1. William L Masterton Cecile N. Hurley http://academic.cengage.com/chemistry/masterton Edward J. Neth University of Connecticut Chapter 8 Thermochemistry

2. Heat (Not in Notes) Heat will flow from a hotter object to a colder object Mix boiling water with ice Temperature of the ice rises after it melts Temperature of the water falls

3. 8.1 Principles of Heat Flow Thermodynamics: the study of energy changes in chemical reactions and the influence of energy on those changes. system: that part of the universe on which attention is focused surroundings: the rest of the universe Heat: thermal energy transferred between the system and the surroundings Temperature: a measure of the average kinetic energy of the particles in a system

4. Figure 6.1 – Systems and Surroundings

5. Chemical Reactions: When we study a chemical reaction, we consider the system to be the reactants and products The surroundings are the vessel (beaker, test tube, flask) and air “surrounding” the reaction “system”

6. Direction of Heat flow: Heat will flow from a hotter object to a colder object Example: Mix boiling water with ice Temperature of the ice rises after it melts Temperature of the water falls

7. Heat is given the symbol, q q is positive when heat flows into the system from the surroundings q is negative when heat flows out of the system into the surroundings

8. Heat flow continued: Endothermic processes have positive q Endothermic reaction – a chemical reaction that absorbs energy as heat from its surroundings H 2 O (s)  H 2 O (ℓ)q > 0 Exothermic processes have negative q Exothermic reaction – a chemical reaction that releases energy as heat into its surroundings CH 4 (g) + 2O 2 (g)  CO 2 (g) + H 2 O (ℓ) q < 0

9. Exothermic and Endothermic Processes

10. Magnitude of Heat Flow (q) – the amount of heat transfer James Joule: (1818-1889) calorimetry calorie: English unit of measure for heat; 1 calorie is equal to the amount of energy required to increase the temperature of a 1g sample of water 1°C Calorie: Nutritional Calorie; 1,000cal = 1Cal

11. Joule: S.I. unit – Joules(J) or kilojoules(kJ) are the metric measures of heat Conversions (calories to Joules): 1 calorie = 4.184 J 1 kilocalorie = 1 Calorie ( nutritional calorie) Nutritional calories are kcal

12. The Calorimetry Equation: q = m x c x  t c (lowercase) is the specific heat: the quantity of heat needed to raise the temperature of one gram of a substance by 1 °C c depends on the identity and phase of the substance

13. Specific Heat The specific heat of a substance, like the density or melting point, is an intensive property that can be used to identify a substance or determine its purity

14. Water Water has an unusually large specific heat A large quantity of heat is required to raise the temperature of water Climate is moderated by the specific heat of water

15. Table 8.1

16. Example 8.1: Compare the amount of heat given off by 1.40g of liquid water when it cools from 100.0°C to 10.0°C to that given off when 1.40g of steam cools from 200.0°C to 110.0°C.

17. Measurement of Heat Flow Calorimeter: a device used to measure the heat flow of a reaction The walls of the calorimeter are insulated to block heat flow between the reaction and the surroundings The heat flow for the system is equal in magnitude and opposite in sign from the heat flow of the calorimeter q reaction = - q calorimeter

18. Figure 8.2

19. Coffee-cup Calorimeter We disregard the (-) in calculations and label q’s sign based on the process being endo or exo.

20. Example 8.2

21. Heat calculation for mixing of solutions: If the concentrations are less than or equal to 1.0M, you can assume the density to be the same as water 1g/mL, and you must add the masses together to get the total mass of the combined solution.

22. Figure 8.3

23. Bomb Calorimeter The bomb calorimeter is more versatile than the coffee-cup calorimeter Reactions involving high temperatures and gases A heavy metal vessel surrounded by water Often used to determine the caloric content of foods

24. Enthalpy The heat flow at constant pressure is equal to the difference in enthalpy (heat content) between products and reactants The symbol for enthalpy is H We measure changes in enthalpy using a calorimeter and a reaction run at constant pressure: ΔH = H products – H reactants The sign of the enthalpy change is the same as for heat flow: ΔH > 0 for endothermic reactions ΔH < 0 for exothermic reactions Enthalpy is a state variable

25. Figure 8.4 – Enthalpy of Reaction

26. Thermochemical Equations A thermochemical equation is a chemical equation with the ΔH for the reaction included Endothermic Example – energy is added to the system NH 4 NO 3 (s)  NH 4 + (aq) + NO 3 - (aq)  H = +28.1 kJ

27. Exothermic Example – energy is released from the system CH 4(g) +2O 2(g)  CO 2(g) + 2H 2 O (l) + 891kJ

28. Figure 8.5 – An Endothermic Reaction

29. Conventions of Thermochemistry: 1. The sign of  H indicates whether the reaction is endothermic or exothermic 2. The coefficients of the thermochemical equation represent the number of moles of reactant and product 3. The phases of all reactant and product species must be stated 4. The value of  H applies when products and reactants are at the same temperature, usually 25 °C

30. Rules of Thermochemistry: 1. The magnitude of  H is directly proportional to the amount of reactant or product 2.  H for the reaction is equal in magnitude but opposite in sign for  H for the reverse of the reaction 3. The value of  H is the same whether the reaction occurs in one step or as a series of steps This rule is a direct consequence of the fact that ΔH is a state variable This rule is a statement of Hess’s Law

31. Example 8.4

32. Enthalpy of Phase Changes Phase changes involve enthalpy There is no change in temperature during a phase change Endothermic: melting or vaporization Exothermic: freezing or condensation Pure substances have a value of ΔH that corresponds to melting (reverse, fusion) or vaporization (reverse, condensation)

33. Hess’s Law: The heat of a reaction (  H ) is constant, whether the reaction is carried out directly in one step or indirectly through a number of steps. The heat of reaction (  H ) can be determined as the sum of the heats of reaction of several steps

34. Example 8.5

35. Example 8.6

36. Enthalpies of Formation The standard molar enthalpy of formation,, is equal to the enthalpy change For one mole of a compound At constant pressure of 1 atm At a fixed temperature of 25 °C From elements in their stable states at that temperature and pressure Enthalpies of formation start on page 4 of the Reference Book

37. The standard enthalpy of formation of a pure element in its standard state at 25 °C is zero The enthalpy of formation of H + (aq) is also zero

38. Calculation of The symbol Σ refers to “the sum of” Elements in their standard states may be omitted, as their enthalpies of formation are zero (standard state only) The coefficients of reactants and products in the balanced equation must be accounted for

39. Example 8.7

40. Example 8.8

41. The First Law of Thermodynamics Thermodynamics Deals with all kinds of energy effects in all kinds of processes Two types of energy Heat (q) Work (w) The Law of Conservation of Energy  E system = -  E surroundings The First Law  E = q + w The total change in energy is equal to the sum of the heat and work transferred between the system and the surroundings

42. Conventions q and w are positive When the heat or work enters the system from the surroundings q and w are negative When the heat or work leaves the system for the surroundings

43. Figure 8.10

44. Example 8.9

45. Spontaneity of Reactions Spontaneous Processes - Everyday process that take place on their own, without outside forces What is a spontaneous process? (examples) An ice cube will melt when added to a glass of water at room temperature A mixture of hydrogen and oxygen will form water when a spark is applied An iron (or steel) tool will rust if exposed to moist air

46. Spontaneity Spontaneity and rate are not connected; FAST DOES NOT = SPONTANEOUS If a reaction is spontaneous in one direction, it will be non-spontaneous in the reverse direction under the same conditions

47. Spontaneity and Equilibrium A spontaneous process is one that moves a reaction system toward equilibrium, and a nonspontaneous process moves away from equilibrium (equilibrium = a state of no net change in a process)

48. Which of the following are spontaneous processes? 1. Snowman melting in the sun 2. Assembling a jigsaw puzzle 3. Rusting of an iron object in humid air 4. Recharging of a camera battery

49. The Energy Factor Many spontaneous reactions are exothermic, but not all! Many spontaneous processes proceed with a decrease in energy Boulders roll downhill Your cell phone battery discharges over time Recall that exothermic reactions proceed with a decrease in energy Spontaneous reactions are typically exothermic

50. Exceptions phase changes H 2 O (s)  H 2 O (l) is endothermic but spontaneous at room temperature Some reactions become spontaneous with a simple increase in temperature CaCO 3 (s)  CaO (s) + CO 2 (g) ΔH = +178.3 kJ ΔH is not the only criterion for spontaneity

51. The Randomness Factor Nature tends to move spontaneously from a state of lower probability (order) to one of higher probability (disorder), or Systems overtime without outside influence, will move toward a condition of maximum probability (disorder).

52. Figure 17.2

53. Entropy, S What is entropy? described as an increase in disorder or randomness; symbol or variable S. A measure of the amount of randomess in a system When water vaporizes the entropy of the gas molecules is greater than the liquid Kinetic molecular theory supports greater motion will result in greater entropy

54. Figure 17.3 – Disorder and Order

55. Factors that Influence the amount of entropy A liquid has higher entropy than the solid from which it formed A gas has higher entropy than the liquid or solid from which it formed Increasing the temperature of a substance increases entropy

56. Figure 17.4

57. The Third Law of Thermodynamics A completely ordered, pure crystalline solid at 0 K has an entropy of zero This is the only time the entropy of a substance is zero Absolute zero has not been reached; it is still a theoretical limit

58. Example 17.1

59. Standard Molar Entropies Unlike enthalpy, molar entropy cannot be directly measured Elements have nonzero molar entropies In calculating the standard molar entropy change, elements must be taken into account Standard molar entropies are always positive numbers, i.e.,  S° > 0 Aqueous ions may have negative S°values

60. ΔS° for Reactions Units are J/mol-K or kJ/mol-K

61. The Second Law of Thermodynamics In a spontaneous process, there is a net increase in entropy, taking into consideration both the system and the surroundings That is, for a spontaneous process:

62. Free Energy, G Gibbs Free Energy If  G is negative, < 0, the reaction is spontaneous If  G is positive, > 0, the reaction is nonspontaneous as written (the reverse reaction is spontaneous) If  G = 0, the system is in equilibrium; no tendency for the reaction to occur in either direction

63. Figure 17.5

64. Relation Among ΔG, ΔH and ΔS The Gibbs-Helmholtz equation  G =  H – T  S Spontaneous reactions generally have  H 0 In specific cases, either term may dominate phase changes,  S is dominant; some reactions,  S is nearly zero and  H will dominate

65. 17.4 The Standard Free Energy Change, ΔG° Standard conditions: Gases are at 1 atm pressure Solutions are 1M for ions or molecules Under standard conditions:  G° =  H° – T  S° Recall that If ΔG° < 0, spontaneous If  G° > 0, nonspontaneous If  G° = 0, equilibrium

66. Free Energy of Formation We can use the Gibbs-Helmholtz equation to calculate the standard free energy of formation for a compound The sign of If negative, the formation of the compound is spontaneous If positive, the formation of the compound is nonspontaneous

67. Example 17.3

68. Effect of Temperature, pressure and concentration on reaction spontaneity. Table in notes

69. Example 17.4 Predict which of the four cases in the table above will apply to the following reactions, and if the reaction will be spontaneous. Use ∆H as given and your estimate of the sign of ∆S. C 6 H 12 O 6 (s) +6O 2(g)  6CO 2(g) +6H 2 O (g) ΔH = -2540 kJ Cl 2(g)  2Cl (g) ΔH is positive

70. Standard Free Energy Change A. Equation at standard conditions: ΔG o = ΔH o - T ΔS o