1. Prentice Hall © 2003Chapter 5 Chapter 5 Thermochemistry CHEMISTRY The Central Science 9th Edition David P. White

2. Prentice Hall © 2003Chapter 5 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. The Nature of Energy

3. Prentice Hall © 2003Chapter 5 Kinetic Energy and Potential Energy Electrostatic potential energy, E el, is the attraction between two oppositely charged particles, Q 1 and Q 2, a distance d apart: The constant  = 8.99  10 9 J-m/C 2. If the two particles are of opposite charge, then E el is the electrostatic attraction between them. The Nature of Energy

4. Prentice Hall © 2003Chapter 5 Units of Energy SI Unit for energy is the joule, J: We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal The Nature of Energy

5. Prentice Hall © 2003Chapter 5 Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe. The Nature of Energy

6. Prentice Hall © 2003Chapter 5 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 (which doesn’t exist-Dr. Hansen) is the transfer of energy between two objects. Energy is the capacity to do work or transfer heat. The Nature of Energy

7. Prentice Hall © 2003Chapter 5 Internal Energy Internal Energy: total energy of a system. Cannot measure absolute internal energy. Change in internal energy, The First Law of Thermodynamics

8. Prentice Hall © 2003Chapter 5 Relating  E 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: The First Law of Thermodynamics

11. Prentice Hall © 2003Chapter 5 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. The First Law of Thermodynamics

12. Prentice Hall © 2003Chapter 5 State Functions State function: depends only on the initial and final states of system, not on how the internal energy is used. The First Law of Thermodynamics

13. State Functions

14. Prentice Hall © 2003Chapter 5 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)  Zn 2+ (aq) + H 2 (g) The work performed by the above reaction is called pressure-volume work. When the pressure is constant, Enthalpy

15. Enthalpy

16. Prentice Hall © 2003Chapter 5 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, Enthalpy

17. Prentice Hall © 2003Chapter 5 Since we know that We can write When  H, is positive, the system gains heat from the surroundings. When  H, is negative, the surroundings gain heat from the system. Enthalpy

18. Prentice Hall © 2003Chapter 5 Enthalpy

19. Prentice Hall © 2003Chapter 5 For a reaction: Enthalpy is an extensive property (magnitude  H is directly proportional to amount): CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ 2CH 4 (g) + 4O 2 (g)  2CO 2 (g) + 4H 2 O(l)  H =  1780 kJ Enthalpies of Reaction

20. Prentice Hall © 2003Chapter 5 When we reverse a reaction, we change the sign of  H: CO 2 (g) + 2H 2 O(l)  CH 4 (g) + 2O 2 (g)  H = +890 kJ Change in enthalpy depends on state: H 2 O(g)  H 2 O(l)  H = -88 kJ Enthalpies of Reaction

21. Prentice Hall © 2003Chapter 5 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. Calorimetry

22. Prentice Hall © 2003Chapter 5 Constant Pressure Calorimetry Atmospheric pressure is constant! Calorimetry

23. Prentice Hall © 2003Chapter 5 Calorimetry Constant Pressure Calorimetry

24. Prentice Hall © 2003Chapter 5 Calorimetry Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion. Bomb Calorimetry (Constant Volume Calorimetry)

25. Prentice Hall © 2003Chapter 5 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: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H = -88 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ Hess’s Law

26. Prentice Hall © 2003Chapter 5 Note that:  H 1 =  H 2 +  H 3 Hess’s Law

27. Prentice Hall © 2003Chapter 5 Example 1. Na(s) + 1/2Cl 2 (g)  Na(g) + Cl(g) ; ΔH = +230kJ Na(g) + Cl(g)  Na + (g) + Cl - (g); ΔH = +147kJ Na(s) + ½ Cl 2 (g)  NaCl(s); ΔH = -411 kJ Calculate ΔH for the reaction Na + (g) + Cl - (g)  NaCl(s)

28. Prentice Hall © 2003Chapter 5 If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation,  H o f. Standard conditions (standard state): 1 atm and 25 o C (298 K). Standard enthalpy,  H o, 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. Enthalpies of Formation

29. Prentice Hall © 2003Chapter 5 Example 2. Fe(s) + Br 2 (l)  FeBr 2 (s); ΔH = -249.8 kJ FeBr 3 (s)  FeBr 2 (s) + ½ Br 2 (l); ΔH = +18.4 kJ Calculate the heat of formation of FeBr 3.

30. Prentice Hall © 2003Chapter 5 Answers -147 -230 -411 kJ = -788 kJ -249.8-18.4 kJ = -268.2 kJ

31. Prentice Hall © 2003Chapter 5 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. Example – carbon’s most stable form is graphite. oxygen’s most stable form is O 2 Enthalpies of Formation

32. Prentice Hall © 2003Chapter 5 Enthalpies of Formation

33. Prentice Hall © 2003Chapter 5 Using Enthalpies of Formation of Calculate Enthalpies of Reaction We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation. Enthalpies of Formation

35. Prentice Hall © 2003Chapter 5 Using Enthalpies of Formation of Calculate Enthalpies of Reaction For a reaction Enthalpies of Formation

36. Prentice Hall © 2003Chapter 5  Hrxn° =  n  Hf° (products) -  n  Hf° (rctnts) n = moles C. Ex: combustion of ethane 2 C 2 H 6 (g) + 7 O 2 (g) --> 4 CO 2 (g) +6 H 2 O(l)  H°rxn =  n  Hf° (prod) -  n  Hf° (rctnts)  H°rxn = 4 mol  Hf°CO 2 + 6 mol  Hf°H 2 O - 2 mol  Hf° C 2 H 6 -7 mol  Hf° O 2 = 4 mol (-393.51 kJ/mol) + 6 mol (-285.83 kJ/mol - 2 mol (-84.68 kJ/mol) - 7 mol (O) = -1574.0 kJ + (-1715.0 kJ) - (-169.4 kJ) - 0 = -3119.6 kJ --> Note: 5 Steps to these problems

37. Prentice Hall © 2003Chapter 5 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: C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O,  H = -2816 kJ Fats break down as follows: 2C 57 H 110 O 6 + 163O 2  114CO 2 + 110H 2 O,  H = -75,520 kJ Foods and Fuels

38. Prentice Hall © 2003Chapter 5 Foods Fats: contain more energy; are not water soluble, so are good for energy storage. Foods and Fuels

39. Prentice Hall © 2003Chapter 5 Fuels In 2002 the United States consumed 1.03  10 17 kJ of fuel. Most from petroleum, natural gas and coal. Remainder from nuclear, wind, and hydroelectric. Fossil fuels are not renewable. Foods and Fuels

40. Prentice Hall © 2003Chapter 5 Foods and Fuels

41. Prentice Hall © 2003Chapter 5 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. Foods and Fuels