Structure

1. Structure Dr. Clower CHEM 2411 Spring 2014 McMurry (8th 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 (s) and pi (p) 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: NaOCH3 • This is a Na+cation and a CH3O- anion • Example: NH4Cl • This is a NH4+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 (d-) • Less EN atom has partial positive charge (d+) • 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) • Dipole-dipole interactions • Hydrogen bonding • 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. C H O C H C H C H O H 3 3 3 2 ethanol, b.p. = 78°C dimethyl ether, b.p. = -25°C C H C H C H N H C H C H N C H H C N C H 3 2 2 3 2 3 3 3 H H C H 3 propylamine, b.p. 49°C ethylmethylamine, b.p. 37°C trimethylamine, b.p. 3.5°C C H C H N H C H C H O H 3 2 2 3 2 ethanol, b.p. = 78°C ethyl amine, b.p. 17°C Boiling Points and Intermolecular Forces

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) CH3CH2CH2CH3 CH3CH2CH2OH (b) CH3CH2NHCH3 CH3CH2CH2NH2 (c) CH3CH2CH2CH2CH3 CH3CH2CH(CH3)2

20. Example • Will each of the following molecules be soluble in water? (a) CH3CO2H (b) CH3CH2CH3 (c) CH3C(O)CH3 (d) CH2=CHCH3

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. Let’s consider carbon… • 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

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: • sp3 • sp2 • sp

25. sp3 hybridization • 4 atomic orbitals → 4 equivalent hybrid orbitals • s+ px + py + pz→ 4 sppp→ 4 sp3 • Orbitals have two lobes (unsymmetrical) • Orbitals arrange in space with larger lobes away from one another (tetrahedral shape) • Each hybrid orbital holds 2e-

26. Formation of methane • The sp3 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.

27. Motivation for hybridization? • Better orbital overlap with larger lobe of sp3 hybrid orbital then with unhybridizedp orbital • Stronger bond • Electron pairs farther apart in hybrid orbitals • Lower energy

28. Another example: ethane • C atoms bond by overlap of an sp3 orbital from each C • Three sp3 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 (s) bonds • Electron density centered between nuclei • Most common type of bond • Pi (p) bonds • Electron density above and below nuclei • Associated with multiple bonds • Overlap between two p orbitals • C atoms are sp2 or sp hybridized

30. Bond rotation • Single (s) bonds freely rotate • Multiple (p) bonds are rigid

31. sp2 hybridization • 4 atomic orbitals → 3 equivalent hybrid orbitals + 1 unhybridizedp orbital • s + px + py + pz→ 3 spp + 1 p = 3 sp2 + 1 p • Shape = trigonal planar (bond angle = 120º) • Remaining p orbital is perpendicular to hybrid orbitals

32. Formation of ethylene (C2H4) • Two sp2-hybridized orbitals overlap to form a C─C s bond • Two sp2 orbitals on each C overlap with H 1s orbitals (4 C ─ H) • porbitals overlap side-to-side to form a  bond • sbond and bond result in sharing four electrons (C=C) • Shorter and stronger than single bond in ethane

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

34. Formation of acetylene (C2H2) • Two sp-hybridized orbitals overlap to form a s bond • One sp orbital on each C overlap with H 1s orbitals (2 C─H) • porbitals overlap side-to-side to form two  bonds • sbond and two bonds result in sharing six electrons (C≡C) • Shorter and stronger than double bond in ethylene

35. Summary of Hybridization

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 sp3 orbitals • One sp3 orbital is occupied by the lone pair • Three sp3 orbitals form bonds to H • H–N–H bond angle is 107.3° • Water • The oxygen atom is sp3-hybridized • The H–O–H bond angle is 104.5°

37. Example • Consider the structure of thalidomide and answer the following questions: • What is the hybridization of each oxygen atom? • What is the hybridization of each nitrogen atom? • How many sp-hybridized carbons are in the molecule? • How many sp2-hybridized carbons are in the molecule? • How many sp3-hybridized carbons are in the molecule? • How many p bonds are in the molecule?

38. Example • Consider the structure of 1-butene: • Predict each C─C─C bond angle in 1-butene. • Which carbon-carbon bond is shortest? • 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: CH2NH2+ 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 p electrons)

40. ResonanceHybrid • 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 • Filled valence shells • More covalent bonds • Least separation of unlike charges (if applicable) • 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

47. Examples

48. Acids and Bases • Two types in organic chemistry • 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 pKa • Stronger acid = lower pKa

49. Acids and Bases • You can predict acid strength without a pKa 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? • H2O or NH3? • HBr or HCl? • CH3OH or CH3CO2H?