1. Chapter 2 Structure and Properties of Organic Molecules Organic Chemistry, 6 th Edition L. G. Wade, Jr. Jo Blackburn Richland College, Dallas, TX Dallas County Community College District  2006,  Prentice Hall

2. Chapter 22 Wave Properties of Electrons Standing wave vibrates in fixed location. Wave function, , mathematical description of size, shape, orientation. Amplitude may be positive or negative. Node: amplitude is zero. =>

3. Chapter 23 Wave Interactions Linear combination of atomic orbitals  on different atoms produce molecular orbitals  on the same atom give hybrid orbitals. Conservation of orbitals. Waves that are in phase add together. Amplitude increases. Waves that are out of phase cancel out. =>

4. Chapter 24 Bonding Region Electrons are close to both nuclei. =>

5. Chapter 25 Sigma Bonding Electron density lies between the nuclei. A bond may be formed by s-s, p-p, s-p, or hybridized orbital overlaps. The bonding MO is lower in energy than the original atomic orbitals. The antibonding MO is higher in energy than the atomic orbitals. =>

6. Chapter 26 Bonding Molecular Orbital Two hydrogens, 1s constructive overlap =>

7. Chapter 27 Anti-Bonding Molecular Orbital Two hydrogens, destructive overlap. =>

8. Chapter 28 H 2 : s-s overlap =>

9. Chapter 29 Cl 2 : p-p overlap => Constructive overlap along the same axis forms a sigma bond.

10. Chapter 210 HCl: s-p overlap Question: What is the predicted shape for the bonding MO and the antibonding MO of the HCl molecule? =>

11. Chapter 211 Pi Bonding Pi bonds form after sigma bonds. Sideways overlap of parallel p orbitals. =>

12. Chapter 212 Multiple Bonds A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. =>

13. Chapter 213 Molecular Shapes Bond angles cannot be explained with simple s and p orbitals. Use VSEPR theory. Hybridized orbitals are lower in energy because electron pairs are farther apart. Hybridization is LCAO within one atom, just prior to bonding. =>

14. Chapter 214 sp Hybrid Orbitals 2 VSEPR pairs Linear electron pair geometry 180° bond angle =>

15. Chapter 215 sp 2 Hybrid Orbitals 3 VSEPR pairs Trigonal planar e - pair geometry 120° bond angle =>

16. Chapter 216 sp 3 Hybrid Orbitals 4 VSEPR pairs Tetrahedral e - pair geometry 109.5° bond angle =>

17. Chapter 217 Sample Problems Predict the hybridization, geometry, and bond angle for each atom in the following molecules: Caution! You must start with a good Lewis structure! NH 2 CH 3 -C  C-CHO =>

18. Chapter 218 => Single bonds freely rotate. Double bonds cannot rotate unless the bond is broken. Rotation around Bonds

19. Chapter 219 Isomerism Same molecular formula, but different arrangement of atoms: isomers. Constitutional (or structural) isomers differ in their bonding sequence. Stereoisomers differ only in the arrangement of the atoms in space. =>

20. Chapter 220 Structural Isomers =>

21. Chapter 221 Stereoisomers Cis - same side Trans - across Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp 2 carbon. No cis-trans isomers possible =>

22. Chapter 222 Bond Dipole Moments are due to differences in electronegativity. depend on the amount of charge and distance of separation. In debyes,  x  (electron charge) x d(angstroms) =>

23. Chapter 223 Molecular Dipole Moments Depend on bond polarity and bond angles. Vector sum of the bond dipole moments. =>

24. Chapter 224 Effect of Lone Pairs Lone pairs of electrons contribute to the dipole moment. =>

25. Chapter 225 Intermolecular Forces Strength of attractions between molecules influence m.p., b.p., and solubility, esp. for solids and liquids. Classification depends on structure.  Dipole-dipole interactions  London dispersions  Hydrogen bonding =>

26. Chapter 226 Dipole-Dipole Forces Between polar molecules. Positive end of one molecule aligns with negative end of another molecule. Lower energy than repulsions, so net force is attractive. Larger dipoles cause higher boiling points and higher heats of vaporization. =>

27. Chapter 227 Dipole-Dipole =>

28. Chapter 228 London Dispersions Between nonpolar molecules Temporary dipole-dipole interactions Larger atoms are more polarizable. Branching lowers b.p. because of decreased surface contact between molecules. =>

29. Chapter 229 Dispersions =>

30. Chapter 230 Hydrogen Bonding Strong dipole-dipole attraction. Organic molecule must have N-H or O-H. The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule. O-H more polar than N-H, so stronger hydrogen bonding. =>

31. Chapter 231 H Bonds =>

32. Chapter 232 Boiling Points and Intermolecular Forces ethanol, b.p. = 78° Cethyl amine, b.p. = 17 ° C

33. Chapter 233 Solubility Like dissolves like. Polar solutes dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents. Molecules with similar intermolecular forces will mix freely. =>

34. Chapter 234 Ionic Solute with Polar Solvent Hydration releases energy. Entropy increases. =>

35. Chapter 235 Ionic Solute with Nonpolar Solvent =>

36. Chapter 236 Nonpolar Solute with Nonpolar Solvent =>

37. Chapter 237 Nonpolar Solute with Polar Solvent =>

38. Chapter 238 Classes of Compounds Classification based on functional group. Three broad classes  Hydrocarbons  Compounds containing oxygen  Compounds containing nitrogen. =>

39. Chapter 239 Hydrocarbons Alkane: single bonds, sp 3 carbons Cycloalkane: carbons form a ring Alkene: double bond, sp 2 carbons Cycloalkene: double bond in ring Alkyne: triple bond, sp carbons Aromatic: contains a benzene ring =>

40. Chapter 240 Compounds Containing Oxygen Alcohol: R-OH Ether: R-O-R ' Aldehyde: RCHO Ketone: RCOR ' =>

41. Chapter 241 Carboxylic Acids and Their Derivatives Carboxylic Acid: RCOOH Acid Chloride: RCOCl Ester: RCOOR ' Amide: RCONH 2 =>

42. Chapter 242 Compounds Containing Nitrogen Amines: RNH 2, RNHR ', or R 3 N Amides: RCONH 2, RCONHR, RCONR 2 Nitrile: RCN =>

43. Chapter 243 End of Chapter 2