Unlock document.
This document is partially blurred.
Unlock all pages and 1 million more documents.
Get Access
Chapter 5. Thermochemistry
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 5.5 Changes in Internal Energy 5.2 The First Law of Thermodynamics
Figure 5.10 Internal Energy Is a State Function, but 5.2 The First Law of Thermodynamics
Heat and Work Are Not
Animations: Section:
Work of Gas Expansion 5.3 Enthalpy
Movies: Section:
Thermite 5.2 The First Law of Thermodynamics
Activities: Section:
Sign Conventions for q and w 5.2 The First Law of Thermodynamics
3-D Models: Section:
Carbon Dioxide 5.8 Fuels and Foods
Oxygen 5.8 Fuels and Foods
Chapter 5
60
Other Resources
Further Readings: Section:
Weight-Loss Diets and the Law of Conservation 5.2 The First Law of Thermodynamics
of Energy
Pictorial Analogies III: Heat Flow, 5.3 Enthalpy
Thermodynamics, and Entropy
Analogical Demonstrations 5.3 Enthalpy
Heat Flow vs. Cash Flow: A Banking Analogy 5.3 Enthalpy
Live Demonstrations: Section:
Evaporation as an Endothermic Process 5.2 The First Law of Thermodynamics
Flaming Cotton 5.4 Enthalpies of Reaction
Thermochemistry
61
Chapter 5. Thermochemistry
Common Student Misconceptions
• Students confuse power and energy.
• Students confuse heat with temperature.
• Students fail to note that the first law of thermodynamics is the law of conservation of energy.
• Students have difficulty in determining what constitutes the system and the surroundings.
• Sign conventions in thermodynamics are always problematic.
Teaching Tips
• Remind students that the values of ∆H°f for the same compound but in a different phase are different;
Lecture Outline
5.1 The Nature of Energy
• Thermodynamics is the study of energy and its transformations.
• Thermochemistry is the study of the relationships between chemical reactions and energy changes
involving heat.
• Definitions:
Chapter 5
62
• Potential energy is the energy an object possesses by virtue of its position or composition.
• Electrostatic energy is an example.
Units of Energy
• SI unit is the joule (J).
• The nutritional Calorie, Cal = 1,000 cal = 1 kcal.
System and Surroundings
• A system is the part of the universe we are interested in studying.
• Surroundings are the rest of the universe (i.e., the surroundings are the portions of the universe that
are not involved in the system).
• Example: If we are interested in the interaction between hydrogen and oxygen in a cylinder, then the
H2 and O2 in the cylinder form a system.
Transferring Energy: Work and Heat
• From physics:
• Force is a push or pull exerted on an object.
• Energy is the capacity to do work or to transfer heat.
FORWARD REFERENCES
• The concepts of the system, surroundings, and universe will be used again in Chapter 19.
• Energy in J will be used in Chapter 6 to express energy of a hydrogen atom and the energy
of an emitted photon.
5.2 The First Law of Thermodynamics
1
• The first law of thermodynamics states that energy cannot be created or destroyed.
• The first law of thermodynamics is the law of conservation of energy.
• That is, the energy of system + surroundings is constant.
Thermochemistry
63
Internal Energy
2
• The total energy, E, of a system is called the internal energy.
• It is the sum of all the kinetic and potential energies of all components of the system.
• Absolute internal energy cannot be measured, only changes in internal energy.
Relating E to Heat and Work
3
• From the first law of thermodynamics:
• When a system undergoes a physical or chemical change, the change in internal energy is given
by the heat added to or liberated from the system plus the work done on or by the system:
Endothermic and Exothermic Processes
• An endothermic process is one that absorbs heat from the surroundings.
• An endothermic reaction feels cold.
• An exothermic process is one that transfers heat to the surroundings.
• An exothermic reaction feels hot.
State Functions
4
,
5
,
6
• A state function depends only on the initial and final states of a system.
• Example: The altitude difference between Denver and Chicago does not depend on whether you
is the same in both cases.
FORWARD REFERENCES
• Equilibria of endothermic vs. exothermic reactions will be differently affected by temperature
changes in Chapter 15.
• Water autoionization is mentioned to be an endothermic process (Chapter 16, section 16.3).
Chapter 5
64
• Solid→liquid→gas phase changes are endothermic, while gas→liquid→ solid phase changes
are exothermic.
• Most spontaneous reactions (Chapter 19) are exothermic.
• State functions will be further discussed in Chapter 19 (entropy and Gibbs free energy).
5.3 Enthalpy
7
,
8
,
9
,
10
,
11
• Chemical and physical changes that occur around us occur under essentially constant pressure of
Earth’s atmosphere.
• Changes may be accompanied by work done by or on the system.
• Changes may involve the release or absorption of heat.
• For most reactions P∆V is small thus ∆H = ∆E
• Heat transferred from surroundings to the system has a positive enthalpy (i.e., ∆H > 0 for an
endothermic reaction).
• Heat transferred from the system to the surroundings has a negative enthalpy (i.e., ∆H < 0 for an
exothermic reaction).
• Enthalpy is a state function.
A Closer Look at Energy, Enthalpy, and P-V Work
12
• Many chemical reactions involve work done on or by the system.
• Work is often either electrical or mechanical work.
• Mechanical work done by a system involving expanding gases is called pressure-volume work
or P-V work.
7
“Heat Flow vs. Cash Flow: A Banking Analogy” from Further Readings
8
“Pictorial Analogies III: Heat Flow, Thermodynamics, and Entropy” from Further Readings
9
Thermochemistry
65
5.4 Enthalpies of Reaction
13
,
14
,
15
,
16
,
17
,
18
,
19
,
20
• For a reaction, ∆Hrxn = Hproducts – Hreactants.
• The enthalpy change that accompanies a reaction is called the enthalpy of reaction or heat of
reaction (∆Hrxn).
• Consider the thermochemical equation for the production of water:
2H2(g) + O2(g) → 2H2O(g) ∆H = –483.6 kJ
• The equation tells us that 483.6 kJ of energy are released to the surroundings when water is
formed.
• ∆H noted at the end of the balanced equation depends on the number of moles of reactants and
products associated with the ∆H value.
FORWARD REFERENCES
• Enthalpies of chemical reactions and physical processes will be used throughout the textbook.
• Enthalpies of reactions will be calculated using average bond enthalpies in Chapter 8.
• Enthalpy changes will be used in Chapter 19 to evaluate Gibbs free energy changes.
5.5 Calorimetry
• Calorimetry is a measurement of heat flow.
• A calorimeter is an apparatus that measures heat flow.
13
“Enthalpy of Solution” Activity from Instructor’s Resource CD/DVD
14
“Formation of Water” Movie from Instructor’s Resource CD/DVD
Chapter 5
66
Heat Capacity and Specific Heat
21
• Heat capacity is the amount of energy required to raise the temperature of an object by 1 °C.
Constant-Pressure Calorimetry
22
,
23
• The most common technique is to use atmospheric pressure as the constant pressure.
• Recall H = qp.
• The easiest method is to use a coffee cup calorimeter.
qsoln = (specific heat of solution) (grams of solution) T = –qrxn
• For dilute aqueous solutions, the specific heat of the solution will be close to that of pure water.
Bomb Calorimetry (Constant-Volume Calorimetry)
24
• Reactions can be carried out under conditions of constant volume instead of constant pressure.
5.6 Hess’s Law
25
,
26
• Hess’s Law: If a reaction is carried out in a series of steps, ∆H for the reaction is the sum of ∆H for
each of the steps.
• The total change in enthalpy is independent of the number of steps.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ∆H = –802 kJ
FORWARD REFERENCES
• Similar rules will apply to ∆G and ∆S in Chapter 19.
21
“Boiling Water in a Paper Cup: Heat Capacity of Water” from Live Demonstrations
22
“A Specific Heat Analogy” from Further Readings
Thermochemistry
67
5.7 Enthalpies of Formation
27
,
28
• If a compound is formed from its constituent elements, then the enthalpy change for the reaction is
called the enthalpy of formation, ∆Hf.
• Standard state (standard conditions) refer to the substance at:
C60.
• The standard enthalpy of formation of the most stable form of an element is zero.
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
29
,
30
• Use Hess’s law!
• Example: Calculate ∆H for C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
FORWARD REFERENCES
• In Chapter 19 (section 19.5) Gibbs free energies of formation, ∆G°f, will used to find ∆G°rxn
in an analogical manner to how ∆H°f values are used to find ∆H°rxn.
27
“Formation of Aluminum Bromide” Movie from Instructor’s Resource CD/DVD
28
Table 5.3 from Transparency Pack
Chapter 5
68
5.8 Foods and Fuels
• Fuel value is the energy released when 1 g of substance is burned.
• The fuel value of any food or fuel is a positive value that must be measured by calorimetry.
Foods
31
,
32
,
33
,
34
,
35
,
36
• Fuel value is usually measured in Calories (1 nutritional Calorie, 1 Cal = 1000 cal).
• Most energy in our bodies comes from the oxidation of carbohydrates and fats.
Fuels
37
,
38
,
39
,
40
,
41
,
42
,
43
,
44
• In 2008 the United States consumed about 1.05 1017 kJ/year (9.4 105 kJ of fuel per person per
day).
• Most of this energy comes from petroleum and natural gas.
• The remainder of the energy comes from coal, nuclear and hydroelectric sources.
Other Energy Sources
45
• Nuclear energy: energy released in splitting or fusion of nuclei of atoms.
• It is used to produce about 21% of the electric power in the US.
• Fossil fuels and nuclear energy are nonrenewable sources of energy.
• Renewable energy sources include:
31
“Calories–Who’s Counting?” from Further Readings
32
“Sucrose” 3-D Model from Instructor’s Resource CD/DVD
33
“Carbon Dioxide” 3-D Model from Instructor’s Resource CD/DVD
Thermochemistry
69
• Solar energy
• Wind energy
FORWARD REFERENCES
• Pollution resulting from fossil-fuel combustion will be discussed in Chapter 18 (section 18.2).
• Pollution-free fuel cell-powered vehicles will be mentioned in Chapter 20 (section 20.7).
Chapter 5
70
Further Readings:
1. John J. Fortman, “Pictorial Analogies III: Heat Flow, Thermodynamics, and Entropy,” J. Chem.
Educ., Vol. 70, 1993, 102–103.
6. Brother Thomas McCullogh CSC, “A Specific Heat Analogy,” J. Chem. Educ., Vol. 57, 1980, 896.
7. Doris R. Kimbrough, “Heat Capacity, Body Temperature, and Hypothermia,” J. Chem. Educ., Vol. 75,
1998, 48–49.
8. JCE Editorial Staff, “JCE Classroom Activity #65: Calories–Who’s Counting?,” J. Chem. Educ.,
Vol. 81, 2004, 1440A.
14. Israel Dostrovsky, “Chemical Fuels from the Sun,” Scientific American, Vol. 265, 1991, 102–107.
15. Richard Monastersky, “The Ice that Burns. Can Methane Hydrates Fuel the 21st Century?” Science
News, Vol. 154, 1998, 312–313. This article explores a potential new source of natural gas.
Live Demonstrations:
1. Bassam A. Shakhashiri, “Evaporation as an Endothermic Process,” Chemical Demonstrations: A
Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989).
pp. 249–251.
Thermochemistry
71
2. Lee R. Summerlin, Christie L. Borgford, and Julie B Ealy, “Flaming Cotton,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
p. 104. A cotton ball sprinkled with sodium peroxide bursts into flame upon addition of water.
5. Lee. R. Summerlin,, Christie L. Borgford, and Julie B. Ealy, “A Chemical Hand Warmer,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
pp.101–102. Oxidation of iron in a plastic baggie is used to prepare a hand-warmer.
6. Lee R. Summerlin and James L. Ealy, Jr. , “Endothermic Reaction: Ammonium Nitrate,” Chemical
Demonstrations, A Sourcebook for Teachers, (Washington: American Chemical Society, 1988), p. 65.
Temperature changes that accompany dissolution of ammonium nitrate in water are measured.
