the overall reaction in a commerical heat pack can be repesented as shown below?

Question

4Fe(s) + 3O2(g) -> 2Fe2O3(s) ΔH= -1652KJ

a. how much heat is released when 4.36mol iron is reacted with excess O2?

b. how much heat is released when 1.04mol Fe2O3 is produced?

c. how much heat is released when 2.26g iron reacted with excess O2?

d. how much heat is released when 10.6g Fe and 1.43g O2 are reacted?

all in KJ

Answer 1

The ΔH given is -1652KJ/mol of iron.

a. 4.36mol iron * -1652KJ/mol of iron =? (calculator)

b. Stoichiometry rules:
2 moles of Fe2O3 = 4 moles of Fe
1. 04 moles of Fe2O3 = ?
moles Fe = 4 *1.04/2 = calculator
Then using the moles of Fe you can calculate the heat released using the method I showed you in part A.

c. Convert mass of iron into moles of iron.
mass= Molar mass x moles
moles= mass/molarmass = 2.26/(check periodictable)

Then using the moles of Fe you can calculate the heat released using the method I showed you in part A.

d. Convert both iron and oxygen into moles in the method shown in part C. Divide moles of iron by 4 and moles of oxygen by 3 to find the limiting reactant (the smallest resulting number). If iron is your limiting reactant, multiply the moles of iron by ΔH as in part A. If oxygen is your limiting reactant, use stoichiometry to find the moles of iron that actually reacted, as in Part B, then multiply the moles of iron by ΔH as in part A.

Hopefully that helps you solve this question, and future questions (I don't want to just give you answers, how will you learn anything?)

Answer 2

Bookkeeping and proportionality.

a) 4.36/4 x 1652 kJ evolved
b) 1.04/2 x 1652

Get the idea?

c) Convert g Fe to moles
d) Convert Fe and O2 to moles. Work out which is limiting reagent. Then proceed as in (a) or (b) depending