Question

4. One side of an aluminum electrode is exposed to a solution of NiCl2. The electrode hasdimensions of 40.0 mm by 200.0 mm. Assuming even electrodeposition, how long, in minutes, would it take to plate a 20.0 μm layer of Ni on the electrode with a current of 2.00 A? Ni has a density of 8.912 g/cm3.

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5. Consider a concentration cell at 25 °C using the following half reaction:

Cu2+(aq) + 2e– → Cu(s) Two copper electrodes, each weighing 20.0 g, and containing 1.00 L of 0.500 M Cu2+ solutionand 1.00 L of 0.150 M Cu2+ solution, respectively, are assembled with two beakers, wire, and asalt bridge. What will the mass of the electrode at the anode be when the cell potential is 0.00750 V?

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6. Consider the cell described below at 298 K:

Co | Co2+ (2.00 M) || Ni2+ (0.500 M) | Ni

Given the anodic and cathodic compartments have equal volumes, what will the concentration of Ni2+ be when the cell is dead?

thanks for help！！

Answer 1

4)
lets find the volume in Ni deposited
V = L*B*H
= 40.0 mm * 200.0 mm * 20.0 μm
= 4 cm * 20 cm *0.002 cm
= 0.16 cm^3

mass = density * volume
= 8.912 g/cm^3 * 0.16 cm^3
= 1.426 g

mol of Ni = mass of Ni / molar mass of Ni
= 1.426 / 58.7 g/mol
= 0.0243 mol

Ni2+   +   2e-    ------> Ni

1 mol of Ni requires 2 mol of electron
1 mol of electron = 96485 C
So,1 mol of Ni requires 192970 C

charge required = number of mol * charge for 1 mol
= 0.0243 mol * 192970 C/mol
= 4688 C

time = Charge/current
= 4688 C / 2.00 A
= 2344 S
= 2344/60
=39.1 minutes

Answer: 39.1 minutes

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