Final answer:
To determine the weight of anthraquinone produced, calculate the charge passed, convert it to moles of electrons, adjust for current efficiency, and then multiply by the molar mass of anthraquinone. The answer is approximately 2.50 grams.
Step-by-step explanation:
The student has asked how much anthraquinone can be produced by the passage of a current of 1 A for 40 minutes, given a current efficiency of 96.5%. First, calculate the total charge passed using the formula Q = It, where I is the current in Amperes and t is the time in seconds. To convert minutes to seconds, multiply by 60. So, Q = 1 A * (40 min * 60 sec/min) = 2400 C.
Next, determine the moles of electrons transferred using Faraday's constant (96500 C/mol of e-). As anthracene to anthraquinone is a two-electron process, we divide the total charge by half of Faraday's constant. The moles of electrons are then 2400 C / (2 * 96500 C/mol of e-) = 0.01244 mol of e-.
Considering the current efficiency of 96.5%, the effective moles of electrons are 0.01244 mol * 0.965 = 0.012005 mol. Since the mole ratio of anthracene to anthraquinone is 1:1, the moles of anthraquinone produced will be the same as the effective moles of electrons.
Finally, to find the weight of anthraquinone, multiply the moles by its molar mass (C14H8O2, with a molar mass of approximately 208.21 g/mol). Weight of anthraquinone = 0.012005 mol * 208.21 g/mol = 2.50 g (rounded to 3 significant figures).