Webpage: http ://blackboard.latech/ Dr. Upali Siriwardane CTH 311 Phone 257-4941

1. Chemistry 100(02) Fall 2009 Webpage: http://blackboard.latech.edu/ Dr. UpaliSiriwardane CTH 311 Phone 257-4941 Upali@chem.latech.edu Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F10:00 - 12:00. Exams: October 5, 2009 (Test 1): Chapter 1 & 2 October 28, 2009 (Test 3): Chapter 3 & 4 November 18, 2009 (Chapter 5 & 6) November 19, 2009 (Make-up test) comprehensive: Chapters 1 through 6 ) 8:00-9:15 AM, CTH 328

2. Chapter 6. Energy and Chemical Reactions • 6.1 The Nature of Energy Page 214 • 6.2 Conservation of Energy Page 217 • 6.3 Heat Capacity Page 222 • 6.4 Energy and Enthalpy Page 227 • 6.5 Thermochemical Expressions Page 233 • 6.6 Enthalpy Changes for Chemical Reactions Page 235 • 6.7 Where Does the Energy Come From? Page 240 • 6.8 Measuring Enthalpy Changes: Calorimetry Page 242 • 6.9 Hess's Law Page 246 • 6.10 Standard Molar Enthalpies of Formation Page 248 • 6.11 Chemical Fuels for Home and Industry Page 253 • 6.12 Foods: Fuels for Our Bodies Page 256 • PORTRAIT OF A SCIENTIST: James P. Joule Page 215 • ESTIMATION: Earth's Kinetic Energy Page 217

3. Chapter 6. KEY CONCEPTSEnergy and Chemical Reactions

4. Kinetic energy and potential energy External or Macroscopic Energy • Potential Energy: Energy of an object as a result of its position • Kinetic Energy: Energy of an object as a result of its motion. Internalor submicroscopic (nano-scale) Energy) • Potential Energy: Energy of an atoms or molecules as a result of its position at nano-scale • Kinetic Energy: Energy of an object as a result of motion of its atoms and molecules at nano-scale. Temperature is directly proportional to kinetic energy (thermal energy) of atoms and molecules • Total Energy = Kinetic + Potential

5. Energy Transformations

6. Thermal Energy

7. Forms of Energy Energy - the ability to do work. Work - when a force is applied to an object. There are several types of energy: Thermal - heat Electrical Radiant - including light Chemical Mechanical - like sound Nuclear

8. Energy units • Kinetic energy was defined as: • kinetic energy = mv2 • m = mass and v = velocity. • Joule (J)- the energy required to move a 2 kg mass at a speed of 1 m/s. It is a derived SI unit. • J = kinetic energy = (2 kg) (1 m/s)2 • = 1 kg m2s-2 • Volume expansion work; P DV • 24.5 L atm x 101. 3 J = 2482 J • 1 L atm 1 2 1 2

9. Energy units and unit conversion 1J = 1 kg m2/sec2 1 cal = 4.184 J 1kcal = 1 Cal thus 1 Cal = 1 kcal = 1000 cal = 4.184 kJ = 4184 J

10. Law of Conservation of Energy “Energy cannot be created or destroyed in a chemical reaction.” During a reaction, energy can change from one form to another. Example. Combustion of natural gas. Chemical bonds can be viewed as potential energy. So during the reaction: 2CH4 (g) + 3O2 (g) 2CO2(g) + 2H2O (l) + thermal energy + light some potential energy is converted to thermal energy and light.

11. First Law of Thermodynamics • the total amount of energy in the universe is a constant • the amount of heat transferred into a system plus the amount of work done on the system must result in a corresponding increase of internal energy in the system

12. Measuring Temperature

13. Hot and Cold Iron

14. Heat Transfer Heat is always transferred from the hotter to the cooler sample

15. State functions • Depend only on the initial and final states of a system. They are independent of how the system gets from one state to another. • State functions include: • Energy • Pressure • Volume • Temperature • Enthalpy (DH)

16. Path Independent Energy Changes

17. Enthalpy Diagram

18. Thermochemistry Terminology Chemical energy – energy associated with a chemical reaction Thermochemistry– the quantitative study of the heat changes accompanying chemical reactions Thermodynamics– the study of energy and its transformations

19. Thermochemistry Terminology State functions  properties which depend only on the initial and final states properties which are path independent Non-state properties  properties which are path dependent state properties  E non-state properties  q & w

20. Stepwise Energy Changes in Reactions

21. Bond Energy and DH • Energy associated with holding 2 atoms together • For example, if the reaction HF(g) -> H(g) + F(g) ;   DH = +565 kJ then we can say that the HF bond energy is 565 kJ/mol.

22. Why is it necessary to divide Universe into System and Surrounding Universe = System + Surrounding System that part of the universe under investigation Surroundings the rest of the Universe

23. Internal Energy • The sum of the individual energies of all nano-scale particles (atoms, ions, or molecules) in that sample • Chemical Energy: Potential energy as stored in bonds • Nuclear energy: E = 1/2mc2 • Thermal Energy: Depends on the temperature • Total Internal Energy: Depends on the type of particles, and how many of them there are in the sample

24. Internal Energy, Heat, and Work Sign Conventions: gain (+). Loss(-)

25. Internal Energy

26. What is the internal energy change (DE) of a system? DE is associated with changes in atoms, molecules and subatomic particles Total Energy= Eexternal + DEinternal DE = work (w) or volume expansion + heat (kinetic energy) + nuclear energy + chemical energy DE = heat (q) + w (work) DE = q + w DE = q -P DV; w =- P DV

27. Volume Expansion Workw = -p DV, Why is negative sign? • Work has a sign: performed (- , loss) or done on the system (+, gain) • volume expansion work: compression and expansion • compression DV = Vf -Vi ;DV is negative expansion DV = Vf -Vi ;DV is positive • compression: w = -p DV; DV = -, w is + expansion: w = -pDV; DV = +, w is -

28. Volume Expansion Type Work w = PDV Expansion w is + compression w is - DV = Vinitial + Vincrease P Vincrease P Vinitial qp = +2kJ

29. Calculations using DE = q + w and DE = q - PDVa) In a process in which 89 J of work is done on a system, 567 J of heat is given off. What is the DE of the system?b) In a process in which a gas expands from 25 L to 50 L against a constant pressure of 0.980 atm and 650 J of heat is absorbed. What is the DE of the system?

30. a) In a process in which 89 J of work is done on a system, 567 J of heat is given off. What is the DE of the system? DE = q + w DE = -567 J +89 J DE = -478 J • Internal energy (DE) of a system is decreased or loss by an amount of 478 J.

31. b) In a process in which a gas expands from 25 L to 50 L against a constant pressure of 0.980 atm and 650 J of heat is absorbed. What is the DE of the system? DV = 50 L- 25 L= 25 L w = -pDV = -0.980 atm x 25 L =- 24.5 L atm to convert L atom to J use conversion factor 1 L atm = 101. 3 J -24.5 L atm x 101. 3 J = - 2482 J 1 L atm -pDV = -2482 J q = 650 DE = q -pDV DE = 650 -2482 J = -1832 J Internal energy(DE) of a system is decreased or loss by an amount of 1832 J

32. Measuring thermal energy changes Thermal energy cannot be directly measured. We can only measure differences in energy. To be able to observe energy changes, we must be able to isolate our system from the rest of the universe. Calorimeter - a device that is used to measure thermal energy changes and provide isolation of our system.

33. Heat capacity vs Specific heat Every material will contain thermal energy. Identical masses of substances may contain different amounts of thermal energy even if at the same temperature. Heat capacity. The quantity of thermal energy required to raise the temperature of an object by one degree. Specific heat.The amount of thermal energy required to raise the temperature of one gram of a substance by one degree.

34. Specific Heats at 25oC, 1 atm Substance DH Al(s) 0.90 Br2 (l) 0.47 C (diamond) 0.51 C (graphite) 0.71 CH3CH2OH (l) 2.42 CH3(CH2)6CH3 (l) 2.23 Substance DH Fe (s) 0.45 H2O (s) 2.09 H2O (l) 4.18 H2O (g) 1.86 N2 (g) 1.04 O2 (g) 0.92 DH = specific heat, J g-1oC-1

35. What basic ideas are used incalorimetric calculations? Heat gain = - heat loss (qlost = - qgained) 1st law of thermodynamics Heat gain/loss = Specific heat x mass x Dt Heat gain/loss = Heat capacity x Dt Dt = final temperature – initial temperature Dt = tf - ti What unknown? Spec. heat, heat cap, tf, mass, DH(coffee cup), DE (bomb calorimeter)

36. EXAMPLEIf 100. g of iron at 100.0oC is placed in 200. g of water at 20.0oC in an insulated container, what will the temperature, oC, of the iron and water when both are at the same temperature? The specific heat of iron is 0.106 cal/goC. (100.g 0.106cal/goC (Tf - 100.)oC) = qlost - qgained = - (200.g  1.00cal/goC (Tf - 20.0)oC) 10.6(Tf - 100.oC) = - 200.(Tf - 20.0oC) 10.6Tf - 1060oC = - 200.Tf + 4000oC (10.6 + 200.)Tf = (1060 + 4000)oC Tf = (5060/211.)oC = 24.0oC

37. What is it? • heat gain/loss? • Heat gain= [Specific heat x mass x Dt] • Heat loss = - [Specific heat x mass x Dt] • specific heat? • amount of heat needed to increase temperature of 1g of a material by 1oC • heat capacity? • heat capacity = [Specific heat x mass] • temperature drop/gain (Dt)? • Dt = tfinal - tinitial

38. Two ways of measuring heat changes • Constant pressure? • qp(open to atmosphere-coffee cup calorimeter) • Constant volume? • qv(bomb calorimeter).

39. Bomb Calorimeter

40. Conduct a reaction and look at the temperature change. Coffee cup calorimeter

41. What exactly is DH? • Heat measured at constant pressure qp • Chemical reactions are exposed to atmosphere and are held at a constant pressure. • Volume of materials or gases produced can change. ie: work = -PDV • DE = qp + w • qp = DE + PDV; w = -PDV • DH = DE + PDV; qp = DH(enthalpy )

42. When 1.00 mole of HCl is reacted with 1.00 mole of NaOH in 1.0 L of water, the temperature of water increases by 13.7oC. Calculate the heat of the reaction for the following thermochemical equation. HCl(aq) + NaOH (aq) ---> NaCl (aq) + H2O(l); DH= ?

43. How do you measure DE? Heat measured at constant volume qv Chemical reactions takes place inside a bomb. Volume of materials or gases produced can not change. ie: work = -PDV= 0 DE = qv + w qv = DE + 0; w = 0 DE = qv = DE(internal energy )

44. Bomb calorimeter(constant volume)

45. DE and DH from First Law of Thermodynamics DE = q + w at constant V, wexpansion = 0 DE = qv + 0 at constant P, wexpansion = PDV DE = qp + PDV qp = DH = DE - PDV

46. EXAMPLE A 1.000g sample of a compound was burned in an oxygen bomb calorimeter. It produced 12.0 kJ of heat. The temperature of the calorimeter and 2000 g of water was raised 4.645oC. How much heat is gained by the calorimeter? heat gained = – heat lost heatcalorimeter + heatwater = heatreaction heatcalorimeter = heatreaction - heatwater

47. EXAMPLEA 1.000g sample of a compound was burned in an oxygen bomb calorimeter. It produced 42.0 kJ of heat. The temperature of the calorimeter and 2000 g of water was raised 4.645oC. How much heat is gained by the calorimeter? heatcalorimeter = heatreaction - heatwater heat = 42.0 kJ - ((2.000kg)(4.645oC)(4.184kJ/kgoC)) = 3.13 kJ

48. Example What is the mass of water equivalent of the heat absorbed by the calorimeter? #g = (3.13 kJ/4.645oC)(1.00kg C/4.184kJ) = 1.61 x 102 g

49. What is it? • heat of fusion • Heat absorbed in this process of breaking intermolecular forces in a solid to become a liquid • heat of evaporation • Heat absorbed in this process of breaking intermolecular forces in a liquid to become a gas

50. ExampleA 1.000 g sample of ethanol (MM 46.07) was burned in the sealed bomb calorimeter described above. The temperature of the water rose from 24.284oC to 27.559oC. Determine the heat for the reaction. m = (2000 + ”161")g H2O Dt = 27.559oC - 24.284oC = 3.275oC q = m  s.h. Dt q= (2161g)(4.184J/goC)(3.275oC) = 29.61 kJ q = (29.61 kJ/1.000g)(46.07 g/mol) = 1364 kJ/mol