Calculate the number of coulombs required to deposit 40 . 5 g of Al when the electrode reaction is Al 3 + + 3 e - → Al ( s ) .

Electrode reaction : Al 3 + + 3 e - → Al ( s ) Given : Weight of aluminium to be deposited = 40 . 5 g Solution : Number of electrons exchanged in 1 mol n = 3 Amount of charge deposited q = n × F Amount of charge required to deposit 1 mol of Al = 3 × 96500 Amount of charge required to deposit 1 g of Al = 3 × 96500 27 . 0 Amount of charge required to deposit 40 . 5 g of Al = 3 × 96500 × 40 . 5 27 = 434250 C Thus, 434250 coulombs are required to deposit 40 . 5 g of aluminium.

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A solution of metal salt was electrolysed for $15$ minutes with a current of $1.5 A$ . The mass of the metal deposited was $0.000783 kg$ . Calculate the equivalent mass of the metal.

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The anodic half-cell of lead-acid battery is recharged using electricity of

0 .05

Faraday. The amount of

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electrolyzed in

g

during the process is: (Molar mass of

{PbSO}_{4} = 303 g {mol}^{- 1}

)

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Two Faraday of electricity is passed through a solution of

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. The mass of copper deposited at the cathode is: (Atomic mass of Cu = 63.5 amu)

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A solution of copper sulphate electrolysed for $20$ minute with a current of $1.5$ Ampere. Calculate the mass of copper deposited at the cathode. $(F = 96500 C)$

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

The number of electrons delivered at the cathode, during electrolysis by a current of

1 ampere in 60 seconds

is (charge on electron

= 1.60 \times 10^{- 19} C

)

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The quantity of electricity needed to separately electrolyse

1 M

solution of

{ZnSO}_{4}, {AlCl}_{3}

and

{AgNO}_{3}

completely is in the ratio of

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During the electrolysis of molten sodium chloride, the time required to produce

0.10 mol

of chlorine gas using a current of

3 amperes

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A steady current of 2 amperes was passed through two electrolytic cells X and Y connected in series containing electrolytes ${FeSO}_{4}$ and ${ZnSO}_{4}$ until $2.8 g$ of $Fe$ deposited at the cathode of cell X. How long did the current flow ? Calculate the mass of $Zn$ deposited at the cathode of cell Y.

$($ Molar mass : $Fe = 56 g {mol}^{- 1} Zn = 65.3 g {mol}^{- 1}, 1 F = 96500 C {mol}^{- 1}$ $)$

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

The density of metal and equivalent weight of a metal are

10.5 g {cm}^{- 3}

and

100

, respectively. The time required for a current of

3 amp

to deposit a

0.005 mm

thick layer of the same metal on an area of

80 {cm}^{2}

is closest to

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

State Faraday's second law of electrolysis.

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

State Faraday's first law of electrolysis.

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

When 9.65 ampere current was passed for 1.0 hour into nitrobenzene in acidic medium, the amount of p-aminophenol produced is:

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

A solution of

N i {(N O_{3})}_{2}

is electrolyzed between platinum electrode

0.1

Faraday electricity. How many mole of

N i

will be deposited at the cathode?

EASY

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

For electroplating,

1.5 amp

current is passed for

250 s

through

250 m L of 0.15 M

solution of

M S O_{4} .

Only

85 %

of the current was utilised for electrolysis. The molarity of

M S O_{4}

solution after electrolysis is closest to [Assume that the volume of the solution remain constant]

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

A constant current (0.5 amp) is passed for 1 hour through (i) aqueous

A g N O_{3}

. (ii) aqueous

C u S O_{4}

and (iii) molten

A l F_{3}

, separately. The ratio of the mass of the metals deposited on the cathode is

[M_{Ag}, M_{Cu}, M_{Al}

are molar masses of the respective metals

]

HARD

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

Calculate the molarity of a solution containing

5 g

N a O H

dissolved in the product of a

H_{2} - O_{2}

fuel cell operated at

1 A

current for

595.1

hours.

(Assume

1 F = 96500 C / m o l

of electrons and molecular weight of

N a O H

40 g m o l^{- 1}

)

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Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

How long (approximate) should water be electrolysed by passing through

100

amperes

current so that the oxygen released can completely burn

27.66 g

of diborane? (Atomic weight of

B = 10.8 u

)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

A current of $10.0 A$ flows for $2.00 h$ through an electrolytic cell containing a molten salt of metal $X$ . This results in the decomposition of $0.250 mol$ of metal $X$ at the cathode. The oxidation state of $X$ in the molten salt is:

$(F = 96, 500 C)$

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

Give reason: On the basis of

E^{\circ}

values,

O_{2}

gas should be liberated at anode but it is

{Cl}_{2}

gas which is liberated during electrolysis of aqueous

NaCl

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

Calculate the number of coulombs required to deposit

40.5 g

Al

when the electrode reaction is

{Al}^{3 +} + 3 e^{-} \to Al (s)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Electrolytic Cells

A solution of

N i {(N O_{3})}_{2}

is electrolyzed between platinum electrodes using a current of

5

amperes for

20

minutes. What mass of

N i

is deposited at the cathode?

A solution of metal salt was electrolysed for 15 minutes with a current of 1.5 A . The mass of the metal deposited was 0.000783 kg. Calculate the equivalent mass of the metal.

Important Questions on Electrochemistry

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A solution of metal salt was electrolysed for $15$ minutes with a current of $1.5 A$ . The mass of the metal deposited was $0.000783 kg$ . Calculate the equivalent mass of the metal.