How many seconds are required to produce 4.94 mg of chromium metal from an acidic solution of potassium dichromate, using a current of 0.234 a?

The half-reaction is as followed: Cr2O72â’(aq) + 14H+ + 6e⒠→ 2Cr3+(aq) + 7H2O From the half-reaction, you can retrieve the following information: 1 mole of potassium dichromate =6 moles of e^- 6 moles of e^-=2 moles of Cr You will also need the following information: 1 mole of e^-=96,485 C and 1 mole of Cr=52.00g Calculate the number of moles that 4.94mg equates too: 4.94 mg=4.94 x 10^-3g of chromium*(1 mol/52.00g)=9.50 x 10^-5 mole of Cr How many moles of electrons are need to produce 9.50 x 10^-5 mole of Cr? Solve for moles of electrons: 9.50 x 10^-5 mole of Cr*(6 moles of e^-/2 mole of Cr)=2.85 x 10^-4 moles of e^- Whats the charge of 2.85 x 10^-4 moles of electrons? Use Faraday's constant: 2.85 x 10^-4 moles of e^-*(96,485 C/1 mole of e^-)=2.750 x 10^1 C Since current (A)=charge (C)/time (s), solve for time: A=C/s C/A=s 2.750 x 10^1 C/0.234 A=time (s) 1.18 x 10^2 s=118s=time <= 3 significant figures

It will fall off and tip over.