1. Chapter 8 Thermochemistry: Chemical Energy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Chapter 8 Thermochemistry: Chemical Energy
Copyright © 2011 Pearson Prentice Hall, Inc.

2. Energy and Its Conservation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Energy and Its Conservation
4/19/2017 7:13:46 PM
Conservation of Energy Law: Energy cannot be created or destroyed; it can only be converted from one form to another.
Energy: The capacity to supply heat or do work.
Kinetic Energy (EK): The energy of motion.
Potential Energy (EP): Stored energy.
This is one way to state the first law of thermodynamics.
Units:
1 cal = J (exactly)
1 Cal = 1000 cal = 1 kcal
Copyright © 2011 Pearson Prentice Hall, Inc.

3. Energy and Its Conservation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Energy and Its Conservation
Copyright © 2011 Pearson Prentice Hall, Inc.

4. Energy and Its Conservation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Energy and Its Conservation
4/19/2017 7:13:46 PM
Thermal Energy: The kinetic energy of molecular motion and is measured by finding the temperature of an object
Heat: The amount of thermal energy transferred from one object to another as the result of a temperature difference between the two
Heat transfers from hot to cold. Hot and cold are relative terms.
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5. Internal Energy and State Functions
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Internal Energy and State Functions
4/19/2017 7:13:46 PM
First Law of Thermodynamics: The total internal energy E of an isolated system is constant
DE = Efinal - Einitial
The first law is essentially the law of conservation of energy.
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6. Internal Energy and State Functions
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Internal Energy and State Functions
CO2(g) + 2H2O(g) kJ energy
CH4(g) + 2O2(g)
DE = Efinal - Einitial = -802 kJ
802 kJ is released when 1 mole of methane, CH4, reacts with 2 moles of oxygen to produce 1 mole of carbon dioxide and two moles of water.
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7. Internal Energy and State Functions
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Internal Energy and State Functions
State Function: A function or property whose value depends only on the present state, or condition, of the system, not on the path used to arrive at that state
Internal energy, E, is a state function. You can also think of climbing a hill. Regardless of the path taken to climb the hill, the height difference is simply the height of the hill.
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8. Expansion Work w = F x d 3CO2(g) + 4H2O(g) C3H8(g) + 5O2(g)
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Expansion Work
w = F x d
3CO2(g) + 4H2O(g)
C3H8(g) + 5O2(g)
6 mol of gas
7 mol of gas
The production of gas (net change is 1 mole of gas from reactants to products) causes the piston to move up.
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9. Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Expansion Work
Expansion Work: Work done as the result of a volume change in the system
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10. Chapter 8: Thermochemistry: Chemical Energy
4/19/2017
8.5 Energy and Enthalpy
DE = q + w
q = heat transferred
w = work = -PDV
q = DE + PDV
Constant Volume (DV = 0):
qV = DE
Constant Pressure:
qP = DE + PDV
This book uses the definition: DE = Efinal - Einitial
“U” is also used for internal energy.
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11. Heat of reaction (at constant pressure)
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Energy and Enthalpy
4/19/2017 7:13:46 PM
qP = DE + PDV = H
Enthalpy change or
Heat of reaction (at constant pressure)
DH = Hfinal - Hinitial
Enthalpy is a state function whose value depends only on the current state of the system, not on the path taken to arrive at that state.
Enthalpy is path-independent.
= Hproducts - Hreactants
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12. 8.6The Thermodynamic Standard State
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017
8.6The Thermodynamic Standard State
3CO2(g) + 4H2O(g)
C3H8(g) + 5O2(g)
DH= kJ
3CO2(g) + 4H2O(l)
C3H8(g) + 5O2(g)
DH = kJ
Thermodynamic Standard State: Most stable form of a substance at 1 atm pressure and at a specified temperature, usually 25 °C; 1 M concentration for all substances in solution.
Standard enthalpy of reaction is indicated by the symbol ΔHo
The conversion from liquid to gas phases requires energy. Therefore, it’s important that the states are specified.
Putting the “°” superscripted next to the DH means standard state.
The standard pressure is actually 1 bar. The enthalpy difference is usually small, however.
3CO2(g) + 4H2O(g)
C3H8(g) + 5O2(g)
DH° = kJ
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

13. Enthalpies of Physical and Chemical Change
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Enthalpies of Physical and Chemical Change
Enthalpy of Fusion (DHfusion): The amount of heat necessary to melt a substance without changing its temperature
Enthalpy of Vaporization (DHvap): The amount of heat required to vaporize a substance without changing its temperature
Enthalpy of Sublimation (DHsubl): The amount of heat required to convert a substance from a solid to a gas without going through a liquid phase
Physical changes.
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14. Enthalpies of Physical and Chemical Change
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Enthalpies of Physical and Chemical Change
Copyright © 2011 Pearson Prentice Hall, Inc.

15. Enthalpies of Physical and Chemical Change
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Enthalpies of Physical and Chemical Change
2Fe(s) + Al2O3(s)
2Al(s) + Fe2O3(s)
DHo = -852 kJ
exothermic
2Al(s) + Fe2O3(s)
2Fe(s) + Al2O3(s)
The enthalpy of reaction is for the reaction in the direction written.
DHo = +852 kJ
endothermic
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16. Examples
Identify each of the following processes as endothermic or exothermic and indicate the sign of ΔHo
Sweat evaporating from your skin
Water freezing in a freezer
Wood burning in fire

17. Example
An LP gas tank in a home barbeque contains 13.2 kg of propane, C3H8. Calculate the heat (in kJ) associated with the complete combustion of all of the propane in the tank
C3H8(g) + CO2(g)  3 CO2(g) + 4 H2O(g)
ΔHo = -2044kJ

18. Example
How much heat (in kJ) is evolved when 5.00g of aluminum reacts with a stoichiometric amount of Fe2O3?
2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s) ΔHo = -852 kJ

19. Calorimetry and Heat Capacity
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Calorimetry and Heat Capacity
Measure the heat flow at constant pressure (DH).
Calorimetry = study of heat flow.
This is sometimes done in undergraduate labs in styrofoam cups.
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20. Calorimetry and Heat Capacity
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Calorimetry and Heat Capacity
4/19/2017 7:13:46 PM
Measure the heat flow at constant volume (DE).
Bomb calorimeter.
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21. Calorimetry and Heat Capacity
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017
Calorimetry and Heat Capacity
Heat Capacity (C): The amount of heat required to raise the temperature of an object or substance a given amount.
q = C x DT
Molar Heat Capacity (Cm): The amount of heat required to raise the temperature of 1 mol of a substance by 1 °C.
q = (Cm) x (moles of substance) x DT
Specific Heat: The amount of heat required to raise the temperature of 1 g of a substance by 1 °C.
DT is the temperature change.
Heat capacity is an extensive property.
Specific heat is an intensive property.
q = (specific heat) x (mass of substance) x DT
Copyright © 2008 Pearson Prentice Hall, Inc.

22. Calorimetry and Heat Capacity
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Calorimetry and Heat Capacity
Molar Heat Capacity (Cm): The amount of heat necessary to raise the temperature of 1 mol of a substance by 1 oC
q = Cm x Moles of substance x DT
Copyright © 2011 Pearson Prentice Hall, Inc.

23. Calorimetry and Heat Capacity
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Calorimetry and Heat Capacity
Note that the values for water are considerably higher than most other substances. Interesting discussion points could be how large bodies of water affect weather, the human body is approximately 60% water and the role it plays in homeostasis, etc.
© 2012 Pearson Education, Inc.
Copyright © 2011 Pearson Prentice Hall, Inc.

24. Examples
Assuming that Coca Cola has the same specific heat as water [4.183 J/goC], calculate the amount of heat (in kJ) transferred when one can (about g) is cooled from 25.0oC – 3.0oC

25. Chapter 8: Thermochemistry: Chemical Energy
8.9 Hess’s Law
4/19/2017
Hess’s Law: The overall enthalpy change for a reaction is equal to the sum of the enthalpy changes for the individual steps in the reaction.
Haber Process:
2NH3(g)
3H2(g) + N2(g)
DH°total = ???
Multiple-Step Process - Given
N2H4(g)
2H2(g) + N2(g)
DH°1 = 95.4 Kj
2NH3(g)
N2H4(g) + H2(g)
DH°2 = kJ
Hydrazine, N2H4, is a reactive intermediate (more on that in Ch 12: Chemical Kinetics). Chemical reactions are far more complicated than we discuss in Ch 3 (stoichiometry).
The overall reaction and step 2 can be measured, but step 1 cannot be measured easily.
2NH3(g)
3H2(g) + N2(g)
DH°total = DH°1 + DH°2
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26. Chapter 8: Thermochemistry: Chemical Energy
4/19/2017
Hess’s Law
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27. Example Find ΔHorxn for the following reaction:
3H2(g) + O3(g)  3 H2O(g) ΔHorxn = ??
Use the following reactions with known ΔH’s
2H2 (g) + O2(g)  2 H2O(g) Δ Ho = kJ
3O2(g)  2 O3 (g) Δ Ho = kJ

28. Example Fin ΔHorxn for the following reaction
C(s) + H2O(g)  CO(g) + H2(g) Horxn = ?
Use the following reactions with known H’s
C(s) + O2(g)  CO2(g) ΔHo = kJ
2CO(g) O2(g)  2CO2(g) Δ Ho = kJ
2H2 (g) + O2(g)  2H2O (g) Δ Ho = kJ

29. Standard Heats of Formation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Standard Heats of Formation
4/19/2017 7:13:46 PM
Standard Heat of Formation (DHof ): The enthalpy change for the formation of 1 mol of a substance in its standard state from its constituent elements in their standard states
Standard states
CH4(g)
C(s) + 2H2(g)
DHof = kJ
1 mol of 1 substance
Students need to remember solid, liquid, gases, and diatomics learned earlier.
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30. Standard Heats of Formation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Standard Heats of Formation
Note the lack of anything that is 0.
Elements (standard state) are defined to have a standard enthalpy of formation of 0.
Copyright © 2011 Pearson Prentice Hall, Inc.

31. Standard Heats of Formation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
Standard Heats of Formation
4/19/2017 7:13:46 PM
Ho = Hof (Products) - Hof (Reactants)
cC + dD
aA + bB
Ho = [c Hof (C) + d Hof (D)] - [a Hof (A) + b Hof (B)]
Using table values.
Products
Reactants
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32. Standard Heats of Formation
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
Standard Heats of Formation
Using standard heats of formation, calculate the standard enthalpy of reaction for the photosynthesis of glucose (C6H12O6) and O2 from CO2 and liquid H2O.
C6H12O6(s) + 6O2(g)
6CO2(g) + 6H2O(l)
DHo = ?
This is Problem 8.16 in the text.
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33. Example
Use the information in Table 8.2 to calculate ΔHo (in kJ) for the reaction of ammonia with oxygen gas to yield nitric oxide (NO) and water vapor, a step in the Ostwald process for the commercial production of nitric acid

34. An Introduction to Entropy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
An Introduction to Entropy
4/19/2017 7:13:46 PM
Spontaneous Process: A process that, once started, proceeds on its own without a continuous external influence
Copyright © 2011 Pearson Prentice Hall, Inc.

35. An Introduction to Entropy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
An Introduction to Entropy
4/19/2017 7:13:46 PM
Entropy (S): The amount of molecular randomness in a system
Copyright © 2011 Pearson Prentice Hall, Inc.

36. An Introduction to Entropy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
An Introduction to Entropy
4/19/2017 7:13:46 PM
Spontaneous processes are
favored by a decrease in H (negative DH).
favored by an increase in S (positive DS).
Nonspontaneous processes are
favored by an increase in H (positive DH).
favored by a decrease in S (negative DS).
Copyright © 2011 Pearson Prentice Hall, Inc.

37. Example
Predict whether ΔSo is likely to be positive or negative for each of the following reactions
H2C=CH2(g) + Br2(g)  BrCH2CH2Br(l)

38. 8.14 An Introduction to Free Energy
Chapter 8: Thermochemistry: Chemical Energy
8.14 An Introduction to Free Energy
4/19/2017
Gibbs Free Energy Change (DG)
DG = DH - T DS
Enthalpy of
reaction
Temperature
(Kelvin)
Entropy
change
Copyright © 2008 Pearson Prentice Hall, Inc.
Copyright © 2008 Pearson Prentice Hall, Inc.

39. An Introduction to Free Energy
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017
An Introduction to Free Energy
Gibbs Free Energy Change (DG)
DG = DH DS - T
DG < 0 Process is spontaneous
DG = 0 Process is at equilibrium
(neither spontaneous nor nonspontaneous)
DG > 0 Process is nonspontaneous
Figure 8.13 nicely demonstrates how enthalpy, entropy, and free energy relate to each other for solid water (ice) and liquid water.
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40. Example
Is the Haber process for the industrial synthesis of ammonia spontaneous or nonspontaneous under standard conditions at 25.0oC. At what temperature (oC) does the changeover occur?

41. An Introduction to Free Energy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
An Introduction to Free Energy
4/19/2017 7:13:46 PM
Gibbs Free Energy Change (DG)
DG = H - T S
Temperature
(Kelvin)
Enthalpy of
reaction
Entropy
change
Copyright © 2011 Pearson Prentice Hall, Inc.

42. An Introduction to Free Energy
Chemistry: McMurry and Fay, 6th Edition
Chapter 8: Thermochemistry: Chemical Energy
4/19/2017 7:13:46 PM
An Introduction to Free Energy
Gibbs Free Energy Change (DG)
DG = H - T S
DG < 0 Process is spontaneous
DG = 0 Process is at equilibrium
(neither spontaneous nor nonspontaneous)
DG > 0 Process is nonspontaneous
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Copyright © 2011 Pearson Prentice Hall, Inc.