Chapter 6 Lecture Problems
6.1 Chemical Hand Warmers
6.2 The Nature of Energy: Key Definitions
6.3 The First Law of Thermodynamics: There is No Free Lunch
6.4 Quantifying Heat and Work
- 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
- 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?
- 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.
6.5 Measuring E for Chemical Reactions: Constant-Volume Calorimetry
- 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)
- Calculate the heat capacity of the calorimeter in kJ oC-1.
- 15.00 g methane = 0.937 moles
- [delta] Hcombustion = 890.3 kJ/mol
- [delta] H = 834.6 kJ
- [delta] T = 2.190 C
- heat capacity = 381.1 kJ/oC
- 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.
6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure
6.7 Constant-Pressure Calorimetry: Measuring Hrxn
- 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)
- Is the process exothermic or endothermic? Endothermic, temperature decreased.
- 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.
- [delta] T = -18.83 oC
- Mass = 325 g
- Assume the specific heat for the solution is the same as water
- Erxn = (4.2 J g-1 C-1) * (18.83 oC) * (325 g)
- E = 25703 J
- 75 g salt, 80.04 g/mole, 0.937 moles, 27 kJ/mole
6.8 Relationships involving Hrxn
- Use energy level diagram to calcuate /\H for the following reaction:
C (s) + 2 H2 -> CH4
|Reaction||[delta] Hrxn (kJ)|
|C (s) + O2 (g) -> CO2 (g)
|H2 (g) + 1/2O2 (g) -> H2O (l)
|CH4 (g) + 2 O2 (g) -> CO2 (g) + 2 H2O (l)
|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)
||[delta] Hof (kJ)
|Multiplier||Compound||[delta] H (kJ)|
|-4||Cl2 (g) ||0|
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)
||[delta] H (kJ)
|N2H4 (l) + O2 (g) -> N2 (g) + 2 H2O (l)
|H2 (g) + 1/2O2 -> H2O (l)
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)
- Find [delta] H for each reaction.
- Balance total reaction and find [delta] H.
- How much energy is required to produce 1000 kg of nitric acid?
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