AP Chapter 11

1. AP Chapter 11 Intermolecular Forces / Liquids and Solids HW: 2 4 6 13 15 16 19 31 39 45 105

2. 11.1 – A Comparison of the States of Matter The fundamental difference between states of matter is the distance between particles. Indefinite volume Indefinite shape Rotational, Vibrational and Translational Particle Motion Definite volume Indefinite shape Rotational and Vibrational Definite volume Definite shape Rotational Motion

3. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

4. Phase • Homogeneous part of a system in contact with other parts of the system but separated from them by a well defined boundary • The state/phase a substance is in at a particular temperature and pressure depends on two properties: • The kinetic energy of the particles • The strength of the attractions between the particles

5. 11.2 - Intermolecular Forces (External Bonds) The attractions between molecules are not nearly as strong as the intramolecular attractions (covalent bonds) that hold compounds together.

6. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

7. 1. Ion-Dipole Interactions • Ion-dipole interactions are an important force in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents • Hydration = Process of water molecules surrounding a solute

8. 2. Intermolecular Forces Intermolecular forces between molecules are referred to as van der Waals forces.

9. van der Waals Forces • a. Dipole-dipole interactions • b. London Dispersion forces • Induced Dipoles (not a true London Dispersion Force) • Instantaneous Dipoles

10. a. Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other. • The larger the dipole moment, the stronger the attraction

11. Dipole-Dipole Interactions The more polar (larger resultant dipole moment) the molecule, the higher is its boiling point because the attractions are stronger.

12. Hydrogen Bonding • Can be considered a unique kind of dipole-dipole interaction • Hydrogen atom in a polar bond externally bonds to the non-bonding electron pair (lone pair) on an electronegative atom • Polar bond usually the H connected to a N, O, or F • We call these interactions hydrogen bonds. • Often seen in biological systems.

13. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Because when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

14. b. London Dispersion Forces(occur between nonpolar substances/atoms)-Induced Dipoles (not all books talk about these. Not a TRUE London Dispersion Force) • Occurs when a nonpolar atom/molecule comes in proximity to charged particle/molecule • Induces a charge on the nonpolar molecule And then is attracted to dipole + ->

15. London Dispersion Force-Instantaneous Dipole While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom. Causes an instantaneous dipole. The same can occur in any molecule…

16. Instantaneous Dipole Looks like this At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

17. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

18. London Dispersion Forces -A London Dispersion Force Occurs -London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

19. London Dispersion Forces • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability.

20. Factors Affecting London Forces • The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). • This is due to the increased surface area in n-pentane. (that’s why oil is thicker than water)

21. Factors Affecting London Forces • The strength of dispersion forces tends to increase with increased molecular weight. • Larger atoms have larger electron clouds, which are easier to polarize.

22. Dipole-Dipole Interactions vs. London Forces • Nonpolar molecules only have London Forces • Dipole molecules have both Dipole-Dipole and London Forces • Which is more important to the property of the molecule?? • When comparing two dipole molecules with different BP, etc.: • If size, shape and # of e-s of the two molecules are very similar, the reason the BPs are different are due to dipole-dipole interactions • If the molecules are of very different size than dispersion forces are the main explanation

23. Dispersion Forces vs. Dipole-Dipole Interactions Why choose these two molecules to compare? Both have identical numbers of electrons, and if you made models you would find that the sizes were similar - as you can see in the diagrams. That means that the dispersion forces in both molecules should be much the same. The higher boiling point of fluoromethane is due to the large permanent dipole on the molecule because of the high electronegativity of fluorine. However, even given the large permanent polarity of the molecule, the boiling point has only been increased by some 10°. (Dipole-Dipole Interactions don’t always make that much of a difference).

24. The boiling points are: Dispersion Forces vs. Dipole-Dipole Interactions Here is another example showing the dominance of the dispersion forces. Trichloromethane, CHCl3, is a highly polar molecule because of the electronegativity of the three chlorines. There will be quite strong dipole-dipole attractions between one molecule and its neighbours. On the other hand, tetrachloromethane, CCl4, is non-polar. The outside of the molecule is uniformly - in all directions. CCl4 has to rely only on dispersion forces So which has the highest boiling point? CCl4 does, because it is a bigger molecule with more electrons. The increase in the dispersion forces more than compensates for the loss of dipole-dipole interactions.

25. 11.3 - Physical Properties of Liquids The strength of the attractions between particles can greatly affect the properties of a substance or solution.

26. Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

27. Surface Tension No molecules pull upward on the surface Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

28. Capillary Action • Liquid in a tube is not at the same level as the liquid in the surrounding container -Cohesion – Intermolecular attraction between liquid molecules -Adhesion – Attraction between unlike molecules (ex – liquid and the tube) Adhesion > Cohesion Cohesion > Adhesion

29. 11.4 – Phase Changes • Transfer from one phase to another by addition or removal of energy • Liquid – Vapor Equilibrium • When molecules have enough energy they escape the surface • Higher T, Higher E, More Evaporation

30. Vapor Pressure • At any temperature, some molecules in a liquid have enough energy to escape. • As the temperature rises, the fraction of molecules that have enough energy to escape increases.

31. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

32. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate. (in a closed container)

33. Vapor Pressure • The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. • The normal boiling point is the temperature at which its vapor pressure is 760 torr. • Molar Heat of Vaporization = Energy needed to vaporize one mole of liquid.

34. Clausius-Clapeyron Equation (p. 456) • Describes the relationship between P (vapor pressure) and T • Plot ln P on y-axis vs 1/T on x-axis • Slope = -DHvap/R • y-intercept = c • R = 8.314 J/molK • Comparing two Temps: ln(Pvap) = -DHvap/(RT) + c Sometimes shown without – and T2 and T1 reversed…

35. Boiling Point • Temp at which VP liquid = P external • Water = Normal BP = 100oC • Normal = P external = 1 atm

36. Critical Temperature and Pressure • Gas -> Liquid occurs by either decrease temp or increase pressure • Critical Temp = Tc = “Last” temperature at which a liquid can exist. At ALL pressures above this temp, only a gas can exist • Critical Pressure = Pc = Minimum pressure that must be applied to bring about liquification at the Critical Temperature

37. Liquid – Solid Equilibrium Heating / Cooling Curves: • Freezing = l->s • Melting/Fusion = s->l • Melting Point / Freezing Point = Temperature at which solid and liquid are at equilibrium • Water = Normal fp = 0oC • Molar Heat of Fusion = Energy required to melt 1 mole of a solid

38. Supercooling • Liquid temporarily cooled below Freezing Point • Happens when removed from heat so rapidly that molecules don’t have time to form solid. • Unstable • Can solidify by seed crystal or scratch on side of container

39. Solid-Vapor Equilibrium • Sublimation = s->g • Deposition = g->s • Molar Heat of Sublimation = Energy required to sublime 1 mole solid

40. Heat Calculations • BEFORE Phase Changes (temp is changing): q = msDT • DURING Phase Changes (temp constant due to all energy going into changing bonds): q = m*1/M*DH moles

41. 11.6 - Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

42. Phase Diagrams • Triple Point = Conditions at which s, l, g all coexist • Water = 0.01oC and 0.006 atm • The AB line is the liquid-vapor interface. • It starts at the triple point (A), the point at which all three states are in equilibrium.

43. Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

44. Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

45. Phase Diagrams • The AD line is the interface between liquid and solid. • The melting point at each pressure can be found along this line.

46. Phase Diagrams • Below A the substance cannot exist in the liquid state. • Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

47. Phase Diagram of Water • Note the high critical temperature and critical pressure: • These are due to the strong van der Waals (both dipole-dipole and London Dispersion) forces between water molecules.

48. Structure and Properties of Water • Water is a very unique compound • Has a VERY high specific heat (can absorb large amounts of heat but with small change in temperature) • Solid is less dense than the liquid • Maximum density occurs at 4oC

49. Phase Diagram of Water • The slope of the solid–liquid line is negative. • This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid. • RARE

50. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.