Electrochemistry

The cell reaction of a cell is $\mathrm{Mg}\left(\mathrm{s}\right) + {\mathrm{Cu}}^{2+}\left(\mathrm{aq}\right) \to \mathrm{Cu}\left(\mathrm{s}\right) + {\mathrm{Mg}}^{2+}\left(\mathrm{aq}\right).$ If the standard reduction potentials of Mg and Cu are - 2.37 and + 0.34V respectively. The emf of the cell is

• 2.03V

• -2.03V

• +2.71V

• -2.71V

In the electrochemical reaction, $2{\mathrm{Fe}}^{3+}+ \mathrm{Zn} \to {\mathrm{Zn}}^{2+}+ 2{\mathrm{Fe}}^{2+}$ increasing the concentration of Fe²⁺

• increases cell emf

• increases the current flow

• decrease the cell emf

• alter the pH of the solution

Distribution law was given by :

• Henry

• vant Hoff

• Nernst

• Ostwald

When Cu reacts with AgNO₃ solution, the reaction takes place is:

• oxidation of Cu

• reduction of Cu

• oxidation of Ag

• reduction of NO${}_3^{-}$

For a cell involving one electron E°_cell = 0.59 V at 298 K, the equilibrium constant for the cell reaction is :[ Given that $\frac{2.303}{\mathrm{F}}= 0.059\mathrm{V} \mathrm{at} \mathrm{T} = 298 \mathrm{K}]$.

• 1.0 × 10⁵

• 1.0 × 10¹⁰

• 1.0 × 10³⁰

• 1.0 × 10²

For the cell reaction

2Fe³⁺ (aq) +2l^-(aq) → 2Fe²⁺(aq) + I₂(aq)

E${}_{\mathrm{cell}}^{⊝}$= 0.24V at 298 K. The standard Gibbs energy (Δ_rG^Θ) of the cell reaction is:

[Given that Faraday constant F = 96500 C mol^–1]

• – 23.16 kJ mol^–1

• 46.32 kJ mol^-1

• 23.16 kJ mol^-1

• - 46.32 kJ mol^-1