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Intermolecular Forces: PowerPoint Presentation
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Intermolecular Forces:

Intermolecular Forces:

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Intermolecular Forces:

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  1. 0 Intermolecular Forces: What holds everything together (Chapter 14)

  2. 0 Intramolecular forces (bonds) • Hold atoms together in molecules • Have high energy associated with them • it’s difficult to break molecules into their individual atoms • Different types based upon what is going on with the electrons (electron clouds)

  3. 0 Types of bonds: • Ionic • attraction between fully charged molecules/ atoms • NaCl, made from Na+ and Cl-; or • Ca(OH)2, made from Ca2+ and 2OH- • Covalent • electrons are shared between atoms, • water (H2O) and • sugar (C6H12O6) • Can be polar or nonpolar • Based on • electronegativity • VSEPR geometry (shape)

  4. 0 Intermolecular forces (IMFs) • Hold molecules together • Much weaker than intramolecular forces • Intramolecular bonds are usually 100x or even 1000x stronger *(kJ are units of energy like Calories; 1Cal= 4.184kJ) • 1000cal= 1Cal • 1cal =4.184J

  5. 0 Figure 14.2: Intermolecular forces exist between molecules. Bonds exist within molecules.

  6. 0 Why do we care? • The strength of the IMFs determine the state of matter • Solid, liquid, or gas* • *Not plasma- intramolecular bonds are broken to get plasmas

  7. Solids, Liquids, and Gases 0 *all at room temperature, ~25C **small variations occur due to temperature changes, very little variable with pressure changes

  8. 0 • Things with strong IMFs tend to be solids at room temperature • Things with weak IMFs tend to be gases at room temperature • Medium IMFs tend to be in between- • liquids, yes, but with varying characteristics • Amorphous solids: long transition between solid and liquid states- gets soft, then melts (like wax) • Crystalline solids: definite, clear melting point (no soft transition- ie: ice)

  9. 0 Types of IMFs • In order of increasing strength: • London dispersion forces • Dipole- dipole • Hydrogen bonds

  10. 0 London dispersion forces • LDFs occur in all molecules, but are the only forces that are present in nonpolar molecules such as diatomic molecules and atomic substances • CO2, N2, He • They occur because the electron clouds around molecules are not always evenly distributed. • When the electron clouds are unevenly distributed, temporary partial charges result

  11. 0 Figure 14.6: Atoms with spherical electron probability. 14.6: The atom on the left develops an instantaneous dipole.

  12. 0 LDFs, con’t • These temporary partial charges are called induced or temporary dipoles • This temporary dipole forming in a nonpolar substance is strong enough to cause a dipole to occur in a neighboring molecule

  13. 0 Figure 14.3: (a) Interaction of two polar molecules. (b) Interaction of many dipoles in a liquid.

  14. 0 LDfs, con’t • Basically, everything lines up temporarily, but long enough to keep everything together • Common in gases

  15. 0 See LDFs at work here • http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-forces.shtml • These dipoles fluctuate; they do not last very long, but they do occur frequently enough to have a significant effect overall

  16. 0 Dipole- dipole forces: • Are stronger than LDfs because they occur in polar molecules that already have permanent dipole moments (in other words, partial charges already exist) • Are AKA as van der Waals interactions at times, but in actuality both induced dipole attractions and dipole-dipole attractions are van der Waals forces

  17. 0 Examples • HCl and other acids* • HCN • NH3 • *except HF, which does something else

  18. 0 • What would happen between polar and nonpolar molecules? (Do forces of attraction exist? Do the molecules repel?) Explain!

  19. 0 Hydrogen bonding • Are stronger than dipole-dipole forces or LDFs • Occurs in only the most polar bonds • between molecules containing H-F, H-O and H-N bonds only • Are the reason that water is so different from any material from similar atoms, like H2S

  20. 0 Figure 14.4: Hydrogen bonding among water molecules.Norton Interactive: IMFs tutorialSelect Hydrogen bonding in water from bottom of list

  21. 0 • http://www.northland.cc.mn.us/biology/Biology1111/animations/hydrogenbonds.html • (note: I am not responsible for the music on the above web site) • Polarity and hydrogen bond formation • Ice at different temperatures

  22. 0 Which is ice? Which is liquid water? Explain. • Ice at different temperatures

  23. 0 Water is special because… • It has a high specific heat, meaning that it takes a lot of energy to raise the temperature of a sample of water by even 1 degree • Specific heat of water (c)= 1 cal/ g°C or 4.184J /g°C • The solid phase is LESS dense than the liquid phase, so ice floats on water • It’s a good solvent for many substances due its polarity • H2O is liquid at RT, where H2S is a gas

  24. 0 Figure 14.5: The boiling points of covalent hydrides.

  25. 0 Water is special • And water would not be special without hydrogen bonding • H bonding plays vital roles in • DNA (holding together the chains of DNA) • Protein shape (and therefore the protein’s function; think hair!)

  26. H bonding in dna1

  27. H bonding in DNA

  28. Amino Acids- they make proteins

  29. Protein Structures

  30. Protein Structure and H Bonding

  31. 0 For the next slides: • Determine polarity of group • Determine type of IMFs are possible in group • Determine if the group will be highly soluble in water

  32. 0 IMFs in proteins

  33. 0 Sickle Cell Anemia • Glu (glutamic acid) replaced by Val (valine)

  34. 0 • What would happen if a molecule capable of H-bonding comes into contact with: • A nonpolar substance • A polar substance that does not H-bond

  35. 0 • Strength increases from left to right; when ions are involved, attractive forces are greater than when they are not involved. • http://cwx.prenhall.com/bookbind/pubbooks/blb/chapter11/medialib/blb1102.html

  36. 0 Dealing with this pic… • Ion- dipole forces • Ionic Bonding • Basically electrostatic attractive forces between positive and negative charges • Strong

  37. 0 IMFs influence… • Boiling point/ Melting Point • Viscocity • Surface Tension • Capillary Action • Vapor pressure/ rate of evaporation • State of Matter (at room temp) • Density falls here, but can vary even within state

  38. 0 IMFs and mass • The mass of a material makes a difference, so yes, mass (size) matters • Larger molecules have stronger forces than similar molecules that are smaller (in terms of mass)

  39. 0 Figure 14.5: The boiling points of covalent hydrides.

  40. 0 Boiling points and masses of noble gases • Helium: -269°C 4.00 g/mol • Neon: -246°C 20.18 g/mol • Argon: -186°C 39.95 g/mol • Krypton: -152°C 83.80 g/mol • Xenon: -108°C 131.3 g/mol • radon -62°C ~222 g/mol Larger atoms have larger e- clouds, which lead to greater polarizability

  41. 0 Saturated Hydrocarbons, or Alkanes As melting point increases, boiling point increases (saturated hydrocarbons are hydrocarbons with as many Hs as possible)

  42. 0 Shape also matters • Butane, bp -0.5 degrees C • 2-methylpropane -11.7 degrees C Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules. Also, the molecules can stack with each other better H

  43. 0 Butane and 2-methylpropane Compare the properties of these two compounds: n-butane . . . . . . . . . . . . . . . . .. . . . . . . . . . . . . 2-methylpropane 0.601 . . . . . . . . . . . . . . . . relative density (liquid) . . . . . . . . . . . . . . . . 0.551 1.348 . . . . . . . . . . . . . . . . refractive index (liquid) . . . . . . . . . . . . . . . .1.351 - 0.5 . . . . . . . . . . .. . . . . . . boiling point (oC) . . . . . . . . . . . . . . . . . . .. . - 11.7 - 138.3 . . . . . . . . . . . . . . . . melting point (oC) . . . . . . . . . . . . . . .. . . . - 159.6 It is clear that the different carbon skeletons make a difference to the properties, especially the melting and boiling points.

  44. 0 Fats v Oils:Saturated v. Unsaturated • Molecular size, bond order, and bond orientation: • How different IMFs result in differences in food molecules

  45. 0 • A carbon exists where two lines intersect • Atoms other than C and H are written in • Hs are not usually written out- • They fill in to complete octets on other atoms

  46. 0 Random cis trans fats • (Omega 3 and Omega 6 fats have the double bonds on the 3 or 6th carbon)

  47. 0 fatty acids and triglycerides • 3 Fatty acid chains (above) join with a glycerol molecule (top right) to form a triglyceride (right, saturated)

  48. 0 Triglyceride formation

  49. 0 Triglycerides • Oils • More unsaturated FAs • Liquid at RT • Fats • More saturated FAs • Solid at RT