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. 2018 Dec 5;123(1):73–81. doi: 10.1021/acs.jpca.8b10530

CO2/O2 Exchange in Magnesium–Water Clusters Mg+(H2O)n

Erik Barwa 1, Milan Ončák 1,*, Tobias F Pascher 1, Thomas Taxer 1, Christian van der Linde 1, Martin K Beyer 1,*
PMCID: PMC6331139  PMID: 30516989

Abstract

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Hydrated singly charged metal ions doped with carbon dioxide, Mg2+(CO2)−(H2O)n, in the gas phase are valuable model systems for the electrochemical activation of CO2. Here, we study these systems by Fourier transform ion cyclotron resonance (FT–ICR) mass spectrometry combined with ab initio calculations. We show that the exchange reaction of CO2 with O2 proceeds fast with bare Mg+(CO2), with a rate coefficient kabs = 1.2 × 10–10 cm3 s–1, while hydrated species exhibit a lower rate in the range of kabs = (1.2–2.4) × 10–11 cm3 s–1 for this strongly exothermic reaction. Water makes the exchange reaction more exothermic but, at the same time, considerably slower. The results are rationalized with a need for proper orientation of the reactants in the hydrated system, with formation of a Mg2+(CO4)−(H2O)n intermediate while the activation energy is negligible. According to our nanocalorimetric analysis, the exchange reaction of the hydrated ion is exothermic by −1.7 ± 0.5 eV, in agreement with quantum chemical calculations.

Introduction

Because of the still increasing consumption of fossil fuels, carbon dioxide is one of the most problematic greenhouse gases produced by humankind. As CO2 is highly thermodynamically stable, it cannot be easily activated in chemical reactions, and its utilization is very limited.1 For the transformation of CO2 to fuels, activation is usually achieved under high temperature and pressure conditions over heterogeneous catalysts in the Sabatier process.2 A promising alternative is electrochemical activation of CO2, and the formation of formic acid in electrochemical cells has been reported as early as in 1870.3 A key intermediate is the carbon dioxide radical anion CO2–, which has attracted growing attention in gas phase studies since Compton and Klots reported its stabilization by solvation.4−6 Carbon dioxide activation in the gas phase was recently reviewed by Weber7,8 and Schwarz.9

Photodissociation of the C–O bond in CO2–(H2O)n has been reported by Sanov and co-workers.10,11 In aqueous solution, spectra of CO2– in the UV7,12,13 have been measured, and the symmetric stretching and bending modes have recently been identified by Raman spectroscopy.14 In gas phase clusters, infrared spectra of CO2–(H2O)n have been obtained with up to two water molecules in the O–H stretch region.7,15 Reactions of hydrated electrons with CO2 directly revealed the process of reductive activation, resulting in the formation of CO2–(H2O)n.16−21 C–H, C–C, and C–S bond formation was observed with a series of reactants.22−26 Uggerud and co-workers have shown in elegant studies that ClMgCO2– complexes can be formed by collision induced dissociation of the oxalic acid complex, and studied the reactivity of reductively activated CO2.27−30 Weber and co-workers investigated reductive CO2 activation in M–(CO2)n systems by infrared spectroscopy.7,8,31 Menges et al. demonstrated the capture of CO2 by a cationic Ni(I) complex and characterized of the activated CO2 molecule by cryogenic infrared spectroscopy.32 The Johnson group recently also characterized radical ion adducts between imidazole and CO2 by vibrational spectroscopy.33

In the present work, we are interested in the influence of metal centers on the reactivity of CO2– in water clusters, choosing magnesium as a well-investigated metal. Magnesium has also atmospheric relevance as roughly 5 tons of Mg as interplanetary dust enters earth’s atmosphere every day.34 By photoionization or charge transfer reactions with NO+ and O2+, Mg+ is formed, which further reacts with its surroundings in the mesosphere and lower thermosphere, inter alia, with CO2 and O2.35

Bond dissociation energies for Mg+ complexes with a series of small molecules, including CO2, as well as the binding energies of the first four water molecules were determined by collision-induced dissociation (CID) experiments by the Armentrout group.36,37 Williams and co-workers studied doubly charged hydrated magnesium by blackbody infrared radiative dissociation (BIRD).38 Singly charged hydrated magnesium ions Mg+(H2O)n have been examined with respect to the influence of blackbody infrared radiation, photodissociation, and reactions with small molecules by FT-ICR mass spectrometry and methods of theoretical chemistry.39−54

Duncan et al. investigated infrared photodissociation spectroscopy of Mg+(CO2)n and Mg+(CO2)n Ar ion–molecule complexes55 as well as Mg+(H2O)Arn56 and Mg+CO2.57 The reactions of hydrated magnesium cations and first-row transition metal ions with CO2 as well as O2 were recently explored by van der Linde et al.54,58 For M = Mg, Cr and Co, charge transfer from metal centers leads to uptake of CO2 by the positively charged hydrated metal ions, forming M2+(CO2)−(H2O)n with low efficiency in collisions of M+(H2O)n with CO2.54,59 Quantum chemical calculations corroborate that charge transfer occurs, resulting in a doubly charged metal center and a negative CO2•−.53,54 Similar reactions are observed with O2 for M = Mg, Cr, Co, Ni, and Zn.

The calculations also show that Mg2+(CO2)−(H2O)n as well as Mg2+(O2)−(H2O)n may exist as either solvent-separated ion pairs (SSIP) or contact ion pairs (CIP) whereby the SSIP is the energetically more stable configuration for n = 16.54 Since the electron is located in the valence shell of CO2 or O2, the hydrogen loss observed for Mg+(H2O)n clusters, 6 < n < 14,60 does not take place. While CO2 reacts 2 – 3 times faster with hydrated electrons than O2,61 a different behavior is observed in the case of hydrated Mg+: O2 is 4 – 5 times more reactive with [Mg(H2O)n]+ than CO2.54 In the reaction of CO2–(H2O)n ions with O2, CO4– is very likely formed as an intermediate,21 as suggested by Weber.7 Previous theoretical calculations54 predicted that for clusters with 16 water molecules attached, the reaction energy of the O2/CO2 exchange reaction is about −1.86 eV.

To test this prediction experimentally, we investigate the CO2/O2 exchange reaction in Mg2+(CO)2–(H2O)n cluster distributions along with nanocalorimetric analysis20 of clusters n ≤ 70. Quantum chemical calculations are used to map possible reaction pathways for both bare and hydrated clusters, respectively, and to monitor the course of the exchange reaction.

Experimental and Theoretical Methods

The experiments are performed on a modified 4.7 T FT–ICR Bruker/Spectrospin CMS47X mass spectrometer62 equipped with a Bruker infinity cell63 and an external laser vaporization source.64,65 A frequency doubled Nd:YAG laser (Continuum Surelite II) is used to generate Mg2+(CO2)−(H2O)n ions by evaporation of isotopically enriched 24Mg from a solid metal target and supersonic jet expansion of a hot plasma in a helium/water/CO2 gas mixture. Twenty laser shots at 10 Hz and approximately 5 mJ pulse energy are used to generate the ions. The ions are rotationally and vibrationally cooled below room temperature due to the supersonic expansion into high vacuum, accelerated downstream from a skimmer and transferred to the ICR cell by a system of electrostatic lenses through several differential pumping stages.66 In the ICR cell, ions are stored at room temperature in an electromagnetic trap under ultrahigh vacuum conditions in a 4.7 T magnetic field.67 O2 is introduced at constant backing pressure, allowing the monitoring of reaction kinetics by taking mass spectra after different reaction delays. For each mass spectrum, 20 experiment cycles are averaged.

Absolute rate coefficients kabs are obtained by analyzing the pseudo-first-order kinetic plots of different experimental runs taken over a range of pressures. The cluster distribution shrinks due to blackbody infrared radiative dissociation (BIRD)68−76 and the exothermicity of the reaction.58,77,78 The error of the rate coefficient was estimated to be about 30% due to the uncertainty of the pressure calibration.79,80 The noise level of the summed intensities was calculated with the Gaussian law of error propagation from the noise level of each peak. It should be noted that the internal temperature of the clusters is given by the interplay between radiative heating and evaporative cooling. Experiments on phase transitions in water clusters place this temperature in the region of 100–200 K.81,82

The collision rates kADO, kHSA and kSCC are calculated using the average dipole orientation (ADO),83,84 hard sphere average dipole orientation (HSA) and surface charge capture (SCC) models, which yield the efficiencies ΦADO = kabs / kADO, ΦHSA = kabs / kHSA, and ΦSCC = kabs / kSCC, respectively.85 Nanocalorimetric analysis is performed by fitting the average cluster size ⟨n⟩ of reactant and product ions over time with a set of differential equations, which yields the average number of water molecules evaporating due to the heat of the reaction.20,21,61 The evaporation of one H2O molecule removes ΔEvap = 0.45 ± 0.03 eV from the cluster.82,86

Selected clusters were optimized using the M06 density functional theory (DFT) functional87 along with the def2TZVP basis set. As the DFT theory might struggle to quantitatively describe the nature of the Mg+/O2 interaction,88 we recalculated the structures using the complete basis set QB3 (CBS-QB3) method.89 This method is able to reproduce values calculated at the coupled cluster level (CCSD(T)/aug-cc-pVQZ) as already noted elsewhere.88 Molecular volume calculations for the HSA and SCC methods and charge analysis within the ChelpG scheme90 were performed at the MP2/def2TZVP level.

Molecular dynamics was run on the M06/6-31+G* potential energy surface with a time step of 30 au (∼0.75 fs). Investigated [Mg(CO2)(H2O)n]+ clusters were first thermalized at 250 K using the Nosé–Hoover thermostat. Then, an O2 molecule was added at a distance of 10 Å with respect to the cluster center of mass and a microcanonical simulation was performed. Twenty trajectories were run for each structure. A dynamics run was stopped when a neutral molecule (O2 or CO2) leaves the cluster by more than 10 Å or after 7 ps (or 12 ps when CO4– was formed to investigate its dissociation). We considered only runs, either reactive or nonreactive, where O2 approached the cluster with a distance shorter than 3 Å with respect to any cluster atom.

All quantum chemical calculations were performed in the Gaussian program,91 the Abin code was used for molecular dynamics.92 For calculation of unimolecular rate constants, a standard RRKM implementation was used.93

Results and Discussion

We investigate the O2/CO2 exchange reaction in the Mg2+(CO2)−(H2O)n ion:

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Measured rate coefficients of reaction 1 for various average cluster sizes are collected in Table 1, calculated reaction energies for n = 0–7 are shown in Table 2.

Table 1. Measured Rate Coefficients kabs and Efficiencies ΦHSA, ΦSCC, and ΦADO for O2/CO2 Exchange Reaction 1 with Different Average Cluster Size n of the Initial Mg2+(CO2)−(H2O)n Cluster and Their Associated Mean Energy Release ΔEnc Calculated from the Evaporated Number of Water Molecules ΔNvap.

n kabs [10–11 cm3 s–1] ΦHSA [%] ΦSCC [%] ΦADO [%] ΔNvap ΔEnc [eV]
0 12.3 – – 21 – –
7 2.2 3.0 1.6 – – –
20 1.2 1.6 0.7 – – –
36 1.6 1.7 0.8 – 4.3 1.9
43 1.9 1.9 0.9 – 3.1 1.4
55 2.4 1.9 0.9 – 3.6 1.6
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Table 2. Reaction Energies (in eV) of Hydration of Mg2+(CO2)−(H2O)n and Mg2+(O2)−(H2O)n Clusters, the CO2/O2 Exchange Reaction, and CO4 Formation, Calculated at the CBS-Q3 Level of Theory.

  n
reaction 0 1 2 3 4 5 6 7
Mg2+(CO2)−(H2O)n + H2O → Mg2+(CO2)− (H2O)n+1 –1.14 –1.19 –1.31 –0.90 –0.77 –0.68 –0.58 -
Mg2+(O2)−(H2O)n + H2O → Mg2+(O2)−(H2O)n+1 –2.19 –1.67 –1.30 –0.92 –0.63 –0.59 –0.63 -
Mg2+(CO2)−(H2O)n + O2 → Mg2+(O2)−(H2O)n + CO2 –0.29 –1.34 –1.82 –1.80 –1.82 –1.69 –1.60 –1.64
Mg2+(CO2)−(H2O)n + O2 → Mg2+(CO4)−(H2O)n –0.22 –1.20 –1.66 –1.67 –1.66 –1.63 –1.64 –1.68
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The Nonhydrated Species Mg+(CO2)

First, we discuss the O2/CO2 reaction for n = 0, i.e., the conversion from Mg+(CO2) to Mg2+(O2)−. This reaction proceeds with a relatively high rate coefficient of kabs = 1.2 × 10–10 cm3 s–1 resulting in an efficiency of ΦADO = 21%. The same reaction was previously examined in a fast flow tube-mass spectrometer at a pressure of 1.2 Torr, with a measured rate coefficient roughly five times smaller than in our experiment,35,94 probably due to the different pressure. In the fast flow tube-mass spectrometer experiment, the [(CO2)Mg(O2)]+ complex was also observed, which got stabilized in collisions with background gas.35 Since the pressure in our chamber is 7 orders of magnitude lower than in the fast flow tube, there are nearly no collisions, resulting in immediate elimination of the CO2 molecule.

The calculated exchange reaction energy is −0.29 eV (at the CBS-QB3 level), the respective reaction profile is shown in Figure 1. Theoretical calculations predict that Mg+(CO2) is linear while Mg2+(O2)− has C2v symmetry, as pointed out before.88 Analysis of atomic charges shows that these two ions have considerably different electronic structures: The bonding in Mg+(CO2) can be best described as an ion-induced dipole interaction, with a charge transfer of −0.29 e from CO2 to Mg+ (qMg = 0.71 e), because the linear CO2 is reluctant to accept an electron. In Mg2+(O2)−, a considerable charge transfer from Mg+ to O2 is observed (qMg = 1.63 e).

Figure 1.

Figure 1

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Calculated reaction profile for reaction of [MgCO2]+ with O2. Calculated energies (in eV) are given at the CBS-QB3 and M06/def2TZVP (in parentheses) levels of theory.

The O2/CO2 exchange reaction on Mg+ is predicted to follow a direct pathway, with adsorption of O2 followed by dissociation of CO2. The energy released during the formation of the [(CO2)Mg(O2)]+ encounter complex (∼1.8 eV) easily induces dissociation of the CO2 unit in the absence of stabilizing collisions. In the [(CO2)Mg(O2)]+ structure, there is already a considerable charge transfer from Mg (qMg = 1.34 e) toward O2 (qO2 = –0.55 e). The high stability of [CO2MgO2]+ is in agreement with its observation in the above-mentioned flow tube experiment.35

The Mg2+(CO4)− structure is a local minimum on the potential energy surface, with the charge on Mg calculated as qMg = 1.68 e. However, its formation requires considerable cluster reorganization, and it is stabilized by only ∼0.2 eV (Figure 1) relative to Mg+(CO2). For energetic as well as mechanistic reasons, we do not expect an Mg2+(CO4)− intermediate to be formed in our experiment.

The course of the Mg+(CO2) + O2 reaction was further studied by molecular dynamics (see Table 3 and the SI). Only two channels were observed during the simulation time (7 ps), viz. formation of [(CO2)Mg(O2)]+ (75%) and scattering of O2 on the Mg+(CO2) ion when O2 approaches the ion from the side of the CO2 (25%). We observed no elimination of CO2 on the time scale of 7 ps. This can be understood considering the high dissociation energy of CO2 from Mg2+(O2)− (Figure 1). According to our RRKM calculations, the rate of CO2 dissociation is about 5 × 105 s–1 when disregarding thermal energy of the cluster, using energetics and frequencies calculated at the M06/def2TZVP level. Thus, on the time scale of the experiment (i.e., seconds), energy redistribution takes place and CO2 dissociates.

Table 3. Reaction Channels (in %) Observed during Molecular Dynamics for Four Selected Mg2+(CO2)−(H2O)n Clustersa.

reaction channel Mg+(CO2) Mg2+(CO2–)(H2O)2 Mg2+(CO2)−(H2O)5, Va Mg2+(CO2)−(H2O)5, Vb
scattering, i.e., [Mg(CO2)(H2O)n]+ + O2 25 60 70 85
Mg2+(O2)−(H2O)n + CO2 0 35 10 0
[(CO2)Mg(O2)(H2O)n]+ 75 0 5 0
Mg2+(CO4)−(H2O)n 0 5 15 15
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a

A total of 20 molecular dynamics runs on the M06/6-31+G* potential energy surface were performed for each isomer, with total time of 7 ps (prolonged to 12 ps when the CO4 moiety is formed). The O2/CO2 exchange reaction proceeds either directly on the Mg+ core (for Mg2+(CO2)−(H2O)2) or through CO4– formation (for Mg2+(CO2)−(H2O)5).

The Hydrated Species Mg2+(CO2)−(H2O)n

When the Mg+(CO2) core is hydrated to Mg2+(CO2)−(H2O)n, different reactivity patterns are observed. Figure 2 shows the mass spectra at an O2 pressure of 6.4 × 10–8 mbar after different delays. The corresponding reaction kinetics can be seen in Figure 3a. After 4 s, more than half of Mg2+(CO2)−(H2O)n ions were converted to Mg2+(O2)−(H2O)n. No reaction with a second oxygen molecule was observed. At the same time, the [(CO2)Mg(O2)(H2O)n]+ complex was not observed, suggesting that CO2 leaves the cluster soon after O2 uptake. The absence of [(CO2)Mg(O2)(H2O)n]+ in the FT-ICR mass spectra places an upper limit of ∼100 ms on the lifetime of these species, but it can be expected that the actual lifetime is significantly shorter.

Figure 2.

Figure 2

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Mass spectra of the reaction of [Mg(CO2)(H2O)n]+ with O2 at a pressure of 6.4 × 10–8 mbar after 0, 2.8, 6 and 10 s. Quantitative formation of Mg(O2)(H2O)n-x+ was observed.

Figure 3.

Figure 3

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(a) Reaction kinetics of reaction 1 extracted from the mass spectra seen in Figure 2. (b) Average cluster size ⟨n⟩ of [Mg(CO2)(H2O)n]+ and [Mg(O2)(H2O)n]+. Clusters shrink due to the exothermicity of the reaction and BIRD.

Figure 3b summarizes the average cluster size ⟨n⟩ of reactant and product as a function of time. A nanocalorimetric fit yields the number of water molecules that evaporate due to the reaction. In the first 10 s, BIRD has a large influence on the average cluster size. The BIRD rate decreases with decreasing size.70,74 When n = 3 is reached, no further dissociation is observed on the time scale of the experiment.

A mean rate coefficient kabs = 1.5 × 10–11 cm3 s–1 has been obtained for n ≥ 20, with similar rate coefficients in the n = 20–55 range (see Table 1). For a cluster with 43 water molecules, reaction efficiency is predicted to be ΦHSA = 1.9% and ΦSCC = 0.9%. The actual efficiency might lie somewhere in between, i.e. about one in 70 collisions is reactive. Compared to the complex without water molecules, the exchange reaction is an order of magnitude slower.

For the quantitative nanocalorimetric analysis, only data were fitted with a starting average cluster size of n ≥ 36 to minimize the influence of the size dependent rate coefficient. The mean energy release of reaction 1 corresponds to ΔEnc(1) = 1.7 ± 0.5 eV, and the mean number of evaporated water molecules m = 3.8 as calculated for clusters with n ≥ 36 and averaged over six measurements. The energy is, within the experimental uncertainty, identical to the value of ΔEnc(2) = 1.5 ± 0.3 eV measured in a previous study for the reaction without the magnesium core, reaction 2.21

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To understand the reaction mechanism, in particular whether formation of a Mg2+(CO4)−(H2O)n intermediate is required, we calculated structures of Mg2+(CO2)−(H2O)n, Mg2+(O2)−(H2O)n and Mg2+(CO4)−(H2O)n clusters for n = 1–7 as shown in Figure 4. The calculated energies of hydration and exchange reactions are collected in Table 2. When Mg+CO2 is hydrated by one water molecule, there are two possible isomers with different electronic structure. The more stable isomer Ia with linear CO2 shows a very limited charge transfer, and electron density is actually transferred from CO2 to Mg+(H2O) (qCO2 = 0.29 e). In isomer Ib, which lies 0.44 eV above Ia, CO2 is coordinated in bidentate fashion to Mg+, and a considerable charge transfer to CO2 takes place (qCO2 = –0.63 e).

Figure 4.

Figure 4

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Calculated structures of selected [Mg(CO2)(H2O)n]+, [Mg(O2)(H2O)n]+ and [Mg(CO4)(H2O)n]+ clusters along with relative energy (in eV) and charge on the CO2 or O2 unit (in e), at the CBS-QB3 level of theory. Charges were retrieved using the CHELPG scheme at the MP2/def2TZVP level of theory.

For two water molecules, CO2 connected to Mg+ already tends to accept an electron, loses its linearity, and binds either in monodentate or bidentate fashion. The structure with linear CO2 (IIb) is also local minimum on the potential energy surface, but is higher in energy. The bidentate motif is predicted to be the most stable one for n = 1–4. However, there is only a small energy difference between monodentate and bidentate structures. Already, for n = 7, the bidentate structure collapses during optimization into a monodentate structure. For n = 6–7, we also optimized two structures where CO2 is not directly attached to Mg+ (VId, VIIb). However, this configuration is less stable by about 0.7 eV.

For all structures with at least 2 water molecules, the charge transfer from Mg+ to CO2 is already substantial, with the CO2 charge of about (–0.9 – –0.6) e for both monodentate and bidentate structure. The structure of hydrated Mg+(CO2) clusters can be thus described as Mg2+(CO2)−(H2O)n. Even in the structure with CO2 not directly attached to Mg+ (VId, VIIb), about −0.5 e is transferred to CO2. This inefficient charge transfer can also explain the relative instability of this structure.

For O2 attached to hydrated Mg+, bidentate binding seems to be more favorable, but the difference relative to the monodentate structure is negligible for n > 3. For n = 6–7, structures without direct O2–Mg+ interaction were also considered (VIc, VIIb). These structures are only about 0.25 eV less stable with respect to structures with Mg+-coordinated O2. At the same time, an HO2 radical might be formed (VIIb), with the proton transfer favored by the strongly polarizing Mg2+ center.

The calculated energies released in the O2/CO2 exchange reaction 1 amount to about (–1.8 – –1.6) eV for n ≥ 2, while n = 0, 1 feature lower exothermicity (see Table 2). In other words, the reaction energy seems to converge already for n = 2, with only limited changes with ongoing hydration, irrespective whether a water molecule is added in the first or second solvation shell. This can be traced to similar hydration energies for clusters with CO2 and O2. The calculated reaction energy agrees within error limits with the experimentally measured value.

In a previous publication,21 a prominent role of the CO4– ion was suggested for reaction 2, as previously predicted by Weber.7 CO4– formation on a hydrated Mg2+(CO2)− center is also strongly exothermic, with a release of 1.6–1.7 eV for n > 1.

For both O2/CO2 exchange reaction and CO4– formation on Mg2+(CO2)−(H2O)n, reaction exothermicity increases considerably for n > 1. However, the exchange reaction on a hydrated cluster is found to proceed much slower in the experiment, compared to reactivity of bare Mg+(CO2). This effect must be caused by water molecules already in the first hydration layer, as a small rate coefficient is observed even for n = 7 where the second hydration layer does not contain more than three water molecules. To determine the reaction mechanism and the role of the Mg+ ion in the process, we investigated the exchange reaction using both time-independent calculations and molecular dynamics simulations.

In Figure 5, we analyze reaction mechanisms through relaxed scans along the potential energy surface. Three reaction mechanisms were considered. For two reactions proceeding through direct charge transfer (a, b), we follow the CO2 angle as the reaction coordinate. For CO2–, an angle of about 135° is expected, neutral CO2 is linear. In Mg2+(CO2)−(H2O)5Va (Figure 5a), the reaction is hindered by the reorganization of the hydrogen bonded network, with a barrier below 0.2 eV. When the CO2 molecule starts to linearize, its charge diminishes, it is pushed out of the cluster and an HO2 moiety is formed. For Mg2+(CO2)−(H2O)6 with CO2 in the second solvation shell (Figure 5b), the reaction barrier further decreases and could not be determined unequivocally. Again, the charge transfer leads to linearization of CO2 and formation of an HO2 moiety.

Figure 5.

Figure 5

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Possible reaction mechanisms for the O2/CO2 exchange calculated as relaxed scans along selected reaction coordinates (CO2 angle or CO4– decomposition) at the M06/def2TZVP level of theory, energy is shown with respect to the entrance channel of the respective structure, i.e., Mg2+(CO2)–(H2O)n + O2. (a) Direct O2/CO2 exchange reaction on the Mg+ center. (b) O2/CO2 exchange reaction remote from the Mg+ center. (c) Formation of Mg2+(O2)– core through CO4– ion decomposition.

The third reaction scenario proceeds through formation of CO4– and direct CO2/O2 exchange on the Mg+ center (Figure 5c). The pathway does not require a high activation energy (∼0.1 eV) and lies well under the entrance channel energy; however, a suitable initial structure with a low O2...Mg distance has to be reached. We also considered oxygen atom exchange between O2 and CO2 on the Mg center, the respective barrier was however found to exceed the entrance channel energy and we do not expect this process to take place.

To understand the role of hydration and the difference between reactivity of singly and doubly coordinated CO2, we study reactivity of Mg2+(CO2)−(H2O)n clusters toward O2 using molecular dynamics. Three structures were selected, IIa, Va, and Vb. Because of the expected low efficiency of the reaction, we performed the dynamics at 250 K, i.e. at higher temperature than the internal temperature of the clusters of about 100–200 K. The results are collected in Table 3, selected videos of MD runs are provided in the Supporting Information.

When we compare reaction channels observed for Mg+(CO2) and Mg2+(CO2)−(H2O)2, a strong influence of hydration can be observed. First, the ratio of scattering reactions increases about twice due to water molecules hindering the approach to the Mg+ core. Second, complexation observed for Mg+(CO2) is here substituted by the exchange reaction that is seen in 35% of cases (note the high reaction energy of about −1.8 eV in Table 2). In one case, formation of a long-living CO4– was observed.

For isomers Va and Vb, we saw similar reaction channels. Scattering of O2 from the ion is the most important channel (70% and 85%, respectively), followed by CO4– formation (15% for both isomers). For isomer Va, dissociation of CO4– to form a Mg2+(O2)− core is observed in two dynamic runs, following the reaction pathway suggested in Figure 5c. Finally, a Mg+(CO2) core with a loosely attached O2 was also observed for Va (5%).

Ligand exchange through CO4– formation is the only reaction channel observed for isomers Va and Vb to produce a Mg2+(O2)−(H2O)5 cluster. Because of the hydration of the Mg+ core, no direct O2/CO2 exchange reaction without CO4– formation is documented. The charge exchange suggested in Figure 5a was not seen, probably due to the short time of the respective nonreactive collision. Molecular dynamics results also explain why the exchange reaction proceeds much slower, although the reaction exothermicity increases considerably. After hydration of Mg+(CO2), the O2 molecule must approach the cluster from a suitable direction to form CO4–. In calculations, we observe an increase in scattering probability by a factor of ∼3 when passing from a bare Mg+(CO2) ion to Mg+(CO2)(H2O)n with n = 5, compared to a factor of ∼5 in the experiment when comparing n = 0 and n ∼ 7.

Another indication of the increased importance of proper orientation of the impacting O2 molecule are the results for CO2–(H2O)n. The metal-free species exhibit a 2–3 times higher rate constant for CO2/O2 exchange21 than Mg2+(CO2)−(H2O)n. Without the Mg+ core, we might expect CO2– located on the cluster surface up to larger cluster sizes, thus increasing the reaction cross section. The nanocalorimetric analysis on the CO2–(H2O)n clusters leads to ΔEnc = 1.5 ± 0.3 eV,21 which is slightly lower compared to the present experiment on the Mg+ core. This supports the assumption that CO2– is located on the surface of the CO2–(H2O)n cluster and bound to a low number of water molecules, resulting in a lower evaporation energy. The measured efficiency is also 50% higher for CO2–(H2O)n than for Mg2+(CO2)−(H2O)n. The presence of Mg+ makes the exchange reaction less effective due to the solvation of the Mg2+(CO2)− core in the clusters.

Conclusion

We studied singly charged hydrated magnesium cations with carbon dioxide, Mg2+(CO2)−(H2O)n, in reaction with neutral molecular oxygen. The reaction is efficient, resulting in an exchange reaction of CO2 to O2. For n = 0, an absolute rate coefficient kabs = 1.2 × 10–10 cm3 s–1 is measured. For hydrated clusters, the reaction is slower, and we obtained an average of kabs = 1.5 × 10–11 cm3 s–1 for hydration by 20–55 water molecules. This behavior can be rationalized by two different reaction scenarios. For n = 0, direct CO2/O2 exchange on the Mg+ core is observed. With hydration, CO4– formation seems to be necessary for the reaction to proceed, as indicated by molecular dynamics simulations. Although CO4– formation and subsequent dissociation to O2– and CO2 on the hydrated Mg+ core are considerably more exothermic compared to the reaction on a bare Mg+ ion, hydration at the same time lowers the reaction cross section for CO4– formation on the cluster. This is also indirectly supported by a higher reaction constant observed for the same reaction without the Mg+ core.21 For the hydrated clusters, we determined the reaction enthalpy by a nanocalorimetric analysis to be 1.7 ± 0.5 eV, in agreement with the calculated value.

Acknowledgments

Financial support from the Austrian Science Fund (FWF), Project No. P 28896, is gratefully acknowledged. M.O. acknowledges support through the Lise Meitner Programme of the FWF, Project No. M2001-NBL. The computational results presented have been achieved using the HPC infrastructure LEO of the University of Innsbruck.

Supporting Information Available

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.8b10530.

  • Cartesian coordinates of all optimized structures (PDF)

  • Videos from molecular dynamics simulations (ZIP)

The authors declare no competing financial interest.

Supplementary Material

jp8b10530_si_001.pdf (128.2KB, pdf)
jp8b10530_si_002.zip (30.9MB, zip)

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