1. STRUCTURE Dr. Clower CHEM 2411 Spring 2014 McMurry (8 th ed.) sections 1.6-1.11, 2.10-2.11, 20.2, 2.4-2.6, 3.5-3.7, 4.3-4.9, 7.2, 7.6

2. Topics Structure Physical Properties Hybridization Resonance Acids and Bases Conformations of Alkanes and Cycloalkanes Unsaturation Alkene Stability

3. Structure Drawing organic structures Sigma (  ) and pi (  ) bonds Single bonds = 2e - = one sigma bond Double bonds = 4e - = one sigma bond and one pi bond Triple bonds = 6e - = one sigma bond and two pi bonds Which bond is shortest? Longest? Weakest? Strongest? Remember formal charges

4. Ionic Structures Be on the lookout for metals (cations) and ions Example: NaOCH 3 This is a Na + cation and a CH 3 O - anion Example: NH 4 Cl This is a NH 4 + cation and a Cl - anion

5. Classification of atoms C atoms can be classified as: Primary (1º) = C bonded to 1 other C Secondary (2º) = C bonded to 2 other C Tertiary (3º) = C bonded to 3 other C Quaternary (4º) = C bonded to 4 other C

6. Classification of atoms H atoms are classified based on the type of carbon to which they are attached

7. Classification of alcohols Alcohols are classified based on the type of carbon to which the -OH is bonded Classify these alcohols as 1º, 2º or 3º:

8. Classification of amines and amides Amines and amides are classified based on the number of C atoms bonded to the N

9. Classification of amines and amides Classify these functional groups:

10. Electronegativity and Bond Polarity Electronegativity Ability of atom to attract shared electrons (in a covalent bond) Most electronegative atom = F Differences in electronegativity determine bond polarity Bond polarity How electrons are shared between nuclei Equal sharing of electrons = nonpolar; unequal = polar

11. Bond Polarity Example: C─O What atom is more electronegative (C or O)? More EN atom has partial negative charge (  - ) Less EN atom has partial positive charge (  + ) Arrow shows direction of polarity Nonpolar bonds Any atom with itself C─H

12. Molecular Dipole Moment Overall electron distribution within a molecule Depends on bond polarity and bond angles Vector sum of the bond dipole moments (consider both magnitude and direction of individual bond dipole moments) Lone pairs of electrons contribute to the dipole moment Symmetrical molecules with polar bonds = nonpolar

13. Intermolecular Forces Strength of attractions between molecules Based on molecular polarity Influence physical properties (boiling point, solubility) 1. Dipole-dipole interactions 2. Hydrogen bonding 3. London dispersions (van der Waals)

14. 1. Dipole-Dipole Interactions Between polar molecules Positive end of one molecule aligns with negative end of another molecule Lower energy than repulsions Larger dipoles cause higher boiling points

15. 2. Hydrogen Bonding Strongest dipole-dipole attraction H-bonded molecules have higher boiling points 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

16. 3. London Dispersion Forces van der Waals forces Exist in all molecules Important with nonpolar compounds Temporary dipole-dipole interactions Molecules with more surface area have stronger dispersion forces and higher boiling points Larger molecules Unbranched molecules

17. Boiling Points and Intermolecular Forces CH 3 OCH 3 CH 3 CH 2 OH dimethyl ether, b.p. = -25°C ethanol, b.p. = 78°C NCH 3 H 3 C CH 3 CH 3 CH 2 CH 2 N H H NCH 3 CH 3 CH 2 H ethyl amine, b.p. 17°C trimethylamine, b.p. 3.5°C propylamine, b.p. 49°C ethylmethylamine, b.p. 37°C CH 3 CH 2 OH CH 3 CH 2 NH 2

18. Solubility and Intermolecular Forces Like dissolves like Polar solutes dissolve in polar solvents Nonpolar solutes dissolve in nonpolar solvents Molecules with similar intermolecular forces will mix freely

19. Example Which of the following from each pair will have the higher boiling point? (a)CH 3 CH 2 CH 2 CH 3 CH 3 CH 2 CH 2 OH (b)CH 3 CH 2 NHCH 3 CH 3 CH 2 CH 2 NH 2 (c)CH 3 CH 2 CH 2 CH 2 CH 3 CH 3 CH 2 CH(CH 3 ) 2

20. Example Will each of the following molecules be soluble in water? (a)CH 3 CO 2 H (b)CH 3 CH 2 CH 3 (c)CH 3 C(O)CH 3 (d)CH 2 =CHCH 3

21. Structure of Organic Molecules Previously: Atomic/electronic structure Lewis structures Bonding Now: How do atoms form covalent bonds? Which orbitals are involved? What are the shapes of organic molecules? How do bonding and shape affect properties?

22. Linear Combination of Atomic Orbitals Bonds are formed by the combination of atomic orbitals containing valence electrons (bonding electrons) Two theories: Molecular Orbital Theory Atomic orbitals of two atoms interact Bonding and antibonding MO’s formed Skip this stuff Valence Bond Theory (Hybridization) Atomic orbitals of the same atom interact Hybrid orbitals formed Bonds formed between hybrid orbitals

23. How many valence electrons? In which orbitals? So, both the 2s and 2p orbitals are used to form bonds How many bonds does carbon form? All four C-H bonds are the same i.e. there are not two types of bonds from the two different orbitals How do we explain this? Hybridization Let’s consider carbon…

24. Hybridization The s and p orbitals of the C atom combine with each other to form hybrid orbitals before they combine with orbitals of another atom to form a covalent bond Three types we will consider: sp 3 sp 2 sp

25. sp 3 hybridization 4 atomic orbitals → 4 equivalent hybrid orbitals s + p x + p y + p z → 4 sppp → 4 sp 3 Orbitals have two lobes (unsymmetrical) Orbitals arrange in space with larger lobes away from one another (tetrahedral shape) Each hybrid orbital holds 2e -

26. The sp 3 hybrid orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds Each C–H bond strength = 439 kJ/mol; length = 109 pm Each H–C–H bond angle is 109.5°, the tetrahedral angle. Formation of methane

27. Motivation for hybridization? Better orbital overlap with larger lobe of sp 3 hybrid orbital then with unhybridized p orbital Stronger bond Electron pairs farther apart in hybrid orbitals Lower energy

28. Another example: ethane C atoms bond by overlap of an sp 3 orbital from each C Three sp 3 orbitals on each C overlap with H 1s orbitals Form six C–H bonds All bond angles of ethane are tetrahedral

29. Both methane and ethane have only single bonds Sigma (  ) bonds Electron density centered between nuclei Most common type of bond Pi (  ) bonds Electron density above and below nuclei Associated with multiple bonds Overlap between two p orbitals C atoms are sp 2 or sp hybridized

30. Bond rotation Single (  bonds freely rotate Multiple (  bonds are rigid

31. sp 2 hybridization 4 atomic orbitals → 3 equivalent hybrid orbitals + 1 unhybridized p orbital s + p x + p y + p z → 3 spp + 1 p = 3 sp 2 + 1 p Shape = trigonal planar (bond angle = 120º) Remaining p orbital is perpendicular to hybrid orbitals

32. Formation of ethylene (C 2 H 4 ) Two sp 2 -hybridized orbitals overlap to form a C─C  bond Two sp 2 orbitals on each C overlap with H 1s orbitals (4 C ─ H) p orbitals overlap side-to-side to form a  bond  bond and  bond result in sharing four electrons (C=C) Shorter and stronger than single bond in ethane

33. 4 atomic orbitals → 2 equivalent hybrid orbitals + 2 unhybridized p orbitals s + p x + p y + p z → 2 sp + 2 p Shape = linear (bond angle = 180º) Remaining p orbitals are perpendicular on y-axis and z-axis sp hybridization

34. Formation of acetylene (C 2 H 2 ) Two sp-hybridized orbitals overlap to form a  bond One sp orbital on each C overlap with H 1s orbitals (2 C─H) p orbitals overlap side-to-side to form two  bonds  bond and two  bonds result in sharing six electrons (C≡C) Shorter and stronger than double bond in ethylene

35. Summary of Hybridization Hybridization of Csp 3 sp 2 sp ExampleMethane, ethaneEthyleneAcetylene # Groups bonded to C432 GeometryTetrahedralTrigonal planarLinear Bond angles~109.5~120~180 Types of bonds to C 44 3  2  C-C bond length (pm)154134120 C-C bond strength (kcal/mol)90174231

36. Hybridization of Heteroatoms Same theory Look at number of e - groups to determine hybridization Each lone pair will occupy a hybrid orbital Ammonia: N’s orbitals (sppp) hybridize to form four sp 3 orbitals One sp 3 orbital is occupied by the lone pair Three sp 3 orbitals form bonds to H H–N–H bond angle is 107.3° Water The oxygen atom is sp 3 -hybridized The H–O–H bond angle is 104.5°

37. Example Consider the structure of thalidomide and answer the following questions: a) What is the hybridization of each oxygen atom? b) What is the hybridization of each nitrogen atom? c) How many sp-hybridized carbons are in the molecule? d) How many sp 2 -hybridized carbons are in the molecule? e) How many sp 3 -hybridized carbons are in the molecule? f) How many  bonds are in the molecule?

38. Example Consider the structure of 1-butene: a) Predict each C─C─C bond angle in 1-butene. b) Which carbon-carbon bond is shortest? c) Draw an alkene that is a constitutional isomer of 1-butene.

39. Resonance Multiple Lewis structures for one molecule Differ only in arrangement of atoms Example: CH 2 NH 2 + ion These are resonance structures/forms Valid Lewis structures (obey Octet Rule, etc.) Same number of electrons in each structure Atoms do not move Differ only in arrangement of electrons (lone pair and  electrons)

40. Resonance Hybrid These structures imply that the C─N bond length and formal charges are different Actually not true; these structures are imaginary Molecule is actually one single structure that combines all resonance forms Resonance hybrid Contains characteristics of each resonance form More accurate and more stable than any single resonance form Lower energy (more stable) because of charge delocalization

41. Electron Movement Electrons move as pairs Can move from an atom to an adjacent bond, or from bonds to adjacent atoms or bonds Use curved arrows to show e - motion (electron pushing) Start where electrons are, end where electrons are going Connect resonance forms with resonance arrow This is not an equilibrium arrow

42. Contribution to Hybrid Structure Resonance forms do not necessarily contribute equally to the resonance hybrid They are not necessarily energetically equivalent More stable structures contribute more 1. Filled valence shells 2. More covalent bonds 3. Least separation of unlike charges (if applicable) 4. Negative charge on more EN atom (if applicable) Which of these is the major contributor to the resonance hybrid?

43. Benzene Resonance structures: Curved arrows? Is one structure more stable (contribute more)? Resonance hybrid: All carbon-carbon bonds are the same length Somewhere between C─C and C=C

44. Acetone Resonance structures: Curved arrows? Which structure is the major contributor? Which is the minor contributor? Are any structures not likely to form? Resonance hybrid:

45. Patterns in Resonance Structures

46. Examples

48. Acids and Bases Two types in organic chemistry 1. Brønsted-Lowry Acid = proton (H + ) donor Base = proton acceptor Some molecules can be both (e.g. water) = amphoteric Reaction will proceed from stronger acid/base to weaker acid/base Acid strength measured by pK a Stronger acid = lower pK a

49. Acids and Bases You can predict acid strength without a pK a value Strong acids have weak conjugate bases Weak conjugate bases are stable structures Have negative charge on EN atom (within a period) Have negative charge on a larger atom (within a group) Negative charge delocalized by resonance

50. Example Which is the stronger acid in each pair? a) H 2 O or NH 3 ? b) HBr or HCl? c) CH 3 OH or CH 3 CO 2 H?

51. Acids and Bases 2. Lewis Acid = electron pair acceptor, “electrophile” Base = electron pair donor, “nucleophile” Lewis acid react with Lewis base form a new covalent bond

52. Lewis Acids Incomplete octet (e.g. CR 3 +, BX 3 ), or Polar bond to H (e.g. HCl), or Carbon with d+ due to polar bond (e.g. CH 3 Cl)

53. Lewis Bases Nonbonded electron pair (anything with O, N, anions)

54. Lewis Bases If there is more than one possible reaction site (more than one atom with a lone pair), reaction occurs so that the more stable product is formed.