1. Intermolecular Forces and Liquids

2. Kinetic Molecular Theory According to the Kinetic Molecular Theory, ALL particles of matter are in constant motion. This theory helps explain the behavior of solids, liquids, and gases.

3. Behavior of gases Particles in a gas are never at rest. Gaseous atoms travel in a straight line until it collides with either another atom or the wall of the container. The constant motion of gas particles allow it to fill a container of any shape or size.

4. Behavior of liquids Particles in a liquid are more closely packed than the particles in a gas. Therefore, attractions between the particles in a liquid do affect the movement of the particles. – It slows them down – less kinetic energy

5. A liquid takes the shape of its container because particles can flow to new locations. The volume is constant because forces of attraction keep the particles close together.

6. Behavior of solids Solids have a definite volume and shape because their particles vibrate around fixed locations. Strong attractions restrict motion and keep each atom in a fixed location relative to its neighbors. Atoms vibrate around its location but it does not exchange places with neighboring atoms.

7. Phase changes The reversible physical change that occurs when a substance changes from one state of matter to another.

8. Forces that hold atoms and molecules together Ionic Bond: due to electrostatic attraction between opposite charges (ionic compounds). Covalent Bond: due to combining of atomic orbitals when electrons are shared. Intermolecular Forces: due to electrostatic attraction between opposite charges (ionic and covalent compounds).

9. Summary of Intermolecular Forces Van der Waals intermolecular forces: Ion-dipole forces Ion-dipole forces Dipole-dipole forces Dipole-dipole forces – Special dipole-dipole force: hydrogen bonds Forces involving nonpolar molecules: induced forces Forces involving nonpolar molecules: induced forces Dispersion or London forces. Dispersion or London forces.

10. Two factors affect the strength of ion or dipole force (Coulomb’s law) Magnitude of charge Distance

11. Generally – order of strength Strongest: Ionic Bonds Ion-dipole bonds Hydrogen Bonding Dipole forces Induced dipole Weakest: Dispersion forces Reflects in properties

12. Comparing Properties Melting Point, Boiling Point, Heat of Fusion, Heat of Vaporization, Surface tension, viscosity - - - all go up as strength of IMF increases Vapor pressure – goes down as strength of IMF increases (higher IMF makes evaporation less likely and less gas molecules mean lower vapor pressure)

13. Ion-Ion Forces for comparison of magnitude Na + —Cl - in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 o C MgO, mp = 2800 o C

14. Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance.

15. Dipole-Dipole Forces Dipole-dipole forces bind molecules having permanent dipoles to one another.

16. Dipole moment Polar molecules or dipoles have positive and negative side. A measure for the polarity is the dipole moment (p. 376). Dipole moment is given in Debye (D).

17. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole- dipole attractions. H-bonding occurs when X and Y are N, O, or F

18. Watch out! Hydrogen must be connected to F, O, or N. Not all hydrocarbons with these atoms present contain hydrogen bonding. Example: Dimethyl ether: CH 3 – O – CH 3 does NOT have hydrogen bonding. Methanol: CH 3 OH does have hydrogen bonding.

19. H-Bonding Between Methanol and Water H-bondH-bondH-bondH-bond

20. Hydrogen Bonding in H 2 O H-bonding is especially strong in water because the O—H bond is very polar the O—H bond is very polar there are 2 lone pairs on the O atom there are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

21. Hydrogen Bonding in H 2 O Ice has open lattice-like structure. Ice density is < liquid. And so solid floats on water.

22. A consequence of hydrogen bonding

23. H-bonding leads to abnormally high specific heat capacity of water (4.184 J/gK) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

24. DNA Base-Pairing through H-Bonds

25. Boiling point of water Water has a very high boiling point compared to other simple hydrogen containing compounds, such as CH 4, H 2 S, and even NH 3 and HF (which exhibit H-bonding as well, but less “extreme” as H 2 O

26. Boiling Points of Simple Hydrogen-Containing Compounds

29. FORCES INVOLVING INDUCED DIPOLES How can non-polar molecules such as O 2 and I 2 dissolve in water? The water dipole INDUCES a dipole in the O 2 electric cloud. Dipole-induced dipole

30. FORCES INVOLVING INDUCED DIPOLES Solubility increases with the mass of the gas Process of inducing a dipole is polarization Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.

31. Dispersion forces Induced dipole – induced dipole Between 2 non-polar molecules. Caused by the movement of the electron cloud. The more electrons (higher molar mass) the stronger the force (higher MP and BP).

32. FORCES INVOLVING INDUCED DIPOLES Formation of a dipole in two nonpolar I 2 molecules. Induced dipole-induced dipole (dispersion) (dispersion)

33. FORCES INVOLVING INDUCED DIPOLES The induced forces between I 2 molecules are very weak, so solid I 2 sublimes (goes from a solid to gaseous molecules).

34. FORCES INVOLVING INDUCED DIPOLES The magnitude of the induced dipole depends on the tendency to be distorted. Higher molar mass = stronger forces MoleculeBoiling Point ( o C) MoleculeBoiling Point ( o C) CH 4 (methane) - 161.5 CH 4 (methane) - 161.5 C 2 H 6 (ethane)- 88.6 C 2 H 6 (ethane)- 88.6 C 3 H 8 (propane) - 42.1 C 3 H 8 (propane) - 42.1 C 4 H 10 (butane) - 0.5 C 4 H 10 (butane) - 0.5

35. Boiling Points of Hydrocarbons CH 4 C2H6C2H6C2H6C2H6 C3H8C3H8C3H8C3H8 C 4 H 10 Note linear relation between bp and molar mass.

36. LiquidsLiquids In a liquid Molecules are in constant motionMolecules are in constant motion There are appreciable intermolecular forcesThere are appreciable intermolecular forces Molecules are close togetherMolecules are close together Liquids are almost incompressibleLiquids are almost incompressible Liquids do not fill the containerLiquids do not fill the container

37. Liquids: Energy and Phase changes

39. Melting and Freezing Melting: add energy to break bonds that keep molecules at fixed position (kinetic energy goes up): endothermic Freezing: energy released (kinetic energy goes down) as particles “settle” in fixed positions: exothermic

40. LIQUID VAPOR Evaporation: Add energy to break IM bonds Condensation: Remove energy to form IM bonds Evaporation and condensation

41. To evaporate, molecules must have sufficient energy to break IM forces. This breaking requires energy, so the process of evaporation is endothermic. Evaporation

42. Condensation When a gas or vapor condensates, the kinetic energy of molecules gets lower, while IM forces get stronger. Energy is released: this process is exothermic. When a gas or vapor condensates, the kinetic energy of molecules gets lower, while IM forces get stronger. Energy is released: this process is exothermic.

43. Distribution of Energy in a Liquid

44. At higher T a much larger number of molecules has high enough energy to break IM forces and move from liquid to vapor state. High E molecules carry away E. You cool down when sweating or after swimming.

45. When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE. EQUILIBRIUM VAPOR PRESSURE is the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation.

46. Liquid in flask evaporates and exerts pressure on manometer. Measuring Equilibrium Vapor Pressure

47. HEAT OF VAPORIZATION HEAT OF VAPORIZATION is the heat required (at constant P) to vaporize the liquid. Compd.∆ vap H (kJ/mol) IM Force H 2 O40.7 (100 o C)H-bonds SO 2 26.8 (-47 o C) dipole Xe12.6 (-107 o C)induced dipole

48. Heat of vaporization How much heat is needed to vaporize 6.51 mL of water? (Δ vap H = 40.7 kJ/mol) Density of water = 1.0 g/mL, so 6.51 mL = 6.51 g 6.51 g/18.02 g/mol = 0.361 mol H 2 O 0.361 x 40.7 = 14.7 kJ

49. Boiling Liquids Liquid boils when its vapor pressure equals atmospheric pressure. Then bubbles of vapor form within the liquid.

50. Boiling point Boiling occurs when a liquid turns to a gas inside the liquid ◦ bubbles are produced Liquid boils when its Vapor Pressure = Atmospheric Pressure ◦ Normal boiling point Larger IMF = lower vapor pressure = high BP Weaker IMF = high vapor pressure = lower BP

51. Molecules at surface behave differently than those in the interior. Molecules at surface experience net INWARD force of attraction. This leads to SURFACE TENSION — the energy required to break the surface. Surface Tension

52. Surface Tension also leads to spherical liquid droplets

53. Surfactants Surface tension can be decreased by adding surfactants (soap, detergents). They interfere with hydrogen bonding.

57. Capillary action IMF also lead to capillary action and to the existence of a meniscus for a water column. This is caused by ADHESIVE FORCES between water and glass

59. Capillary Action Movement of water up a piece of paper depends on H-bonds between H 2 O and the OH groups of the cellulose in the paper.

60. Viscosity Resistance to flow – Goes up as IMF increases Also goes up as length of hydrocarbon chain increases as molecules get tangled up and don’t flow easily

61. Phase diagram We use these diagrams to relate the process that occur when a substance changes from one phase to another. Substances are in the following states when in certain locations on the diagram: – Solid – left side of diagram – Liquid – middle of diagram – Gas/Vapor – right side of diagram

63. When either the temp or pressure is changed, you can identify the process that is taking place and identify the phase change. – Ex (from diagram on last slide) – At 1 atm if you increase the temperature from 90 o C to 200 o C, the process you are undergoing is vaporization or boiling (liquid to gas).

65. Triple point The change of state occurs right on the equilibrium line. Triple point identifies the conditions when you have all 3 states in dynamic equilibrium with one another.

66. Normal MP and BP Tm  normal melting point – The point at 1 atm or 101.3 kPa when solid turns to liquid. Tb  normal boiling point – The point at 1 atm or 101.3 kPa when a liquid turns to a vapor

67. Critical point Critical point – you are no longer able to distinguish between gas and liquid phases past this point.

68. CO 2 Phase diagram