Energy

1. Energy • Energy transfer occurs between a system and its surroundings. • System = piece of the universe.

2. Energy • Two types: heat and work.

3. Energy • Heat = energy transfer that is the result of contact between two substances of differing temperatures.

4. Energy • Energy can also be transferred between a system and its surroundings.

5. Energy • The system may do work on the surroundings, or the surroundings may do work on the system.

6. Energy • Work = F*d (a force is used to move an object). • Ex. Cu strip being hammered. • Ex. Gas engine.

7. Energy • The state of a system is defined by its composition, temperature, and pressure.

8. Energy • State Property: quality that is dependent only on the state of the system and not how that system reached that state.

9. Energy • Example: reactions under constant volume and pressure do not do expansive work - their q values are different!

10. Conversions • The joule is the SI unit of energy (1J = 1N*m)

11. Conversions • 1 calorie = 4.184 J • Dietary calorie = C = 1000 c = 4184 J • 1 kJ = 1000 J

12. Exo vs Endothermic • Exothermic: reaction releases heat into the surroundings. • Endothermic: reaction absorbs heat.

13. Exo vs Endothermic • Exothermic: products formed are at lower energy than the reactants. (usually spontaneous reaction)

14. Exo vs Endothermic • Endothermic: products formed are at higher energy than the reactants (usually non-spontaneous).

15. Exo vs Endothermic • For a reaction at constant P and T: endothermic: q = delta H > 0 • Exothermic q = delta H < 0

16. Activation Energy • A reaction requires a collision among reactants. Activation energy is the minimal energy required to make this happen.

17. Activation Energy • The net energy transfer is what determines whether the reaction is exothermic or endothermic.

18. Calorimetry • The study of heat transfer using a calorimeter (a device that isolates a system and measures its temperature change)

19. Calorimetry • Specific heat: the heat needed to raise 1 gram of a substance 1 degree Celsius. • Different substances change temperature differently.

20. Calorimetry • Specific heat of water = 4.184 J/g*C • q = mass*delta T*specific heat

21. Calorimetry • Ex. Suppose 652 J of heat is added to 15.0 g of water originally at 20.0o C. What is the final t? • t = 30.4o C

22. Calorimetry • In a coffee-cup calorimeter, delta H reaction = -q water. Heat given off by the reaction is absorbed by the water in the cup.

23. Calorimetry • Ex. Suppose heat is absorbed by 412 g of water, increasing its t from 20.12 to 29.86oC. What is delta H? • delta H = -16.8 kJ

24. Calorimetry • In bomb calorimeters, some heat is absorbed by the metal as well as the water.

25. Calorimetry • q reaction = -q calorimeter = -(C calorimeter) x delta t • C calorimeter = total heat capacity of the bomb + water.

26. Calorimetry • Ex. Suppose combustion of 1.60 g of methane in a bomb calorimeter raises the t by 5.14o C. What is the q reaction (C = 17.2 kJ/oC)? • q reaction = -88.4 kJ

27. Phase changes • Heat lost = heat gained • Process: melting to boiling • Heat of fusion, Heat of vaporization • Q = mass * heat of fusion

28. Phase changes • Hfof water = 333 J/g • Hv of water = 2260 J/g • Specific heat of ice = 2.06 J/g*C • Steam = 2.02 J/g*C

29. Phase Change How much energy is required to raise the temperature of 2.50g of water from -3.00 C to 108.00 C? qtot = 7.58 kJ

30. Thermochem • Rule 1: H is directly proportional to the amount of reactants or products. • (stoichiometry)

31. Thermochem • Rule 2: H for a reaction is equal in magnitude but opposite in sign to H for the reverse reaction or process. • A(s) --> A(l) kJ • A(l) --> A(s) kJ

32. Thermochem • Rule 3: Reactions occur in steps. Hess’ Law: If equation 1 + equation 2 = equation 3 (the steps result in the overall net equation), then H3 = H1 +H2 .

33. Thermochem Ex: step 1 - C(s) + 1/2 O2(g) --> CO(g); step 2 - CO(g) + 1/2 O2(g) --> CO2(g) H = -283.0 kJ; Hnet = (-393.5 kJ). What is the net equation and the H of the first step? H1 = -110.5 kJ, C(s) + O2(g) --> CO2(g)

34. Heats of Formation • Hof of a compound = H when one mole of the compound is formed from its elements in their stable state.

35. Thermochem • Ex. 2 Ag(s) + Cl2(g) --> 2 AgCl(s) Ho = -254.0 kJ • What is Hof for silver chloride? • Hof = -127.0 kJ

36. Thermochem • Ex. HgO(s) --> Hg(l) + 1/2 O2(g) Ho = +90.8 kJ • What is Hof for mercury(II) oxide? • Hof = -90.8 kJ

37. Heats of Formation • For any thermochemical equation: • Ho = f products) - f reactants) • Note: the heat of formation of an element in its stable state = 0

38. Thermochem • Ex. (use Table 8.3) • Calculate the standard enthalpy change (o) for the combustion of methane. • Ho = -890.3 kJ

39. Heats of Formation • Can also apply to ions if you take the standard heat of formation of the hydrogen ion to be zero. (see Table 8.3 bottom)

40. Thermochem • Ex. Calculate the standard enthalpy change (o) for Zn(s) + 2H+(aq) --> Zn2+ (aq) + H2(g). • Ho = -153.9 kJ

41. Bond Energies • Defined as H when one mole of bonds is broken in the gaseous state. • Ex. Cl2(g) --> 2Cl(g) • H = B.E. Cl-Cl = 243 kJ

42. Bond Energies • In general, since multiple bonds involve more bonding pairs, they tend to be stronger than single bonds.

43. Bond Energies • Ex. Use bond energies to to determine if the following is an exothermic or endothermic reaction: • 2 CO(g) + O2(g) --> 2CO2(g). • exothermic

44. Thermodynamics • First law: E = q + w where the change in energy of a system, q = heat flow into the system, and w = work done on the system.

45. Thermodynamics • Esystem = -Esurroundings. • At constant volume: w=0, qv=E • At constant pressure: w= -PV; q = H = E + ngRT