Question

how much time is required to do the electroplating of Ag layer on a coffee tray(30cm x 15cm) to a thickness of 1mm using a constant current of 1.0 A. given that the density of Ag is 10.5 g/cm^3. a) 7720 s b)120 min c) 772 s d) 77.2 s

Answer 1

Answer:

mark as bianlist

Explanation:

Volume of Ag metal required to coat the coffee tray = 30 x 15 x 2 x 0.1 = 90 cm–3 Number of Faraday passed = 0.08 Number of Coulomb's passed = 0.08 x 96500 = 7720 C time = 7720/how-much-time-is-required-to-do-the-electroplating-of-ag-layer-on-a-coffee-tray-30-cm-x-15-cm

Answer 2

Answer:

7720 sec

Explanation:

Volume of plate=30*15*2*0.1=90cm^3

Then, mass used:-

Mass=volume/density

=>90/10.5=8.57gms

Gram equicalent weight if SILVER(Ag) used is =8.57/108=0.08

No.of Faradays=0.08

then time taken =nF

                        =0.08*96500

                        =965*8

                        =7720sec

Therefore, total time taken = 7720sec