1. UNIT IX SOLIDS, LIQUIDS HEAT PROBLEMS CHAPTER 16 PART1 AND CHAPTER 14

2. INTERMOLECULAR FORCES  Forces between molecules  Not as strong as within molecules (covalent and ionic)

3. van der Waals Forces (Intramolecular Force)  Dispersion Forces (London Forces)  Exists between non-polar molecules  weakest I.M.F.  Due to temporary shifts in electron cloud density  Examples CH 4 O 2

4.  Dipole-Dipole Forces slightly polar Example: CHCl 3

5.  HYDROGEN BONDING VERY polar Strongest Examples NH 3 (N -- H) H 2 O (O -- H) HF (F-- H) HCl (Cl -- H)

6. SOLIDS AND LIQUIDS

7. SOLIDS  Orderly rigid and cohesive  Particles that vibrate around fixed points

8. SOLIDS  CRYSTAL true solids particles are arranged in an orderly repeating 3-D pattern

9. SOLIDS CRYSTALS (cont) – consists of a MEMBER o one particle (ion, atom, molecule

10. SOLIDS  several members together make up UNIT CELL simplest repeating unit retains its shape

11. SOLIDS several unit cells together make up a CRYSTAL LATTICE  3-D arrangement of unit cells repeated over and over

12. SOLIDS Vocab- ANHYDROUS (without water) - compound containing no water of hydration HYDRATE-compound with water molecules attached (CuSO4 * 6H2O)

13. SOLIDS AMORPHOUS – solid – no definite repeating pattern – no true melting point – no plateau EXAMPLES: glass, butter, tar, plastic

14. LIQUIDS DEFINITION – particles vibrate around a moving point – non-orderly, non-rigid, cohesive – more space between particles than a solid – exert a vapor pressure – Fluid – ability to flow

15. LIQUIDS UNITS Temperature  average kinetic energy (KE) °C °F K (Kelvin)

16. LIQUIDS VAPOR PRESSURE Definition  pressure exerted by vapor molecules above a liquid when dynamic equilibrium is reached

17. LIQUIDS Pressure  measure of force with which gas molecules hit the side of container  normal atmospheric pressure at sea level  Standard Pressure Units =760 torrs = 760 mmHg = 101.3 kilopascals (kPa)

18. LIQUIDS VAPOR PRESSURE Dynamic equilibrium - 2 opposite processes occurring at same time and same rate VAPOR LIQUID

19. LIQUIDS VAPOR PRESSURE Dynamic Equilibrium depends upon:  Temperature - increase temperature, increase vapor pressure  T  VP VAPOR LIQUID

20. LIQUIDS Strength of inter-molecular forces; hydrogen bonding(such as water) is strongest. increase forces; decrease vapor pressure  IMF  VP VAPOR LIQUID VAPOR LIQUID WATER ALCOHOL

21. LIQUIDS  Viscosity - measure of resistance to flow (how thick) Example – Molasses (syrup) has a high viscosity  Volatility - how easily a liquid evaporates

22. LIQUIDS  Very volatile: high vapor pressure low IMF low boiling point EXAMPLES: alcohol, perfume VAPOR LIQUID ALCOHOL

23. LIQUIDS  Not volatile low vapor pressure high IMF high boiling point Examples: molasses, water VAPOR LIQUID WATER

24. CHANGES IN STATE OR PHASES  Sublimation- – solid changes directly into gas without going through the liquid state Examples: solid iodine, solid air fresheners, "dry" ice

25. CHANGES IN STATE OR PHASES Melting / Freezing – goes from solid to liquid or liquid to solid

26.  Vaporization - evaporation occurs only on the surface at room temperature cooling process Sweat boiling occurs throughout the liquid requires energy CHANGES IN STATE OR PHASES

27.  Boiling Point: vapor pressure = atmospheric (outside) pressure (for any boiling point) normal boiling point vapor pressure = standard pressure standard pressure = 1 atm, 760 torrs, 760 mm Hg,101.3 kPa CHANGES IN STATE OR PHASES

28.  Boiling Point: different altitudes higher altitudes have lower air pressures Denver has a lower boiling point 95 °C than Houston has (100 °C) Foods take longer to cook in Denver than Houston.

29. VAPOR PRESSURE DIAGRAMS 1000 900 800 700 600 500 400 300 200 100 760 20 4060 80 100 CHLOROFORM ETHYL ALCOHOL WATER Temperature ( °C) Vapor pressure (mm Hg)

30. PHASE DIAGRAMS Graphs that show conditions (temperature and pressure) under which a substance will exist as a solid, liquid, or gas.

31. PHASE DIAGRAMS 700 600 500 400 300 200 100 760 80 120160 Temperature ( ° C) Pressure (mm Hg) 800 40 60 100140180 X Z X - Triple point All three states are in equilibrium at this temperature and pressure. X-Y line - These are sublimation points. Z - Critical temp. and pressure. A gas can't be liquified above this point.

32. PHASE DIAGRAMS 700 600 500 400 300 200 100 760 80 120160 Temperature ( ° C) Pressure (mm Hg) 800 40 60 100140180 X Z SOLID LIQUID GAS Lines represent 2 phases in equilibrium.

33. PHASE DIAGRAMS 700 600 500 400 300 200 100 760 80 120160 Temperature ( ° C) Pressure (mm Hg) 800 40 60 100140180 X Z Normal boiling point (condensation) occurs when standard pressure crosses liquid / gas line Normal boiling point (condensation) occurs here.

34. PHASE DIAGRAMS 700 600 500 400 300 200 100 760 80 120160 Temperature ( ° C) Pressure (mm Hg) 800 40 60 100140180 X Z Normal melting point (freezing) occurs where standard pressure crosses liquid / solid line. Normal melting point (freezing) occurs here

35. PHASE DIAGRAMS 700 600 500 400 300 200 100 760 80 120160 Temperature ( °C) Pressure (mm Hg) 800 40 60 100140180 X Z Freezing or melting point Boiling or condensation point Deposition or sublimation point

36. UNIQUE PROPERTIES OF WATER STRONG HYDROGEN BONDING CAUSES: – high boiling point and melting point – high specific heat capacity – high surface tension needle floats – Water droplets are spherical

37. HEAT VS. TEMPERATURE  Energy transferred from one body to another because of a difference in temperature  Average Kinetic Energy  Written as KE

38. HEAT VS. TEMPERATURE  UNITS – calories (c) – kCal - C (1000 calories) – Joules - J energy for one heartbeat – 1 cal = 4.18 J – 1 kCal = 4180 J  UNITS – °C - celsius – °F -Fahrenheit – K - kelvin (no degree sign!)

39. HEAT VS. TEMPERATURE  Measured by: – indirectly by a calorimeter  Measured by: – thermometer

40. HEAT VS. TEMPERATURE  DEPENDS UPON – mass more mass means more heat – Cp (S) - specific heat type of matter some hold heat better than others –  T - change in temperature  DEPENDS UPON – amount of movement of the particles in the substance

41. HEAT VS. TEMPERATURE  FORMULA q=energy (J) m=mass (g)  q = (m) (  T) (Cp)  q = (m) (T 2 -T 1 ) (Cp)

42. Specific Heat or Heat Capacity  Amount of heat needed to raise 1 gram of a substance 1 degree Celsius  Units – (J/g o C) – (cal/g o C)  Examples – water --- 4.18 J/g o C or 1 cal/g o C – Au --- 0.129 cal/g o C – alcohol --- 2.45 J/g o C

43. Calorie  Amount of heat needed to raise one gram of water one degree of celsius  It takes one calorie to raise one gram of water one degree of Celsius

44. Heat of Fusion - Hf  Amount of heat needed to melt one gram of a substance at its melting point  Units (cal/g)  Examples – water (Hf) = 334 J/g or 76.4 cal/g – Ag = 88 J/g

45. HEAT OF VAPORIZATION - Hv  Amount of heat needed to vaporize one gram of a substance at its boiling point  Examples – water (Hv) = 2260 J/g or 539 cal/g – Pb = 858 J/g

46. PHASE CHANGE DIAGRAMS SOLID LIQUID GAS TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 Heat of fusion - Melting point - substance is becoming a liquid Heat of vaporization Boiling point- substance is becoming a liquid WATER

47. Heat Calculations - Formulas The state remains the same and there is no change in temperature. q= joules m=grams Cp=J/g or J/c q= (m) (Cp) q = Heat

48. Example of Non-Changing State  Melting/freezing at melting point  Vaporizing/condensing at boiling point How much energy does it take to melt 55g of gold at its melting point? Cp = 64.5 J/g q= (m) (Cp) = (55g)(64.5 J/g) = 3547.5 J

49. HEAT EQUATION  One substance with a temperature change q=joules (J) m= mass (g) Cp = specific heat capacity (J/g °C) (J/c °C) T 2 = final temperature T 1 = initial temperature q = (m) (Cp) (T 2 -T 1 )

50. HEAT EQUATION EXAMPLE ***Heating or cooling with no change in state*** How much energy is released as 33 g of solid silver cools from 95 °C to 60°C? Cp of silver = 0.236 J/g °C

51. HEAT TRANSFER EQUATION How a substance changes the temperature of another substance used in calorimeter calculations (m 1 ) (Cp 1 ) (T 2 -T 1 ) = (m 2 ) (Cp 2 ) (T 2 -T 1 ) Warm substance losing energy Cool substance gaining energy Energy LOST = Energy GAINED

52. HEAT TRANSFER EQUATION EXAMPLE A piece of metal is dropped into a beaker of boiling water whose temperature is 95 C. The 5g piece of metal is put into 100g of cold water at 20 C. The temperature of the water rises to 30 C. What is the specific heat of the metal? Cp (water) = 4.18 J/g C o o o o

53. EQUATION FOR CHANGING TEMPERATURE AND STATES Draw the phase change diagram

54. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 Use the following equations: q = (m) (Cp) q = (m) (Cp) (T 2 -T 1 )

55. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 1. Heat solid to melting point q = (m) (Cp) (T 2 -T 1 )

56. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 2. Melting solid to liquid q = (m) (Cp)

57. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 3. Heat liquid to boiling point q = (m) (Cp) (T 2 -T 1 )

58. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 4. Change liquid to gas q = (m) (Cp)

59. CHANGING STATES AND TEMPERATURE TEMPERATURE ( C) o HEAT (cal/g) OR TIME 100 0 5. Heating gas q = (m) (Cp) (T 2 -T 1 )

60. CHANGING STATES AND TEMPERATURES 1. Heat solid to melting point : KE 2. Melt solid to liquid: PE 3. Heat liquid to boiling point: KE 4. Change liquid to gas: PE 5. Heat gas: KE q = (m) (Cp) (T 2 -T 1 ) q = (m) (Cp) q = (m) (Cp) (T 2 -T 1 ) q = (m) (Cp) q = (m) (Cp) (T 2 -T 1 ) When to use which equations:

61. CHANGING TEMPERATURE AND CHANGING STATES EXAMPLE How much energy is needed to change 30g of ice at -5 °C to steam at 120 °C?