Electrochemistry · Chemistry · AP EAPCET

In a cell a copper electrode was used as a cathode. What is the electrode potential (in V) of the copper electrode dipped in $0.1 \mathrm{M} \mathrm{Cu}^{2+}$ solution at 298 K ?

$$ \left(E_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\ominus}=0.34 \mathrm{~V} ; \frac{2.303 R T}{F}=0.06 \mathrm{~V}\right) $$

Observe the following statements about dry cell

I. It is a primary battery.

II. Zinc vessel acts as cathode.

III. A paste of moist $\mathrm{NH}_4 \mathrm{Cl}, \mathrm{MnO}_2$ and $\mathrm{ZnCl}_2$, is present between two electrodes

IV. The potential of this cell is 1.5 V .

The correct statements are

$$ \text { Match the following } $$

$$ \begin{array}{llll} \hline & \begin{array}{l} \text { List-I (Symbol of } \\ \text { electrical property) } \end{array} & & \text { List-I (Units) } \\ \hline \text { (A) } & \Lambda_{\mathrm{m}} & \text { (I) } & \mathrm{Scm}^{-1} \\ \hline \text { (B) } & \mathrm{G} & \text { (II) } & \mathrm{m}^{-1} \\ \hline \text { (C) } & \mathrm{K} & \text { (III) } & \mathrm{Scm}^2 \mathrm{~mol}^{-1} \\ \hline \text { (D) } & \mathrm{G}^* & \text { (IV) } & \mathrm{S} \\ \hline \end{array} $$

The correct answer is

Consider the following cell reaction

$$ 2 \mathrm{Fe}^{3+}(a q)+2 \mathrm{I}^{-}(a q) \rightleftharpoons 2 \mathrm{Fe}^{2+}(a q)+\mathrm{I}_2(s) $$

At 298 K , the cell emf is 0.237 V . The equilibrium constant for the reaction is $10^x$. The value of $x$ is $\left(F=96500 \mathrm{C} \mathrm{mol}^{-1} ; R=8.3 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}\right)$.

When the lead storage battery is in use (during discharge) the reaction that occurs at the anode is

$$ \text { Match the following } $$

$$ \begin{array}{cccc} \hline & \begin{array}{c} \text { List-I } \\ \text { (Transition metal, M) } \end{array} & & \begin{array}{c} \text { List-II } \\ \left(E_{M^{2+} / M}^{\ominus}\right) \end{array} \\ \hline \text { (A) } & \mathrm{Ni} & \text { (I) } & -1.18 \\ \hline \text { (B) } & \mathrm{Mn} & \text { (II) } & -0.91 \\ \hline \text { (C) } & \mathrm{Fe} & \text { (III) } & -0.25 \\ \hline \text { (D) } & \mathrm{Cr} & \text { (IV) } & -0.44 \\ \hline \end{array} $$

The correct answer is

At 298 K , the following reaction takes place for a cell at the hydrogen electrode

$$ \mathrm{H}^{+}(a q)+e^{-} \longrightarrow \frac{1}{2} \mathrm{H}_2 \text { (1 bar) } $$

The solution pH is 10.0 . What is the hydrogen electrode potential in volts?

$$ \left(\frac{2303 R T}{F}=0.06 \mathrm{~V}\right) $$

The resistance of a conductivity cell filled with 0.1 M KCl solution is $100 \Omega$. If the resistance of the same cell when filled with 0.2 M KCl solution is $520 \Omega$, the molar conductivity of 0.02 M solution (in $\mathrm{S} \mathrm{cm}^2 \mathrm{~mol}^{-1}$ ) is (Given: conductivity of 0.1 M KCl solution $=1.29 \mathrm{Sm}^{-1}$ )

For which of the following the $E^{\ominus}\left(M^{3+} / M^{2+}\right)$ is negative?

If $E_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.441 \mathrm{~V}$ and $E_{\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}}^{\circ}=0.771 \mathrm{~V}$, the standard emf of the cell reaction $\mathrm{Fe}(s)+2 \mathrm{Fe}^{3+}(a q) \longrightarrow 3 \mathrm{Fe}^{2+}(a q)$ is

The specific conductance of 0.05 M NaOH solution is $0.0115 \mathrm{~S} \mathrm{~cm}^{-1}$ What is its molar conductance ( $\wedge_{\mathrm{m}}$ ) in $\mathrm{Scm}^2 \mathrm{~mol}^{-1}$ ?

What is $E_{\text {cell }}$ (in V) of the following cell at $298 \mathrm{~K} ?$

$$ \begin{aligned} & \left(E_{\mathrm{Zn}^{2+} / \mathrm{Zn}}^{\ominus}=-0.76 \mathrm{~V} ; E_{\mathrm{Ni}^{2+} / \mathrm{Ni}}^{\ominus}=-0.25 \mathrm{~V} ; \frac{2.303 R T}{F}=0.06 \mathrm{~V}\right) \\ & 1(s) \mathrm{Zn}^{2+}(0.01 \mathrm{M}) \mathrm{Ni}^{2+}(0.1 \mathrm{M}) \mathrm{Ni}(s \end{aligned} $$

At 300 K , the $E_{\text {cell }}^{\ominus}$ of

$$ A(s)+B^{2+}(a q) \rightleftharpoons A^{2+}(a q)+B(s) $$

is 1.0 V . If $\Delta_r S^\theta$ of this reaction is $100 \mathrm{JK}^{-1}$. What is $\Delta_r H^{\ominus}$ (in $\mathrm{kJ} \mathrm{mol}^{-1}$ ) of this reaction?

$$ \left(\mathrm{F}=96500 \mathrm{C} \mathrm{~mol}^{-1}\right) $$

Consider the cell reaction at 300 K .

$$ A(s)+B^{2+}(a q) \rightleftharpoons A^{2+}(a q)+B(s) $$

Its $E^{\ominus}$ is 1.0 V . The $\Delta_r H^{\ominus}$ of the reaction is $-163 \mathrm{kJmol}^{-1}$.

What is $\Delta_r s^{\ominus}$ (in $\mathrm{JK}^{-1}$ ) of the reaction?

$$ \left(F=96500 \mathrm{C} \mathrm{~mol}^{-1}\right) $$

In which of the following Galvanic cells emf is maximum?

(Given, $E_{\mathrm{Mg}^{2+} \mid \mathrm{Mg}}^{\circ}=-2.36 \mathrm{~V}$

and $E_{\mathrm{Cl}_2 \mid 2 \mathrm{Cl}^{-}}^{\circ}=+136 \mathrm{~V}$ )

Consider the following standard electrode potentials ( $E^{\circ}$ in volts) in aqueous solution.

$$ \begin{array}{|c|c|c|} \hline \text { Element } & M^{3+} / M & M^{+} / M \\ \hline \mathrm{Al} & -1.66 & +0.55 \\ \hline \mathrm{TI} & +1.26 & -0.34 \\ \hline \end{array} $$

Based on this data. which of the following statements is correct?

The standard reduction potentials of $2 \mathrm{H}^{+} / \mathrm{H}_2, \mathrm{Cu}^{2+} / \mathrm{Cu}, \mathrm{Zn}^{2+} / \mathrm{Zn}$ and $\mathrm{NO}_3^{-}, \mathrm{H}^{-} / \mathrm{NO}$ are 0.0 0.34 . -0.76 and 0.97 V respectively. Identify the correct statements from the following.

I. $\mathrm{H}^{+}$does not oxidise Cu to $\mathrm{Cu}^{2+}$

II. Zn reduces $\mathrm{Cu}^{2+}$ to Cu

III. $\mathrm{NO}_3^{-}$oxidises Cu to $\mathrm{Cu}^{2+}$

The standard reduction potentials of $2 \mathrm{H}^{+} / \mathrm{H}_2, \mathrm{Cu}^{2+} / \mathrm{Cu}, \mathrm{Zn}^{2+} / \mathrm{Zn}$ and $\mathrm{NO}_3^{-}, \mathrm{H}^{-} / \mathrm{NO}$ are 0.0 , +0.34 . -0.76 and 0.97 V respectively. Observe the following reactions

I. $\mathrm{Zn}+\mathrm{HCI} \rightarrow$

II. $\mathrm{Cu}+\mathrm{HCl} \rightarrow$

III. $\mathrm{Cu}+\mathrm{HNO}_3 \rightarrow$

Which reactions does not liberate $\mathrm{H}_2(g)$ ?

Aqueous $\mathrm{CuSO}_4$ solution was electrolysed by passing 2 amp of current for 10 min . What is the weight (in g) of copper deposited at cathode ?

$$ \left(\mathrm{Cu}=63 \mathrm{u} ; F=96500 \mathrm{C} \mathrm{~mol}^{-1}\right) $$

The $E^{-}$of $M\left|M^{2+} \| \mathrm{Cu}^{2+}\right| \mathrm{Cu}$ is 0.3 V .

At what concentration of $\mathrm{Cu}^{2+}\left(\mathrm{in} \mathrm{mol} \mathrm{L} \mathrm{L}^{-1}\right)$, the $\mathrm{E}_{\mathrm{cel}}$ value becomes zero ? $\left(\frac{2.303 R T}{F}=0.06\right)$

(Conc. of $\mathrm{M}^{2+}=0.1 \mathrm{M}$ )

96.5 amperes current is passed through the molten $$\mathrm{AlCl}_3$$ for 100 seconds. The mass of aluminium deposited at the cathode is (atomic weight of $$\mathrm{Al}=27 \mathrm{u}$$)

38.6 amperes of current is passed for 100 seconds through an aqueous $$\mathrm{CuSO}_4$$ solution using platinum electrodes. The mass of copper consumed from the solution and volume of gas liberated at STP are respectively (molar mass of $$\mathrm{Cu}=63.54 \mathrm{~g} \mathrm{~mol}^{-1}$$).

The reduction potential of hydrogen electrode at $$25^{\circ} \mathrm{C}$$ in a neutral solution is ($$p_{\mathrm{H}_2}=1$$ bar)

In the electrolysis of a CuSO$$_4$$ solution, how many grames of Cu are plated out on the cathode, in the time that is required to liberate 5.6 L of O$$_2$$(g), measured at 1 atm and 273 K, at the anode?

If hydrogen electrons dipped in two solutions of pH = 3 and pH = 6 are connected by a salt bridge, the emf of the resulting cell is

At $$291 \mathrm{~K}$$, saturated solution of $$\mathrm{BaSO}_4$$ was found to have a specific conductivity of $$3.648 \times 10^{-6} \mathrm{ohm}^{-1} \mathrm{~cm}^{-1}$$ and that of water being used is $$1.25 \times 10^{-6} \mathrm{ohm}^{-1} \mathrm{~cm}^{-1}$$. If the ionic conductances of $$\mathrm{Ba}^{2+}$$ and $$\mathrm{SO}_4^{2-}$$ are 110 and $$136.6 \mathrm{ohm}^{-1} \mathrm{~cm}^2 \mathrm{~mol}^{-1}$$ respectively. The solubility of $$\mathrm{BaSO}_4$$ at $$291 \mathrm{~K}$$ will be [Atomic masses of $$\mathrm{Ba}=137, \mathrm{~S}=32, \mathrm{O}=16]$$

Find the emf of the following cell reaction. Given, $$E_{\mathrm{Cr}^{3+} / \mathrm{Cr}^{2+}}^{\Upsilon}=-0.72 \mathrm{~V}$$ and $$E_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\Upsilon}= -0.42 \mathrm{~V}$$ at $$25^{\circ} \mathrm{C}$$ is $$\mathrm{Cr}\left|\mathrm{Cr}^{3+}(0.1 \mathrm{M})\right| \mid \mathrm{Fe}^{2+} (0.1 \mathrm{M}) \mid \mathrm{Fe}$$

For $$\mathrm{C{r_2}O_7^{2 - } + 14{H^ + } + 6{e^ - }\buildrel {Yields} \over \longrightarrow 2C{r^{3 + }} + 7{H_2}O,{E^\Upsilon } = 1.33}$$ V at $$[C{r_2}O_7^{2 - }] = 4.5$$ millimole, $$[C{r^{3 + }}] = 1.5$$ millimole and $$E = 1.067$$ V, then calculate the pH of the solution.

Assertion (A) Sodium acetate on Kolbe’s electrolysis gives ethane.

Reason (R) Methyl free radical is formed at cathode.

When a current of 10 A is passes through molten AlCl$$_3$$ for 1.608 minutes. The mass of Al deposited will be

[Atomic mass of Al = 27 g]

The molar conductivities $$\left(\lambda_{\mathrm{m}}^{\Upsilon}\right)$$ at infinite dilution of $$\mathrm{KBr}, \mathrm{HBr}$$ and $$\mathrm{KNH}_2$$ are 120.5, 420.6 and $$90.48 \mathrm{~S} \mathrm{~cm}^2 \mathrm{~mol}^{-1}$$ respectively. Find the value of $$\lambda_{\mathrm{m}}^\Upsilon$$ for $$\mathrm{NH}_3$$.