the following cell at 250C

H2(1 atm) | 0.5 M HCOOH || 1 M CH

3COOH | H

2

(1 atm)

The Ka for HCOOH and CH

3COOH are

1.77 × 10–4 and 1.80 × 10–5 respectively.

4. The following Galvanic cell Zn | Zn++ (1 M, 100 mL)|| Cu++ (1 M, 100 mL) | Cu was operated as anelectrolytic cell as; Cu as anode and Zn as cathode.0.48 A current was passed for 10 hours and thenthe cell was allowed to function as galvanic cell.Calculate e.m.f. of the cell at 250C, assuming thatonly electrode reactions occuring were thoseinvolving Cu/Cu++ and Zn/Zn++.

10

Zn/Zn

0

Cu/Cumolg5.65Zn,V76.0E,V34.0E .

5. An aqueous solution of NaCl on electrolysis givesH

2(g), Cl

2(g) and NaOH according to the reaction :

2Cl–(aq)

+ 2H2O = 2OH—

(aq) + H

2(g) + Cl

2(g).

A direct current of 25 amperes with a currenteffeciency of 62% is passed through 20 litres of NaClsolution (20% by weight). Write down the reactionstaking place at the anode and the cathode. How longwill it take to produce 1Kg of Cl

2 ? What will be the

molarity of the solution with respect to hydroxideion ? (Assume no loss due to evaporation)

6. An acidic solution of Cu2+ salt containing 0.4 g ofCu2+ is electrolysed until all the copper isdeposited. The electrolysis is continued for sevenmore minutes with the volume of solution kept at100 ml and the current at 1.2 amp. Calculate thevolume of gases evolved at NTP during the entireelectrolysis. (At. wt. Cu = 63.6)

7. The standard reduction potential at 250C of the

reaction, 2H2O + 2e— H

2 + 2OH— is – 0.8277

V. Calculate the equilibrium constant for the

reaction 2H2O H

3O+ + OH— at 250C.

8. Find the solubility product of a saturated solutionof Ag

2CrO

4 in water at 298 K if the emf of the cell

Ag/Ag+ (satd. Ag2CrO

4 soln.) // Ag+ (0.1 M)/Ag is

0.164 V at 298 K.

9. The standard reduction potential for Cu2+ / Cu is+0.34 V. Calculate the reduction potential atpH = 14 for the above couple. K

sp of Cu(OH)

2 is

1.0 × 10–19.

10. An excess of liquid mercury is added to anacidified solution of 1.0 × 10–3 M Fe3+. It is foundthat 5% of Fe3+ remains at equilibrium at 250C.

Calculate Hg/Hg0

2E , assuming that the only

reaction that occurs is

2Hg + 2Fe3+ Hg22+ + 2Fe2+

]V77.0EGiven[ 23 Fe/Fe0

11. The standard reduction potential of the Ag+/Agelectrode at 298 K is 0.799 V. Given that for AgI,K

sp = 8.7 × 10–17, evaluate the potential of the

Ag+/Ag electrode in a saturated solution of AgI. Alsocalculate the standard reduction potential of theI—/AgI/Ag electrode.

12. The E.M.F. of the following cell at 298 K is 1.0495V.

Pt/H2(1 atm) |LiOH (0.01 mol dm–3)| |LiCl (0.01)

mol/dms)| AgCl (s) / Ag. Determine the ionicproduct of water.

(b) Calculate the quantity of electricity thatwould be required to reduce 12.3 g ofnitrobenzene to aniline, if the currentefficiency for the process is 50 per cent.If the potential drop across the cell is3.0 volts, how much energy will beconsumed ?

33. A cell, Ag | Ag+ || Cu2+ | Cu, initially contains 1 MAg+ and 1 M Cu2+ ions. Calculate the change in thecell potential after the passage of 9.65 A of currentfor 1 h.

CE – 17


New Delhi – 110 018, Ph. : 9312629035, 8527112111

ANSWERS SUBJECTIVE (INITIAL STEP EXERCISE)

1. 208.7 mL 2. 0.8026 amp. 3. (i) 6.2 litre(ii) 0.823 litre

4. E0cell

= – 1.10 V6. E

cell = 0.4575 V

6. (i) E0cell

= 0.40 V (ii) Cd(s) | Cd2+ is negative electrone(iii) E

cell = 0.4295 V

7. 2Zn/ZnE = – .79 volt 8. (a) – 195.895 kJ mol–1

(b) H298k

= – 167.7 kJ mol–1

(c) S298

= – 95.0 Jk–1 mol–1

9. [Ag+] = 1.47 × 109M 10. n = 411. (b) 120 c (c) .036 gm (d) 5 hr 21 min.

12. 2.89 × 10–4 cm 13. 0

Ni/Ni2E = .22 VV

14. 212.3 kJ 15. 3.103 V, Ecell

will increase to 3.02 V16. 1.69 17. 1.02 × 104 cm2

18. 19.06 gm 19. 0.20, 8360.4 mL20. (i) Zn + Cu2+ Zn2+ + Cu (ii) 1.113 V

(iii) spontaneous21. 705.801 kJ 22. 1.035 × 1021, yes23. 5.08 24. 0.3105 V, 2.95 × 10–12 M25. 0.1589 gm, 56 mL, 0.31218 amperes 26. 55.97, 227. 0.131 A, 20 gm 28. 110.88 gm mol–1

29. –2.171 × 102 kJ30. Pt, SO

2|H

2SO

4|| HCl | Cl

2, Pt, 40.2 days 31. [Ni2+] = 5.6 × 10–18 M

32. (a) [Ag+] = 1.477 × 10–9 M (b) 115800 coulomb, 347.4 kJ33. 0.01 V

ANSWERS SUBJECTIVE (FINAL STEP EXERCISE)

1. SnCl2 : SnCl

4 = 18.26 : 0.26

2. (s) Ag | AgBr(s), kBr || KCl, AgCl(s) | Ag(s), Ecell

= –0.337 V3. –.0204 V 4. 1.1375. 2Cl– Cl

2 + 2e– (Anode)

2e– + 2H2O 2OH– + H

2 (cathode), 48.7 hr, 1.408 mol/lit

6. 158.15 mL 7. 9.35 × 10–15

8. 2.287 × 10–12 9. –0.22 V10. 0.792 V

11. V325.0EAg/Ag

V1485.0EAg/AgI/I

12. 10–14

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CE – 1 ELECTROCHEMISTRYeinsteinclasses.com/Electrochemistry.pdf · It is defined as the conducting power of all the ions present in one mol of an ... thickness of 10–2 cm