A. In this reaction we see that 1652 kJ of energy is released when 4 moles of Fe is oxidized, as per the balanced equation. So that's the answer to A.
B. In the 1st place, this is worded incorrectly as Fe is not produced but consumed. Assuming this is a mere "oversight", we could say that since 1652 kJ is released when 4.00 mol of Fe is reacted (see A), then 1/4 of that would be released when 1 mole reacts. Namely, 412 kJ.
C. Here we have either a limiting reactant problem to do first before we can answer the question or else there are two separate things being asked. Namely, the amount of heat released when 10.0 g of Fe react AND (separate answer) the amount released when 2.00 g O2 react. {In other words, this is a rater poorly worded question!}
Let's look at the 1st possibility: We will determine the LR.
1) If Fe were the LR, then: (1652 kJ/4 mol Fe)(1 mol Fe/ 55.85 g Fe) 10.0 g Fe = 73.9 kJ released
If O2 were the LR, then: (1652 kJ/3 mol O2)(1 mol O2/32.0 g O2) 2.00 g O2 = 34.4 kJ released.
So - based on the interpretation thatyou have BOTH 10.0 g Fe and 2.00 g O2, only 34.4 kJ of heat would be released as there is an excess of Fe.
2) Now, if the question is interpreted as two separate possibilities, then you have BOTH of the answers shown in the above calculation.
