Chemistry Question?

Question

How do you solve this problem?

Calculate the energy change in kJ for the formation of 1.23 mol of LiF(s) given the following information:
Li(g) → Li(s) : −159.4 KJ/mol
F2(g) → 2F(g) : 157.98 KJ/mol
Li(g) → Li+(g) + e− : 520 KJ/mol
F(g) + e− → F−(g) : −328 KJ/mol
Lattice energy for LiF = −1037 KJ/mol

Answer 1

If you know the heat of a reaction, just reverse the sign to get the heat of the reverse reaction. Also, the heat of formation is the heat you would get from reaction of the elements in their standard states. The standard state of Li is Li(s) and of F, F2(g). To get them to react, you have to convert them to gaseous atoms:

Li(s) ===> Li(g) 159.4kJ/mol
F2(g) ===> 2F(g) 157.88kJ/mol
1/2F2(g) ===> F(g) 78.9kJ/mol because you only need one F
Total: 238.3kJ/mol to get Li(g) and F(g)

Next, ionize the atoms:

Li(g) ===> Li+(g) + e- 520kJ/mol
F(g) + e- ===> Fi-(g) -328kJ/mol
Total: 192kJ/mol to get Li+ and F-

Finally, put the ions together:

Li+ + F- ===> LiF -1037kJ/mol

238.3kJ/mol + 192kJ/mol + (-1037kJ/mol) = -607kJ/mol

You wanted the heat to get 1.23 moles LiF, and that would be (1.23)(-607) = 747kJ