1. Thermochemistry Chapter 5 Thermochemistry John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Hmwk pg. 188-196, # 7, 12, 15, 19, 21, 23, 29, 33, 41, 45, 51, 53, 59, 61, 67, 71, 77, 83, 97, 103

2. Thermochemistry Energy _______________________________.  _____: Energy used to cause an object that has mass to move.  _____: Energy used to cause the temperature of an object to rise.

3. Thermochemistry Potential Energy Energy an object possesses by virtue of its position or chemical composition.

4. Thermochemistry Kinetic Energy Energy an object possesses by virtue of its motion.

5. Thermochemistry Units of Energy The SI unit of energy is the joule (J). An older, non-SI unit is still in widespread use: The calorie (cal). 1 cal = 4.184 J Calories that we eat are actually __________ are abbreviated as _____ 1 J = 1  kg m 2 s2s2

6. Thermochemistry System and Surroundings The _____ includes the molecules we want to study (here, the hydrogen and oxygen molecules). The _____ are everything else (here, the cylinder and piston).

7. Thermochemistry Work Energy used to move an object over some distance. w = F  d,

8. Thermochemistry Heat Energy can also be ________________. Heat flows from _____ objects to _____ objects. Recall, cold is simply a ________________.

9. Thermochemistry Transferal of Energy a)The potential energy of this ball of clay is ________ when it is moved from the ground to the top of the wall. b)As the ball falls, its potential energy is converted to _____ energy. c)When it hits the ground, its kinetic energy falls to _____ (since it is no longer moving); some of the energy does work on the ball, the rest is ________________.

10. Thermochemistry First Law of Thermodynamics Energy ________________________________. In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa. Use Fig. 5.5

11. Thermochemistry Internal Energy The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E. Use Fig. 5.5

12. Thermochemistry Internal Energy By definition, the change in internal energy,  E, is the final energy of the system minus the initial energy of the system: Use Fig. 5.5

13. Thermochemistry Changes in Internal Energy If  E > 0, E final > E initial  Therefore, the system absorbed energy from the surroundings.  This energy change is called ________.

14. Thermochemistry Changes in Internal Energy If  E < 0, E final < E initial  Therefore, the system released energy to the surroundings.  This energy change is called ________.

15. Thermochemistry Changes in Internal Energy When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w). That is

16. Thermochemistry  E, q, w, and Their Signs

17. Thermochemistry Exchange of Heat between System and Surroundings When heat is absorbed by the system from the surroundings, the process is ________.

18. Thermochemistry Exchange of Heat between System and Surroundings When heat is released by the system to the surroundings, the process is ________.

19. Thermochemistry State Functions Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.

20. Thermochemistry State Functions However, we do know that the internal energy of a system is independent of the path by which the system achieved that state.  In the system below, the water could have reached room temperature from either direction.

21. Thermochemistry State Functions Therefore, internal energy is a ____________. It depends only on ______________________, not on the path by which the system arrived at that state. And so,  E depends only on E initial and E final.

22. Thermochemistry State Functions However, q and w are _____ state functions. Whether the battery is shorted out or is discharged by running the fan, its  E is the same.  But q and w are different in the two cases.

23. Thermochemistry State Functions Yes! Your bank account balance No! The distance traveled from Penncrest to K o P mall.

24. Thermochemistry Work When a process occurs in an __________, commonly the only work done is a change in volume of a gas pushing on the surroundings (or being pushed on by the surroundings).

25. Thermochemistry Work We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston.

26. Thermochemistry Enthalpy If a process takes place at ____________ (as the majority of processes we study do) and the only work done is this ______________, we can account for heat flow during the process by measuring the ________ of the system. Enthalpy is the internal energy plus the product of pressure and volume:

27. Thermochemistry Enthalpy When the system changes at constant pressure, the change in enthalpy,  H, is  H =  (E + PV) This can be written  H =  E + P  V

28. Thermochemistry Enthalpy Since  E = q + w and w = −P  V, we can substitute these into the enthalpy expression:  H =  E + P  V So, at constant pressure the change in enthalpy is the heat gained or lost.

29. Thermochemistry Endothermicity and Exothermicity A process is endothermic when  H is _____. A process is exothermic when  H is _____.

30. Thermochemistry Endo or Exo? State whether ΔH is positive or negative The combustion of fuel in a car engine Salting winter roads Making snow on the ski trails Skiing down the trails Snow-covered trails melt in the spring time. Placing an ice pack on an injury

31. Thermochemistry Enthalpies of Reaction The ________ in enthalpy,  H, is the enthalpy of the products minus the enthalpy of the reactants:

32. Thermochemistry Enthalpies of Reaction This quantity,  H, is called the enthalpy of reaction, or the _______________.

33. Thermochemistry The Truth about Enthalpy 1.Enthalpy is an ________ property. 2.  H for a reaction in the forward direction is equal in size, but __________, to  H for the reverse reaction. 3.  H for a reaction depends on the state of the products and the state of the reactants.

34. Thermochemistry Enthalpy Calculation Sucrose is oxdized to carbon dioxide and water. The enthalpy change can be measured in the laboratory as shown C 12 H 22 O 11 + 12O 2 → 12CO 2 + 11H 2 O ΔH = -5645 kJ What is the enthalpy change for the oxidation of 5.00 g of sugar?

35. Thermochemistry Calorimetry Since we cannot know the ________________ ________________, we measure  H through calorimetry, the measurement of ________________.

36. Thermochemistry Heat Capacity and Specific Heat The amount of energy required to raise the temperature of a substance by 1 K (1  C) is its _________________. We define _________________ as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.

37. Thermochemistry Heat Capacity and Specific Heat Specific heat, then, is Specific heat = heat transferred mass  temperature change

38. Thermochemistry Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

39. Thermochemistry Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/mol-K), we can measure  H for the reaction with this equation:

40. Thermochemistry Using Specific Heat Capacity A 192-g piece of copper is heated to 100.0 C in boiling water and then dropped into a beaker containing 751 g of water at 4.0 C. What is the final temperature of the copper and water after thermal equilibrium is reached if the specific heat of copper is 0.385 J/gK? (Assume no heat is lost to warm the beaker or the surroundings.)

41. Thermochemistry Another Example A 15.5-g piece of chromium, heated to 100.0°C, is dropped into 55.5 g of water at 16.5°C. The final temperature of the metal and water is 18.9°C. What is the specific heat capacity of chromium? (Assume no heat is lost to warm the container or the surroundings.)

42. Thermochemistry Bomb Calorimetry Reactions can be carried out in a sealed _______ such as this one, and measure the heat absorbed by the water.

43. Thermochemistry Bomb Calorimetry Because the _________ in the bomb calorimeter is constant, what is measured is really the change in internal energy, ______________ For most reactions, the difference is very _____.

44. Thermochemistry Using a Bomb Calorimeter Methylhydrazine (CH 6 N 2 ) is commonly used as a liquid rocket fuel. The combustion of methylhydrazine is as follows: 2CH 6 N 2 + 5O 2 → 2N 2 + 2CO 2 + 6H 2 O When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00°C to 39.50°C. In a separate experiment the heat capacity of the calorimeter is measured to be 7.794 kJ/°C. What is the heat of reaction for the combustion of a mole of CH 6 N 2 in this calorimeter? Was the reaction endo- or exothermic?

45. Thermochemistry Hess’s Law  H is well known for many reactions, and it is ___________________  H for every reaction in which we are interested. However, we can _______  H using  H values that are _______ and the properties of enthalpy.

46. Thermochemistry Hess’s Law Hess’s law states that “If a reaction is carried out in a series of steps,  H for the overall reaction __________________ __________________ ________________.”

47. Thermochemistry Hess’s Law Because  H is a state function, the ___________________ depends only on the ______________________ ______________________ _____________________.

48. Thermochemistry Enthalpies of Formation An enthalpy of formation,  H f, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.

49. Thermochemistry Standard Enthalpies of Formation Standard enthalpies of formation,  H f, are measured under standard conditions. 

50. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

51. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

52. Thermochemistry Calculation of  H Imagine this as occurring in 3 steps: C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l)

53. Thermochemistry C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) C 3 H 8 (g)  3 C (graphite) + 4 H 2 (g) 3 C (graphite) + 3 O 2 (g)  3 CO 2 (g) 4 H 2 (g) + 2 O 2 (g)  4 H 2 O (l) C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) Calculation of  H The sum of these equations is:

54. Thermochemistry Calculation of  H We can use Hess’s law in this way: where n and m are the stoichiometric coefficients. 

55. Thermochemistry C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) Calculation of  H  H= [3(-393.5 kJ) + 4(-285.8 kJ)] - [1(-103.85 kJ) + 5(0 kJ)] = [(-1180.5 kJ) + (-1143.2 kJ)] - [(-103.85 kJ) + (0 kJ)] = (-2323.7 kJ) - (-103.85 kJ) = -2219.9 kJ Notice, that by definition, __________________ __________________ __________________ _________________, and therefore, the standard enthalpy of O 2 is _____.

56. Thermochemistry Another Example Calculate ΔH for this reaction 2C (s) + H 2 (g) → C 2 H 2 (g) Given the following reactions and their respective enthalpy changes C 2 H 2 (g) + 5/2O 2 (g) → 2CO 2 (g) + H 2 O (l) ΔH = -1299.6 kJ C (s) + O 2 (g) → CO 2 (g) ΔH = -393.5 kJ H 2 (g) + 1/2O 2 (g) → H 2 O (l) ΔH = -1299.6 kJ

57. Thermochemistry You Try This One Suppose you want to know the enthalpy change for the formation of methane, CH4, from solid carbon (as graphite) and hydrogen gas: C(s) + 2H2 (g) → CH4 (g)ΔH = ? The enthalpy change for this reaction cannot be measured in the laboratory because the reaction is very slow. We can, however, measure the enthalpy changes for the combustion of carbon, hydrogen and methane. C (s) + O 2 (g) → CO 2 (g) ΔH = -393.5 kJ H 2 (g) + ½ O 2 (g) → H 2 O (l) ΔH = -285.8 kJ CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O (l) ΔH = -890.3 kJ Use these energies to obtain ΔH for the formation of methane.

58. Thermochemistry Energy in Foods Most of the fuel in the food we eat comes from carbohydrates and fats.

59. Thermochemistry Fuels The vast majority of the energy consumed in this country comes from fossil fuels.

60. Thermochemistry Thermochemistry is Clearly Important to Our Everyday Lives