Electro chemistry practice · Hard

The cell reaction for the given cell is spontaneous if :

$\mathop {\mathrm{P}{\mathrm{t}_{\mathrm{C}{\mathrm{l}_2}}}}\limits_{{\mathrm{P}_1}} \left| {\mathrm{C}{\mathrm{l}^ - }\left( {1M} \right)} \right|\left| {\mathrm{C}{\mathrm{l}^ - }\left( {1M} \right)} \right|\mathop {\mathrm{P}{\mathrm{t}_{\mathrm{C}{\mathrm{l}_2}}}}\limits_{{\mathrm{P}_2}}$

When electric current is passed through acidified water for 1930 s, 1120mL of ${\mathrm{H}_2}$ gas is collected (at STP) at the cathode. What is the current passed in amperes?

${\mathrm{I}_2}\left( s \right)|{\mathrm{I}^ - }\left( {0.1\mathrm{M}} \right)$ half-cell is connected to a ${\mathrm{H}^ + }\left( {aq} \right)\left| {{\mathrm{H}_2}\left( {1\mathrm{bar}} \right)} \right|$ Pt half-cell and emf is found to be 0.7714 V. If $E_{{I^2}/{I^ - }}^ \circ = 0.535\mathrm{V}$, find the $\mathrm{pH}$ of ${\mathrm{H}^ + }/{\mathrm{H}_2}$ half-cell

At 298 K, the standard electrode potentials of $C{u^{2 + }}/Cu,Z{n^{2 + }}/Zn,F{e^{2 + }}/Fe\,\,\mathrm{and}\,\,A{g^ + }/Ag$ are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V, respectively.
On the basis of standard electrode potential, predict which of the following reaction cannot occur?

The standard reduction potential ${E^ \circ }$ for the half reactions are as
$\mathrm{Zn} \to \mathrm{Z}{\mathrm{n}^{2 + }} + 2{e^ - },{E^ \circ } = 0.76V$
$\mathrm{Cu} \to \mathrm{C}{\mathrm{u}^{2 + }} + 2{e^ - },{E^ \circ } = 0.34V$
The emf for the cell reaction,
$\mathrm{Zn} + \mathrm{C}{\mathrm{u}^{2 + }} \to \mathrm{Z}{\mathrm{n}^{2 + }} + \mathrm{Cu}$

Assertion: Copper sulphate solution cannot be stored in iron-vessel.

Reason: Copper atoms are oxidized by iron ions.

The sum of the two transport number of ions for an electrolyte is always equal to :

Following cell has EMF $0.7995 \mathrm{~V}$.

$\mathrm{Pt}\left|\mathrm{H}_2(1 \mathrm{~atm})\right| \mathrm{HNO}_3(1 \mathrm{M}) \| \mathrm{AgNO}_3(1 \mathrm{M}) \mid \mathrm{Ag}$

If we add enough $\mathrm{KCl}$ to the $\mathrm{Ag}$ cell so that the
final $\mathrm{Cl}^{-}$is $1 \mathrm{M}$. Now the measured emf of the cell is $0.222 \mathrm{~V}$. The $\mathrm{K}_{s p}$ of $\mathrm{AgCl}$ would be-

Two electrolytic cells one containing acidified ferrous chloride and another acidified ferric chloride are connected in series. The ratio of iron deposited at cathodes in the two cells when electricity is passed through the cells will be

${\mathrm{I}_2}\left( s \right)|{\mathrm{I}^ - }\left( {0.1\mathrm{M}} \right)$ half-cell is connected to a ${\mathrm{H}^ + }\left( {aq} \right)\left| {{\mathrm{H}_2}\left( {1\mathrm{bar}} \right)} \right|$ Pt half-cell and emf is found to be 0.7714 V. If $E_{{I^2}/{I^ - }}^ \circ = 0.535\mathrm{V}$, find the $\mathrm{pH}$ of ${\mathrm{H}^ + }/{\mathrm{H}_2}$ half-cell

$\mathrm{MnO}_4^ - + 8{\mathrm{H}^ + } + 5{e^ - } \to \mathrm{M}{\mathrm{n}^{2 + }} + 4{\mathrm{H}_2}\mathrm{O};{E^ \circ } = 1.51\mathrm{V}$

$\mathrm{Mn}{\mathrm{O}_4} + 4{\mathrm{H}^ + } + 2{e^ - } \to \mathrm{M}{\mathrm{n}^{2 + }} + 2{\mathrm{H}_2}\mathrm{O};{E^ \circ } = 1.23\mathrm{V}$

$E_{\mathrm{MnO}_4^ - |\mathrm{Mn}{\mathrm{O}_2}}^ \circ$ is

Consider the following cell

$\mathrm{Pt}\left|\mathrm{H}_2(\mathrm{g})\right| \mathrm{HCl}(\mathrm{aq}) \| \mathrm{Cl}^{-}(\mathrm{aq})|\mathrm{AgCl}(\mathrm{s})| \mathrm{Ag}(\mathrm{s})$

E.M.F of the above cell is 0.355 V at $25^{\circ} \mathrm{C}$ and 0.305 V at $30^{\circ} \mathrm{C}.$ The enthalpy change per mole of the reaction taking place at $25^{\circ} \mathrm{C}$ is

[Consider 1 mole of Ag is formed]

Assertion : For cell reaction; at $298 \mathrm{KZn}(\mathrm{s})+\mathrm{Cu}^{+2}(1 \mathrm{M}) \rightleftharpoons \mathrm{Zn}^{+2}(1 \mathrm{M})+\mathrm{Cu}(\mathrm{s})$ $\mathrm{E}_{\text {cell }}^{\circ}=0$

Reason: At standard condition ; cell potential is always zero.

Calculate equilibrium constant (approx) at $25^{\circ} \mathrm{C}$ for the cell reaction

$\mathrm{Fe}_{(\mathrm{s})}+\mathrm{CuSO}_4 \rightleftharpoons \mathrm{FeSO}_4+\mathrm{Cu}_{(\mathrm{s})}$
Given $\begin{aligned} & \mathrm{E}_{\mathrm{Fe} / \mathrm{F} \mathrm{c}^{+2}}^{\mathrm{o}}=+0.453 \text { volt } ; \ \mathrm{E}_{\mathrm{Cu} / \mathrm{Cu}^{+2}}^{\mathrm{o}}=-0.3435 \text { volt }\end{aligned}$

$\mathrm{C}{\mathrm{u}^ + }\left( {aq} \right)$ is unstable in solution and undergoes simultaneous oxidation and reduction, according to the reaction
$2\mathrm{C}{\mathrm{u}^ + }\left( {aq} \right)\leftrightharpoons\mathrm{C}{\mathrm{u}^{2 + }}\left( {aq} \right) + \mathrm{Cu}\left( s \right)$
choose correct ${E^ \circ }$ for the above reaction if
$E_{\mathrm{C}{\mathrm{u}^{2 + }}/\mathrm{Cu}}^ \circ = 0.34\mathrm{V},E_{\mathrm{C}{\mathrm{u}^{2 + }}/\mathrm{C}{\mathrm{u}^ + }}^ \circ = 0.15\mathrm{V}$

Statement-I: The voltage of mercury cell remains constant for long period of time.

Statement - II: The net cell reaction does not involve active species(ions) .

Specific conductance of $10^{-4} \mathrm{~M} \mathrm{n}$-Butyric acid aqueous solution is $1.9 \times 10^{-9} \mathrm{~S} \mathrm{~m}^{-1}.$ If molar conductance of $\text{n}$-Butyric acid at infinite dilution is $380 \times 10^{-} \mathrm{S} \mathrm{m}^2 \mathrm{~mol}^1,$ then $\mathrm{K}_\text{a}$ for $\mathrm{n}$-Butyric acid is:

In which case $\left(E_{\text {cell }}-E_{\text {cell }}^{\circ}\right)$ is zero

$\mathrm{Sn}^{4+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Sn}^{2+} E^{\circ}=0.13 \mathrm{~V}$

$\mathrm{Br}_2+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Br}^{-} \quad E^{\circ}=1.08 \mathrm{~V}$

Calculate $K_{\mathrm{eq}}$ for the cell formed by two electrodes

Consider the following reaction :

$2\mathrm{A}{\mathrm{g}^ + }{\mathrm{C}_6}{\mathrm{H}_{12}}{\mathrm{O}_6} + {\mathrm{H}_2}\mathrm{O} \to 2\mathrm{Ag}\left( s \right) + {\mathrm{C}_6}{\mathrm{H}_{12}}{\mathrm{O}_7} + 2{\mathrm{H}^ + }$

When ammonia is added to the solution, pH is raised to 11. Which half-cell reaction is affected by pH and by how much?

$\mathrm{A}{\mathrm{g}^ + } + {e^ - } \to \mathrm{Ag};E_{\mathrm{red}}^ \circ = 0.8\mathrm{V}$

${\mathrm{C}_6}{\mathrm{H}_{12}}{\mathrm{O}_6} + {\mathrm{H}_2}\mathrm{O} \to {\mathrm{C}_6}{\mathrm{H}_{12}}{\mathrm{O}_7} + 2{\mathrm{H}^ + } + 2{e^ - };E_{\mathrm{red}}^ \circ = - 0.05\mathrm{V}$

$\mathrm{Ag}\left( {\mathrm{N}{\mathrm{H}_3}} \right)_2^ + + {e^ - } \to \mathrm{Ag}\left( s \right) + \mathrm{N}{\mathrm{H}_3};E_{\mathrm{red}}^ \circ = 0.337\mathrm{V}$

Given, $\mathrm{E}^{\circ}=-0.27 \mathrm{~V}$ for the $\mathrm{Cl}^{-}\left|\mathrm{PbCl}_2\right| \mathrm{Pb}$ couple and $0.12 \mathrm{~V}$ for the $\mathrm{Pb}^{2+} \mid \mathrm{Pb}$ couple, determine $\mathrm{K}_{\mathrm{sp}}$ for $\mathrm{PbCl}_2$ at $25^{\circ} \mathrm{C}$ ? (Take $\frac{2.303 R T}{F}=0.06$ )

A $0.200 \mathrm{~M} \mathrm{~KOH}$ solution is electrolysed for $1.5 \mathrm{~h}$ using a current of $8.00 \mathrm{~A}$. How many moles of $\mathrm{O}_2$ were produced at the anode?

Small quantities of solutions of compounds $TX,TY$ and $TZ$ are put into separate test tubes containing X, Y and Z solutions. $TX$ does not react with any of these. $TY$ reacts with both X and Z. $TZ$ reacts with X. The decreasing order of ease of oxidation of the anions ${X^ - },{Y^ - },{Z^ - }$ is

How much charge (in $\mathrm{F}$) must flow through solution during electrolysis of aq. $\mathrm{Na}_2 \mathrm{SO}_4$ at $0^{\circ} \mathrm{C}$ and 1 atm to produce 33.6 L of product gases at $50 \%$ current efficiency?

The cell, $\mathrm{Zn}\left| {\mathrm{Z}{\mathrm{n}^{2 + }}\left( {1\mathrm{M}} \right)} \right|\left| {\mathrm{C}{\mathrm{u}^{2 + }}\left( {1\mathrm{M}} \right)} \right|\mathrm{Cu}$

($E_{\mathrm{cell}}^ \circ =$1.10 V), was allowed to be completely discharged at 298 K. The relative concentration of

$\mathrm{Z}{\mathrm{n}^{2 + }}$ to $\mathrm{C}{\mathrm{u}^{2 + }}\left( {\frac{{\left[ {\mathrm{Z}{\mathrm{n}^{2 + }}} \right]}}{{\left[ {\mathrm{C}{\mathrm{u}^{2 + }}} \right]}}} \right)$ is

Find the solubility product of a saturated solution of $\mathrm{Ag}_2 \mathrm{CrO}_4$ in water at $298 \mathrm{~K}$ if the emf of the cell $\mathrm{Ag} | \mathrm{Ag}^{+} \text{(satd. }\mathrm{Ag}_2 \mathrm{CrO}_4 \text{ soln.)}\| \mathrm{Ag}{ }^{+}(0.1 \mathrm{~M}) | \mathrm{Ag}$ is $0.164 \mathrm{~V}$ at $298 \mathrm{~K}.$

Which colourless gas evolves when $\mathrm{N}{\mathrm{H}_4}\mathrm{Cl}$ reacts with zinc in A dry cell battery?

A certain quantity of electricity is passed through aqueous solution of $\mathrm{AgN}{\mathrm{O}_3}$ and $\mathrm{CuS}{\mathrm{O}_4}$ connected in series. If Ag (at.wt.108) deposited at the cathode is 1.08 g then Cu deposited at the cathode is (at. wt. of Cu is 63.53 ):

If the ${H^ + }$ concentration is decreased from 1 M to ${10^{ - 4}}$ M at ${25^ \circ }C$  for the couple $MnO_4^ - /M{n^{(2 + )}}$, then the oxidising power of the $MnO_4^ - /M{n^{(2 + )}}$ couple decreases by

A $0.5 {M} {NaOH}$ solution offers a resistance of $31.6 {ohm}$ in a conductivity cell at room temperature. What shall be the approximate molar conductance of this ${NaOH}$ solution if cell constant of the cell is $0.367 {~cm}^{-1}$.