Introduction

1. BIOENERGY. The sugars in sugarcane, produced from CO2, H2O, and sunshine via photosynthesis, can be converted into ethanol, which is used as an alternative to gasoline. In certain climate, such as that in Brazil, the sugarcane crop replenishes itself rapidly, making cane-based ethanol a sustainable fuel source.

3. Introduction • Energy and Chemistry • Energy is necessary for all life. • Plants, sunlight, and photosynthesis • Human beings derive energy from plants and animals • Energy changes accompany chemical reactions • All of the energy we use are derived from chemical reactions • Thermodynamics is the study of energy and its transformations

4. 5.1 The Nature of Energy • Definition of Energy • Energy is the capacity to do work or to transfer heat • Work is the energy used to cause an object with mass to move against a force • Heat is the energy used to cause the temperature of an object to increase

5. 5.1 The Nature of Energy • Kinetic Energy and Potential Energy • Objects that have mass and motion possess kinetic energy • Potential energy is the stored energy that arise from the attractions and repulsions an object experiences in relation to other objects

6. 5.1 The Nature of Energy • 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), the amount of energy required to raise the temperature of 1 g of water by 1 °C 1 cal = 4.184 J

7. 5.1 The Nature of Energy • System and Surroundings • The system is the portion we single out for study and typically includes the molecules we want to study (here, the hydrogen and oxygen molecules). • The surroundings are everything else (here, the cylinder and piston). • Open systems: matter & heat exchangeable • Closed systems: only heat exchangeable • Isolated systems: no exchange

8. 5.1 The Nature of Energy • Transferring Energy: Work and Heat 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. 5.1 The Nature of Energy • Transferring Energy: Work and Heat • Energy can also be transferred as heat. • Heat flows from warmer objects to cooler objects.

10. 5.1 The Nature of Energy • Transferring Energy: Work and Heat

11. 5.2 The First Law of Thermodynamic • Energy is Conserved 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.

12. 5.2 The First Law of Thermodynamic • Internal Energy The internal energy (E) of a system is the sum of all kinetic and potential energies of all components of the system E = Ek + Ep

13. 5.2 The First Law of Thermodynamic • 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 = Efinal − Einitial

14. 5.2 The First Law of Thermodynamic • Internal Energy • If E > 0, Efinal > Einitial • Therefore, the system absorbed energy from the surroundings.

15. 5.2 The First Law of Thermodynamic • Internal Energy • If E < 0, Efinal < Einitial • Therefore, the system released energy to the surroundings.

16. 5.2 The First Law of Thermodynamic • Relating E to Work and Heat • 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.

17. 5.2 The First Law of Thermodynamic • Relating E to Work and Heat

18. 5.2 The First Law of Thermodynamic • Endothermic and Exothermic Processes

19. 5.2 The First Law of Thermodynamic • State Function Consider chemical changes of H2O: The value of a state function (internal energy, E, in this case) depends only on the present state of the system, not on the path the system took to reach that state

20. 5.2 The First Law of Thermodynamic • State Function • 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.

21. 5.3 Enthalpy • Enthalpy • Consider a chemical reaction: w= -PV

22. 5.3 Enthalpy • Enthalpy • Enthalpy is a measure of the total energy of a system • Enthalpy accounts for heat flow in processes occurring at constant pressure when no forms of work are performed other than P-V work • Enthalpy (H) is expressed by: H = E + PV (state function) • When the system changes at constant pressure, the change in enthalpy, H, is H = (E + PV) = E + PV = (qp+w) − w = qp(state function?) • H equals the heat gained or lost at constant P

23. 5.3 Enthalpy • Enthalpy • When H is positive: endothermic process • When H is negative: exothermic process

25. At constant V, the change in E is equal to the heat gained or lost At constant P, the change in His equal to the heat gained or lost The difference between E and H is the amount of P-V work done by the system when the process occurs at constant P, −PV In many reactions, V is close to zero, which makes PV (the difference between E and H) small Generally H can be used to measure energy changes during most chemical processes

26. 5.3 Enthalpy • H > 0, endothermic • H < 0, exothermic

27. 5.4 Enthalpies of Reaction • Enthalpy of Reaction thermochemical equation enthalpy diagram

28. 5.4 Enthalpies of Reaction This tragedy, in Lakehurst, New Jersey, on May 6, 1937, led to the discontinuation of hydrogen as a buoyant gas in such craft

29. 5.4 Enthalpies of Reaction • Enthalpy of Reaction • Enthalpy is an extensive property. • H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction. • H for a reaction depends on the state of the products and the state of the reactants. for 1 mol of CH4

30. 5.4 Enthalpies of Reaction

31. 5.4 Enthalpies of Reaction

32. 5.5 Calorimetry • Heat capacity and specific heat • Heat capacity, C • The amount of heat required to raise temperature by 1 K (˚C) • Molar heat capacity, Cm • The heat capacity of one mole of a substance • Specific heat, Cs • The heat capacity of one gram of a substance

33. 5.5 Calorimetry • Heat capacity and specific heat

34. 5.5 Calorimetry Sample Exercise 5.5 Relating Heat, Temperature Change, and Heat Capacity (a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 ˚C (about room temperature) to near its boiling point, 98 ˚C? The specific heat of water is 4.18 J/g-K. (b) What is the molar heat capacity of water?

35. 5.5 Calorimetry • “Are you running a fever?” • Difficult to maintain T • Maintaining T in our body • High heat capacity of water • Optimal for muscle function and biochemical rxns in our body at 35.8-37.2 ˚C • Hypothalamus controls body T • Radiation, convection, and evaporation (humidity and water replenishment) • Shivering and reddish skin

36. 5.5 Calorimetry • Constant-pressure Calorimetry • Calorimetry is the measurement of heat flow • Consider a system: • A reaction in a coffee-cup • System: reactants and products • Surrounding: water and calorimeter

37. 5.5 Calorimetry Sample Exercise 5.6 Measuring ΔH Using a Coffee-Cup Calorimeter When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0 ˚C to 27.5 ˚C. Calculate the enthalpy change for the reaction in kJ/molHCl, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g-K. Hmol = -2.7 kJ/0.050 mol = -54 kJ/mol

38. 5.5 Calorimetry • Constant-volume Calorimetry • Bomb calorimeter • Designed for measuring T by combustion reactions • Inlet valve: O2 supply • Electrical contact: initiation • Surrounding: water and calorimeter • The heat capacity of the calorimeter is measured by combusting a standard sample

39. 5.5 Calorimetry Sample Exercise 5.7 Measuring qrxn Using a Bomb Calorimeter Methylhydrazine (CH6N2, MW 46.1 g/mol) is used as a liquid rocket fuel. The combustion of methylhydrazine with oxygen produces N2(g), CO2(g), and H2O(l): 2 CH6N2(l) + 5 O2(g) → 2 N2(g) + 2 CO2(g) + 6 H2O(l) 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. Calculate the heat of reaction for the combustion of a mole of CH6N2.

40. 5.6 Hess’s Law • Combustion of CH4(g) • Consider the combustion reaction of CH4(g) • We can think of the reaction as a two-step process

41. 5.6 Hess’s Law • Hess’s Law • 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 • Based on the fact that enthalpy is a state function • Useful for calculating H that are difficult to measure directly (Ex. Carbon to CO)

42. 5.6 Hess’s Law

43. 5.6 Hess’s Law

44. 5.6 Hess’s Law

45. 5.6 Hess’s Law

46. 5.7 Enthalpies of Formation • An enthalpy of formation, Hf, is defined as the enthalpy change for the formation of a compound from its constituent elements • The conditions of T, P, and state (aq, s, l,…) should be defined. • Standard enthalpy change, H° • H when all reactants and products are in their standard state (1 atm, 273 K or 25 ˚C)

47. 5.7 Enthalpies of Formation • Standard Enthalpy of Formation • The standard enthalpy of formation of a compound, H°f, is the enthalpy change for the reaction that forms one mole of the compound from its elements in their standard states: • By definition, H°fof the most stable form of any element is zero The enthalpy change of the reaction

48. 5.7 Enthalpies of Formation • Standard Enthalpy of Formation

49. 5.7 Enthalpies of Formation • Calculation of reaction enthalpies

50. 5.7 Enthalpies of Formation • Calculation of reaction enthalpies