CH 302 Spring 2005
Chapter 15: Chemical Thermondynamics
1. Write the correct term for each definition given below:
a. The energy necessary to break one mole of bonds in a gaseous substance.
Bond energy
b. The most stable state of a substance under standard pressure and temperature.
Thermochemical standard state
c. The amount of heat required to raise the temperature of an object one degree C.
Heat capacity
d. The study of energy change in chemical systems.
Thermodynamics
e. Parameters that define the current state of a chemical system.
State functions
f. ∆H when reactants in standard states are converted to products in standard states.
Standard enthalpy change
2. For each of the following reactions, (a) does the enthalpy increase or decrease; (b) is
Hreactants > Hproducts or is Hreactants < Hproducts; (c) is ∆H positive or negative?
Al2O3 (s) → 2Al(s) + 3/2 O2(g) (endothermic)
(a) Decrease
(b) ∆Hreactants < ∆Hproducts
(c) Positive
Sn(s) + Cl2(g) → SnCl2(s)
(exothermic)
(a) Increase
(b) ∆Hreactants > ∆Hproducts
(c) Negative
3. Calculate the enthalpy change when 35.0 g of iron reacts with oxygen at 1 atm and
25°C.
∆H = -824 kJ/mol
4Fe(s) + 3O2(g) → 2Fe2O3(s)
∆H= (35.0g)(1mol Fe/55.8g)(1 mol rxn/4 mol Fe)(-824 kJ/mol rxn)
= -129 kJ/mol released to surroundings
4. 2.9 g of fuel is burned in a calorimeter that contains 200 g of water at 20.00 °C. After
combustion, the temperature of the water raises to 23.12 °C. The heat capacity of the
calorimeter is 85.5 J/°C. How much heat is evolved per gram of fuel burned?
∆Hwater= mC∆T = (200 g)(4.184 J/g°C)(3.12°C) = 2611 J
∆Hcalorimeter = c∆T= (85.5 J/°C)(3.12°C) = 267 J
∆Hrxn = ∆Hwater + ∆Hcalorimeter = 2611 J + 267 J = 2878 J
Heat/gram = 2878 J/2.9 g = 992 J/g
5. If 50,000 J is added to 500 g of benzene initially at 25°C, what will the final
temperature be? The specific heat of benzene is 1.74 J/g°C.
∆T = ∆H/mC = (50,000 J)/(500 g )(1.74 J/g°C) = 57.5°C
Tfinal = ∆T + Tinitial = 57.5°C + 25°C =82.5°C
6. Given that
S(s) + O2(g) → SO2(g)
∆H = -296.8 kJ/mol
S(s) + 3/2 O2(g) → SO3(g)
∆H = -395.6 kJ/mol
determine the enthalpy change for the decomposition reaction
2SO3(g) → 2SO2(g) + O2(g)
2SO3 → 2S + 3O2
+ 2S + 2 O2 → 2SO2
2SO3 → 2SO2 + O2
∆H = 2(-296.8 kJ/mol) – 2(-395.6 kJ/mol) = 197.8 kJ/mol
7. Given that
2H2 + O2 → 2H2O
C3H4 + 4O2 → 3CO2 + 2H2O
C3H8 + 5O2 → 3CO2 + 4H2O
determine the heat of the reaction
C3H4 + 2H2 → C3H8.
∆H = -571.6 kJ/mol
∆H = -1937 kJ/mol
∆H = -2220 kJ/mol,
C3H4 + 4O2 → 3CO2 + 2H2O
+ 3CO2 + 4H2O → C3H8 + 5O2
+ 2H2 + O2 → 2H2O
C3H4 + 2H2 → C3H8
∆H = -571.6 kJ/mol – 1937 kJ/mol – (- 2220 kJ/mol) = -289 kJ/mol
8. Use the given bond energies to estimate the enthalpy of reaction for the reaction
C3H8 + Cl2 → C3H7Cl + HCl.
C-H 413 kJ/mol
Cl-Cl 242 kJ/mol
H-Cl
C-Cl
432 kJ/mol
339 kJ/mol
∆H = [∆HC-H + ∆HCl-Cl] – [∆HC-Cl + ∆HH-Cl] = [413 + 242] – [339+432] =
-116 kJ/mol
9. 10.0 g of a solid are combusted at 1100 K in a bomb calorimeter. A container that
surrounds the calorimeter holds 1.5 L of water that increases in temperature by 20°C.
Assume water has a density of 1 g/mL and a specific heat of 1 cal/g°C. Determine ∆H in
kcal for the 10.0g sample.
(1 cal/g°C)(20°C)(1g/mL)(1500 mL) = 30,000 cal = 30 kcal