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Abstract

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies

Reviewed herein are equilibria between halocarbenes and halocarbanions, carbenes and carbene complexes and carbenes and O-ylides. The transient species were visualized by laser flash photolysis coupled with UV–Visible spectroscopy. This methodology enabled the determination of equilibrium constants and the extraction of associated thermodynamic parameters. Parallel computational studies provided anticipated structures and energies for the transient species, as well as electronic absorption wavelengths and oscillator strengths.


Carbene–Carbanion Equilibria

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies

The most well-known equilibrium in carbene chemistry is arguably the second step of Hine's classic mechanism for the hydrolysis of chloroform, Eq. (2) in the three-step sequence Eqs. (1)-(3) [1].

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Although step (2) is shown as a (rate determining) forward process, Hine later demonstrated that dichlorocarbene (CCl2) could react with Cl, and with other halide ions Br or I, so as to either reform the trichloromethide carbanion or afford the new trihalomethide carbanions CBrCl2 or CICl2 [2]. Thus, step (2) should be written as an equilibrium and we can ask what is the equilibrium constant, K.

To determine K for carbene–carbanion equilibria it is necessary to visualize these species for quantitation. A combination of laser flash photolysis (LFP) for carbene generation, coupled with UV–Vis spectroscopy to visualize the resulting short-lived intermediates, enables the quantitative study of carbene–carbanion equilibria, provided that the key species have accessible spectroscopic signatures.

CCl2 may be generated by LFP of dichlorodiazirine, Eq. (4) [3]. When this reaction is carried out in 1:1 CH2Cl2–MeCN solution containing

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0.9 m tetrabutylammonium (TBA) chloride, a weak absorption appears at 328 nm, assigned to CCl3; the absorption is calculated (TD-B3LYP/6-311 + G(d)//PBE/6-311 + G(d), in CPCM-simulated MeCN) at 344 nm [4]. The appearance of CCl3 suggests the operation of the reverse of Eq. (2); cf. Eq. (4). However, it is not possible to directly determine K for the Eq. (2)/Eq. (4) equilibrium because the very weak absorption of CCl2, observed at 460–530 nm in an argon matrix [5, 6], cannot be detected under our LFP conditions.

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A system that is easier to study is CCl2/CCl2Br, obtainable by the reaction of CCl2 with Br, Eq. (5) [4].

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The bromodichloromethide carbanion absorbs at 388 nm (computed at 376 nm) in 1:1 CH2Cl2–MeCN with 0.6 m TBABr [4]. Although CCl2 itself is still not visualized, the strong absorbance of CCl2Br makes possible an indirect estimation of K for Eq. (5), extracted from the competitive reactions shown in Scheme 1 [4].

Here, LFP of diazirine 1 generates CCl2 which is either trapped by Br to give CCl2Br (k1) or captured by cyclohexene (k2) to yield 7,7-dichloronorcarane 4. Carbanion CCl2Br can revert to CCl2 and Br (k−1) or add to acrylonitrile (k4), ultimately affording cyclopropane 3 via Michael intermediate 2. Finally, CCl2 itself can add directly to acrylonitrile (k3), yielding cyclopropane 3. All the rate constants of Scheme 1 can be measured, except for k−1. Thus, k1 (2.1 × 107 m−1 s−1) is obtained from the slope of a correlation of the observed rate constants for the rise of CCl2Br vs the concentration of Br [4]. The rate constant k2 (6.4 × 107 m−1 s−1) for the addition of CCl2 to cyclohexene is obtained using pyridine to visualize CCl2 as the pyridine ylide [3, 7, 8]. The rate constant for the addition of CCl2 to acrylonitrile (k3 = 4.9 × 105 m−1 s−1) is obtained from competitive CCl2 additions to acrylonitrile vs cyclohexene [9]. Finally, k4 (4.1 × 106 m−1 s−1) is derived from the slope of a correlation of the observed rate constants for the decay of CCl2Br absorbance vs the concentration of acrylonitrile [4].

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Scheme 1. Reactions of dichlorocarbene and bromodichloromethide carbanion.

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With rate constants k1 − k4, and an apparent value of k3/k2 = 0.18 in the presence of 0.28 m Br, we can extract K = k1/k−1 ~ 10 m−1 from Scheme 1 [4]. Therefore, k−1 ~ 2.1 × 106 s−1 for the reversion of CCl2Br to CCl2 and Br. A computational study at the PBE/6-311 + G(d) level in simulated acetonitrile gives ΔH° = −7.7 kcal mol−1 and ΔG° = −2.5 kcal mol−1, leading to ~ 66 m−1 for the CCl2 + Br image CCl2Br equilibrium at 25°C [4].

Although CCl2 lacks a useful UV absorbance in our LFP experiments, its phenyl derivative phenylchlorocarbene (PhCCl) exhibits a strong π[RIGHTWARDS ARROW]p absorption at ~300 nm, together with a weak σ[RIGHTWARDS ARROW]p absorption at long wavelengths [10]. The UV–Vis spectrum of phenylbromocarbene (PhCBr) is analogous. Accordingly, PhCCl and PhCBr support equilibria with halide ions and reactions with acrylonitrile that parallel CCl2 in Scheme 1 [11]. This is illustrated for PhCBr in Scheme 2.

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Scheme 2. Reactions of phenylbromocarbene and phenyldibromomethide carbanion.

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Here, PhCBr is generated by LFP of phenylbromodiazirine 5 in the presence of bromide ion (TBABr), setting up an equilibrium between PhCBr and PhCBr2. The carbanion absorbs strongly at 430 nm. Analogously to Scheme 1, a competition ensues between PhCBr additions to tetramethylethylene (k2) or acrylonitrile (k3), and the addition of PhCBr2 to acrylonitrile (k4). From the measured rate constants k1 − k4, and the product ratio for 7/8 as a function of bromide concentration, we derive K = k1/k−1 = 2.8 m−1 and k−1 = 7.9 × 107 s−1, where k1 = 2.2 × 108 m−1 s−1 [11]. As we show below, a subsequent direct determination of K is in excellent agreement with the value extracted from Scheme 2.

An important practical aspect of the reactions in Scheme 2 is bromide ion catalysis of the addition of PhCBr to acrylonitrile. Thus, PhCBr is an electrophilic carbene that adds well to an electron-rich alkene like tetramethylethylene (k2 = 3.9 × 107 m−1 s−1), but poorly to an electron-poor alkene like acrylonitrile (k3 = 1.7 × 106 m−1 s−1) [12]. Therefore, the rapid conversion of PhCBr to PhCBr2 (k1 = 2.2 × 108 m−1 s−1) and the efficient capture of PhCBr2 by acrylonitrile (k4 = 2.9 × 107 m−1 s−1), result in enhanced conversion of PhCBr to cyclopropane 7 via Michael addition intermediate 6.

For example, in the absence of Br, the relative rate of addition of PhCBr to acrylonitrile vs tetramethylethylene (k3/k2) is 0.042. In the presence of 0.58 m TBABr, krel increases 28-fold to 1.19 due to operation of the k1/k4 pathway [11]. There is an effective umpolung of PhCBr from an electrophilic to a nucleophilic reagent via PhCBr2. Similar results are obtained with the PhCCl/PhCCl2 system where krel for addition to acrylonitrile vs tetramethylethylene increases from 0.045 in the absence of Cl to 0.62 in the presence of 0.54 m TBACl, corresponding to a 13.8-fold enhancement via PhCCl2 [11].

A similar catalysis operates in the addition of phenylfluorocarbene (PhCF) to acrylonitrile. The relative rate of addition of electrophilic PhCF to acrylonitrile vs tetramethylethylene is 0.080 in a 1 m MeCN–THF solvent. In the presence of 0.49 m TBABr, however, krel increases to 1.40, a 17.5-fold increase [13]. Bromide ion catalysis operates through conversion of PhCF to PhCFBr, which adds rapidly to acrylonitrile. The mechanism is analogous to that shown for PhCBr in Scheme 2. In the case of PhCF, generated by LFP of phenylfluorodiazirine, measurements of the corresponding rate constants lead to an estimate of K ~ 27 m−1 for the equilibrium of Eq. (6), where k1 = 2.46 × 107 m−1 s−1 and k 1 = 9.12 × 105 s−1 [13].

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Thus far, our determinations of the equilibrium constants for carbene–carbanion equilibria have been indirect. However, when both species exhibit strong UV–Vis absorptions, K can be determined directly. Let us revisit the PhCBr–PhCBr2 system of Scheme 2. LFP of phenylbromodiazirine 5 in the presence of 0.1 m TBABr in dichloroethane (DCE) gives the calibrated spectrum shown in Fig. 1 [14].

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Figure 1. UV–Vis spectrum acquired 80 ns after LFP of phenylbromodiazirine with 0.10 m TBABr in DCE (after calibration). Absorptions of PhCBr are at 316 and 620 nm; absorption of PhCBr2 is at 428 nm. Reprinted with permission from ref. [14], copyright 2012, American Chemical Society.

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Here, the π[RIGHTWARDS ARROW]p absorption of PhCBr appears at 316 nm and its σ[RIGHTWARDS ARROW]p absorption falls at 620 nm. The PhCBr2 carbanion absorbs at 428 nm. From measurements of the absorbance ratio, A316/A428 as a function of bromide concentration, we obtain K = 3.0 m−1 for the equilibrium of Eq. (7), in excellent agreement with the “indirect” value (K = 2.8 m−1) derived from Scheme 2 [11, 14]. We also measured k1 = 3.55 × 107 m−1 s−1, so that k−1 = 1.18 × 107 s−1.

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Three qualifications should be noted. (1) Carbene–carbanion equilibria are not stable on the ms time scale; both species are reactive and the equilibrium will be drained from both ends. However, the A316/A428 absorbance ratio is relatively constant from 200 to 300 ns after the laser flash so that an average value of the ratio can be determined over the 200–250 ns postflash time interval [14]. (2) The extinction coefficients of the carbene and carbanion, necessary to transform the absorbance ratio into a concentration ratio, are unknown. Instead, we used the computed oscillator strengths (f) of PhCBr and PhCBr2 at the appropriate wavelengths (f = 0.4326 and f = 0.1370, respectively; TD-B3LYP/6-311 + G(d)//M06-2X/6-311 + G(d) in simulated DCE solvent). We estimate that the f ratio reproduces the desired ratio of extinction coefficients to within a factor of 2 [14]. It should be emphasized that computational studies were crucial throughout our studies, providing anticipated absorption wavelengths and structures for the various transient species, as well as computed f values. [3] The determination of K ignores possible cation–anion interactions between TBA+ and Br or TBA+ and PhCBr2. Potential aggregation of cation–anion pairs is also not considered. Factors (2) and (3) introduce some uncertainty into the calculation of K, so that the agreement between the “indirect” value (2.8 m−1) and “direct” value (3.0 m−1) is comforting.

An analogous determination of K for Eq. (8), based on absorptions of PhCCl at 292 nm and PhCCl2 at 404 nm, gave K = 4.01 m−1 [14]. The value of k1 was independently measured as 1.97 × 108 m−1s−1, so that k−1 = 4.91 × 107 s−1 [14].

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Determination of K as a function of temperature (260–309 K) permitted the extraction of the thermodynamic parameters governing Eqs. (7) and (8). For PhCBr/PhCBr2, we found Δ = −1.9 kcal mol−1, ΔS° = −4.3 eu and ΔG° = −0.62 ± 0.03 kcal mol−1 at 298 K. For PhCCl/PhCCl2, the analogous values were ΔH° = −5.7 kcal mol−1, ΔS° = −17 eu and ΔG° = −0.71 kcal mol−1 [14]. Computed values of the thermodynamic parameters for Eqs. (7) and (8) are more negative than the experimental values. At the M06-2X/6-311 + G(d) level, for example, ΔH° = −10.4 (PhCCl) or −8.4 (PhCBr) kcal mol−1, while ΔG° = −3.7 or −2.0 kcal mol−1, respectively. ΔS° ~ −21 eu in both systems [14]. Differences between the experimental and computed values may stem from neglect in the calculations of halide and halocarbanion interactions with the ammonium cations.

We see that although the formation of PhCX2 from PhCX + X is favorable (ΔH < 0), it is opposed by the unfavorable entropy associated with converting two reactants into one product. As a result, Δ is only slightly negative and K for Eqs. (7) and (8) is only 3–4 m−1. Put another way, k1 for the formation of PhCX2 is very rapid when X = Cl or Br, but the reversion of the carbanion to the carbene is similarly fast.

Following the direct measurement of K for the PhCCl/PhCCl2 system, we studied the equilibria between a series of ring-substituted ArCCl [9] and ArCCl2 [10]; cf. Eq. (10) [15].

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    Structure 1. 

All the carbenes and carbanions of Eq. (10) were readily observed, and values of K, k1 and k−1 were obtained. These constants are collected in Table 1, together with analogous data for the previously discussed carbene–carbanion equilibria.

Table 1. Equilibrium and kinetic data for carbene–carbanion systems
CarbeneCarbanionK (m−1)k1 (m−1 s−1)k1 (s−1)References
  1. *Estimated value; see text.

CCl2CCl2Br10a*2.1 × 1072.1 × 106* [4]
PhCFPhCFBr27a*2.5 × 1079.1 × 105* [13]
PhCClPhCCl24.02.0 × 1084.9 × 107 [14]
PhCBrPhCBr23.03.6 × 1071.2 × 107 [14]
9b10b313.2 × 1081.0 × 107 [15]
9c10c2405.2 × 1082.2 × 106 [15]
9d10d3.01.0 × 1083.4 × 107 [15]
9e10e563.8 × 1086.9 × 106 [15]

For Eq. (10), equilibrium constants are a sensitive function of the aryl substituent: K ranges from a low of 3.0 m−1 with a p-fluoro substituent to a high of 240 m−1 with a p-trifluoromethyl substituent [15]. The substituent dependence of K is nicely illustrated by a Hammett correlation of K vs σp where ρ = +3.18 (using σp+ for p-F and σm for m-Cl) [15]. The positive ρ value indicates that electron-withdrawing substituents stabilize the carbanion and destabilize the carbene, shifting the equilibrium toward the carbanion and increasing K.

A Hammett correlation of computed values of K (wB97XD/6-311 + G(d) in simulated DCE) vs σp gave ρ = +12.0 [15]. The sign and significance of ρ are analogous to that proffered for correlation of the experimental K values. However, the huge value of ρ for the computed K values reflects a full unit negative charge on the carbanions. In the experimental systems, cation–anion interactions between Bu4N+ and carbanion 10 will reduce the effective negative charge and decrease ρ. Also, aggregation of Bu4N+ArCCl2 might lead to further dispersal of the carbanion negative charge and an additional decrease in ρ [15].

The forward reaction of Eq. (10) also exhibits ρ > 0 for Hammett correlation of k1. Values of +1.05 in DCE [15] or +0.86 in MeCN–THF and CCl4 [16] are observed. Electron-withdrawing groups that stabilize the carbanion and destabilize the carbene accelerate carbanion formation and increase k1. However, because there is only a partial negative charge on the carbanionic carbon in the transition state for k1 of Eq. (10), the magnitude of ρ is only about a third of ρ for the correlation of K, where the negative charge on the carbanionic carbon must be considerably greater [15].

For the reverse reaction of Eq. (10), correlation of k−1 vs σp gives ρ = −2.63 [15]. The change in the sign of ρ, from positive for k1 to negative for k−1, signifies that in the reverse direction electron donating groups destabilize the carbanion, stabilize the carbene and accelerate the reaction [16]. The Hammett correlations of K, k1 and k−1 elucidate the sense and magnitude of the electronic effects that modulate the equilibria of Eq. (10).

Carbene–Carbene Complex Equilibria

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies

LFP-generated halocarbenes such as methylchlorocarbene, p-nitrophenylchlorocarbene and dichlorocarbene react with electron-rich aromatic ethers (e.g. 1,3,5-trimethoxybenzene, TMB) to form transient π-complexes and O-ylides [17-19]. These reactions are akin to the solvation of carbenes [20]. Indeed, Platz et al. estimated K = 222 m−1 and ΔG° = −3.2 kcal mol−1 for O-ylide formation between a dioxane-solvated chlorocarbene amide and dioxane, Eq. (11) [21].

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    Structure 2. 

Earlier, carbene–alkene complexes had been suggested by Turro and Moss to account for the observation of negative activation energies in the additions of PhCCl to alkenes [22]. However, the negative activation energies could be explained by a carbene–alkene addition model that eschewed carbene–alkene complexes, and such complexes did not appear to be minima on computed reaction energy surfaces [23].

We reported the first direct spectroscopic analysis of a carbene–carbene complex equilibrium [24]. LFP-generated PhCCl reacted with TMB to afford the UV–Vis spectrum shown in Fig. 2.

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Figure 2. Calibrated LFP UV–Vis spectrum of PhCCl in 0.75 mm TMB/pentane solution 50 ns after laser flash; PhCCl absorptions at 324 and 596 nm, PhCCl-TMB complex at 484 nm.

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Here, the π[RIGHTWARDS ARROW]p and σ[RIGHTWARDS ARROW]p absorptions of PhCCl appear at 324 and 596 nm, respectively. The absorption of the PhCCl–TMB π-complex appears at 484 nm and represents charge transfer from the π-system of TMB to the vacant carbenic p-orbital [24]. The two most stable computed complexes (B97D/6-311 + G(d)) are illustrated in Fig. 3. They are sandwich-type species featuring substantial overlap of the aromatic rings and interaction of the carbenic carbon with C2 of TMB [24].

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Figure 3. Two computed complexes formed between PhCCl and TMB (in perspective, carbene on top; green, chlorine; red, oxygen; gray, carbon; white, hydrogen). Reprinted with permission from ref. [24], copyright 2010, American Chemical Society.

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From measurements of A324/A484 as a function of TMB concentration at 294 K, we obtained K = 1264 m−1 for the equilibrium of Eq. (9). We used computed oscillator strengths (f) in place of the unknown extinction coefficients of PhCCl and the PhCCl/TMB complex [24]. The experimental Ks are determined as K = Kapp multiplied by the computed oscillator strength ratio for the carbene and the complex. The computed ratios range up to 3.6 for the various complexes. Thus, errors in the calculated ratios can induce only small errors in the proffered equilibrium constants.

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From the dependence of K on temperature, we obtained ΔH° = −7.1 kcal mol−1, Δ = −10.2 eu and Δ = −4.1 kcal mol−1 [22]. Comparison of K and Δ for Eq. (9) with analogous data for the halocarbene–halocarbanion equilibria, Eqs. (7) and (8), reveals that the π-complex formation of Eq. (9) is significantly more thermodynamically favorable than the carbanion formation. In particular, comparison of data for Eqs. (9) and (8) indicates that conversion of PhCCl to the π-complex with TMB is favored by 3.4 kcal mol−1 in Δ, relative to carbanion formation with Cl.

Given that π-complex formation between PhCCl and TMB is largely driven by charge transfer from the electron-rich TMB to the electrophilic PhCCl, we would expect that increasing the electrophilicity of the carbene would enhance complex formation and increase K. This is indeed the case as demonstrated by the TMB complexation of pentafluorophenylchlorocarbene, F5-PhCCl, Eq. (13) [25].

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The pentafluorophenyl group is more electron withdrawing than the phenyl group as manifested by σI or σR+ values ([26-28], M. Charton, personal communication), and F5-PhCCl should therefore bind more strongly to TMB than PhCCl.

From absorptions of F5-PhCCl at 300 nm and its F5-PhCCl/TMB π-complex at 516 nm, we derived K = 3.21 × 105 m−1 at 294 K, representing an increase of 250-fold over K for the analogous PhCCl equilibrim of Eq. (9) [25]. A study of the temperature dependence of K for Eq. (13) afforded ΔH° = −10.2 kcal mol−1, ΔS° = −9.5 eu and ΔG° = −7.4 kcal mol−1. Comparison with the thermodynamic parameters of Eq. (9) shows that π-complexation of F5-PhCCl by TMB is favored by ~3 kcal mol−1 in ΔG° and ~2 kcal mol−1 in ΔH°, relative to π-complexation of PhCCl [24, 25]. Note that complex formation here is between an excellent donor, trimethoxybenzene, and an unusually electron-poor acceptor, F5-PhCCl, comprised of a perfluorinated benzene ring and a singlet carbene with a formally vacant p-orbital on the carbenic carbon.

A more general demonstration of the manipulation of K by modulating the electronic character of the phenyl ring of PhCCl can be obtained from a Hammett study of the substituent dependence of the equilibrium constant of Eq. (9) [28]. We examined the π-complexation equilibria of p-X-PhCCl with TMB, where X = NO2, CF3, Cl, H, Me and MeO. K ranged from 134 000 m−1 for X = NO2 to 164 m−1 for MeO (cf. Table 2).

Table 2. Equilibrium constants for carbene–π-complexation with TMB*
X in p-X-PhCClK (m−1)
  1. *From ref. [28], except for F5-PhCCl, ref. [25]; This value differs from K = 1260 m−1 (above and ref. [24]) because of a change in f values; Pentafluorophenylchlorocarbene.

NO2134 000
CF325 800
Cl5400
H1450
Me1080
MeO164
F5-PhCCl321 000

We obtained a good Hammett correlation of log K with σp, with ρ = +2.48 [28]. As in dihalomethide carbanion formation, Eq. (10), where ρ = +3.18, the positive sign indicates that complexation is favored by electron-withdrawing substituents on phenyl. These substituents destabilize the carbene, but enhance charge transfer in the carbene–TMB complex [28]. Donating substituents (Me, MeO) result in low values of K. Note that K for F5-PhCCl exceeds K for p-O2N-PhCCl (Table 2); pentafluorophenyl is a strong electron-withdrawing moiety [25].

Carbene-O-Ylide Equilibria

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies

p-Nitrophenylchlorocarbene (PNPCC) forms spectroscopically well-characterized O-ylides with simple ethers, e.g. diethyl ether, Eq. (14) [18].

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    Structure 3. 

PNPCC also forms an O-ylide with THF [29, 30]. We were able to demonstrate the reversibility of O-ylide formation with diethyl ether (Et2O), di-n-propyl ether (Pr2O) and THF, and determine the associated equilibrium constants and thermodynamic parameters [31].

Consider PNPCC and THF as an example. In heptane, the carbene displays a strong π[RIGHTWARDS ARROW]p absorption at 308 nm and a weaker σ[RIGHTWARDS ARROW]p signal at 628 nm. In the presence of added THF (in heptane), an absorbance of the carbene O-ylide appears at 452 nm (computed at 456 nm) [31]. From the dependence of the intensities of these absorptions on the concentration of THF, K = 7.5 m−1 at 295 K for Eq. (14) with R2O = THF [31]. Analogous studies with Et2O and Pr2O gave K = 0.14 m−1 and 0.10 m−1, respectively. Note that K is reduced for the open-chain Et2O and Pr2O vs the cyclic THF. Perhaps steric problems associated with the “floppy” alkyl groups of the open-chain ethers reduce the stability of their ylides compared with the THF ylide.

Determinations of K as a function of temperature afforded the thermodynamic parameters recorded in Table 3 [31]. Clearly, O-ylide formation is enthalpically favorable for the reactions of PNPCC with THF and Pr2O. However, the entropies of formation are quite negative at ca −30 eu, and significantly impact Δ. Thus, K is >1 for THF, but <1 for Pr2O.

Table 3. Equilibria of PNPCC and alkyl ethers, Eq. (14)
EtherK (m−1)ΔHo (kcal mol−1)ΔS° (eu)ΔG° (kcal mol−1)*
  1. *At 295 K; Not determined.

THF7.5−11−32−1.2
Et2O0.14
Pr2O0.10−6.5−271.3

As one might expect for an equilibrium like Eq. (14), where uncharged reactants afford a zwitterionic product, the equilibrium constant varies with solvent polarity. For example, we examined O-ylide formation between PNPCC and THF or tetrahydropyran (THP) [32].

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    Structure 4. 

In each case, K was obtained for the equilibrium formation of the respective O-ylide using the LFP-spectroscopic methodology described above [31]. The dependence of K on the polarities of various pentane/1,2-dichloroethane solvent blends was then determined. Thus, with THF and THP, K varied from 4.4 to 153 m−1 and from 2.1 to 99 m−1, respectively, as the solvent's dielectric constant increased from 1.84 (pure pentane) to 5.4–5.5 (0.42–0.43 mole fraction of DCE in pentane).

Conclusion

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies

LFP-spectroscopy makes possible the direct visualization and determination of halocarbene–halocarbanion, carbene–carbene complex and carbene–O-ylide equilibria. Methodology and results are reviewed for (e.g.) equilibria between phenylhalocarbenes and phenyldihalomethide carbanions, phenylhalocarbenes and phenylhalocarbene–trimethoxybenzene π-complexes and PNPCC and p-nitrophenylchlorocarbene–O-ylides. The temperature dependence of the determined equilibrium constants permits the extraction of the associated thermodynamic parameters.

References

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies
  • 1
    Hine, J. (1950) Carbon dichloride as an intermediate in the basic hydrolysis of chloroform. A mechanism for substitution reactions at a saturated carbon atom. J. Am. Chem. Soc. 72, 24382445.
  • 2
    Hine, J. and A. M. Dowell Jr (1954) Carbon Dihalides as intermediates in the basic hydrolysis of haloforms. III. Combination of carbon dichloride with halide ions. J. Am. Chem. Soc. 76, 26882692.
  • 3
    Moss, R. A., J. Tian, R. R. Sauers, D. H. Ess, K. N. Houk and K. Krogh-Jespersen (2007) The synthesis of dichlorodiazirine and the generation of dichlorocarbene: spectroscopy and structure of dichlorocarbene ylides. J. Am. Chem. Soc. 129, 51675174.
  • 4
    Moss, R. A., M. Zhang and K. Krogh-Jespersen (2009) The trichloromethide and bromodichloromethide carbanions. Org. Lett. 11, 57025704.
  • 5
    Jacox, M. E. and D. E. Milligan (1970) Matrix-isolation study of the vacuum-ultraviolet photolysis of methyl chloride and methylene chloride. Infrared and ultraviolet spectra of the free radicals CCl, H2CCl, and CCl2. J. Chem. Phys. 53, 26882701.
  • 6
    Milligan, D. E. and M. E. Jacox (1967) Matrix-isolation study of the reaction of carbon atoms with chlorine. The electronic and vibrational spectra of the free radical CCl2. J. Chem. Phys. 47, 703707.
  • 7
    Chateauneuf, J. E., R. P. Johnson and M. M. Kirchoff (1990) Absolute kinetics of dichlorocarbene in solution. J. Am. Chem. Soc. 112, 32173218.
  • 8
    Moss, R. A., J. Tian, R. R. Sauers, C. Skalit and K. Krogh-Jespersen (2007) Chlorofluorocarbene: first UV observation of a dihalocarbene in solution. Org. Lett. 9, 40534056.
  • 9
    Moss, R. A., M. Zhang and K. Krogh-Jespersen (2009) Latent nucleophilicity of dichlorocarbene. Org. Lett. 11, 19471950.
  • 10
    Pliego, J. R. Jr, W. B. De Almeida, S. Çelebi, Z. Zhu and M. S. Platz (1999) Singlet-triplet gap and the electronic vibrational spectra of chlorophenylcarbene: a combined, theoretical, and experimental, study. J. Phys. Chem. A 103, 74817486.
  • 11
    Moss, R. A. and J. Tian (2005) Concurrent cyclopropanation by carbenes and carbanions. J. Am. Chem. Soc. 127, 89608961.
  • 12
    Moss, R. A. (2002) Carbenic philicity. In Carbene Chemistry: From Fleeting Intermediates to Powerful Reagents. (Edited by G. Bertrand), pp. 57101. Dekker, New York City, NY.
  • 13
    Moss, R. A. and J. Tian (2006) Modulation of selectivity in a fluorocarbene cyclopropanation reaction: a catalytic role for bromide ion. Tetrahedron Lett. 47, 34193421.
  • 14
    Wang, L., R. A. Moss and K. Krogh-Jespersen (2012) Directly observed halocarbene – halocarbanion equilibration. J. Am. Chem. Soc. 134, 1745917461.
  • 15
    Wang, L., R. A. Moss and K. Krogh-Jespersen (2013) Hammett analyses of halocarbene – halocarbanion equilibria. Org. Lett. 15, 20142017.
  • 16
    Moss, R. A. and J. Tian (2006) Hammett studies of aryldichloromethide carbanion reactions. Org. Lett. 8, 12451247.
  • 17
    Moss, R. A., J. Tian, R. R. Sauers and K. Krogh-Jespersen (2007) Tracking invisible alkylchlorocarbenes by their σ[RIGHTWARDS ARROW]p Absorptions: dynamics and solvent interactions. J. Am. Chem. Soc. 129, 1001910028.
  • 18
    Moss, R. A., L. Wang, E. Weintraub and K. Krogh-Jespersen (2008) The solvation of carbenes: π and O-ylidic complexes of p-nitrophenylchlorocarbene. J. Phys. Chem. A 112, 46514659.
  • 19
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  • 20
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  • 21
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  • 22
    Turro, N. J., G. F. Lehr, J. A. Butcher Jr, R. A. Moss and W. Guo (1982) Temperature dependence of the cycloaddition of phenylchlorocarbene to alkenes. Observation of “Negative Activation Energies.”. J. Am. Chem. Soc. 104, 17541756.
  • 23
    Houk, K. N., N. G. Rondan and J. Mareda (1985) Theoretical studies of halocarbene cycloaddition selectivities A new interpretation of negative activation energies and entropy control of selectivity. Tetrahedron 41, 15551563.
  • 24
    Moss, R. A., L. Wang, C. M. Odorisio and K. Krogh-Jespersen (2010) A Carbene – carbene complex equilibrium. J. Am. Chem. Soc. 132, 1067710679.
  • 25
    Wang, L., R. A. Moss and K. Krogh-Jespersen (2011) Trimethoxybenzene complexes of pentafluorophenylchlorocarbene. J. Phys. Chem. A 115, 81138118.
  • 26
    Charton, M. (1981) Electrical effect substituent constants for correlation analysis. Prog. Phys. Org. Chem. 13, 119251; cf. pp. 143 and 149.
  • 27
    Ehrenson, S., R. T. C. Brownlee and R. W. Taft (1973) A generalized treatment of substituent effects in the benzene series. A statistical analysis by the dual substituent parameter equation (1). Prog. Phys. Org. Chem. 10, 180.; cf. p. 13.
  • 28
    Wang, L., R. A. Moss, J. Thompson and K. Krogh-Jespersen (2011) Hammett analysis of a family of carbene – carbene complex equilibria. Org. Lett. 13, 11981201.
  • 29
    Çelebi, S., M.-L. Tsao and M. S. Platz (2001) Laser flash photolysis studies of the reaction of arylhalocarbenes with tetramethylethylene as a function of solvent. J. Phys. Chem. A 105, 11581162.
  • 30
    Tsao, M.-L., Z. Zhu and M. S. Platz (2001) Matrix and time-resolved infrared spectroscopy of chloro-p-nitrophenylcarbene and related species. J. Phys. Chem. A 105, 84138416.
  • 31
    Hoijemberg, P. A., R. A. Moss and K. Krogh-Jespersen (2012) Reversible O-ylide formation in carbene/ether reactions. J. Phys. Chem. A 116, 358363.
  • 32
    Hoijemberg, P. A., R. A. Moss and K. Krogh-Jespersen (2012) Solvent polarity effects on carbene/ether – O-ylide equilibria. J. Phys. Chem. A 116, 47454750.

Biographies

  1. Top of page
  2. Abstract
  3. Carbene–Carbanion Equilibria
  4. Carbene–Carbene Complex Equilibria
  5. Carbene-O-Ylide Equilibria
  6. Conclusion
  7. Acknowledgements
  8. References
  9. Biographies
  • Image of creator

    Robert A. Moss is Research Professor and Louis P. Hammett Professor Emeritus at Rutgers University, The State University of New Jersey. A 1960 graduate of Brooklyn College, he received his doctorate from the University of Chicago in 1963 under the direction of Professor G.L. Closs. After a NAS-NRC Postdoctoral Fellowship with Professor Ronald Breslow at Columbia University, he joined the faculty at Rutgers University in 1964. He has remained at Rutgers with the exception of research leaves at M.I.T., the University of Oxford, the Weizmann Institute of Science, and the Hebrew University of Jerusalem. Professor Moss and his colleagues have authored more than 400 scientific publications, including eight coedited books. These contributions have dealt with reactive intermediates, reactions in organic aggregates, and the decontamination of toxic organic phosphorous compounds. In 2010 he received an Arthur C. Cope Scholar Award from the American Chemical Society.

  • Image of creator

    Lei Wang was born in 1978 in Nanjing, P. R. China. He gained his Ph.D. in 2005 from Nanjing University, China, working on photoinduced cycloaddition reactions of heterocyclic compounds. Since 2006, he is a Postdoctoral Research Fellow at the Rutgers University with Professor Robert A. Moss. He is engaged in methodological and mechanistic studies of reactive intermediates, especially carbenes.

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    Pablo A. Hoijemberg received his title of Licenciado in Chemical Sciences and then his Ph.D. in Chemistry from the Universidad de Buenos Aires (Buenos Aires, Argentina) in 2003 and 2009, respectively, working on the topic “Cage effect in supercritical fluids.” After that he moved for a year (2009–2010) to Mulhouse, France, as a Postdoctoral Research Fellow at the Department of General Photochemistry (now LPIM), at the Université de Haute Alsace, performing research on “Photopolymerization in miniemulsions.” Then, since 2011, he joined the Department of Chemistry and Chemical Biology at Rutgers University (New Jersey, USA) as a Postdoctoral Research Associate, to work on Carbene chemistry, specifically on the formation of complexes and ylides with ethers, crown compounds and other substrates.

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    Karsten Krogh-Jespersen is currently a Professor in the Department of Chemistry and Chemical Biology at Rutgers University. Danish by birth, he came to the United States in 1972 for graduate studies in chemistry at New York University. After receiving his Ph.D. in 1976 (with Mark A. Ratner), he did postdoctoral research at the University of Erlangen-Nurnberg (with Paul. V. R. Schleyer) and at Rutgers University (with Lionel Goodman). He joined the faculty at Rutgers University in 1979. His research interests focus on computational investigations of molecular electronic structure, in particular in areas of physical inorganic and organic chemistry.