Redox Reactions and Electrochemistry

The quantity of charge required to deposit one mole of copper from an aqueous solution of $\mathrm{CuSO_4}$ is ($F$ = one faraday)

500 mL of 0.2 M $\mathrm{MnO_4^-}$ solution in basic medium when mixed with 500 mL of 1.5 M KI solution, oxidises iodide ions to liberate molecular iodine. This liberated iodine is then titrated with a standard $x$ M thiosulphate solution in presence of starch till the end point. If 300 mL of thiosulphate was consumed, then the value of $x$ is ______.

Consider the following two half-cell reactions along with the standard reduction potential given : $\mathrm{CO_2+6H^++6e^-\rightarrow CH_3OH+H_2O}$, $E^\circ_{red}=0.02$ V; $\frac{1}{2}\mathrm{O_2+2H^++2e^-\rightarrow H_2O}$, $E^\circ_{red}=1.23$ V. A fuel cell was set up using the above two reactions such that the cell operates under the standard condition of 1 bar pressure and 298 K temperature. The fuel cell works with 80% efficiency. If the work derived from the cell using 1 mol of $\mathrm{CH_3OH}$ is used to compress an ideal gas isothermally against a constant pressure of 1 kPa, then the change in the volume of the gas, $\Delta V$ = ______ m$^3$. (nearest integer) Given : F = 96500 C mol$^{-1}$

In order to oxidise a mixture of 1 mole each of $\mathrm{FeC_2O_4}$, $\mathrm{Fe_2(C_2O_4)_3}$, $\mathrm{FeSO_4}$ and $\mathrm{Fe_2(SO_4)_3}$ in acidic medium, the number of moles of $\mathrm{KMnO_4}$ required is

An electrochemical cell, consist of the following two redox couples, $\mathrm{M^{x+}(aq)/M(s)}$ [$E^\ominus_{red}=+0.15$ V] and $\mathrm{Fe^{3+}(aq)/Fe(s)}$ [$E^\ominus_{red}=-0.036$ V]. The cell EMF ($E_{cell}$) is recorded to be 0.2057 V. If the reaction quotient of the electrochemical reaction is found to be $10^{-2}$, then the value of $x$ is ______. (Nearest integer) [Given: M is a p-block metal and $\frac{2.303RT}{F}=0.059$ V]

One half cell in a voltaic cell is constructed by dipping silver rod in $\mathrm{AgNO_3}$ solution of unknown concentration, other half cell is Zn rod dipped in 1 molar solution of $\mathrm{ZnSO_4}$. A voltage of 1.60 V is measured at 298 K for this cell. What is the concentration of $\mathrm{Ag^+}$ ions used in terms of $\log x$ ($x=[\mathrm{Ag^+}]$) ? $E^\ominus_{\mathrm{Zn^{2+}/Zn}}=-0.76$ V, $E^\ominus_{\mathrm{Ag^+/Ag}}=+0.80$ V, $\frac{2.303RT}{F}=0.059$ V

At 298 K, the molar conductivity of $x$% (w/w) MX solution (aqueous) is 123.5 S cm$^2$ mol$^{-1}$. The conductance of same solution is $1.9\times10^{-3}$ S. The value of $x$ is ______ $\times10^{-2}$. (Given: cell constant = 1.3 cm$^{-1}$; molar mass of MX is 75 g mol$^{-1}$, density of aqueous solution of MX at 298 K is 1.0 g mL$^{-1}$)

Given at 298 K : $E^\ominus_{\mathrm{Fe^{2+}/Fe}}=X$ Volt, $E^\ominus_{\mathrm{Fe^{3+}/Fe}}=Y$ Volt. The $E^\ominus_{\mathrm{Fe^{3+}/Fe^{2+}}}$ in Volt at 298 K is given by :

A solution of aluminium chloride is electrolysed for 30 minutes using a current of 2A. The amount of the aluminium deposited at the cathode is ______. [Given: molar mass of aluminium and chlorine are 27 g mol$^{-1}$ and 35.5 g mol$^{-1}$ respectively. Faraday constant $=96500$ C mol$^{-1}$]

Given below are two statements : Statement (I) : Corrosion is an electrochemical phenomenon in which pure metal acts as an anode and impure metal as a cathode. Statement (II) : The rate of corrosion is more in alkaline medium than in acidic medium. In the light of the above statements, choose the correct answer from the options given below :