Counter-Intuitive Gas-Phase Reactivities of [V2 ]+ and [V2 O]+ towards CO2 Reduction: Insight from Electronic Structure Calculations.

[V2 O]+ remains "invisible" in the thermal gas-phase reaction of bare [V2 ]+ with CO2 giving rise to [V2 O2 ]+ ; this is because the [V2 O]+ intermediate is being consumed more than 230 times faster than it is generated. However, the fleeting existence of [V2 O]+ and its involvement in the [V2 ]+ → [V2 O2 ]+ chemistry are demonstrated by a cross-over labeling experiment with a 1:1 mixture of C16 O2 /C18 O2 , generating the product ions [V2 16 O2 ]+ , [V2 16 O18 O]+ , and [V2 18 O2 ]+ in a 1:2:1 ratio. Density functional theory (DFT) calculations help to understand the remarkable and unexpected reactivity differences of [V2 ]+ versus [V2 O]+ towards CO2 .

Rolf Huisgen has shaped physical organic chemistry in a nearly unparalleled fashion.1 His extensive, seminal work on, for example, 1,3‐dipolar cycloadditions1b forms a pillar of contemporary mechanistic organic chemistry, not to speak of their role in the synthesis of heterocyclic compounds. By using the then available experimental repertoire, the Huisgen school unraveled the finest mechanistic details of numerous chemical processes, and their work gives testament for chemist generations to come.2

A familiar kinetic scheme of a multi‐step process is depicted in Equation 1, and one problem here concerns the identification of “intermediate” [B]. This holds true in particular when the rate constants k 1 and k 2 differ significantly, that is, k 2 ≫ k 1.

A typical example is provided by the gas‐phase chemistry of the [Pd]+/CH3I couple which gives rise to the formation of [Pd(CH2I)]+.3 In this case, the product ion is not formed by the entropically favored direct cleavage of the C−H bond of CH3I.3a Rather, in the first step a short lived [Pd(CH3)]+ intermediate is generated, which reacts further with CH3I in a very fast metathesis process to form [Pd(CH2I)]+ and CH4.3b Recently, in the context of combined experimental/computational studies on the timely topic of CO2 activation,4 we came across a similar, unexpected situation in the thermal gas‐phase reactions of bare diatomic [V2]+ with CO2.

The cluster [V2]+ was formed by supersonic expansion of helium into a vanadium plasma generated by laser ablation/ionization of a vanadium disk using a Nd:YAG laser, operating at 532 nm inside the external cluster source of a Fourier‐transform ion cyclotron resonance (FT‐ICR) mass spectrometer as described previously (for details, see the Supporting Information).5 A fraction of the ion population was then guided by a static ion optical system into the ICR cell. Next, in a sequence of pulses, argon was admitted to the ICR cell such that the ions collide on average about 1×105 times with argon. This procedure ensures thermalization of hot ions and quenching of excited electronic states. After mass‐selection of the thermalized [V2]+ ions, they were reacted with CO2 or C18O2 at a constant pressure low enough to ensure single collision conditions. The elementary compositions of the charged species have been confirmed by exact mass measurements using high resolution mass spectrometry.

The result of the reaction of [V2]+ with CO2 at a partial pressure of 1.0×10−7 mbar and a reaction time of 10 s inside the ion trap of the FT‐ICR machine is shown in Figure 1 b. In addition to the signal of the starting material [V2]+ (signal A, Figure 1), a signal D appears, which has been assigned the formula [V2O2]+ by exact mass measurements; also, very small amounts of atomic [V]+ are present (signal B). Formally, the formation of [V2O2]+ corresponds to the loss of one carbon atom. If C18O2 acts as a reaction partner, signal D from Figure 1 b is shifted by four mass units (Figure 1 c; signal E). Clearly, oxygen atom transfer (OAT) from CO2 to the vanadium cluster takes place. However, since a carbon atom represents an extremely poor leaving group, the suggestion was raised as to whether the formation of [V2O2]+ from the [V2]+/CO2 couple is the result of a multi‐step process, although the spectra in Figure 1 b,c do not exhibit any clue for the involvement of conceivable intermediates. The assumption for the operation of a multi‐step process is corroborated by a crossover‐labeling experiment using a 1:1 mixture of C16O2 and C18O2: The appearance of [V2 16O18O]+ (signal F in Figure 1 d) together with the other two isotopomers in a 1:2:1 ratio verifies this hypothesis. According to Equations (2) and (3) and in line with Figure 1 d, one has to conclude that [V2]+ reacts consecutively with two CO2 molecules via an OAT by the release of one CO molecule each time; the monoxide [V2O]+ serves as a short‐lived intermediate (Scheme 1). Further oxidation processes according to Equation (4) in Scheme 1 do not occur at ambient temperature.

Figure 1

Representative mass spectra: a)–e) are for the thermal reactions of [V2]+ with a) Ar at 1.0×10−7 mbar, b) C16O2 at 1.0×10−7 mbar, c) C18O2 at 1.0×10−7 mbar, and d) a 1:1 mixture of C16O2 and C18O2 at 2.0×10−7 mbar after a reaction time of 10 s, respectively; e) C18O2 at 1.0×10−7 mbar and a reaction time of 1s; f) and g) represent the thermal reactions of [V2O]+ with f) Ar at 4.0×10−9 mbar, g) C18O2 at 4.0×10−9 mbar, after a reaction time of 2s, respectively. All x‐axes are scaled in m/z, and the y‐axes are normalized relative ion abundances.

Scheme 1

Multi‐step activation of CO2 by [V2]+ with [V2O]+ serving as a fleeting intermediate.

In the [V2]+/CO2 couple, the supposed intermediate [V2O]+ can only be detected experimentally as an extremely weak signal under highly optimized conditions of both pressure and reaction time (signal G in Figure 1 e). However and fortunately, [V2O]+ can be synthesized independently and in sufficient quantities in the external ion source by adding traces of O2 to the helium buffer gas (Figure 1 f). As shown in Figure 1 g, the so‐formed [V2 16O]+ indeed reacts with C18O2 to produce [V2 16O18O]+ by the expulsion of C18O. The collision efficiency6 φ of that reaction is close to 100 %. The high reactivity of [V2O]+ explains why, [V2O]+ from the reaction of [V2]+ with CO2, remains elusive. As soon as the monoxide has formed within the ICR cell, it immediately reacts with another CO2 molecule. The high reactivity of the [V2O]+ intermediate, generated from the [V2]+/CO2 couple, is further enhanced due to the substantial exothermicity of the process (see below) and the fact that, owing to the absence of a heat bath, part if not most of the excess energy is stored in the ion by partitioning the energy released between [V2O]+ and CO. To distinguish possible reactions of the charged particles with omnipresent background gases from those of the selected substrate, spectra with argon were recorded as well. As shown in Figure 1 a,f side reactions with background gases like trace amounts of oxygen are limited to the generation of small quantities of products typical for reactions of [V2]+ with O2,9e which are more pronounced in [V2O]+, since O2 has been added to the He buffer gas.

The rate constants k 4([V2]+/CO2) and k 5([V2O]+/CO2) for reactions shown in Equations (2) and (3) in Scheme 1 were measured and they amount to 3.8×10−12 (efficiency φ=0.5 %) and 8.8×10−10 cm3 molecule−1 s−1 (φ=1.15), respectively; these numbers translate to a ratio of k 5/k 4>230. Owing to an uncertainty in the determination of the absolute pressure of CO2, an error of ±30 % is associated with the rate measurements.5b For the relative rate constants, however, the error is much smaller, typically around ±5 %.

The considerable reactivity difference of [V2]+ versus [V2O]+ is unexpected and counter‐intuitive on the ground that the reactions in Equations (2) and (3) constitute redox processes in which electrons are transferred from the metal‐containing cluster ions to the CO2 molecule. Therefore, one would expect that [V2]+ reduces CO2 to CO more readily than the already oxidized and electron‐depleted [V2O]+. Consequently, k 4 should be larger than k 5; this is not the case. Why does the reactivity of [V2]+ and [V2O]+ not reflect their respective oxidation states? Further, does [V2O]+ react mechanistically the same way with CO2 as [V2]+? These questions have been addressed by density functional theory (DFT) calculations.

Earlier theoretical work had already shown that a description of vanadium dimers constitutes a notoriously difficult case to be dealt with by DFT calculations.7 This dilemma holds true as well for [V2O2]+.8 While for [V2O2]+ the gas‐phase IR spectrum suggests a planar, four‐membered V‐O‐V‐O ring structure of C s symmetry,8b an unambiguous assignment of the electronic ground state (2A′ or 6A′) remained unresolved at the DFT levels employed.8b, 8c The problems are compounded by the observation that the gas‐phase chemistry of vanadium ions9 is often marred by a two‐ or multi‐state reactivity scenario.10 As to the bare [V2]+ cluster, due to its multireference character7, 11 we applied the ZORA‐NEVPT2(9e,12o)/ZORA‐def2‐QZVPP Scheme (Table S1 in the Supporting Information).12 According to these high‐level calculations, the ground state of [V2]+ corresponds to 4Σg −. This assignment agrees well with experimental data and other, less sophisticated calculations.13 The lowest excited state of [V2]+ is a doublet (2Δ) and exceeds the ground state by only about 3 kJ mol−1 (Table S2). Therefore, it is very likely that under the given experimental conditions a mixture of two electronic states, 2,4[V2]+, is present, in which the quartet may dominate. We mention in passing that low‐energy excited states of the neutral V2 dimer were identified by Hübner and Himmel using matrix isolation spectroscopy.11 To construct a simplified potential energy surface (PES) for the [V2]+/[V2O]+/CO2 systems, extensive quantum chemical calculations were performed at the ZORA‐M06L/ZORA‐def2‐QZVPP//M06L/def2‐TZVP level of theory. Electronic structure calculations using the multi‐reference ZORA‐NEVPT2 scheme are not feasible technically. Rather, the M06L functional was chosen based on the exhaustive testing carried out by Truhlar and co‐workers.7 This functional performed best also among all the density functionals we benchmarked (see Tables S2, S6 and S7). In addition, M06L delivers fairly good results for the electronic states of [V2O]+ and [V2O2]+ (see Tables S3 and S4).

Figure 2 displays the PESs of the most favorable pathways for the reduction of CO2 with vanadium ions having doublet spin multiplicity as well as selected structural parameters of stationary points along the reaction coordinate. This particular spin state was chosen on the ground that along the whole transformation [V2]+ → [V2O2]+ the spin multiplicity was conserved. It should be mentioned at this point, however, that because of the strongly correlated electrons, the energetic and geometric properties of some stationary points with different spin multiplicities are located fairly close to each other. Therefore, it is quite possible that there is a superposition of the doublet with the quartet state and vice versa, facilitating a multi‐state reactivity scenario pending on the efficiency of spin‐orbit coupling (see below for a further comment).10 However, our aim is not to provide a quantitatively correct description of the whole PESs; rather, we aim at offering a qualitative understanding of the amazing reactivity behavior in the [V2]+/[V2O]+/CO2 systems. As the reactivity pattern proved to be quite complex, herein we limit ourselves to the main results. For further details, see Figures S2 and S3.

Figure 2

Schematic potential energy surfaces (ΔH 298K in kJ mol−1) obtained at the ZORA‐M06L/ZORA‐def2‐QZVPP//M06L/def2‐TZVP level of theory for the reactions of a) 2[V2]+ and b) 2[V2O]+ with CO2, respectively. Key structures with selected geometric parameters and symmetries are also provided. Bond lengths [Å]; angles [°]. The relative energies of 2,4[V2]+ were obtained at the ZORA‐NEVPT2(9e,12o)/ZORA‐def2‐QZVPP level of theory. For details, see the Supporting Information.

As displayed in Figure 2 a, at the doublet surface the exothermic OAT14 in the system [V2]+/CO2 commences with the barrier‐free formation of an encounter complex 1 a by coordinating the incoming ligand to only one of the two vanadium atoms in an end‐on (η 1) mode. Subsequently, intermediate 2 a is generated by surmounting transition state TS1 a, which is accompanied by a substantial elongation of the V−V bond from about 1.6 Å to 2.5 Å; this bond length does not change that much further in the remaining steps. In contrast, significant structural reorganization happens for CO2 at these stages. The initially linear substrate gets bent with an O‐C‐O angle of 141° in 2 a, and the C−O bond length increases from 1.17 to 1.27 Å; clearly, CO2 is already markedly activated in this initial step. Next, the complete scission of one C−O bond of the CO2 moiety occurs along the route 2 a → TS2 a → 3 a. The reaction is completed by the release of CO to generate 2[V2O]+. The energetically submerged transition state TS1 a, as part of the reduction sequence of the first CO2 molecule, is with 29 kJ mol−1 located below the entrance asymptote and represents the rate‐limiting step of the entire reaction.

As to the exothermic reaction of [V2O]+ with CO2 and as displayed in Figure 2 b, 2[V2O]+ is approached by carbon dioxide as well, to then also form an η 1 encounter complex 1 b; subsequently, this complex has sufficient internal energy to easily surmount transition state TS1 b and generate 2 b. Next, the complete cleavage of the activated C−O bond takes place along the sequence 2 b → TS2 b → 3 b, and release of CO from 3 b provides 2[V2O2]+. In contrast to the [V2]+/CO2 couple, here it is transition state TS2 b (−48 kJ mol−1) in which the C−O bond of the CO2 molecule is broken and which represents the rate‐limiting step in the reduction of CO2 to CO.

At first glance, the reductions of the two CO2 molecules by [V2]+ and [V2O]+ do not seem to differ too much in terms of structural features of the intermediates. So, why do they exhibit quite pronounced reactivity differences? As will be shown, these features can be attributed to both intrinsic structural and electronic properties. As suggested by a reviewer, one reason may be due to the spin situation. For the [V2O]+/CO2 couple the OAT is both exothermic and spin‐allowed (Figure 2 b); the whole reaction is confined to the doublet surface. This must not necessarily be the case for the [V2]+/CO2 system. While for electronically excited 2[V2]+ the reaction with CO2 is also both exothermic and spin‐allowed (Figure 2 a), for the ground state 4[V2]+ a quartet → doublet spin change has to occur along the overall exothermic reaction coordinate (Table S5). If this electronic reorganization is inefficient, as has been reported for the analogous exothermic [V]+/CO2 conversion,9f as well as for the [V]+/N2O couple,9a the reaction of [V2]+ with CO2 will be slowed down as well. Further, while thermochemical arguments on the endo‐ versus exothermicity of a reaction,7 based only on DFT calculations, may be considered with caution, the insensitivity of the [V2]+/CO2 reactivity to intentional excitation of the [V2]+ projectile, strongly suggests that process (2) is exothermic and that its much reduced reactivity is not due to a thermodynamic impediment.

Indeed, there are other factors to be considered to explain the much higher reactivity of the [V2O]+/CO2 couple. Let us begin by looking at the charge distribution in the ionic actors. According to Figure 3 a, the reactive vanadium atoms in 2[V2O]+ carry more positive charge as compared to 2[V2]+. It is therefore expected that [V2O]+ will act as a worse electron donor for the reduction of CO2 compared to [V2]+.15 However, this is not the whole truth. In fact, the presence of the bridging oxygen atom in [V2O]+ actually plays a supportive role in the CO2 activation which overcompensates the unfavorable charge situation.

Figure 3

a) NBO‐calculated charges at the vanadium atoms coordinated to CO2 and b) development of the V−V bond length [Å] in the reactions of 2[V2]+ and 2[V2O]+ with CO2.

First, the presence of the oxygen atom induces a structural rearrangement of the active site, thus, providing a “preorganized” structure, which reduces substantially the energy demand to reach transition state TS1 b; this effect is well‐known from enzymatic processes.16 The calculated Wiberg Bond Index (WBI)17 of 4.5 indicates a rather strong and short V−V bond for [V2]+ (1.63 Å); a much smaller value of the WBI of 1.2 and an elongated V−V bond of 2.53 Å holds true for 2 a (Figure 2 and Figure 3 b). As a consequence, stretching of the V−V bond in [V2]+ in the step 1 a → 2 a is quite energy‐demanding (Figure 2 a). In contrast, the corresponding transformation 1 b to 2 b in [V2O]+ requires much less energy as the V−V bond is already sufficiently elongated due to the presence of the oxygen bridge. Thus, [V2O]+ possesses a “prepared” structure which greatly enhances the reduction of CO2. In the remaining stages of the reaction there is no further dramatic change in the properties of the V−V bond, as indicated by the rather small oscillations of the WBI's around a value of 1. The moderate movements of the two vanadium atoms relative to each other do not require that much energy; rather, it is this structural change of the “active” site which assists in lowering TS1 b as compared with TS1 a.

There is yet another aspect to be considered. In 2[V2O]+ the presence of an oxygen atom raises the energy level of the electron‐donating d‐orbitals. As shown in Figure 4, in TS1 a of [V2]+ the overlap between the lobes of the doubly‐occupied πd/d‐orbital of the metal center and the vacant π*‐orbital of the carbon atom of CO2 has yet to be established. In contrast, for the early transition state TS1 b, the doubly‐occupied πd/d‐orbital is already delocalized quite extensively to the carbon atom of CO2, and the root cause for this difference can be traced back to orbital energy differences. The introduction of an oxygen bridge raises the πd/d‐orbital energy from −10.5 eV in [V2]+ to −9.5 eV in [V2O]+; this greatly eases the transfer of the d‐electrons into the empty π*‐orbital of CO2 thus favoring the reduction of [V2O]+.

Figure 4

The frontier orbitals of the transition states of TS1 a,b from the doublet surface of the CO2 activation. Arrows and rounded rectangles indicate the overlap regions between the doubly occupied π‐electron‐donating orbital of the metals and the lowest empty orbital of CO2.

As shown repeatedly,16, 18 the presence or absence of “prepared states” are crucial in controlling the structural, electronic and energetic features of transition states. While there exist quite a number of examples for this scenario in the area of gas‐phase C−H bond activation,19 to our knowledge the present study describes for the first time this effect in the context of CO2 activation.

In conclusion, in the present experimental/computational work we report and explain quite a few unexpected observations:

A fleeting [V2O]+ intermediate is involved in the multistep oxidation of [V2]+ by CO2.

The counter‐intuitive oxidation behavior of [V2]+ versus [V2O]+ and the quite enormous reactivity differences of the two cluster ions have been traced back to the supportive role of the oxygen bridge in [V2O]+. This cluster oxide, in contrast to bare [V2]+, exhibits structural and electronic features which aid in the activation of CO2. While thermochemical aspects are not likely to cause the relative “inertness” of [V2]+, an inefficient quartet → doublet spin change in the [V2]+/CO2 couple may well matter in slowing down its reactivity.

Conflict of interest

The authors declare no conflict of interest.

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Supplementary

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