Lewis Structures: 5 steps

1. Lewis Structures: 5 steps • Count valence e- available If an ANION, add charge to # valence e- If a CATION, subtract charge from # valence e- • Draw skeleton Least electronegative atom is the central atom (usually first element in formula, never H) Place surrounding atoms around the central atom add dashes to show bonds from central atom to surrounding atoms

2. Lewis Structures-cont’d 3. Satisfy the octet rule for each surrounding atom using valence electrons. Note that H can only have 2 e- 4. Any remaining valence e- go on the central atom. 5. Check final structure If the octet rule is not satisfied for the central atom, borrow e- pairs to make double or triple bonds. It is possible for the central atom to have an expanded octet, (eg. SF6) or a contracted octet, (eg. BF3). Theses are exceptions to the octet rule.

3. Lewis Structures– H2O • # val e- available = 2 H = 2(1) = 2 1 O = 1(6) = 6 = 8 e- available 2. Skeleton (O is central…) H O H

4. Lewis Structures– H2O Each dash is 2 e- 3. Add dots so that each has 8 e- 4. Count e- shown : 2 dashes = 4 e- 4 dots = 4 e- = 8 e- total 5. Check: 8 e- shown = 8 e- available So OK!!

5. Lewis Structure, C2H4 • # e- available = 2 C = 2(4) = 8 4 H = 4(1) = 4 = 12 e- available 2. Skeleton (C---C is central…)

6. Lewis Structure, C2H4, cont’d 3. Add dots so that each has 8 e- Remember each dash = 2 e- 4. Count e- shown: 4 C-H bonds = 4 (2) = 8 e- 1 C-C bond = 1 (2) = 2 e- 4 dots = 4 e- total = 14 e-

7. Lewis Structure, C2H4, cont’d 5. Compare # e- shown to # e- available 14 e- shown >12 e- available So must make a double bond: • Remove 1 electron from each C • Combine remaining 2 e- in a covalent bond, forming a double bond

8. Lewis Structure, C2H4, cont’d Check: # e-shown = # e- available 12 = 12 OK!! 4 C-H bonds = 4 (2) = 8 e- 1 C=C bond = 1 (4) = 4 e- total = 12 e- (or count # lines = 6 6 lines *2e- each = 12 e-)

9. Lewis Structures Exercise • Get in pairs • Molecules assigned • Give “model kits” • Each group use model to make Lewis Structure • Show model, draw LS on board

10. Lewis Structures Exercise- compounds CH4 N2O NO3– NO2 PCl5 NH4 + SO2 SF6 CO2 SO42- NH3 SO32- XeF4 NO2–

11. Molecular Shape – VSEPR Theory • Can relate Lewis Structure to 3-D shape of a molecule Valence Shell Electron Pair Repulsion Theory e- pairs arrange themselves around a nucleus to minimize -/- repulsions e- pairs get as far away from each other as possible and as close to the nucleus as possible

12. VSEPR Theory • Use “AXnEm” designation A = central atom X = atoms bonded to A n = # atoms bonded to E = unsharede- PAIRS on A (lone pairs) m = # unshared e- pairs on A • Get from Lewis Structure!!

13. AXE designation A = “Oxygen Ex— water, H2O X = H n = 2 E = pairs of dots on O m = 2 pairs AX2E2

14. Use AXE to give shape

15. Use AXE to give shape

16. Use AXE to give shape

17. n+m = 5 family

18. Use AXE to give shape

19. n+m = 6 family

20. Be able to… • Draw a Lewis structure for any assigned molecule • Based on the Lewis Structure, give the AXE designation. • Determine molecule shape, bond angles, and polarity.

21. Molecule Polarity • Tell if molecule has one “side” that the electrons like to congregate… • Based on molecule shape and bond polarity

22. Bond Polarity • Even though electrons are shared between two nuclei in covalent bonds, often the sharing is NOT • One atom often has greater affinity for e- than other • Look at differences in electronegativity (EN) • The more EN the atom, the more it “hogs” the shared e-

23. Bond Polarity F H Which is more EN? F or H? F Make the “bond” an arrow pointing toward the more EN atom Put a “+” across the tail (other end) of the arrow So e- spend more time around F So F has a “partial negative” charge d- H has a “partial positive” charge, d+ F H F H d- d+ F H

24. Bond Polarity • The 2 shared e- between H and F tend to spend more time around F

25. Molecule Polarity • Consider both bond polarity AND shape!

26. Is Water a polar molecule? 1. Draw the Lewis Structure 2. Get shape: AX2E2…so BENT

27. Is Water a polar molecule? 3. Look at each OH bond—determine bond polarity O H d- d+ d- d+ d+

28. Is Water a polar molecule? 4. Rotate to look end on / smash into page 5. If different (+/-), then POLAR (if same +/+ or -/-, then NONPOLAR) d- d- d+ d+ d+

29. Are the following molecules polar? • CO2 • CCl4 • CHCl3

30. Intermolecular Forces(IMF) • Attractive forces between 2 or more molecules • Need to consider molecule shape and polarity

31. 3 types of IMF • Dispersion forces: between NONPOLAR molecules (weakest) 2. Dipole: between POLAR molecules (intermediate) 3. Hydrogen bonds: special case of dipole forces, between H in one molecule and O, N, or F in another (strongest)

32. Dipole Forces • +/- attraction between POLAR molecules Partial + on one molecule attracted to the partial – on a neighboring molecule

33. Hydrogen Bonds • Special case of dipole attractions • In each molecule, must have H bonded to O, N, or F Partial + (H) on one molecule attracted to the partial – (O, N, or F) on a neighboring molecule

34. Hydrogen bonds • Individually, each H-bond is weak (compared to a covalent bond) • Collectively, H-bonds are VERY strong, especially in large molecules like proteins or DNA

35. Dispersion Forces • Between 2 non-polar molecules • “temporary dipole – induced dipole” “Temporary Dipole”

36. Dispersion Forces The “temporary dipole” now “induces” a neighboring molecule to become a dipole (d- pushes e- away from it in the neighbor, making that end of the neighbor d+) Dispersion force = +/- attraction between temporary dipole and induced dipole

37. IMFs and States of Matter • Think of the ability of a material to change phases as a measure of the strength of IMFs

38. Solids

39. Melting (s  l) • As Temperature increases, molecules vibrate/move/bounce more and more • Gain enough Kinetic energy (KE) to overcome some IMFs • Molecules can now slide around one another

40. Liquids

41. Boiling (l  g) • As temperature increases, molecules gain more and more energy • Soon overcome all IMF • Molecules no longer “attached” to each other • Escape to the gas phase

42. Gas

43. Solutions • Homogenous mixture of 2 or more substances • Solute = material being dissolved, • Usually a solid • Present in least amount • Solvent = material in which the solute is being dissolved • Usually a liquid • Present in greatest amounts

44. Aqueous Solutions • Solvent = water • Solute = some solid

45. Dissolving Process • Think of what happens on a molecular level

46. Before mixing… Interact w/each other medium IMFs Liquid- so molecules still slide around Interact w/each other Strong IMFs Close packed materials (solid)

47. During Mixing… Replace IMFs from like molecules with IMFs from others Now each solute has an IMF interaction with a solvent molecule

48. “Like Dissolves Like” • If the IMFs between a solute and a solvent are similar, then the solute will dissolve in the solvent! • Ex- NaCl in H2O • NaCl is polar (ionic) H2O is polar (H-bonds) • Similar IMFs, so NaCl will dissolve in water

49. Will CH4 dissolve in H2O? • CH4 is nonpolar… IMF = dispersion forces • H2O is polar…..IMF = H-bonds • Different IMFs so CH4 will NOT dissolve in H2O