Chapter 5 Homework These Equations Are Called Thermochemical Equations

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

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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

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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:

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• 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

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• 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.

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Internal Energy

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• 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

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• 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

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• 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).

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• 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

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• 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

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• 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.

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“Heat Flow vs. Cash Flow: A Banking Analogy” from Further Readings

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“Pictorial Analogies III: Heat Flow, Thermodynamics, and Entropy” from Further Readings

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5.4 Enthalpies of Reaction

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• 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.

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“Enthalpy of Solution” Activity from Instructor’s Resource CD/DVD

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“Formation of Water” Movie from Instructor’s Resource CD/DVD

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Heat Capacity and Specific Heat

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• Heat capacity is the amount of energy required to raise the temperature of an object by 1 °C.

Constant-Pressure Calorimetry

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• 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)

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• Reactions can be carried out under conditions of constant volume instead of constant pressure.

5.6 Hess’s Law

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• 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.

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“Boiling Water in a Paper Cup: Heat Capacity of Water” from Live Demonstrations

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“A Specific Heat Analogy” from Further Readings

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5.7 Enthalpies of Formation

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• 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

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• 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.

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“Formation of Aluminum Bromide” Movie from Instructor’s Resource CD/DVD

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Table 5.3 from Transparency Pack

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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

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• 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

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• 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

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• 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:

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“Calories–Who’s Counting?” from Further Readings

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“Sucrose” 3-D Model from Instructor’s Resource CD/DVD

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“Carbon Dioxide” 3-D Model from Instructor’s Resource CD/DVD

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• 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).

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