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No 1 (2016.5.06) 有機反応化学 ( Advanced Organic Reaction)(H28年度前期�古田) Text: Molecular Orbitals and Organic Chemical Reactions Ian Fleming (Wiley, 2009) 1. Chap.1: Molecular Orbital Theory (5/6) 2. Chap.2: Molecular Orbitals and the Structures of Organic Molecules (5/13) [Question] Hydroperoxide ion P.9 is much less basic than hydroxide ion P.10. Why, then, is it so much more nucleophilic? 3. Chap.3: Chemical Reactions (5/20) 4. Chap.4: Ionic Reaction- Reactivity (5/27) 5. Chap.5: Ionic Reaction- Stereochemistry (6/03, 10) 6. Chap.6: Thermal Pericyclic Reactions (6/24) 7. Chap.7: Radical Reactions (7/15) 8. Chap.8: Photochemical Reactions (7/22) 6/17(集中講義)、7/01(休講)、08(休講)、補講日程未定 参考書: 大学院講義�有機化学I, II (東京化学同人) �マーチ有機化学 Modern Physical Organic Chemistry (E.V.Anslyn & D. A. Dougherty) 有機反応機構の書き方(奥山�格:丸善) �������演習で学ぶ有機反応機構(化学同人) [Question] Why does diazomethane P.15 add to methyl acrylate P.16 to give the isomer P.17 in which the nitrogen end of the dipole is bonded to the carbon atom bearing the methoxycarbonyl group, and not the other way round P.14? 1. Molecular Orbital Theory 1.1 The Atomic Orbitals of a Hydrogen Atom φ: wave function φ2dτ: the probability of finding the electron in the volume dτ ∫φ2dτ = 1 E1s = -13.60 eV 1.4 Å ~ 90 % , 2 Å ~ 99%. Pictureofelectrondistribution The van der Waals radius at 1.2 A has no theoretical significance-it is an empirical measurement from solid state structures, being one-half of the distance apart of the hydrogen atom in a C-H bond and the hydrogen atom in the C-H bond of an adjacent molecule. Bohr radius 0.529 Å H N P O S F Cl Br I (Å) 1.2 1.5 1.9 1.40 1.85 1.35 1.80 1.95 2.15 1.2 Molecules Made from Hydrogen Atoms 1.2.1 The H2 Molecule σ = c1φ1 + c2φ2 1.1 σ: the molecular orbital φ1: the atomic 1s wave function on atom 1 φ2: the atomic 1s wave function on atom 2 c1, c2 : a measure of the contribution which the atomic orbital is making to the molecular orbital σ2 = (c1φ1 + c2φ2)2 = (c1φ1)2 + (c2φ2)2 + 2c1φ1c2φ2 1.2 c1c2 > 0: the electron population between the two atoms is increased c1c2 < 0: a low electron population in the space between the nuclei, By making a bond between two hydrogen atoms, we create two new orbitals, σ and σ*, which we call the molecular orbitals; the former is bonding and the latter antibonding. A lowering of energy is equal to twice the value of Eσ. antibonding bonding S12 = ∫φ1φ2dτ S12: overlap integral 1.3 Energy of an electron in a bonding molecular orbital Bonding (σ) and Antibonding (σ*) Orbitals Energy of an electron in an antibonding molecular orbital α: Coulomb integral β: resonance integral H: Hamiltonian Coeficient (C1, C2) In each orbital the sum of the squares of all the c–values must equal one, since only one electron in each spin state can be in the orbital. Since |c1 | must equal |c2 | in a homonuclear diatomic like H2, the values of c1 and c2 in the bonding orbital must be, namely 1/√2 = 0.707: If all molecular orbitals were filled, then there would have to be one electron in each spin state on each atom, and the sum of the squares of all the c values on any one atom in all the molecular orbitals must also equal one. Thus the σ*-antibonding orbital of hydrogen will have c-values of 0.707 and –0.707, because these values make the whole set fit both criteria. 1.2.2 The H3 Molecule Symmetry rule: combine the orbitals that have the same symmetry Triangle H3 . H3+,H3 stable H3 unstable Straight line H3 1.2.3 The H4 'Molecule' (tetrahedral) Symmetry-(xz, yz) less bonding the same energy as 1s less bonding Two H2 molecules do not combine to form an H4 molecule. 1.3 C-H and C-C Bonds C: (1s)2(2s)2(2p)2 E2p = -10.7 eV 1.3. 1 The Atomic Orbitals of a Carbon Atom C: (1s)2(2s)2(2p)2 E2s = -19.50 eV (vs. E1s = -13.60 eV for H) node 0.9 Å 1.5 Å Effective overlap Atomic Orbital Energies 1.3.2 Methane (CH4) Atomic Orbitals H4 + C Symmetry-(xz, yz) Molecular Orbitals of CH4 1.3.3 Methylene (:CH2) 3 orbital interaction H2 + C Symmetry xz yz two-fold z (H-C-H) Linear combination of 3 atomic orbitals Ic1I=Ic3I bent form (more stable) SSS SSS SSS 1.3.4 Hybridisation Because of the presence of the inner sphere in the 2s orbital, the nucleus is actually inside the back lobe, and a small proportion of the front lobe reaches behind the nucleus (Review) CH4 1.3.5 C-Cσ Bonds and π Bonds: Ethane (C2H6) Energy level and electron distribution sp3 hybridization C:8e H:6e Since C-C single bonds are typically about 1.54 Å long, the overlap integral at this distance for π bonding is a little less than half that for bonding. π bonds are therefore much weaker. The C-C bond would have been seen as the bonding overlap of sp3 hybridised orbitals on carbon with each other which would have used different proportions of s and p orbitals, and would have been labelled sp3. 1.3.6 Ethylene: C=C π-bond β = 140 kJ/mol (= 1.45 eV = 33 kcal/mol) => 2β = 280 kJ/mol (π bond energy) β Electron in the Box 1.4 Conjugation-Hückel Theory Allyl System resonance σ-framework Orbital Energy where k is the number of the atom along the sequence of n atoms Place dummy atoms at both ends From 3 orbitals to 4 orbitals? NBMO 2 x 0.414β (116 kJ/mol) We cannot create four orbitals from three, because we cannot use the p orbital separately twice. 1.4.2 Butadiene Frontier Orbitals HOMO: the highest occupied molecular orbital LUMO: the lowest unoccupied molecular orbital SOMO: the singly occupied molecular orbitals allyl cation LUMO:ψ2 HOMO:ψ1 LUMO allyl anion allyl radical LUMO:ψ3* HOMO:ψ2 SOMO:ψ2 HOMO SOMO giving the coefficient Cjr for atom j in molecular orbital r of a conjugated system of n atoms (so that j and r = 1, 2, 3, • • ., n) Conjugated systems are often, but not always, lower in energy than unconjugated systems. 1.4.3 Longer Conjugated System HOMO LUMO 1.5 Aromaticity 6 π orbitals of benzene 1.5.1 Aromatic Systems Ψ1~Ψ6 Hückel aromaticity: (4n+2) π-electrons ------ special stability n=1 n=0 n=3 n=4 Where does this special stability come from? C6H6 hexatriene Total π bonding energy of benzene 2 x 4β = 8β Frost circle hexatriene 2 x (1.802 + 1.247 + 0.445)β = 7β From Hexatriene to Benzene 1.7 Heteronuclear Bonds, C-M, C-X and C=O Electronegativity and Atomic Orbital Energies Pauling s χ χA − χB = Δ 23.06 Δ = D(A − B) − D(A − A)D(B − B) D:bonddissociationenergy € down € up down Ex.NMRmethod 1.7.1 Atomic Orbital Energies and Electronegativity Inamotoetal,Chem.Lett.1982,1003. 3D Electronegativity Scale C-H is not strongly polarised. 2.54 2.30 L. C. Allen JACS 1989, 111, 9003 Valence Atomic Orbital Energies (eV) Atomic Orbital Energies for Hybrid Orbitals (eV) 1.7.2 C-X σ Bonds H3C-Cl C-Cl σ bond Weakerbond? C (-10.7eV) (-13.7eV) D(C-C):347kJ/mol D(C-Cl):352kJ/mol Na-F Na+ F- When two orbitals of unequal energy interact, the lowering in energy is less than when two orbitals of very similar energy interact. Ion pair Orbital overlap is negligible When orbitals of identical energy interact, the energy lowering is roughly proportional to S. When they are significantly different in energy, it is roughly proportional to S2. They are inversely proportional to the energy difference Ei. (-5.2eV) (-18.6eV) 1.7.3 C-M σ Bonds C-Cl C-F C-Li C-Li (-5.4eV) (-10.7eV) C-Li CH3Li HOMO 1.7.4 C=O π Bonds HCHO (-10.7eV) (-15.9eV) LUMO Report No1 (5/06) Back-side attack of nucleophile 1. In the discussion about the H3 molecule, we could have combined the three atoms in a straight line. Show that there would be less bonding both in σ1 and σ2* and less antibonding in the σ* orbital. 2. Calculate the coefficients for the frontier orbitals of the heptatrienyl cation. hint: use equation or its geometrical equivalent