Organic Chemistry - 246A

1. Organic Chemistry - 246A Homework DUE Friday, 5 Sept Problems in McMurry 1.24; 1.28; 1.31; 1.45; 1.46; 1.47 => (1.48—1.52 BONUS Problems)

2. Summary of 1st Week’s Lectures • Carbon atoms can be oxidized or reduced (oxidation state) • Electronic structure of an atom described by wave equations (Y) • Electrons occupy orbitals around the nucleus • 1st row elements require 8 e— to achieve a “filled shell” octet • s orbitals are spherical (o) • p orbitals are dumbell-shaped (∞) • Covalent bonds => the bonding e— pair is shared between atoms • Sigma bonds (s) are circular in cross-section and are formed by head-on interaction of s, sp, sp2 or sp3 hybridized orbitals • Pi bonds (p) are “dumbell” in shape and are formed from side-side interaction of two p orbitals (88) • s and p orbitals can be combined to form multiple bonds • Double bonds use sp2 orbitals, have trigonal (120°) geometry (s + p) • Triple bonds use sp orbitals, have digonal (180°) geometry (s + 2p) • N and O have lone pairs (:) that can accept a proton

3. Molecular Orbital Theory (MOT) • Robert Mulliken in the 1940’s proposed that molecules have orbitals (MOs) that are the result of combinations of atomic orbitals (AOs) • A molecular orbital describes a region of space in a molecule where electrons are most likely to be found (e—s no longer associated with an individual atom) • AO combinations are additive and subtractive to form MOs • + combination (AO + AO = MO) forms MOs that is lower in energy (bonding) • — combination (AO — AO = MO) that is higher in energy (anti-bonding)

4. MO’s for Ethylene • Bonding and antibonding molecular orbitals result from combination of two patomic orbitals in ethylene • The bonding MO (+ combination) has no nodebetween nuclei • The antibonding MO (— combination) has a node between nuclei • Only the bonding MO is occupied (with 2 e—)

5. Greatest Least Electronegativity • Electronegativity (EN): ability of an atom to attract the shared e— in a covalent bond • F is most electronegative (EN = 4.0), Cs is least (EN = 0.7) • Differences in EN produce bond polarity

6. Dipoles and Polar Covalent Bonds • The greater the electronegativity difference between two bonded atoms, the more polarized the bond • Polarization gives rise to partial charges, and a dipole, m • If the electronegativity difference of the two atoms is greater than 2, the bond usually ionizes

7. Inductive Effect • For a polar covalent bond C—X, where X is an electronegative atom (e.g. Cl in CH3CH2Cl) the partial negative charge (d—) on the chlorine atom induces a partial positive charge (d+) on the neighboring carbon atom • The partial charges cause a dipole to be formed (m = Q x r) • Experimentally, it is easy to measure m, and if we know the bond length, r, then the partial charge Q can be calculated Chloroethane (Ethyl Chloride)

8. Dipoles Can Cancel Each Other • Even though each C—Cl bond has a strong dipole, the tetrahedral symmetry causes them to oppose each other • Resulting molecular dipole, m = 0

9. Charge Distribution Calculations

10. Resonance (VBT Concept) • Some molecules are have structures that cannot be shown with a single valence bond representation (e.g. carbonate, CO3=) • In these cases resonance contributors are drawn, which differ in the position of the  bond(s) and lone pair(s), and all contribute to the real structure of the molecule • The resonance forms are connected by a double-headed arrow <—> • In this case, we say that the negative charges on CO3= are delocalized

11. Resonance Hybrids • A structure with resonance forms does not alternate between the forms • Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid • Benzene (C6H6) has two resonance forms with alternating double and single bonds, yet all the bonds are of equal length- intermediate in length between single and double bond length

12. Rules for Resonance • Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing all the contributors can you visualize the actual structure) • Resonance forms differ only in the placement of their  e— and/or nonbonding e— • Different resonance forms of a substance don’t have to be equivalent (e.g. amide H2N—C=O <—> +H2N=C—O—), and can contribute to different degrees • Resonance forms must be valid Lewis structures: the octet rule has to be satisfied • The resonance hybrid is more stable than any individual resonance form (resonance stablization)

13. Curved Arrows (Sir Robert Robinson) • We can imagine that electrons move in pairs to convert from one resonance form to another • Of course electrons are delocalized and this is only a bookkeeping device

14. Delocalization • Radicals (•), anions (-) and cations (+) can also be delocalized • Reactions can take place at either end of the intermediate

15. Acid-Base Theories • The Arrhenius concept that acids are solutions containing “H+” and bases are solutions containing “OH—” is not very useful in organic chemistry • Brønsted–Lowry theory defines acids as H+ donors, and bases as H+ acceptors • Lewis theory defines acids as lone pair acceptors, and bases as lone pair donors

16. Acetone Anion Displays Resonance • Acetone can give up a proton to a strong base to form an anion • Acetone anion can react at O or at C

17. Acetyl Acetone Is More Acidic than Acetone • 2,4-Pentanedione (Acetyl Acetone, or AcAc) is deprotonated even by very weak bases • It has 3 resonance forms

18. pKa Deprotonated Protonated