Chemistry help?

Question

1. The total volume of hydrogen gas needed to fill the Hindenburg was 2.00 x 10^8L at 1.00 atm and 25.0°C. How much energy was evolved when it burned?

2. What are the signs (positive or negative) of q (heat) and w (work) for the following exothermic process at P=1 atm and T=370K
H2O(g)-->H2O(l)

Then one problem using Hess' Law, I can't get it to work out correctly:

3. Using the following data, calculate the standard heat of formation of the compound ICl (g) in kJ/mol
Cl2(g) --> 2Cl(g) ∆H° = 242.3kJ
I2(g) --> 2I(g) ∆H° = 151.0kJ
ICl(g) --> I(g) + Cl(g) ∆H° = 211.3kJ
I2(s) --> I2(g) ∆H° = 62.8kJ

Answer 1

(1) Use ideal gas law to find the number of moles of H2:
n = PV/RT = 2.00x10^8/(0.08206*298.2) = 8.17x10^6 (mole)
Std enthalpy change of formation of water: ΔfH(gas): −241.83 kJ/mol
Since 1 mole of H2 reacts with O2 to form 1 mole of water, the energy evolved when it burned is:
8.17x10^6mole*(−241.83 kJ/mol) = 1.98x10^9 kJ.

(2) w: negative
q: negative

(3) Cl2(g) --> 2Cl(g) ∆H° = 242.3kJ...(1)
I2(g) --> 2I(g) ∆H° = 151.0kJ.............(2)
ICl(g) --> I(g) + Cl(g) ∆H° = 211.3kJ...(3)
I2(s) --> I2(g) ∆H° = 62.8kJ................(4)
0.5[(1)+(4)+(2)]-(3):
0.5Cl2(g) + 0.5 I2(s) ==> ICl(g), ∆H° = 0.5*(242.3+151.0+62.8)-211.3 = 16.75 kJ/mol.