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Intermolecular Forces:

0. Intermolecular Forces:. What holds everything together (Chapter 14). 0. Intra molecular forces (bonds). Hold atoms together in molecules Have high energy associated with them it’s difficult to break molecules into their individual atoms

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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

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