Reactive Intermediate Chemistry

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approximately equal. This factor is why such a good correlation is seen in Figure 3.2 for the hydrocarbon species. Although the Kd values for the lithium and cesium salts of hydrocarbon-derived carbanions are different in THF,133 the lithium values are uniformly 100 times larger due to the stronger interaction of the cation with THF. This effect cancels out when both scales are arbitrarily referenced to the DMSO value for fluorene (22.9 on a per hydrogen basis). Note that the situation is somewhat more complicated in THF because two types of ion pairs are formed, contact ion and solvent-separated ion pairs. In the former, the cation directly interacts with the carbanion whereas in the later, the cation is fully solvated and it is the cation–solvent complex that interacts with the carbanion. Cesium generally prefers contact ion pairs, but lithium prefers solvent-separated ion pairs with delocalized hydrocarbon carbanions because it allows for strong electrostatic interactions of the lithium cation with the oxygens of THF.

For an ion-pair acidity value to differ significantly from the free-ion value, Scheme 3.3 indicates that there must be a large difference between Kd1 and Kd2. In other words, the species of interest must have an ion-pair dissociation constant that is very different from the reference hydrocarbon used in the measurement (R1H). For example, when R2 is an enolate, Kd2 will be very small for lithium because there will be a strong electrostatic interaction with the highly charged oxygen of the enolate and the salt will resist dissociation. As a result, KIP will be unusually large and the parent acid, R2H, will appear to be unusually acidic under the ion-pairing conditions. The 9.2 pK difference between the lithium and cesium ion-pair acidities of p-biphenyl isopropyl ketone indicates that the formation of the lithium ion pair (relative to the cesium ion pair) provides roughly nine orders of magnitude of stabilization to the enolate. This is an overstatement of the true effect because lithium has a higher dissociation constant than cesium with the hydrocarbon reference species, but it still implies that Kd is nearly 10 million times smaller for lithium cations than cesium cations with this enolate. Obviously, there is a very strong electrostatic interaction between the small, densely charged lithium cation and the oxygen of the enolate (in this case, the lithium must be adopting a contact ion pair). The large cesium cation doesn’t offer the high charge density of the lithium cation and therefore ion-pairing has little effect on the stability of the enolate relative to the delocalized carbanions used to develop the acidity scale (i.e., those that fit the correlation line in Figure 3.2), and hence the cesium ion-pair pK is close to the DMSO value.

Finally, it is worth noting that large differences in acidity can occur in switching to solvents capable of hydrogen bonding. As expected, the effect is closely related to nature of the atom bearing the bulk of the charge in the anion. Cox and Stewart134 developed an acidity function that uses data in aqueous DMSO mixtures to extrapolate to a dilute aqueous reference state. With this approach, the pKa of fluorene is nearly the same as in pure DMSO (22.1 vs. 22.6). However, acetone’s pK value is >6 units lower in aqueous solution100 than DMSO. Clearly, the localization of charge on the oxygen is allowing the acetone enolate to benefit more from the access to hydrogen bonding and it is greatly stabilized by the solvent relative to a delocalized carbanion derived from a hydrocarbon. Nitroalkanes also exhibit

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closer together with cyclopentadiene (0.0 kcal/mol) > indene ( 3.1 kcal/mol) > fluorene ( 3.7 kcal/mol).135 Despite all of them yielding highly delocalized hydrocarbon carbanions, the change in the size across this series has a marked effect on their solvation energies in DMSO. The effect is not limited to this solvent and a similar trend is observed for the cesium ion pairs of these carbanions in cyclohexylamine. Clearly, the larger carbanions benefit from greater charge delocalization in the gas phase, but this also reduces their solvation energy significantly and effectively destabilizes them in solution. Third, field effects caused by the interaction of an internal dipole with a carbanion are generally attenuated in solution. For example, perfluoro substitution on fluorene causes a 26.7 kcal/mol decrease in the gasphase free energy of deprotonation, but only a 16.1 kcal/mol decrease in DMSO. The internal dipoles effectively provide another mode of charge delocalization, and therefore reduce the solvation energy of the carbanion. However, there are examples where the opposite is true. A para nitro group enhances the gas-phase acidity of toluene by 28.4 kcal/mol, but the effect in DMSO is 31.2 kcal/mol. In other words, the nitro group increases the solvation energy of the resulting benzylic carbanion. This increase suggests that the solvent must have favorable electrostatic interactions with the charge-bearing oxygens of the nitro group that can overcome the effects of delocalizing the charge (Scheme 3.4).

Scheme 3.4

4.3.4. Kinetic Acidities in the Condensed Phase. For very weak acids, it is not always possible to establish proton-transfer equilibria in solution because the carbanions are too basic to be stable in the solvent system or the rate of establishing the equilibrium is too slow. In these cases, workers have turned to kinetic methods that rely on the assumption of a Brønsted correlation22 between the rate of proton transfer and the acidity of the hydrocarbon. In other words, log k for isotope exchange is linearly related to the pK of the hydrocarbon (Eq. 13). The a value takes into account the fact that factors that stabilize a carbanion generally are only partially realized at the transition state for proton transfer (there is only partial charge development at that point) so the rate is less sensitive to structural effects than the pK. As a result, a values are expected to be between zero and one. Once the correlation in Eq. 13 is established for species of known pK, the relationship can be used with kinetic data to extrapolate to values for species of unknown pK.

As noted above, internal return can be a problem in kinetic acidity measurements, but Streitwieser et al.136 showed that this is not a major impediment in cyclohexyl-

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negative charge, limiting the effects of resonance. In the product, the charge is fully delocalized and the effects of resonance stabilization are fully realized.

Scheme 3.5

In some cases, anomalous values have been obtained in rate–equilibrium relationships involving proton transfers. The most famous one involves nitroalkanes

where substituents cause a greater effect on the relative rate of proton transfer than the position of the equilibrium.140–144 In other words, a plot of log k versus

pK gives a slope >1. Bordwell and Hughes141 suggested an explanation much like the situation shown in Scheme 3.5. In the rate-determining step, hydrogenbonding causes charge to be localized at the carbanion center rather than on the nitro group’s oxygens. If the substituents are better able to interact with the carbanion center than the nitro group, their impact is greatest at the transition state (i.e., more charge is localized at the carbanion in the transition state). After the proton transfer is complete, the charge relocates onto the oxygens of the nitro group and the substituent’s effect is diminished. Recent computational work supports this view.145 However, it appears that nitroalkanes do not give anomalous protontransfer behavior in the gas phase so the importance of a high degree of charge delocalization onto the nitro’s oxygens seems to be a solvation phenomenon.146 As noted in Section 4.3.3, solvation has marked effects on the stability of nitrosubstituted carbanions.

In a novel kinetic approach, Dorfman et al.147 developed methods for rapidly generating very reactive carbanions such as the benzyl anion in solvent mixtures containing water and alcohols. With pulsed radiolysis techniques, they have been able to study the fast and very exothermic reactions of carbanions with these solvents. The studies have shown that despite the high exothermicity, the protonation is not diffusion controlled and depends on the nature of the carbanion’s counterion.

4.4. Carbon Acidities and Bond Strengths

A very useful thermodynamic cycle links three important physical properties: homolytic bond dissociation energies (BDE), electron affinities (EA), and acidities. It has been used in the gas phase and solution to determine, sometimes with high

accuracy, carbon acidities (Scheme 3.6).148 For example, the BDE of methane has been established as 104:9 0:1 kcal/mol149 and the EA of the methyl radical, 1:8

0:7 kcal/mol, has been determined with high accuracy by photoelectron spectroscopy (PES) on the methyl anion (i.e., electron binding energy measurements).55 Of course, the ionization potential of the hydrogen atom is well established, 313.6 kcal/ mol, and as a result, a gas-phase acidity ( Hacid) of 416:7 0:7 kcal/mol has been

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Scheme 3.6

determined for methane.54 The cycle can also be used to calculate BDEs using acidity and EA data. Accurate bond strengths in fundamental species such as acetylene150 and benzene151 have been determined in this way. Finally, the same set of equations can be applied to condensed-phase data with the terms taking slightly different meanings (e.g., instead of EA, the reduction potential of the radical is used). This approach has been widely exploited by Bordwell and co-workers152 who have used their extensive set of acidity measurements in DMSO as a foundation for determining a wide range of C H BDE values. Note that earlier, Breslow153 applied a related set of equations and used electrochemical data to estimate the acidity of very weak carbon acids.

5. REACTIVITY

Carbanions play critical roles in a wide variety of reaction pathways. As stated in the Introduction, this chapter will not focus on the synthetic utility of carbanions, but will instead focus on their mechanistic significance. In this section, a sample of important reaction mechanisms that involve transient or relatively short-lived carbanion intermediates will be introduced. As you will see, the key element in these mechanisms is the ability to form a carbanion that is reasonably stable, and often the kinetics of the reactions are dominated by carbanion stability. The role of carbanion intermediates in elimination reactions will be presented in some detail as a way to illustrate some of the methods that have been developed to probe for carbanion intermediates in reaction mechanisms. Other processes including additions and rearrangement reactions will be presented in less detail, but the role of carbanion stability in these reactions will be outlined.

5.1. Carbanion Intermediates in Elimination Reactions

It is well known that base-induced elimination reactions can proceed either by

a single, concerted step (E2), or by two steps, proton transfer and leaving group expulsion, with a carbanion intermediate (E1cB) to yield an alkene.154–156 The

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deciding factor between the two pathways is whether the carbanion intermediate is stable enough such that the leaving group expulsion step has a barrier on the potential energy surface. In simple alkyl halide systems (e.g., ethyl bromide), the carbanion is a high-energy species and the leaving group expulsion step does not have a barrier so the two steps, merge into one and an E2 mechanism is observed. However, if powerful carbanion stabilizing groups are introduced at the b-carbon, it is possible to shift the system to an E1cB pathway.

Scheme 3.7

As indicated in Scheme 3.7, the first step of an E1cB mechanism can be reversible and therefore deprotonation at the b-carbon does not always lead to product formation. By applying a steady-state approximation to the carbanion concentration, the following rate law is obtained for an E1cB reaction:

The carbanion can be destroyed in two ways, k2 or k 1, and two limiting types of behavior are expected for E1cB mechanisms. If k2 k 1[BH], then the carbanion always decomposes to the alkene product and the rate law simplifies to kobs ¼ k1. In other words, the rate is only dependent on carbanion formation and the rate law has a form that is identical to what would be obtained in a concerted E2 reaction. This has been referred to as an E1cBIRR mechanism, where IRR indicates that deprotonation is irreversible. On the other hand, if k2 k 1[BH], then carbanion formation rarely leads to the product and the rate law simplifies to the following:

These reactions often have been run in buffer solutions where [B ] and [BH] represent the concentrations of the buffer’s components. In Eq. 15, it can be seen that the rate depends on the concentration of the proton donor in the buffer solution. The