1. Thermochemistry © 2009, Prentice-Hall, Inc. Chapter 4 Thermochemistry  Thermodynamics Dr.Imededdine Arbi Nehdi Chemistry Department, Science College, King Saud University

2. Thermochemistry © 2009, Prentice-Hall, Inc. The study of energy and its transformations is known as thermodynamics The relationships between chemical reactions and energy changes is an aspect of thermodynamics called thermochemistry

3. Thermochemistry © 2009, Prentice-Hall, Inc. Energy Energy is the ability to do work or transfer heat. –Energy used to cause an object that has mass to move is called work. –Energy used to cause the temperature of an object to rise is called heat.

4. Thermochemistry © 2009, Prentice-Hall, Inc. Potential Energy Potential energy is energy an object possesses by virtue of its position or chemical composition. Fig. 4.1 A bicycle at the top of a hill (left) has a high potential energy. As the bicycle proceeds down the hill (right), the potential energy is converted into kinetic energy; the potential energy is lower at the bottom than at the top.

5. Thermochemistry © 2009, Prentice-Hall, Inc. Kinetic Energy Kinetic energy is energy an object possesses by virtue of its motion. 1 2 E k =  mv 2 where m is mass, v is the speed, and E k is the kinetic energy.

6. Thermochemistry © 2009, Prentice-Hall, Inc. 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 1 J = 1  kg m 2 s2s2

7. Thermochemistry © 2009, Prentice-Hall, Inc. Definitions: System and Surroundings The system includes the molecules we want to study (here, the hydrogen and oxygen molecules). The surroundings are everything else (here, the cylinder and piston). If the hydrogen and oxygen react to form water, energy is liberated. 2 H 2 (g) + O 2 (g) 2 H 2 O (l) + energy The system has not lost or gained mass; it undergoes no exchange of matter with its surrounding. However, it does exchange energy with its surroundings in the form of work and heat.

8. Thermochemistry © 2009, Prentice-Hall, Inc. Definitions: Work Energy used to move an object over some distance is work. w = F  d where w is work, F is the force, and d is the distance over which the force is exerted.

9. Thermochemistry © 2009, Prentice-Hall, Inc. Heat Energy can also be transferred as heat. Heat flows from warmer objects to cooler objects.

10. Thermochemistry © 2009, Prentice-Hall, Inc. We can now provide a more precise definition of energy: Energy is the capacity to do work or transfer heat.

11. Thermochemistry © 2009, Prentice-Hall, Inc. Conversion of Energy Energy can be converted from one type to another. For example, the cyclist above has potential energy as she sits on top of the hill.

12. Thermochemistry © 2009, Prentice-Hall, Inc. Conversion of Energy As she coasts down the hill, her potential energy is converted to kinetic energy. At the bottom, all the potential energy she had at the top of the hill is now kinetic energy.

13. Thermochemistry © 2009, Prentice-Hall, Inc. First Law of Thermodynamics Energy is neither created nor destroyed. 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.

14. Thermochemistry © 2009, Prentice-Hall, Inc. 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.

15. Thermochemistry © 2009, Prentice-Hall, Inc. Internal Energy By definition, the change in internal energy,  E, is the final energy of the system minus the initial energy of the system:  E = E final − E initial

16. Thermochemistry © 2009, Prentice-Hall, Inc. Changes in Internal Energy If  E > 0, E final > E initial –Therefore, the system absorbed energy from the surroundings. –This energy change is called endergonic.

17. Thermochemistry © 2009, Prentice-Hall, Inc. Changes in Internal Energy If  E < 0, E final < E initial –Therefore, the system released energy to the surroundings. –This energy change is called exergonic.

18. Thermochemistry © 2009, Prentice-Hall, Inc. 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,  E = q + w.

19. Thermochemistry © 2009, Prentice-Hall, Inc.  E, q, w, and Their Signs

20. Thermochemistry © 2009, Prentice-Hall, Inc. Exchange of Heat between System and Surroundings When heat is absorbed by the system from the surroundings, the process is endothermic.

21. Thermochemistry © 2009, Prentice-Hall, Inc. Exchange of Heat between System and Surroundings When heat is absorbed by the system from the surroundings, the process is endothermic. When heat is released by the system into the surroundings, the process is exothermic.

22. Thermochemistry © 2009, Prentice-Hall, Inc. State Functions Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex problem.

23. Thermochemistry © 2009, Prentice-Hall, Inc. 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.

24. Thermochemistry © 2009, Prentice-Hall, Inc. State Functions Therefore, internal energy is a state function. It depends only on the present state of the system, not on the path by which the system arrived at that state. And so,  E depends only on E initial and E final. The value of state function does not depend on the particular history of the sample, only on its present condition.

25. Thermochemistry © 2009, Prentice-Hall, Inc. State Functions However, q and w are not 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. a) A battery shorted out by a wire loses energy to the surroundings as heat; no work is performed. b) A battery discharged through a motor loses energy as work (to make the fan turn|). Some heat will be released to the surroundings.

26. Thermochemistry © 2009, Prentice-Hall, Inc. Work Usually in an open container the only work done is by a gas pushing on the surroundings (or by the surroundings pushing on the gas).

27. Thermochemistry © 2009, Prentice-Hall, Inc. 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. w = -P  V

28. Thermochemistry © 2009, Prentice-Hall, Inc. Enthalpy If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system. Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV Enthalpy is a state function

29. Thermochemistry © 2009, Prentice-Hall, Inc. 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

30. Thermochemistry © 2009, Prentice-Hall, Inc. Enthalpy Since  E = q + w and w = -P  V, we can substitute these into the enthalpy expression:  H =  E + P  V  H = (q+w) − w  H = q So, at constant pressure, the change in enthalpy is the heat gained or lost by the system when the process occurs under constant pressure.  H = H final – H initial = q p

31. Thermochemistry © 2009, Prentice-Hall, Inc. Endothermicity and Exothermicity A process is endothermic when  H is positive.

32. Thermochemistry © 2009, Prentice-Hall, Inc. Endothermicity and Exothermicity A process is endothermic when  H is positive. A process is exothermic when  H is negative.

33. Thermochemistry © 2009, Prentice-Hall, Inc. Enthalpy of Reaction The change in enthalpy,  H, is the enthalpy of the products minus the enthalpy of the reactants:  H = H products − H reactants - It is found experimentally that 890 kJ of heat is produced when 1 mol of CH 4 is burned in a constant-pressure system. -The combustion of 2 mol of CH 4 with 4 mol of O 2 -releases twice as much heat, 1780 kJ.

34. Thermochemistry © 2009, Prentice-Hall, Inc. Enthalpy of Reaction This quantity,  H, is called the enthalpy of reaction, or the heat of reaction. Fig. a)A candle is held near a balloon filled with H 2 gas and O 2 gas. b) The H 2 (g) ignites, reacting with O 2 (g) to form H 2 O (g). The resultant explosion produces the yellow ball of flame. The system gives off heat to its surroundings. c) The enthalpy diagram of this reaction. (a) (b) (c)

35. Thermochemistry © 2009, Prentice-Hall, Inc. The Truth about Enthalpy 1.Enthalpy is an extensive property. 2.  H for a reaction in the forward direction is equal in size, but opposite in sign, to  H for the reverse reaction. 3.  H for a reaction depends on the state of the products and the state of the reactants.

36. Thermochemistry © 2009, Prentice-Hall, Inc. Calorimetry Since we cannot know the exact enthalpy of the reactants and products, we measure  H through calorimetry, the measurement of heat flow by measuring the temperature change it produces. Calorimeter

37. Thermochemistry © 2009, Prentice-Hall, Inc. Heat Capacity and Specific Heat The amount of energy required to raise the temperature of a substance by 1 K (1  C) is its heat capacity.

38. Thermochemistry © 2009, Prentice-Hall, Inc. Heat Capacity and Specific Heat We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.

39. Thermochemistry © 2009, Prentice-Hall, Inc. Heat Capacity and Specific Heat So if we known (q) we can calculate specific heat: C s = q m  Tm  T The amount of heat (q) transferred (gained or loosed) by a masse (m) of substance with a specific heat capacity (Cs) and temperature change of  T is: q = Cs x m x  T The molar heat capacity (Cm) of a substance is: Cm = Cs x M M: molar mass of the substance (g/mol)

40. Thermochemistry © 2009, Prentice-Hall, Inc. 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. -This calorimeter is used for the reactions occurring in solution. -Because the calorimeter is not sealed, the reaction occurs under atmospheric pressure. Fig. Coffee-cup calorimeter, in which reactions occur at constant pressure.

41. Thermochemistry © 2009, Prentice-Hall, Inc. Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/g-K), we can measure  H for the reaction with this equation: q soln = m  C s   T q soln = - q rxn q soln : the heat gained (loosed) by the solution q rxn : the heat gained (loosed) by the reaction m : masse of solution Cs : specific heat of solution (  4.184 J/g-K)

42. Thermochemistry © 2009, Prentice-Hall, Inc. Bomb Calorimetry Reactions can be carried out in a sealed “bomb” such as this one. The heat absorbed (or released) by the water is a very good approximation of the enthalpy change for the reaction. Combustion reactions are the most conveniently studied using bomb calorimeter. After the sample (organic compound) has been placed in the bomb, the bomb is sealed and pressurized with oxygen. Fig. Cutaway view of a bomb calorimeter, in witch reactions occur at constant volume.

43. Thermochemistry © 2009, Prentice-Hall, Inc. Bomb Calorimetry Combustion reactions are the most conveniently studied using bomb calorimeter. After the sample (organic compound) has been placed in the bomb, the bomb is sealed and pressurized with oxygen. It is then placed in the calorimeter, which is essentially an insulated container, and covered with an accurately measured quantity of water. When all the components within the calorimeter have come to the same temperature, the combustion reaction is initiated by passing an electrical current through a fine wire that is in contact with the sample. When the wire gets sufficiently hot, the sample ignites. Fig. Cutaway view of a bomb calorimeter, in witch reactions occur at constant volume.

44. Thermochemistry © 2009, Prentice-Hall, Inc. Bomb Calorimetry Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy,  E, not  H. For most reactions, the difference is very small. So the heat of the reaction (q rxn) is : Ccal : Heat capacity of the calorimeter  T:the temperature change of calorimeter q rxn = - C cal x  T

45. Thermochemistry © 2009, Prentice-Hall, Inc. Hess’s Law  H is well known for many reactions, and it is inconvenient to measure  H for every reaction in which we are interested. However, we can estimate  H using published  H values and the properties of enthalpy.

46. Thermochemistry © 2009, Prentice-Hall, Inc. Hess’s Law Hess’s law states that “if a reaction is carried out in a series of steps,  H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”

47. Thermochemistry © 2009, Prentice-Hall, Inc. Hess’s Law Because  H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.

48. Thermochemistry © 2009, Prentice-Hall, Inc. 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. 2C (graphite) + 3H 2 (g) + ½ O 2 (g) C 2 H 5 OH (l)  f H  = -277.7 kJ

49. Thermochemistry © 2009, Prentice-Hall, Inc.

50. Thermochemistry © 2009, Prentice-Hall, Inc. Standard Enthalpies of Formation Standard enthalpies of formation,  H f °, are measured under standard conditions (25 °C and 1.00 atm pressure). By definition, the standard enthalpy of formation of the most stable form of any element is zero (  H f (O 2 g) = 0, …)

51. Thermochemistry © 2009, Prentice-Hall, Inc. Calculation of  H of reaction using enthalpies of formation Imagine this as occurring in three 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 © 2009, Prentice-Hall, Inc. Calculation of  H Imagine this as occurring in three 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 © 2009, Prentice-Hall, Inc. Calculation of  H Imagine this as occurring in three 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)

54. Thermochemistry © 2009, Prentice-Hall, Inc. 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:  H 1 = -  H f °(C 3 H 8 (g)),  H 2 = 3  H f °(CO 2 (g)),  H 3 = 4  H f °(H 2 O(l)),  H rxn =  H 1 +  H 2 +  H 3 = 2220 kJ

55. Thermochemistry © 2009, Prentice-Hall, Inc. Calculation of  H We can use Hess’s law in this way:  H =   n  H f ° products –   m  H f ° reactants where n and m are the stoichiometric coefficients.

56. Thermochemistry © 2009, Prentice-Hall, Inc.  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 C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O (l) Calculation of  H