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 4Fe + 3O2 -> 2Fe2O3 + 1.65E3 kJ 

C) what is the enthalpy change for the formation of 23.6 g of iron(iii) oxide? 

I'm trying to understand thermochemical reactions and stoichiometry, so could you please tell me if this is correct? And please explain the parts that I clearly don't understand: 

The question, is it the same thing as asking, "how much energy was released if iron(iii) oxide had a mass of 23.6 g?" 
According to the equation: 
4 moles of Fe and 3 moles of O2 react to release -1.65E3 kJ. ....then 2 moles of Fe "contains" -1.65E3 kJ?? 

To calculate how much kJ of energy was released when the mass of Fe2O3 was 23.6g, we need the amount of moles: 
23.6 g are in 1 mole of Fe2O3, and every mole of Fe2O3 is 160g, the molar mass. (??) which, I think??, explains why this works : 
23.6g Fe2O3 x 1 mol Fe2O3/160g 
And Since there's -1.65E3 kJ per 2 moles of Fe2O3: -1.65E3/2 mol Fe2O3 

Putting it altogether: 
Enthalpy = 23.6 g Fe2O3 x 1 mol Fe2O3/160g x -1.65E3kJ/2 mol 2Fe2O3 
= 1.22E2 kJ 

I'm hoping I've got the units all right too... 

Also, what does it mean if "there's -1.65E3 kJ per 2 moles of Fe2O3"? Then there's -1.65E3 kJ 'contained' in the Fe2O3??? Because we wouldn't say -1.65E3 kJ is RELEASED when Fe2O3 reacts, because Fe2O3 is the product... So what exactly is going on (between Fe2O3 and the energy or enthalpy change or whatever.......)??

  • Chemistry -

    I saw your post and the same kind of questions with Bob Pursley. I suppose you're still struggling. I think you are making this a lot harder than it is. Your 122E3 kJ is correct and the equation you used looks ok to me. Here is how I do them
    1.65E3 kJ energy is released when 2*160g (that's 320 g) Fe2O3 is formed. How much is released to form 23.6g Fe2O3?. That's
    1.65E3 kJ x (23.6g/320g) = 122 kJ.

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