Review

source materials: Carey, "Advanced Organic Chemistry;" March, 3rd Ed., Acids and Bases, Elimination Reactions, SN1 and SN2 reactions, Addition Reactions

Class Notes

  1. Fundamental concepts

    1. Acid-base theories

      1. Arrhenius

      2. Bronsted-Lowry

      3. Lewis

    2. Lewis acids/bases in organic chemistry

      1. Nucleophile - shares an electron pair with another atom to form a new covalent bond: a Lewis base

      2. Electrophile - gains an electron pair from another atom to form a new covalent bond: a Lewis acid

      3. Factors that affect the Lewis basicity of substances (a. - d. in order of strength of effect; e. and f. of variable influence)

        1. Resonance: can increase or decrease electron density at the atom that behaves as a base

        2. Polarizability: a measure of the ease of distortion of an atom/molecule's electron cloud

          1. hard: high electronegativity, small size = electrons held tightly

          2. soft: low electronegativity, large size = electrons held loosely

          3. hard bases donate more easily to hard acids

          4. soft bases donate more easily to soft acids

        3. Electronegativity: the tendency of an atom to attract electrons to itself

        4. Inductive effect

          1. Electrons in C-C bonds are more polarizable than electrons in C-H bonds

          2. Lone charges do not exist in the real world, require stabilization, either through solvent effects or through stabilization from within the molecule

          3. Alkyl groups are electron-donating; replacing H atoms around a carbocation with alkyl groups results in its stabilization

          4. The electron donating or withdrawing effect of a group transmitted through C-C sigma bonds is called the inductive effect

        5. Charge: everything being equal, a base with a negative charge is a better nucleophile than when uncharged - electron cloud held less tightly, more polarizable

        6. Hybridization: the greater the degree of s character, the closer the bonding electrons are held; i.e., increasing s character increases the apparent electronegativity of an atom

    3. Hybridization and hybrid orbitals

      1. Orbital shapes - s, p, d

      2. Orbital overlap is essential for chemical bonding to occur; the greater the overlap the stronger the bond

        1. Show overlap between hydrogen and oxygen in water

        2. Overlap is an important factor in bond strength but certainly not the only factor

      3. If carbon has four valence electrons (2s2 2p2), how can it form four bonds?

        1. In the 1930s Linus Pauling suggested that one of the 2s electrons is promoted to the empty 2p orbital, resulting in the formation of four hybrid orbitals that are intermediate in energy and shape

        2. Diagram of orbitals vs. energy

        3. This hybridization theory can be easily and accurately used to describe bonding in many molecules

        4. atom
          molecule
          valence configuration
          hybridization
          remaining orbitals
          Be
          BeH2
          2s2
          sp
          2 p orbitals
          B
          BF3
          2s2 2p1
          sp2
          1 p orbitals
          C
          CH4
          2s2 2p2
          sp3
          none
          P
          PCl5
          3s2 3p3
          dsp3
          4 d orbitals
          S
          SF6
          3s2 3p4
          d2sp3
          3 d orbitals

    4. Bonding in carbon

      1. Carbon is tetravalent

      2. Tetrahedral geometry

      3. Covalent bonds: shared electrons

      4. The extent to which the bonding electrons are shared is a function of dEN of the two bonding atoms; bond polarity

        1. regions of a molecule that are partially positively charged are susceptible to attack from electron-rich reactants

        2. regions of a molecule that are partially negatively charged are susceptible to attack from electron-poor reactants

      5. Multiple bonds

        1. Double bonds - consist of one sigma bond and one pi bond

          1. Sigma bonds

            1. Cylindrical probability distribution around the bond axis

            2. Sigma bonds are created by the overlap of hybridized orbitals

          2. Pi bonds

            1. Probability out of the plane (above and below) the bond axis

            2. Pi bonds are created by the overlap of the unhybridized p orbitals

        2. Triple bonds consist of one sigma bond and two pi bonds

        3. Pi bonds are electron sources, behave as nucleophiles in acid-base reactions

        4. Bond strength (Carey: 343, Table 9.1)
          ethane
          368 kJ/mol
          		  
          ethylene
          611 kJ/mol
          243 kJ/mol (66%)
          acetylene
          820 kJ/mol
          209 kJ/mol (57%)

    5. Lewis structures and resonance: "success in organic chemistry depends on writing correct Lewis structures"

      1. Resonance: when a molecule can be represented by two or more possible Lewis structures in which the positions of the atoms are constant and which differ only in the distribution of bonding and nonbonding electrons, the actual molecule is probably best represented by a weighted average of all of the possible Lewis structures
        (39% x 2, 7.3% x 3)

      2. Each contributing Lewis structure must have the same total number of electrons, the same net charge, and the same number of nbp

      3. Each structure contributes in proportion to its stability

        1. Rules (see March: 33, Carey: 24f, Table 1.5): deal with numbers of covalent bonds, charge separation, and formal charges

      4. Single Lewis structures often incorrectly depict electrons as being localized; resonance structures correctly depict electron delocalization

      5. Electron delocalization results in stabilization of the molecule

      6. A molecule in which there are delocalized electrons is more stable than can be implied by any individual Lewis structure

      7. The degree of stabilization is greatest when contributing Lewis structures are of equal stability

      8. The difference in energy between the actual energy of the molecule and the energy of the lowest energy Lewis structure is its resonance energy

    6. Hyperconjugation

      1. Caused by overlap of C-H sigma bond orbital with empty p orbital in sp2 hybridized carbocation

      2. The resulting delocalization results in stabilization of the carbocation

    7. Other structural features

      1. Steric strain (total strain): the sum of the steric hindrance, torsional strain, angle strain, and bond length distortion

      2. Steric hindrance (van der Waal strain): hindrance caused by physical space constraints

      3. Torsion strain: the destabilization that results from eclipsed bonds on adjacent angles caused by electron pair repulsion of the bonding electrons

      4. Angle strain: strain resulting from deviation from ideal bond angles; can result in diminished orbital overlap and weakened bonds

      5. Bond length distortion: destabilization that results when a bond length deviates from its normal value

  2. Elimination reactions (Carey: 181-198, 348-9)

    1. General

      1. Involve elimination of portions of molecule on adjacent carbon atoms with the resulting formation of a new pi bond

      2. Order of reactivity 3° > 2° > 1° reflects the relative stabilities of both the transition states and the resulting alkenes

      3. Factors affecting alkene stability

        1. Degree of substitution: tetrasubstituted > trisubstituted > disubstituted > monosubstituted > unsubstituted

        2. sp2 hybridized carbons are electron-attracting; electron-releasing groups - such as alkyl groups - tend to stabilize alkenes

        3. Steric strain: trans-disubstitued are generally more stable than cis-disubstituted

      4. Eliminations are regioselective (preferentially occur in one direction) and follow the Zaitsev Rule: 1,2-elimination reactions yield the most highly substituted alkene as the major reaction product (true for both E1 and E2 reactions)

    2. E1 reactions - elimination unimolecular - typified by the two-step acid-catalyzed dehydration of alcohols

      2. First order kinetics: rate = k[alkyl halide]

      3. Can be observed in alkyl halides, but typically only observed in 3° and some 2° alkyl halides and when the base is weak or [base] is low

      4. Rearrangements can occur during E1 eliminations with the resulting formation of a more stable carbocation

    3. E2 reactions - elimination bimolecular - typified by single-step base-catalyzed dehydrohalogenation of alkyl halides

      2. Second order kinetics: rate = k[base][alkyl halide]

      3. Partial double-bond develops in transition state

      4. The rate of elimination depends on the halogen and the strength of the C-X bond

        1. Order of bond strength: R-I < R-Br < R-Cl < R-F

        2. Order of reactivity: R-I > R-Br > R-Cl > R-F

      5. Three fundamental requirements for the E2 reaction

        1. A good leaving group

        2. A strong base

        3. A hydrogen atom on a carbon atom adjacent to the carbon atom with the leaving group

  3. Substitution reactions (Carey Ch. 8)

    1. General

      1. Alkyl halides can experience replacement (substitution) of the halogen atom

      2. Substitutions occur under somewhat similar circumstances to those of elimination reactions

      3. Halogen atom leaves as an anion; the pair of electrons in the bond are abstracted by the leaving halogen atom (a function of greater EN than the carbon atom to which the halogen is bonded)

      4. In general substitutions only affect halogens bonded to sp3 hybridized carbon atoms

      5. Leaving groups

        1. Order of reactivity: R-I > R-Br > R-Cl > R-F

        2. As with elimination reactions, reactivity is a function of the strength of the C-X bond

      6. Nucleophiles and nucleophilicity

        1. Nucleophilicity - a measure of how fast a Lewis base displaces a leaving group compared to the reaction of methyl iodide in methanol

        2. May be either anions or electrically neutral

        3. Neutral Lewis bases - water, alcohols, carboxylic acids - are generally weak nucleophiles and weaker Lewis bases than their conjugate bases (which are usually anions)

        4. The more basic the nucleophile, the more reactive it is (remember: the stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger its conjugate base)

        5. Basicity increases as one moves from left to right across a period (i.e., as EN decreases)

        6. Solvent-nucleophile bonds impact nucleophilicity; the stronger the nucleophile-solvent bonds, the harder it is for the nucleophile to "escape" the solvent

        7. Polarizability also favorably affects nucleophilicity as it increases

      7. Solvent effects

        1. SN1: polar solvents (high dielectric constant) are better able to stabilize the ions formed during the reaction

        2. SN2: polar aprotic solvents tend to make it easier for the nucleophile to "escape" the solvent (no "hydrogen bonds")

    2. SN1 - substitution nucleophilic unimolecular - an ionization mechanism
      1. SN1 mechanism

      2. First order rate kinetics: rate = k[alkyl halide]

      3. For SN1 reactions, the order of reactivity: 3° > 2° > 1° > methyl due to stabilization of the carbocation

      4. Rearrangements can occur during SN1 substitutions with the resulting formation of a more stable carbocation

    3. SN2 - substitution nucleophilic bimolecular - the "direct displacement" mechanism

      1. SN2 mechanism

      2. Second order kinetics: rate = k[base][alkyl halide]

      3. Inversion of configuration due to steric considerations during the reaction; the base must approach the alkyl halide from the side opposite to the bond with the leaving group

      4. For SN2 reactions, the order of reactivity: 3° < 2° < 1° < methyl due to steric hindrance of the nucleophile attack

      5. Steric hindrance can also occur when the adjacent carbon is 3°

    4. Substitution vs. elimination

      1. ". . . . the characteristic reaction of alkyl halides with Lewis bases is elimination, and. . . . substitution predominates only under certain special circumstances." (Carey: 323)

      2. The two most important factors are the structure of the alkyl halide and the basicity of the nucleophile

      3. If the carbon is not crowded (steric effects) and not stabilized by substituents (e.g. primary alkyl halides), SN2 will occur in preference to E2

      4. If the base is a weaker base than OH- it may react with primary and secondary alkyl halides to give substitution products

      5. Tertiary alkyl halides only undergo substitutions in the absence of anionic Lewis bases i.e., when neutral solvent molecules behave as Lewis bases

      6. As temperature increases, the likelihood of elimination increases

      7. SN2 reactions can be planned around sterically unhindered alkyl halides, weakly basic nucleophiles, and low temperatures

      8. SN1 reactions generally only occur when E1 reactions are impossible

  4. Addition reactions

    1. General

      1. A reaction in which two molecules combine to yield a single product molecule; the reagent is simply added to the substrate molecule; resulting in the breaking of a pi bond and the formation of two sigma bonds

      2. Pi electrons are not held as tightly, serve as electron source

      3. Alkenes/alkynes can be attacked by electrophiles or free radicals

    2. Electrophilic addition

      2. Step 2 is very much like second step in Sn1 reaction

      3. Regioselectivity of hydrogen halide addition is described by Markonikov's Rule: when an unsymmetrically substituted alkene reacts with a hydrogen halide the hydrogen adds to the carbon with the greatest number of hydrogens, and the halogen atom adds to the more highly substituted carbon

      4. Since a carbocation is formed, rearrangements can take place

    3. Free radical additions cannot be described in terms of Lewis acid-base chemistry

  5. Summary

    1. How can many organic reactions be described in terms of Lewis acid-base chemistry?

    2. What factors exert the greatest influence on the rate of a reaction?