Components of Organic Chemistry Reactions: synthesis Organic compounds Structure bonding, conformation, analysis, stereochem. Reactivity interaction with other molecules: mechanism, dynamic stereochem. CareySundbergA-Chap1 1
그저익숙하도록읽는것뿐이다. 글을읽는사람이, 비록글의뜻은알았으나, 만약익숙하지못하면읽자마자곧잊어버리게되어, 마음에간직할수없을것은틀림없다. 이미읽고난뒤에, 또거기에자세하고익숙해질공부를더한뒤라야비로소마음에간직할수있으며, 또흐뭇한맛도있을것이다. - 퇴계이황 ( 금장태著 ) CareySundbergA-Chap1 2
Chapter 1. Bonding & Structure Valence Bond Theory: hybridization & resonance bonding electron pairs are localized between two atoms simple & effective but over-simplified & many exceptions Molecular Orbital Theory: Schrödinger equation bonding electrons are distributed over the entire molecule more accurate & close to the real structure but difficult to calculate & many assumptions in calculation CareySundbergA-Chap1 3
Valence Bond Theory (I) Lewis in 1916 & eitler-london in 1927 ( 2 : ) chemical bonding results from sharing of the electrons between the two atoms: localized bonding electron pairs ybridization: complex molecules (C 4 ) directed valence: Pauling in 1931; 4 bot. & 5 4 bonds rather than 2 & more effective overlap due to the highly directional 4 sp 3 orbitals: σ bonds 2 C=C 2 3 sp 2 & 1 2p or 4 sp 3, trigonal; C C 2 sp & 2 2p or 4 sp 3, digonal: 5 ( 2 C=C=C 2?) π bonds from 2 sp 3 : bent bonds CareySundbergA-Chap1 4
Valence Bond Theory (II) ybridization of strained molecules: 3-ring more p for C-C (17%, sp 5 )& more s for C- (33%, sp 2 ) bent bonds: less overlap; 7 Fig. 1.5 & 5 Fig. 1.4 electronegativity change: more electronegative (s character) C in strained molecules; 7 Fig. 1.6 propellanes: 8 Fig. 1.7; [4.4.4] sp 3 ; [2.2.2] sp 2 ; [1.1.1] inverted carbon reactive: 8 middle & bottom CareySundbergA-Chap1 5
Valence Bond Theory (III) Resonance theory: more than one Lewis structure possible for complex molecule Rule of thumb: 9 middle (a d); acrolein O + O O - C- II I C+ IV acrolein O + O - C - III V C + CareySundbergA-Chap1 6
Valence Bond Theory (IV) Application of resonance theory to acrolein a. Acrolein has properties of 5 resonance structures The real structure is their resonance hybrid b. All resonance structures satisfy the octet rule c. More stable resonance structures resemble more closely the real structure (major contributors): maximum No. of covalent bonds (I), minimum separation of unlike charges (I), negative charge on the more electronegative atom (IV/V) d. Usually, delocalization of electrons enhances stability relatively to a single localized structure: an energy barrier to rotation; 10 top (allyl cation) & 11 top (amide) CareySundbergA-Chap1 7
OEt + + OEt + OEt more stable C - - C vs O O - C - CareySundbergA-Chap1 8
Valence Bond Theory (V) Experimental proof for the resonance hybrid the weaker C=O bond: 1690 vs 1730 cm -1 (IR) the deshielded β-carbon in 1 & 13 C NMR: 115.9 136.2 18.7 136.4 137.1 CO 136.0 192.1 128.0 O 25,7 197.5 7.0 35.2 O 27.5 207.6 the chemical reactivity: 1,2-/1,4-addition of nucleophiles CareySundbergA-Chap1 9
Structural Properties of Chemical Bonds (I) Bond length: 13 Table 1.2 nearly constant for a hybridization type of carbon Bond energy: variable; 14 Table 1.3 (homolytic) eterolytic ΔE dis : more sensitive; 15 Table 1.4 eat of formation: isomers; 16 Table 1.5 stable: branched alkane & more substituted tans-alkene Polarity: different electronegativity; 17 Table 1.6 dipole moment = Σ bond dipoles; 17 Table 1.7 Mulliken electronegativity (χ) = (I + A) / 2; 18 CareySundbergA-Chap1 10
Structural Properties of Chemical Bonds (II) Polarity of hydrocarbons: 18 Scheme 1.1 electronegativity of C: sp > sp 2 > sp 3 ; C for normal sp 3 Polarity transmission: polar effect; reactivity inductive effect: successive polarization through bonds field effect: through-space interactions of the electric dipoles substituent effect on reactivity: 19 Table 1.8 Polarizability: size; 21-23 Table 1.9 & Fig. 1.8 response of electrons to nearby charges: SAB principle hardness: difficult distortion, softness: easy distortion δ δ + CareySundbergA-Chap1 11
Field Effect: Through-space Electrostatic Interaction Cl Cl Cl Cl CO 2 CO 2 pk a = 6.07 pk a = 5.67 CareySundbergA-Chap1 12
Molecular Orbital (MO) Theory: Overview (I) ψ = Eψ : accuracy vs amount of computation LCAO approximation: ψ = c 1 φ 1 + c 2 φ 2 + + c n φ n minimum basis set: combination of AO chosen 2s, 2p x, 2p y, 2p z for C, N, O and 1s for semiempirical calculations: use of experimentally determined parameters [ET, CNDO, MINDO-3, MNDO, AM1, MM2, PM3]; faster, simpler but limited applications ab initio calculations: absence of adjustable parameters (fewer assumptions) using SCF [STO-3G, 4-31G, 6-31G]; more reliable, accurate but more complex, time-consuming comparison between the two calculations: 29 Table 1.12 CareySundbergA-Chap1 13
MO Theory: Overview (II) Results obtainable from MO calculations the energy of each MO & charge distributions total electronic energy of the molecule relative to the atoms the calculated molecular energy relative stabilities of isomeric molecules conformational effects (the total energy as a function of molecular geometry) the minimum energy: the most favorable molecular structure the coefficients of the AOs contributing each MO applicable to the situation in the gas phase (on a single molecule) CareySundbergA-Chap1 14
MO Theory: Applications (I) Charge distribution of a molecule the electron density (q) at each atom r : C 3+ ; 27 Table 1.10 q r = n j c jr2, n=no. of e -, c j =coefficient at jth MO j 7 MOs from 3 1s & C 2s, 2p x, 2p y, 2p z 3 occupied MOs with 2 e - for each q C =3.565 e -, q =0.812 e - : total charge = (+0.435) [(4-3.565)] + 3(+0.188) [3(1-0.812)] = +1.000 LUMO: p z (localized purely on the carbon atom) cf.: 3 sp 2 & 1 p z from VB theory CareySundbergA-Chap1 15
MO Theory: Applications (II) eat of reaction: Δ f = Δ reactant - Δ product isodesmic reactions: test of reliability; 28-9 Table 1.11-12 the same No. of formal bonds of each type on each side Structure & energy: C 3 +, C 3, C 3 - ; 4-31G basis set Fig. 1.9, 29: C 3 + & C 3, planar; C 3 -, non-planar Substituent effects: π-donor vs inductive effects X-C + 2 vs -C + 2 : 29 middle & 30 Table 1.13 X-C - 2 vs -C2 - : 31 Table 1.14 vinyl: rotational barrier of an allyl cation & anion CareySundbergA-Chap1 16
ückel Molecular Orbital (MO) Theory (I) Simple but useful for the conjugated compounds assumption: the π-system can be treated independently of the σ framework in conjugated planar molecules (orthogonality) & mainly determines the chemical and spectroscopic properties Energy levels for each MO E=α+m j β, m j =2cos[jπ/(n+1)]: linear polyenes; 32 Table 1.15 α: Coulomb integral, a constant for all carbon atoms β: resonance integral, 0 for nuclei of nonbonding distance DE = E polyene - E ethylene = (6α+6.988β) - (6α+6β) = 0.988β, β 18 kcal/mol; [2(α+1.802β)+2(α+1.247β)+2(α+0.445β)] CareySundbergA-Chap1 17
ückel Molecular Orbital (MO) Theory (II) Coefficients of 2p AO of atom r for each MO: 32 bot. a node between different signs: antibonding; 34 Fig. 1.10 bonding between the similar size in concerted reactions ückel s rule: planar monocyclic conjugated polyenes aromatic: [4n+2] e - in the π system; antiaromatic: [4n] e - benzene-aromatic vs cyclobutadiene-antiaromatic E=α+m j β, m j =2cos(2jπ/n): 35 Fig. 1.11 (annulenes) DE benz = 2β vs DE cybu = 0; (aromatic) vs (antiaromatic) Frost s circle: a mnemonic; 35 Fig. 1.12 Charged C 3 3 & C 5 5 systems: 36 Fig. 1.13 CareySundbergA-Chap1 18
Calculation of Total Electron Energy E Cyclobutadiene: E cybu = 4α + 4β j=0, m 0 = 2cos(0/4) = 2; j =±1, m ±1 = 2cos(±2π/4) = 0; j=+2, m +2 = 2cos(4π/4) = -2; E 1 = α + 2β, E 2 = E 3 = α, E 4 = α -2β (empty) DE cybu = E cybu E ethylene = [2(α+2β)+1α+1α] (4α+4β) = 0 Benzene: E benz = 6α + 8β j=0, m 0 = 2; j =±1, m ±1 = 2cos(±2π/6) = 1; j=±2, m ±2 = 2cos(±4π/6) = -1; j=+3, m +3 = 2cos(+6π/6) = -2; E 1 = α + 2β, E 2 = E 3 = α + β, E 4 = E 5 = α - β (empty), E 6 = α -2β (empty) DE benz = E benz E ethylene = [2(α+2β)+4(α+β)] (6α+6β) = 2β CareySundbergA-Chap1 19
MO Energy Level of Acyclic Polyenes Allyl cation: E = 2α + 2.83β; anion: E = 4α + 2.83β Ψ 1 [E = α + 1.414β], Ψ 2 [E = α], Ψ 3 [E = α 1.414β] Butadiene: E = 4α + 4.47β Ψ [E = α + 1.618β], Ψ 1 2 [E = α + 0.618β] Ψ [E = α - 0.618β], Ψ 3 4 [E = α - 1.618β] Pentadienyl cation: E = 4α + 5.46β Pentadienyl anion: E = 6α + 5.46β Ψ 1 [E = α + 1.732β], Ψ 2 [E = α + 1.000β], Ψ 3 Ψ 4 [E = α - 1.000β], Ψ 5 [E = α 1.732β] [E = α] CareySundbergA-Chap1 20
Qualitative Application of MO Theory(I) Rules for construction of MO energy level diagram total No. of MOs = No. of AOs; aufbau principle the symmetry of MOs = the symmetry of the molecule symmetric: the same sign, antisymmetric: opposite sign orthogonal orbitals do not interact: p x, p y, p z the energy of more electronegative atoms is lower the more the No. of nodes, the higher the energy Diatomic molecules with 1s AO: 37 Fig. 1.14 2 +, 2, e 2 +, e 2, e + : 61, 103, 60, (0), 43 kcal/mol CareySundbergA-Chap1 21
Qualitative Application of MO Theory(II) Energy level diagram of CO: 38 Fig. 1.15 total No. of MOs = 10 (C & O: 1s, 2s, 2p x, 2p y, 2p z ; 14 e - ) the MOs 1s e - : negligible due to the large energy gap 10 e - in 5 MOs from 8 MOs: 4 bonding & 1 antibonding MO energy diagram of C 4 : 40 Fig. 1.18 ab initio calculation results: no 4 equivalent MOs qualitative analysis of energy diagram frame of reference: a cube; 37 Fig. 1.14 3 C 2 axes (x, y, z): symmetrical or antisymmetrical CareySundbergA-Chap1 22
Application of MO Theory to Reactivity (I) Perturbation MO (PMO) theory for new MOs change of the MO pattern with a change in structure the changes are small & the new MO pattern would be similar to known MOs of the similar system strongest interactions between MOs with close energy FMO theory: important interactions between OMO of one reactant & LUMO of the other; relative energy only MOs of matching symmetry can interact Reactivity difference between C 2 =C 2 & C 2 =O ethylene with E + & formaldehyde with Nu: 48 Fig. 1.25 CareySundbergA-Chap1 23
Reactivity Difference: C 2 =C 2 & C 2 =O LUMO π OMO π E + Nu: π LUMO π OMO 2 C C 2 2 C O CareySundbergA-Chap1 24
Application of MO Theory to Reactivity (II) Substituent effects on the reactivity of double bonds ethylene with a π-donor: reactive to E + ; 49 top figure allyl anion: reactive site at the terminal atoms; β-c & N ethylene with a π-acceptor: reactive to Nu; 49 middle butadiene: larger coefficient at the β-c of LUMO prediction with VB theory: resonance; 49 bottom Concerted reactions & symmetry: 53 bottom allyl cation & ethylene: forbidden; 51 Fig. 1.27 allyl cation & butadiene: allowed; 52 Fig. 1.28 CareySundbergA-Chap1 25
Application of MO Theory to Reactivity (III) yperconjugation of allylic systems interaction between σ & π bonds: σ bonds not on the nodal plane of π-system; 54 middle eclipsed conformation favored: 1.5-2.0 kcal/mol repulsive & attractive interactions: 55 top & Fig. 1.30 slightly longer C 3 - & C 1 =C 2 but shorter C 3 -C 2 prediction with VB theory: no-bond resonance; 55 Rotational barrier: 3 kcal/mol, ethane; 55 middle hindered rotation due to π character of some MOs: 56 Fig. 1.31 CareySundbergA-Chap1 26
Other Quantitative Methods Molecular graphs & critical points: 58 Fig. 1.32 Partial structures from MO quantitative calculations partition of total electron density among atoms applications: 59-62 Table 1.17 & Figs. 1.33-5 Density functional theory (DFT): the B3LYP method simpler calculation & rather accurate: 63 Tables 1.18-9 electronegativity: V = n / r : 61 top & 64 Table 1.20 n: No. of valence-shell e -, r : effective atomic radius hardness & V: acidity & stability; 64 & 22 Table 1.21 & 9 Quantitative VBT: localized MOs & delocalization CareySundbergA-Chap1 27