Question

A thin layer of gold can be applied to another material by an electrolytic process. The surface area of an object to be gold plated is 49.7 cm2 and the density of gold is 19.3g/cm3. A current of 3.30 A is applied to a solution that contains gold in the +3 oxidation state.

Part A

Calculate the time required to deposit an even layer of gold 1.20×10⁻³ cm thick on the object.

t =

nothing

s

Answer 1

First we calculate the volume of gold to be deposited.

Volume = Surface area*Thickness = 49.7*1.2*(10^-3) cm³ = 5.964*(10^-2) cm^3

Next we calculate the mass of gold required.

Mass = Density*Volume = 19.3*5.964*(10^-2) = 1.151 g

Next we calculate moles of Au in this much mass.

Moles = Mass/MW = 1.151/197 = 5.84*(10^-3) moles

Now, 1 mole of Au will be deposited by using 3 moles of e-.

So,

Moles of e- required = 3*5.84*(10^-3) = 1.752*(10^-2) mol e-

Now, 1 mole e- carries 96500C charge.

So, charge carried by these many moles of e- = 1.752*(10^-2)*96500 = 1691 C

Now, in order to get this much charge, the time for which current should be applied is calculated by:

Time = Charge/Current = 1691/3.30 = 512.42 s