A molten oxide of an unknown element W is electrolyzed in an electrolytic cell…
CHEMISTRY · Multi-part · AP · Chemistry · Thermodynamics and Electrochemistry · Thermodynamics and Electrochemistry
Problem
A molten oxide of an unknown element W is electrolyzed in an electrolytic cell equipped with inert electrodes. A constant current of 1.40 A is passed through the melt for exactly 1.50 h. During the electrolysis, oxygen gas evolves at the anode and metallic W deposits at the cathode. When the experiment is complete, the following data are recorded:
• Volume of O₂(g) collected at 1.00 atm and 273 K: 246 mL
• Mass of metallic W obtained: 1.26 g
The atomic molar mass of W is 184 g mol⁻¹.
Assume no side reactions occur and that all charge is used solely for the production of the two observed products.
(a) Calculate the total number of moles of electrons that passed through the cell during the 1.50 h electrolysis.
(b) Determine the number of moles of O₂(g) produced, based on the measured gas volume.
(c) Using your answer from part (b), calculate the number of moles of electrons involved in the oxidation process at the anode that generated the O₂(g). Explicitly relate this electron count to the balanced anode half-reaction.
(d) The electrons released at the anode are consumed at the cathode to produce metallic W. Using the difference between the total electrons from part (a) and the electrons calculated in part (c), determine the number of moles of electrons involved in the reduction of Wⁿ⁺ to W(s).
(e) From the data in part (d) and the measured mass of W, determine the integer value of n (the oxidation state of W in the molten oxide).
(f) Write the empirical formula of the oxide WₘOₚ that was present in the melt before electrolysis, and verify that the charges in your formula balance.