Chapter 6 Lecture Problems

6.4 Quantifying Heat and Work

1. How much energy is required to heat 250 mL of water from 15 C to 90 C to make a cup of hot chocolate given that the specific heat capacity of water is 4.184 J g-1 K-1

2. Use the same amount of energy that was required to heat the water to heat 250 g of Gold (0.128 J g-1 K-1). What is the final temperature?
3. Heat 50 g piece of copper to 500 C. Then Place in 1 liter of water at 20 C. What is the final Temperature? You can solve this problem algebraically or using successive approximations.
4. Solutions

6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry

1. Burning methane to find heat the capacity of a calorimeter. Burning a 15.00 g sample of methane, CH4, in a bomb calorimeter increases the temperature in the calorimeter by 2.190 oC. The heat of combustion of methane is 890.3 kJ mol-1. (Question from Kask, U.; Rawn, D. General Chemistry (W.C. Brown, Dubuque, IA, 1993)
1. Calculate the heat capacity of the calorimeter in kJ oC-1.
1. 15.00 g methane = 0.937 moles
2. [delta] Hcombustion = 890.3 kJ/mol
3. [delta] H = 834.6 kJ
4. [delta] T = 2.190 C
5. heat capacity = 381.1 kJ/oC
2. A 0.1258 g sample of butane is burned in the same calorimeter. The heat of combustion for butane is 2877 kJ/mol. The initial temperature of the calorimeter is 17.57 oC. Calculate the final temperature for the calorimeter.

6.7 Constant-Pressure Calorimetry: Measuring Hrxn

1. Dissolving salt in water. Dissolving 75.0 g of ammonium nitrate in 250 g of water at 25.72 oC decreases the temperature to 6.89 oC in a calorimeter. Calculate delta H for the reaction. (Question from Olmsted and Williams Chemistry, 3rd ed (Wilen and SOns, New York)
1. Is the process exothermic or endothermic? Endothermic, temperature decreased.
2. Assuming that no heat is gained or lost by the calorimeter walls, calculate the heat of solution of the salt in joules per mole of the salt.
1. [delta] T = -18.83 oC
2. Mass = 325 g
3. Assume the specific heat for the solution is the same as water
4. Erxn = (4.2 J g-1 C-1) * (18.83 oC) * (325 g)
5. E = 25703 J
6. 75 g salt, 80.04 g/mole, 0.937 moles, 27 kJ/mole

6.8 Relationships involving Hrxn

1. Use energy level diagram to calcuate /\H for the following reaction:
C (s) + 2 H2 -> CH4
2. Data:
 Reaction [delta] Hrxn (kJ) C (s) + O2 (g) -> CO2 (g) -393.51 H2 (g) + 1/2O2 (g) -> H2O (l) -285.83 CH4 (g) + 2 O2 (g) -> CO2 (g) + 2 H2O (l) -890.37

Calculations:
 Multiplier Reactants Products /\H (kJ) 1 C + O2 CO2 -393.51 2 2 H2 + O2 2 H2O -571.66 -1 CO2 (g) + 2 H2O (l) CH4 (g) + 2 O2 (g) 890.37 Overall C (s) + O2 (g) CO2 (g) -74.8

6.9 Enthalpies of Reaction from Standard Heats of Formation

• [delta] H for production of 500 g CCl4 based upon [delta] Hf and the reaction:
CH4(g) + 4 Cl2 (g) -> CCl4 (l) + 4 HCl (g)

Data:
 Compound [delta] Hof (kJ) CH4 (g) -74.8 Cl2 (g) 0 CCl4 (l) -135.4 HCl (g) -92.3

Calculations:
 Multiplier Compound [delta] H (kJ) -1 CH4 (g) 74.8 -4 Cl2 (g) 0 1 CCl4 (l) -135.4 4 HCl (g) -369.2 Overall -429.8

• Calculate [delta] Hrxn for the production of hydrazine rocked fuel and the amount of energy required to produce 1.00 kg of hydrazine formed from the reaction: (Ebbing page 252)
N2 (g) + 2H2 (g) -> N2H4 (l)

 Reaction [delta] H (kJ) N2H4 (l) + O2 (g) -> N2 (g) + 2 H2O (l) -622.2 H2 (g) + 1/2O2 -> H2O (l) -285.8

• Calculate [delta] Hrxn for the combustion of gasoline and determine how much energy is released by burning 1 gallon of gasoline? (Density 0.7025 g/cm3)
2 C8H18 + 25 O2 -> 16 CO2 + 18 H2O
• Calculate how much energy is required to produce 150 g of glucose ([delta]Hf = -1274 kJ mole-1, Aitkins) from CO2 and H2O
6 CO2 (g) + 6 H2O (l) -> C6H12O6 (s) + 6 O2 (g)

• The Ostwand process for nitric acid involves the following steps:
4 NH3 (g) + 5 O2 (g) -> 4 NO (g) + 6 H2O (g)
2 NO (g) + O2 (g) -> 2 NO2 (g)
3 NO2 (g) + H2O (l) -> 2 HNO3 (aq) + NO (g)
1. Find [delta] H for each reaction.
2. Balance total reaction and find [delta] H.
3. How much energy is required to produce 1000 kg of nitric acid?
solutions