1. Phases of Matter Liquids, Solids (Crystals) & Solutions Colligative Properties Dr. Ron Rusay Intermolecular Forces I

2. Changes of State –Phase transitions –Phase Diagrams Liquid State –Pure substances and colligative properties of solutions, which depend upon the ratio of the number of solute particles to the number of solvent molecules in a solution. They are independent of the nature of the solute particles. Intermolecular Forces: Phases of Matter & Colligative Properties

3. Solid State Classification of Solids by Type of Attraction between Units Crystalline solids; crystal lattices and unit cells Structures of some crystalline solids Determining the Crystal Structure by X-ray Diffraction Intermolecular Forces: Phases of Matter

4. Phase Transitions Melting: change of a solid to a liquid. Freezing: change a liquid to a solid. Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called Sublimation. Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called Deposition. H 2 O(s)  H 2 O(l) H 2 O(l)  H 2 O(s) H 2 O(l)  H 2 O(g) H 2 O(s)  H 2 O(g) H 2 O(g)  H 2 O(l) H 2 O(g)  H 2 O(s)

5. Phases of Matter / Intermolecular Forces

6. Phase Changes

7. QUESTION

8. Bonds vs. Intermolecular Forces 16 kJ/mol 431 kJ/mol (150 - 1000 kJ/mol) (Ionic bond 700-4,000 kJ/mol)

9. Ion-Dipole Forces (40-600 kJ/mol) Interaction between an ion and a dipole (e.g. NaOH and water = 44 kJ/mol ) Strongest of all intermolecular forces. Intermolecular Forces

10. Dipole-Dipole Forces (permanent dipoles) Intermolecular Forces 5-25 kJ/mol

11. Dipole-Dipole Forces Intermolecular Forces

12. London or Dispersion Forces An instantaneous dipole can induce another dipole in an adjacent molecule (or atom). The forces between instantaneous dipoles are called London or Dispersion forces ( 0.05-40 kJ/mol). Intermolecular Forces

13. London Dispersion Forces Intermolecular Forces Which has the higher attractive force?

14. London Dispersion Forces Intermolecular Forces

15. Gecko: toe, setae, spatulae 6000x Magnification http://micro.magnet.fsu.edu/primer/java/electronmicroscopy/magnify1/index.html Geim, Nature Materials (2003) Glue-free Adhesive 100 x 10 6 hairs/cm 2 Full et. al., Nature (2000) 5,000 setae / mm 2 600x frictional force; 10 -7 Newtons per seta

16. Boiling Points & Hydrogen Bonding

17. Boiling Points & Hydrogen Bonding

18. Hydrogen bonds, a unique dipole-dipole (10- 40 kJ/mol).

19. Fig. 1 STM and AFM measurements J Zhang et al. Science 2013;342:611-614 Published by AAAS Visualizing Intermolecular Hydrogen Bonds

20. (Left) Fig. 2 AFM measurements of 8-hq assembled clusters on Cu(111) (Right) Fig. 4 AFM measurements of coordination complexes J Zhang et al. Science 2013;342:611-614 Visualizing Intermolecular Hydrogen Bonds

22. Which pure substances will not form hydrogen bonds? I) CH 3 CH 2 OHII) CH 3 OCH 3 III) H 3 C−NH−CH 3 IV) CH 3 F A) I and II B) I and III C) II and III D) II and IV QUESTION

23. Hydrogen Bonding Intermolecular Forces

24. DNA: Size, Shape & Self Assembly http://www.umass.edu/microbio/chime/beta/pe_alpha/atlas/atlas.htm Views & Algorithms 10.85 Å

25. Intermolecular Forces

26. Protein Shape: Forces, Bonds, Self Assembly, Folding 10-40kJ/mol 700-4,000kJ/mol 150-1000kJ/mol 0.05-40kJ/mol Ion-dipole (Dissolving) 40-600kJ/mol

27. Predict which liquid will have the strongest intermolecular forces of attraction (neglect the small differences in molar masses). A) CH 3 COCH 2 CH 2 CH 3 (molar mass = 86 g/mol) B) CH 3 CH 2 CH 2 CH 2 CH 2 OH (molar mass = 88 g/mol) C) CH 3 CH 2 CH 2 CH 2 CH 2 CH 3 (molar mass = 86 g/mol) D) HOH 2 C−CH=CH−CH 2 OH (molar mass = 88 g/mol) QUESTION

28. Vapor Pressure on the Molecular Level Vapor Pressure Would water have a higher or lower vapor pressure @ the same temperature? ( bp H 2 O > CH 3 CH 2 OH; bp = o C when Vapor Pressure = Atmospheric Pressure )

29. Explaining Vapor Pressure on a Molecular Level Vapor Pressure

30. Volatility, Vapor Pressure, and Temperature Vapor Pressure

31. Solid Liquid Gas Vapor Pressure

32. QUESTION

33. Temperature & Vapor Pressure The boiling point (b.p.) of a pure liquid is the temperature at which the vapor pressure above the liquid equals the external pressure. Could water boil @ 0 o C?

34. Temperature Dependence of Vapor Pressures The vapor pressure above the liquid varies exponentially with changes in the temperature. The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. (R = 8.314 J K −1 mol −1 )J K mol

35. Clausius – Clapeyron Equation A straight line plot results when ln P vs. 1/T is plotted and has a slope of  H vap /R. Clausius – Clapeyron equation is true for any two pairs of points.

36. QUESTION

37. Heating Curve http://chemconnections.org/general/movies/HeatingCurves.swf

38. Energy (Heat) and Phase Changes Heat of vaporization: heat needed for the vaporization of a liquid. H 2 O(l)  H 2 O(g)  H = 40.7 kJ/mol Heat of fusion: heat needed for the melting of a solid. H 2 O(s)  H 2 O(l)  H = 6.01 kJ/mol Temperature does not change during the change from one phase to another. 50.0 g of H 2 O(s) and 50.0 g of H 2 O(l) were mixed together at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water.