Ch. 6 Thermochemistry

1. Ch. 6 Thermochemistry • The relationship between chemistry and energy • Basic concept of thermodynamics • Energy conversion: • Energy: the capacity to do work or to produce heat • Law of conservation of energy: energy can be converted from one form to another but can be neither created nor destroyed.

2. Basic concepts of thermochemistry Energy is classified as: Potential energy: energy due to position or composition Kinetic energy: energy due to the motion of an object KE = ½ mv2

3. Figure 6.1 The Total Energy of the Universe is Constant

4. Basic concepts of thermochemistry Two ways of the energy transfer: Heat: energy transferred between two objects due to a temperature difference between them. Work: force acting over a distance (w = f x d) Ex. Gas expansion and compression P = Work = force x distance = F xΔh Since P = F/A or F = P x A, then Work = P x A xΔh For a cylinder device, ΔV= finial volume – initial volume = A x Δh Then, Work = P x A x Δh = PΔV

5. Figure 6.4 The Piston, Moving a Distance Against a Pressure P, Does Work On the Surroundings Work=PΔV W= -PΔV

6. Figure 6.2 Exothermic Process(Energy flows out of the system) The energy gained by the surroundings must be equal to the energy lost by the system. In any exothermic reaction, some of the potential energy stored in the chemical bonds is being converted to the thermal energy via heat.

7. System: the object being studied at a given moment. a macroscopic system a closed system an open system an adiabatic system an isolated system Surroundings: The portion of universe that is outside the system.

8. Figure 6.3 Endothermic Process(heat flow is into the system)

9. Basic concepts of thermochemistry Unit of energy: SI unit of energy: joule = KE = ½ (mv2)=1/2 x 2.0Kg (1m/s)2 (if m = 2.0 Kg and v = 1 m/s) = 1.0 = 1.0 joule 1 Calorie = the heat required to raise the temperature of 1.00 g pure water 1.00C (from 14.5 to 15.50C) 1 cal = 4.184 joules 1Kcal = 1000 calories

10. Heat capacity C = • ∴ q = CΔT • Specific heat capacity (s) : the heat capacity is given per gram of substance. (J/oC·g or J/K·g) • Molar heat capacity : the heat capacity is given per mole of substance. (J/oC·mol or J/K·mol) • = s/ molar mass • Using specific heat capacity: energy exchanged = C· m ·T • ex.

12. Figure 6.5 A Coffee-Cup Calorimeter Made of Two Styrofoam Cups

13. Thermodynamics: the study of energy and its interconversions. Thermochemistry: the relationship between chemical reaction and energy change. The 1st law of thermodynamics: the energy of universe is constant. (能量守恆定律) The internal energy (E) of a system is the sum of KE and PE of all particles in the system. The internal energy of a system can be changed by a flow of work, heat, or both. i.e. ΔE = q + w ΔE: change of E, q: heat, w: work Thermodynamic quantities always consist of two parts: a number, giving the magnitude of the change a sign, indicating the direction of the flow Ex. A quantity of energy flows into the system via heat (an endothermic process), +x, and -x for an exothermic process.

14. Ex 6.1, 6.2, 6.3

15. Figure 6.4 The Piston, Moving a Distance Against a Pressure P, Does Work On the Surroundings w =PΔV w = -PΔV • ΔV > 0, expansion → w < 0, system does work to the surroundings • ΔV < 0, compression → w > 0, work has done on the system

16. QUESTION 1 The combustion of a fuel is an exothermic process. This means… 1. the surroundings have lost exactly the amount of energy gained by the system. 2. the potential energy of the chemical bonds in the products should be less than the potential energy of the chemical bonds in the reactants. 3. q must be positive; w must be negative 4. DE would have a + overall value because the surroundings have gained energy.

17. QUESTION 2

18. QUESTION 3

19. State function: A property of a system that is determined by the state of condition of the system and not by the how it got to the state; its value is fixed when temperature, pressure, composition, and physical form are specified; P, V, T, E, and H are state functions. Because E is a state function, ΔE depends only on the initial and finial states of the system and not how the change occurs. Ex. of textbook: Denver (elevation 5180 ft) , Chicago (elevation 674 ft) Δh = 5280-674 = 4606 ft Elevation is a state function and distance is not. State function: A quantity whose changed value is determined only by its initial and finial vales.

20. Figure 6.6 A Bomb Calorimeter.(constant volume) E = q + w = q = qv

21. Figure 6.6 A Bomb Calorimeter

22. Heat transfer at const. V: Bomb calorimeter ΔE = qv • Heat transfer at const. P : Enthalpy ΔE = qp + w = qp –PextΔV ΔE = qp –PΔV, (Pext = Pinternal = Psystem) qp = ΔE + PΔV ; PΔV= Δ(PV) for const. P qp = Δ(E + PV) Enthalpy (焓) H = E + PV  qp = Δ(E + PV) = ΔH ΔH = qp = ΔE + PΔV (only at constant pressure) ∵ E, P, and V are state functions; H is a state function • The enthalpy changed,  H, is the heat added to (or lost by) a system. • For a chemical reaction, the enthalpy change is:  H = Hproducts - Hreactants

23. A Group of Firewalkers in Japan

24. Thermochemistry Energy: 1. kinetic energy 2. potential energy 3. internal energy (E) E: Total energy of the system; due to K.E. due to P.E. due to chemical energy stored in chemical bond: thermochemistry Reaction enthalpies: qreaction (at cont. P) = ΔH = Hproducts – Hreactants ΔH > 0; endothermic reaction ΔH < 0; exothermic reaction CO(g) + O2 (g) → CO2 (g)ΔH = -283.0 kJ CO2 (g) → CO(g) + O2 (g)ΔH = +283.0 kJ 2CO2 (g) → 2 CO(g) + O2 (g)ΔH = +566.0 kJ

25. Hess’s Law

26. Hess’s Law C(s) + O2 (g)→ CO2 (g)ΔH = -393.5 kJ CO2 (g) → CO(g) + O2 (g)ΔH = +283.0 kJ -------------------------------------------------------------------- C(s) + O2 (g) → CO(g)ΔH = ? equation = equation (1) + equation (2) +ΔH = ΔH1 + ΔH2 +  Note: The sign of ΔH indicates whether the reaction is endothermic or exothermic reaction. ΔH change sign when a reaction is reversed. The magnitude of ΔH is directly proportional to the quantities of reactants and products, the coefficients represent number of mole. (ΔH  amount of reactants or products) The phases (physical states) of all species must be specified, using the symbol of (s), (l), and (g).

27. Given H2 (g) + O2 (g)→ H2O (l)ΔH = -285.8 kJCalculate ΔH for the equation 2H2O (l) → 2H2 (g) + O2 (g)ΔH = ? H2O(s)→ H2O(l)ΔHfus = +6.007 kJ mol-1 H2O(l) → H2O(s)ΔHfreez = -6.007 kJ mol-1 ΔHfus = -ΔHfreez

28. Graphite

29. Standard-State Enthalpies (ΔHo) Absolute Energies and absolute enthalpies cannot be measured and calculated. Only ΔH and Δ , sometime, the measurement of ΔH is impossible, ΔH can be obtained by calE can be measured. Standard-States for chemical substances: For compounds For pure solids and liquids (condensed states), pure liquid or solid at 1 atm and specified temp (298.15 oK) For gases, 1 atm and specified temp. For dissolved species, 1 M solution at 1 atm and specified temp. For elements The standard states of elements are the forms which the elements exit under conditions of 1 atm and 25oC. The standard state for oxygen is O2(g); the standard state for sodium is Na(s); the standard state for mercury is Hg(l); etc.

30. Standard enthalpy of formation (ΔHfo) def: The enthalpy change for the reaction that produces one mole of the compound from its elements in their stable states. . H2 (g) + O2 (g)→ H2O (l)ΔH = -285.83 kJ ΔHfo(H2 O(l)) = -285.83 kJ mol-1 H2 (g) → H(g)ΔH = +217.96 kJ ΔHfo(H (g)) = +217.96 kJmol-1

31. Table 6.2 Standard Enthalpies of Formation for Several Compounds at 25°C

32. Figure 6.8 Pathway for the Combustion of Methane

33. Figure 6.9 Schematic Diagram of Energy Changes

34. Figure 6.10 A Pathway for the Combustion of Ammonia

35. Figure 6.11 Energy Sources Used in the United States

39. Figure 6.12 The Atmosphere Traps Some Light Energy, Keeping the Earth Warmer than it Otherwise Would Be

40. Figure 6.14 Coal Gasification