Calculate the emf of the cell in which of the following reaction takes place:\$Ni(s) + 2Ag^+ (0.002 M) \rightarrow Ni^{2+} (0.160 M) + 2Ag(s)\$[Given that \$E^{\circ}_{cell} = 1.05 V\$]

Given,\$[Ag^+]=0.002M\$\$[Ni^{2+}]=0.160M\$\$n=2\$According to Nernst equation:\$E_{cell}=E^0-\dfrac{0.0591}{2}\log \dfrac{[Ni^{2+}]}{[Ag^+]^2}\$\$\therefore E_{cell}=1.05-\dfrac{0.0591}{2}\log \dfrac{0.160}{(0.002)^2}\$\$\therefore E_{cell}=1.05-0.02955\, \log{(4\times 10^4)}\$\$\therefore E_{cell}=0.91V\$Hence, the correct option is \$\text{B}\$

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>>Class 12
>>Chemistry
>>Electrochemistry
>> $Electrolytic Cells and Electrolysis$

Electrolytic Cells and Electrolysis

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Electrolytic Cells and Electrolysis

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Faraday's Law

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Products of Electrolysis

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Q: How would you determine the standard electrode potential of the \$Mg^{2+}\$ \$|Mg\$?

To determine the electrode potential of \$mg^{+2}/mg^{1}\$ we use hydrogen electrode (anode)The standard hydrogen electrode potential is zero.\$\therefore E^{o}cell = E^{o} right+ E^{o} left\$\$E^{o} cell= E^{o} mg^{+2}/mg^{-1} E^{o}n_{2}/n_{1}\$ \$E^{o} mg^{+2}/mg= E^{o} cell\$

Q: How much charge is required for the following reductions?(i) \$1\ mol\$ of \$Al^{3+}\$ to \$Al\$(ii) \$1\ mol\$ of \$Cu^{2+}\$ to \$Cu\$(iii) \$1\ mol\$ of \${MnO_4}^{-}\$ to \$Mn^{2+}\$

\$Al^{3+} + 3e^-\rightarrow Al\$The reduction will require 3 Faraday of electricity or \$\displaystyle 3 \times 96500 = 2.895 \times 10^5 \: C\$.\$Cu^{2+} + 2e^-\rightarrow Cu\$The reduction will require 2 Faraday of electricity or \$\displaystyle 2 \times 96500 = 1.93 \times 10^5 \: C\$.\${MnO_4}^{-} + 5e^- \rightarrow Mn^{2+} \$The reduction will require 5 Faraday of electricity or \$\displaystyle 5 \times 96500 = 4.825 \times 10^5 \: C\$.Note: 1 mole of electron has a charge of 1 faraday.

Q: How many electrons are there in one coulomb of electricity?

Change on \$1e^-=1.602\times 10^{-19}\$ \$C\$If total charge \$=1\$ \$C\$\$\therefore\$ Total no. of \$e^-=\cfrac{1}{1.602\times 10^{-19}}=6.24\times 10^{18}e^-\$Hence, the correct option is \$(C)\$.

Q: How much electricity in terms of Faraday is required to produce:(i) \$20\ g\$ of \$Ca\$ from molten \$CaCl_{2}\$(ii) \$40\ g\$ of \$Al\$ from molten \$Al_{2}O_{3}\$[Given : Molar mass of calcium & Aluminium are 40 g \$mol^{-1}\$ & 27 g \$mol^{-1}\$ respectively]

\$(1)\$ \$\displaystyle Ca^{2+} + 2e^- \rightarrow Ca\$\$1\$ mole (\$40\$ g) of \$Ca\$ will require \$2\$ F of electricity.\$20\$ g (\$0.5\$ mole) of \$Ca\$ will require \$1\$ F of electricity.(2) \$\displaystyle Al^{3+} + 3e^- \rightarrow Al \$\$1\$ mole (\$27\$ g) of \$Al\$ will require \$3\$ moles of electrons.\$40\$ g of \$Al\$ will require \$\displaystyle \dfrac {3 \times 40F}{27} = 4.4\$ F

Q: A solution of \$Ni(NO_{3})_{2}\$ is electrolysed between platinum electrodes using a current of \$5\$ amperes for \$20\$ minutes. What mass of \$Ni\$ is deposited at the cathode?

Given\$I = 5A\$\$Time = 20 \times 60 = 1200\,\,s\$\$\therefore Charg e = current \times time\$\$\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\, = 5 \times 1200\$\$\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\, = 6000\,\,C\$According to the reaction.\$N{i^{2 + }}\left( {aq.} \right) + 2e \to Ni\,\left( s \right)\$\$Nickel\,\,deposite\,by\, (2 \times 96487)\,C = 58.7\,\,g\$\$\therefore Nickel\,deposite\,by\ 6000\,C = \dfrac{{58.7 \times 6000}}{{2 \times 96487}}\$\$\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\, \quad= 1.825\,\,g\$\$\therefore \$ Hence \$1.825\,g\$ of nickel will be deposited at the cathode.

Q: Three electrolytic cells \$A, B, C\$ containing solutions of \$ZnSO_{4}, AgNO_{3}\$ and \$CuSO_{3}\$, 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?

Cell \$B\$: \$\displaystyle Ag^+ + e^- \rightleftharpoons Ag\$ at cathode.\$1\$ mole (\$108\$ g) of Ag is deposited by \$96500\$ \$C\$.\$1.45\$ g of Ag will be deposited by \$\displaystyle \dfrac {96500 \times 1.45}{108} = 1295.6 \: C\$.Now,\$\displaystyle Q = It \$\$\displaystyle 1295.6 = 1.5 \times t \$\$\displaystyle t = 864 \: s\$Cell \$A\$: \$\displaystyle Zn^{2+} + 2e^- \rightarrow Zn\$\$2\$ moles of electrons (\$\displaystyle 2 \times 96500\$ C of current) produces \$1\$ mole (\$63.5\$ g) of zinc.\$1295.6\$ \$C\$ of electricity will deposit \$\displaystyle \dfrac {65.3}{2 \times 96500} \times 1295.6 = 0.438\$ g of zincCell \$C\$: \$\displaystyle Cu^{2+} + 2e^- \rightarrow Cu\$\$2\$ moles of electrons (\$\displaystyle 2 \times 96500\$ \$C\$) of current will produce \$1\$ mole (\$63.5\$ g) of \$Cu\$. \$1295.6\$ \$C\$ of current will deposit \$\displaystyle \dfrac {63.5 \times 1295.6}{2 \times 96500} = 0.426 \: g\$ of copper

Q: During the electrolysis of molten sodium chloride, the time required to produce 0.10 mol of chlorine gas using a current of 3 amperes is:

\$2Cl^{-}\rightarrow Cl_2+2e^{-}\$\$W=\dfrac{E}{96500}\times t\$\$0.1\times71= \dfrac{35.5}{96500}\times 3\$= \$6433.33 \ sec\$\$= 107.22\ min= 110\ min\$Hence, option \$D\$ is correct.

Q: In electrolysis of water: (a) Name the gas collected at cathode and anode.(b) Why is volume of one gas collected at one electrode is double of anode?(c) Why are few drops of dilute \$H_2SO_4\$ added to water?

a) At cathode, \$H_2\$ is evolved and at anode \$O_2\$ is evolved.b) At cathode, \$H^+\$ ion takes two electron to convert itself into \$H_2\$ gas. \$2\$ moles of \$H^+\$ gives \$1\$ mole of \$H_2\$.\$2H^+(aq)+ 2e^- \rightarrow H_2(g)\$At anode, \$OH^-\$ ion releases two electrons to convert into water and \$O_2\$. \$2\$ mole \$OH^-\$ gives \$0.5\$ mole \$O_2\$, i.e why volume of one gas collected at one electron is double of another.c) Acid is added to make the water conduct electricity as the distilled water is a non-conductor of electricity.

Introduction to Electrolysis

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Faraday's First Law of Electrolysis

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Faradays Second Law of Electrolysis

25 mins

Numerical on Faraday's First Law of Electrolysis

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Numerical on Faraday's Second Law of Electrolysis

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Electrolysis of Molten NaCl

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Quick Summary With Stories

Introduction to Electrolysis

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Faraday Second Law of Electrolysis

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Important Questions

Calculate the emf of the cell in which of the following reaction takes place:

$N i (s) + 2 A g^{+} (0.002 M) \to N i^{2 +} (0.160 M) + 2 A g (s)$

[Given that $E_{c e l l}^{\circ} = 1.05 V$ ]

Medium

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How would you determine the standard electrode potential of the $M g^{2 +}$ $∣ M g$ ?

Hard

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How much charge is required for the following reductions?

(i) $1 m o l$ of $A l^{3 +}$ to $A l$
(ii) $1 m o l$ of $C u^{2 +}$ to $C u$
(iii) $1 m o l$ of $M n O_{4}^{-}$ to $M n^{2 +}$

Medium

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How many electrons are there in one coulomb of electricity?

Medium

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How much electricity in terms of Faraday is required to produce:

(i) $20 g$ of $C a$ from molten $C a C l_{2}$
(ii) $40 g$ of $A l$ from molten $A l_{2} O_{3}$

[Given : Molar mass of calcium & Aluminium are 40 g $m o l^{- 1}$ & 27 g $m o l^{- 1}$ respectively]

Medium

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A solution of $N i (N O_{3})_{2}$ is electrolysed between platinum electrodes using a current of $5$ amperes for $20$ minutes. What mass of $N i$ is deposited at the cathode?

Medium

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Three electrolytic cells $A, B, C$ containing solutions of $Z n S O_{4}, A g N O_{3}$ and $C u S O_{3}$ , 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?

Hard

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A solution of $C u S o_{4}$ is electrolysed for 10 minutes with a current of 1.5 amperes. What is the mass of copper deposited at the cathode?

Medium

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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 is:

Hard

NEET

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In electrolysis of water:

(a) Name the gas collected at cathode and anode.
(b) Why is volume of one gas collected at one electrode is double of anode?
(c) Why are few drops of dilute $H_{2} S O_{4}$ added to water?

Medium

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Revise with Concepts

Electrolytic Cells and Electrolysis

Faraday's Law

Products of Electrolysis

Learn with Videos

Quick Summary With Stories

Introduction to Electrolysis

Faraday Second Law of Electrolysis

Calculate the emf of the cell in which of the following reaction takes place: Ni(s)+2Ag+(0.002M)→Ni2+(0.160M)+2Ag(s) [Given that Ecell∘​=1.05V]

How would you determine the standard electrode potential of the Mg2+ ∣Mg?

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+

How many electrons are there in one coulomb of electricity?

How much electricity in terms of Faraday is required to produce: (i) 20 g of Ca from molten CaCl2​ (ii) 40 g of Al from molten Al2​O3​ [Given : Molar mass of calcium & Aluminium are 40 g mol−1 & 27 g mol−1 respectively]

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?

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?

During the electrolysis of molten sodium chloride, the time required to produce 0.10 mol of chlorine gas using a current of 3 amperes is:

In electrolysis of water: (a) Name the gas collected at cathode and anode. (b) Why is volume of one gas collected at one electrode is double of anode? (c) Why are few drops of dilute H2​SO4​ added to water?

Calculate the emf of the cell in which of the following reaction takes place:

$N i (s) + 2 A g^{+} (0.002 M) \to N i^{2 +} (0.160 M) + 2 A g (s)$

[Given that $E_{c e l l}^{\circ} = 1.05 V$ ]

How would you determine the standard electrode potential of the $M g^{2 +}$ $∣ M g$ ?

How much charge is required for the following reductions?

(i) $1 m o l$ of $A l^{3 +}$ to $A l$
(ii) $1 m o l$ of $C u^{2 +}$ to $C u$
(iii) $1 m o l$ of $M n O_{4}^{-}$ to $M n^{2 +}$

How much electricity in terms of Faraday is required to produce:

(i) $20 g$ of $C a$ from molten $C a C l_{2}$
(ii) $40 g$ of $A l$ from molten $A l_{2} O_{3}$

[Given : Molar mass of calcium & Aluminium are 40 g $m o l^{- 1}$ & 27 g $m o l^{- 1}$ respectively]

A solution of $N i (N O_{3})_{2}$ is electrolysed 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 $C u S o_{4}$ is electrolysed for 10 minutes with a current of 1.5 amperes. What is the mass of copper deposited at the cathode?

In electrolysis of water:

(a) Name the gas collected at cathode and anode.
(b) Why is volume of one gas collected at one electrode is double of anode?
(c) Why are few drops of dilute $H_{2} S O_{4}$ added to water?