1. Chapter 13 Mrs. Ward AP Chemistry

2. Good things to know: pg. 426 Kinetic Molecular Theory pg. 448 Characteristics of phases

3. Intermolecular vs. Intramolecular Intermolecular Forces (IMF): the forces present between individual particles Intramolecular Forces: the forces within a molecule that hold it together (covalent/ionic bonds) –Intermolecular forces are generally weaker than intramolecular forces

4. Intermolecular Forces A substance's state at a certain temperature depends on the particle’s ability to overcome intermolecular forces. Gases : kinetic energy of the particles is significantly greater than the intermolecular forces of the particles.

5. Intermolecular Forces Liquid: kinetic energy of the particles is greater than their intermolecular forces but not great enough for the particles to completely break free of each other.

6. Intermolecular Forces Solid: kinetic energy of the particles is less than the intermolecular forces of the particles.

7. Intermolecular Forces Intermolecular forces are weaker than ionic/covalent bonds. A. Less energy is needed to boil or melt a substance than to actually break bonds. B. These forces influence boiling points. C. Introducing: van der Waals forces (intermolecular forces). D. At higher pressures and lower temperatures, these forces become more important. (Remember real gases!?)

8. Ion-ion Interactions Ionic bonds= strong= high melting points When molten can conduct electricity As a liquid IMF is weaker than if solid, but still higher than neutral species Ionic substances w/multiply charged ions have higher melting/boiling points than single charged ions (pg. 450)

9. Dispersion Forces (London) Positive nuclei of one atom and electron cloud of another molecule are attracted Induces temporary dipole Occurs in ALL molecules but it is the only one present in nonpolar molecules.

10. As cloud becomes larger, electrons attracted less strongly by own nuclei and are more easily polarized (Therefore, dispersion forces are stronger for larger molecules)

11. Shape Matters Shape matters! More exposed surface = more contact = more induced dipoles!

12. Dipole-dipole Forces 2. Dipole-dipole forces: a. Occurs in polar molecules b. Larger the dipole moment (electronegativity difference), the stronger the force. c. These are stronger than dispersion forces. Short distances necessary because only partial charges in dipole

13. Dipole-dipole Forces Higher temperature= more movement (kinetic energy)= dipole interactions less important

14. Hydrogen Bonding A type of dipole-dipole interaction a. Occurs in polar molecules that have an H-F, H-O or H-N bond (H-FON!) b. Stronger than dipole-dipole and dispersion forces. c. Because of the attraction b/w the H and the lone pair on a small electronegative atom (F, O, N!)

15. Relative strength: A. Hydrogen bonding >>Dipole-dipole forces >> London dispersion forces B. If the type of intermolecular force is the same, molecular size will influence the boiling point. (more polarizable!) C. If type and size is the same, shape will influence the boiling point. (more surface!)

16. Identify the types of intermolecular forces present in condensed phase of each. Then identify major type of force: H 2 O I 2 NO 2 Water: H-bonds & London forces…H- bonds stronger Iodine: Nonpolar so London forces only Nitrogen dioxide: Polar= dipole-dipole interactions & London forces… dipole interactions stronger

17. Liquid State (pg. 455) Viscosity is the resistance to flow High= honeylow=gasoline The stronger IMF existing makes liquid more viscous –H-bonds usually have highest viscosities An increase in size/surface area of molecule means higher viscosity (more dispersion forces present) Decreases as temperature rises (more kinetic energy allows molecules to overcome IMF)

18. Which is more viscous? C 5 H 12 or C 12 H 26 WHY???

19. Surface Tension Measure of inward forces that must be overcome to expand surface area of liquid. Stronger IMF = high surface tension

20. Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion attracted to glass Cohesion attracted to each other

21. Evaporation (Vaporization) Liquid  Gas Molecules must gain enough kinetic energy to overcome IMF (increases as temperature increases) Condensation If gas molecule collides w/liquid, can be “captured” Dynamic equilibrium evaporation/condensation at same rates in closed container –(More evap. = more molecules in air = more likely to collide w/liquid and condense) H 2 O (l) H 2 O (g)

22. Le Chatelier’s Principle A system at equilibrium, or changing towards equilibrium, responds in the way that relieves stress placed on it.

23. Vapor Pressure (vp) Partial pressure of vapor molecules above the surface of liquid at equilibrium at a given temperature (increases as temperature increases since vaporation increases as temperature increases) Easily vaporized liquids= volatile = high vapor pressure (Strong IMF holds molecules in liquid state means not very volatile so larger molecules and h- bonds have low vapor pressure)

24. The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm. If you add heat to a liquid at its boiling point, the temperature remains constant, because the energy is used to overcome the IMF (represented by a plateau on a heating curve).

25. Heat transfer w/liquids Molar heat capacity the amount of heat that must be added to (a stated mass of) liquid to raise its temperature by one degree Celsius J/mol * C If heat added under constant pressure, temp will rise until hits boiling point, then temp remains constant… Molar heat of vaporization (enthalpy) amount of heat to be added to one mole of the liquid at its boiling point to convert it to vapor w/no change in temperature (Can be expressed in kJ/mol or J/g) Why do you think that the units of heat of vap. Does NOT include temp? –At boiling there is no change in temperature.

26. Heats of vaporization increase as boiling points (and IMF) increase and as vapor pressures decrease

27. Liquids can evaporate below their boiling points (perspiration) Perspiration absorbs heat (2.41 kJ per gram of water) from body and is evaporated away, thus we feel cooler. Heat of condensation amount of heat to be removed from vapor to condense it (w/no change in temp)

28. Arrange in order of lowest to highest boiling point. H 2 S, H 2 O, CH 4, H 2, KBr Answer: H 2 CH 4 H 2 S H 2 O KBr KBr is ionic, so highest… Water has strong h-bonds, so it boils at next highest H 2 S is also polar covalent so boils below H 2 O but above other 2 CH 4 and H 2 both nonpolar (weakest IMF), CH 4 is larger (more dispersion forces) so will boil at higher temp than H 2

29. Pg.492 # 21, 25, 27, 28, 29, 33

30. The Solid State Melting point (freezing point) is the temperature at which its solid and liquid phases coexist in equilibrium Normal melting point is melting point at 1 atmosphere of pressure (though changes in pressure have small effects on melting points) Molar heat of fusion is the energy required to melt 1 mole of a solid substance. H 2 O (s) H 2 O (l)

32. Phase Diagrams A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas Only applies to closed systems (so matter doesn’t escape to surroundings) The triple point is where all 3 phases meet.

33. Some important points: 1. Lines (AD, AB, AC) represent points where phase changes or phase equilibria occur.

34. 2. Triple point (A): P and T where all three phases are in equilibrium. 3. Critical point (B): Created by the critical temp. and critical press. Above this point, liquid and vapor are indistinguishable.

35. 4. Melting points: Line AD. 5. Boiling points: Line AB. 6. Sublimation points: Line AC.

36. A note about the melting point line. When the slope of this line is: 1. Positive, the solid phase is more dense than the liquid phase. 2. Negative, the liquid phase is more dense than the solid phase.

37. Steps for Drawing a Phase Diagram 1) Draw axes/label (pressure vs. temp) –Include units!!! (should be atm for pressure) 2) Plot triple point and normal melting point. Connect these…that’s the melting point line. Use the slope & draw the line past this point continuing the trend. 3) Draw lines from the triple point to the normal b.p. for boiling point line and from the triple point down for the sublimation line. 4) Label solid, liquid, and gas phases

38. A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A unit cell is the basic repeating structural unit of a crystalline solid. lattice point Unit Cell Unit cells in 3 dimensions At lattice points: Atoms Molecules Ions

41. Types of Crystals Ionic Crystals – Ion-Ion interactions are the strongest (including the “intermolecular forces” (H bonding, etc.) Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsClZnSCaF 2

42. Types of Crystals Covalent Crystals – Stronger than IM forces but generally weaker than ion-ion Lattice points occupied by atoms Held together by covalent bonds Hard, high melting point Poor conductor of heat and electricity diamond graphite carbon atoms

43. Types of Crystals Molecular Crystals Lattice points occupied by molecules Held together by intermolecular forces Soft, low melting point Poor conductor of heat and electricity Example: P 4, Cl 2

44. Types of Crystals Cross Section of a Metallic Crystal nucleus & inner shell e - mobile “sea” of e - Metallic Crystals – Typically weaker than covalent, but can be in the low end of covalent Lattice points occupied by metal atoms Held together by metallic bonds Soft to hard, low to high melting point Good conductors of heat and electricity

45. Crystalline quartz (SiO 2 ) Non-crystalline quartz glass An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing

46. Polarity and Solubility I. “Like dissolves like” A. polar dissolves polar and nonpolar dissolves nonpolar. B. Therefore, polar does NOT dissolve nonpolar and vice versa. C. When answering these questions, ALWAYS state polarity!