1. Intermolecular Forces (IMFs) and States of Matter GPS 18

2. Phases of Matter SOLID Definite Shape Definite Volume LIQUID Shape varies depending on container Definite Volume GAS Takes on the shape and volume of the container ENERGY INCREASES Simulation PLASMA A gas heated to extremely high temp. Atoms ionize to result in electrons and Ions

3. Phase Changes Phase ChangePhasesEnergy absorbed or released for phase change to occur? MeltingSolid to liquid FreezingLiquid to solid EvaporatingLiquid to gas CondensingGas to liquid SublimatingSolid to gas DepositingGas to solid

4. Phase Changes of Water (H 2 O)

5. Heating/Cooling Curve: Particle-View

7. Heating Curve for Water “Horizontals”: potential energy is increasing (particles are spread farther apart to change phases) so there is no increase in temperature Simulation

8. What phase or phases are present at D? ____________________ What letter corresponds to only gas being present? ___________ What is happening at B? __________________________________

9. What is the melting point of this substance? ______________ About how many minutes did it take for this substance to begin to boil? _______________

10. What is the boiling point of this substance?

11. Cooling Curve What phase(s) is/are present from A to B? _______________ What phase(s) is/are present from B to C? _______________ What phase(s) is/are present from C to D? _______________ What phase(s) is/are present from D to E? _______________

12. Cooling Curve What is the melting point for this substance? _____________ What is the boiling point for this substance? _____________

13. Phase Diagrams The state of a substance not only depends on temperature, but also depends on pressure.

14. General Phase Diagram

15. Phase Diagrams: Points of Interest At all points along each of the lines dividing the phases, those two phases exist in equilibrium with each other Example: CO 2 (l) ⇄ CO 2 (g) At the triple point, all three phases are present at equilibrium with each other Example: CO 2 (s) ⇄ CO 2 (l) ⇄ CO 2 (g) Beyond the temperature and pressure of the critical point, liquid and gas phase are no longer distinguishable (supercritical fluid formed)

16. S ⇄ L L ⇄ G S ⇄ G S⇄L⇄GS⇄L⇄G

17. A substance at its Triple Point The picture (and video) show a substance at its triple point, which is the specific temperature and pressure at which solid, liquid, and gas phases are in equilibrium solid ⇄ liquid ⇄ gas http://www.techeblog.com/index.php/te ch-gadget/mind-bending-video-shows- liquid-that-boils-and-freezes-at-the- same-time

18. Boiling Water at Room Temperature (about 20˚C) You can boil water at room temperature (about 20˚C) if you decrease the pressure of the air surrounding the water (normal air pressure is around 1 atm)

19. Boiling Water at Room Temperature (about 20˚C) https://www.youtube.com/watch?v=glLPMXq6yc0 A vacuum pump may be used to remove air from the bell jar, decreasing the air pressure inside the bell jar. Air pumped OUT Lower Pressure

20. Melting Ice by Increasing Pressure You can melt ice by applying pressure to it (without changing the temperature)

21. Melting Ice by Increasing Pressure The pressure applied by the wire melts the ice and the ice then refreezes around the wire, stringing the ice on the wire.

22. Phase Diagrams Practice What represents the triple point? What phase is present at 2? What phase is present at 3? What phase change occurs from 1 to 2? ________________ What phase change occurs from 1 to 3? ________________

23. Intermolecular Forces (IMFs) Intermolecular forces (IMFs) forces of attraction between particles Intramolecular forces forces of attraction within molecules

24. Intermolecular Forces (IMFs)

25. Relative Strength of Intramolecular forces

26. Relative Strength of Intermolecular forces

27. Relative Bond Strength Intramolecular forces: Ionic bonds Covalent bonds Metallic bonds Intermolecular forces (van der Waal’s forces): Hydrogen bonds Dipole-dipole forces London dispersion forces Increasing Bond Strength Order of strength may vary depending on the exact compound Order of strength when comparing molecules of similar size (by mass)

28. Dipole-Dipole Forces Forces of attraction between polar molecules (dipoles) Dipole-dipole forces between HCl molecules δ+δ+ δ+δ+ δ-δ- δ-δ-

29. Dipole-Dipole Forces: Polar molecules Recall that trigonal pyramidal and bent molecules are always polar For all other shapes see the examples below: All fluorines have same electronegativity, individual bond polarities cancel out, electrons are shared equally, nonpolar molecule. Cl is less electronegative than F, individual bond polarities do NOT cancel out, electrons are NOT shared equally, polar molecule.

30. Hydrogen Bonding A special case of dipole-dipole forces Strongest type of intermolecular force, if comparing molecules of similar size. Occurs between molecules of compounds containing H-N, H-F, or H-O bonds. Also occurs between hydrocarbons containing –OH groups Example: occurs between molecules of H 2 O, NH 3, HF, C 2 H 5 OH

31. Hydrogen Bonding between Water Molecules Hydrogen bonding is the attraction between the hydrogen in one molecule and a highly electronegative atom (oxygen, nitrogen, or fluorine) from another molecule

32. Hydrogen Bonding between Water Molecules

33. Hydrogen bonding between HF molecules Hydrogen bond

34. Hydrogen bonding between Organic Molecules Organic compounds, such as alcohols and carboxylic acids, that contain –OH are attracted to each other through hydrogen bonds Ethanol, C 2 H 5 OH Acetic acid, CH 3 COOH

35. Hydrogen bonding: in DNA

37. London Dispersion Forces (LDFs) London dispersion forces (LDFs) The weakest type of intermolecular force (when comparing molecules of similar size) Occurs between nonpolar atoms or molecules like Ne or H 2 Animated example of LDFs between two nonpolar atoms

38. London Dispersion Forces

39. London Dispersion Forces: Geckos Geckos are able to climb almost any surface at any angle because of the London Dispersion forces between its feet and the surface. A gecko can hang on a glass surface by just one toe. Stanford University researchers recently developed a gecko-like robot which uses synthetic setae to climb walls!

41. Intermolecular Forces: Practice For the following substances, identify the type of intermolecular forces between molecules in a sample of each, then rate the relative strength of the IMFs in each sample: NH 3 F 2 PF 3

42. Breaking Bonds: Intramolecular vs Intermolecular Process 1: CO 2 (g) → C(s) + O 2 (g) Process 2: CO 2 (s) → CO 2 (g) Identify whether intermolecular or intramolecular bonds are broken in each process above. Process 1: _________________________________ Process 2:__________________________________ Specify which type of bond is broken in each process above. Process 1: _________________________________ Process 2:__________________________________ Which process would require more energy to complete? ______

43. IMFs: Effect on Physical Properties The stronger the intermolecular forces (↑ IMFs): More likely a solid or liquid at room temperature Higher surface tension (↑ S.T.) Lower vapor pressure (↓ VP) Higher boiling point (↑ BP)

44. Intermolecular Forces: State of Matter The state of a substance at room temperature depends on the strength of its IMFs. Particles highly attracted to each other. Stronger IMFs Particles weakly attracted to each other. Weaker IMFs Solid Substance X Liquid Substance YGaseous Substance Z

45. Intermolecular Forces: State of Matter Which do you think has the weakest intermolecular forces, when compared at the same temperature: A(g), B(l) or C(s)? Justify your answer.

46. Stronger IMFs, Greater Surface Tension The stronger the intermolecular forces (↑ IMFs), the higher the surface tension (↑ S.T.)

47. Stronger IMFs, Greater Surface Tension Water has high surface tension due to relatively strong IMFs (hydrogen bonding) between molecules. This allows water bugs to “walk” on water.

48. Vapor Pressure: Liquid-Gas Equilibrium In a closed container, a liquid evaporates until a high enough pressure is established, then some of the gas condenses back into a liquid. At this point, an equilibrium is reached between the liquid phase and gas phase InitiallyEquilibrium

49. Vapor Pressure: Liquid-Gas Equilibrium The equilibrium is dynamic, meaning the number of molecules entering and leaving the liquid and gas phases is the same

50. Vapor Pressure The pressure of the vapor (gas) found directly above a liquid in a closed container, once equilibrium is established.

51. Stronger IMFs, Lower Vapor Pressure The particles in A are more strongly attracted to each other (stronger IMFs), therefore less of the particles have vaporized (lower vapor pressure). The particles in B are not as strongly attracted to each other (weaker IMFs), therefore more of the particles have vaporized (higher vapor pressure). A and B are two liquids in closed containers at room temperature:

52. Vapor Pressure: Volatile Liquids Volatile liquids Readily evaporates at room temperature Nonvolatile liquids Does not readily evaporate at room temperature

53. Vapor Pressure: Interpreting Graphs Which substance has the greatest vapor pressure at 20 o C? Which substance will evaporate the fastest at room temperature? Which substance has the strongest IMFs?

54. Vapor Pressure: Practice Predict whether F 2 or HBr has the higher vapor pressure at room temperature. Justify your answer.

55. Vapor Pressure: Practice Predict whether acetone, a volatile liquid, or glycerol, a nonvolatile liquid, has the higher vapor pressure at room temperature. Justify your answer.

56. Vapor Pressure: Boiling Point As temperature increases, vapor pressure increases. Normal boiling point is the temperature at which the liquid boils at 1 atm (or 760 mmHg, or 760 torr)

57. Boiling Point: Interpreting Graphs Using the vapor pressure diagram above: What is the boiling point of ethanol?

58. Boiling Point When the vapor pressure equals the atmospheric pressure, the gas escapes the liquid. This temperature is called the boiling point.

59. Boiling Point Physical changes, like boiling, only break intermolecular forces (NOT intramolecular forces). For example, when water boils hydrogen bonds between molecules are broken, NOT covalent bonds inside of each water molecule.

60. Stronger IMFs, Higher Boiling Point Hydrogen bonding requires more energy to separate particles, so higher boiling point temperature

61. Boiling Point: Practice Explain the following in terms of intermolecular forces: O 2 boils at -183˚C, HCl boils at -85˚C, and H 2 O boils at 100˚C.

62. Energy Required to Change Phases: Heat of Fusion (∆H fus ) and Heat of Vaporization (∆H vap )

63. Heat of Vaporization and Heat of Fusion Heat of Vaporization (∆H vap ) The amount of energy required to vaporize a liquid to the gas phase. Heat of Fusion (∆H fus ) The amount of energy required to melt a solid to the liquid phase. ∆H vap and ∆H fus are different for each substance because of varying strength of IMFs SubstanceHeat of Vaporization (kJ/mol) Heat of Fusion (kJ/mol) Benzene (C 6 H 6 ) 44.39.84 Bromine (Br 2 ) 29.510.8 Mercury (Hg) 59.42.33 Methane (CH 4 ) 9.20.84 Water (H 2 O)40.76.00

64. Heat of Vaporization and Heat of Fusion What amount of energy is required to change 32.5 g of liquid Br 2 to gaseous Br 2 ? 6.00 kJ SubstanceHeat of Vaporization (kJ/mol) Heat of Fusion (kJ/mol) Benzene (C 6 H 6 )44.39.84

65. Heat of Vaporization and Heat of Fusion What amount of energy is required to change 12.9 g of solid Benzene (C 6 H 6 ) to liquid Benzene? 1.63 kJ SubstanceHeat of Vaporization (kJ/mol) Heat of Fusion (kJ/mol) Benzene (C 6 H 6 )44.39.84

66. Energy (q) required to heat a single phase q =

67. Energy (q) released/absorbed q surroundings = m C ∆T q sur = the measured energy lost or gained during the reaction (energy of the surroundings) as heat m = mass, in grams (if solutions mixed together, add for total mass) ∆T is change in temperature (T final – T initial ) in ˚C C = specific heat capacity (J/g˚C)

68. Specific Heat Capacity (C) The amount of energy (or heat) needed to raise the temperature of 1 gram of a substance by 1 Celsius (ºC) degree Describes a substance’s capacity to store energy Water has a relatively high heat capacity compared to other materials Aluminum foil C = 0.897 J/g˚C Copper penny C = 0.385 J/g˚C Glass marble C = 0.840 J/g˚C Water C = 4.184 J/g˚C

69. Energy (q) of a Reaction Calorimeter Device that measures the heat (q) lost or gained during a reaction or any process that involves heat exchange (such as dissolving a salt in water) Heat absorbed by the surroundings was given off by the reaction: q surroundings = - q reaction

70. Energy (q) of a Reaction If 3580 joules of energy are needed to raise the temperature of 3.5 grams of solid copper from 25 ⁰ C to 125 ⁰ C, what is the specific heat capacity of copper?

71. Energy (q) of a Reaction In a coffee cup calorimeter, 100.0 mL of NaOH solution and 100.0 mL of HCl solution are mixed. Both solutions were originally at 24.6˚C. After the reaction, the final temperature is 31.3˚C. Assuming that all solutions have a density of 1.0 g/mL and a specific heat capacity of 4.184 J/g˚C, calculate the amount of heat that was emitted during the neutralization of HCl by NaOH. Assume that no heat is lost to the surroundings or the calorimeter apparatus.

72. References http://images.google.com

73. SOLs CH 4 f CH 5 a, b, c, d, e