1. University of North Carolina, Wilmington
Thermochemistry
David P. White
University of North Carolina, Wilmington
Chapter 5
Copyright 1999, PRENTICE HALL
Chapter 5 1 1 1 1

2. The Nature of Energy Kinetic and Potential Energy From Physics:
Force is a push or pull on an object.
Work is the product of force applied to an object over a distance:
w = F  d
Energy is the work done to move an object against a force.
Kinetic energy is the energy of motion:
Copyright 1999, PRENTICE HALL
Chapter 5

3. The Nature of Energy Kinetic and Potential Energy
Potential energy is the energy an object possesses by virtue of its position.
Potential energy can be converted into kinetic energy. Example: a ball of clay dropping off a building.
Copyright 1999, PRENTICE HALL
Chapter 5

4. The Nature of Energy Energy Units 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
Copyright 1999, PRENTICE HALL
Chapter 5

5. The Nature of Energy Systems and Surroundings
System: part of the universe we are interested in.
Surroundings: the rest of the universe.
Copyright 1999, PRENTICE HALL
Chapter 5

6. First Law of Thermodynamics
Internal Energy
Internal Energy: total energy of a system.
Cannot measure absolute internal energy.
Change in internal energy, DE = Efinal - Einitial
Copyright 1999, PRENTICE HALL
Chapter 5

7. 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:
DE = q + w
Copyright 1999, PRENTICE HALL
Chapter 5

8. First Law of Thermodynamics
Relating DE to Heat and Work
Copyright 1999, PRENTICE HALL
Chapter 5

9. First Law of Thermodynamics
Relating DE to Heat and Work
Copyright 1999, PRENTICE HALL
Chapter 5

10. First Law of Thermodynamics
Endothermic and Exothermic Processes
Endothermic: absorbs heat from the surroundings.
Exothermic: transfers heat to the surroundings.
An endothermic reaction feels cold.
An exothermic reaction feels hot.
Copyright 1999, PRENTICE HALL
Chapter 5

11. 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.
Copyright 1999, PRENTICE HALL
Chapter 5

12. First Law of Thermodynamics
State Functions
Copyright 1999, PRENTICE HALL
Chapter 5

13. DH = Hfinal - Hinitial = qP
Enthalpy
Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
Can only measure the change in enthalpy:
DH = Hfinal - Hinitial = qP
Copyright 1999, PRENTICE HALL
Chapter 5

14. Enthalpies of Reaction
For a reaction
DHrxn = H(products) - H (reactants)
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
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
Copyright 1999, PRENTICE HALL
Chapter 5

15. q = (specific heat)  (grams of substance)  T.
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.
q = (specific heat)  (grams of substance)  T.
Be careful of the sign of q.
Copyright 1999, PRENTICE HALL
Chapter 5

16. Calorimetry Heat Capacity and Specific Heat
Copyright 1999, PRENTICE HALL
Chapter 5

17. Calorimetry Constant-Pressure Calorimetry
Atmospheric pressure is constant!
DH = qP
qrxn = -qsoln = -(specific heat of solution)  (grams of solution)  DT.
Copyright 1999, PRENTICE HALL
Chapter 5

18. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry)
Reaction carried out under constant volume.
Use a bomb calorimeter.
Usually study combustion.
Copyright 1999, PRENTICE HALL
Chapter 5

19. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry)
qrxn = -CcalorimeterT.
Copyright 1999, PRENTICE HALL
Chapter 5

20. 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
Copyright 1999, PRENTICE HALL
Chapter 5

21. Hess’s Law In the above enthalpy diagram note that H1 = H2 + H3
Copyright 1999, PRENTICE HALL
Chapter 5

22. 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.
If there is more than one state for a substance under standard conditions, the more stable one is used.
Copyright 1999, PRENTICE HALL
Chapter 5

23. Enthalpies of Formation
Standard enthalpy of formation of the most stable form of an element is zero.
Copyright 1999, PRENTICE HALL
Chapter 5

24. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.
Hrxn = H1 + H2 + H3
Copyright 1999, PRENTICE HALL
Chapter 5

25. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
For a reaction:
Copyright 1999, PRENTICE HALL
Chapter 5

26. 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
Fats: contain more energy; are not water soluble, so are good for energy storage.
Copyright 1999, PRENTICE HALL
Chapter 5

27. Foods and Fuels
Foods
Copyright 1999, PRENTICE HALL
Chapter 5

28. Foods and Fuels Fuels U.S.: 1.0 x 106 kJ of fuel per day.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
Copyright 1999, PRENTICE HALL
Chapter 5

29. 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.
Copyright 1999, PRENTICE HALL
Chapter 5

30. Thermochemistry End of Chapter 5 Copyright 1999, PRENTICE HALL54 54 54 54