‌E‌cell ‌=E°‌cell ‌−‌

0.059

n

‌log‌

[Mg2+]

[Cu2+]

‌E‌cell ‌=2.71−‌

0.059

2

‌log‌

[0.001]

[0.01]

‌E‌cell ‌=2.71−(−0.0295)=2.74V
(i) When external opposite applied voltage is less than 2.71, it is less than E°‌cell ‌, therefore, the electrons will flow from the anode to the cathode, and current will flow from cathode (copper electrode) to anode (magnesium electrode).

(ii) When external opposite applied potential is greater than 2.71, it is greater than E°‌cell ‌, therefore, the reaction will be reversed, and the current will flow from anode to cathode.
OR
(a) Charge required to deposite 2.8g‌Fe :

mol‌ ‌Fe=‌

‌ mass ‌

‌ molar mass ‌

=‌

2.8g

56g⋅mol−1

=0.05‌mol
2F charge is required to discharge 1‌mol of Fe2+ ions as Fe, therefore deposition of 0.05‌mol‌Fe will need
0.05×2=0.1F=0.1F ×‌

96500C

F

=9650C
The quantity of charge is related to current as
Q=It
Therefore, the time needed to deposit 2.8g Fe is:

t=‌

Q

I

=‌

9650C

2A

=4825 s
So, the current flowed through the cells for 4825 seconds.

The amount of Zn deposited in cell Y can be calculated using Faraday's second law:

‌‌

‌ mass of ‌Zn

‌ mass of ‌Fe

=‌

‌ Eq.wt of ‌Zn

‌ Eq wt of ‌Fe

‌‌ mass of ‌Zn=2.8g×‌

65.3g∕2

56g∕2

=3.265g≈3.3g
Therefore, the mass of Zn deposited in cell Y in the same time is 3.3g.

(b) (i) Molar conductivity of strong electrolytes increases linearly as the square root of the concentration decreases; therefore, electrolyte A is a strong electrolyte. Molar conductivity of weak electrolytes increases non-linearly as square root of concentration decreases; therefore, electrolyte B is a weak electrolyte.
(ii) As concentration of strong electrolyte approaches zero, the molar conductivity of the plot intercepts the molar conductivity axis, giving the limiting value of molar conductivity Em0. The plot of molar conductivity of weak electrolyte tends to infinity as its concentration approaches zero; it does not intersect the molar conductivity axis.