D. A. Evans An Introduction to Frontier Molecular Orbital Theory-1 Chem 206 http://www.courses.fas.harvardedu/-chem206/ ■ Problem of the Day The molecule illustrated below can react through either Path A or Path B to i form salt 1 or salt 2. In both instances the carbonyl oxygen functions as the Chemistry 206 nucleophile in an intramolecular alkylation. What is the preferred reaction pat for the transformation in question Advanced Organic Chemist Lecture number 1 Path B Introduction to FMo Theory ■ Your Answer General Bonding considerations The H2 Molecule Revisited (Again!) a Donor& Acceptor Properties of Bonding Antibonding States L Hyperconjugation and"Negative"Hyperconjugation Anomeric and related Effects Reading Assignment for week A. Carey& Sundberg: Part A; Chapter 1 B. Fleming, Chapter 1&2 C Fukui, Acc. Chem. Res. 1971, 4, 57 D. O.J. Curnow J Chem. Ed. 1998. 75. 910 E J 1. Brauman. Science. 2002. 295 2245 Matthew d shair Wednesday, eptember 18, 2002
D. A. Evans Chem 206 Matthew D. Shair Wednesday, September 18, 2002 http://www.courses.fas.harvard.edu/~chem206/ ■ Reading Assignment for week: A. Carey & Sundberg: Part A; Chapter 1 B. Fleming, Chapter 1 & 2 C. Fukui,Acc. Chem. Res. 1971, 4, 57. D. O. J. Curnow, J. Chem. Ed. 1998, 75, 910. E. J. I. Brauman, Science, 2002, 295, 2245. Chemistry 206 Advanced Organic Chemistry Lecture Number 1 Introduction to FMO Theory ■ General Bonding Considerations ■ The H2 Molecule Revisited (Again!) ■ Donor & Acceptor Properties of Bonding & Antibonding States ■ Hyperconjugation and "Negative" Hyperconjugation ■ Anomeric and Related Effects An Introduction to Frontier Molecular Orbital Theory-1 ■ Problem of the Day The molecule illustrated below can react through either Path A or Path B to form salt 1 or salt 2 . In both instances the carbonyl oxygen functions as the nucleophile in an intramolecular alkylation. What is the preferred reaction path for the transformation in question? + + Br – Br – 1 2 Path A Path B Br N H O Br O O Br N O H O N O H Br ■ Your Answer
D. A. Evans An Introduction to Frontier Molecular Orbital Theory-1 Chem 206 Universal Effects Governing Chemical Reactions Stereoelectronic Effects There are three ■ Steric Effects Geometrical constraints placed upon ground and transition states by orbital overlap considerations Nonbonding interactions (Van der Waals repulsion) between substituents within a molecule or between reacting molecules Fukui postulate for reaction During the course of chemical reactions, the interaction of the highest filled(HOMO) and lowest unfilled(antibonding) J I Brauman, Science, 2002, 295, 2245 molecular orbital (LUMO)in reacting species is very important to the stabilization of the transition structure General Reaction Types Radical reactions(-10%):A·+B·—A-B Electronic Effects(Inductive Effects) Polar Reactions(90%): A(+ B(+)-A--B rate decreases as r becomes more electronegative Lewis Base Lewis Acid Inductive Effects: Through-bond polarization FMO concepts extend the donor-acceptor paradigm to Field Effects: Through-space polarization non-obvious families of reactions Your thoughts on this transformation ■ Examples to consider H2+2Li()→2LH R3Sio Lewis acid EtO2 C\ CH3-1+ Mg(O)- CH3-MgBr OSiR3 Danishefsky, JOC 1991, 56, 38 diastereoselection >94.6
RO H O C Br Me R R SN2 C R R Me Br Me2CuLi O OSiR3 R3SiO EtO C Me R R RO H O Me H C R R Me Nu RO H O H Me OSiR3 OSiR3 EtO2C H2 CH3–I A(:) A· B(+) B· 2 LiH CH3–MgBr A B A B D. A. Evans An Introduction to Frontier Molecular Orbital Theory-1 Chem 206 + Br:– minor major Br: – Nu: Nonbonding interactions (Van der Waals repulsion) between substituents within a molecule or between reacting molecules ■ Steric Effects Universal Effects Governing Chemical Reactions There are three: ■ Electronic Effects (Inductive Effects): + SN1 rate decreases as R becomes more electronegative Inductive Effects: Through-bond polarization Field Effects: Through-space polarization Danishefsky, JOC 1991, 56, 387 Lewis acid diastereoselection >94 : 6 Your thoughts on this transformation "During the course of chemical reactions, the interaction of the highest filled (HOMO) and lowest unfilled (antibonding) molecular orbital (LUMO) in reacting species is very important to the stabilization of the transition structure." Geometrical constraints placed upon ground and transition states by orbital overlap considerations. ■ Stereoelectronic Effects Fukui Postulate for reactions: ■ General Reaction Types Radical Reactions (~10%): + Polar Reactions (~90%): + Lewis Base Lewis Acid FMO concepts extend the donor-acceptor paradigm to non-obvious families of reactions ■ Examples to consider + 2 Li(0) + Mg(0) J. I. Brauman, Science, 2002, 295, 2245
D. A. Evans The H2 Molecular Orbitals Antibonds Chem 206 The H2 Molecule(again!!) Linear Combination of Atomic Orbitals(LCAO): Orbital Coefficients Let's combine two hydrogen atoms to form the hydrogen molecule ■ Rule two Mathematically, linear combinations of the 2 atomic 1s states create two new orbitals, one is bonding, and one antibonding Each MO is constructed by taking a linear combination of the individual atomic orbitals(AO) Bonding mo Rule one: a linear combination of n atomic states will create n Mos Antibonding MO 0*=C 1V+C2Y2 1s ■ Rule three(c12+(c2 The squares of the C-values are a measure of the electron population AE in neighborhood of atoms in question I Rule Four: bonding(C-+ antibonding(C1=1 In lCao method, both wave functions must each contribute one net orbital Let's now add the two electrons to the new mo, one from each h atom Consider the pi-bond of a C=o function: In the ground state pi-C-0 o*(antibonding is polarized toward Oxygen. Note(Rule 4) that the antibonding MO is polarized in the opposite direction 1s G(bonal Note that△E1 is greater than△E2.Why
D. A. Evans Chem 206 The H2 Molecule (again!!) Let's combine two hydrogen atoms to form the hydrogen molecule. Mathematically, linear combinations of the 2 atomic 1s states create two new orbitals, one is bonding, and one antibonding: Energy 1s 1s s* (antibonding) ■ Rule one: A linear combination of n atomic states will create n MOs. DE DE Let's now add the two electrons to the new MO, one from each H atom: Note that DE1 is greater than DE2 . Why? s (bonding) s (bonding) DE2 DE1 s* (antibonding) 1s 1s y2 y2 y1 y1 Energy s = C1y1 + C2y2 Linear Combination of Atomic Orbitals (LCAO): Orbital Coefficients Each MO is constructed by taking a linear combination of the individual atomic orbitals (AO): Bonding MO Antibonding MO C* s* =C*1y1– 2y2 The coefficients, C1 and C2, represent the contribution of each AO. ■ Rule Three: (C1 )2 + (C2 )2 = 1 antibonding(C*1 ) = 1 2 bonding(C + 1 )2 ■ Rule Four: Energy p* (antibonding) p (bonding) Consider the pi -bond of a C=O function: In the ground state pi-C–O is polarized toward Oxygen. Note (Rule 4) that the antibonding MO is polarized in the opposite direction. C C O C O The H2 Molecular Orbitals & Antibonds The squares of the C-values are a measure of the electron population in neighborhood of atoms in question In LCAO method, both wave functions must each contribute one net orbital ■ Rule Two: H H H H O
D. A. Evans Bonding generalizations Chem 206 I Bond strengths(Bond dissociation energies)are composed of a Orbital orientation strongly affects the strength of the resulting bond covalent contribution (o Ecow and an ionic contribution( Eionid For g Bonds Bond Energy(BDE)=o Covalent+ 8 Ionic( Fleming, page 27) A○ OBo teoA○B When one compares bond strengths between C-C and C-x, where X is some other element such as O, N, F, Si, or S, keep in mind that covalent and ionic contributions vary independently. Hence, the mapping of trends is not a trivial exerci A-B Useful generalizations on covalent bonding This is a simple notion with very important consequences. It surfaces in OVerlap between orbitals of comparable energy is more effective the delocalized bonding which occurs in the competing anti(favored) than overlap between orbitals of differing energy syn (disfavored)E2 elimination reactions. Review this situation For example, consider elements in Group IV, Carbon and Silicon. We An anti orientation of filled and unfilled orbitals leads to better overlap know that C-C bonds are considerably stronger by Ca 20 kcal mol This is a corrollary to the preceding generalization than c-si bonds There are two common situatio ⊙⊙→(t(⑥( Case-1: Anti Nonbonding electron pair&C-x bond 土 X G'C-X lone pa 00 LUMO HOMO LUMO C-SP CSP. 土 cC-C H3C-CH3 BDE =88 kcal/mol H3C-SiH3 BDE -70 kcal/mol Case-2: Two anti sigma bonds Bond length =1.534 A Bond length= 1.87A This trend is even more dramatic with pi-bonds F"C-X C-C= 65 kcal/mol C-si= 36 kcal/mol Si-si= 23 kcal/mol 你8m 0 LUMO a Weak bonds will have corresponding low-lying antibonds Formation of a weak bond will lead to a corresponding low-lying antibonding orbital. Such structures are reactive as both nucleophiles electrophile
A B A A C A C A Y C X A C X X X ·· lone pair HOMO s* C–X LUMO s* C–X LUMO lone pair HOMO C C C C C Si C-SP3 C-SP3 C-SP3 C Si Si-SP3 Y C C X A B A B Y C C B X D. A. Evans Bonding Generalizations Chem 206 ■ Weak bonds will have corresponding low-lying antibonds. p C–C = 65 kcal/mol p C–Si = 36 kcal/mol p Si–Si = 23 kcal/mol This trend is even more dramatic with pi-bonds: s* C–Si s* C–C s C–Si s C–C Bond length = 1.534 Å Bond length = 1.87 Å H3C–CH3 BDE = 88 kcal/mol H3C–SiH3 BDE ~ 70 kcal/mol Useful generalizations on covalent bonding When one compares bond strengths between C–C and C–X, where X is some other element such as O, N, F, Si, or S, keep in mind that covalent and ionic contributions vary independently. Hence, the mapping of trends is not a trivial exercise. Bond Energy (BDE) = Ecovalent + Eionic (Fleming, page 27) ■ Bond strengths (Bond dissociation energies) are composed of a covalent contribution ( Ecov) and an ionic contribution ( Eionic). better than For example, consider elements in Group IV, Carbon and Silicon. We know that C-C bonds are considerably stronger by Ca. 20 kcal mol-1 than C-Si bonds. ■ Overlap between orbitals of comparable energy is more effective than overlap between orbitals of differing energy. Formation of a weak bond will lead to a corresponding low-lying antibonding orbital. Such structures are reactive as both nucleophiles & electrophiles Better than Better than Case-2: Two anti sigma bonds s C–Y HOMO s* C–X LUMO s* C–X LUMO Case-1: Anti Nonbonding electron pair & C–X bond ■ An anti orientation of filled and unfilled orbitals leads to better overlap. This is a corrollary to the preceding generalization. There are two common situations. Better than For p Bonds: For s Bonds: ■ Orbital orientation strongly affects the strength of the resulting bond. Better than This is a simple notion with very important consequences. It surfaces in the delocalized bonding which occurs in the competing anti (favored) syn (disfavored) E2 elimination reactions. Review this situation. s C–Y HOMO
D. A. Evans Donor-Acceptor Properties of Bonding and Antibonding States Chem 206 Donor Acceptor Properties of C-c&C-O Bonds Hierarchy of Donor Acceptor States Consider the energy level diagrams for both bonding& antibonding Following trends are made on the basis of comparing the bonding and orbitals for C-c and c-o bonds antibonding states for the molecule CH3-x where X=C, N,o,F,&H gc-C I o-bonding States: (C-x) c。 C-SP 上csp土 CH3-NH2 decreasing o-donor capacity C-C poorest donor gC-O a The greater electronegativity of oxygen lowers both the bonding g-anti-bonding States:(C-X) antibonding C-O states. Hence oc-c is a better donor orbital than gc-o CHH Io'C-O is a better acceptor orbital than gC-c CH3CH3 CH3-NH2 Donor Acceptor Properties of CsP3-CsP3& CsP3-CsP2 Bonds CH3-OH CH3-F Increasing o'-acceptor capacity best acceptor C-SE The following are trends for the energy levels of nonbonding states of several common molecules. Trend was established by photoelectron spectroscopy C-SP2 Nonbonding States gC-C W 中p卅 a The greater electronegativity of Csp? lowers both the bonding antibonding C-C states. Hence Io CSP3-CsP3 is a better donor orbital than g CsP3-CSP2 do'CSP3-CsP2 is a better acceptor orbital than oCSP3-CSP3 decreasing donor capacity poorest donor
D. A. Evans Donor-Acceptor Properties of Bonding and Antibonding States Chem 206 ■ s *CSP3-CSP2 is a better acceptor orbital than s *CSP3-CSP3 C-SP3 C-SP3 s* C–C s C–C C-SP3 s C–C s* C–C C-SP2 Donor Acceptor Properties of CSP3-CSP3 & CSP3-CSP2 Bonds ■ The greater electronegativity of CSP2 lowers both the bonding & antibonding C–C states. Hence: ■ s CSP3-CSP3 is a better donor orbital than s CSP3-CSP2 ■ s *C–O is a better acceptor orbital than s *C–C ■ s C–C is a better donor orbital than s C–O ■ The greater electronegativity of oxygen lowers both the bonding & antibonding C-O states. Hence: Consider the energy level diagrams for both bonding & antibonding orbitals for C–C and C–O bonds. Donor Acceptor Properties of C-C & C-O Bonds O-SP3 s* C-O s C-O C-SP3 s C-C s* C-C better donor better acceptor decreasing donor capacity Nonbonding States poorest donor The following are trends for the energy levels of nonbonding states of several common molecules. Trend was established by photoelectron spectroscopy. best acceptor poorest donor Increasing -acceptor capacity s-anti-bonding States: (C–X) s-bonding States: (C–X) decreasing -donor capacity Following trends are made on the basis of comparing the bonding and antibonding states for the molecule CH3–X where X = C, N, O, F, & H. Hierarchy of Donor & Acceptor States CH3–CH3 CH3–H CH3–NH2 CH3–OH CH3–F CH3–H CH3–CH3 CH3–NH2 CH3–OH CH3–F HCl: H2O: H3N: H2S: H3P:
D. A. Evans Hybridization vs Electronegativity Chem 206 Electrons in 2S states"see" a greater effective nuclear charge than electrons in 2P states There is a linear relationship between %S character This becomes apparent when the radial probability functions for S and P-states are examined the radial probability functions for the hydrogen atom s&p states are show 100% <sORbital oa06∝ X2 S Orbital <2S Orbital C2POrbital A s-Character 3so己→∧ 3POrbitalO/ There is a direct relationship between %S character hydrocarbon acidity S-states have greater radial penetration due to the nodal properties of the wave function. Electrons in S-states""a higher nuclear charge Above observation correctly implies that the stability of nonb pairs is directly proportional to the of S-4 the doubly occupied orbital CH64) Least stable·==-- Most stable (DCSP3 (DCSP2 (DCsp PhCC-H(2 The above trend indicates that the greater the of S-character at a given atom, the greater the electronegativity of that atom S-Character
2 2.5 3 3.5 4 4.5 5 Pauling Electronegativity 20 25 30 35 40 45 50 55 % S-Character CSP3 CSP2 CSP NSP3 NSP2 NSP 25 30 35 40 45 50 55 60 Pka of Carbon Acid 20 25 30 35 40 45 50 55 % S-Character CH4 (56) C6 H6 (44) PhCC-H (29) CSP3 CSP2 CSP 1 S Orbital 2 S Orbital 3 S Orbital 2 S Orbital 2 P Orbital 3 P Orbital D. A. Evans Chem 206 This becomes apparent when the radial probability functions for S and P-states are examined: The radial probability functions for the hydrogen atom S & P states are shown below. Electrons in 2S states "see" a greater effective nuclear charge than electrons in 2P states. Above observation correctly implies that the stability of nonbonding electron pairs is directly proportional to the % of S-character in the doubly occupied orbital Least stable Most stable The above trend indicates that the greater the % of S-character at a given atom, the greater the electronegativity of that atom. There is a direct relationship between %S character & hydrocarbon acidity There is a linear relationship between %S character & Pauling electronegativity Hybridization vs Electronegativity Å Radial Probability 100 % 100 % Radial Probability Å S-states have greater radial penetration due to the nodal properties of the wave function. Electrons in S-states "see" a higher nuclear charge
D.A. Evans Hyperconjugation: Carbocation Stabilization Chem 206 a The interaction of a vicinal bonding orbital with a p-orbital is referred to as hyperconjugation Physical Evidence for Hyperconjugation This is a traditional vehicle for using valence bond to denote charge delocalization a Bonds participating in the hyperconjugative interaction, e.g. C-R, will be lengthened while the C(+)-C bond will be shortened First X-ray Structure of an Aliphatic Carbocation The graphic illustrates the fact that the C-R bonding electrons can delocalize"to stabilize the electron deficient carbocationic center Note that the general rules of drawing resonance structures still hold the positions of all atoms must not be change Stereoelectronic Requirement for Hyperconjugation 1006°1608A Syn-planar orientation between interacting orbitals The Molecular Orbital Description FC-R FC-R T. Laube, Angew. Chem. Int Ed. 1986, 25, 349 The Adamantane Reference (MM-2) 1528A FC-R I Take a linear combination of o C-R and CSP2 p-orbital The new occupied bonding orbital is lower in energy. When you stabilize the electrons is a system you stabilize the system itself
+ + + + C C R H H H H C H H R C H H H C H C H H Me Me Me H [F5Sb–F–SbF5 ]– Me Me Me C D. A. Evans Chem 206 The Adamantane Reference (MM-2) T. Laube, Angew. Chem. Int. Ed. 1986, 25, 349 First X-ray Structure of an Aliphatic Carbocation 110 ° 100.6 ° 1.530 Å 1.608 Å 1.528 Å 1.431 Å ■ Bonds participating in the hyperconjugative interaction, e.g. C–R, will be lengthened while the C(+)–C bond will be shortened. Physical Evidence for Hyperconjugation The new occupied bonding orbital is lower in energy. When you stabilize the electrons is a system you stabilize the system itself. ■ Take a linear combination of s C–R and CSP2 p-orbital: s C–R s* C–R s C–R s* C–R The Molecular Orbital Description Syn-planar orientation between interacting orbitals Stereoelectronic Requirement for Hyperconjugation: The graphic illustrates the fact that the C-R bonding electrons can "delocalize" to stabilize the electron deficient carbocationic center. Note that the general rules of drawing resonance structures still hold: the positions of all atoms must not be changed. + ■ The interaction of a vicinal bonding orbital with a p-orbital is referred to as hyperconjugation. Hyperconjugation: Carbocation Stabilization This is a traditional vehicle for using valence bond to denote charge delocalization. +
D. A. Evans Negative" Hyperconjugation Chem 206 I Delocalization of nonbonding electron pairs into vicinal antibonding Syn Orientation antibonding a* C-R orbitals is also possible R R P hybrid orbital 1H This decloalization is referred to as"Negative"hyperconjugation;Anti Orientation antibonding o* c-R Since nonbonding electrons prefer hybrid orbitals rather that P:R orbitals, this orbital can adopt either a syn or anti relationship x hybrid orbital to the vicinal c-r bond The Molecular Orbital Description a Overlap between two orbitals is better in the anti orientation as 0*C-R stated in"Bonding Generalizations"handout The Expected Structural Perturbations Nonbonding e pair Change in Structure Spectroscopic Probe Shorter c-x bond H ■ Longer C-R bond As the antibonding C-R orbital C-R decreases in energy, the magnitude ■ Stronger C-X bond nfrared Spectroscopy of this interaction will increase Weaker C-R bond Infrared Spectroscopy Note that o C-R is slightly destabilized ■ Greater e-density at R NMR Spectroscopy ■ Less e-density at NMR Spectroscopy
■ Greater e-density at R NMR Spectroscopy ■ Less e-density at X NMR Spectroscopy ■ Longer C–R bond X-ray crystallography ■ Weaker C–R bond Infrared Spectroscopy ■ Stronger C–X bond Infrared Spectroscopy ■ Shorter C–X bond X-ray crystallography Change in Structure Spectroscopic Probe The Expected Structural Perturbations As the antibonding C–R orbital decreases in energy, the magnitude of this interaction will increase s C–R ●● s* C–R The Molecular Orbital Description ■ Delocalization of nonbonding electron pairs into vicinal antibonding orbitals is also possible D. A. Evans "Negative" Hyperconjugation Chem 206 X Since nonbonding electrons prefer hybrid orbitals rather that P orbitals, this orbital can adopt either a syn or anti relationship to the vicinal C–R bond. C X R H H H H X H H H C H R ●● This decloalization is referred to as "Negative" hyperconjugation antibonding s* C–R ■ Overlap between two orbitals is better in the anti orientation as stated in "Bonding Generalizations" handout. + – Anti Orientation filled hybrid orbital filled hybrid orbital antibonding s* C–R Syn Orientation – C X + H H C X H H C H H C H R X H R X C X H H C X H H R: R: Nonbonding e– pair ●● ●● ●● ●● ●● Note that s C–R is slightly destabilized R R
D. A. Evans Lone Pair Delocalization: N2 F2 Chem 206 The interaction of filled orbitals with adjacent antibonding orbitals can have an ordering effect on the structure which will stabilize a particular The trans Isomer Now carry out the same analysis with the same 2 orbitals present in the trans isomer geometry. Here are several examples ase 1: N2F2 Tras s islemels can exist a N-SP 0 antibonding N-F g N-F N-SP, HOMO) There are two logical reasons why the trans isomer should be more In this geometry the"small lobe"of the filled N-SP2 is required to overlap with the large lobe of the antibonding C-F orbital. Hence, when the new MOs are generated the new bonding orbital is not as stabilizir a The nonbonding lone pair orbitals in the cis isomer will be destabilizing: as for the cis isomer due to electron-electron repulsion Conclusions The individual C-F dipoles are mutually repulsive(pointing in same direction) in the cis isomer a Lone pair delocalization appears to override electron-electron and dipole-dipole repulsion in the stabilization of the cis isomer Let's look at the interaction with the lone pairs with the adjacent c-f; to better orbital overlap calization is stronger in the cis isomer due In fact the cis isomer is favored by 3 kcal/ mol at 25C antibonding orbitals Important Take-home Lesson he cis isomer Orbital orientation is important for optimal orbital overlap N-F antibonding g N-F A-b forms stronger pi-bond than filled (HOMO) AOB° rms stronger。ABo This is a simple notion with very important consequences. It surfaces in I Note that two such interactions occur in the molecule even though the delocalized bonding which occurs in the competing anti(favored) only one has been illustrated syn(disfavored)E2 elimination reactions. Review this situation
N F N N N F N N F F F (HOMO) N F N F A A B B F A A B B Chem 206 The cis Isomer ■ Note that two such interactions occur in the molecule even though only one has been illustrated. ■ Note that by taking a linear combination of the nonbonding and antibonding orbitals you generate a more stable bonding situation. s* N–F filled N-SP2 antibonding s* N–F filled N-SP2 In fact the cis isomer is favored by 3 kcal/ mol at 25 °C. Let's look at the interaction with the lone pairs with the adjacent C–F antibonding orbitals. This molecule can exist as either cis or trans isomers The interaction of filled orbitals with adjacent antibonding orbitals can have an ordering effect on the structure which will stabilize a particular geometry. Here are several examples: D. A. Evans Lone Pair Delocalization: N2F2 Case 1: N2F2 The trans Isomer Now carry out the same analysis with the same 2 orbitals present in the trans isomer. filled N-SP2 antibonding s* N–F ■ In this geometry the "small lobe" of the filled N-SP2 is required to overlap with the large lobe of the antibonding C–F orbital. Hence, when the new MO's are generated the new bonding orbital is not as stabilizing as for the cis isomer. filled N-SP2 (HOMO) s* N–F Conclusions ■ Lone pair delocalization appears to override electron-electron and dipole-dipole repulsion in the stabilization of the cis isomer. (LUMO) (LUMO) .. .. .. .. There are two logical reasons why the trans isomer should be more stable than the cis isomer. ■ The nonbonding lone pair orbitals in the cis isomer will be destabilizing due to electron-electron repulsion. ■ The individual C–F dipoles are mutually repulsive (pointing in same direction) in the cis isomer. ■ This HOMO-LUMO delocalization is stronger in the cis isomer due to better orbital overlap. Important Take-home Lesson Orbital orientation is important for optimal orbital overlap. forms stronger pi-bond than forms stronger sigma-bond than This is a simple notion with very important consequences. It surfaces in the delocalized bonding which occurs in the competing anti (favored) syn (disfavored) E2 elimination reactions. Review this situation
D. A. Evans The Anomeric Effect and Related Issues Chem 206 Infrared evidence for lone pair delocalization into vicinal antibonding orbitals The Anomeric effect The N-H stretching frequency of cis-methyl diazene is 200 cm-1 lower It is not unexpected that the methoxyl substituent on a cyclohexane ring than the trans isomer prefers to adopt the equatorial conformation H antibonding OMe AG°=+0.6 kcal/mol OMe vN-H= 2188 cm i What is unexpected is that the closely related 2-methoxytetrahydropyr prefers the axial conformation v N-H= 2317 cm-1 N-SP2 AG°=-06 kcal/mol OMe a The low-frequency N-H shift in the cis isomer is a result of N-H a That effect which provides the stabilization of the axial OR conformer which overrides the inherent steric bias of the substituent is referred to as bond weakening due to presence of the anti lone par on the vicinal the anomeric effect nitrogen which is interacting with the N-H antibonding orbital. Note that the orbital overlap is not nearly as good from the trans Principal HOMO-LUMO interaction from each conformation is llustrated N C. Craig& co-workers JACS 1979, 101, 2480 dehyde C-H Infrared Stretching Frequencies The IR C-H stretching frequency for aldehydes is lower than the closely related olefin C-H stretching frequency. For years this observation has i axial O lone paire>0*C-H axial o lone paire>0*C-o gone a Since the antibonding C-o orbital is a better acceptor orbital than the antibonding C-H bond, the axial OMe conformer is better stabilized by this interaction which is worth ca 1.2 kcal Other electronegative substituents such as cl, sr etc anomeric stabilization. H vC-H=3050 cm-1 1781A vC-H=2730 cm I We now conclude that this is another example of the vicinal lone pair effect This conformer 1819AC preferred by 1.8 kcal/mol Why is axial C-Cl bond longer?
N N Me N N Me H N H N Me N N Me O H OMe O H OMe C H C R R R Cl O H O O H Cl H Cl H OMe OMe O H H OMe OMe H O H H C R O H Chem 206 ■ We now conclude that this is another example of the vicinal lone pair effect. D. A. Evans The Anomeric Effect and Related Issues filled N-SP2 Infrared evidence for lone pair delocalization into vicinal antibonding orbitals. n N–H = 2188 cm -1 n N–H = 2317 cm -1 filled N-SP2 antibonding s* N–H .. antibonding s* N–H The N–H stretching frequency of cis-methyl diazene is 200 cm-1 lower than the trans isomer. N. C. Craig & co-workers JACS 1979, 101, 2480. n C–H = 3050 cm -1 n C–H = 2730 cm -1 Aldehyde C–H Infrared Stretching Frequencies The IR C–H stretching frequency for aldehydes is lower than the closely related olefin C–H stretching frequency. For years this observation has gone unexplained. The Anomeric Effect It is not unexpected that the methoxyl substituent on a cyclohexane ring prefers to adopt the equatorial conformation. D G° = +0.6 kcal/mol D G° = –0.6 kcal/mol What is unexpected is that the closely related 2-methoxytetrahydropyran prefers the axial conformation: ■ That effect which provides the stabilization of the axial OR conformer which overrides the inherent steric bias of the substituent is referred to as the anomeric effect. axial O lone pair«s* C–H axial O lone pair«s* C–O Principal HOMO-LUMO interaction from each conformation is illustrated below: ■ Since the antibonding C–O orbital is a better acceptor orbital than the antibonding C–H bond, the axial OMe conformer is better stabilized by this interaction which is worth ca 1.2 kcal/mol. Other electronegative substituents such as Cl, SR etc. also participate in anomeric stabilization. This conformer preferred by 1.8 kcal/mol 1.819 Å 1.781 Å Why is axial C–Cl bond longer ? ●● ●● ●● ●● ●● ●● ■ The low-frequency N–H shift in the cis isomer is a result of N–H bond weakening due to presence of the anti lone par on the vicinal nitrogen which is interacting with the N–H antibonding orbital. Note that the orbital overlap is not nearly as good from the trans isomer