A student carries out the following reaction: 2 Fe (s) + 3 S (g) → Fe2S3 (s). The student determines that 2.67 moles of iron (III) sulfide were produced, resulting in a percent yield of 63%. Prior to conducting the reaction of iron with sulfur, the student produced the gaseous sulfur that was used via heating solid sulfur in a syringe.

S8 (s) + heat → 8 S (g)

Assume all of the sulfur was converted to the gaseous form, and the entirety of the gaseous sulfur was used in the reaction. What mass of solid sulfur was used in producing the gaseous sulfur?

A.
408 g

B.
913 g

C.
484 g

D.
8547 g

Answerthe half-reaction having the greatest tendency to occur at the anode:     i₂(aq) + 2e⁻  ⇄ 2i⁻  ;   e = +0.53 v======oxidation occurs at the anode. we must therefore find the half-reaction with the greatest tendency to oxidize.the given values of e are reduction potentials (the form of the half-reactions are of reduction).the lower the reduction potential, the higher the potential to oxidize. the half-cell with the lowest reduction potential will have the greatest tendency to oxidize at the anode.the half-reaction having the greatest tendency must be    i₂(aq) + 2e⁻  ⇄ 2i⁻    e = +0.53 vbecause its reduction potential is the lowest out of the three.