1. 10.1 – 10.2 10.8 – 10.9 Intermolecular Forces Go over Tests and Turkey Questions and Read P. 442, 444-456: Monday 12/1 PPT: Tuesday 12/2 – Wednesday 12/3

2. In Class Discussion: Ch. 10 # 3 – 10, 33, 104 Homework after 10.1-10.2, 10.8-10.9 Ch. 10 #35-43 odd, 87, 91

4. Intermolecular Forces (IM Forces) – The forces that hold one molecule next to another. – Between molecules – Related to heat of fusion and vaporization (phase changes) – Dispersion, Dipole-Dipole (H-bonding), ion-dipole, ion- induced dipole, dipole-induced dipole, Ion-ion forces could be called inter and/or intra … arguable Intramolecular Forces (Chemical Bonds) – Forces that hold the atoms together inside of a molecule. – Inside a molecule – Related to heat of reaction (bonds breaking and forming) – Ionic, Covalent, Metallic

5. Intermolecular Forces are generally weaker than Intramolecular Forces. – The attraction between an H and an O in one water molecule is greater than the attraction between the H in one water molecule and the O in another. – The energy needed to boil water is less than the energy needed to break apart a water molecule for a chemical reaction.

10. Kinetic Molecular Theory of Gases (Ch. 5) says that we can neglect the interactions between molecules. – Gases experience negligible IM forces. – Remember Ideal Gas vs. Real Gas Liquids and Solids properties differ from those of gases due to large intermolecular forces.

12. Gas: – KE >> IM Forces – Compressible; expands; volume of container – Fluid (flows quickly, shape of container) Liquid: – KE similar to IM Forces – Condensed phase (incompressible; retains volume) – Fluid (flows, although slowly; shape of container) Solid: – KE << IM Forces – Condensed phase (incompressible; retains volume) – Not a fluid (doesn’t flow; retails shape)

13. Phase changes occur due to changes in Temperature (KE), or Pressure Increasing KE can cause a change from solid to liquid or liquid to gas, based on relationship between KE and IM Forces. – Example: H 2 O is solid ice at -10°C and liquid at 20°C. Increasing Pressure can cause a change from gas to liquid or liquid to solid, acting with IM Forces against particle’s KE – Example: H 2 O boils at 100 0 C at sea level atmospheric pressure. If you decrease the pressure, it will boil at a lower temperature. – Less pressure pushing molecules together, less KE needed to overcome the IM Forces.

14. The Temp and/or Pressure at which phase changes occur depend on the particle’s IM Forces The phase a substances is at room temperature (a given KE) depends on the substance’s IM Forces – Greatest IM Forces = solid – Least IM Forces = gas – Example: I 2 has stronger IM forces than Br 2, which has stronger IM forces than Cl 2. At room temp, I 2 is solid, Br 2 is liquid, and Cl 2 is gas. Boiling Points: I 2 (365°F), Br 2 (138°F), and Cl 2 (-29°F).

15. All Intermolecular Forces are electrostatic, involving attractions between positive and negative species or areas. – Remember: Electrostatic Forces increase with increased charge. Electrostatic Forces decrease with increased distance between charges.

16. Types of Intermolecular Forces (1) Dispersion Forces (2) Dipole-Dipole Forces – Hydrogen Bonding being the strongest (3) Ion – Dipole Forces (4) Induced Forces – Ion induced dipole – Dipole induced dipole

17. Dispersion Forces (P. 447 Brown) Electrostatic attraction between all molecules. Electrons move randomly. Instantaneous dipole moments form during any given second – More e- on one end than the other, causing partial positive and partial negative ends – See Fig 11.4 in Brown and 10.5 in Zumdall

19. Dispersion Forces (P. 447 Brown) Polarizability: the ease with which the charge distribution is distorted; how easy it is for the particle to develop a temporary dipole. Greater Polarizability = stronger dispersion forces = higher boiling point – See Figure 11.5 in Brown more electrons (usually determined by Molar Mass) = greater polarizability – Ex: Br 2 has a higher boiling point than Cl 2 More surface area of electrons = greater polarizability – Ex: isomers of C 5 H 12 have different strengths of dispersion forces; See Figure 11.6 in Brown

22. Remember, ALL particles experience dispersion forces. Larger polarizability (larger particle w/greater surface area) = greater dispersion forces

23. Dipole-Dipole Forces (P. 448-452 Brown) Some particles have a permanent dipole. – These particles are called “polar”. The attraction between the partial negative end of one polar molecule and the partial positive end of a second polar molecule. Greater polarity = stronger dipole-dipole forces = higher boiling point

26. See Figure 11.8 in Brown Lets draw the Lewis Structure of each of these four molecules to analyze their polarity and evaluate what makes one polar than another. Remember: more polar = stronger dipole-dipole forces = higher boiling point – Assuming all have similar dispersion forces. – “For molecules of approximately equal mass and size, the strength of intermolecular attractions increases with increasing polarity … boiling point increases as the dipole moment increases.” Brown P. 449

28. Greater differences in Electronegativities = greater polarity = stronger dipole-dipole forces

29. Dipole-dipole Force strength: BrF > ClF > F 2

30. See Figure 11.9 Brown or 10.4 Zumdall Each color is its own group on the periodic table.

32. See Figure 11.9 Brown or 10.4 Zumdall Each color is its own group on the periodic table. Why do molecules with elements in group 6A have greater boiling points than molecules with elements of similar molar mass in Group 5A? Why do all these molecules (containing elements in group 6A and group 5A bonded to H) have higher boiling points than molecules of H bonded to Group 4A elements?

34. See Figure 11.9 Brown or 10.4 Zumdall Each color is its own group on the periodic table. Why is the boiling point of H 2 Te greater than that of H 2 Se greater, and H 2 Se’s boiling point greater than that of H 2 S? Why is the boiling point of HI greater than that of HBr, and HBr’s boiling point greater than that of HCl?

36. See Figure 11.9 Brown or 10.4 Zumdall Each color is its own group on the periodic table. Why do H 2 O, NH 3, and HF all have higher boiling points than any of the other molecules on this graph? – NOTICE THE DIFFERENCE IN PATTERN WITH THESE 3

37. Hydrogen Bonding (type of dipole-dipole) Attraction between a hydrogen atom that is bonded to a highly electronegative atom (F, O, or N) of one molecule and the lone pair(s) of a highly electronegative atom (F, O, or N) in another molecule. See Figure 11.10 of Brown or Figure 10.3 of Zumdall

40. Hydrogen Bonds Causes 3 small molecules (H 2 O, NH 3, and HF) to be liquid when other molecules with similar molar masses are gas at room temperature. Gives water its extra high specific heat (changes temperature slower than other molecules)

41. Hydrogen Bonding plays a major role in biochemistry. Stablizing structures of protiens. Causes DNA to be double helixed and fold over itself.

50. Hydrogen Bonding is what causes solid H 2 O to be less dense than liquid H 2 O. See Figure 11.11 in Brown or Figure 10.12 in Zumdall Ice floats on top of liquid water. Most solids sink in their own liquid.

54. Ion-Dipole Forces Typically involved in aqueous (water) solutions of ionic solutes.

59. Ion-induced dipole and dipole-induced dipole A permanent dipole (polarity) can be induced in an otherwise non-polar molecule if it is placed next to a polar molecule or an ion. Stronger than dispersion forces of similar sized molecules. Weaker than actual dipole-dipole or ion-dipole forces of similar sized molecules.

61. Properties of Liquids Due to Intermolecular Forces. Greater IM Forces = increase in each Surface Tension – A liquid’s resistance to increasing its surface tension. – A liquid’s desire to keep its surface area to a minimum. Capillary Action – Spontaneous rising of a liquid in a narrow tube – Causes a concave mensicus (See Figure 10.7) Viscosity – A liquid’s resistance to flow – Syrup has greater viscosity than water, causing it to pour slower than water.

62. Stronger IM Forces also Increases boiling point (bp) Decreases vapor pressure (opposite of bp) Increases melting point (mp) Increases specific heat (c)

63. Compare IM Forces Large molecules have greater dispersion forces than small molecules. Polar molecules have dipole-dipole forces when nonpolar molecules do not.

64. Compare IM Forces A nonpolar molecule can have stronger IM forces than a polar molecule if the nonpolar molecule is much larger. – Ex: C 16 H 34 has stronger dispersion forces than H 2 O has both dispersion and dipole-dipole forces.

65. Compare IM Forces A smaller molecule can have stronger IM forces than a larger molecule if the smaller molecule is much more polar. – Ex: 1-propanol (CH 3 CH 2 CH 2 OH) has a boiling point of 97°C when water (H 2 O) has a boiling point of 100°C. They are both polar, but the relative polarities are very different.

66. Figure 10.40

70. Heat Curve

72. Phase Diagram

74. If solid is more dense than liquid, solid-liquid slope is positive. If solid is less dense than liquid (water), solid- liquid slope is negative.

75. In Class Discussion: Ch. 10 # 3 – 10, 33, 104 Homework after 10.1-10.2, 10.8-10.9 Ch. 10 #35-43 odd, 87, 91