Electrochemistry Lab Help

I have to do the following calculations and I don't know how to do them. This is the data:
mass of copper anode before electrolysis= 23.0425g
mass of copper anode after electrolysis= 22.9041g
time of electrolysis= 603 seconds
current= 0.73 ampere
Volume of gas collected in the burette = 54.50mL
Atmospheric pressure(corrected to standard conditions)= 758.0 mmHg
Room Temperature= 20.0 C(celsius)
Vapor pressure of water at 20.0 C= 17.5 mmHg

a.) Write the balanced equation for the reaction taking place in this electrolysis
b.) Calculate the ratio of the experimental moles of hydrogen produced to moles of copper consumed and compare this ratio to the theoretical ratio expected from the balanced chemical equation.
c.) Calculate the experimental value for the Faraday(the amount of charge carried by one mole of electrons) and compare it to the accepted value by calculating the percentage error.
d.) Calculate the experimental value of Avagadro's number using your experimentally determined value for the Faraday and the charge of one electron (1.602^-19 Coulomd).

Any help would be appreciated..even just one calculation. I just want to know how to do these calculations, Thankyou.

I believe this electrochemical reaction refers to electrolysis of a solution of strong acid with a Copper Anode and a Platinum cathode.

At the Anode, Copper is oxidized to Cu2+ and goes into solution
At the Cathode, H+ is reduced to H2 and this is the gas

The anodic rxn: Cu---------------->>Cu2+ + 2e
cathodic rxn : 2H + 2e --------------->>H2
So overall reaction : Cu + 2H+ ---------------------------> Cu2+ + H2

n(H2) = pV/RT where n = # of moles of H2......................................…
partial Pressure of H2 + partial Pressure H2O = 758 so you subtract pressure of water from atmospheric to get partial pressure of H2
p = (758 - 17.5)/760 atm; V = 0.0545 L; T = 273 + 20 = 293 K