Electrochemistry

The pressure of \ce{H2} required to make the potential of the \ce{H2} electrode zero in pure water at $298\ \text{K}$ is:

The molar conductivity of a $0.5\ \text{mol dm}^{-3}$ solution of \ce{AgNO3} with electrolytic conductivity of $5.76 \times 10^{-3}\ \text{S cm}^{-1}$ at $298\ \text{K}$ is

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

If the $E^\circ_{cell}$ for a given reaction has a negative value, which of the following gives the correct relationships for the values of $\Delta G^\circ$ and $K_{eq}$?

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)

Zinc can be coated on iron to produce galvanized iron but the reverse is not possible. It is because

In the electrochemical (Daniell) cell \ce{Zn | ZnSO4(0.01 M) || CuSO4(1.0 M) | Cu}, the emf is $E_1$. When the concentration of \ce{ZnSO4} is changed to $1.0\ \text{M}$ and that of \ce{CuSO4} to $0.01\ \text{M}$, the emf changes to $E_2$. Which one of the following is the correct relationship between $E_1$ and $E_2$? (Given $\frac{RT}{F} = 0.059$)

Consider the change in oxidation state of bromine corresponding to the standard reduction (electrode) potentials shown for each successive step: \ce{BrO4^-} $\xrightarrow{+1.82\ \text{V}}$ \ce{BrO3^-} $\xrightarrow{+1.50\ \text{V}}$ \ce{HBrO} $\xrightarrow{+1.595\ \text{V}}$ \ce{Br2} $\xrightarrow{+1.0652\ \text{V}}$ \ce{Br^-} Then the species undergoing disproportionation is

For the cell reaction \ce{2Fe^3+(aq) + 2I-(aq) -> 2Fe^2+(aq) + I2(aq)}, $E^\circ_{cell} = 0.24$ V at $298$ K. The standard Gibbs energy ($\Delta_r G^\circ$) of the cell reaction is:

For a cell involving one electron $E^\circ_{cell} = 0.59$ V at $298$ K, the equilibrium constant for the cell reaction is: (Given that $\frac{2.303\,RT}{F} = 0.059$ V at $T = 298$ K)

Following limiting molar conductivities are given as $\Lambda_m^\circ(\ce{H2SO4}) = x\ \text{S cm}^2\,\text{mol}^{-1}$ $\Lambda_m^\circ(\ce{K2SO4}) = y\ \text{S cm}^2\,\text{mol}^{-1}$ $\Lambda_m^\circ(\ce{CH3COOK}) = z\ \text{S cm}^2\,\text{mol}^{-1}$ $\Lambda_m^\circ$ (in $\text{S cm}^2\,\text{mol}^{-1}$) for \ce{CH3COOH} will be :

The standard electrode potential ($E^\circ$) values of \ce{Al^3+/Al}, \ce{Ag+/Ag}, \ce{K+/K} and \ce{Cr^3+/Cr} are $-1.66\ \text{V}$, $0.80\ \text{V}$, $-2.93\ \text{V}$ and $-0.74\ \text{V}$, respectively. The correct decreasing order of the reducing power of the metals is:

On electrolysis of dil. sulphuric acid using platinum (\ce{Pt}) electrodes, the product obtained at the anode will be:

The number of Faradays ($F$) required to produce $20$ g of calcium from molten \ce{CaCl2} (Atomic mass of \ce{Ca} $= 40$ g mol$^{-1}$) is:

The molar conductances of \ce{NaCl}, \ce{HCl} and \ce{CH3COONa} at infinite dilution are $126.45$, $426.16$ and $91.0\ \text{S cm}^2\,\text{mol}^{-1}$ respectively. The molar conductance of \ce{CH3COOH} at infinite dilution is. Choose the right option for your answer.

The molar conductivity of $0.007\ \text{M}$ acetic acid is $20\ \text{S cm}^2\,\text{mol}^{-1}$. What is the dissociation constant of acetic acid? Choose the correct option. $\lambda^\circ(\ce{H+}) = 350\ \text{S cm}^2\,\text{mol}^{-1}$, $\lambda^\circ(\ce{CH3COO-}) = 50\ \text{S cm}^2\,\text{mol}^{-1}$

Given below are the half-cell reactions: \ce{MnO4^- + 8H+ + 5e- -> Mn^2+ + 4H2O}, $E^\circ_{\ce{MnO4^-/Mn^2+}} = -1.510\ \text{V}$ \ce{1/2 O2 + 2H+ + 2e- -> H2O}, $E^\circ_{\ce{O2/H2O}} = +1.223\ \text{V}$ Will the permanganate ion, \ce{MnO4^-}, liberate \ce{O2} from water in the presence of an acid?

At 298 K, the standard electrode potentials of \ce{Cu^2+/Cu}, \ce{Zn^2+/Zn}, \ce{Fe^2+/Fe} and \ce{Ag+/Ag} are $0.34\ \text{V}$, $-0.76\ \text{V}$, $-0.44\ \text{V}$ and $0.80\ \text{V}$, respectively. On the basis of standard electrode potential, predict which of the following reactions cannot occur?

Find the emf of the cell reaction \ce{Ni(s) + 2Ag+(0.001 M) -> Ni^2+(0.001 M) + 2Ag(s)} (Given $E^\circ_{cell} = 1.05\ \text{V}$, $\frac{2.303RT}{F} = 0.059$ at $298\ \text{K}$)

Molar conductance of an electrolyte increases with dilution according to the equation: $\Lambda_m = \Lambda_m^\circ - A\sqrt{c}$. Which of the following statements are true? (A) This equation applies to both strong and weak electrolytes. (B) Value of the constant $A$ depends upon the nature of the solvent. (C) Value of constant $A$ is same for both \ce{BaCl2} and \ce{MgSO4}. (D) Value of constant $A$ is same for both \ce{BaCl2} and \ce{Mg(OH)2}.

The correct value of cell potential (in volt) for the reaction that occurs when the following two half cells are connected is: \ce{Fe^2+(aq) + 2e- -> Fe(s)}, $E^\circ = -0.44\ \text{V}$ \ce{Cr2O7^2- + 14H+ + 6e- -> 2Cr^3+ + 7H2O}, $E^\circ = +1.33\ \text{V}$

The $E^\circ$ values are given as: $E^\circ_{\ce{Al+/Al}} = +0.55\ \text{V}$ and $E^\circ_{\ce{Tl+/Tl}} = -0.34\ \text{V}$; $E^\circ_{\ce{Al^3+/Al}} = -1.66\ \text{V}$ and $E^\circ_{\ce{Tl^3+/Tl}} = +1.26\ \text{V}$. Identify the incorrect statement.

The correct option for a redox couple is

The conductivity of centimolar solution of \ce{KCl} at $25\,^\circ\text{C}$ is $0.0210\ \text{ohm}^{-1}\text{cm}^{-1}$ and the resistance of the cell containing the solution at $25\,^\circ\text{C}$ is $60\ \text{ohm}$. The value of cell constant is

Match List-I with List-II: List-I (Conversion) (A) $1$ mol of \ce{H2O} to \ce{O2} (B) $1$ mol of \ce{MnO4-} to \ce{Mn^2+} (C) $1.5$ mol of \ce{Ca} from molten \ce{CaCl2} (D) $1$ mol of \ce{FeO} to \ce{Fe2O3} List-II (No. of Faraday required) (i) $3\,F$ (ii) $2\,F$ (iii) $1\,F$ (iv) $5\,F$ Choose the correct match:

Mass in grams of copper deposited by passing $9.6487$ A current through a voltmeter containing copper sulphate solution for $100$ seconds is: (Given: Molar mass of \ce{Cu} $= 63$ g mol$^{-1}$, $1\,F = 96487$ C)

If the molar conductivity ($\Lambda_m$) of a $0.050\ \text{mol L}^{-1}$ solution of a monobasic weak acid is $90\ \text{S cm}^2\,\text{mol}^{-1}$, its extent (degree) of dissociation will be [Assume $\lambda^\circ_+ = 349.6\ \text{S cm}^2\,\text{mol}^{-1}$ and $\lambda^\circ_- = 50.4\ \text{S cm}^2\,\text{mol}^{-1}$]

The standard electrode potential ($E^\circ$) for the half-cell reaction $\ce{Fe^{3+} + e^- -> Fe^{2+}}$ at 298 K is: [Given $E^\circ(\ce{Fe^{3+}/Fe})=-0.04\,$V and $E^\circ(\ce{Fe^{2+}/Fe})=-0.44\,$V]

A solution of copper sulphate is electrolysed for $10$ minutes with a current of $1.5$ amperes. The mass of copper deposited at the cathode is: (Given: Molar mass of \ce{Cu} $= 63$ g mol$^{-1}$; $1\,F = 96487$ C mol$^{-1}$)

Calculate the emf of the half cell given below: \ce{Pt(s) | H2(g, 2 atm) | HCl(aq, 0.02 M)} $E^\circ_{\ce{H+/H2}} = 0\ \text{V}$ (Given: $\frac{2.303RT}{F} = 0.059$, $\log 2 = 0.3010$)

For a strong electrolyte salt $XY$, the plot of $\Lambda_m$ versus $\sqrt{c}$ has slope $-90.0\,\text{S cm}^2\text{mol}^{-3/2}\text{L}^{1/2}$ at 298 K. At $0.01\,$M, $\Lambda_m=145.0\,\text{S cm}^2\text{mol}^{-1}$. The limiting molar conductivity of $Y^-$ ion, $\lambda^0_{Y^-}$ (in S cm$^2$mol$^{-1}$), is: [Given $\lambda^0_{X^+}=74.0\,\text{S cm}^2\text{mol}^{-1}$]