Electrochemistry

1- What mass of Cu(s) is electroplated by running 28.5A of current through a Cu 2+ (aq) solution for 4.00 h ?Express your answer to three significant figures and include the appropriate units.

2-How many minutes will it take to electroplate 60.1g of gold by running 5.00 A of current through a solution of Au + (aq) ?Express your answer to three significant figures and include the appropriate units.

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3-A Cu/Cu 2+ concentration cell has a voltage of 0.22V at 25 ∘ C . The concentration of Cu 2+ in one of the half-cells is 1.4×10−3M .What is the concentration of Cu 2+ in the other half-cell? Express your answer using two significant figures.

4- Calculate the equilibrium constant for each of the reactions at 25 ∘ C .O 2 (g)+2H 2 O(l)+2Cu(s)→4OH − (aq)+2Cu 2+ (aq) Express your answer using one significant figure.

5-Use tabulated electrode potentials to calculate ΔG ∘ rxn for each reaction at 25 ∘ C .2Fe 3+ (aq)+3Sn(s)→2Fe(s)+3Sn 2+ (aq) Express your answer using two significant figures.

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Electrochemistry Part II
1- What mass of Cu(s) is electroplated by running 28.5 A of current through a Cu 2+ (aq) solution for 4.00 h ?Express your answer to three significant figures and include the appropriate units.
Cu2+(aq) + 2e -> Cu(s)

Charge = ampere x second = 28.5 A * 1.44 x 10^4 s = 4.104 x 10^5 C
= 4.104 x 10^5 / 9.65 x 10^4 Faraday = 4.253 Faraday
As 1 mole corresponds to 2 Faradays here, number of moles = 4.253 / 2 = 2.13 moles
So the mass of Cu = number of moles x MW = 2.13 mol x 63.55 g /mol = 135.362 g

2-How many minutes will it take to electroplate 60.1g of gold by running 5.00 A of current through a solution of Au + (aq) ? Express your answer to three significant figures and include the appropriate units.
Au+ + e -> Au(s)

Moles of Au = mass / MW = 60.1 / 196.967 = 0.305
As 1 mole corresponds to 1 Faraday here, so current = 0.305 Faraday = 0.305 x 96500 C …

Solution Summary
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