Need help with this chemistry problem please?

Question

What heat is given off when 2.5g of iron rusts according to the following reaction? The reaction has a delta H of -1648 kJ/mole.

4Fe(s) + 6O2(g) ----> 2Fe2O3(s)

Answer 1

The equation isn't correctly balanced though it doesn't really matter since we don't need the equation.

4Fe(s) + 3O2(g)→2Fe2O3(s)

2.5 g x -1648 KJ/mol x 1 mol/55.8 g=-73.8 KJ

=-74 KJ(2 sig figs)

Answer 2

The heat of a reaction is calculated from the difference in formation heats of the products and reactants. Fortunately for you, pure elements do not have heats of formation, since they do not form; they just exist. As such, the delta H you have IS the heat released by the reaction.

You are missing an important piece of information though that would be useful in solving the problem; we need to know what the delta H is in terms of; moles of iron or moles of iron oxide.

Assuming it is in terms of iron, all you need do now is find out how many moles of iron there are in 2.5 grams, then multiply that molar amount by the delta H you were given to get the heat (in kJ) released by the reaction.

If it is in terms of iron oxide the problem is a tiny bit more complex; you still need to find the moles of iron. However, you now need to convert that to moles of iron oxide (half as much iron oxide in terms of moles) before you multiply it by the delta H.

Hope that helps!