Albright, Thomas A. Orbital interactions in chemistry / Thomas A. Albright, Jeremy K. Burdett,. Myung-Hwan Whangbo. – 2nd edition. pages cm. Includes index. Share. Email; Facebook; Twitter; Linked In; Reddit; CiteULike. View Table of Contents for Orbital Interactions in Chemistry. Orbital Interactions in Chemistry begins by developingmodels and reviewing molecular orbital theory. Next, the bookexplores orbitals in the organic-main group.

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The thirty years that have elapsed since the acclaimed first edition of this book have witnessed a spectacular evolution of the field of theoretical and computa-. Chapters 6±11 describe applications of orbital interaction theory to various chemical systems in order to show how familiar concepts such as. studying the orbital interaction between reagent and reactant in terms of a chemical bonds to study the origin of intermolecular bond formation and the.

Both TS and TB orbital interactions occur between close orbitals: the distance between orbital centers is usually less than 3. Hydrogen bonds are formed between two molecules with strongly contrasting electronegativities, one of which is terminated by a hydrogen atom This type of bonding has been studied for more than a century, and remains to be an active topic in contemporary scientific research.

A: MO Theory: Orbital Interactions

Different from conventional hydrogen bonds, the interaction between benzene and methane has a dual nature in that both dispersion and electrostatic terms contribute to the interaction energy Results Orbitals distribution As many experimental and theoretical results have proved that benzene-methane has the on-top type isomer configuration 28 , 29 , 30 , we take this configuration as our model.

All the occupied and front unoccupied orbitals of isolated benzene, isolated methane and benzene-methane complex were checked Figs 1S to 3S. The calculated orbital composition reveals that these orbitals consist of both benzene and methane Table S1. The results indicate that the energy level of the benzene-methane complex is not a simple superposition between benzene and methane. The size of the contributing carbon orbital should be drawn a little smaller than the uninteracted carbon orbital, and the size of the out-of-phase oxygen orbital smaller still.

Consider now the pz orbital on the left. It can overlap and therefore interact, only with one orbital on the right, the pz orbital of oxygen. The interaction is of p type. Because of the smaller overlap, the p-type interaction is intrinsically smaller than a s-type interaction between similar orbitals.

As a rough guide, you may take De L for the p-type interaction to be about one-half of De L for the s-type interaction, similarly for the energies of destabilization. These are transferred to the middle of the diagram unchanged in energy or shape, although, for completeness, they should be redrawn in place; they are nonbonding orbitals of the carbonyl group, nO0 and nO.

Lastly, one must occupy the MOs with the correct number of electrons. A neutral dicoordinated carbon atom has two valence electrons and a neutral uncoordinated oxygen atom has six, for a total of eight. Place electrons into the MOs two at a time. It is predominantly a 2p orbital lying in the plane of the molecule. Within the local C2v point group, it transforms as the b2 irreducible representation. Within the local C2v point group, it transforms as the b1 irreducible representation.

Of the four occupied MOs, two are bonding and two are nonbonding, resulting in a net bond order of 2, that is, a double bond. Both the s and p bonds are polarized toward oxygen, the p more than the s because of the smaller intrinsic overlap p-type overlap is smaller than s-type overlap.

Dipole Moment. A large bond dipole moment is expected, with the negative end at oxygen and the positive end at carbon. The combination of s and p bonds forces coplanarity of the oxygen atom, the carbon atom, and the other two atoms attached to the carbon atom. One cannot say anything about the magnitude of the ionization potential from just one interaction diagram. The transition to this state is symmetry forbidden it is quadrupole allowed but dipole forbidden and will be expected to be weak.

This is indeed the case. The photochemical a cleavage is called the Norrish type I reaction, and the rearrangement is called the Norrish type II reaction. Both are discussed in Chapter Since two reagents are involved, the initial orbital overlaps are small and the interactions will be governed by the principles of weak interactions.

In general, the orbital interaction diagram for the interaction of two reagents is only appropriate for the initially formed van der Waals or hydrogen-bonded complex. The trajectory of approach and the orbitals involved will almost always imply the product, but the interaction diagram will not show the bonding in the product.

If the orbital donating or accepting electrons is nonbonding, the newly formed bond will not greatly perturb the other bonds in the fragment. However, if a s or p bond acts as an electron donor, that bond is weakened since it loses electrons and, in a chemical reaction, will be broken altogether.

Similarly, if, as is usually the case, an empty antibonding orbital is the electron acceptor, the corresponding bond is also weakened, and if the reaction goes all the way, the bond is completely broken. The bonding of the products of a chemical reaction must be deduced with separate interaction diagrams since these are in the regime of strong interactions.

In short, the weak-interaction regime and the strong-interaction regime must not be mixed in the same diagram. This can be accomplished by means of orbital correlation diagrams, as seen in Chapter Basicity or Nucleophilicity. This is constructed along the same principles as for the carbonyl group itself. On the left-hand side we place some or all of the orbitals derived for the carbonyl group in Figure 3.

As stated above, only the HOMO will be required, but let us choose all of them to make the point. On the right hand side, we place the single 1s orbital of the hydrogen proton , sH. Since H is slightly less electronegative than C, sH should be a little above the energy level of the 2 pC orbital of the CH2 group prior to interaction see Figure 3. What are the possible interactions between the left- and righthand sides?

Because the proton is a separate free-roaming species, the sH orbital is not precluded by symmetry from interacting with any of the orbitals of the carbonyl group since it can always approach from a direction that avoids nodal surfaces. The most probable direction is the one which leads to the greatest gain in energy. Since the sH orbital is empty, only the occupied orbitals of the carbonyl group need be considered, and it is usually safe to focus only on the HOMO since this has the smallest energy separation with an empty orbital, the LUMO sH.

In Figure 3. Notice that we have reoriented the molecular framework and all of the orbitals into the same orientation from the one in Figure 3.

To make this point, Figure 3. This diagram shows only the initial interaction as would be suitable for formation of a hydrogen bond. Since the donor orbital is nonbonding, addilower in energy and the sOH tion of the Lewis acid to the oxygen is not accompanied by loss of either the s or p bond.

We can also crudely estimate the basicity of the carbonyl oxygen atom. Acidity or Electrophilicity. The LUMO is also relatively low in energy. Orbital interaction arguments are applied to deduce trajectories of attack of a nucleophile to the carbonyl group by showing the HOMO of the nucleophile ammonia in this case and the LUMO of the carbonyl.

Trajectory of nucleophilic attack on a carbonyl group deduced from orbital interaction theory. The application of orbital interaction theory to understand the electronic structures of molecules, as illustrated in the case of the carbonyl group in Figure 3. The second broad area of application of orbital interaction theory is in the area of intermolecular interactions, from which many aspects of chemical reactivity may be inferred.

Unfortunately, the one-electron theoretical foundation for this kind of long-range interaction involving electron transfer is much less sound. The reason for this is easy to understand and is illustrated in Figure 3. Prior to interaction Figure 3. At intermediate separations Figure 3.

After bond formation Figure 3. In the intermediate stage Figure 3. In terms of valence bond structures, MO theory places comparable emphasis on the three structures R1 , R2 , and R3 , whereas, in reality, resonance structure R2 alone most accurately represents the true situation. The reactivity of B as a donor nucleophile is expected to be correlated with its ionization potential [76]. The MO description remains as a powerful and conceptually simple means for understanding the bonding even in the region of the transition state, and such MO descriptions are derivable by the simple rules of orbital interaction theory.

Particulary for the intermolecular case illustrated by Figure 3. In a solvent of high dielectric constant, these are greatly ameliorated, permitting orbital interactions to dominate. Indeed, it has been shown by ab initio calculations in the presence of a medium modeled as a continuum that groups become less electronegative and less hard with increasing dielectric constant [79].

Molecular orbital calculations also do not provide an easily interpretable picture in terms of group orbital interactions for several reasons. First, the basis set does not consist of group orbitals but rather AO-like functions from which group orbitals are also constructed.

As a case in point, the structures of 1-methylcyclohexyl carbocation as predicted by ab initio MO calculations are readily interpreted in terms of a two-orbital interaction diagram such as Figure 3.

However, evidence of the bonding interaction is completely obscured among the many occupied orbitals of the molecule.

One needed to seek evidence of the bonding by examining the antibonding interactions in the more highly localized LUMO [80, 81]! The interaction diagram is illustrated in Figure 4. This is in recognition that much cancellation will occur due to the node in the middle. Not much can be said about the absolute magnitudes of De L and De U , but we are concerned here with trends. In each of the two interaction diagrams shown and the two implied in between, the energy of the sp n orbital of the C atom is drawn at the same level.

The halogen ends of the bonds are constructed from 2p, 3p, 4p, and 5p orbitals, respectively. The sCX decreases along the series and the sCX bond increases. The predicted trend in sCX bond energies is observed in the bond ionization potentials [82]. No attempt has been made to show the increased size or nodal character of the p orbitals of Cl, Br, or I. In an absolute sense, of course, heterolytic cleavage is not a likely process for any of these bonds in the absence of other factors, as discussed below.

Variation in orbital size is approximately proportional to n 2 , where n is the principal quantum number of the valence shell. It should be noted that because of the smaller overlaps, the last approximation in equation 4. The interaction matrix element takes on double importance since the stabilization energy varies as the square of it. One must bear in mind that heterolysis occurs only in polar media where additional bond forming to solvent takes place.

Heterolytic cleavage of neutral molecules in the gas phase is never observed. The energy required to separate charges, of the order of the ionization potential of the H atom However, it is readily compensated by solvation in polar solvents, especially water. Since solvation energies are very large, the simple theory proposed to this point could only rationalize gross trends. Heterolytic cleavage of charged species in mass spectrometric and negative-ion cyclotron resonance experiments is commonly observed.

Homolytic Cleavage of s Bonds Involving C or H The energy change associated with homolytic bond dissociation is given by the equation Figure 4. Peroxides and the halogens are important reagents in organic chemistry. Orbital interaction theory will indicate how they react and why. The resulting orbitals are not shown; they are not polarized since the interacting orbitals are degenerate. The twisted conformations of hydrazines and peroxides arise from this avoidance.

The most dramatic observation in Figure 4. As a consequence, the oxy and halo compounds are strong Lewis acids subject to nucleophilic attack and strong oxidizing agents can readily accept electrons. The interhalogen s bonds are compared in Figure 4. In this case, the intrinsic interaction matrix element decreases rapidly along the series F, Cl, Br, I, as does the electronegativity. A dramatic exception to this generalization is in their interaction with coordinatively unsaturated transition metal complexes, which may be simultaneously excellent donors and acceptors.

The reactivity of s bonds will be discussed in greater detail in other chapters, but general principles can be expounded here. Diagrams for the interaction of two adjacent s bonds only the two adjacent sp n hybrid ends are shown: In the following discussion, the s bond will be represented by the sp 3 hybrid orbital of the C atom. Interaction of s bonds with neighboring s bonds depends on the orientation of one bond relative to the other Figure 4. It is clear that the interaction is minimized but not zero when the two bonds are perpendicular to each other Figure 4.

It is not immediately obvious, however, whether the coplanar anti Figure 4. In fact, a large body of experimental data has been interpreted to show that the coplanar anti arrangement represents the strongest interaction between adjacent s bonds [6, 84]. As well as the relative orientation of the bonds, the extent of interaction will depend on the polarizations of the bonds.

This situation is depicted in Figure 4. This is the case if the substituent is a metal. Diagrams for the interaction of two adjacent s bonds with C bonded to a a less electronegative group and b a more electronegative group. However, it does not follow that the intrinsic interaction between two sp n hybridized orbitals i.

The through-space interaction of the two p bonds of norbornadiene was presented in Chapter 3 as exemplifying a four-electron, two-orbital interaction.

The interaction of the nonconjugated p bonds of 1,4-cyclohexadiene cannot be treated in the same way: The planar molecular skeleton reduces the direct interaction between the two p bonds while increasing the interaction of the p bonds with the intervening CH2 groups. The interaction diagram is shown in Figure 4.

The consequences of such through-space and Figure 4. The spectroscopic and chemical properties of 1,4-diazabicyclo[2. The two lowest IPs are 7.

Compare these to the IPs of trimethyl amine 8. Whether the bond acts as a donor in an interaction with a low-lying virtual orbital or as an acceptor in an interaction with a high-lying occupied orbital, the consequences for the bond are the same, a reduction of the bond order and a consequent weakening of the bond.

The extreme consequence is a rupture of the bond, as occurs in a hydride transfer or a nucleophilic substitution SN 2. The fact that both consequences are usually observed has suggested an inverse relationship between bond strength as measured by the force constant and the bond length [91]. As a s Acceptor The general features orbital distribution and energy for a s bond between C and a more electronegative element or group, X, are shown in Figures 4.

If the interaction falls short of proton abstraction, the attractive interaction is called a hydrogen bond. Both aspects are discussed further in Chapter Carried to the extreme, an elimination reaction E1cb ensues, as discussed in Chapter If the interaction falls short of elimination, the attractive interaction is called negative hyperconjugation [67].

Referring to Figure 4. So are bonds to H from metals or metal-centered groups e. As a s Donor The general features orbital distribution and energy for a s bond between C and a less electronegative element or group, M, are shown in Figure 4. The energy of the s orbital will be relatively high and the orbital is polarized toward carbon. Optimum interaction between the s orbital and a localized empty orbital of a Lewis acid will occur between the C and M, closer to the carbon end, as shown in Figure 4.

Strong orbital interaction in a weak CH-π hydrogen bonding system

If the interaction falls short of abstraction, a hydride bridge may be formed. As a p Donor A s bond between C and a less electronegative element or group, M, is polarized toward carbon. Optimum interaction between the s orbital and an adjacent localized empty orbital of an electron p acceptor group like carbonyl will occur if the s orbital and the empty p orbital of the neighboring group are coplanar, as shown in Figure 4.

The two-orbital, two-electron interaction is accompanied by charge transfer from the s orbital and consequent reduction of the bond order, as well as partial p bond formation between C and the adjacent group.

Where the adjacent group is a carbocationic center, distortion of the donor s bond toward the acceptor site or migration of the bond i. If the interaction falls short of elimination or migration, the attractive interaction is called hyperconjugation. The tendency to alter the geometry or change the conformation so as to maximize the attractive interaction is well documented [96].

The bonding in car boranes is another example [99]. If the LUMO is the sp hybrid orbital at the C end of a polarized s bond, such as to a halide, geometric distortion also occurs, particularly a lengthening of the receiving s bond. Carried to the extreme, an elimination reaction occurs E1cb, as discussed in Chapter Migration of M to the adjacent group does not occur.

Each carbon atom is considered to be sp 2 hybridized Figure 4. Bonding in cyclopropane: W5 and W6. Orbitals W2 , W3 , and W4 , are closer together simply because the 2p orbitals overlap poorly in the orientation imposed by the three-membered ring geometry. In the absence of substituents, the orientation of the nodal surfaces of the pairs of degenerate MOs is arbitrary, but the orientation will be as shown in Figure 4.

The cyclopropyl ring will therefore be expected to act both as a good p donor and a good p acceptor. The reference energy, a, and the energy scale in units of b are introduced. Each center is sp 2 hybridized and has one unhybridized p orbital perpendicular to the trigonal sp 2 hybrid orbitals.

Hydrogen atoms are part of the framework and are not counted. The overlap integral between two parallel p orbitals is small and is approximated to be exactly zero. Equation 5. The subsequent steps are precisely those which were followed in Chapter 3.

The energy is expressed as an expectation value of the MO [equation 5. If the two atoms are not nearest neighbors, then hAB is set equal to zero. It is usual to divide each row of the determinant by b. While the determinant was expanded in Chapter 3 to yield the secular equation, it is more convenient in general to diagonalize the determinant using a computer.

An interactive computer program, SHMO, has been written to accompany this book []. The MOs are displayed as linear combinations of 2p atomic orbitals seen from the top on each center, with changes of phase designated by shading. The SHMO orbitals of allyl, butadiene, and pentadienyl.

The vertical scale is energy in units of jbj, relative to a. The orbitals which are near a will be of most interest in various applications. In the pentadienyl system, p3 plays the same role. In butadiene, the HOMO is p2. Since the energy of the HOMO is higher than the energy of the HOMO of ethylene, one might conclude that butadiene is more basic than ethylene and more reactive toward electrophilic addition.

As is often the case in orbital interaction theory, one must resort to experimental observations to evaluate the relative importance of opposed factors. Several points may be noted. Each ring has degenerate pairs of MOs as a consequence of the three- or higher-fold axis of symmetry.

SHMO orbitals for cyclopropenyl, cyclobutadiene, cyclopentadienyl, and benzene. The energies are in units of jbj relative to a. Two alternative but equivalent representations are shown for the degenerate p orbitals of cyclobutadiene.

The orientation of the nodal surfaces of the degenerate MOs is entirely arbitrary. Two equivalent orientations are shown for cyclobutadiene.

A perturbation at one of the vertices such as by a substituent will rotate the nodes of the degeneate set so that one node passes through that vertex. The orientations shown are those which should be adopted for the purposes of interaction diagrams involving a single substituent on the ring. Conjugated p systems which do not contain any odd-membered rings are called alternant, and provided all the atoms are the same, alternant systems have a symmetrical distribution of orbital energies about the mean a for C.

These features are readily apparent in the orbitals portrayed in Figures 5. Of the four cyclic conjugated p systems shown in Figure 5. Cyclobutadiene is a special case. With two electrons in the degenerate HOMOs, one would expect that the electrons would separate and that the ground state would be a triplet. However, a distortion of the geometry from square to rectangular would eliminate the degeneracy and permit a singlet ground state.

Theoretical computations suggest that the higher value may be correct []. Thus, we may assume that the following relations hold: A measure of the electron population on each center is easily obtained as below. Verify equations 5. Exercise 3. Verify that the net charge at each carbon atom of each of the neutral ring systems shown in Figure 5. A positive value indicates bonding. Small negative values of BAB may result. These are indications of antibonding or repulsive interaction between the centers concerned.

Show that the bond order for benzene is 0. Molecular orbitals are built up by the interaction of the atomic orbitals of these elements held together at bonding separations.

Computational Chemistry in Drug Design

The energies of electrons in 2p orbitals of N and O, normally found in p bonding environments i. These are suitable values for a pyridine N and a carbonyl O. The energies of electrons in 2p orbitals of N and O in normal saturated bonding environments i. Thus a tricoordinated N is more electronegative than a dicoordinated N, and similarly for dicoordinated versus monocoordinated O.

The change in hN is 0: The change in hO is larger, 1: Within the same molecule, the coordination number of N or O may readily be changed by the process of protonation or deprotonation, as 1 0 1 2 1 2 2 1 1 2 1 2 2 Number of Electrons 0.

This value is probably too high. In a molecular environment, this value is expected to be somewhat less where the presence of other nuclei may stabilize p orbitals relative to s.

The last value is similar to his estimate for NH2. Boyd and Edgecombe have placed all three values near 2. Reed and Allen, using their bond polarity index, have assigned values of 0. Nevertheless, the energy criterion may be applied to deduce gross features.

In Figure 5. The lower the LUMO, the more reactive. Carbocations, with LUMO near a, are the most powerful acids and electrophiles, followed by boranes and some metal cations. Carbonyls, imines, and nitriles exemplify this group. The higher the HOMO, the more reactive. Carbanions, with HOMO near a, are the most powerful bases and nucleophiles, followed by amides and alkoxides. The neutral nitrogen amines, heteroaromatics and oxygen bases water, alcohols, ethers, and carbonyls will only react with relatively strong Lewis acids.

These will be discussed in subsequent chapters. Polyenes are the dominant organic examples of this group. A second class of compounds also falls in this category, coordinatively unsaturated transition metal complexes. These systems are treated separately in Chapter Nucleophilic attack on the p bond even by the strongest Lewis bases has not been reported.

The reaction is undoubtedly facilitated by active participation of the lithium cation as a Lewis acid []. The normal course of reaction of alkenes involves addition of Lewis acids electrophiles yielding an intermediate carbocation which is trapped by a weak nucleophile []. The most common electrophilic addition reactions are summarized in Figure 6.

All reactions are regioselective. Polarization of the HOMO away from the point of attachment of the X substituent directs electrophilic attack to that carbon. Attack may also be directed to X itself in certain cases, although this is usually reversible and may have no net consequences. The X: The R's may be H or alkyl, or aryl, or even acyl.

The weak electrophile, the carbonyl of acetaldehyde, adds at the distal C atom. Polarization of the LUMO away from the point of attachment of the Z substituent directs nucleophilic attack to that carbon.

Associated Data

Attack may also be directed to Z itself in certain cases, and this may be irreversible, providing an alternate pathway for the reaction. The R's may be H, alkyl, aryl, or even X: Thus, the Michael addition is the best example: Twisting of the two ends of the double bond relative to each other has the consequence of reducing the p overlap and hence the resonance integral is less than jbCC j.

The increased susceptibility of twisted strained alkenes toward electrophilic attack has been demonstrated experimentally for MCPBA meta-chloroperbenzoic acid epoxidation [] of a variety of strained alkenes. The rate enhancement was attributed to relief of strain in the transition state, but a correlation was noted with ionization potential, and hence the energy of the HOMO.

Conjugated alkenes which have high HOMO energies also were less reactive than expected on the basis of the correlation with twisted monoalkenes. The ionization potential, 8. First, the carbon atoms are dicoordinated, with the consequence that the p orbitals are at higher energy than aC.

Additional factors such as increased coulombic repulsion of the two electrons in each p MO due to the shorter separation may further destabilize the p MOs. In fact, the reactivity of alkynes toward electrophilic attack is rather similar to that of alkenes when the electrophile would be expected to form an acyclic intermediate e. However, addition of nucleophiles is faster in general than to alkenes [, pp. The case of organometallic bonding is treated in Chapter The intrinsic interaction matrix elements decrease in the following series: SHMO description of p bonds: Indeed, these species are in general not isolable but have been postulated as reactive intermediates [, ].

Silaethene, CH2 SiH2 , has been isolated in an argon matrix []. In both silaethenes and disilenes, the silicon atom is not planar, but fairly strongly pyramidal [].

CH2 ] The major carbon centered reaction intermediates in multistep reactions are carbocations carbenium ions , carbanions, free radicals, and carbenes. Formation of most of these from common reactants is an endothermic process and is often rate determining. By the Hammond principle, the transition state for such a process should resemble the reactive intermediate.

CH2 ]. Carbocations Figure 7. A carbocation is strongly stabilized by an X: Thus, the more X: Carbocationic center interacting with a an X: Thus, an alkyl group may be considered to be an X-type substituent.

The donor abilities of s bonds were discussed in Chapter 4. A carbocation is only weakly stabilized by a Z substituent Figure 7. The interaction is weak because the p bond of a Z substituent is very low in energy and polarized away from the cationic center. The Z-substituted cations are being increasingly reported as intermediates in solvolysis reactions []. A prime example is triphenyl carbocation. Intermolecular Reactions of Carbocations Carbocations are strong Lewis acids which occur as intermediates in reactions following the SN 1 Chapter 9 or E1 Chapter 10 mechanistic routes.

The most obvious and. CH2 ] common reaction is recombination with a nucleophile a Lewis base to form a s bond: If the nucleophilic site HOMO involves a nonbonded pair of electrons path a , a stable covalently bonded complex will form.

If the HOMO is a s bond, direct reaction is unlikely unless the bond is high in energy and sterically exposed, as in a three-membered ring, but if the bond is to H, hydride abstraction may occur path b, steps 1 and 2 or a hydride bridge may form path b, step 1.

The last two possibilities are discussed further in Chapter Intramolecular Reactions of Carbocations Intramolecular reactions of carbocations are shown in the following scheme: Each of the cationic species may react intermolecularly, as shown in the previous scheme at sites labeled A. Of course, step 3 would not occur in this case. From the point of view of kinetic or thermodynamic stability of silyl cations, what can be deduced from orbital interaction theory?

Silicon is less electronegative than carbon Table A. Thus, silyl cations are thermodynamically stable but kinetically very reactive. Carbanions Except for the most highly stabilized carbanions, carbanion chemistry in solution is always complicated by the presence of the counterion, usually a metal, which is a Lewis acid and almost invariably is involved in the course of the reaction.

In recent years, much information has been gathered about carbanion stabilities, structures, and reactiv-.

Carbanionic center interacting with a an X: Figure 7. A carbanion is destabilized by an X: The Lewis basicity and nucleophilicity are greatly increased. Carbanions are easily oxidized and may spontaneously autoionize in the gas phase. Methyl carbanion is stable in the gas phase, but ethyl, 2-propyl, and tert-butyl carbanions have not been observed in the gas phase [].

Accordingly, radical intermediates should always be suspected in X: Carbanions next to two sulfur atoms are common. Halogen substitution also favors carbanion formation []. A carbanion is strongly stabilized by a Z substituent Figure 7. In the last sense, these groups are acting as Z-type substituents.

The synthetic potential of trigonal boron as a Z-type substituent in stabilizing carbanionic centers has been demonstrated []: The very important reactive intermediate, the enolate ion, is an example of a Z-substituted carbanion. The oxygen atom bears the majority of the negative charge, but the HOMO has the highest contribution from the carbon atom.

Accordingly, charged electrophiles hard electrophiles should preferentially add to the oxygen atom and neutral soft electrophiles would be expected to add to the carbon atom. The gas-phase reactions of acyclic enolate ions have been studied by Fourier transform ion cyclotron spectroscopy []. The C or O selectivity was shown to depend on the HOMO energy as measured by electron detachment threshhold energies and frontier orbital interactions as well as the charge distribution.

Carbon Free Radicals Free radicals are molecules with an odd number of electrons. In our simple theory, all electrons are considered to be paired up in molecular orbitals, leaving one orbital with a single electron. Often the SOMO is strongly localized. If the localization is to a tricoordinated trigonal carbon atom, then the radical species is described as a carbon free radical.

The methyl radical is planar and has D3h symmetry. The planarity of the methyl radical has been attributed to steric repulsion between the H atoms []. The C center may be treated as planar for the purpose of constructing orbital interaction diagrams. Carbon free radical center interacting with a an X: Substituent Effects and Reactivity.

If the SOMO is relatively low in energy, the principal interaction with other molecules will be with the occupied MOs three-electron, two-orbital type, Figure 3. In this case the radical is described as electrophilic. If the SOMO is relatively high in energy, the principal interaction with other molecules may be with the unoccupied MOs one-electron, two-orbital type, Figure 3. In this case the radical is described as nucleophilic. A free radical is stabilized by an X: The nucleophilicity of the radical is greatly increased.

A carbon free radical is stabilized by a Z substituent Figure 7. The SOMO is lowered in energy and the free radical is more electrophilic as a consequence. Free-radical polymerization of a 1: It has a low-lying LUMO which is not polarized due to the symmetrical substitution. Because of the low-lying LUMO, dimethyl fumarate is susceptible to nucleophilic attack.

Since the HOMO is also raised in energy as a result of the interaction, vinyl acetate would be particularly susceptible to electrophilic attack, the preferred site being the C atom b to the substituent. Addition of a radical to dimethyl fumarate generates a Z-substituted carbon free radical. Addition of a radical to vinyl acetate generates an X: The situation is depicted in Figure 7. And so on. Interactions which determine the relative reactivities of carboxyalkyl left and acyloxy radicals.

It is related to the RSE by. The RSE values for a number of radicals of the. The trends among the p donor groups are generally as expected by orbital interaction considerations.

There is a large uncertainty in the value for SH, with theoretical computations [] favoring a lower value for RSE than the experimental value []. On the basis of the same electronegativity and intrinsic matrix element arguments as were applied for OH versus Cl, one would expect an RSE of the SH group to be less than that of the NH2 group. The listing permits the placement of the CN and NO2 groups relative to the carbonyls.

As stated above, simple considerations imply that stabilization increases with increasing numbers of stabilizing substituents, but by how much? Table 7. The greatest stabilization ensues when both a Z-type and an X: Carbenes Carbenes are species which contain a dicoordinated carbon atom formally with two valence electrons.

The possible electronic structures of the parent carbene, methylene: CH2 , are shown in Figure 7. The 2p orbital of the dicoordinated C atom is placed above a by about 0: This is incorrect. The coulomb repulsion is most severe when two electrons are constrained to the same small MO. Problems Molecular orbital is the formation of molecular orbitals resulting from the overlap of atomic orbitals. The two new orbitals which are formed from the interaction of two atomic orbitals, are antibonding and bonding orbital.

The antibonding orbital is destabilized and has higher energy than stabilized bonding orbital. How do they interact in the molecule orbitals? And what are the factors that affect the orbital interaction? Molecular orbitals Atomic orbitals are the quantum states of the individual electrons forming the electron cloud and moving around an atom.

In the same orbitals, electrons have the same energy and each orbital contains a maximum of two electrons. Similar to the way electrons occupy the atomic orbitals, electrons occupy the molecular orbitals surrounding the molecule. The formations of molecular orbitals are from the combination of atomic orbitals or more specific, from the wave interaction of atomic orbitals.

Orbital interaction- general concepts: The interaction between two atomic or molecular orbitals will form two new orbitals. One new orbital is antibonding orbital which has the higher energy than the original molecule orbital. The other new orbital is the bonding orbital which is lower in energy than the initial one. The stabilization of the bonding molecular orbital and destabilization of the antibonding can increase when the overlap of two orbitals increases.

Jean, Yvesand , page 14 In the molecular interaction, there are the two important orbitals that interact each other. One is the highest energy occupied molecular orbital is called HOMO. The other one is the lowest energy unoccupied molecular orbital LUMO. When a pair of electrons filled in one of the molecule orbital and no electron occupy in the other orbital, this interaction is very stable and called filled- empty interaction.

Therefore, the interaction between them is very strong. The picture 1 shows the filled-empty interaction of HOMO-LUMO interaction: picture 1 Factors affect the strength of orbital interaction: There are several factors that contribute to the overlap of two atomic orbitals; as a consequence, they affect the strength of the orbital interaction: the symmetry of the orbitals, the energy difference between them, their orbital sizes, and distance between the orbitals First, the orbital interactions are depending on their symmetry.

We stated that orbital interactions are allowed if the symmetries of the atomic orbitals are compatible with one another. Based on the symmetry, the orbital interactions from bonding result in three different types.The book is intended for students of organic chemistry at the senior undergraduate and postgraduate levels and for chemists in general seeking qualitative understanding of the often quantitative data produced by modern computational chemists [8].

Some are, of course — the noble gases come to mind. Case 4: Compare all of the groups and faces of the trans-3,4-dimethylcyclopentanones below, by both internal comparison and external comparison.

The lowest energy and its associated wave function correspond to the ground state of the molecule. Within the local C2v point group, it transforms as the b1 irreducible representation. The interaction of triplet-coupled electron pairs is repulsive and does not lead to bond formation.

The chemical consequences of whether Figure 7. One new orbital is antibonding orbital which has the higher energy than the original molecule orbital.