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Physical Organic Chemistry PHYSICAL ORGANIC CHEMISTRY Yu-Tai Tao (陶雨台） Tel: (02)27898580 E-mail: [email protected] Website:http://www.sinica.edu.tw/~ytt Textbook: “Perspective on Structure and Mechanism in Organic Chemistry” by F. A. Corroll, 1998, Brooks/Cole Publishing Company References: 1. “Modern Physical Organic Chemistry” by E. V. Anslyn and D. A. Dougherty, 2005, University Science Books. Grading: One midterm (45%) one final exam (45%) and 4 quizzes (10%) homeworks Chap.1 Review of Concepts in Organic Chemistry § Quantum number and atomic orbitals Atomic orbital wavefunctions are associated with four quantum numbers: principle q. n. (n=1,2,3), azimuthal q.n. (m= 0,1,2,3 or s,p,d,f,..magnetic q. n. (for p, -1, 0, 1; for d, -2, -1, 0, 1, 2. electron spin q. n. =1/2, -1/2. § Molecular dimensions Atomic radius ionic radius, ri：size of electron cloud around an ion. covalent radius, rc：half of the distance between two atoms of same element bond to each other. van der Waal radius, rvdw：the effective size of atomic cloud around a covalently bonded atoms. - Cl Cl2 CH3Cl Bond length measures the distance between nucleus (or the local centers of electron density). Bond angle measures the angle between lines connecting different nucleus. Molecular volume and surface area can be the sum of atomic volume (or group volume) and surface area. Principle of additivity (group increment) Physical basis of additivity law: the forces between atoms in the same molecule or different molecules are very “short range”. Theoretical determination of molecular size：depending on the boundary condition. Boundary is a certain minimum value of electron density. Molecular volume (1 au = 6.748e/Å3 ), 0.001au 0.02au expt’l CH4 25.53 19.58 17.12 C2H6 39.54 31.10 27.34 C3H8 53.64 42.76 37.57 C4H10 67.64 44.13 47.80 § Heats of formation and reaction Heat of formation：Difference in enthalpy between the compound and starting elements in their standard states obtained indirectly from other components of known ΔHf° correct for necessary phase change (such as vaporization, sublimation) correct for ΔH at different T by heat capacity experimental measurement by calorimeter m C ＋ n/2 H C H ΔH ° (gr) 2 (g) m n O f To calculate the heat of formation of (g) O O 6 C ＋ 4 H ＋ O ΔH ° (gr) 2(g) 2(g) (g) f O O ＋ 7 O 6 CO ＋ 4 H O (s) 2(g) 2(g) 2 (g) O ΔHcomb= -735.9 Kcal/mol 6 C(gr) ＋ 6 O2(g) 6 CO2(g) ΔHcomb= -94.05 (Kcal/mol C) × 6 4 H2(g) ＋ O2(g) 4 H2O(g) ΔHcomb= -68.32 (Kcal/mol H2) O O ΔH = 21.46 Kcal/mol (s) (g) subl O O ΔHf°= 6 × (-94.05) ＋ 4 × (-68.32) ＋735.9 ＋21.46 = -80.22 (Kcal/mol) Relative difference in heat of formation O OCCF3 + CF3COOH ∆H= -10.93 Kcal/mol O OCCF3 + CF3COOH ∆H= -9.11 Kcal/mol ∆H= -1.82 Kcal/mol H2 CH3CH＝CHCH3 CH3CH2CH2CH3 ΔH=-28.6Kcal/mol cis H2 CH3CH＝CHCH3 CH3CH2CH2CH3 ΔH=-27.6Kcal/mol trans The heat of hydrogenation is much smaller than the heat of combustion. Both will give the difference of the stability of the two isomers. § Bond increment calculation of heat of formation Principle of additivity：The property of a large molecule can be approximated by adding the contribution of its component. H H H H H For butane H H HH H (-3.83) × 10 ＋ 3× (2.73) ＝ -30.11 Kcal/mol § Group increment calculation of heat of formation 3 × (-10.08) ＋ 2 × (-4.95) ＋ 1 × (-1.90) ＝ -42.04 Kcal/mol (obs. -41.66) CH3 CH3 4 × (-10.08) ＋ 2 × (-1.90) ＝ -44.12 Kcal/mol (obs. -42.49) CH3 CH3 Further refinements correct for van der Waal strain, angle strain…. The electrostatic model for the stability of branched alkane +1 +1 +1 +1 +1 H -10 +1 +1 H H Charge on H：0.278× 10 esu H H HH -3 -2 -2 -1 -2 -2 -2 -2 Charge on C：neutralizing charge HH H HH HH+1 HH Branched more stable +1 +1 +1 +1 +1 +1+1 +1 JACS 1975, 97, 704. Homolytic & Heterolytic Dissociation Energies homolysis ΔH°hom：standard homolytic A－B(g) A． (g) ＋B． (g) bond dissociation energy heterolysis electron transfer + - A (g)＋B (g) ΔH°het：standard heterolytic bond dissociation energy In gas phase：ΔH°het >ΔH°hom In solution： solvation of ions can lower ΔH°het, so that heterolysis becomes favorable. Use bond dissociation energy to calculate reaction ΔH°r e.g. CH3－HCH3．＋H．ΔH°r＝ +104 Kcal/mol Cl－Cl Cl．＋Cl．ΔH°r＝ +58 Kcal/mol Cl．＋CH3． CH3Cl ΔH°r＝ -84 Kcal/mol ＋） Cl．＋H． H－Cl ΔH°r＝ -103 Kcal/mol CH3－H ＋ Cl－Cl CH3Cl ＋ HCl ΔH°r＝ -25 Kcal Hess Law：The difference in enthalpy between products and reactants is independent of the path of the reaction. by heat of formation ΔH°r＝ΣΔH°f(prd) －ΣΔH°f(pre) = -23.7 Kcal/mol at 300℃ § Bond length and bond energy Bond length is nearly a constant property between molecules. § Polarizability The ability of electron cloud to distort in response to external field, defined as the magnitude of dipole induced by one unit of field gradient. ¾Polarizability decreases across a row of the periodic table, . (C>N>O>F, CH4>NH3>H2O) ¾ Polarizability increases along a column,(S>O, P>M, H2S>H2O) C-I bond is more polarizable than C-Cl ¾alkanes are more polarizable than alkenes, may due to Electronegativity. sp2 carbons are more electronegative than sp3 carbons. § Bonding Model Valence Bond Theory (VB) G.N. Lewis 1916 Chemical bonds result from the sharing of electron pairs between two approaching atoms. The bond is localized. means two + bonding electrons in HH H－H the region H：H between two atom The region is orbital. Cl+ Cl Cl Cl each achieve a filled shell HCl+ H Cl σ bond from S and P σ bond by P and P σ：cylindrical symmetry For complex molecules, hybridization and resonance are used to describe molecules in terms of orbitals which are mainly localized between two atoms. Hybridization Theory： For carbon 1s22s22p2 3 the 2s, 2px, 2py, 2pz hybridize to form four equivalent sp 4 bonds can be formed on carbon highly directional sp3 orbital provide for more efficient overlap. + 4 methane CH4 3 sp sp3 H H H 109.5° C C H H H H H ethene π bond, plane sym. H H HHs + 2p → 3sp2 120° or H H HHone p remaining acetylene s + p → 2sp HHor HH two p remaining H-C-C linear No. of Hybri Geometry ligand d 3 4 sp Tetrahedr CH4, CCl4, CH3OH, al 2 3 sp Trigonal CH2=CH2, H2C=O, C6H6, = CO3 , CH3． 2 sp linear HC≡CH, CO2, HCN, H2C=C=CH2 Resonance Theory： An extension of valence bond theory for molecules that more than one Lewis structure can be written. Useful in describing electron delocalization, in conjugate system and reactive intermediates. (a)If more than one Lewis structure can be written, which has nuclear positions constant, but differ in assignment of electrons, the molecule is described by a combination of these structure (a hybrid of all). (b)The most favorable (lowest energy) resonance structure makes the greatest contribution to the true structure. determining energy：maximum number of covalent bond, minimum separation of unlike charge, placement of negative charge on most electronegative atom (vice versa). (c)Those with delocalized electrons are usually more stable than single localized structure. H 2 C H 2 C H 2 C H 2 C CCH2 CCH2 CO CO H 3 C H 3 C H 3 C H 3 C more stable two equivalent structure H H H H charge located equally on two C’s C C C C H C H H C H H H the allyl cation is planar for maximum p interaction H H H H restricted rotation around single bond H H H H C O H C O H C O C C C C C C H H H H H H majorsignificant minor § Dipole moment：the vector quantity that measures the separation of charges. 0.1 e 0.1e +q -q 1 electron charge = 4.8x10-10 esu d = 1.5x10-8cm bond dipole = q × d = 0.1× (4.8× 10-10esu) × 1.5× 10-8 cm = 0.72× 10-18esu．cm = 0.72 D (1 Debye = 10-18esu．Cm) Molecular dipole is the vector sum of various “bond dipoles”. It provides information about molecular structure and bonding. e.g. CH3F μ= 1.81 D H 1.81× 10-18 H C F q = = 0.27 e- -10 H 1.385× 4.8× 10 1.385Å For dichlorobenzene Cl μ= 2.30, 1.55, 0 Cl Cl Cl μ= 1.61 D Cl Cl Cl From trigonometry, the calculated angles between two bond “dipole moment” are 89° , 122° , 180°. support the concept that benzene is planar. The dipole moment results from unequal sharing of the electron. due to different attraction for electron electronegativity Polar bond = [covalent bond] +λ [ionic bond] λ= weighing factor 2 % ionic character = λ × 100 % (1 +λ2) HCl +0.17 electron charge on H． - 0.17 electron charge on Cl． λ = 0.45 § Electronegativity & Bond Polarity Electronegativity：The power of an atom in a molecule to attract electrons to itself. Pauline 1932 χp：based on the difference in bond energy of AB and A-A + B-B other scale of electronegativity, 2 more related to atomic properties.
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