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Ch 10: Intermolecular Forces and Types of Solids

Ch 10: Intermolecular Forces and Types of Solids. AP Chemistry Sequoyah High School. 10.1: Intermolecular Forces (IMF). IMF < intramolecular forces (covalent, metallic, ionic bonds) IMF strength: solids > liquids > gases

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Ch 10: Intermolecular Forces and Types of Solids

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  1. Ch 10: Intermolecular Forces and Types of Solids AP Chemistry Sequoyah High School

  2. 10.1: 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. 10.1: Types of IMF • Electrostatic forces: act over larger distances in accordance with Coulomb’s law http://dwb4.unl.edu/ChemAnime/attractive_forces.htm a.Ion-ion forces: strongest; found in ionic crystals (i.e. lattice energy) http://chemmovies.unl.edu/ChemAnime/LICLD/LICLD.html

  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). http://wps.prenhall.com/wps/media/objects/439/449969/Media_Portfolio/Chapter_14/Dissolution_NaCl_Water.MO <

  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)? http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-forces.shtml http://usm.maine.edu/~newton/Chy251_253/Lectures/CarbonylReduction/AldehydesKetones.html

  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. http://www.visionlearning.com/library/module_viewer.php?mid=57 ( ice and water simulation) http://www.youtube.com/watch?v=LGwyBeuVjhU ( H bond movie) http://www.youtube.com/watch?v=lkl5cbfqFRM&feature=related ( H bonding in water)

  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. http://www.youtube.com/watch?v=PcoiLAsUvqc&feature=related http://www.youtube.com/watch?v=gmjLXrMaFTg&feature=related http://en.wikipedia.org/wiki/Water_%28molecule%29#Density_of_water_and_ice

  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 • http://dwb4.unl.edu/ChemAnime/LONDOND/LONDOND.html • 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. 10.2: Properties resulting from IMF • Viscosity: resistance of a liquid to flow • Surface tension: energy required to increase the surface area of a liquid

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

  19. 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)

  20. 10.8: Phase Changes Processes: • Endothermic: melting, vaporization, sublimation • Exothermic: condensation, freezing, deposition I2 (s) and (g) Microchip

  21. Water: Enthalpy diagram or heating curve

  22. 10.8: Vapor pressure A liquid will boil when the vapor pressure equals the atmospheric pressure, at any T above the triple point. http://glencoe.com/sites/common_assets/advanced_placement/chemistry_chang9e/animations/chang_2e/vapor_pressure.swf http://dwb4.unl.edu/ChemAnime/VPTEMPD/VPTEMPD.html http://dwb4.unl.edu/ChemAnime/VP3LIQD/VP3LIQD.html Pressure cooker ≈ 2 atm Normal BP = 1 atm 10,000’ elev ≈ 0.7 atm 29,029’ elev (Mt. Everest)≈ 0.3 atm

  23. 10.9: 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.)

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

  25. *

  26. *

  27. 10.3-7: Structures of solids • Amorphous: without orderly structure Ex: rubber, glass • Crystalline: repeating structure; have many different stacking patterns based on chemical formula, atomic or ionic sizes, and bonding

  28. Types of crystalline solids (Table 11.6)

  29. Carbon dioxide (dry ice) Sucrose Ice

  30. Graphite Diamond SiO2

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