Evidence for Electron Transfer in the Reactions of Hydrated

Article Cite This: J. Phys. Chem. A XXXX, XXX, XXX−XXX

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Evidence for Electron Transfer in the Reactions of Hydrated Monovalent First-Row Transition-Metal Ions M(H2O)n+, M = V, Cr, Mn, Fe, Co, Ni, Cu, and Zn, n < 40, toward 1‑Iodopropane Ina Gernert†,‡ and Martin K. Beyer*,†,‡ †

Institut für Physikalische Chemie, Christian-Albrechts-Universität zu Kiel, Olshausenstraße 40, 24098 Kiel, Germany Institut für Ionenphysik und Angewandte Physik, Universität Innsbruck, Technikerstraße 25, 6020 Innsbruck, Austria

‡

S Supporting Information *

ABSTRACT: Hydrated metal ions in the gas phase serve as model systems to investigate the impact of hydration on the chemistry of monovalent transition-metal centers. As a prototypical organometallic reaction involving electron transfer, the reactions of M(H2O)n+, M = V, Cr, Mn, Fe, Co, Ni, Cu, and Zn, n < 40, with C3H7I are studied by Fourier transform ion cyclotron resonance mass spectrometry. While no reaction was observed for vanadium, three different reactions were observed with the other metals, two of them involving the oxidation of the metal ion. Ligand exchange occurs for all metals except zinc. This reaction is sensitive to the size of the solvation shell and is observed predominantly for small cluster sizes. For Cr, Co, and Zn, the metal center is oxidized with formation of MI+ ions. The formation of [MC3H6(C3H7I)2]+, M = Co+, Ni+, proceeds most likely via oxidative addition of C3H7I to the metal ion via insertion into the C−I bond, followed by reductive elimination of HI. For Cu+, this reaction seems to stop after the insertion of the metal into the C−I bond, resulting in Cu(C3H7I)(H2O)n+. The reactions are compared with earlier studies on electron transfer involving hydrated metal centers.

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increase the costs for water treatment.21 To remove the pollutants, metals, especially iron, can be used.21,22 In a different environmental context, transition-metal ions affect the photochemistry in the atmosphere.23 Because of the erosion of the earth’s crust,23 iron ions in various oxidation states are deposited in remote areas of the oceans, and they are present in atmospheric aerosols.24 As a consequence, aqueous-phase photochemical reactions in atmospheric water droplets take place, initiated by the photoreduction of Fe(III) leading to the formation of Fe(II) radicals.24 Hydrated transition-metal ions, that is, a central metal ion embedded in a solvation shell of water molecules, provide an interesting medium for studies of aqueous transition-metal chemistry in the unusual oxidation state of +I.25 Experiments of transition-metal cations M(H2O)n+, M = Cr, Mn, Fe, Co, Ni, Cu, and Ag,26,27 in water clusters with HCl show the uptake of HCl. Furthermore, loss of HCl is determined at a specific number of water molecules left in the cluster. However, Zn(I) cations are oxidized to Zn(II), and a hydrogen atom is released upon uptake of a second HCl molecule.28 Monovalent hydrated zinc cations are also oxidized by acetonitrile in the gas phase.29 In studies of hydrated metal ions with small molecules like oxygen, nitrous oxide, carbon dioxide,30 and nitric oxide,31 different reactions depending on the metal and reactant were

INTRODUCTION Ions solvated by a defined number of ligand molecules present a useful and interesting gas-phase system for solvation studies. 1−11 These studies give information about the thermodynamics and kinetics of the reactions and provide valuable insight into the role of metal ions as catalysts.12 Several experiments13−15 about the chemistry of iodopropane and other organic molecules on metal surfaces show decomposition of iodopropane and formation of propyl groups due to the adsorption of the alkyl group on the surface. These model systems help to develop an understanding of elementary steps in metal-catalyzed reactions involving organic compounds. Catalytic processes were examined with bare metal ions in the gas phase.16 For example, the catalytic cycle of the oxidation of ethane by Fe+/N2O was examined in the gas phase.17 Furthermore, for Fe+, Co+, and Ni+, dehydrogenation of linear alkanes larger than propane by insertion into the C−C bond was studied by collision-induced dissociation.18 Especially for Co+ and Ni+, olefin formation was observed in the reaction with alkyl halides and alcohols.19,20 There is a large number of examples for the activation of C−C or C−H bonds by bare metal ions.16 It is well-established that the efficiency of the reactions of metal ions is affected by ligands, with the consequence that primary products of the bare metal ion reaction do not necessarily undergo further reactions.16 Metals are also interesting for the reduction of chlorinated hydrocarbons like tetrachloromethane or trichloroethene. Small amounts of these carcinogenic substances in groundwater © XXXX American Chemical Society

Received: August 22, 2017 Revised: November 3, 2017 Published: November 27, 2017 A

DOI: 10.1021/acs.jpca.7b08385 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

Figure 1. Mass spectra of the reaction of Cr(H2O)n+ with C3H7I at a pressure of 5 × 10−9 mbar after 0 (a), 3 (b), and 7 (c) s. Oxidation of Cr+ with formation of CrI+ is observed.

number of evaporating water molecules is obtained by fitting the temporal development of the average cluster size for the product and the reactant cluster distribution.33 Whereas the reactant cluster distribution is only shrinking due to roomtemperature blackbody infrared radiative dissociation (BIRD),42−47 the behavior of the product cluster distribution is additionally influenced by the reaction.33 Both effects are accounted for by a set of differential equations describing the average cluster size as a function of time.33 The fits were obtained by fitting several data sets simultaneously. ΔNvap is the average number of water molecules evaporating due to the exothermicity of the reaction. The reaction enthalpy ΔEnc is derived from the product of ΔNvap with the energy required for the evaporation of one water molecule from a water cluster, which amounts to 43.3 ± 3.1 kJ mol−1 in this cluster size range.48,49 A small thermal correction of 4.3 ± 0.5 kJ mol−1 accounts for the heat of the room-temperature reactant molecule and the contribution to the heat capacity of the newly formed product ion.33 Nanocalorimetry works if the reaction rate is independent from cluster size as well as internal energy content of the clusters. Since the clusters are constantly heated by roomtemperature blackbody radiation and respond by evaporative cooling, the width of the internal energy distribution is at least the binding energy of a water molecule, that is, 43.3 kJ mol−1.48,49 Their internal vibrational temperature lies in the range of 100−150 K.48,50 Although the fits look good, we do not have a direct way to verify whether these assumptions are fulfilled. In particular, metal-ion-doped water clusters tend to exhibit strongly size-dependent reactivity.30,31 The absolute numbers must therefore be taken with care, but they give an idea about the relative exothermicity of the subsequent reaction steps.

observed. On the basis of the thermochemistry of the reactions it was possible to show that cobalt, nickel, and zinc are forming M2+/O2− ion pairs with oxygen involving charge transfer, while chromium forms a dioxygen complex in the late stage of the reaction.30 To learn more about the reactions of the transition-metal cations with alkyl halides in aqueous environment and the electron transfer between metal ion and organic compounds, the interaction of M(H2O)n+, M = V, Cr, Mn, Fe, Co, Ni, Cu, and Zn, n < 40, with C3H7I was investigated by Fouriertransform ion cyclotron resonance (FT-ICR) mass spectrometry.

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EXPERIMENTAL DETAILS All experiments were performed on a modified Bruker/ Spectrospin CMS47X FT-ICR mass spectrometer.32−35 The instrument is equipped with 4.7 T superconducting magnet, Bruker infinity cell, an APEX III data station, TOPPS ion optics power supply, and an ICC2 Infinity Cell Controller with BCH preamplifier. The M+(H2O)n ions were generated with an external laser-vaporization source36−38 and were transferred by an electrostatic lens system to the ICR cell. The vaporization laser (Nd:YAG laser Continuum Surelite II, operated at 532 nm) and the frequency doubling crystal were heated by 20 laser shots for minimizing the drift of the initial cluster size distribution, followed by 20 laser shots at 10 Hz and typically 5 mJ pulse energy to fill the cell. The reaction delay was measured relative to the end of the fill cycle. Therefore, some reaction products were observed at nominal 0 s reaction delay. Isotopically enriched targets31 were used where applicable: 52Cr (99.9%), 56Fe (99.7%), 58Ni (99.9%), 63Cu (99.3%), and 64Zn (99.4%) (STB Isotope Germany GmbH). 1-Iodopropane (Sigma-Aldrich, 99%) was introduced into the ultrahigh vacuum region by a leak valve at a constant pressure in the range from 5 × 10−9 to 2 × 10−7 mbar. The reactions were monitored by recording mass spectra at different delays. The average number of water molecules that are evaporating due to the reaction was determined with nanocalorimetry39 as described before.33,40,41 The information about the average

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RESULTS AND DISCUSSION Vanadium. V(H2O)n+, starting with n = 20−40, is unreactive toward iodopropane. The measurements were done in the pressure range of 3 × 10−8 to 2 × 10−7 mbar, and no reaction was observed. Only the already known B


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The Journal of Physical Chemistry A intracluster reactions51 activated by infrared photons were observed for clusters with n ≤ 20. The loss of atomic and molecular hydrogen with formation of V(II)OH(H2O)m+ and V(III)(OH)2(H2O)p+ is determined. Both reactions are strongly dependent on cluster size.51 However, no further reaction is observed with the oxidation products except shrinking of the cluster distribution due to BIRD. Hydrated monovalent vanadium ions are showing low reactivity toward other substances. Also with other small molecules like N2O, O2, CO2, and NO, no reactions were observed in our experiment.30,31 Chromium. Hydrated monovalent Cr(H2O)n+, n = 20−45, react with iodopropane by formation of [CrI(H2O)m]+, reaction 1, which implies oxidation of the metal center. The mass spectra, shown in Figure 1 for an experiment at a pressure of 5 × 10−9 mbar, reveal formation of [CrI(H2O)m]+ already at nominally 0 s reaction delay, when the 2 s fill cycle of the cell was completed. As mentioned before, shrinking of the cluster distribution is caused by BIRD, immediately evident in Figure 1b,c. Cr(H 2O)n+ + C3H 7I → [CrI(H 2O)m ]+ + C3H 7· + (n − m)H 2O

(1)

This reaction, however, seems to require a minimum cluster size, since for small clusters, slow ligand exchange is observed instead, at delays ∼30 s, reaction 2. Since ligand-exchange products are observed only containing at most two water molecules, it seems to set in only around n ≈ 3−5, when empty coordination sites become available at the metal center. Cr(H 2O)n+ + C3H 7I → [CrC3H 7I(H 2O)p ]+ + (n − p)H 2O

p≤2

(2)

For small clusters, also the primary product [CrI(H2O)m]+ undergoes ligand exchange with iodopropane, ∼20 s reaction delay, reaction 3.

Figure 2. Kinetics (a), average cluster size ⟨n⟩ (b), and difference of the cluster size Δ⟨n⟩ (c) of the reaction of Cr(H2O)n+ with C3H7I at a pressure of 5 × 10−9 mbar. The difference fit (c) deviates systematically from the data points, which suggests a higher reactivity of smaller clusters.

CrI(H 2O)m+ + C3H 7I → [CrIC3H 7I(H 2O)k ]+ + (m − k)H 2O

k≤4

Table 1. Absolute Rate Constants k in 1 × 10−10 cm3 s−1 and Nanocalorimetric Reaction Enthalpy ΔEnc in kJ mol−1 for the Primary Reaction of [M(H2O)n]+ with C3H7I

(3)

After 80 s reaction delay the signal of [CrI(H2O)3]+ has the highest intensity. Two more signals are present, [CrI(C3H7I)(H2O)2]+ and [Cr(C3H7I)2]+, due to further ligand exchange reactions. After 80 s reaction time increase of the intensity of the two clusters and decrease of the [CrI(H2O)3]+ until 500 s reaction delay are observed; see Figure S1 for a mass spectrum at 300 s. A quantitative kinetic analysis, Figure 2a, yields the rate constant for the oxidation reaction. The summed intensities of the reactant and product cluster distributions are fluctuating slightly due to problems to keep the pressure constant. The rate constants for all metals are summarized in Table 1. With k1 = 3 × 10−10 cm3 s−1, reaction 1 is quite efficient. A nanocalorimetric fit of the data is shown in Figure 2b,c. The closing of the gap between reactant and product cluster size in Figure 2b indicates that smaller clusters are reacting a bit faster than the larger ones. A loss of on average 2.4 water molecules for the oxidation reaction is determined. However, the cluster size dependence causes significant error. Cr+ has a half-filled 3d shell with the stable configuration 3d5 of five parallel spins.52 Despite this stability, the oxidation reaction at nominally 0 s reaction delay to Cr2+ is observed with

M Cr Mn Fe Co Ni Cub Zn

k

ΔEnc

3 × 10−10

−102 ± 31a

2.6 × 10−10

−36 ± 22

1.5 × 10−10 4.1 × 10−10

−132 ± 24a −75 ± 23

a Probably too high due to size-dependent reactivity. bLigand exchange without elimination of C3H7.

iodopropane. This is in line with our earlier studies, where the observed slow uptake of O2 as well as CO2 required electron transfer from the Cr+ center to the reactant.30 Also in reactions with acetonitrile and NO, evidence for electron transfer was obtained.29,31 Ligand exchange of small clusters, reactions 2 and 3, can be rationalized by the high polarizability of 11.5 Å3 and strong dipole moment of 2.04 D for iodopropane, compared to 1.45 Å3 and 1.85 D for water.53 With decreasing cluster size the C


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probably occurs first. This was corroborated by repeating the experiments with a smaller cluster distribution of [Mn(H2O)n]+, n = 10−20, performed at the same pressure, to make sure that [HMnOH(H2O)n−1]+ are formed before the uptake of the first iodopropane molecule. Loss of all water molecules was also observed here. Mn+ has a half-filled 3d shell and one electron in the 4s shell.57 Thus, the oxidation involves transfer of the 4s electron. Both ligand exchange as well as oxidation of the metal were observed by Uppal and Staley57 with singly charged manganese cations undergoing sequential reactions with various alkyl chlorides in the gas phase. Iron. Similar to manganese, monovalent hydrated Fe(H2O)n+ ions, starting with n = 15−40, exhibit a strong size dependence in their reactivity. Only after 5 s ligand exchange, reaction 8, is observed, albeit with very low efficiency; see Figure S3. Formation of [FeI(H2O)p]+ is observed only after 9 s, reaction 9, but even after 20 s, only traces of this product are present.

probability increases that iodopropane reaches a free coordination site of the metal. Oxidation reactions of chromium with alkyl halides have been studied to our knowledge only with Cr2+, since Cr+ is not stable in aqueous environment.25,27,54 Kochi and Powers55 analyzed the reduction of alkyl halides with different Cr(II) complexes. For most reactants they observed a reduction of the alkyl halide and the formation of Cr(III) complexes with the halide, in line with our present results for Cr(I). Manganese. The reaction of hydrated Mn(I) ions, starting with n = 20−40, with iodopropane is strongly size-dependent. Only at 28 s reaction delay, Figure 3, obtained at a pressure of 3

Fe(H 2O)n+ + C3H 7I → [FeC3H 7I(H 2O)m ]+ + (n − m)H 2O

+ (n − p)H 2O

Mn(H 2O)n+ + C3H 7I → [MnC3H 7I(H 2O)m ]+ (4)

k≤2

+ (m − k)H 2O

(5)

As in the experiments with acetonitrile, complete loss of water molecules is observed after long delays. After 30 s reaction delay, traces of an oxidation product are observed, reaction 6, followed after 80 s by ligand exchange, reaction 7.

+ (p − q)H 2O

(6)

[MnI(H 2O)p ]+ + C3H 7I → [MnIC3H 7I(H 2O)q ]+ + (p − q)H 2O

q≤2

(10)

q≤3

(11)

As shown in Figure S5, the summed intensities of the products do not follow pseudo-first-order kinetics. The slow onset of product formation after 5 s corresponds to an increased reactivity of small clusters. Comparison of the product intensities with the average cluster size shows that the ligand exchange reaction 8 occurs for clusters with n < 20 and that the oxidation reaction 9 occurs for n < 15, although no precise upper limit can be given. A similar cluster size dependence was observed for the reaction of Fe(H2O)n+ with nitric oxide.31 The initial uptake of NO and formation of HNO occur in a size regime around n = 15−25. It seems that the solvation shell is hindering the reactions of Fe+. One may speculate that, since oxidative addition of iodopropane requires two coordination sites, the reaction is starting in a smaller cluster size regime than for nitric oxide, which occupies only one coordination site. For FeO+, Baranov et al.58 observed the sequential addition of three ligand molecules with H2O, CO2, and N2O in the gas phase, similar to our observation of FeI(C3H7I)3+. However, this analogy must not be taken too far, since in FeO+ iron is in oxidation state of +III, while FeI+ contains Fe(II). Our results are consistent with the measurements of Allison and Ridge.20

Mn(H 2O)n+ + C3H 7I → [MnI(H 2O)p ]+ + C3H 7· p≤3

k≤3

[FeI(H 2O)p ]+ + C3H 7I → [FeIC3H 7I(H 2O)q ]+

29

+ (n − p)H 2O

(9)

[FeC3H 7I(H 2O)m ]+ + C3H 7I → [Fe(C3H 7I)2 (H 2O)k ]+

MnC3H 7I(H 2O)m+ + C3H 7I → [Mn(C3H 7I)2 (H 2O)k ]+ + (m − k)H 2O

p ≤ 10

The reaction becomes more efficient when the hydration shell has shrunken to n = 5, and formation of an FeI+ core becomes dominant, possibly in secondary reactions. In the mass spectrum at 50 s, Figure S4, all water molecules are exchanged, and additional ligand exchange occurs, reactions 10 and 11. The most intense product at this delay is FeI(C3H7I)3+.

× 10−8 mbar, the first uptake of iodopropane by very small clusters is observed, reaction 4, followed by a second uptake at later times, reaction 5; see Figure S2.

m≤4

(8)

Fe(H 2O)n+ + C3H 7I → [FeI(H 2O)p ]+ + C3H 7·

Figure 3. Mass spectra of the reaction of Mn(H2O)n+ with C3H7I at a pressure of 3 × 10−8 mbar after 0 and 28 s. Ligand exchange with iodopropane is observed.

+ (n − m)H 2O

m ≤ 12

(7)

Further ligand exchange leads to the products observed after 300 s, [MnI(C3H7I)2]+ and [MnI(C3H7I)3]+, which dominate the mass spectrum; see Figure S2. Both reactions occur at long reaction delays, when the cluster size is smaller than six water molecules, and free coordination sites become available. An earlier study of the reactions of [Mn(H2O)n]+ with D2O showed that [Mn(H2O)n]+, n ≤ 20, undergo an intracluster redox reaction with formation of [HMnOH(H2O)n−1]+.56 The complete loss of water molecules observed here implies that the intracluster reaction leading to [HMnOH(H2O)n−1]+ is reversible, possibly mediated by the presence of additional ligands, in this case C3H7I. Since the reaction takes place at long delays, the BIRD-activated hydride−hydroxide formation D


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The Journal of Physical Chemistry A They examined the chemistry of bare metal cations Fe+ with alkyl halides and alcohols in the gas phase. Formation of FeI+ was observed in the reaction of the bare Fe+ with iodomethane, while the FeCH3I+ product was only observed via ligand exchange with Fe(CO)+. Cobalt. In the reaction of Co(H2O)n+, starting with n = 15− 40, product peaks CoI(H2O)m+ with very low intensity are already present at nominally 0 s reaction delay, reaction 12; see Figure S6. After 10 s, that is, for small clusters, also ligand exchange, reaction 13, is starting. Both products undergo further ligand exchange with up to three molecules of C3H7I.

[Co(C3H 7I)2 X]+ + C3H 7I → [CoC3H6(C3H 7I)2 ]+ + HI +X

(12)

Co(H 2O)n+ + C3H 7I → [CoC3H 7I(H 2O)p ]+ + (n − p)H 2O

p≤5

(14)

The kinetic and nanocalorimetric analysis of the first 5 s, where only reaction 12 occurs, Figure S7, does not show any signs of a strong cluster size dependence. Therefore, the value for the reaction enthalpy should be reliable; see Table 1. The reaction is with k12 = 2.6 × 10−10 cm3 s−1 quite efficient, but with ΔEnc,12 = −36 ± 22 kJ mol−1 only weakly exothermic. The HI elimination reaction 14 has precedent in the literature. Allison and Ridge observed for the reaction of bare Co+ and Ni+ with 2-chloropropane in the gas phase a metal insertion, followed by formation of MC3H6+ and HCl.20 For the ligand exchange product [Co(NO)C3H7Cl]+ they report the formation of [C3H6Co(NO)C3H7Cl]+ during the reaction with a further molecule i-C3H7Cl, which closely resembles reaction 14. These reactions require the presence of a β-hydrogen, since for chloromethane no HCl elimination was observed. Nickel. Monovalent hydrated nickel cations Ni(H2O)n+, starting with n = 15−40, at 4 × 10−8 mbar undergo ligand exchange with iodopropane, reaction 15, as evident from the mass spectra in Figure S8. After 6 s the sequential uptake of up to four molecules of C3H7I is observed; see reaction 16. The measured summed intensities and the average cluster size for the ligand exchange reaction 15 are shown in Figure S9. For the first two seconds the intensity of the product clusters is low, and only a weak increase of the intensity is observed. After 3 s, when the average cluster size drops below ⟨n⟩ = 20, the intensity of the products is increasing faster. Interestingly, the average cluster size of the product remains constant after the first few seconds. This indicates that large clusters with more than ∼15 water molecules do not react, consistent with the observed maximum of 12 water molecules attached to the product of reaction 15.

Co(H 2O)n+ + C3H 7I → [CoI(H 2O)m ]+ + C3H 7· + (n − m)H 2O

X = H 2O, C3H 7I

(13)

Somewhat unexpectedly, at 15 s reaction delay [CoC3H6(C3H7I)2]+ is formed, which besides [CoI(C3H7I)3]+ becomes the second major product of the reaction. Looking at the intensities at delays up to 500 s, Figure 4b, indicates that

Ni(H 2O)n+ + C3H 7I → [NiC3H 7I(H 2O)m ]+ + (n − m)H 2O

m ≤ 12

(15)

[NiC3H 7I(H 2O)m ]+ + C3H 7I → [Ni(C3H 7I)2 − 4 (H 2O)k ]+ (m − k)H 2O

k≤6

(16) +

Already at 15 s, Ni(C3H7I)4 has the highest intensity. Interestingly, after 10 s the formation of [NiC3H6(C3H7I)2]+ is observed. From the mass spectra shown in Figure 5, one can deduce that this ion is formed exclusively from [Ni(C3H7I)3]+, most likely in the collision with another molecule C3H7I via reaction 17, which resembles reaction 14. As for Co+, also for bare Ni+ several experiments are known, where metal insertion into a C−X bond is followed by formation of metal−olefin complexes in the gas phase.18−20

Figure 4. (a) Mass spectrum of the reaction of Co(H2O)n+ with C3H7I at a pressure of 4 × 10−9 mbar after 200 s. Quantitative formation of [CoC3H6(C3H7I)2]+ is observed. Small peaks caused by reactions with impurities are noticed at longer reaction delays. (b) Product intensity as a function of reaction delay (15−500 s).

[Ni(C3H 7I)3 ]+ + C3H 7I → [NiC3H6(C3H 7I)2 ]+ + HI + C3H 7I

(17)

Copper. starting with n = 15−40, react at 4 × 10−8 mbar with sequential uptake of up to five molecules C3H7I, reaction 18, as immediately evident from the mass spectra shown in Figure 6. In contrast to the earlier discussed metals, the first two steps are efficient and occur for relatively large cluster sizes. At long reaction delays, however, the iodopropane molecules are lost again from the cluster due to BIRD. Interestingly, after 500 s reaction delay at 4 × 10−9 mbar Cu(H2O)n+,

[CoC3H6(C3H7I)2]+ is formed from the precursors [Co(C3H7I)2H2O]+ or [Co(C3H7I)3]+, reaction 14, most likely via oxidative addition of C3H7I to the metal center, that is, insertion of Co+ into the C−I bond, followed by reductive elimination of HI. E


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addition. Precedence for such a behavior is available for bare Cu+ in the literature. Jones and Staley59 analyzed the chemistry of Cu+ with alkyl chlorides in the gas phase. They observed that Cu+ reacts with alkyl halides by dehydrochlorination with retention of HCl or the alkene by Cu(I). Furthermore, Lang et al.60 demonstrated the amination of aryl halides using Cu(I) catalysis in solution. With Cu2O as a catalyst, iodobenzene was converted to aniline with 74% yield. The authors suggested that the first step of the catalytic reaction is insertion of Cu(I) in the C−I bond. Interestingly, in the reaction of Cu(H2O)n+ with C3H7I studied here, reductive elimination of HI is not observed. One may speculate that the intermediate formed following βhydrogen shift is not stable for copper. Zinc. The reaction of iodopropane with Zn(H2O)n+ is the fastest reaction. At nominally 0 s reaction delay at 4 × 10−8 mbar, the intensity of the product ZnI(H2O)m+ is already higher than that of the reactant, Figure S12. Already at 6 s all reactant clusters have been converted to ZnI(H2O)m+ via reaction 19. From 8 s on, ligand exchange of ZnI(H2O)m+ is observed, reaction 20; see Figure S13 for a mass spectrum at 20 s. After 60 s reaction delay ZnI(C3H7I)3+ is the most intense product, but at longer reaction delays further ligand loss is observed due to BIRD.

Figure 5. Mass spectra of the reaction of Ni(H2O)n+ with C3H7I at a pressure of 4 × 10−8 mbar after 20 (a), 30 (b), and 50 (c) s. Formation of [Ni(C3H7I)2C3H6]+ is observed. Small peaks are due to impurities.

the products of the reaction are Cu(C 3 H 7 I) 2 + and Cu(C3H7I)H2O+, Figure S10; that is, one water molecule resists ligand exchange. This is different from all other metals, for which the final exchange products are water-free.

Zn(H 2O)n+ + C3H 7I → [ZnI(H 2O)m ]+ + C3H 7· + (n − m)H 2O

+

[Cu(C3H 7I)x (H 2O)n ] + C3H 7I → [Cu(C3H 7I)x + 1(H 2O)m ]+ + (n − m)H 2O x=0−4

(19)

ZnI(C3H 7I)x (H 2O)m+ + C3H 7 → ZnI(C3H 7I)x + 1 (18)

(H 2O)k + + (m − k)H 2O

To obtain quantitative data, the kinetic and nanocalorimetric analyses were performed for a data set obtained at a reactant pressure of 4 × 10−9 mbar; see Figure S11. With k18 = 1.5 × 10−10 cm3 s−1, uptake of the first molecule is a little slower than the oxidation reactions observed with the other metals, but it is still efficient. The reaction seems to be strongly exothermic, with ΔEnc,18 = −132 ± 24 kJ mol−1. This is in fact too exothermic for a simple ligand exchange, and it suggests that Cu+ inserts into the C−I bond via oxidative

x=0−2

(20)

The kinetic and nanocalorimetric analysis is shown in Figure 7. The intensity behavior of the reactant cluster distribution in the first half second of the fit is not completely linear in the semilogarithmic plot, which points again to a small cluster-size dependence and a slightly increased reactivity for smaller clusters. The rate is with k19 = 4.1 × 10−10 cm3 s−1 relatively high; the reaction enthalpy lies with ΔEnc,19 = −75 ± 23 kJ mol−1 between the values for chromium and cobalt. In

Figure 6. Mass spectra of the reaction of Cu(H2O)n+ with C3H7I at a pressure of 4 × 10−8 mbar after 0 (a), 2 (b), and 4 (c) s. The ligand exchange with two iodopropane molecules was observed. F


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available for bonding in V+. It is possible that V(H2O)n+ contains a HVOH+ ionic core, like in the case of Al(H2O)n+.61 This would completely change the coordination chemistry and reactivity of the vanadium center. The vanadium species were also unreactive against NO, O2, and CO2. Formation of MI(H2O)m+ with oxidation of the metal center to +II oxidation state was observed for chromium, manganese, iron, cobalt, and zinc. Only monovalent copper and nickel cations are not showing the formation of MI+ core. Cu+ has filled d orbitals and an empty 4s orbital;52 therefore, the reactivity of hydrated monovalent copper ions is low.29,31 Although Ni+ has an electron configuration of 3d9 and one partly filled d orbital is available,62 it shows for most reactions only ligand exchange.29−31 On the one hand, this might be due to its relatively high second ionization energy (IE), more than 1 eV higher than that of cobalt. On the other hand, the second IE of zinc is even higher, and still Zn+(H2O)n are the most reactive species. The efficient ZnI+(H2O)m formation corresponds to the reactivity with acetonitrile, where Zn+(H2O)n is the only species for which hydroxide formation ZnOH+(H2O)m is observed.29 Chromium, cobalt, and zinc are reacting with very similar total rates, but they differ in the determined nanocalorimetric formation enthalpy. The reaction with Co(H2O)n+ is less exothermic than with zinc and chromium. The high reaction enthalpy was also observed for the reaction of Cr+ with acetonitrile29 and oxygen.30 For Zn+, compared to the usual divalent state, the additional electron is located in the 4s orbital. Formation of ZnI+ with transfer of the 4s electron to iodine dramatically reduces the ionic radius of Zn, leading to significantly increased hydration energy, which rationalizes the high exothermicity for Zn+, despite a high second IE. For Cr+ and Co+, the difference in thermochemistry almost exactly matches the difference of the second ionization energies, which seems plausible. The reactivity of manganese and iron depends strongly on cluster size. Manganese is reacting in a narrow cluster-size regime around n = 4−6. This shows a high dependence on the solvation shell of the reactivity. Exchange experiments with D2O together with BIRD experiments for the same metals showed that only Mn(H2O)n+ undergoes an intracluster reaction to form [HMnOH(H2O)n−1]+ for n ≤ 20.56 The reason for the late start of the reaction could be the requirement to reverse this intracluster reaction; that is, the oxidation reaction only works for Mn(H2O)n+ with intact H2O molecules, which would be consistent with the complete loss of water ligands as discussed above.

Figure 7. Kinetics (a), average cluster size ⟨ n ⟩ (b), and difference of the cluster size Δ⟨n ⟩ (c) of the reaction of Zn(H2O)n+ with C3H7I at a pressure of 4 × 10−8 mbar for the first 3 s. The kinetic fits follow a pseudo-first-order behavior.

summary, reaction 19 follows textbook chemistry, since it brings zinc to its preferred oxidation state. Metals in Comparison. The reactivity of the studied systems is quite diverse; therefore, the key behavior is summarized in Table 2 for each metal and compared with earlier works on NO, O2, and CO2,30,31 as well as the second ionization energy and ionic radius of the M2+ ion.39 Some patterns emerge, but also astounding discrepancies are recognized. Only for vanadium no reaction with iodopropane was observed, although empty and partly filled d orbitals are

Table 2. Summary of the Observed Reaction Pathways in the Reaction of [M(H2O)n]+ with C3H7I and Comparison with Earlier Data on Reactivity with NO, O2, CO2, Second Ionization Energy (eV), and Ionic Radius (Å) of Hexacoordinated M2+ M V Cr Mn Fe Co Ni Cu Zn a

MI(H2O)m+

[M(C3H7I)x(H2O)m]+

all very small small all

small very small small small small alld

[MC3H6(C3H7I)2]+

NOa

O2b fast

slow

yes yes

fast very slow fast, hydroxide fast fast

fast slow

very slow

fast

fast

all

CO2b

2nd IEc

r(M2+)c

14.66 16.49 15.64 16.19 17.08 18.17 20.29 17.96

0.79 0.73 0.67 0.61 0.65 0.69 0.73 0.74

Ref 31. bRef 30. cRef 53. dEvidence for oxidative addition. G


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The Journal of Physical Chemistry A Scheme 1. Possible Metal Insertion Mechanisma for Cobalt and Nickel for the Formation of MC3H6+

a

On the basis of ref 20.

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The ligand exchange reaction of the M(H2O)n+ species is observed for all metals except zinc. This reaction is strongly dependent on the cluster size and mostly occurs at small cluster sizes. This behavior correlates with the low solubility of iodopropane in aqueous environment, with 1.14 g/kg H2O.53 In addition, the hydration shell blocks the free coordination sites of the metal ion. Zn(H2O)n+ is not reacting by ligand exchange, since all clusters react by formation of [ZnI(H2O)n]+ before small cluster sizes of the reactants are reached. For copper, the exothermicity of the ligand exchange reaction points toward oxidative addition of C3H7I via insertion of the metal into the C−I bond. For the other metals, this insertion is also possible. In particular, for Ni(H2O)n+ and Co(H2O)n+, the formation of [MC3H6(C3H7I)2]+ observed for longer reaction delay most likely starts with oxidative addition, as illustrated in Scheme 1, followed by a β-hydrogen shift and reductive elimination of HI.

ACKNOWLEDGMENTS Financial support from the Deutsche Forschungsgemeinschaft, Grant No. BE2505/4-3, is gratefully acknowledged.

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CONCLUSIONS The studied hydrated metal ions exhibit a rich variety of reaction pathways with 1-iodopropane. Most metals are oxidized with formation of the hydrated metal iodide and release of a propyl radical, but for manganese and iron, this is observed only for very small hydration shells. Ligand exchange was observed for all metals except zinc and vanadium and may go along with oxidative addition via insertion of the metal into the C−I bond. This reaction is sensitive to the cluster size and starts in most cases in a narrow cluster size range. For cobalt and nickel the formation of [MC3H6(C3H7I)2]+ was observed, which requires metal insertion into the C−I bond followed by loss a β-hydrogen shift to afford reductive elimination of HI and formation of a metal−propene complex. There is no apparent correlation of the reactivity with the second ionization energy or the ionic radius, which underlines the complex interplay of different factors governing transition-metal chemistry.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b08385. Mass spectra, intensity data, cluster size data, kinetics data (PDF)

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Martin K. Beyer: 0000-0001-9373-9266 Notes

The authors declare no competing financial interest. H


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