Given $E_{Fe^{3+}/Fe^{2+}}^{\circ} = +0.76 \ V$ and $E_{I_{2}/I^{-}}^{\circ} = +0.55 \ V$. The equilibrium constant for the reaction taking place in a galvanic cell consisting of the above two electrodes is (Given $\frac{2.303 \ RT}{F} = 0.06 \ V$)

$(i)$ When a copper rod is dipped in $1 \ M$ silver nitrate solution,silver crystals are obtained. $(ii)$ Silver metal does not react with $1 \ M$ zinc nitrate solution. $(iii)$ When zinc metal is dissolved in $1 \ M$ copper sulfate solution,copper metal is displaced. Arrange the three metals in the decreasing order of their reducing power.

Redox reactions play a pivotal role in chemistry and biology. The values of standard redox potential $(E^{\circ})$ of two half-cell reactions decide which way the reaction is expected to proceed. $A$ simple example is a Daniel cell in which zinc goes into solution and copper gets deposited. Given below are a set of half-cell reactions (acidic medium) along with their $E^{\circ}$ ($V$ with respect to normal hydrogen electrode) values.
$I_2 + 2e^{-} \rightarrow 2I^{-} \quad E^{\circ} = 0.54 \ V$
$Cl_2 + 2e^{-} \rightarrow 2Cl^{-} \quad E^{\circ} = 1.36 \ V$
$Mn^{3+} + e^{-} \rightarrow Mn^{2+} \quad E^{\circ} = 1.50 \ V$
$Fe^{3+} + e^{-} \rightarrow Fe^{2+} \quad E^{\circ} = 0.77 \ V$
$O_2 + 4H^{+} + 4e^{-} \rightarrow 2H_2O \quad E^{\circ} = 1.23 \ V$
$1.$ Among the following,identify the correct statement.
$(A)$ Chloride ion is oxidized by $O_2$
$(B)$ $Fe^{2+}$ is oxidized by iodine
$(C)$ Iodide ion is oxidized by chlorine
$(D)$ $Mn^{2+}$ is oxidized by chlorine
$2.$ While $Fe^{3+}$ is stable,$Mn^{3+}$ is not stable in acid solution because
$(A)$ $O_2$ oxidizes $Mn^{2+}$ to $Mn^{3+}$
$(B)$ $O_2$ oxidizes both $Mn^{2+}$ and $Fe^{2+}$ to $Fe^{3+}$
$(C)$ $Fe^{3+}$ oxidizes $H_2O$ to $O_2$
$(D)$ $Mn^{3+}$ oxidizes $H_2O$ to $O_2$
$3.$ Sodium fusion extract,obtained from aniline,on treatment with iron$(II)$ sulphate and $H_2SO_4$ in presence of air gives a Prussian blue precipitate. The blue color is due to the formation of
$(A)$ $Fe_4[Fe(CN)_6]_3$
$(B)$ $Fe_3[Fe(CN)_6]_2$
$(C)$ $Fe_4[Fe(CN)_6]_2$
$(D)$ $Fe_3[Fe(CN)_6]_3$
Give the answer for questions $1, 2$ and $3.$

The rusting of iron takes place as follows. Calculate $\Delta G^o$ for the net process in $kJ \ mol^{-1}$.
$2H^{+} + 2e^- + \frac{1}{2}O_2 \longrightarrow H_2O_{(l)} ; E^o = +1.23 \ V$
$Fe^{2+} + 2e^- \longrightarrow Fe_{(s)} ; E^o = -0.44 \ V$

On the basis of standard electrode potential values $(E^{\circ})$, suggest which of the following reactions would take place spontaneously? Given: $E^{\circ}_{Zn^{2+}/Zn} = -0.76 \ V$, $E^{\circ}_{Cu^{2+}/Cu} = +0.34 \ V$, $E^{\circ}_{Mg^{2+}/Mg} = -2.37 \ V$, $E^{\circ}_{Fe^{2+}/Fe} = -0.44 \ V$, $E^{\circ}_{Br_{2}/Br^{-}} = +1.09 \ V$, $E^{\circ}_{Cl_{2}/Cl^{-}} = +1.36 \ V$, $E^{\circ}_{Cd^{2+}/Cd} = -0.40 \ V$.

The solubility product of $AgCl$ under standard conditions of temperature is given by