html5-img
1 / 63

Structure and Properties of Organic Molecules

Highland Hall Biochem Block Handout #2. Structure and Properties of Organic Molecules. Adapted from Chapter 2 of Organic Chemistry by L. G. Wade, Jr. Molecular Orbitals (MO’s). Definition: interaction between orbitals on different atoms

edibiase
Télécharger la présentation

Structure and Properties of Organic Molecules

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Highland Hall Biochem Block Handout #2 Structure and Propertiesof Organic Molecules Adapted from Chapter 2 of Organic Chemistry by L. G. Wade, Jr.

  2. Molecular Orbitals (MO’s) • Definition: interaction between orbitals on different atoms • Generated by Linear Combination of Atomic Orbitals (LCAO) • Conservation of orbitals • # of MO’s created = # of AO’s combined • Interactions can be constructive or destructive. => Chapter 2

  3. Bonding MO : formed by constructive combination of AO’s with high e- density between the two nuclei. Results in MO with lower energy. Bonding MO (between two 1S orbitals) Chapter 2

  4. Antibonding MO: MO formed by destructive combination of AO’s with low e- density between the two nuclei Results in MOs higher in energy than the AO’s. Antibonding MO(between two 1S orbitals) Chapter 2

  5. Two atomic orbitals will always overlap so as to give 1 bonding MO and 1 antibonding MO. Example: H2 MO Energy Level Diagram => Chapter 2

  6. 1S 1S The σ bond • Sigma (σ) bond : single bond formed by constructive overlap of orbitals in which the e- density is centered about a line connecting the nuclei (i.e. cylindrically symmetrical) • Constructive orbital overlap → σ bonding MO • Destructive orbital overlap → σ* anti-bonding MO Chapter 2

  7. π Bonding Pi (π) bond : formed by constructive side-to-side overlap of parallel p orbitals. Chapter 2 7

  8. π Bonding • Not cylindrically symmetrical • Parallel overlap required for π bonds • π bonds cannot rotate with respect to each other • π bonds formed in multiple bonds • π bonds form after a σ bond. Chapter 2

  9. => 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. Chapter 2

  10. σ bonds vs π bonds Chapter 2

  11. Hybrid Orbitals • Hybrid orbitals : formed by mixing orbitals on the same atom. • Mixtures of s and p orbitals • Why do we need hybrids? • Bond angles cannot be explained with simple s and p orbitals, if so we would predict carbon to have 90o angles because 2px, 2py, and 2pz are orthogonal. • Actual angles are ~ 109o, 120o or 180o

  12. VSEPR • Valence Shell Electron Pair Repulsion (VSEPR) theory : e- pairs will repel each so as to attain a maximal distance of separation. • The observed bond angles are not possible using only s & p orbitals thus we use hybrids. • Hybridized orbitals are lower in energy because electron pairs are farther apart. • Hybridization is LCAO within one atom, just prior to bonding. Chapter 2

  13. Hybridization Rules 1) Both σ bonding e- and lone pairs occupy hybrid orbitals: # hybrid orbitals = (# σ bonds) + (# lone pairs) 2) The observed hybridization will be that which gives maximum separation of σ bonds and lone pairs 3) For multiple bonds, one bond is a σ bond using hybrid orbitals, while additional bonds are π bonds formed between p orbitals. Chapter 2

  14. sp Hybrid Orbitals(add an s and p orbital together) • 2 atomic orbitals = 2 hybrid orbitals • Linear electron pair geometry • 180° bond angle Chapter 2

  15. sp Hybrid Orbitals(add an s and p orbital together) Example: BeH2 Bond angle is 180o Chapter 2 15

  16. sp2 Hybrid Orbitals(add an s and 2 p orbitals together) • 3 atomic orbitals = 3 hybrid orbitals • Trigonal planar e- pair geometry • 120° bond angle Chapter 2

  17. sp2 Hybrid Orbitals(add an s and 2 p orbitals together) • Example: BeH3 • Bond angle is 120o Chapter 2 17

  18. sp3 Hybrid Orbitals(add an s and 3 p orbitals together) • 4 atomic orbitals = 4 hybrid orbitals • Tetrahedral e- pair geometry • 109.5° bond angle • Example : methane (CH4) Chapter 2

  19. Summary of s-p type orbitals

  20. 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! NH2NH2 CH3-CC-CHO Chapter 2

  21. Isomers • Isomers : molecules which have the same molecular formula, but differ in the arrangement of their atoms • Constitutional (or structural) : isomers which differ in their bonding sequence. • Stereoisomers : molecules that differ only in the arrangement of the atoms in space. Chapter 2

  22. Structural Isomers Chapter 2

  23. Conformations • Conformations : structures differing only in rotations about a single bond. • Example: ethane • Cylindrical symmetry of the σ bond allows rotation about a single bond “eclipsed” “staggered” side view side view

  24. Rotation around Bonds • σ bonds freely rotate. • π bonds cannot rotate unless the bond is broken. • only one conformation • leads to stereoisomers Chapter 2

  25. Stereoisomers Trans - across Cis - same side Stereoisomer : distinct compounds differing only in the spatial arrangement of atoms attached to the C=C double bond. Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp2 carbon. No cis-trans isomers possible Chapter 2 25

  26. Dipole Moment (μ) • Polar bond : bond in which e- are shared unequally. • Non-polar : bond in which e- are shared equally. • Dipole moment (μ) : measure of polarity. • are due to differences in electronegativity. • depend on the amount of charge and distance of separation. • In debyes (D),  = 4.8 x (electron charge) x d(angstroms)

  27. Bond Dipole Moments(μ)

  28. Molecular Dipole Moments • Depends on bond polarity and bond angles. • Vector sum of the bond dipole moments. • Lone pairs of electrons contribute to the dipole moment. μ = 1.9 D μ = 2.3 Dμ = 0 D Chapter 2

  29. More examples: Chapter 2 29

  30. Intermolecular Forces • Strength of attractions between molecules influence m.p., b.p., and solubility; especially for solids and liquids. • Classification depends on structure. • Dipole-dipole interactions • London dispersions • Hydrogen bonding => Chapter 2

  31. 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. Chapter 2

  32. Dipole-Dipole => Chapter 2

  33. London Dispersion • Induced between non-polar molecules • Temporary dipole-dipole interactions • Forces are larger for molecules with large surface area. • large molecules are more polarizable • Branching lowers b.p. because of decreased surface area contact between molecules. Chapter 2

  34. Dispersions => Chapter 2

  35. 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 there is stronger hydrogen bonding. Chapter 2 35

  36. H Bonds => Chapter 2

  37. Hydrogen Bonding & BP Higher bp’s for NH3 & H2O which can H-bond H2O bp is so high because… Chapter 2 37

  38. => Boiling Points and Intermolecular Forces Chapter 2

  39. Solubility • The solubility rule: “Like dissolves like” • Ionic & polar solutes dissolve in polar solvents. • Nonpolar solutes dissolve in nonpolar solvents. • Molecules with similar intermolecular forces will mix freely. Chapter 2

  40. Ionic Solute with Polar Solvent Hydration releases energy. Entropy increases. => Chapter 2

  41. Ionic Solute withNonpolar Solvent Attraction for solvent much weaker than ion-ion interactions Chapter 2

  42. Nonpolar Solute withNonpolar Solvent Small attractive forces for solvent are still strong enough to over come weak ion-ion intermolecular forces in solid. Chapter 2

  43. Nonpolar Solute with Polar Solvent Dissolving solute would require breaking up organized solvent structure. (crisco in water) Chapter 2

  44. Classes of Compounds • Classification based on functional group • Three broad classes • Hydrocarbons • Compounds containing oxygen • Compounds containing nitrogen => Chapter 2

  45. Hydrocarbons • Compounds containing only C and H. • Includes alkane, cycloalkane, alkene, cycloalkene, alkyne, and aromatic hydrocarbons • Typically hydrophobic Chapter 2

  46. Alkanes Hydrocarbons that contain only single bonds, sp3 carbons (CH3-CH2-CH3) Can be cyclic : then called cycloalkane Alkanes are very unreactive Alkane portion of molecule is called an alkyl group Chapter 2 47

  47. Alkenes Hydrocarbons that contain C=C double bonds, sp2 carbons. More reactive than alkanes Show geometric isomerism Cycloalkene: double bond in ring Chapter 2 48

  48. Alkynes Hydrocarbons with C—C triple bond, sp carbons More reactive than alkanes & alkenes Chapter 2 49

  49. Aromatic Hydrocarbons Hydrocarbons with alternating C—C single bond and C=C double bonds. Have unusual stability compared to alkenes and cycloalkenes Chapter 2 50

More Related