1. Chapter 5 Thermochemistry
CHEMISTRY The Central Science 9th Edition
Chapter 5 Thermochemistry
David P. White
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Chapter 5

2. The Nature of Energy Kinetic Energy and Potential Energy
Kinetic energy is the energy of motion:
Potential energy is the energy an object possesses by virtue of its position.
Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.
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3. The Nature of Energy Kinetic Energy and Potential Energy
Electrostatic potential energy, Ed, is the attraction between two oppositely charged particles, Q1 and Q2, a distance d apart:
The constant  = 8.99  109 J-m/C2.
If the two particles are of opposite charge, then Ed is the electrostatic repulsion between them.
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4. The Nature of Energy Units of Energy
SI Unit for energy is the joule, J:
We sometimes use the calorie instead of the joule:
1 cal = J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
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5. The Nature of Energy Systems and Surroundings
System: part of the universe we are interested in.
Surroundings: the rest of the universe.
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6. The Nature of Energy Transferring Energy: Work and Heat
Force is a push or pull on an object.
Work is the product of force applied to an object over a distance:
Energy is the work done to move an object against a force.
Heat is the transfer of energy between two objects.
Energy is the capacity to do work or transfer heat.
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7. The First Law of Thermodynamics
Internal Energy
Internal Energy: total energy of a system.
Cannot measure absolute internal energy.
Change in internal energy,
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8. The First Law of Thermodynamics
Relating DE to Heat and Work
Energy cannot be created or destroyed.
Energy of (system + surroundings) is constant.
Any energy transferred from a system must be transferred to the surroundings (and vice versa).
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 absorbed by the system plus the work done on or by the system:
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9. The First Law of Thermodynamics

11. The First Law of Thermodynamics
Exothermic and Endothermic Processes
Endothermic: absorbs heat from the surroundings.
Exothermic: transfers heat to the surroundings.
An endothermic reaction feels cold.
An exothermic reaction feels hot.
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12. The First Law of Thermodynamics
State Functions
State function: depends only on the initial and final states of system, not on how the internal energy is used.
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13. State Functions

14. Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
Enthalpy
Chemical reactions can absorb or release heat.
However, they also have the ability to do work.
For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work.
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
The work performed by the above reaction is called pressure-volume work.
When the pressure is constant,
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15. Enthalpy

16. Enthalpy
Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
Enthalpy is a state function.
If the process occurs at constant pressure,
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17. Enthalpy Since we know that We can write
When DH, is positive, the system gains heat from the surroundings.
When DH, is negative, the surroundings gain heat from the system.
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18. Enthalpy
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19. Enthalpies of Reaction
For a reaction:
Enthalpy is an extensive property (magnitude DH is directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = kJ
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20. Enthalpies of Reaction
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ
Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -88 kJ
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21. Calorimetry Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise the temperature of an object (by one degree).
Molar heat capacity = heat capacity of 1 mol of a substance.
Specific heat = specific heat capacity = heat capacity of 1 g of a substance.
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22. Calorimetry Constant Pressure Calorimetry
Atmospheric pressure is constant!
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23. Calorimetry Constant Pressure Calorimetry Prentice Hall © 2003
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24. Calorimetry Bomb Calorimetry (Constant Volume Calorimetry)
Reaction carried out under constant volume.
Use a bomb calorimeter.
Usually study combustion.
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25. Hess’s Law
Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.
For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ
2H2O(g)  2H2O(l) H = -88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ
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26. Hess’s Law
Note that:
H1 = H2 + H3
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27. Enthalpies of Formation
If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof .
Standard conditions (standard state): 1 atm and 25 oC (298 K).
Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state.
Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.
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28. Enthalpies of Formation
If there is more than one state for a substance under standard conditions, the more stable one is used.
Standard enthalpy of formation of the most stable form of an element is zero.
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29. Enthalpies of Formation
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30. Enthalpies of Formation
Using Enthalpies of Formation of Calculate Enthalpies of Reaction
We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.
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32. Enthalpies of Formation
Using Enthalpies of Formation of Calculate Enthalpies of Reaction
For a reaction
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33. Foods and Fuels
Foods
Fuel value = energy released when 1 g of substance is burned.
1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.
Energy in our bodies comes from carbohydrates and fats (mostly).
Intestines: carbohydrates converted into glucose:
C6H12O6 + 6O2  6CO2 + 6H2O, DH = kJ
Fats break down as follows:
2C57H110O O2  114CO H2O, DH = -75,520 kJ
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34. Foods and Fuels
Foods
Fats: contain more energy; are not water soluble, so are good for energy storage.
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35. Foods and Fuels
Fuels
In 2000 the United States consumed 1.03  1017 kJ of fuel.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
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36. Foods and Fuels
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37. Foods and Fuels
Fuels
Fuel value = energy released when 1 g of substance is burned.
Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g.
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38. End of Chapter 5: Thermochemistry
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