1. 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.

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

4. Antibonding MO (between two 1S orbitals)
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.
Chapter 2

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

6. 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
1S S
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 σ bond π bond 1) overlap head-to-head side-to-side
2) e- density
cylindrically symmetric about bond axis
maximum in regions about & below internuclear axis
3) # of bonds
only one can exist between 2 atoms
1 or 2 bonds between atoms
4) free rotation
yes no
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
Hybrid S
character P
Bond
Angle
Geometry
Bond Formed sp
1/2
180o
Linear σ
sp2
1/3
2/3
120o
Trigonal planar
sp3
1/4
3/4
109o
Tetrahedral

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
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.
Cis - same side
Trans - across
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 D μ = 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…
Compound
bpoC
CH4
-162
NH3
-33
H2O
+100
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 with Nonpolar Solvent
Attraction for solvent much weaker than ion-ion interactions
Chapter 2

42. Nonpolar Solute with Nonpolar 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. Solute Solvent Soluble?
Polar
Yes
Non-polar No

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

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

47. 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

48. 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

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

50. 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

51. Compounds Containing Oxygen
Alcohol: R-OH
Ether: R-O-R’
Aldehyde: RCHO
Ketone: RCOR'
Chapter 2

52. Alcohol Organic compounds that contian the hydroxyl group (-OH)
Common solvents
Have polar & non-polar parts
Chapter 2 52

53. Ethers
Organic compounds that containing 2 alkyl groups bonded to a bridging O atom
More polar than hydrocarbons because of the O atom.
Chapter 2 53

54. Aldehydes + Ketones Compounds containing the C=O carbonyl group
Somewhat polar because of high μ of C=O bond.
Aldehyde : has 1 alkyl group
and one H atom on each side
of the carbonyl (RCOH)
Ketone : has 2 alkyl groups (RCOR’)
R = alkyl groups
Chapter 2 54

55. Carboxylic Acids and Their Derivatives
Organic compounds containing the carboxyl group
Example: CH3COOH =
These are weak acids, often used in buffers
Chapter 2

56. Carboxylic Acids and Their Derivatives
Carboxylic Acid: RCOOH
Acid Chloride: RCOCl
Ester: RCOOR'
Amide: RCONH2
=>
Chapter 2 56

57. Carboxylic Acid Derivatives
Have the formula RCOL
All react with H2O to give RCOOH L
Compound
—OH
acid —X
X = halide
acid/halide
—OR
ester
— NH2
amide
Chapter 2 57

58. Nitrogen Compounds Amines: RNH2, RNHR', or R3N
Amides: RCONH2, RCONHR, RCONR2
Nitrile: RCN
Chapter 2 58

59. Amines Alkylated derivatives of ammonia (NH3)
General formulas RNH2, RNHR', or R3N
Polar because of C—N and N—H bond.
Example:
ethylamine = CH3CH2-NH2
1,5-pentanediamine

60. Amides General formulas : RCONH2, RCONHR, RCONR2
The amide C-N linkage in proteins is known as the pepide bond
Chapter 2 60

61. Wave Function Wave Equation: H  = E 
Solution gives  , the wave function
2 = amplitude of  at a given point in space
electron density at a given point in space
is a mathematical description of size, shape, orientation, can be + or - + _ -
=>
Chapter 2

62.  for a 1s orbital
Chapter 2

63.  for a 2p orbital Px, Py and Pz are energetically equivalent
Chapter 2