Chapter 11 Intermolecular Forces

1. Chapter 11 Intermolecular Forces

2. 11.1: Intermolecular Forces (IMF) • IMF < intramolecular forces (covalent, metallic, ionic bonds) • IMF strength: solids > liquids > gases • Boiling points and melting points are good indicators of relative IMF strength.

3. 11.2: Types of IMF • Electrostatic forces: act over larger distances in accordance with Coulomb’s law • Ion-ion forces: strongest; found in ionic crystals (i.e. lattice energy)

4. d+ d+ d+ d+ d- d- d+ d+ d- d- Cl- S2- d+ d+ d+ d+ d- d- d+ d+ • Ion-dipole: between an ion and a dipole (a neutral, polar molecule/has separated partial charges) • Increase with increasing polarity of molecule and increasing ion charge. Ex: Compare IMF in Cl- (aq) and S2- (aq). <

5. Dipole-dipole: weakest electrostatic force; exist between neutral polar molecules • Increase with increasing polarity (dipole moment) of molecule Ex: What IMF exist in NaCl (aq)?

6. Hydrogen bonds (or H-bonds): • H is unique among the elements because it has a single e- that is also a valence e-. • When this e- is “hogged” by a highly EN atom (a very polar covalent bond), the H nucleus is partially exposed and becomes attracted to an e--rich atom nearby.

7. H-bonds form with H-X•••X', where X and X' have high EN and X' possesses a lone pair of e- • X = F, O, N (since most EN elements) on two molecules: F-H O-H N-H :F :O :N

8. * There is no strict cutoff for the ability to form H-bonds (S forms a biologically important hydrogen bond in proteins). • * Hold DNA strands together in double-helix Nucleotide pairs form H-bonds DNA double helix

9. H-bonds explain why ice is less dense than water.

10. Ex: Boiling points of nonmetal hydrides Conclusions: • Polar molecules have higher BP than nonpolar molecules • ∴ Polar molecules have stronger IMF • BP increases with increasing MW • ∴ Heavier molecules have stronger IMF Boiling Points (ºC) • NH3, H2O, and HF have unusually high BP. • ∴ H-bonds are stronger than dipole-dipole IMF

11. Inductive forces: • Arise from distortion of the e- cloud induced by the electrical field produced by another particle or molecule nearby. • London dispersion:between polar or nonpolar molecules or atoms • * Proposed by Fritz London in 1930 • Must exist because nonpolar molecules form liquids Fritz London(1900-1954)

12. How they form: • Motion of e- creates an instantaneous dipole moment, making it “temporarily polar”. • Instantaneous dipole moment induces a dipole in an adjacent atom • * Persist for about 10-14 or 10-15 second Ex: two He atoms

13. * Geckos! • Geckos’ feet make use of London dispersion forces to climb almost anything. • A gecko can hang on a glass surface using only one toe. • Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb walls http://en.wikipedia.org/wiki/Van_der_Waals%27_force

14. London dispersion forces increase with: • Increasing MW, # of e-, and # of atoms (increasing # of e- orbitals to be distorted) Boiling points: Effect of MW: Effect of # atoms: pentane 36ºCNe –246°C hexane 69ºCCH4   –162°C heptane 98ºC ??? effect: H2O 100°C D2O 101.4°C • “Longer” shapes (more likely to interact with other molecules) C5H12 isomers: 2,2-dimethylpropane 10°C pentane 36°C

15. Summary of IMF Van der Waals forces

16. Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points.

17. 11.3: Properties resulting from IMF • Viscosity: resistance of a liquid to flow Viscosity depends on: -the attractive forces between molecules -the tendency of molecules to become entangled -the temperature

18. 11.3: Properties resulting from IMF • Surface tension: energy required to increase the surface area of a liquid

19. 3. Cohesion:attraction of molecules for other molecules of the same compound 4. Adhesion:attraction of molecules for a surface

20. Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forces Ex: Hg H2O (cohesion rules) (adhesion rules)

21. Phase Changes • Surface molecules are only attracted inwards towards the bulk molecules. • Sublimation: solid  gas. • Vaporization: liquid  gas. • Melting or fusion: solid  liquid. • Deposition: gas  solid. • Condensation: gas  liquid. • Freezing: liquid  solid. Energy Changes Accompanying Phase Changes • Energy change of the system for the above processes are:

22. Intermolecular Forces Bulk and Surface

23. Phase Changes Energy Changes Accompanying Phase Changes • Sublimation: Hsub > 0 (endothermic). • Vaporization: Hvap > 0 (endothermic). • Melting or Fusion: Hfus > 0 (endothermic). • Deposition: Hdep < 0 (exothermic). • Condensation: Hcon < 0 (exothermic). • Freezing: Hfre < 0 (exothermic). • Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: • it takes more energy to completely separate molecules, than partially separate them.

24. Phase Changes Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions (e.g. water sublimes when snow disappears without forming puddles). • The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. • The sequence cool gas  condense  cool liquid  freeze  cool solid is exothermic.

25. Phase Changes Energy Changes Accompanying Phase Changes

26. Phase Changes Heating Curves • Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. • These points are used to calculate Hfus and Hvap. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

27. Phase Changes Heating Curves

28. Heating Curve Illustrated

29. Phase Changes Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction.

30. Critical Temperature, Tc

31. Transition to Supercritical CO2

32. Supercritical CO2 Used to Decaffeinate Coffee

33. Vapor Pressure Explaining Vapor Pressure on the Molecular Level • Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. • These molecules move into the gas phase. • As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. • After some time the pressure of the gas will be constant at the vapor pressure.

34. Gas-Liquid Equilibration

35. Vapor Pressure Explaining Vapor Pressure on the Molecular Level • Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. • Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

36. Vapor Pressure Volatility, Vapor Pressure, and Temperature • If equilibrium is never established then the liquid evaporates. • Volatile substances evaporate rapidly. • The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

37. Liquid Evaporates when no Equilibrium is Established

38. Vapor Pressure Volatility, Vapor Pressure, and Temperature

39. Vapor Pressure Vapor Pressure and Boiling Point • Liquids boil when the external pressure equals the vapor pressure. • Temperature of boiling point increases as pressure increases. • Two ways to get a liquid to boil: increase temperature or decrease pressure. • Pressure cookers operate at high pressure. At high pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. • Normal boiling point is the boiling point at 760 mmHg (1 atm).

40. Phase Diagrams • Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. • Given a temperature and pressure, phase diagrams tell us which phase will exist. • Features of a phase diagram: • Triple point: temperature and pressure at which all three phases are in equilibrium. • Vapor-pressure curve: generally as pressure increases, temperature increases. • Critical point: critical temperature and pressure for the gas. • Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. • Normal melting point: melting point at 1 atm.

41. Phase Diagrams • Any temperature and pressure combination not on a curve represents a single phase.

42. Phase Diagrams The Phase Diagrams of H2O and CO2 • Water: • The melting point curve slopes to the left because ice is less dense than water. • Triple point occurs at 0.0098C and 4.58 mmHg. • Normal melting (freezing) point is 0C. • Normal boiling point is 100C. • Critical point is 374C and 218 atm. • Carbon Dioxide: • Triple point occurs at -56.4C and 5.11 atm. • Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes it does not melt.) • Critical point occurs at 31.1C and 73 atm.

43. Phase Diagrams The Phase Diagrams of H2O and CO2

44. 11.4: Phase Changes Processes: • Endothermic: melting, vaporization, sublimation • Exothermic: condensation, freezing, deposition I2 (s) and (g) Microchip

45. Water: Enthalpy diagram or heating curve

46. 11.5: Vapor pressure • A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point. Pressure cooker ≈ 2 atm Normal BP = 1 atm 10,000’ elev ≈ 0.7 atm 29,029’ elev (Mt. Everest)≈ 0.3 atm

47. 11.6: Phase diagrams: CO2 • Lines: 2 phases exist in equilibrium • Triple point: all 3 phases exist together in equilibrium (X on graph) • Critical point, or critical temperature & pressure: highest T and P at which a liquid can exist (Z on graph) Temp (ºC) • For most substances, inc P will cause a gas to condense (or deposit), a liquid to freeze, and a solid to become more dense (to a limit.)

48. Phase diagrams: H2O • For H2O, inc P will cause ice to melt.

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