Main group coordination chemistry at low temperatures: A review of matrix isolated Group 12 to Group 18 complexes

Accepted Manuscript Title: Main Group Coordination Chemistry at Low Temperatures: A Review of Matrix Isolated Group 12 to Group 18 Complexes Author: Nigel A. Young PII: DOI: Reference:

S0010-8545(12)00260-3 doi:10.1016/j.ccr.2012.10.013 CCR 111650

To appear in:

Coordination Chemistry Reviews

Received date: Revised date: Accepted date:

4-9-2012 22-10-2012 22-10-2012

Please cite this article as: N.A. Young, Main Group Coordination Chemistry at Low Temperatures: a Review of Matrix Isolated Group 12 to Group 18 Complexes, Coordination Chemistry Reviews (2010), doi:10.1016/j.ccr.2012.10.013 This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

Main Group Coordination Chemistry at Low Temperatures: a Review of Matrix Isolated Group 12 to Group 18 Complexes.

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Nigel A. Young Department of Chemistry The University of Hull Kingston upon Hull HU6 7RX UK [email protected] Tel: 44 1482 465442 Fax 44 1482 466410

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Abstract

The available data on main group (groups 12 to 18) Lewis acid-base, donor-acceptor, charge-

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transfer and van der Waals complexes stabilised by matrix isolation techniques are presented, tabulated and evaluated in conjunction with data from complementary gas phase experiments

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Matrix isolation Coordination complexes Lewis acid-base donor-acceptor charge transfer van der Waals

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Keywords

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and computational chemistry.

Dedication

This review is dedicated to the memory of Steve Ogden, who introduced me to, and taught me so much about, the art of matrix isolation and the science of vibrational spectroscopy.

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Introduction The definition of coordination chemistry this review will utilise is the formation of a discrete new entity (complex, adduct) formed as a result of a Lewis acid-base (acceptor-donor, charge transfer) interaction between two species, with the formation of a new dative bond. This is

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shown schematically in Figure 1 for the formation of the ammonia borane (H3N•BH3) complex. The review will focus on the complexes of main group compounds with donor and

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acceptor atoms from Groups 13 to 18. Whilst it might be thought that “the only difference between a covalent bond and a dative bond is where you think the electrons come from” [1],

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Haaland has presented an elegant distinction between covalent and dative bonding based on the nature of the fragments formed when the relevant bond is broken either heterolytically or

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homolytically [1]. Rupture of a dative bond (such as in the ammonia borane complex) yields either two species without net charge or spin (H3N and BH3), or two species with both net charge and net spin (H3N•+ and BH3•−). In contrast, rupture of the covalent bond in the

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isoelectronic H3C-CH3 yields either species with net spin (H3C•) or species with net charge (H3C+ and H3C−). As the minimum energy rupture in either dative or covalent bonding will

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yield neutral species, this indicates that dative bond rupture proceeds heterolytically, whilst covalent bond rupture proceeds homolytically. Haaland also comments that the strength of

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the dative, or donor-acceptor, bond is very sensitive to inductive effects, especially at the acceptor (acid) atom. He also notes that the strength of dative bonds does not normally exceed half of that of a normal, covalent, bond between the same atom pair, with the result that the bond distances are longer in dative bound complexes than their covalent analogues [1].

A wide variety of notations are used in the literature for representing Lewis acid-base complexes. To aid clarity and discussion they will be presented in this review as base•acid, with the connecting atoms adjacent to the • e.g. H3N•BH3. (Where there is ambiguity, or for hydrogen bonding, this will not be specified and there will just be a space.) There are also a number of different terminologies used to describe such species, including Lewis-acid base, donor-acceptor, charge-transfer, van der Waals etc, which imply differing levels of interaction between the components, but these are not always specified by the authors. Mulliken introduced a sophisticated nomenclature for describing such complexes or adducts, with the most general terms being inner type complexes where there is significant charge 2 Page 2 of 213

redistribution and outer type complexes where the interaction is weak with very little charge transfer [2, 3]. Reed et al. have shown that for the n-σ* interactions to overcome the n-σ repulsion, the acceptor orbitals must protrude more into space than the corresponding filled orbitals on the acid. This results in the donor-acceptor interactions becoming sufficiently significant so that the acceptor and donor species are pulled within the van der Waals contact distance [4]. They developed a criterion that for there to be significant donor-acceptor

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character in the bonding, the van der Waals penetration distance, dp, should be ≥ 0.1 Å, where

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dp = RvdW −Req

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and RvdW is the sum of the van der Waals radii of the appropriate atoms, Req is the equilibrium

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interatomic distance in the complex.

χ = (RvdW – Req)/(RvdW – Rcov)

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They also developed a covalency ratio, χ

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where Rcov denotes the sum of the atomic covalent radii for the atoms involved in the closest

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contact in the complex. Combining this with the previous equation (assuming that the van der Waals radius of an atom is 0.8 Å greater than the covalent radius) means that χ can be expressed as χ ≈ dp/1.6

For hydrogen bonding, values of χ ≈0.3 are common, but are known up to 0.9 for the bifluoride ion [4].

Likewise different phrases such as binding energies, stabilisation energies and interaction energies, as well as dissociation energies are used by different authors to describe the thermochemical parameters defining the strength of the dative bond. In addition, as the complexes are often weakly bound, these are very sensitive to the computational methodology and basis sets used, and may or may not have been corrected for zpe, bsse and temperature in a variety of ways. Therefore, it is not always easy to make meaningful comparisons between reports from different groups. In general, the authors’ original 3 Page 3 of 213

terminology will be used to describe the complex, and the results of their thermochemical calculations. Weakly acidic and basic species generally form complexes, which may or may not be induced to participate in subsequent thermal or photochemical reactions. Jonas et al. have provided some calibration for the strength of interaction in these type of complexes [5, 6].

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They regarded values of the dissociation energy (D0) of ca. 100-200 kJ mol−1 as calculated for OC•BH3, H3N•BH3, Me3N•BH3, H3N•BF3, H3N•BCl3, and Me3N•BCl3 as indicating a

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strong interaction. Weakly bound, van der Waals, complexes such as OC•BF3, HCN•BF3, MeCN•BF3, Me2O•BF3, OC•BCl3 and MeCN•BCl3 have smaller D0 of around 20 – 40 kJ

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mol−1 [6]. Complexes with even smaller values have been calculated.

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Main group coordination chemistry involving dative bond formation, is commonly carried out in solution or the gas phase, and this approach has resulted in the spectroscopic and structural characterisation of many different types and examples of coordination complexes

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[7]. Reviews of gas phase complexes of the Group 13-15 complexes have been recently published [7, 8]. However, there are also occasions when the proposed the complex may not

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be observable under these conditions.

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There are three general causes of this: (i)

weakly bound complexes

(ii)

complexes with intermediate stoichiometry

(iii)

intermediate complexes of highly reactive reagents.

The first of these is where the coordinate bond is so weak with small binding energies that entropy works against complex formation at room temperature, and as the dative bond is broken by thermal energy they are likely to be fully dissociated. However, if the experiments are carried out at low temperatures, complex formation can be stabilised. Probably the best examples of this are the equilibria between the nitrogen oxide monomers and the dimers which have weak N-N bonds that are only stable at low temperature. For example the white N2O4 solid stable at low temperatures reverts to the brown NO2 gas at room temperature, and the intense blue N2O3 solid and liquid dissociates to NO2 and NO on warming to room temperature. The second and third causes are related and involve slowing down or catching reactions before they go to completion. In (ii) the aim would be to stabilise a 1:1 complex, if a 1:2 or 2:1 complex was known at room temperature, and in (iii) the objective would be to 4 Page 4 of 213

trap an initial complex before elimination or decomposition resulted in the conventional reaction product. Paradoxically, lowering the temperature can simultaneously solve the problems of both thermal fragility and extreme reactivity observed at ambient conditions, so that the rate of product formation can be increased in all three cases. Whilst it is possible to use neat thin

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films at 77 K, the information content of the spectra is much higher if the reagents and products are diluted and trapped in cryogenic, inert, matrices at ca. 4 - 14 K, as the inter-

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molecular interactions are kept to a minimum. This is known as matrix isolation, and has allowed for the characterisation of a very wide selection of otherwise intractable, unstable

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and highly reactive species since its development by Pimentel in the 1950s [9]. In essence it is similar to preparing a KBr disc for IR spectroscopy, except that the KBr is replaced by a

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cryogenic matrix or solid such as neon, argon or nitrogen, which are usually regarded as inert. If the interaction energy is predominantly van der Waals then the complex may be poorly defined geometrically, leading to broad IR spectral features. Ault has commented that if the

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binding energy is higher than ca. −10 kJ mol−1 then the complexes will be well defined with sharp IR spectral features [10]. In addition to stabilising the complexes for spectroscopic

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characterisation, controlled annealing or photolysis of the matrix allows for the investigation

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of subsequent reactions. A very beneficial consequence of the reduced inter-molecular interactions in the cryogenic matrices is that the line widths in IR spectra are often around 1 cm−1 or less, which allows for enhanced speciation and the observation of individual isotopomeric features which are invaluable in structure determination. E.g. a different isotope pattern will be observed for end-on and side-on bonded N2 if a mixture of 14N2, 14N15N and 15

N2 are used [11]. Likewise, by use of 16O/18O isotopic enrichment it is possible to identify if

ozone is bound via the central or terminal oxygen atom in a Lewis acid-base complex [12]. More reactive solid hosts such as O2 or CO are also employed. A more recent development is the use of solid para-hydrogen (p-H2) which can be regarded as a quantum solid of spherical particles [13-16]. An alternative is to use a frozen organic glass such as adamantane which only requires liquid nitrogen as a cryogen, and a particularly powerful adaptation of this is the use of rotating cryostats [17]. The solid organic hosts will usually have a considerable number of IR spectral features in the regions of interest, which can make spectral interpretation challenging. Although EPR combined with rotating cryostats is very useful to study the reactivity of atoms [18-21] it is less suited to closed shell systems as usually encountered in main group coordination chemistry. Other techniques for lowering the sample 5 Page 5 of 213

temperature, but which also require serious cryogenic apparatus include: cryosolutions [22] of liquid xenon [23], krypton [24], argon [25] and nitrogen [14]; helium nanodroplets [2630], supersonic expansion [31-33] as well as combined supersonic jet matrix isolation experiments [34]. Within the definition of coordination chemistry outlined at the start, this review will be

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limited to the discussion of discrete complexes formed as a result of a Lewis acid-base interaction between two molecular species. Whilst hydrogen bonding could be included, this

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is a whole field on its own [35] and these interactions have been largely excluded to keep the scope manageable, and to concentrate on what might be considered “real” coordination

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chemistry. However, a recent review outlining the use of a pKa slide rule for predicting H bond strengths from acid-base molecular properties should be noted [36]. The above

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definition also rules out the use of atoms as starting materials. These were reviewed previously in 1985 [37] and 2002 [38] and whilst there has been an explosion of main group, transition metal and lanthanoid atom matrix chemistry with the advent of laser ablation [39],

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there is not space to cover these in this review. Therefore, the review will cover the interaction of main group compounds to give complexes. Although there have been a few

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reviews of this field [10] they have been limited in scope so the coverage will include

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material from the 1980s through to the present (mid 2012). The coverage of the experimental matrix data is intended to be (reasonably) comprehensive, as is that of the complimentary gas phase experiments (which usually followed the matrix work), however, that for the supporting computational data is much more selective. As previously noted, it needs to be borne in mind that as a significant number of the complexes are weakly bound, the computational results may well be very dependent on the methodology (MP2, DFT etc) and basis sets employed, and that comparison between different groups is not always straightforward.

After a brief outline of the matrix isolation methodology, together with its advantages and disadvantages, the Lewis acid-base complexes formed by compounds from Groups 12 to 18 will be reviewed. The structure chosen has been to consider and tabulate the complexes according to the Lewis acids of each group, and their interaction with other main group Lewis bases. The data for the common bases (NH3, Me2O etc) have also been tabulated to enable easy comparison. To save undue repetition, unless otherwise stated it can be assumed that the

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complexes discussed are trapped/isolated in solid argon matrices at 10-14 K, and that they have been characterised by IR spectroscopy. Matrix Isolation As outlined above matrix isolation is an experimental method for stabilising species long enough for them to be spectroscopically characterised. The most common approach is to use condensed argon as the matrix host with the species of interest diluted in it at ratios of the

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order of 1:100 to 1:1000, but more concentrated and dilute mixtures have also been used. The more dilute the mixture, the higher the quality of the spectra in terms of bandwidth, but this

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has downsides in terms of detection. In the application to coordination complexes which may be produced in-situ, there is the issue of mobility of the acid and base within the rigid matrix.

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Atoms and small molecules can diffuse through the solid argon host on deposition and/or after annealing. The details of the experimental set-up required for matrix isolation has been

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well described by Dunkin [40]. In essence, there is a cryostat with a spectroscopically suitable substrate held at a low enough temperature to condense the matrix gas, which is often found to be half the melting point of the matrix solid. The most common forms of cryostat are

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He closed cycle systems, rather than flow cryostats, and with these it is now possible to reach ca. 4 K. It should be noted that the early work employed liquid hydrogen cryostats. The

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cryostat is mounted in a vacuum chamber which has a number of ports to allow for

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spectroscopy, sample inlet, photolysis etc. The quality of the matrix deposit is very dependent on the deposition conditions used, and each laboratory has its own preferred method, but in general the matrix gas is bled into the vacuum chamber at a few mmol hr−1 until the spectra are of sufficient intensity and quality. Pulsed deposition has been favoured by some workers. The most common spectroscopic technique used is infrared as this gives excellent speciation and product identification. The sharp bands (fwhm ~1 cm−1 or better) that result from trapping in a non-interacting medium allow for the use of both natural and induced isotopes that are crucial to determining the number of atoms involved, as well as the bond angles and mode of bonding. Although it is attractive to obtain Raman data from matrices, this has always been much more of a struggle due to sensitivity and radiation damage. UV-vis spectroscopy is useful for atomic work, but because of the broad bands in molecular species is usually not sufficient without other supporting data. EPR is very powerful for open shell systems such as organic and main group radicals and transition metal compounds with unpaired electrons. NMR has never really taken off. There are niche specialists using Mössbauer [41-43] and X-ray absorption spectroscopy [44-50]. 7 Page 7 of 213

Sample preparation is usually the key as the materials of interest need to be introduced into the vacuum chamber with an appropriate vapour pressure. For the vast majority this requires either heating or cooling or pre-mixing with argon. In this review the vast majority of the materials have sufficient vapour pressure at room temperature so that they can be pre-mixed with the argon using standard manometric techniques before deposition. In order to control

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the extent of reaction prior to deposition on the cold substrate, Ault has developed the singlejet, merged-jet, twin-jet and concentric-jet terminology to describe the experimental

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configuration [51, 52]. In a single-jet experiment, the reagents and argon are mixed in a single container and bled into the vacuum chamber through an appropriate needle valve. This

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pre-supposes that either there is no significant reaction at room temperature or that the product(s) are volatile. In the merged-jet experiment the individual reagents are diluted in the

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matrix gas in their own containers, but there is a mixing/reaction zone of up to 100 cm, that can also be heated (sometimes called a pyrolysis zone), to enable some reaction on the millisecond timescale to occur before deposition. In contrast, in the twin-jet approach, the

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individual reagents diluted in argon are introduced into the vacuum chamber as close as possible to the cold substrate (few mm) in order to limit any reaction before condensation.

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The concentric-jet allows some in-situ control between merged- and twin-jet methodologies.

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Clearly, the more reactive reagents are going to require the twin-jet approach if the early complexes are going to be identified, whereas the merged-jet is more appropriate for the less reactive systems, and the single-jet where it is believed that the complex is only formed at low temperature. In the vast majority of case described below both merged- and twin-jet were attempted, and the optimal one chosen. In nearly all cases subsequent photochemistry was carried out, but as this usually results in elimination reactions, they will not be discussed in detail. A number of reviews [53-62] and books [40, 63-67] have been published covering the matrix isolation field over the years.

Lewis acid base chemistry at low temperature Molecular Lewis acid-base complexes are generally identified in matrices using IR spectroscopy. The signature of complex formation is vibrational modes that are shifted or perturbed (slightly) from those of the parent donor and acceptor molecules. The observation of new spectral features well removed from those of the parents is much more likely to be 8 Page 8 of 213

characteristic of compound formation via elimination etc. In general, it has been noted that complex formation is characterised by more substantial perturbation to the acid sub-unit (as shown in Figure 1, where the BH3 unit loses its planarity) with a relatively minor perturbation to the base sub-unit [68-70]. It also needs to be borne in mind when the complex formation is actually occurring. For many complexes this will be a gas phase reaction prior to deposition/condensation on the cold window, and this can be controlled by use of merged-,

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twin- and concentric-jet methods. For others it will occur during the condensation process itself as the temperature is reduced to allow the bond enthalpy term to dominate over the

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entropy term. It should also be borne in mind that there is the possibility of the formation of apparent complexes due to the species being trapped adjacent to each other in the matrix

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cage, which does not have a gas phase analogy. The shifts in the vibrational modes in the complexes compared to the parents can either be to higher wavenumber (blue-shifted) or

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lower wavenumber (red-shifted). Whilst this can sometimes be related to the change in population of bonding, non-bonding and anti-bonding orbitals, there are often other effects

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such as change of mass and bond angle that make this simplistic approach dangerous. Some of the bases are ubiquitous (e.g. NH3), but many have been used in matrix studies, and

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and proton affinities [71, 72].

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Table 1 collects together the most important bases, together with their gas phase basicities

Group 12

1.1.

Introduction

Although zinc chemistry is often considered together with that of the bona fide members of the 3d transition series, the Group 12 elements are traditionally regarded as part of the main group as the divalent oxidation state has a full d-manifold. There has been both experimental and computational work to try to identify the existence of +3 or +4 oxidation state compounds of Cd and Hg. The stability of HgIV, and in particular that of HgF4, has been indicated by computational work for nearly twenty years [73-80]. The higher oxidation states are stabilised by relativistic effects, especially on the Hg 6s electrons. A report in 1976 of the electrochemical generation of a short-lived HgIII intermediate [81], has been regarded as inconclusive as the reported ESR spectrum was not that expected for a d9 HgIII compound [78]. Despite the significance of HgIV on the understanding of the periodic table there was a distinct lack of experimental evidence for HgIV apart from some unpublished activity by a 9 Page 9 of 213

number of workers. In 2007 Wang et al. published a paper describing the first experimental IR characterisation of HgF4 in neon and (possibly) argon matrices [82]. This was supported by computational data and provided the first experimental evidence for Hg(IV) [82]. However, Rooms et al. published the results of a detailed matrix investigation using IR, UVvis-NIR and Hg L3-edge XAFS techniques which found no convincing evidence for HgF4 in argon matrices [50]. This did identify a new Hg...F2 complex, as well as providing the first

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experimental bond length of 1.94(2) Å for monomeric HgF2. It needs to be borne in mind that calculations involving mercury are very challenging because of the large number of electrons

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and relativistic effects and therefore, the calculated values of the vibrational modes of HgF4 are very dependent on the method and basis set employed. More recently a mercury(II)

Complexes of dialkyl zinc, cadmium and mercury

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1.2.

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oxyfluoride has been reported [83].

The dialkyl compounds of the Group 12 compounds are common precursors in the

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fabrication of II-VI semiconductors, and have therefore been the subject of many studies [8486]. Whilst the vast majority of the studies described in this review utilise matrix isolation techniques, the Group 12 dialkyls have been studied by both matrix and thin film methods to

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with alumina surfaces [87].

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provide greater insight. The matrix data has provided insight into the reaction of dimethylzinc

1.2.1.

Me2Zn complexes

The reactions between Me2Zn and group 15 and 16 hydrides, NH3, PH3, AsH3, H2O, H2S, H2Se were studied using both matrix and thin film methods [88]. Due to the differences in reactivity, the H2S, H2Se and NH3 experiments were carried out using the twin-jet method, whereas for PH3 and AsH3 the merged-jet technique was used with a 120 cm reaction zone. The reaction with water was a result of trace levels of H2O in the system. (Trace levels of water are ubiquitous in essentially all matrix experiments because of the cryopumping provided by the cryostat and it is usually not possible or feasible to bake the vacuum system.) The reaction between Me2Zn and NH3, PH3 and AsH3 led to the formation of H3N•ZnMe2, H3P•ZnMe2 and H3As•ZnMe2 1:1 complexes in argon matrices. In the case of high concentrations of NH3 there was also evidence for a 2:1 complex, (H3N)2•ZnMe2. Although there was evidence of complex formation with both H2S and H2Se in argon matrices, it was not possible to identify the stoichiometry. In the thin film experiments there was evidence for 10 Page 10 of 213

both 1:1 and 2:1 complexes for NH3 and PH3, but only a 1:1 complex in the case of AsH3. As for the matrix experiments, there was evidence for a single product formed between Me2Zn and H2S and H2Se, but its stoichiometry could not be determined [88]. For Me2S and Me2Se there was no observable reaction with Me2Zn using the twin-jet matrix approach, but product bands were observed with a 120 cm merged-jet apparatus for Me2S, Me2Se, Me2O, Me2CO and Me3N [89]. In argon matrices there was evidence for the formation of a single product

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(1:1 complex) for Me2S, Me2Se and Me2O, but two products (1:1 and 2:1 complexes) for Me2CO and Me3N. The mole ratio of the two reactants could be more easily controlled in thin

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films and two distinct sets of product bands were observed for Me2S, Me2Se, Me2CO and Me3N (no thin film data was presented for Me2O). When the Me2Zn/Me2S and Me2Zn/Me2Se

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films were warmed to 90 K, the product bands sharpened and intensified, but by about 180 K these bands began to disappear leaving a spectrum of solid Me2Zn. For Me3N and Me2Zn

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very intense product bands were observed on deposition, but their temperature dependence was less marked due to the high reactivity prior to deposition, except that they were sufficiently stable to persist up to room temperature. Room temperature pyrolysis of Me2Zn

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and MeOH or MeSH using merged-jet and both argon matrices and thin films resulted in the formation of Me(H)O•ZnMe2 and (Me(H)O)2•ZnMe2, with vibrational modes shifted slightly

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from the parent molecules [90]. In the case of MeSH no new bands were observed in argon

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matrices after room temperature pyrolysis in merged-jet mode, but Me(H)S•ZnMe2 and (Me(H)S)2•ZnMe2 were both observed in thin films below 100 K, but only (Me(H)S)2•ZnMe2 was stable above 100 K. Higher temperature pyrolysis during merged-jet deposition with both MeOH and MeSH gave rise to further products [90]. Whilst there is computational data on Me2Zn [91], the published computational literature on the complexes appears limited to the H2Se•ZnMe2 complex [92]. This indicates that the H2Se•ZnMe2 complex is 5.4 - 13.8 kJ mol−1 less stable than the precursors and that the equilibrium population of the complex is ca. 0.4% at room temperature. The Se-Zn bond in the complex is a result of donation from the lone pair on the Se into an empty 4p orbital on the Zn, with the H2Se unit being approximately perpendicular to the Se-Zn bond. The 2:1 adduct is even less stable (ca. +80 kJ mol−1) [92]. Shown in Figure 2 is the calculated (G09, B3LYP, 6-311G+dp) structure of the H3N•ZnMe2 complex which has a N-Zn bond length of 2.438 Å and C-Zn-C bond angle of 163.2°. (All the other structural representations in this review have been optimised at this level, and are in broad agreement with the published data.)

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On initial deposition of Me2Zn and O3 using a twin-jet approach, an initial cage pair complex was observed, together with the oxidation product, MeOZnMe [52]. On UV photolysis, the former decreased, whilst the latter increased. Other than its reporting, there was very little discussion of the structure of the initial complex, other than it was formed by the two reagents, and could be either a be due to a distinct chemical interaction, or simply cage “late” products) were observed including H2CO, MeOH and C2H6.

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pairing in the matrix. When merged-jet deposition was used, different products (known as

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The data for the Me2Zn complexes are collected together in Table 2, together with the proton affinities [71, 72] of the bases employed. From the plots in Figure 3(a) it is clear that in the

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case of the argon matrix data there is a good correlation between the ZnC2 asymmetric stretching mode in the complex and the proton affinity of the base, even with a selection of

1.2.2.

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reasonable correlation, but there is more scatter.

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2nd, 3rd and 4th period donor atoms [89]. In the case of the thin film data, there is still a

Me2Cd and Et2Cd complexes

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Almond has carried out a number of investigations of complex formation involving Me2Cd

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and Et2Cd under a variety of conditions [93-102]. Of particular relevance are those involving Me2Cd and Et2Cd which formed weakly bound complexes of the form (Et2Te)x•CdMe2 and (Me2Te)x•CdEt2 with TeMe2 and TeEt2 in the gas phase before being trapped in argon matrices. Whilst there was reasonable evidence for the formation of complexes with Me2Cd, it was much more compelling with Et2Cd where many more of the bands belonging to the complex were identified. Although it was not possible to tell whether x = 1 or 2 it was possible to confirm that the C-Cd-C unit was non-linear due to the change in the intensity ratio of the symmetric and asymmetric νC-Cd-C modes [85]. Using thin films at 77 K Almond has also shown the formation of complexes between Me2Cd and

the

potentially

chelating

Me2NCH2CH2N(Me)CH2CH2NMe2

ligands (PMDETA),

Me2NCH2CH2NMe2 MeOCH2CH2OMe

(TMEDA), and

MeOCH2CH2OCH2CH2OMe (diglyme). The work on TMEDA built on previous gas phase experiments [95]. For the phosphorus donor, Me2PCH2CH2PMe2, and the cyclic oxygen donors, 1,4-dioxane and 1,4-thioxane and, there was a much weaker interaction as there was 12 Page 12 of 213

for the sulfur donors, diethyl sulfide, ethylene sulfide (thiirane), 1,4-dithiane and propylene sulfide [100]. It was thought that all the complexes had 1:1 stoichiometries. For propylene sulfide the weak interaction was only observed in the cold solid (77 K) and not in the gas or liquid phases [101]. Almond also found no interaction between Me2Cd and OCS in a 1:1 molar ratio either in the gas phase at room temperature, or in the solid phase at 77 K [100, 102]. At higher temperatures gaseous products (CH4, CO, C2H6) were observed, and in a

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commercial MOCVD apparatus high quality layers of CdS were obtained from Me2Cd-OCS

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mixtures[102]. These data are summarised in Table 3.

Ault has carried out studies involving Me2Cd with H2S, H2Se, NH3, PH3 and AsH3 [103],

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Me2S, Me2Se, Me2CO and MeOH [104] and O3 [105], analogous to the Me2Zn experiments discussed above [88-90]. 1:1 complexes were identified for all of the hydrides in both argon

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matrices and thin films, with some evidence for 2:1 complexes involving NH3. The shift in asymmetric CdC2 mode in the H3N•CdMe2 complex (Table 3) indicated that this was weaker bound than the H3N•ZnMe2 complex. This was also confirmed by the relatively small shift in

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ν2 mode of NH3, and correlates well with the proton affinities of the bases used [104]. The shift in the asymmetric CdC2 mode was smaller for both H2S and H2Se than NH3, and for the

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PH3 and AsH3 complexes, no asymmetric CdC2 modes could be identified for the complexes.

Ac ce pt e

Whilst a shifted hydride stretching mode could be identified for the H2S, PH3 and AsH3 complexes this was not the case for the H2Se complex due to overlapping bands. For H3N•CdMe2 the use of thin films allowed for the formation of 1:1 and 1:2 complexes at low temperature (14 K) the ratio of which could be controlled by varying the thin film composition. On warming the thin film, the product concentration increased until 170 K when dissociation started, and by 200 K only a spectrum of Me2Cd was observed. For H3N•CdMe2 essentially the same spectrum was obtained from both merged and twin-jet deposition. Only product bands close to the parent hydrides were observed for the reaction between Me2Cd and PH3 or AsH3 using matrices or thin films in merged-jet mode [103]. For Me2S, Me2Se, Me2CO and MeOH, 1:1 complexes were observed with Me2Cd in all cases in both argon matrices and thin films [104]. For Me2S and Me2Se there was also evidence of 2:1 complexes in argon matrices with relatively high base concentrations. It is clear from both Almond’s and Ault’s work that although Me2Cd is a very weak Lewis acid, the complexes are sufficiently stable at low temperatures both in solid argon and as thin films to be spectroscopically characterised. The perturbed νCO mode of acetone shifted 3 cm−1 13 Page 13 of 213

in the Me2Cd complex [104] compared to 11.3 cm−1 for the 1:1 Me2Zn complex [89] indicating that Me2Cd is a weaker Lewis acid than Me2Zn. The data for Me2Cd complexes in both argon matrices and as thin solid films are collected together in Table 3. The plots in Figure 4 show that as for the Me2Zn complexes, there is a good correlation between the position of the asymmetric CdC2 stretching mode and the

ip t

proton affinity of the base for the Me2Cd complexes, although as there are only a few

cr

examples of each sort of base, this needs to be interpreted with a little caution.

1.3.

us

Group 13

Introduction

There has been a recent renaissance of interest in Group 13 chemistry, of which

an

cryochemistry, and in particular matrix isolation has played an important role. This includes the identification of gallane [106, 107], alane [108, 109], indane [110], thallium hydrides

M

[111], as well as boranes [112]. There is a substantial body of work on group 13 chemistry, much of it has been reviewed, with some containing more matrix work [113] than others

d

[114-117].

Ac ce pt e

As outlined in the introduction the emphasis of this review is on the complexes/adducts formed between two molecules via a coordinate/dative bond rather than compound formation following reaction or elimination. The first examples will involve halides, before considering hydride and methyl complexes of the group 13 elements.

1.4.

Boron halide complexes

BF3 is a very well known Lewis acid and there have been many experimental as well as computational studies of the complexes it forms with a wide range of bases. One of the advantages of matrix isolation in this case is that it enables the study of a wide range of complexes under very similar conditions, which allows for the evaluation of the Lewis acidbase behaviour, and the data for these is collected in Table 4. The BF3 Lewis acid-base complexes have found use in

10

B enrichment/separation plants within the nuclear industry.

For example Me2O•BF3 was developed for 10B/11B separation during the Manhattan project, with the nitromethane complex receiving more recent interest [118]. 14 Page 14 of 213

BF3 can also act as a base towards itself, with the formation of dimeric species which complicate the spectra, especially in argon matrices. The structure of the BF3 dimer was initially believed to be isostructural with diborane with tetrahedral B and D2h symmetry [119, 120]. However, more recent computational [121] and experimental [122] work indicates that the dimer has C2h symmetry, (where the planarity of the BF3 units is maintained) and that the

ip t

D2h structure was a transition state connecting the C2h dimers. The relatively small shifts in the vibrational modes between the monomer and dimer (Table 4) indicate a weak electron-

cr

donor interaction, which was computed to be about −4 kJ mol−1 [121].

us

The OC•BF3 and N2•BF3 complexes in solid argon have νBF modes red-shifted by 9 and 4 cm−1 respectively, indicating a weak interaction [123] as expected from the B-C and B-N

an

bond lengths of 2.886 and 2.875 Å [124]. More recent work has identified the asymmetric νBF modes in the OC•BF3complex at 1488.3 (11B) and 1435.9 (10B, with site effect at 1432.2 cm−1) in argon and at 1490.1/1487.6 cm−1 (11B) and 1436.7/1434.2 cm−1 in solid nitrogen

M

[125]. The νCO mode in OC•BF3 in the original work was at 2150.5 cm−1 blue-shifted by 12 cm−1 from that of “free” CO [123]. Later work has identified this mode at 2154.1 cm−1 (with

d

a shoulder at 2152.1 cm−1) in nitrogen and 2152.4 cm−1 (with a shoulder at 2151.1 cm−1) in argon matrices [125]. There is a subsequent computational report confirming weak bond

Ac ce pt e

energies of 4 -10 kJ mol−1 [126] in good agreement with the experimental value of 7.5 kJ mol−1 obtained from liquefied argon studies [127]. The absence of any new bands in matrices containing BF3 and CO2 was interpreted as implying that the interaction between them is too weak and that it cannot compete with a second BF3 molecule for the interaction site with the first BF3 unit [128]. Likewise, there was no evidence for complex formation between cyclopropane and BF3, BCl3 or BBr3 [129]. 1.4.1.

Boron halide complexes with Group 15 donor ligands

The H3N•BF3 complex is often regarded as the archetypal coordination compound containing a dative bond, and its structure is given in Figure 5. Despite the BF3 amine complexes being well known room temperature solids [130-136], there was a dearth of vapour phase data until 1982 when H3N•BF3 was first characterised by matrix isolation techniques [137]. The spectra of the H3N•BF3 complex in a KBr disc and an argon matrix were compared and whilst there were some shifts observed, especially in the very sensitive NH3 symmetric deformation mode and the BF3 anti-symmetric stretching mode, these were explicable in terms of different 15 Page 15 of 213

environments [137]. The large red-shift of ca. 200 cm−1in the BF3 asymmetric stretching mode and the large blue shift of 380 cm−1 in the symmetric deformation mode of NH3 were indicative of a strong interaction between the NH3 and BF3. It has been noted that the calculated interaction energies of the H3N•BF3 and Me3N•BF3 complexes vary widely depending on the computational method used, but Ford has shown that there is a linear relationship between interaction energy and calculated shift in the asymmetric stretching

ip t

mode for the amine and methyl substituted [138]. A gas phase microwave study in 1971 and 1977 had given a value of 1.636(4) Å for the B-N bond in Me3N•BF3 [139, 140].

cr

Subsequently, a fast-mixing pulsed-nozzle FT microwave experiment provided gas phase data in for H3N•BF3 1991 with a B-N bond length of 1.59(3) Å [141], but this was later

us

revised to 1.673 Å after consideration of shrinkage effects to [142]. The solid state value is 1.60(2) Å [133]. Stark effect measurements on the gas phase complexes have yielded the

an

following dipole moments: H3N15•BF3, 5.9027(93) D; Me315N•BF3, 6.0157(76); HC15N•BF3, 4.1350(73) D [143].

M

The matrix data [137] were then used to characterise the H3N•BCl3 and H3N•BBr3 complexes for the first time, confirming that they are the short-lived intermediates formed prior to

d

hydrogen halide elimination which is observed at room temperature [144]. The H3N•BX3

Ac ce pt e

yield decreased with increasing mass of halogen. The Lewis acidity of the boron halide was found to correlate well with the position of the symmetric (umbrella) deformation mode of NH3 as the larger the blue shift, the stronger the Lewis acid the NH3 is coordinated to. This changed from 970 cm−1 in free NH3 in argon to 1309 cm−1 for the H3N•BF3 complex in an argon matrix (1438 cm−1 in the solid state) and 1352 cm−1 for H3N•BCl3 complex. No feature could be assigned to this vibrational mode in the H3N•BBr3 data, either because of its inherent weakness, or more likely being masked by the overtone bands of BBr3 in the 13001400 cm−1 region [137]. There was no evidence of association between H2 and BBr3 in an argon matrix prior to UV photolysis which resulted in the formation of HBBr2, H2BBr and HBr [145]. The data for the matrix isolated BCl3 and BBr3 complexes are collected together Table 5. The crystal structure of H3N•BCl3was reported in 1995, with a N-B bond length of 1.579(4) Å [146]. The IR spectra of pyridine, quinoline and acridine BCl3 complexes have been reported in KBr discs [147]. The H3N•BCl3 and H5C5N•BCl3 complexes have subsequently been investigated computationally by fragment mode analysis [148].

16 Page 16 of 213

The boron trifluoride nitrile complexes have been referred to as partially bonded molecules, and lie midway between the H3N•BF3 complex with short B-N bonds (ca. 1.6 Å) and a near tetrahedral boron geometry, and N2•BF3 which is a true van der Waals molecule with a long B-N distance (2.88 Å) with a planar BF3 unit [149]. In addition the B-N bond lengths have been found to be dependent on the phase of the sample, with relatively short B-N distances observed for solid state samples and longer B-N distances for gas phase samples. As the local

ip t

environment of the trapped complex can be changed by changing the matrix gas, and there were no “gas phase” vibrational data, the nitrile complexes have been reasonably well studied

cr

by matrix methods. When MeCN•BF3 was initially investigated a band observed at 2380 cm−1 was assigned to the νCN mode in D3CCN•BF3 (CD3CN was used to try to ameliorate

us

some of the spectral complexity) [150]. As this was close to the solid state value of 2373 cm−1 and that for observed in benzene solution, this was interpreted to mean that in relatively

an

BF3 rich argon matrices the MeCN•BF3 complex was present with a short B-N bond. However, Phillips has re-investigated the MeCN•BF3 complex in argon [151], neon [152] and nitrogen [152] matrices, as well as by single crystal X-ray diffraction [153] and quantum

M

chemical methods [153, 154]. In these studies the νCN mode in MeCN•BF3 was assigned to a band at 2365 cm−1 in Ar (2363 cm−1 in D3CCN•BF3) [151], a doublet at 2352 and 2356 cm−1

d

in Ne [152] and a band at 2368 cm−1 in N2 [152] which were ca. 17 cm−1 red-shifted from

Ac ce pt e

those reported previously [150]. Phillips claims that the 2380 cm−1 band is a spectral component of BF3/Ar rather than part of the complex [151]. It should be noted that the intensity of the bands due to the MeCN•BF3 complex in N2 matrices were lower than in either Ar or Xe, presumably because of the competition from the weak van der Waals N2•BF3 complex [123, 155]. Hattori et al. have also carried out a series of matrix isolation experiments with MeCN and BF3 mixtures in Ar, Xe and N2 matrices [156], and these are in agreement with those of Phillips, [151, 152] rather than the initial work [150]. The supporting calculations on MeCN•BF3 [157] revealed two solutions on a very flat potential, with B-N distances of 1.8-1.9 and 2.3 – 2.4 Å, with the former resulting from basis sets with diffuse functions, and the latter from basis sets without diffuse character. More sophisticated calculations have yielded a B-N distance of 1.95 Å [158] in better agreement with the gas phase value of 2.011(7) Å [159]. It is commented that the earlier discrepancy may be due to large amplitude vibrational motion in the B-N stretching coordinate. Calculations on the related HCN•BF3 complex [154] give a binding energy of −23.8 kJ mol−1 and dipole moments in good agreement with the gas phase data [143], but it was noted that there is an absence of gas phase (or matrix) vibrational data for this complex. Phillips et al. have also 17 Page 17 of 213

investigated the C6H5CN•BF3 and Me3CCN•BF3 complexes using matrix IR spectroscopy, crystallography and computationally [160]. For C6H5CN•BF3 the νCN mode was observed at 2320 cm−1 for both 10BF3 and 11BF3 complexes which were ca. 8 cm−1 lower than those of the cold solid. The Me3CCN•BF3 complex displayed bands at 2323/2327 and 2323 cm−1 for the 11

B and

10

B components which were ca. 15 cm−1 lower than those for the solid, indicating

that the medium in which the complex is formed has an effect on the B-N bond but not as

ip t

much as for acetonitrile complexes [160]. Calculations indicated that the larger organic substituents stabilised the B-N potential close to 1.8 Å, so that the condensed phase effects on

cr

their structural properties were much less pronounced than for MeCN•BF3. Previous computational work of Ford had B-N distances of 2.4 – 2.6 Å for HCN, FCN, NCCN and

us

HC3N complexes with BF3 with interaction energies of −15.5, −13.5, −11.4 and −18.4 kJ mol−1, respectively [161]. The binding energy for the N2•BF3complex was calculated as

an

−2.27 kJ mol−1 [161]. A summary of the values of the νCN modes for these complexes is given in Table 6.

M

A combined matrix-isolation mass-spectroscopy study of the reaction of BF3 with Me3N•A1H3 and Me3N•GaH3, showed that in the gas phase the reactions proceeded via a

d

hydride-halide exchange with the formation of BF2H and NMe3 and no evidence for AlH3 or

Ac ce pt e

GaH3, rather than through a Lewis acid displacement reaction [162]. The phosphorus halide-boron halide complexes, X3P•BY3 (X = Cl, Br, I; Y = Br, I), are sufficiently stable to allow for their characterisation by Raman and IR spectroscopy, solid state 31P NMR spectroscopy, and the molecular structure of Br3P•BBr3 was determined by Xray diffraction, supported by DFT calculations [163].

1.4.2.

Boron halide complexes with Group 16 donor ligands

The interaction of H2O and BF3 is of great technological importance, and hence there have been a number of studies, including a crystal structure [164, 165]. The matrix isolated H2O•BF3 complex has been studied in Ar [166, 167], N2 [128] and Ne matrices [168] over the years, and the data are given in Table 4. From these it is clear that the N2 data of Ford et al. [128] are different from the rest and whilst some of this may be due to matrix shifts, they have also admitted subsequently that some of the data was miss-assigned [169]. In the initial report of the argon matrix data only the values of the asymmetric B-F stretching modes at 18 Page 18 of 213

1313, 1285 and 1241 cm−1 were included [166]. The 1285 and 1241 cm−1 bands were assigned to the ν3 or ν11 modes of the

11

B complex (the double degeneracy of the parent ν3

mode is lifted if the complex no longer has a three-fold axis), with the 1313 cm−1 band belonging to a 10B complex (the second mode was thought to lie under the 1285 cm−1 band). A subsequent report included many more vibrational modes [167], including four bands (1314.6, 1283.9, 1264.7 and 1251.9 cm−1) in the asymmetric B-F stretching region. Whilst

ip t

these were similar to those of Ault [166] it is not clear to which isotopomer or vibrational mode (ν3 or ν11) of the H2O•BF3 complex they belong. Whilst these were red-shifted ca. 150

cr

cm−1 from the parent modes, the symmetric deformation mode (ν6) for both 10B and 11B were reported to be very close (essentially identical) in the H2O•BF3 complex and the parent BF3,

us

which was interpreted to mean that the interaction was very weak. However, there was a band at 654.9 cm−1 which was not readily assigned. This lack of accuracy in assigning the

an

vibrational modes to the appropriate isotopomer is carried over to the subsequent computational report [170]. It appears that the calculated difference between the ν3 and ν11 modes is very similar to that between the

10

B and

11

B components in Aults’s original data

M

[166]. These data has been re-evaluated and used in subsequent reports [168, 169] where the 1314.6 and 1283.9 cm−1 bands have been assigned to the ν3 and ν11 modes of the 11B complex

d

to help in the assignment of the Ne matrix data [168]. In the Ne matrix experiments, no

Ac ce pt e

values of the BF3 symmetric deformation were given, as this work was much more focused on obtaining the first vibrational data for BF2OH after the passage of BF3 and H2O through a long stainless steel line [168]. Yeo and Ford subsequently assigned the four νBF peaks in Ball’s data [167] to the ν3 and ν11 of both 10B and 11B complexes [169]. Whilst this appears sensible on the basis of Ault’s data [166] it does raise the question that despite the importance and apparent simplicity of the H2O•BF3 complex, the matrix IR data is possibly not as secure as it could be. The most recent calculations indicate that H2O•BF3 adopts an eclipsed conformation with the O interacting with the B and the H and F aligned and a binding energy of −56 kJ mol−1 [169]. H2S•BF3 however is staggered with a S•B interaction energy of −6 kJ mol−1 [169].

As was the case for the H3N•BF3 complex, the Me2O•BF3 complex was well known, but the corresponding BCl3 and BBr3 complexes were not, with only brief mentions, prior to the matrix study [166]. The Me2O•BF3 matrix data was in good agreement with earlier thin film data [171], but sharper bands (1 - 3 cm−1 fwhm compared to 50 cm−1 for the thin film) were observed indicating the benefit of using matrix isolation. This also represented the first 19 Page 19 of 213

spectroscopic characterisation of Me2O•BCl3 and Me2O•BBr3 [166]. As in the NH3 complexation studies, the Lewis acidity of the boron halide correlated well with the symmetric C-O-C stretching mode of the coordinated ether. For Me2O•BF3 the C-O-C modes were obscured by other BF3 bands, but deuterium enriched experiments indicated that the symmetric C-O-C mode was expected around 950 cm−1. For Me2O•BCl3 the C-O-C modes were at 999 cm−1 (asymmetric) and 903 cm−1 (symmetric), and in Me2O•BBr3 they were at

ip t

980 cm−1 (asymmetric) and 887 cm−1 (symmetric). A recent crystallographic report gives the B-O bond length in Et2O•BCl3 as 1.543(2) Å (longer than that in the tetraalkoxyborates), with

cr

the B atom tetrahedrally coordinated [172]. In addition, the H2O•BF3 and H2S•BF3 complexes were identified, even though the Lewis basicity of these is lower than Me2O and NH3, and at

us

room temperature hydrolysis readily occurs [166]. In Ford’s investigation of the Me2O•BF3 and Me2S•BF3 complexes [173, 174], the experimental data were in good agreement with

an

those of Ault for Me2O•BF3 [166] and these together with the calculations indicate that the interaction between BF3 and Me2O is substantially greater than that between BF3 and Me2S

M

[173, 174].

In the N2O•BF3 complex the experimental data [175] was consistent with the formation of a

d

1:1 complex as observed in the gas phase [176], and the calculations indicated that the N2O

Ac ce pt e

base was bound via the O and that the preferred geometry was eclipsed with the ONN unit aligning with one of the BF bonds in a nearly parallel arrangement with a binding energy of ca. −5 kJ mol−1 [161, 175, 177]. The OSO•BF3 complex had been studied by microwave spectroscopy and the computational methods [178] prior to the matrix and computational study [179]. The calculational results indicated that the geometry was either cis-trans or trans-cis with respect to the arrangement of the atoms in the F-B-O-S-O chain. The B···O was essentially perpendicular to the BF3 plane, and in the cis-trans isomer the S atom lies above one of the Fs, whereas in the trans-cis isomer the S lies above, but between the other 2 F atoms. The matrix experimental data indicated that the trans-cis conformer was a better candidate for the preferred structure [179], which is in keeping with the calculated binding energy of −9.5 kJ mol−1 which is 1 kJ mol−1 more favourable than the cis-trans isomer [178], although tunnelling was observed between the two structures in the gas phase microwave spectrum which yielded a value of ca. 2.6 Å for the B-O distance [178]. Perfluoroalkylethers have found application in lubrication of both computer magnetic media and spaceflight components. The possible decomposition routes when interacting with Lewis 20 Page 20 of 213

acids has been investigated by matrix isolation techniques [167, 180]. When BF3 was allowed to react with ethers and fluoroethers in argon matrices, the IR spectra clearly indicated the formation of a Et2O•BF3 complex, as well as a weak (F3CH2C)2O•BF3 complex, but there was no evidence for the formation of a (C2F5)2O•BF3 complex [167, 180]. Supporting calculations indicated that the fluorine atoms decreased the electron localisation on the oxygen atom, showing that the interaction between BF3 and O(CH2CF3)2 (−33 kJ mol−1) was

ip t

about half of that in Et2O•BF3 (−80 kJ mol−1), and that there was no interaction between BF3

cr

and O(C2F5)2 [167, 180].

Recently, the Me2CHF•BF3 complex has been studied both experimentally (in solid Ne) and

us

computationally with a 2.3 Å B-F bond and a binding energy of −27 kJ mol−1 [181].

transfer technique was not successful [182].

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The preparation of the [BF3O]2- anion via the reaction of BF3and Tl2O using the oxide

M

It should be noted that Ford has carried out a detailed experimental and theoretical study of the boron halides spanning a period of over twenty years. The experimental work includes the

d

study of the BF3 dimer [120, 122], H2O•BF3 [128], OC•BF3 [125], CO2•BF3 [183], H3N•BF3

Ac ce pt e

[184], Me2O•BF3 and Me2S•BF3 [173, 174], N2O•BF3 [175], and OSO•BF3 [179] complexes. These data are collected together in Table 4. This experimental work has been supported by extensive and detailed calculations [121, 138, 155, 161, 169, 177, 185-192] with one of the conclusions being that the geometry of the lowest energy complex can be dependent on basis set employed [169], even though some calculations used BSSE [155]. As some of the complexes are fairly weakly bound, this is not too surprising, but needs to be borne in mind when assessing the computational results. The large body of computational work includes: the structures and spectral motifs of BF3, BClF2, BCl2F and BCl3 [185] and the BF3 dimer [121, 193]; the structures and energetics of BF3 complexes with linear nitrogen donors such as N2, N2O, HCN, FCN, C2N2 HCCCN [155, 161, 177]; complexes with HF, HCl, HBr, F2, Cl2 Br2, ClF, BrF and BrCl [192]; the H3N•BF3 complex and related methylamine complexes [138, 184]; H3P•BF3 and related methyl complexes [189]; the F3N•BF3 complex [194]. A series of calculations have also been reported involving the complexes of BF3 with Group 16 donors such as CO2 [187], Me2O and Me2S [186], H2O and H2S [169], MeOH, MeSH and related bases [188] and H2CO and its analogues [191]. There are also calculations on the MeF•BF3 and MeCl•BF3 complexes [190]. 21 Page 21 of 213

A graphical summary of the interaction energies from these calculations is shown in Figure 6 and these can be compared to the experimental IR data in Table 4 and the proton affinities and gas phase basicities in Table 1. The advantage of Ford’s work is that it is directly related to the matrix work and uses a consistent computational approach, which makes direct comparison more straight-forward. For the BF3complexes it is clear that the largest

ip t

interaction energy is for the amine complexes and that this increases from NH3 through the partially substituted amines to NMe3 [138]. Similar behaviour is observed for the analogous

cr

phosphine bases, except that there is a much larger reduction in interaction energy for PH3 compared to the methyl substituted bases. However, his comment “Previously reported

us

estimates of the interaction energies of H3N•BF3 and Me3N•BF3 vary widely, depending on the method of computation, so comparison of our values with others is not instructive” [138]

an

needs to be remembered. For example for H3N•BF3 he quotes an interaction energy of −160.35 kJ mol−1, a reaction enthalpy of −150.47 kJ mol−1 compared to a reported experimental reaction enthalpy of -170.8 kJ mol−1 [138], although he had previously reported

M

a value of -60.7 kJ mol−1 for the interaction energy in H3N•BF3 [184], but there does not seem to be any explanation given for this large difference. For the Me3N•BF3 complex the

d

values were -217.23, -207.85 and -112.3 kJ mol−1 [138]. In contrast Leopold quoted values of gas phase binding energies of 80.2 and 137.5 kJ mol−1 for H3N•BF3 and Me3N•BF3,

Ac ce pt e

respectively [149]. Frenking gave bond energies (D0) of 81.9 and 127.5 kJ mol−1 for H3N•BF3 and Me3N•BF3, respectively [6].

Figure 7 shows the plots of the BF3 asymmetric stretching and symmetric deformation modes for the 11B complexes isolated in argon matrices (except for the N2O complex where only N2 data was available) versus the proton affinity of the base. In the uncomplexed BF3 the ν3 (asymmetric stretching) mode was at ca. 1443 cm−1. The value of the ν2 (symmetric deformation) mode is reported from 654 to 674 cm−1 by the same group [122, 173-175, 183, 184], but a value of 676 cm−1 has been observed by others [123, 195]. For the asymmetric stretching mode derived from the ν3 mode, two sets of data are presented, those that retain a C3 axis and hence a single degenerate mode, whereas for complexes without a C3 axis the mean of the two modes are plotted. Perhaps surprisingly this seems to work. There appear to be two sets of complexes within the asymmetric stretching mode data, one set with small shifts in the asymmetric stretching mode, and a second set with much larger shifts. The former include the complexes with N2, CO, CO2, N2O, SO2 and Me2S and can be regarded as 22 Page 22 of 213

having a weak interaction. The second set with much larger shifts include the NH3, H2O, Me2O, Et2O, H2S, nitrile and isonitrile complexes. Although in each case there is a reasonable correlation between the shifted vibrational mode and the proton affinities of the base, it is clear that proton affinity alone cannot be used to predict whether a weak or strong interaction is expected. The Me2S•BF3 complex appears particularly anomalous as although it maintains the correlation between shift and proton affinity for the weak interaction complexes, the

ip t

proton affinity itself would indicate a much larger interaction and hence larger shift in the asymmetric stretching mode. One possibility is that as this involves a S donor ligand, and all

cr

the other data apart from the H2S complex involve second row donor atoms, proton affinity alone may be less important than hard-soft concepts in predicting the extent of interaction

us

where the hardness of Me2O is 772 kJ mol−1 whilst that of Me2S is 579 kJ mol−1 [186]. The data for the symmetric deformation mode only includes the complexes that retain a C3 axis,

an

and in this case there is again reasonable correlation between the shifts and the proton affinity, apart from the data for the H3N•BF3 and MeNC•BF3 complexes. Both of these have substantial shifts of ca. 230 cm−1 in the asymmetric stretching mode indicating substantial

M

interaction, but the reported shifts in the symmetric deformation modes are very small, especially for H3N•BF3 where they are smaller than those of the N2 and CO complexes in Ar

d

and it was blue-shifted in nitrogen [184]. In the solid state [196] the asymmetric stretching mode of H3N•BF3 was at 1148 cm−1 (11B) and the symmetric deformation mode was at 334

Ac ce pt e

cm−1 with an overtone close to the value reported in the matrix data. Assigning features in complicated spectra of complexes is challenging, especially involving mixed modes (the calculations indicate that this mode is probably about 75% BF3 symmetric deformation coupled with 25% symmetric NH3 deformation [184]), but it does seem that there is a possibility this has been misassigned. The position of the MeNC•BF3 complex [195] also appears anomalous as this shift is smaller than that observed for the MeCN complex [151], although the shifts in the asymmetric mode were larger confirming a greater level of interaction. Figure 8 gives a plot of the

11

BF3 asymmetric stretching mode versus the gas

phase experimental boron-donor distance derived from microwave studies. There is a very good linear correlation and although there is a large gap in the centre, it could be used to estimate the boron-donor distances in the matrix data for which no gas phase structural determination is known. Jonas et al. have demonstrated that there is a very good linear relationship between the calculated interaction energy and calculated bond lengths for a range of BH3 and BF3 23 Page 23 of 213

complexes [6]. Ford has also shown a reasonable linear relationship between the calculated interaction energies and calculated vibrational shifts for the asymmetric stretching mode in the H3-nMenN•BF3 complexes [138]. Whilst it would be tempting to extend this to all the experimental data as a way of predicting interaction energies from matrix IR spectra similar to that for the interaction distances in Figure 8, the variation in the reported values of the calculated interaction/binding/bond energies noted above reduces the accuracy. There is the

ip t

added complication that different authors present slightly different parameters, corrected in

cr

different ways.

Therefore, whilst matrix isolation has made a significant contribution, there are some

us

uncertainties in the data, and very matrix little work has been done so far on bases with

1.4.1.

Application of cryosolutions

an

heavier donor atoms, although these complexes are known in the solid state [163].

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In addition to the matrix isolation work, van der Veken has used liquid cryosolutions of Ar, Kr and N2 to study complex formation between a wide range of acids and bases. A particular area of interest has been the complexes of BF3 which are shown in Table 7, and these

d

complement those observed in cryogenic matrices described above and given in Table 4. A

Ac ce pt e

distinct advantage of using cryosolutions is that variable temperatures allow for an experimental determination of the complexation enthalpies via van’t Hoff plots. However, converting these to complexation energies for comparison with those derived from computational studies requires further processing and some assumptions [197] making direct comparisons not always possible. However, it does appear that the theoretical values are often over estimates.

In liquid argon, BF3 was simply solvated, but when nitrogen was introduced there was evidence for the formation of both 1:1 and 2:1 complexes between N2 and BF3, which were also observed in liquid nitrogen [198]. The complexation enthalpy for the 1:1 complex was −4.8(5) kJ mol−1. OC•BF3 and (OC)2•BF3 complexes were observed in liquid argon together with a tentative assignment of a CO•BF3 complex, with complexation enthalpy of −7.6(0.3) and −14.5(1) kJ mol−1 for the 1:1 and 2:1 complexes, respectively [127, 199]. Calculations indicated that OCS and BF3 can form complexes via either the O or S atom of the OCS unit, however the experimental data from liquid Ar and N2 only showed evidence for coordination 24 Page 24 of 213

via the O atom with the νBF mode for

10

B at 1481 cm−1 [200]. For the OCS•BF3 complex

these νBF modes were observed at 1489 and 1438 cm−1 with a complexation enthalpy of −8.0(5) kJ mol−1 [200]. MeF and BF3 form a 1:1 complex, MeF•BF3, which van der Veken described as van der Waals, with the asymmetric νBF modes at 1472 and 1419 cm−1 for 10B and 11B, respectively

ip t

[201]. The complexation enthalpy was −16.8(5) kJ mol−1 in liquid Ar and −18.0(3) kJ mol−1 in liquid Kr. The MeF•BF3 complex has also been studied in the vapour phase by van der

cr

Veken [202]. A detailed study of MeF•BFxCl3−x (x =1, 2) complexes in liquid argon has shown that these weak complexes have the reverse order of Lewis acidity compared to more

us

strongly bound complexes. This was attributed to the relatively poor overlap between the base donor ligand orbitals and the boron pz orbitals, leading to predominantly electrostatic subject of a recent computational report [181].

an

bonding [203]. For BCl3 no adducts were detected. The MeF•BF3 complex has been the

M

In addition, weak complexes between BF3 and a large number of organic species have also been observed in liquid argon, krypton or nitrogen cryosolutions. BF3 and ethylene or

d

propylene form 1:1 complexes with complexation enthalpies for C2H4•BF3 of −10.0(0.2) and −5.4(0.3) kJ mol−1 in liquid argon and nitrogen, respectively [204]. For C3H6•BF3 the values

Ac ce pt e

were −11.8(0.2) and −6.9(0.3) kJ mol−1. 1:1 complexes between BF3 and cyclopropane were observed in liquid argon and liquid nitrogen with a complexation enthalpy of −7.8 kJ mol−1, compared to calculated gas phase complexation energy of −12.0 kJ mol−1 [205]. (BF3 cyclopropane complexes were not detected in the matrix experiments [129].) The 1:1 complexes between BF3 and OCF2 were found to have complexation enthalpies of −11.8(3), −10.6(3) and −7.8(3) in liquid argon, krypton and nitrogen, respectively [206]. The 1:1 complex formed between BF3 and allene was observed in liquid argon and liquid nitrogen, and the complexation enthalpy was determined as −9.2(3) kJ mol−1 in liquid argon [207]. The 1:1 complex formed between BF3 and vinyl fluoride has been studied in both argon cryosolutions and argon matrices with experimental complexation enthalpy of −11.2(3) kJ mol−1 and a calculated complexation energy of −15.2 kJ mol−1 [208]. Although calculations indicated that both σ (via F atom) or π (via C=C) interactions were possible for the BF3 vinyl fluoride complex, the experimental data from argon cryosolutions indicated formation of a σ complex, whereas the argon matrix data also contained weak bands due to a π complex [208]. For BF3 acetylene and propyne complexes the 1:1 complexes predominated, but there was 25 Page 25 of 213

evidence for 1:2 complexes at higher BF3 concentrations [209, 210]. The complexation enthalpies in liquid argon were −9.3(4) and −13.2(4) kJ mol−1 for the C2H2 and C3H4 1;1 complexes, respectively [209]. 1:1 complexes were also formed in liquid argon between BF3 and methylenecyclopropane (complexation enthalpy −10.7(3) kJ mol−1) [211], and cis-2butene (complexation enthalpy −12.5(2) kJ mol−1), trans-2-butene (complexation enthalpy −12.4(2) kJ mol−1), 2-methylpropene (complexation enthalpy −12.7(2) kJ mol−1) [212] and

ip t

cyclopentene (complexation enthalpy −16.1(9) kJ mol−1) [25]. There were two BF3 methylcyclopropane complexes in liquid argon; an asymmetric one where the BF3 interacted

cr

with the C-C bond adjacent to the methyl group (with a complexation enthalpy −9.5(4) kJ mol−1) and a symmetric one in which the interaction occurred with the opposite C-C bond

us

(with a complexation enthalpy −8.3(3) kJ mol−1) [213]. The BF3 cyclopropene 1:1 complex in liquid argon and nitrogen was present as two isomers: a π-type complex with interaction

an

between the BF3 and the double bond of cyclopropene; and a σ complex following interaction with one of the C-C single bonds [214]. The experimental complexation enthalpy for the π-

1.4.2.

M

type complex was −7.4(3) kJ mol−1, with a calculated complexation energy of −12.7(8) [214].

Aluminium halide complexes

d

The N2•AlCl3 complex was discovered by accident when the initial report on molecular AlCl3

Ac ce pt e

claiming it was non-planar [215] was subsequently found to be due to a small air leak in the vacuum chamber which resulted in the formation of N2•AlCl3 [216-218], and that AlCl3 itself was indeed planar [219-221].

Calculations on the H2O•AlF3 and H2O•AlCl3 complexes indicate that the structures are qualitatively similar to those of H2O•BF3 where the H2O molecules leans towards the Lewis acid so that it eclipses two of the Al-F or Al-Cl bonds. The binding energies of −129 kJ mol−1 for H2O•AlF3 and −120.4 kJ mol−1 for H2O•AlCl3 [222] were higher than those predicted for H2O•BF3 (ca. 40 kJ mol−1) [169, 170]. Subsequent work has expanded this to include the (H2O)2•AlX3 and (H2O)3•AlX3 (X = F, Cl) complexes [223]. Boutalib has carried out a series of calculations on H3Y•AlX3 (X = F, Cl, Br) where the staggered conformation was found to be the most stable for all except H3N•AlF3, where the eclipsed conformation was favoured [224]. These calculations also showed that the stability of the complexes decreased from nitrogen to arsenic, and that the AlF3 complexes were more stable than the AlCl3 complexes, which in turn were more stable that the AlBr3 complexes. This variation in stability was not 26 Page 26 of 213

correlated with the extent of charge transfer, but the shortening of the P-H and As-H bond lengths upon complexation was due to an increase in “s” character in these bonds. Subsequent calculations on the complexes between AlCl3 and the hydrogen halides (HX), halogens (X2) and interhalogens (XY) (X, Y = F, Cl, Br) predicted that the staggered conformation represented the stable minima (for all the complexes except those of HCl and HF), and that the eclisped conformers were transitions states. For the HX complexes the

ip t

strength of interaction was found to be HF > HCl > HBr, for the X2 complexes it was Br2 > F2

cr

> Cl2, whereas for XY it was FBr > ClBr > FCl > BrCl > BrF > ClF [225].

AlCl in its 1Σ ground state is not very reactive, so most reactions only proceed after

us

photolysis. AlCl reacts with O2 after photolysis to give the peroxo complex, ClAlO2, and if the oxygen content was high enough the bis superoxo, ClAlO4, complexes were formed [226-

Gallium halide complexes

M

1.4.3.

an

228].

There do not appear to be any reports of gallium halide complexes in the matrix isolation literature, even though Ball has carried out calculations on H2O•GaF3 and H2O•GaCl3

d

complexes in which he encouraged experimental effort [229], much as he did for the

Ac ce pt e

analogous AlF3 and GaCl3 complexes [222]. 1.4.4.

BF3 complexes - computational considerations

Given the significance of the boron halide group 15 complexes there has also been considerable computational effort, especially directed to understanding the order of Lewis acidity of the boron halides. An ab initio investigation of the structures and stabilities of the donor-acceptor complexes H3P•BF3 and Me3P•BF3 was carried out to resolve uncertainties about their stability and structure [230]. These were in good agreement with the experimental data for Me3P•BF3 [231-233] but as there was disagreement with the experimental reports for H3P•BF3 [234-238] it was suggested that further experimental work was required, but this does not appear to have been taken up. More recent calculations have shown that the estimated free molecular enthalpies of dissociation indicate that all the X3P•BY3 (X, Y = Cl, Br, I) complexes are thermodynamically labile species which are stabilised by lattice energies in the solid state. The BI3 complexes have shorter donor-acceptor bond lengths and are more thermodynamically and kinetically stable than the BBr3 and BCl3 complexes as they have larger charge-transfer and smaller destabilising interactions [239]. These calculations were 27 Page 27 of 213

coupled with IR, Raman,

31

P MAS-NMR and X-ray diffraction experiments, which did not

require matrix techniques [163]. Quantum chemical investigations by Frenking [240, 241] of the Y3P•BX3 and Y3P•AlX3 (X = H, F, Cl; Y=F, Cl, Me, CN) complexes identified that the calculated bond dissociation energies, De, for both the borane and alane Lewis acids showed the trend PMe3 > PCl3 ~ PF3 > P(CN)3. The order of the Lewis acid strength of the boranes was dependent on the phosphine Lewis base. For PMe3 and PCl3 the boranes show the Lewis

ip t

acid strength trend BH3 > BCl3 > BF3, but with PF3 and P(CN)3 the order was BH3 > BF3 > BCl3. The bond energies of the alane complexes always showed the trend AlCl3 ≥ AlF3 >

cr

AlH3. It was generally not possible to correlate the trend of the bond energies with one single factor which determines the bond strength. P(CN)3 was identified as a weaker Lewis base

us

than PF3, PCl3 and PMe3 mainly because of its weaker electrostatic attraction, and the complex (NC)3P•BH3 was predicted to have a bond dissociation energy Do of 61.8 kJ mol−1

an

and it was suggested that it could be the ideal candidate to synthesise the first adduct with P(CN)3 as a ligand [240], however, no one seems to have taken up the challenge yet. Comparison of Lewis acid strength of BF3and BCl3 towards strong bases showed that the

M

deformation energy of BCl3 and BF3 was nearly the same, indicating that the stronger bond in H3N•BCl3 compared to H3N•BF3 was due to the enhanced covalency in the former, arising

d

from the lower lying LUMO of BCl3 [242]. Plumley and Evanseck reported that coordinate

Ac ce pt e

covalent bond strengths are not an adequate measure of Lewis acidity, and this should be gauged by the boron atom’s valence deficiency (ability to accept an electron pair) [243]. They also contend that substituent electronegativity can explain boron Lewis acidity for second period donor atoms, but that both electronegativity and atom size are required for third period atoms. In addition, for group 14 donor atoms, substituent electronegativity was sufficient, but for group 16 and 17 an equal balance of electronegativity and atomic size was necessary. The atomic size effects are a result of the variations in σ bond overlap [243]. Another recent computational report has highlighted that understanding the bonding in X•BH3-nFn and X•BH3-nCnn complexes (X = N2, HCN, LiCN, H2CNH, NF3 and n = 0-3), and especially the order of Lewis acidity (BF3 < BCl3 < BBr3) is a challenging task [244]. The high-level ab initio calculations indicated that the usual explanations such as π-donation from the halogen lone pair to the empty p(π) orbital on the B, deformation energies, charge capacities, or LUMO energies cannot independently explain the Lewis acidity. Rather, the binding energies in these complexes result from a combination of at least three factors: a decrease in the electron-accepting ability of B following π-donation by the halogen; the increase in the electron acceptor capacity of B due to deformation and the large increase in 28 Page 28 of 213

deformation energy of the acid with increasing halogen substitution. Of these the dominant effects are those derived from the electronic changes associated with deformation. The complexes formed can either be classified as covalent or van der Waals complexes on the basis of the binding energies and B-N distances. Whilst these are readily available from computational output, extracting them from experimental studies is not always straightforward. It was highlighted earlier that Jonas et al. had indicated that dissociation energies

ip t

(D0) of ca. 100-200 kJ mol−1 are indicative of a strong interaction, whereas for weakly bound, van der Waals, complexes the values of D0 are smaller and around 20 – 40 kJ mol−1 [6]. As

cr

the BCl3 and NH3 reaction is relevant to the production of BN, this has also attracted

us

computational interest [245, 246].

Theoretical dissociation energies of MX3-D complexes (M = Al, Ga, In; X = F, Cl, Br, I; D =

an

YH3, PX3, X−; Y = N, P, As) were found to decrease in the order F > Cl > Br > I; Al > Ga < In; and N >> P ≥ As and the calculated dissociation energies for the ammonia adducts were on average 7 kJ mol−1 higher than the experimental values. No correlation was found between

M

the dissociation energy and the degree of charge transfer. Complexes of ammonia and the metal fluorides have mostly ionic metal-donor bonds, while the other donor-acceptor adducts

d

are mostly covalently bonded. In addition, significant charge redistribution between the

Ac ce pt e

terminal atoms leads to further electrostatic stabilization of the ammonia adducts. Whilst Coulomb interactions destabilize X3P•MX3 complexes, and despite some experimental indications, the existence of these particular complexes in the gas phase was thought to be improbable [247]. This work was subsequently extended to include nitriles and isonitriles [248]. Calculations on the complexes formed between sulfinimines and BF3 indicate that whilst complexation via the O atom results in the lowest energy conformation, the energy difference between this and N complexation is not significant, however, complexation via the S atom can be excluded [249]. Ogawa has carried out a computational study of the Lewis acidity of gallium halides [250]

1.5.

Complexes involving boranes

1.5.1.

Introduction

The recent renaissance in group 13 chemistry has largely involved the hydrides. Whilst matrix isolation has played its part in this, much of this chemistry is outside the scope of this review, but is reviewed elsewhere [113-117]. 29 Page 29 of 213

1.5.2.

Borane and N-donor bases

Ammonia-borane (H3N•BH3) (Figure 1) has found use as a potential chemical hydrogen storage material and reports on calculations at high pressure contain useful reviews [251, 252]. H3N•BH3 is commercially available, and as for the BF3 complexes the B-N distance

ip t

varies from the gas phase (1.6576(16) Å, [253]) to the solid state (1.565(7) Å (X-ray) and 1.58(2) Å (neutron) [254]). The H3N•BH3 complex (as well as D3N•BD3 and D3N•BH3) is an early example where matrix isolation techniques (using liquid hydrogen cryostats) allowed

cr

for a much improved assignment of the IR spectra and confirmation of a C3v point group for

us

the complex [255]. Some of this early data has subsequently been re-assigned using INS data, but this confirmed that the matrix data was different to that of the solid state [256]. Dillen also claims to have ended a thirty year controversy by reassigning some of the original argon

an

matrix bands and the apparent discrepancies in the solid state structural data [257, 258].

M

In the laser ablation of boron atoms in the presence of H2 there was evidence for both a H2•BH and a H2•BH3 complex in addition to BH and B2H6 [112].

d

The final product of high-temperature pyrolysis of gas phase B2H6 and NH3 is BN, which is

Ac ce pt e

an important ceramic material [259-261]. Whilst gas phase FTIR has been used [262], matrix isolation techniques allow for the elucidation of the mechanism. Ault has carried out a series of investigations on the complex formation and pyrolysis of BH3 and B2H6 with NH3 as well as a range of other N, O and S donor atom bases. As these were primarily pyrolysis investigations, merged-jet techniques were employed, but twin-jet tests were also carried out. For mixtures of B2H6 and NH3 in argon, no reaction was observed with twin-jet deposition, or merged-jet (40 cm and 145 cm) deposition at room temperature or up until 80°C with a 145 cm mixing region. However, new product bands were observed from 180°C which increased in intensity up to 250°C [51]. This high temperature pyrolysis of B2H6 and NH3 with products trapped in argon matrices led to the formation of H2B=NH2 [51]. Although, the infrared spectrum of gaseous aminoborane, H2B=NH2, obtained from a flow IR study of the decomposition of H3N•BH3 was known previously [263], the matrix data allowed a more complete assignment. There was no evidence for either BH3 or H3N•BH3 in these experiments, although they were implicated in the reaction mechanism. Pyrolysis of the H3N•BH3 complex itself between 65 and 300 °C, resulted in decreased intensity of the IR 30 Page 30 of 213

bands of the complex and an increase in the formation of H2B=NH2 via H2 elimination [264]. Calculations indicated that the activation barrier for the formation of H2B=NH2 from H3N•BH3 was over 170 kJ mol−1, primarily because the reaction is forbidden by the Woodward-Hoffman rule [265]. Twin-jet deposition of B2H6 and methylamines (Me3N, Me2NH and MeNH2) in argon

ip t

matrices gave no new infrared absorptions, whilst merged-jet experiments gave numerous new bands [266]. The merged-jet pyrolysis of B2H6 with methylamines gave 1:1 complexes

cr

of the form Me3-xHxN•BH3. For the mono- and dimethylamines the H2 elimination reaction yielded the previously known H2B=NMe2, but H2B=NHMe was identified for the first time

us

[266]. For trimethylamine, only the 1:1 Me3N•BH3 adduct was observed, presumably because of the absence of a facile H2 elimination process, thus indicating that elimination of CH4 is

an

more difficult than H2. This work helped clarify the position of the νBN modes in BH3 amine complexes which had previously been the subject of some dispute and uncertainty. More recently, the methylamine-borane molecules have been studied by X-ray diffraction, gas-

M

phase electron diffraction, as well as quantum chemical calculations [267].

d

In a merged-jet copyrolysis study involving N-heterocyclic bases, pyrrolidine was observed

Ac ce pt e

to form a C4H9N•BH3 1:1 complex between ambient and 200°C, with the formation of the H2 elimination product, H2B=NC4H8 occurring above 50 °C. With pyridine only the C5H5N•BH3 complex was formed (which was known previously), and no reaction was found between pyrrole and B2H6 up to pyrolysis temperatures of 380°C. [268] This behaviour was rationalised in terms of the role the nitrogen lone pair plays in the aromaticity of pyrrole. Therefore, whilst the initial 1:1 complex between BH3 and NH3 was not observed, 1:1 complexes were observed between BH3 and all of the methylamines as well as pyrrolidine and pyridine, but not for pyrrole.

Recent calculations on the N2•BH3 complex indicate that it has a short B-N distance close to 1.6 Å (similar to that of H3N•BH3), but that the binding energy is low at −24 kJ mol−1 more reminiscent of a non-bonded complex [269]. Calculations on MeCN•BH3, HCN•BH3, FCH2CN•BH3 and F3CCN BH3 indicate short B-N distances (<1.6 Å), and with appreciable binding energies of around −80 kJ mol−1 [270].

31 Page 31 of 213

1.5.3.

Boranes and O-donor bases

The pyrolysis of B2H6 with MeOH, EtOH or Me2O was investigated by some twin-jet experiments, but mainly with a merged-jet system utilising a 145 cm, 6 mm OD copper tubing with a final 90 cm that could be heated to 400°C [271]. Twin-jet pyrolysis of B2H6 and MeOH gave no new bands, neither were any observed for merged-jet deposition in the absence of heating. Heating to 150°C yielded some new bands, and after heating to 300°C

ip t

more bands were observed, and the parent bands of MeOH had disappeared, and those of B2H6 were very weak. Heating beyond 340°C resulted in reduction of product bands of

cr

methoxyborane, H2B=OMe. This was the first observation of its IR spectrum, although its microwave spectrum was known [272]. Lesser yields of H2B=OMe were observed for Me2O,

us

and ethoxyborane was observed in the pyrolysis of B2H6 and ethanol. No evidence for complex formation was observed for all three of the bases, although a Me(H)O•BH3 complex

an

was postulated as a reaction intermediate, which lost H2 to yield H2B=OMe. This can then eliminate CH4 (via a HBO intermediate) to form H3B3O3 (boroxin), or react with a second

M

MeOH to form Me2O [271].

In merged-jet copyrolysis of B2H6 with either H2O or O2 there was no observable reaction at

d

room temperature or up to 150°C, but new product bands due to H3B3O3 were observed at

Ac ce pt e

200°C and continued to increase up to 350°C [273] rather than H2BOH as found in a microwave study [274]. At the highest temperature bands due to HBO were also observed in good agreement with earlier data [275]. However, no bands assignable to H2BOH were observed, despite this species being identified in microwave pyrolysis studies [274]. The cause of this apparent discrepancy may be due to the higher sensitivity of microwave spectroscopy to H2BOH compared to IR, and that H3B3O3 has no permanent dipole moment, rendering it microwave silent.

1.5.4.

Boranes and S-donor bases

The merged-jet copyrolysis reaction of B2H6 and H2S led to the formation of H2BSH, but only after heating to 260°C. In addition HBS and a small amount of boroxin (due to reaction with residual O2) were observed. The position of the B-S stretching mode in H2BSH suggests single bond character, whereas in analogous oxygen and nitrogen containing compounds, there is considerable double bond character [276]. Merged-jet copyrolysis of B2H6 with MeSH or EtSH at 260°C led to H2BSMe (methylmercaptoborane) and H2BSC2H5, 32 Page 32 of 213

respectively, with no evidence for reactions at lower temperatures. Likewise there was no evidence for Me(H)S•BH3 or Et(H)S•BH3 complexes as the H2 elimination was too rapid [277]. The observation of HBS and CH4 at higher temperatures (analogous to that observed for MeOH [271]) indicates decomposition of H2BSMe. No reaction was observed with Me2S, even up to 400°C [277]. A computational report published not long after Ault’s experimental investigation indicated that the formation of H2BSMe could proceed via either a H2S•B2H6 or

Boranes and P-donor bases

cr

1.5.5.

ip t

H2S•BH3 initial intermediate [278].

us

In contrast to Ault’s previous work with nitrogen, oxygen and sulfur bases, the merged-jet copyrolyisis of B2H6 with PMe3, PH3, or PF3 only led to the formation of Me3P•BH3 with no elimination products, and no H3P•BH3 or F3P•BH3 complexes [279]. Me3P•BH3 was first

an

observed in 1940 [280] and was prepared separately, but on attempted sublimation only the parent decomposition products were observed [279]. Previous results indicated that the entire

M

system had to be maintained below -77°C for it to be observed [281]. F3P•BH3 has also been prepared previously [282]. The matrix results were taken to imply that the equilibrium constant for complex formation between the complex and its precursors is very small at room

d

temperature, and that this behaviour was consistent with the Lewis basicity of the phosphorus

Ac ce pt e

donors.

1.5.6.

Boranes and C-C multiple bonds

No reaction between B2H6 and C2H4 was observed under merged-jet conditions until the reaction zone reached 300°C [283]. No bands due to alkylboranes were observed, despite being known to form under equilibrium conditions [284]. Instead, reaction of B2H6 with C2H4 yielded C2H6, and C2H2 gave C2H4. Bands due to boroxin were also observed due to reaction with trace impurity levels of O2 or H2O but under richer ethylene conditions their intensity could be reduced greatly.

1.6.

Aluminium, gallium and indium hydride complexes

Whilst the application of borane chemistry is mostly directed to ceramics, that of the heavier Group 13 hydrides is more focussed on semiconductor applications [285]. Although less extensive than boron chemistry, there is a considerable body of knowledge on the complexes and compounds of the heavier Group 13 hydrides which has been reviewed [114-116, 286, 33 Page 33 of 213

287], especially in a recent book [117]. It is generally found that the complexes or adducts formed between the heavier Group 13 hydrides and nitrogen, phosphorus and oxygen bases have greater thermal stability than the parent hydrides [288]. Many of the complexes have been well characterised and found use in a variety of synthetic and semiconductor fabrication environments, rather than relying on matrix isolation for their stabilisation [289]. It has been observed that a phosphine binds much more effectively to GaH3 than to AlH3, whereas the

ip t

reverse order is found with an amine. The Al P interaction is sufficiently weak that adducts containing sterically non-hindered phosphines decompose to polymeric [AlH3]n under

cr

relatively mild conditions [290]. Chaillet et al. have carried out a detailed sets of calculations on the H3E•AH3 (E = N, P, As, A = B, Al, Ga, In) complexes [291]. In general they found

us

that the bond energies decreased with the size of the donor and acceptor atoms, but the changes from row to row of the periodic table were irregular. In particular, NH3 formed a

an

slightly stronger bond to AlH3 than BH3 in contrast to the other donors. Whilst the parent hydrides can be prepared in matrices (vide infra), the lack of stable precursors has limited the

M

matrix chemistry reported so far.

Although solid state “AlH3” with an AlF3 structure has been known for a long time, the

d

molecular, vapour phase, Al, Ga and In hydrides have proved much more elusive. AlH3 was

Ac ce pt e

isolated for the first time when Al atoms were photolysed in presence of H2 in Ar or Kr matrices [292]. Al2H6 was formed after the photolysis of laser ablated Al atoms in pure H2 matrices in 2003 [108, 109]. Pulsed laser ablation of Al atoms in the presence of H2/Ar gave AlH, AlH2, AlH3 and variety of Al2H4 species, as well H2 complexes with the major products when condensed in argon matrices [293]. Pulsed laser ablation results in hyper thermal aluminium atoms which have enough energy to overcome the activation barriers which was in contrast to the very inefficient reaction between thermally evaporated of Al atoms and H2 in a Kr matrix, but photolysis did yield AlH and AlH2 [294]. For the H2 complexes the νAlH stretching modes were blue shifted compared to parent hydride. However, it was noted that for the H2•AlH complex these spectral features may arise from no more than H2 perturbations of AlH in the matrix cage, without a well defined structure. For H2•AlH2, the spectral features of the complex were well defined, and alluded to for H2•AlH3 [293]. AlH2 can also be prepared from molecular hydrogen and 2S or 2D excited aluminium atoms [295], whilst AlH3 can be prepared from Al atoms and H atoms formed in a microwave discharge [296].

34 Page 34 of 213

The gas phase structure of Me3N•AlH3 has been determined by both gas phase electron diffraction [297] and pulsed beam FT microwave spectroscopy [298] and yield a Al-N bond length of 2.063 Å. Whilst Ball has carried out some calculations on the H2O•AlH3 complex [222], Boutalib has carried out an extensive set of calculations on alane complexes including: H3E•AlH3 (E = N,

ip t

P, As) [299]; Me3E•AlH3 (E = N, P, As) [Me3D•AlH3]− (D = C, Si, Ge) [300]; [Me2E•AlH3]− and Me2Y•AlH3 (E = N, P, As, Y = O, S, Se) [301]; [H3E•AlH4]+ (E = N, P, As) and

cr

[H2Y•AlH4]+ (Y = O, S, Se) [302]; H3-nXnN•AlH3 (n = 0-3, X = F, Cl) [303]; X3E•AlH3-nMen (X = H, F, Cl, E = N, P, n = 0-3) [304]; and X3As•AlH3-nMen (X = H, F, C, Br, n = 0-3)[305].

us

A recent computational report indicates that the phosphine alane complexes are much stronger Lewis acids than the parent phosphines, and that the enhancement is greater for the

an

alane complexes than borane analogues [306]. The trimers and tetramers of both MH and MH3 (M = Al, Ga) have been investigated by both DFT and coupled cluster calculations

M

[307].

Gallane (Ga2H6) was first identified in 1989 [107] and has been recently updated with a more

d

accurate gas phase structure [308], together with a laser ablation study yielding GaH, GaH2,

Ac ce pt e

GaH2− and Ga2H2 on deposition, with Ga2H4, Ga2H5 and Ga2H6 formed after photolysis [309]. Indium [110, 310] and thallium [111] hydrides analogous to the Ga system were formed from laser ablation of the metal with trapping in pure H2 and D2 as well as H2/Ne matrices. AlH3, GaH3 and InH3 can also be prepared in matrices using hydrogen atoms (from a microwave discharge) and metal atoms [296]. The higher reactivity of gallium and indium dimers towards H2 has been the subject of detailed experimental and computational investigations [311-313].

The Me3N•GaH3 complex was matrix isolated from the vapour above the solid complex at 228 K, and the resultant spectrum more closely resembled that of the vapour phase complex than the solid state [85]. Ball has also carried out calculations on the H2O•GaH3 complex in addition to those on H2O•GaF3 and H2O•GaCl3 [229]. The binding energy of the H3As•GaH3 complex has been calculated as −67 kJ mol−1, with the comment that it should be observable in the gas phase 35 Page 35 of 213

[314]. However, subsequent reports reduced the stability to −30.5 kJ mol−1 [315]. A number of other computational reports have compared the complexes of AlH3 and NHMe2, PHMe2, AsHMe2 Me2O, Me2S or Me2Se [301]; and BH3, AlH3 and GaH3 with NH3, PH3, AsH3, H2O, H2S or H2Se [316].

Trimethyl aluminium, gallium and indium complexes

ip t

1.7.

As bases with “active” hydrogen may result in H2 or CH4 elimination, matrix isolation can be

cr

instrumental in the detection of initial complexes. Coates has reported that the Lewis acid strength of the Group 13 trimethyl compounds towards Me3N is B < Al > Ga > In > Tl [317].

us

Frenking has presented a computational review, which includes the Group 13 trimethyl

1.7.1.

an

complexes [241].

Me3Al complexes

A OC•AlMe3 complex in solid argon has a νCO mode at 2185 cm−1, 47 cm−1 higher than

M

“free” CO, which shifted to 2134 cm−1 when 13CO was used [318]. Attempts to record a room temperature IR spectrum were unsuccessful. In the gas phase, OC•BH3 has νCO at 2165 cm−1

Ac ce pt e

d

[319]. These νCO modes higher than “free” CO indicate a limited π bonding interaction. Merged-jet copyrolysis (50 cm) with the deposition line at room temperature of Me3Al with NH3 with the products trapped in solid argon resulted in the formation of H3N•AlMe3 confirmed by isotopic enrichment experiments [320]. No reaction was observed using twinjet methodology. There was some discrepancy between the solid state IR data obtained at 77 K [321] and the matrix data, with the most likely explanation being that there is some change in structure and geometry between the matrix and the solid. In the case of the matrix data it was argued that this was best explained in terms of a strong interaction between the Me3Al and NH3 sub-units, with a significant distortion away from planarity of the former, and towards planarity for the latter. In the case of Me3N the complex with Me3Al agreed with the literature data [322]. Müller reassigned some of Ault’s bands for H3N•AlMe3 on the basis of RMP2(fc)/6-31G* calculations, and observed the formation of aminodimethylalane, Me2AlNH2, after photolysis of the H3N•AlMe3 complex in argon matrices at 210 nm [323].

36 Page 36 of 213

The gas phase decomposition of H3N•MMe3 (M = Al, Ga, In) has been monitored by gas phase IR spectroscopy between 298 and 573 K using reactant concentrations compatible with those for Al(Ga, In)N MOCVD [324]. At 473-523 K H3N•AlMe3 decomposed quantitatively to Me2AlNH2 (the majority being dimeric) and CH4. For Me3Ga/NH3 and Me3In/NH3 mixtures at temperatures below 543 K, the chemistry is dominated by reversible adduct formation-dissociation, with no detectable CH4. At 574 K limited decomposition was

ip t

observed in the Me3Ga/NH3 mixtures.

cr

Merged-jet codeposition of Me3Al with PH3, Me3P, Me2O, Me2CO or Me2S in argon matrices [325] resulted in the formation of the previously known Me3P•AlMe3, Me2O•AlMe3 and

us

Me2S•AlMe3 complexes, some reassignment of the IR data for the latter two which was originally collected as neat liquids at room temperature [326-329] and the first IR data for

an

Me3P•AlMe3 [325]. In addition, H3P•AlMe3 and Me2CO•AlMe3 were reported for the first time. The yields varied dramatically, with high complex yields for Me2O and Me2CO, and the Me3Al dimerisation reaction.

M

very low yields with Me2S, Me3P and PH3, which was believed to be due to competition with

d

Reaction of Me3Al with methanol in solution is well known to yield dimeric, trimeric and

Ac ce pt e

polymeric alkylaluminium alkoxides. Matrix isolation can give valuable information on the initial species formed. In twin-jet mode, no product bands were observed with Me3Al and MeOH, but in merged-jet (with 40 cm reaction zone) substantial reaction was observed [330]. These reactions go to completion, but some intermediates were also observed. The least stable product was identified as the CH4 elimination product, Me2AlOMe, rather than the 1:1 adduct, Me3(H)O•AlMe3. The identity of the second product, stable at up to 350°C pyrolysis conditions, was thought to be either the dimer [Me2AlOMe]2 or the reaction of Me2AlOMe with a second molecule of MeOH to form MeAl(OMe)2, but it was not possible to determine which product was formed unambiguously. Interestingly, and perhaps surprisingly, no new product bands were observed for the reaction of Me3Al and H2O, using either twin or merged-jet approaches, but there was clear reduction in the intensity of the parent bands and evidence of CH4. One possibility was the condensation of any reaction products in the deposition line, prior to matrix formation [330]. The reaction of Me3Al and O2 required a flow gas reactor with constricted nozzle and the species was identified as monomeric Me2AlOMe, based on oxygen isotopic labelling [331], 37 Page 37 of 213

and also from comparison of features in the reaction of Me3Al and MeOH [330]. Higher temperatures in the reactor led to one or more secondary species, most likely [Me2AlOMe]. The reaction mechanism was postulated to involve the reaction of O2 across the four membered Al-C-Al-C ring in the dimer [Me3Al]2. Although the dimer fraction was usually small, the need for a constricted nozzle and resultant higher back pressure supported this argument. In none of the experiments was there any evidence for either the O2 coordination

ip t

complex, Me3AlOO, or the peroxide insertion product, Me2AlOOMe, which would have very

cr

distinctive 16O/18O isotope patterns [331].

The reaction of Me3Al with H2S led to the formation of Al2S3, which represented its first

us

identification as a vapour phase species [332]. Formation of this involves at least two Me3Al units and three H2S units, with the reaction taking place during the flight time through the

an

pyrolysis zone (milliseconds) but no intermediates were observed in the argon matrix spectrum. When methyl and ethyl mercaptan were used Me3AlSR (R = Me, Et) were the products [332]. Although there was some evidence for the formation of 1:1 complexes of

M

Me3Al with these S donors, the low intensity of the peaks precluded definitive assignment.

d

The shift in the Me3Al asymmetric stretching mode has been shown to be a sensitive probe

Ac ce pt e

for the extent of interaction in Me3Al complexes, and these are given for the above Me3Al data in Table 8 and in Figure 9(a). Whilst this indicates a much stronger interaction for Ndonor bases than P-donor bases and O-donor bases, there is barely enough data to be able to identify whether there is any correlation between these shifts and proton affinity of the base.

1.7.2.

Me3Ga complexes

Trimethylgallium has been an important precursor in the formation of gallium containing semiconductors via metal organic chemical vapour deposition (MOCVD) routes. The GaC3 anti-symmetric stretch has been used as a qualitative indicator of the strength of the acid-base interaction, with shifts of 5 – 35 cm−1 being observed for the complexes compared to the parent Me3Ga. Where the C3 axis is lost, this degeneracy is lifted and a doublet is observed. These matrix data together with those for Me3Al and Me3In are given in Table 8.

38 Page 38 of 213

The H3N•GaMe3 complex is stable in the gas phase and its FTIR spectrum has been observed [333] as has its decomposition products [324]. Me3N•GaMe3 has been characterised by gas phase electron diffraction [288, 334] and single crystal X-ray diffraction [288]. Twin-jet deposition of Me3Ga and AsH3 led to a number of new infrared bands compared to parent ones, and with the use of merged-jet these increased in intensity [335, 336]. As the

ip t

bands were in close proximity to the parent bands, this implied a 1:1 complex or adduct of the form H3As•GaMe3. The GaC3 symmetric and asymmetric modes were at 520 and 565 cm−1 in

cr

the complex compared to 526 and 575 cm−1 in the parent Me3Ga. The observation of the GaC3 symmetric stretch indicated a loss of planarity of the GaC3 unit, whilst the lack of

us

splitting of the asymmetric GaC3 stretching mode implied the presence of a C3 axis in the complex. When experiments were conducted in the absence of an argon or nitrogen matrix,

an

the spectra of the thin film was in broad agreement with that from a matrix, but the conclusions were more tentative due to the significant bandwidths. No bands were observed when an ambient temperature gas cell was used. The calculated binding energy of the

M

H3As•GaMe3 complex has been reported as −17.1 kJ mol−1 [315] or −25.5 kJ mol−1 [337] and

d

these are 13.4 and 21.7 kJ mol−1 less favourable than those for the H3As•GaH3 complex.

Ac ce pt e

There has also been a report on the matrix isolation of Me3N•GaH3, Me3N•GaMe3, and Me2HN•GaMe3 complexes from the vapour above the cooled solids [85]. They also showed that thermal decomposition of Me3N•GaH3 at 60°C yielded NMe3 and a Ga mirror, implying that the breaking of the Ga-N bond is the primary step in the reaction. In contrast, the thermal and photochemical decomposition of the gaseous adducts, Me3N•GaMe3, and Me2HN•GaMe3 results in the formation of methane, implying that the Ga-N bond remains intact in the primary decomposition step. Subsequent, more detailed experiments showed that Me2HN•GaMe3 could be prepared as a white solid at room temperature with a vapour pressure of about 1 Torr [338]. Congruent sublimation meant that it could be structurally characterised by gas phase electron diffraction. Heating to ca. 170°C under argon or nitrogen led to CH4 and [Me2Ga(µ-NMe2)]2, the structure of which was also determined by gas phase electron diffraction [338]. More recent room temperature, higher partial pressure, FTIR experiments indicated the existence of a complex where a second NH3 molecule was hydrogen bonded to the conventional H3N•MMe3 complexes (M = Al, Ga) which can be formulated as H3N···H3N•MMe3 rather than as a genuine 2:1 complex involving NH3 and Me3M [339]. The thermolysis of the single source precursor, Me2N(CH2)3Ga(N3)2 above 39 Page 39 of 213

450°C results in the formation of equal amounts of the azides HN3 and Ga(N3) as well as NH3, H2C=CNMe and HCN according to matrix IR spectra. These reaction products ruled out a β-hydrogen elimination as the initial fragmentation step [340, 341]. A merged-jet study of Me3Ga with PH3, Me3P, SbH3, PF3 or NF3 resulted in the formation of the 1:1 complexes in matrices (Ar and N2) as well as in thin films in each case [342]. The

ip t

asymmetric GaC3 stretching mode was red shifted from the band in the parent mode at 575 cm−1. In the case of NF3 this resulted in a shoulder at 570 cm−1, but for Me3P an intense band

cr

at 549 cm−1 was observed. In addition a band near to 520 cm−1 for each of the complexes was assigned to the symmetric stretching mode, which became activated upon complexation,

us

implying a loss of planarity of the GaC3 unit. Me(H)O•GaMe3, Me(H)S and H2S•GaMe3 were found to be less strongly bound than the Me2O•GaMe3 complex [343]. The

an

Me2O•GaMe3, Me2S•GaMe3 and Me2Se•GaMe3 complexes have nearly identical shifts in the asymmetric GaC3 stretching mode [344]. There was possible weak evidence for the formation of higher stoichiometry complexes. A mass spectrometric investigation of the reaction

M

between Me3Ga and H2Se in a conventional atmospheric pressure MOCVD reactor indicated that the reaction could be described simply without the need for a stable Lewis acid-base

d

complex, but such a complex was not ruled out [345]. Supporting calculations gave a binding

Ac ce pt e

energy of ca. + 69.4 kJ mol−1 (compared to the precursors) for the H2Se•GaMe3 complex indicating its instability. However, the H2Se•GaH3 complex was calculated to be 27.2 kJ mol−1 more stable than its precursors [346]. Gas phase IR spectroscopy has been used to determine an equilibrium constant (Kp) of 2.84 x 102 at 100° C for the association reaction between Me3Ga and NH3, which enabled the determination of ΔH = −68.1 kJ mol−1 and ΔS = −135.4 J K−1 mol−1. [324, 347] The values of the asymmetric GaC3 stretching mode from the matrix data are given in Table 8 and the correlation that Ault noted between these and the proton affinity of the bases is clearly demonstrated in Figure 9(b), although the data point for NMe3 is a little way off. For PH3, AsH3 and SbH3 this was found to be in good agreement with the HX hydrogen bonded complexes of [348]. The largest shifts are for NMe3 and PMe3 and the smallest for PF3 and NF3, with NH3 and PH3 intermediate between these. Whilst the proton affinities of Me2O and Me2S were known, it allowed a prediction of the proton affinity of 815 kJ mol−1 to be made

40 Page 40 of 213

for Me2Se [344], but given the similarity of the GaC3 asymmetric modes this is no great surprise. 1.7.3.

Trimethylindium complexes

In the H3P•InMe3, H3As•InMe3 and H3Sb•InMe3 complexes in argon matrices [349] the shift of the vibrational modes in the complexes compared to the parent sub-units was very similar

ip t

and much smaller than those observed for the analogous Me3Ga complexes [335, 336, 342]. Subsequently the H3N•InMe3 complex was studied (this was first reported by Coates in 1956, but without any spectroscopic characterisation [317]) and the InC3 asymmetric stretching

cr

mode was also very close to all the others [350]. The data for the InC3asymmetric mode are

us

collected together in Table 8. Although the shifts are much smaller, there is still a correlation between them and the proton affinity of the base as shown in Figure 9(c). However, as there is only one representative for each of the group 15 donor atoms, and the shifts are very small,

an

this should not be over interpreted. The relatively small and similar shifts, coupled with the very weak activation of the InC3 symmetric stretching mode indicate that the indium Group

M

15 hydride complexes are quite weakly bound. These data confirm that the order of interaction for Me3Ga is N > P > As > Sb, but that for Me3In as all the shifts are very similar this indicates a similar perturbation of the Me3In sub-unit by all of the Group 15 hydrides. In

d

addition the much smaller shifts for In indicate that the complex formation was stronger for

Ac ce pt e

the case of Ga as compared to In [350].

More recently the reaction of Me3In and O3 (as part of an investigation on the use of ozone to prepare thin metal oxide films) using twin-jet deposition with short reaction times resulted in the observation of small amounts of H3COInMe2, together with a low yield of CH2O. The yield of H3COInMe2 was increased following UV irradiation. With merged-jet deposition, high yields of H2CO, MeOH, and C2H6 were obtained. Attempts at using the concentric-jet method to probe for additional intermediates were not successful in this case [351]. As for the gallium complexes, gas phase IR spectroscopy has been used to determine the equilibrium constant (Kp) of 1.41 x 102 at 100° C for the association reaction between Me3In and NH3, which gave values of ΔH = −62.7 kJ mol−1 and ΔS = −126.6 J K−1 mol−1. [324, 347] There is a detailed computational report on a wide range of indium compounds and their complexes of interest as organometallic precursors for MOCVD [352]. 41 Page 41 of 213

1.7.4.

Trimethylthallium complexes

Although the matrix IR spectrum of Me3Tl and Et3Tl have been reported [353], there do not appear to be any other reports involving complexes although in the case of NH3 and Me3N they have been known for a considerable time and that trimethylthallium is a weaker acid

ip t

than trimethylindium [317].

Group 14

Introduction

cr

1.8.

SiF4 complexes were reviewed in 1989 [354] and 2007 [355] and silicon and germanium

us

halide complexes were reviewed in 2011 [356] all of which included some matrix isolation

1.9.

an

data. CO2 Lewis acid-base complexes were reviewed in 1997 [187].

Silicon and germanium tetrahalides

M

Computational work indicated that the Lewis acidity of the group 14 tetrahalides in the gas phase reduces Sn > Ge > Si (based on the dissociation enthalpies of the complexes), and this was in line with the increasing reorganization energy of the acceptor, but that the donor-

d

acceptor bond energies were Ge < Si < Sn [357]. The authors state that the values of the gas

Ac ce pt e

phase dissociation enthalpy should not be used to estimate the metal-nitrogen bond energy. On the basis of these theoretical studies, the most promising candidates for stable gaseous complexes of group 14 halides were predicted to be complexes of tin tetrahalides with bidentate donor ligands, such as SnF4(phen) [357] which are known in the solid state. Matrix isolation has made a considerable contribution to the knowledge of the SiX4 and GeX4 complexes (X = F, Cl), and this is collected together in Table 9.

1.9.1.

SiF4 and GeF4 complexes

The structure of the SiF4 dimer changes from C2h to D2d on going from the gas phase to Ar, Xe or N2 matrices. The Si-F stretching mode was split into a doublet with one component blue shifted by 6.4 cm−1 and one red shifted by 17.8 cm−1 [358].

42 Page 42 of 213

1.9.1.1.

SiF4 and GeF4 complexes with N donors

SiF4 adducts/complexes with a range of bases have been known for many years, indeed (H3N)2•SiF4 is the oldest known compound with a Si-N bond and the gas phase equilibrium between SiF4, NH3 and (H3N)2•SiF4 has been well studied [356]. This is of industrial relevance as the reactions between SiF4 and NH3 give (eventually) α-Si3N4 [359, 360]. The Si-F stretching modes in the complexes can be used to identify the stoichiometry as they are

ip t

higher in the penta-coordinate geometries than in an octahedral environment.

cr

Prior to Ault’s matrix study in 1981 [361], there was no experimental evidence for a 1:1 complex between SiF4 and NH3, although the 2:1 complex had been known for 200 years.

us

When SiF4/Ar and NH3/Ar mixtures were deposited in twin-jet mode, two sets of IR bands were observed, one with sharp intense features at 409, 456, 706, 854 and 1253 cm−1 (for the

an

natural isotopes), the other set had broader weaker bands at 416, 463, 839 and 1260 cm−1. When higher dilution was employed, the weaker, broader features diminished in intensity. In the absence of a matrix a spectrum was observed with dominant bands at 730 and 1390 cm−1,

M

together with weaker features between 400 and 500, a broad band at 930 cm−1 and a shoulder at 850 cm−1. A Nujol mull spectrum of a pure sample of (H3N)2•SiF4 displayed characteristic

d

features at 730 and 1390 cm−1, but did not contain the other spectral features observed in the non-matrix deposition experiment. On the basis of the concentration dependence as well as 14

N/15N isotopic experiments, Ault assigned the sharp intense bands in the matrix

Ac ce pt e

H/D and

spectrum as arising from an isolated, distinct reaction product, which was assigned to a H3N•SiF4 1:1 complex. The weaker, broader features in the matrix sample spectrum were assigned to aggregates or perturbed species. The position of the Si-F stretching modes indicated a trigonal bipyramidal geometry, and on the basis of a simple normal coordinate analysis, the C3v structure with an axial NH3, rather than a C2v structure with an equatorial NH3 was preferred, but not definitively. Although this maybe unexpected as it breaks the “Muetterties rule” that the most electronegative ligand should occupy the axial site [362], subsequent experimental [363] and computational work [364-366] has confirmed Ault’s initial conclusions that the C3v geometry with an axial NH3 was indeed the minimum on the potential energy surface of H3N•SiF4 and a representation of this structure is given in Figure 10. Subsequent rotational spectroscopy confirmed Ault’s original structure and gave a Si-N distance of 2.090 Å and a measured dipole moment of 5.61 D, which is an enormous enhancement on the sum of the monomer moments of 1.47 D [367].

43 Page 43 of 213

Subsequent experiments have shown that H3N•SiF4 is stable up to −50°C, and attempts to isolate it as a pure compound resulted in the formation of the well known and more stable (H3N)2•SiF4 complex via elimination of SiF4 [368]. SiF4 also formed 1:1 complexes with MeNH2, Me2NH and Me3N (this was already known) in cryogenic matrices with the amine in the axial position [363].

ip t

For SiF4 and PH3 (weaker base than NH3) there was no compelling evidence for the formation of a distinct adduct in these experiments [361]. Although Beattie presented

us

appears to be no further reports of SiF4 phosphine complexes.

cr

tensiometric titration evidence for a complex between SiF4 and PMe3 in 1969 [369], there

Therefore, SiF4 forms complexes with amines, but there are no characterised phosphine

an

complexes, whilst SiCl4 reacts with NH3 to form SiCl3NH2, but not an intermediate complex. Although there was the expected reduction in the νSiF modes on going from NH3 to NMe3 (Table 9), it is not linear, with a larger drop for the mixed intermediate species than NMe3.

M

This behaviour has also been observed in solution and is believed to be due to solvation or

d

steric effects [361, 363].

Ac ce pt e

A detailed series of computational papers has been published on the adducts formed between silicon halides and nitrogen containing donors, including their enthalpies of sublimation [370374]. These investigations confirmed that the H3N•SiX4 complexes have C3v geometry. Whilst the Si-N distance increases from 2.09 Å for H3N•SiF4 to 2.19 Å for H3N•SiCl4 and H3N•SiBr4 it still remains smaller than the sum of the van der Waals radii [370]. The bond dissociation energies (ΔH°298) were calculated as: H3N•SiF4, 42.7 kJ mol−1; H3N•SiCl4, 0.8 kJ mol−1; H3N•SiBr4, −4.7 kJ mol−1 [370].

The formation of 1:1 complexes, H3N•GeF4, MeH2N•GeF4, Me3N•GeF4, rather than 2:1 complexes or hydrogen halide elimination products was observed in experiments involving the codeposition of GeF4 and NH3, NH2Me and NMe3 [375]. As in the case of the SiF4 amine complexes [361, 363], the IR data indicated that the geometry was trigonal bipyramidal [375], with the amine base in an axial position, again breaking the “Muetterties rule” [362]. Twin-jet experiments using pyridine and SiF4 or GeF4 yielded intense features characteristic of the 1:1 complexes, C5H5N•SiF4 and C5H5N•GeF4, with the most likely structure being 44 Page 44 of 213

trigonal bipyramidal with the pyridine in the axial position [376]. There was no evidence for the formation of 2:1 complexes in the matrix. The shift of the vibrational modes of the SiF4 and GeF4 sub-units indicate a substantial interaction, and as the shift was larger for GeF4 than SiF4 this was interpreted as confirming the greater Lewis acidity of GeF4 compared to SiF4. Calculations for the pyridine complexes, indicated that for C5H5N•SiCl4 and C5H5N•SiBr4 the axial isomer showed no evidence of complex formation [373] and that in contrast to the NH3

ip t

complexes the equatorial configuration was preferred. For C5H5N•SiF4 both axial and equatorial geometries demonstrated the formation of complexes, but that the equatorial

cr

isomer was preferred. For the equatorial C5H5N•SiF4, C5H5N•SiCl4, and C5H5N•SiBr4 complexes the bond dissociation enthalpies (ΔH°298) were 11.7, −46.3 and −51.1 kJ mol−1

us

with bond energies of 225 ±10, 169 ± 9 and 149 ±9 kJ mol−1, respectively [373].

an

Twin-jet codeposition of HCN or MeCN with SiF4 or GeF4 resulted in the formation of 1:1 complexes HCN•SiF4, HCN•GeF4, MeCN•SiF4 and MeCN•GeF4 [377]. Although HCN could act as either a base or acid (via hydrogen bonding) the IR spectra were all interpreted in

M

terms the HCN being coordinated through the N as a base, with essentially linear Si-N-C and Ge-N-C linkages involving “end-on” bonding of the base. The overall structure was not

d

discussed in any detail, but the similarity of the spectra to those obtained from the amines

Ac ce pt e

implied a trigonal bipyramidal structure with the HCN base in the axial position. The shifts of the IR frequencies (Table 9) in the complexes from those in the parent species indicated a relatively weak interaction, especially for SiF4. The relative proportion of complex formation was higher for GeF4 than SiF4. With the stronger bases (e.g. amines) there was a much less pronounced difference in yield, and this was taken to indicate that the difference in acidity of SiF4 and GeF4 was more noticeable with the weaker bases [377]. The first GeF4 phosphine complexes were only reported in 2008 [378] and whilst GeCl4 reacts with PMe3 at low temperature in the absence of solvent to form [GeCl4(PMe3)2], this rapidly decomposes in solution [378-380].

1.9.1.2.

Complexes of SiF4 and GeF4 with O donors

Margrave studied the gas phase reactions of SiF2 and SiF4 with MeOH and H2O [381, 382]. In all cases there was no adduct or complex formation reported, and only the hydrogen

45 Page 45 of 213

fluoride elimination products, MeOSiF [381] and F2SiOSiF2 and F3SiOSiF3, were observed [382]. Although Ault had failed to observe the Me2O•SiF4 complex in previous work [361], on reinvestigation of the SiF4 reactions with Me2O, H2O and MeOH using the twin-jet matrix isolation method he observed the formation of the 1:1 complexes, H2O•SiF4, Me(H)O•SiF4

ip t

and Me2O•SiF4 with no evidence for 2:1 complexes [68]. Whilst it was not possible to identify the geometry unambiguously, it was thought that the distortion of the tetrahedral SiF4

cr

unit on complex formation was towards trigonal bipyramidal, but only reached the limit in the amine complexes discussed above [361, 363]. A linear relationship between proton

us

affinity of the base and the position of the Si-F stretching mode (increasing base strength, decrease in frequency) in the complex was interpreted to mean that SiF4 was acting as a π

an

acceptor. A much more recent matrix isolation study of the hydrolysis of SiF4 at higher spectroscopic resolution [383] confirmed the earlier work of Ault [68] and the assumptions of the formation of a molecular complex, followed by trifluorosilanol and eventually

d

387].

M

hexafluorodisiloxane based on gas phase IR data [384, 385] as well as calculations [386,

Ac ce pt e

A series of matrix experiments were carried out with GeF4and O donor bases, and the following 1:1 complexes were observed: H2O•GeF4, D2CO•GeF4; Me(H)O•GeF4, Me2O•GeF4 and Me2CO•GeF4 [388]. The IR data were again consistent with a trigonal bipyramidal structure with the base ligand in an axial position, but a definitive assignment could not be made. In the case of Me2O, the symmetric O-C-O stretching mode only shifted 7 cm−1 in the SiF4 complex, but 33 cm−1 in the GeF4 complex, indicating the different level of interaction. In the quite strongly bound Me2O•BBr3 complex the shift was 40 cm−1 [166]. For the methanol complexes, no perturbed vibrational modes of the base were observed in the case of SiF4, but for the GeF4 analogue both perturbed C-O and O-H stretching modes were observed. In the case of the acetone complex this displayed a shift of over 40 cm−1 in the νCO mode of acetone with GeF4, but in the corresponding SiF4 experiment, there was no evidence for a perturbed νCO mode. Taken together, this evidence is consistent with GeF4 forming substantially stronger complexes than SiF4. The yield of the Me2O•GeF4 molecular complex was very similar in both single-jet and twin-jet modes of deposition, suggesting that the 1:1 gas phase complex was not involved (to any appreciable degree) in the room temperature equilibrium, but that the complex was only formed on the matrix surface, once sufficient 46 Page 46 of 213

thermal energy had been removed from the reactants so that the 1:1 complex could be stabilised [388]. The optimised geometries for the H2O•SiF4, Me(H)O•SiF4 and Me2O•SiF4 1:1 complexes are based on a distorted trigonal bipyramidal structure with the base in an axial position and the three equatorial fluorines displaced towards the base, indicating attack of the base at a face

ip t

rather than edge of the SiF4 tetrahedron [386, 387]. The Si-O bond lengths are basis set dependent but in the region of 2.8 Å for H2O•SiF4, and 2.6 Å for H3C(H)O•SiF4 and

cr

Me2O•SiF4 [386, 387]. Further calculations on the hydrolysis of SiF4 provide some indication that the reaction proceeds through an initial five coordinate transition state perhaps with the

us

formation of a binary pre-reaction complex [385] such as that observed by Ault [68], but with no evidence in the IR spectra [383, 385]. Thus indicating the usefulness of matrix isolation

an

techniques to the study of the initial species formed.

Twin-jet deposition of SiF4 or GeF4 with a selection of oxygen containing heterocycles

M

yielded the following 1:1 complexes: (CH2)2O•SiF4; (CH2)3O•SiF4; (CH2)4O•SiF4; (CH2)2O•GeF4; (CH2)3O•GeF4; and (CH2)4O•GeF4 [389]. Whilst a furan 1:1 complex,

d

(CH)4O•GeF4, was observed for GeF4 with a small shift in the GeF4 modes from the parent,

Ac ce pt e

no νSi-F mode for a complex could be identified, suggesting that no SiF4 complex was formed, or that the interaction was so weak that it was not detectable [389]. In the IR spectrum following single-jet deposition of SiF4 and O(CH2)2, the spectrum contained no trace of O(CH2)2 and only a weak absorption due to SiF4 indicating substantial reaction in the vacuum line prior to deposition. Although it was not possible to identify the structure of the complexes definitively, the sharpness of the bands of both the perturbed acid and base indicated the formation of a well-defined complex (rather than a van der Waals type interaction). The similarities of the spectra to those of the amine complexes [361, 363] was suggestive of similar structures in all cases [389]. The proton affinity of the saturated oxygen heterocycles increases with ring size, and whilst this is mirrored to a certain extent in the νSi-F modes (Table 9), the reverse is observed for the νGe-F modes (Table 10). When SiF4 was added to a condensate of S2O at 77 K and the mixture slowly warmed to room temperature, the IR spectrum of the reaction products indicated the formation of a five coordinate silicon species [390]. The intensity of the bands associated with the reaction products did not increase with time, indicating that the reaction had gone to completion 47 Page 47 of 213

before room temperature was reached. When S2O at room temperature was treated with SiF4 and the mixture cooled to 77 K and then warmed back to room temperature, there was no evidence of any reaction products, presumably as the S2O had decomposed at room temperature before coming into contact with SiF4. As the decomposition products of S2O are SO2 and S (via SO), a series of reactions involving SO2 and SiF4 was also carried out at different temperatures between ambient and 77 K for between 1 and 16 hours, but there was

ip t

no sign of any reaction. The IR data indicates that the most likely reaction product was OS•SiF4 [390]. Stable complexes between SiF4 and aminocyclophosphazenes have been

cr

prepared by the same group, where the stoichiometry was dependent on the ring size and

us

substituents [391].

an

1.9.1.3. GeF4 and GeCl4 complexes with F− and Cl− donors When GeF4 and CsF were codeposited into an Ar matrix the Cs+[GeF5]− ion pair was formed, together with a smaller amount of the well known [GeF6]2− at higher concentrations [392]. When CsF and GeCl4 or CsCl and GeF4 were codeposited with argon, the mixed chloro-

M

fluoro species Cs+[GeClF4]− and Cs+[GeCl4F]− were formed, but in lower yields. The reaction of CsCl and GeCl4 was not successful, which was taken to indicate either the marginal

d

stability of [GeCl5]− or the limitation of the experimental protocol employed. Although three

Ac ce pt e

IR active Ge-F stretches were observed, this was interpreted as a distortion of the trigonal bipyramidal [GeF5]− by the Cs+ counter ion, rather than a lower symmetry of the [GeF5]− unit itself.

1.9.1.4. SiF4 and GeF4 complexes with cyclopropane There was no evidence for the formation of complexes between cyclopropane and either SiF4 or GeF4 [129].

1.9.1.5. SiF4 and GeF4 summary Recent calculations have indicated that the calculated ground state conformers for H3N•SiF4 and H3P•SiF4 are staggered rather than eclipsed [393]. For H2O•SiF4 the preferred conformer is eclipsed with a staggered transition state, whereas for H2S•SiF4 both the eclipsed and staggered are true minima, with the staggered form lower in energy. In each case the base approaches via a face of the SiF4 tetrahedron as noted previously. The energy of interaction decreases sharply NH3 > H2O > HF, with a similar, but more gradual change in the third row PH3 ≈ H2S ≈HCl and this is shown graphically in Figure 11 [393]. 48 Page 48 of 213

The experimental data for the matrix isolated SiF4 and GeF4 complexes are collected together in Tables 9 and 10. Figure 12 displays these graphically for the 1:1 complexes and is based on that of Ault [10, 68], but now includes all the available data. In all cases the highest νSi-F or νGeF mode was used as the triply degenerate mode in SiF4 can yield either 2 or 3 modes in the complexes depending on the symmetry. For comparison the t2 mode of SiF4 is at ca. 1024

ip t

cm−1 and that of GeF4 is at ca. 797 cm−1. The linear relationship between the stretching mode and proton affinity for the complexes of both SiF4 and GeF4 is very good, even for the

cr

N2•SiF4 complex which has a very small shift. This has been taken to imply that SiF4 and GeF4 are acting as π acceptors in these complexes [361, 363]. Therefore for these second row

us

group 15 and 16 donor atom complexes, the base proton affinity is a very good guide to the strength of complex formed. Once complexes of phosphines have been reported it will be

an

interesting to see if the linearity extends to third row elements as well, or whether HSAB

1.9.2.

M

effects come into play.

SiCl4 and HSiCl3 complexes

Ault’s study on SiF4 and NH3 also included experiments on SiCl4 and NH3 and in this latter

d

case the product bands were weak and only observed at relatively high concentrations [361].

Ac ce pt e

Ault assigned these to the 1:1 H3N•SiCl4·complex rather than the hydrogen chloride elimination product, SiCl3NH2, on the basis of isotopic shifts. Whilst product bands were also obtained from the codeposition of SiCl4 with Me2NH, indicating the formation of a complex, it was not possible to identify its structure. No product bands were observed from the reaction of SiCl4 and Me3N indicating that SiCl4 is a weaker Lewis acid than SiF4, and that Me3N is much less reactive than the partially methylated amines [363]. Ammonolysis of SiCl4 is important in the formation of Si3N4 although the initial product formed is polymeric silicon diimide, which requires further thermal processing to convert it to amorphous Si3N4 [394]. Himmel has shown using mass spectrometry as well as gas phase and matrix isolation IR that SiCl3NH2 is an intermediate in the ammonolysis of SiCl4 [395]. Whilst calculations were carried out on both H3N•SiCl4 and (H3N)2•SiCl4, which are postulated intermediates in the formation of SiCl3NH2, no experimental evidence for their existence was presented. The bands reported by Ault in his study of SiCl4 and NH3 using

49 Page 49 of 213

twin-jet deposition [361] do not correspond with those assigned to SiCl3NH2 by Himmel [395], indicating that they are probably due to H3N•SiCl4. The twin-jet deposition of HSiCl3 with NH3, NHMe2, NMe3 and Me2O led to the formation of the 1:1 H3N•SiHCl3, Me2HN•SiHCl3, Me3N•SiHCl3 and Me2O•SiHCl3 complexes where the base donor atom interacts with the silicon centre [396]. The Si-H stretching modes were

ip t

blue shifted, whilst the Cl-Si-H bending mode and Si-Cl stretching modes were red shifted. In line with the other SiF4 complexes the structure appears to be distorted trigonal bipyramidal

cr

with the base ligand in an axial position and H in the equatorial plane. With the amines, there was some evidence for the formation of a second, isomeric form of the complex in which the

us

base was hydrogen bonded to the Si-H. The twin-jet deposition of MeSiCl3 and Me3N or NH3 led to very little complex formation, as expected from the reduction in Lewis acidity when a

an

methyl group is introduced into the Lewis acid.

Both the twin-jet and merged-jet deposition of HSiCl3 and MeOH resulted in the formation of

M

the hydrogen halide elimination product, HCl2SiOMe [397]. The greater yields in the merged-jet experiments indicated that this was a gas phase rather than matrix reaction. In

d

contrast, no new product bands were observed in either twin-jet or merged-jet experiments

Ac ce pt e

involving SiCl4 and MeOH even when the merged-jet reaction zone was heated to 200°C [397]. However, subsequent mass spectrometric and matrix isolation experiments have been carried out by Almond and Ogden [398] who used both gas phase and higher temperature pyrolysis. If SiCl4 and MeOH were allowed to react for three hours prior to analysis by mass spectrometry, then peaks characteristic of HCl elimination products such as SiCl3OMe were observed, which were absent when a merged-jet approach was used. If the sample that had been mixed for three hours was then passed through a tube at 650°C, no peaks from SiCl3OH were observed, but peaks most likely due to MeCl were observed. This suggests that SiCl3OMe decomposes thermally to eliminate MeCl. When merged-jet deposition was used in conjunction with matrix isolation, no product/complex peaks were observed. However, if the premixed sample was used, bands due to HCl (most likely present as Me(H)O···ClH) and SiCl3OMe were observed. When the SiCl4, MeOH and Ar mixture was passed through a pyrolysis tube at 350°C no change in the spectrum was observed. When the pyrolysis temperature was increased to 750°C new IR peaks were observed, indicating the formation of one or several Si-Cl containing products. Unfortunately, they could not be assigned due to low intensity, but no MeCl was observed. In the analogous ethanol experiments it was found 50 Page 50 of 213

that much longer mixing times (18 rather than 3 hours) were required, but the product was the hydrogen chloride elimination product, SiCl3OC2H5 [398]. When the experiments were repeated with Si2Cl6 the same products were observed, but via Si-Si bond cleavage.

1.9.3.

SiBr4 and SiI4 complexes

ip t

There appear to be no reports of matrix coordination chemistry involving SiBr4 and SiI4,

cr

indicating the weakness of their Lewis acidity.

1.10. SiH4, GeH4, SnH4, PbH4 complexes

us

Codeposition of SiH4 with ammonia and other amines in solid argon led to the first experimental identification of the 1:1 complexes H3N•SiH4, MeH2N•SiH4, Me2HN•SiH4 and Me3N•SiH4 [399]. These displayed an increasingly red-shifted νSiH mode from 2177 (and

an

2215 cm−1) in SiH4 [400] to 2175 cm−1 (NH3) through 2151 cm−1 (NH2Me), 2136 cm−1 (NHMe2) to 2111 cm−1 (NMe3) with increasing proton affinity of the base. There was also

M

some evidence for formation of (H3N)x•SiH4·where x was most likely to be 2 [399]. Calculations [365] of the structures of the 1:1 complexes indicate the NH3 approaches the Si atom along a C3 axis, leading to a trigonal bipyramidal structure, as found for the H3N•SiF4

d

complex [361, 363]. Calculations indicate the lack of a donor acceptor bond in H3N•SiH4 as

Ac ce pt e

the Si-N distance between the molecules was 3.2-3.4 Å, with a bond dissociation energy (ΔH°298) of 1.4 kJ mol−1 [370].The most recent calculations indicated that there were three ways that the initially weakly bonded molecular complex H3N•SiH4 eliminates the first hydrogen molecule [401]. The interaction between SiH4 and the amines is relatively weak with only small shifts (ca. 30 cm−1) of the perturbed vibrational modes in the complex compared to the parent compounds. The rotational structure observed for the “umbrella” mode of monomeric NH3 was maintained for the complex in Ar but not N2 matrices. When F2 was deposited with SiH4 and GeH4 in argon matrices, only weak hydrogen bonded HF complexes were observed [402]. The codeposition of nitrous acid in argon matrices with SiH4 and GeH4 led to the formation of well defined 1:1 complexes with trans- and cis-HONO. For the SiH4 and GeH4 complexes there appear to be weak Si···O and Ge···O electron donor interactions with the OH group of HONO, rather than weak hydrogen bonded interactions as found for CH4 [403]. 51 Page 51 of 213

1.11. Silicon dihalides and related carbene analogues A general review of divalent silicon, germanium, and tin compounds with bonds to heteroatoms was published in 2004 [404]. 1.11.1.

Silicon, germanium, tin and lead dihalides

cr

bases such as N2, CO, NO, H2, O2 were reviewed in 2004 [405].

ip t

The complexes of carbene analogues EX2 (E = Si, Ge, Sn, Pb; X = F, Cl, Br, I) with Lewis

The reaction of SiF2 and SiCl2 with O2 in Ar matrices only yielded difluorodioxasilirane,

us

F2Si(O2), and dichlorodioxasilirane, Cl2Si(O2), after photolysis, indicating their low thermal

an

reactivity [406].

A N2•SiCl2 donor-acceptor complex between N2 and SiCl2 in nitrogen matrices had red shifts of about 1 cm−1 for both the symmetric and asymmetric SiCl stretching modes [407]. This

M

compared to 5 and 6 cm−1 for N2•SiF2 [408], 4.3 cm−1 for N2•SnCl2 [409] and 5 cm−1 for N2•PbF2 [409]. No νNN modes were observed, presumably because of the weakness of the

d

interaction, although νNN has been observed for Cp*Si(N2) at 2046 and 2053 cm−1 in liquid

Ac ce pt e

Xe and liquid nitrogen, respectively, using long path length cuvettes [410]. In addition to the 1:1 complexes, there was also evidence for 2:1 complexes, (N2)2•SiCl2 [407], (N2)2•SnF2 [408] and (N2)2•SnCl2 [409] which demonstrated red shifts about twice those of the 1:1 complexes. In addition to N2 complexes, Tevault and Nakamoto also studied CO complexes of SnCl2, PbF2, PbCl2, PbBr2 and PbI2 and NO complexes of SnCl2 and PbF2 [409]. The calculated structures for the mono and bis complexes are shown in Figure 13 [405]. For SiCl2 the N2 unit was 3° off normal to the SiCl2 plane for both the 1:1 and 2:1 complexes, as it was for the 1:1 complex with GeCl2. For the 2:1 complex this had increased to 6°. For SnCl2 the angles were 10° and 13° from normal for the 1:1 and 2:1 complexes respectively. For SnF2 these increased to 13° and 27°. The reaction of SiCl2 (from pyrolysis of Si2Cl6 at 600-800 K) with N2O, NO, O2 yielded SiCl2O polymers but with no experimental evidence for the formation of complexes [411]. The complexes formed between SnF2 and hept-1-yne [412], methyl-chloride [413], benzene, chlorobenzene and toluene [414] trapped in cryogenic matrices have been reported. 52 Page 52 of 213

1.11.2.

SiH2 GeH2 and SnH2 complexes

The complexes of H2Si and Me2Si with CO were reviewed in 2004 [405]. More recently, Maier et al. have reported that SiH2 forms a donor-acceptor end-on complex, N2•SiH2, in the presence of dinitrogen under matrix photolysis of diazidosilane or by combination of flash

ip t

pyrolysis of 1,1,1-trimethyldisilane with co-condensation of the products with N2 at 10 K [415]. The analogous CO complex has also been reported [416].

cr

121 nm vac-UV photolysis of silane in a nitrogen matrix has resulted in the characterisation

us

of a number of transient species, SiN2, Si(N2)2, HSiN2, and H2SiN2 [417].

The reaction of germane and ozone yields the H2O•GeH2 complex amongst six other species

an

[418]. A review of germylenes and stannylenes, including some matrix data was published in

1.11.3.

Organic Silylenes

M

1991 [419].

d

The reactivity of organic silylenes arises from their combined Lewis acid and base potential because of the presence of both a lone pair and a vacant p orbital on the silicon atom. Indeed,

Ac ce pt e

one way to tame the reactivity of the silylenes has been via the formation of complexes with Lewis bases.

The first observation of bright yellow Me2Si in hydrocarbon glasses and inert gas matrices was made in 1979, and it was noted that it could be used as a preformed reagent as it was stable indefinitely in hydrocarbon glasses at 77 K [420]. Until 1994 silylenes were only stable in matrices [421]. Although the reactions of dimethylsilylene with amines in low-temperature matrices have been studied, they only appear to have been reported in conference abstracts [422] although bis-(amino)silylene has been prepared in argon matrices via the photolysis of silicon diazide [423]. As matrix carbene chemistry has been reviewed in 1992 [424], 1993 [425] 1999 [426], 2010 [427] and 2011 [428], it will not be covered in anymore detail here.

53 Page 53 of 213

1.12. Silicon oxide complexes Laser ablated silicon atoms codeposited with Ar/O2 mixtures yield SiO2 and SiO as the primary products [429]. If the matrix was doped with NH3 then on annealing a H3N•SiO2 complex was formed, which on broad band photolysis decomposed to give H2NSiOOH. There was tentative assignment of a weak feature at 1241.1 cm−1 to HSiONH2 on the basis of isotopic shifts and DFT calculations. There was no evidence for H3N•Si or HSiNH2 in these

ip t

experiments [430]. When water was used instead of ammonia, silicic acid (H2SiO3) was produced spontaneously on annealing [431]. Although the calculations predicted that a

cr

H2O•SiO2 complex would be stable, there was no experimental evidence for it, and this was explained as being due to the low energy barrier and high exothermicity of the hydrolysis

us

reaction. SiCO4 was formed in the analogous reactions with CO2, with no evidence for an intermediate complex. However, the calculated structure of SiCO4, could be considered as

an

O=Si(CO3) with the carbonate acting as a bidentate ligand [432].

1.13. Carbon suboxide complexes

M

C3O2 is important in astrochemistry and astrophysics as it is found in dense molecular clouds [433], and has been studied in ice analogues [433, 434]. Experimental (single-jet) and

d

computational investigations of the C3O2 HCl 1:1 complex (as well as those of related oxocumulenes) indicated a T-shaped structure with the H interacting with the central, Cβ (C3),

Ac ce pt e

of C3O2 [435, 436]. This was then extended to include the interaction of C3O2 with amorphous surfaces of HCl hydrates [437]. For the 1:1 H2O and NH3 complexes of C3O2, the experimental data and calculations revealed a structure described as T-shaped, which was indicative of nucleophilic attack by the O or N atom on the C4 of the C3O2moiety, with a N-C distance of 3.08 Å [438]. The ammonia complex is stabilised by an N-C2 distance of 3.16 Å and a 3.01 Å hydrogen bond between the terminal O of C3O2 and H of NH3. These geometries were taken to be suggestive of the formation of pre-reactive intermediates of 1,4 nucleophilic addition [438]. For the pyridine C3O2 1:1 complex, a T-shaped structure was again found from both experimental and computational work, and these indicated that the N atom of the pyridine was attached to the C2 and C4 atoms of C3O2 via electrostatic interactions with distances of 3.28 Å [439]. A similar situation was found for the thiazole 1:1 complex with C3O2 with a N interacting with C2 together with hydrogen bonds between the oxygen atoms of C3O2 and H(C) adjacent to the N on the thiazole, again indicative of a prereactive intermediate of nucleophilic addition [440]. This work was developed further with ammonia, dimethylamine, trimethylamine and 4-(dimethylamino)pyridine forming 1:1 54 Page 54 of 213

complexes with binding energies in the range of −10 to −20 kJ mol−1. In addition it was found that by warming the deposit from 70 to 190 K (depending on the amine) the complexes could be converted into C3O2-amine zwitterions [441]. For the C2H2 C3O2 1:1 complex the calculations indicated three geometries at similar energies, and although the T-shaped complex with the acetylenic H interacting with the central C3O2 carbon was not the lowest in energy (by 3.4 kJ mol−1), the experimental IR spectra fitted this much better than the other

ip t

two. For NCCCH, there is the possibility of nucleophilic attack by the N at Cα (C2) of C3O2 or electrophilic attack by the H at the central Cβ (C3) carbon of C3O2. The best fit between

cr

calculation and experiment was for the latter one. Therefore, for both acetylene and cyanoacetylene, the 1:1 complexes are indicative of electrophilic attack at the central carbon

us

of C3O2 [442].

an

Both experimental [443, 444] and computational [445] data have been reported for the

M

H2O•C3 complex.

1.14. Mössbauer spectroscopy of matrix isolated tin compounds Whilst the vast majority of the data presented in this review are based on IR experiments, it

d

should be noted that for tin, and especially

119

Sn, Mössbauer spectroscopy is also a

Ac ce pt e

possibility. Whilst there do not appear to be any coordination complexes reported in matrices, data is available for the parent compounds [446-455].

Group 15

1.15. Introduction

In the vast majority of Lewis acid-base complexes with linkages involving group 15 atoms, the group 15 atom is acting as a Lewis base, although there are some exceptions where the group 15 atom can function either as a Lewis base or be amphoteric.

1.16. Complexes of Group 15 hydrides NH3 is one of the archetypal Lewis bases. The ν2 symmetric deformation mode of NH3, often referred to as the “umbrella” mode, has been used by many workers as a probe of the strength of the interaction between NH3 and the Lewis acid. Many of the matrix isolated NH3 complexes have been dealt with in other sections, and the values of the ν2 modes for these 55 Page 55 of 213

complexes are listed in Table 11 and 12. It should be noted that NH3 in argon matrices gives rise to complex spectra because of inversion and hindered rotation [456], in addition to dimerisation which is also encountered in other matrices. The spectra in N2 matrices are much simpler as the rotational motion is quenched. A tabulation of the ν2 modes of the heavier group 15 hydrides is given in Table 13 and as for NH3 they are dealt with in other sections. Figure 14 shows a plot of the NH3 symmetric deformation versus the nitrogen –

ip t

acceptor bond lengths derived from gas phase (usually microwave) studies. The correlation from the strongly bound H3N•BF3 and H3N•BH3 complexes through to the more weakly

cr

bound halogen complexes, including the intermediate H3N•SO3 and H3N•SiF4 complexes appears to be good. However, as the points for the H3N•BF3 and H3N•BH3 complexes are

us

expected to be offset from the others as boron is smaller, the correlation involving the third row acceptor atoms is very good and means that estimates of the donor-acceptor bond length

an

can be made for the matrix data in Table 13 for which gas phase structural data are not available.

M

The structures of the NH3 and PH3 dimers have been the subject of a computational study [457] and the coordination chemistry of stibine and bismuthine ligands were reviewed in

d

2006 [458]. Many hydrogen bonded examples in matrices are known and these include: NH3

Ac ce pt e

and the hydrogen halides [459-463]; PH3, AsH3 and SbH3 with hydrogen halides [348]; alkyl amines with HF [464, 465]; NH3 and HCN [466]; NH3 and HONO [467]; the NH3-HOO complex [468]; as well as N2H4 with HF, F2 and O3 [469], CO [470], and HC3N [471]. The formation of NH3 from N atoms and H atoms [472] and N2 and H atoms [473] has been reported, as well as the heavier hydrides (Bi, Sb) via laser ablation [474]. The weakly hydrogen bonded complexes of PH3 and AsH3 were reviewed in 1994 [475].

1.17. NH3 complexes with CO2 and CO CO2 and NH3 are important constituents of the interstellar medium, and have been identified in the icy mantle of interstellar grains, which are also subject to VUV radiation. It is known that low temperature irradiation of NH3 CO2 ice results in the formation of carbamic acid [476, 477], and photolysis of methylamine and CO2 leads to glycine [478]. Bossa et al. [479] building on earlier work of Fredin and Nelander [480] have identified the T-shaped H3N•CO2 complex (Figure 15) with the nitrogen bound to the carbon atom. A (H3N)2•CO2 complex where the second NH3 subunit bridges the T of the (NH3)2•CO2 unit via hydrogen bonding 56 Page 56 of 213

was also observed. Whilst there was evidence for a H3N•(CO2)2 complex, the data was not sufficient for any conclusions about its structure to be made. This is in contrast to the complex formed between NH3 and CO, after photolysis of formamide in solid argon, where surprisingly both the H2NH•CO and H2NH•OC hydrogen bonded complexes were observed with binding energies of −14.6 and −7.1 kJ mol−1 and the oxygen bonded complex dominated

ip t

the spectrum [481].

1.18. ClNO complexes

ClNO is an intriguing molecule with an interesting structure which shows considerable ionic

cr

boding, (ON)+Cl−, and it can act as both a Lewis base and as a Lewis acid [482, 483].

us

Pimentel et al. studied the Lewis base properties of ClNO using HX and X2 (X = Cl, Br) [484]. There were small shifts in the HX modes, indicating weak H-bonds of 4 - 12 kJ mol−1. However, the blue- shifts in the N-O stretching modes were more substantial (≈ 30 cm−1 for

an

HCl, ≈ 20 cm−1 for HBr, ≈ 16 cm−1 for Cl2 and ≈ 24 cm−1 for Br2 ) and these were attributed to a donation of electron density from an anti-bonding orbital associated with NO to the acid.

M

[484]. A subsequent experimental and computational study indicated that the lowest energy structure of the ClNO HCl complex involves simultaneous hydrogen bonds between the H and all three atoms of ClNO (r(H...O) 3.0 Å, r(H...Cl) 2.44 Å, r(H...N) 3.3 Å) as well as a longer

d

interaction (ca. 3 Å) between the Cl and the O of ClNO [485]. The ONCl H2O 1:1 complex is

Ac ce pt e

weakly hydrogen bonded, and the lowest energy structure is where the H2O molecule simultaneously interacts with both the ClN and NO bonds [486]. Photolysis of this complex using 266 nm light leads to formation of trans- and cis-HONO HCl 1:1 hydrogen bound complexes [487]. The dimer is centrosymmetric in both argon and nitrogen matrices [483, 488].

Ault investigated the Lewis acid properties of ClNO using both lone pair donors (Me2CO, Et2O, Me2Se, Me2S,) and π electron donors (benzene, toluene, bromobenzene, chlorobenzene, hexafluorobenzene and hexamethylbenzene) [482]. The νNO mode in the complexes was blueshifted from that at 1805 cm−1 in argon indicating a reduction in the π* occupancy of the NO bond which is counter intuitive to that expected if ClNO is acting as a Lewis acid, and is in the same direction as that when ClNO is acting as a Lewis base [484]. The νNCl mode was red-shifted (from 585 cm−1) as was the bending mode. Ault suggested that the cause of the blue-shift in the νNO mode could be due to simultaneous redistribution within the ClNO from the ON to NCl bond following donation of electron density from the base to ClNO. Other 57 Page 57 of 213

possibilities included a change in the ClNO bond angle and polarisation of the ClNO unit by the base. ClNO is a bent triatomic molecule with a 113° bond angle at the central nitrogen [489], which is thought to carry a net positive charge and so this was proposed as the site of the Lewis acidity [482]. A representation of the structure is given in Figure 16. In the case of the benzene π-electron donors it was thought that the ClNO would be sitting above the plane

The experimental vibrational data is collected together in Table 14.

cr

1.19. NF3 and NCl3 complexes

ip t

with the N directed towards the π system.

us

NF3 is extensively used in the electronics industry for etching and cleaning and it can be considered as a bifunctional Lewis base with coordination via either the N or F. Calculations have provided the first theoretical evidence for BH3 and BF3 complexes demonstrating both

an

modes of coordination [490]. The indications are that at the highest level of theory, the Fcoordinated and N-coordinated complexes are practically degenerate and that the small bond

M

dissociation energies of 4-8 kJ mol−1 indicate that they will only be stable at low temperatures. Other than its HF hydrogen bonded complex [491] there do not appear to be any reports of its Lewis acid base behaviour in the matrix literature. Vacuum-ultraviolet

d

photolysis of NF3 in an argon matrix yields the NF2 radical and in a CO matrix FCO is

Ac ce pt e

formed [492]. NF3 has also found use as an F atom source for matrix studies when passed through a microwave discharge [493].

Although there are reports on matrix isolated NCl3 and its photochemistry [494-496], there do not appear to be any concerned with complex formation in cryogenic matrices. Boutalib has carried out calculations on the complexes of NF3 and NCl3 with alane and methylalanes [304].

1.20. Phosphorus trihalides as Lewis acids Whilst the phosphorus trihalides would normally be considered as Lewis bases, and their complexes are covered in other sections and Table 13, they can also display some Lewis acid behaviour. A study of the gas phase reaction between CsF and PF3 with subsequent trapping in an argon matrix allowed for the first identification of the tetrafluorophosphoranyl anion, [PF4]− paired with a Cs+ cation. The IR data indicated a structure of the anion with no higher 58 Page 58 of 213

than C2v symmetry, with the Cs cation in axial monodentate coordination. Analogous spectra of Cs+[PClF3]− were obtained from the reaction of CsCl and PF3, but no reaction was observed between CsCl and PCl3. These experiments indicate that PCl3 is a weaker Lewis acid than PF3 [497]. No complex formation was observed between PF3 or PCl3 and OCS but photolysis yielded

ip t

CO and SPF3 or SPCl3 [498]. Although CO was observed no other product was observed for

cr

NF3, PH3, AsF3 or HCl [498].

Although not strictly Lewis acid base complexes, a recent computational paper re-assigning

us

bands in OPCl, OPBr, OPI, OPI3, OSbF3, OSbCl3, OPCN, OPSCN, NSCN, and NSSCN is

an

worth noting [499].

1.21. Complexes involving PF5, AsF5, SbF5

M

The H3N•PF5 complex, produced in 8% yield from NH3 and PF5 is sufficiently stable to allow for its single crystal structure to be determined [500]. AsF5 and SbF5 are widely used Lewis acids, and although their matrix spectra have been reported [501], there are no further matrix

d

reports involving acid-base complexes, but there is some dispute over their structure [502,

Ac ce pt e

503]. However, Group 15 halides have been used as ligands in matrix experiments involving AsF3 with Mo and W [504, 505], NF3, PF3, AsF3 with Th [506] Ti, Zr, Hf [507] and U [508].

Group 16

1.22. Introduction

Many of the common bases contain group 16 donor atoms, so have been discussed in detail in other sections. The coordination chemistry of seleno and telluroethers has been reviewed [509, 510].

1.23. H2O complexes As with NH3 discussed in the previous section, H2O is ubiquitous in acid-base chemistry. Depending on the pH it can act as either a Brønsted acid or base in solution, and also as either a Lewis acid or base. The complexes involving H2O acting as a Lewis base are largely 59 Page 59 of 213

covered in other sections, whilst the vast majority of those involving it as a Lewis acid are essentially examples of hydrogen bonding and are outside the scope of this review. Earlier reviews of computational work [511, 512] will provide a source of information. For a review of water monomers and higher multimers in matrices see [513, 514]. The H2O•CO2 1:1 complex is T shaped with the C atom bound to O of water (Figure 17), and

ip t

the matrix IR data indicated there were two isomers, one C2v, the other Cs. Calculations indicated that the stabilisation energy was −8.2 kJ mol−1, and the difference between the two

cr

isomers was less than 0.05 kJ mol−1 [515]. In contrast the HOO CO2 complex has a cyclic structure which is essentially an H bonding interaction with the peroxy radical hydrogen

us

bonded to the lone pair of an oxygen on CO2 and the terminal oxygen as close as possible to C of CO2 [516, 517]. The CO H2O complex involves a hydrogen bond to the C of CO [518].

an

The N2 H2O hydrogen bonded complex has νNN at 2331.5 cm−1 in Ne [519] compared to 2329 cm−1 for N2 in an argon matrix [520]. The structure has a quasi-linear hydrogen bond, OH..N-N, in argon [521-523]. The O2 and H2O complex is also hydrogen bonded and matrix

M

isolation provided the first spectroscopic characterisation of this atmospherically important complex which had been predicted for many years [524, 525]. The H2O Xe and Kr 1:1

d

complexes have non-hydrogen bonded geometries in argon matrices [526]. H2O•CF4 has a

Ac ce pt e

non-hydrogen bonded structure with the oxygen atom of water oriented toward the carbon atom of CF4, but the oxygen atom does not lie on the C–F bond axis [527]. The H2 (ortho and para) complexes with H2O have been investigated by a number of authors in Ne [528], pH2 [528-530] or oD2 matrices [530].

Garden et al. have calculated counterpoise corrected geometries of the hydrated complexes H2O•CO2, H2O•CS2, H2O•OCS, H2O•SO2 and H2O•SO3 [531]. H2O•CO2 is T-shaped, but the H2O•CS2, H2O•OCS complexes would be better described as H2O•SCS and H2O•SCO, as the water oxygen atom is interacting with the sulfur in each case. The H2O•SO2 and H2O•SO3 complexes have double-decker, eclipsed structures but the two planes are tilted with respect to each other by around 10°. The binding energies were calculated as: H2O•CO2, −7.9 kJ mol−1; H2O•SCS, −4.5 kJ mol−1; H2O•SCO, −5.4 kJ mol−1; H2O•SO2, −12.1 kJ mol−1; and H2O•SO3 −26.7 kJ mol−1 [531].

60 Page 60 of 213

1.24. H2S complexes Many of these have been considered elsewhere and the weakly hydrogen bonded complexes of H2S and H2Se were reviewed in 1994 [475]. Photochemistry (193 nm) of H2S/SO2 mixtures in solid Kr gave evidence of HSO2 and cis-HOSO, but no intermediate complex was reported [532]. Matrix isolation and ab initio studies of the H2S CO hydrogen bonded complex indicated two stable minima with nearly linear hydrogen bonds between the units

ip t

with binding energies of −5.22 and −1.54 kJ mol−1 for HSH•CO and HSH•OC, respectively [533]. The experimental data were consistent with the C bound complex. The photo-oxidation

cr

of H2S in solid O2 yields the H2O•SO3 complex [534]. However for O3 H2S mixtures there was no evidence for an intermediate complex before the photolytic formation of HSOH

us

[535].

1.25. H2O2 complexes

an

Most H2O2 complexes simply involve hydrogen bonds, however whilst the H2O2•SO2 complex does have a hydrogen bond between H and O of SO2, there is also a Lewis acid-

M

base interaction between electron lone pair of the other O in H2O2 and the sulfur atom in SO2 [536]. Photolysis of this complex gives the H2O•SO3 complex [536]. The H2O2 dimer is

d

cyclic with both components acting as hydrogen bond donors and acceptors [537]. The weakly hydrogen bound H2O2•O3complex gives the higher oxide, H2O3, on photolysis which

Ac ce pt e

is important for the understanding of hydrogen oxygen radical chemistry [538, 539]. Other examples of hydrogen bonded H2O2 complexes contain the following: H2O [540], O3 [539], Me2O [541], NH3 and NMe3 [542], PH3, PMe3, H2S and Me2S [543], HF, HCl and HBr [544] and CO [545]. UV photolysis of H2O2 in solid Ar, Kr or Xe results in a complex between H2O and a ground state 3P oxygen atom as well as OH• radicals [546]. Calculations indicate that there are two interactions in ozone hydrogen peroxide complexes, one a conventional hydrogen bond, the other involving an O...O contact [547].

1.26. O3 complexes Ozone can act as a Lewis acid when coordinated by the apex oxygen to give a symmetrical structure, or as a Lewis base via one of the terminal oxygen atoms resulting in an asymmetric structure. The mode of coordination can be determined by use of

16

O/18O (1:1) scrambled

ozone where coordination via the apex results in a 1:2:1:1:2:1 sextet for the asymmetric ν3 stretching mode, as found for the F3P•O3 complex [12] where the O3 is acting as a Lewis 61 Page 61 of 213

acid. In contrast, when the coordination is via a terminal oxygen acting as an electron donor, such as in O3•ICF3, an octet with equal intensity components is observed [548]. Ozone complexes were briefly reviewed by Andrews in 1989 [549]. Many of the complexes have been discussed in earlier sections, and Table 15 collects together the available data on the O3 ν3 and ν2 stretching modes in these and the complexes discussed below.

ip t

For many ozone complexes long wavelength (red >500 nm) photolysis can be used by making use of a charge transfer mechanism, rather than accessing the O3 Hartley continuum

cr

band that requires much shorter wavelengths (ca 250 nm) [12, 550].

us

The ozone dimer has two ν3 bands, one higher than the monomer, one lower, and this doublet was taken to mean that the complex was not centrosymmetric, and because the shifts were

an

small the interaction was weak [551].

Nord studied complexes of ozone with ammonia, water, formaldehyde, hydrogen bromide

M

and ethylene in nitrogen matrices by means of infrared spectroscopy. In the lone pair donor complexes, H3N•O3, H2O•O3, H2CO•O3, both the ν3 and ν2 modes of the O3 sub-unit were

d

blue shifted compared to “free” O3. In contrast in the O3•HBr complex (where HBr acts as a σ

Ac ce pt e

acceptor), the ν3 mode was red-shifted, whilst the blue-shift of the ν2 mode was greater than for the lone pair donor complexes. For the π donor H2CCH2•O3 complex the ν3 mode was red-shifted [552]. Although it is tempting to use these different perturbations of the O3 modes to identify the nature of the complex, Nord advised against this [552]. This caution has been reinforced more recently, where there did not seem to be a correlation between red- or blueshifted O3 ν3 modes and the nature of the bonding (hydrogen bonding or charge transfer), with a comment that accurate calculations are required to identify the structural type present [551].

1.26.1.

Ozone hydrogen bonded complexes

Hydrogen bonded complexes between O3 and HF, [553] HCN [554]; H2O [555], the •OH radical, [556] and H2O2 [539, 547] among others have been reported.

1.26.2.

Ozone group 12 complexes

These are considered in section 1.2. 62 Page 62 of 213

1.26.3.

Ozone group 13 complexes

These are considered in section 1.7. 1.26.4.

Ozone group 14 complexes

There was no evidence of an intermediate complex in matrices containing ozone and SiH4, which was used to explain the fact that full output of the Hg lamp (220 -1000 nm) was

ip t

required, rather than red light for other well established complexes [550]. In the matrix reaction of O3 with CS2 (or OCS) there was no evidence presented for an intermediate

cr

complex in the formation of OCS, SO2, CO (or CO2, CO, SO3, SO2) [557].

us

Ault has used matrix isolation techniques to study the complexes and reactions between ozone and alkenes in the search for the elusive Criegee intermediates in the reaction of ozone with alkenes in the troposphere [558-562]. The Criegee mechanism involves the initial

an

formation of a primary ozonide via a 1,3 polar addition across the alkene double bond, resulting in a 1, 2, 3-trioxolane species. As this reaction is fairly exothermic, further reactions

M

and decomposition forming an aldehyde and the proposed carbonyl oxide, Criegee intermediate. These may then recombine to form a secondary ozonide (1, 2, 4-trioxolane) and further reaction products. Using twin-jet codeposition of ozone and either cyclopentadiene or

d

cyclopentene Ault presented evidence for the long sought after Criegee intermediate [559].

Ac ce pt e

Subsequent work has identified the Criegee intermediates in the reaction of ozone with cis-2butene [560], 2,3-dimethyl-2-butene [561]. In the case of propene, merged-jet deposition led to “later” products consistent with Criegee mechanism, however in the twin-jet experiments no reaction was observed unless the matrices were photolysed [562]. Deng et al. have continued this work with an investigation of the reaction of ozone with methacrolein and isoprene, which also supported the Criegee mechanism [563]. A selective vibrational excitation of nearest neighbour reactant pairs of O3 and C2H2, C2H4, NO2 or NO (also F2 and C2H4) resulted in the assignment of features to “nearest neighbour pairs”, but as the perturbed O3 ν3 modes were similar to many of those in Table 15 they have been included for completeness [564]. As part of a study of the excitation of O3 via vibrational transfer from the CO fundamental at 2138 cm−1 in solid argon (ν1 + ν3 for O3 is only 30 cm−1 below this), a OC•O3 complex was observed with νCO bands at 2140.44 cm−1, 2 cm−1 above the free CO value indicating a relatively weakly bound complex [565]. The ozone ν3 mode was also only very slightly blue-shifted by a couple of wavenumbers. 266 nm 63 Page 63 of 213

photolysis resulted in the formation of CO2, which it was shown was only produced from the OC•O3 complex. At room temperature, CO reacts thermally with O3 to produce CO2 [565]. 1.26.5.

Ozone group 15 complexes 1.26.5.1. NH3, PH3, AsH3 SbH3 ozone complexes

In the H3N•O3 complex all the ozone sub-molecule vibrational modes were blue shifted from

ip t

isolated ozone, as was the ν2 mode of NH3 [552, 566]. Photolysis (290-1000 nm) resulted in the destruction of this complex and formation of NH2OH, together with the hydrogen bonded

cr

H3N•HONH2 complex. The calculated structure of the H3N•O3 complex was the same as that for the H3N•SO2 complex where the O3 unit is nearly perpendicular to the C3 axis of NH3,

us

with the central O slightly off this axis (Figure 18). The dissociation energy of the H3N•O3 complex was 9.4 kJ mol−1, but this was thought to be about half the true value because of the

an

minimum basis used [567].

PH3 and O3 deposited from separate spray on lines (twin-jet deposition) resulted in the

M

formation of a H3P•O3 complex. This was partially photobleached by 590 – 1000 nm irradiation and completely photobleached by 515 – 1000 nm irradiation [12], indicating that

d

complex formation enhances lower energy photolysis. On annealing to 30 K the complex was re-formed. The IR spectra of the isotopically substituted species and comparison with the

Ac ce pt e

calculated structure of H3N•O3 [567] indicated that the H3P•O3 complex was symmetric, with the ozone molecule lying in a plane perpendicular to the C3 axis of the PH3 unit and with apical oxygen of the ozone subunit interacting with the phosphorus. The ozone bond angle was the same as in free ozone, within experimental error. The structure was explained on the basis of the phosphorus lone pair interacting with the π* orbital on the apex oxygen of ozone. Analogous structures were observed for the H3N•O3 [12, 552, 566] and the H3As•O3 [568] complexes. In contrast, the structure of the CF3I•O3 complex is asymmetric because of the increased number of lone pairs on the iodine. There is little evidence for charge transfer in the H3P•O3 complex as the vibrational modes were only slightly shifted from those of the parents. The photochemistry involving red visible photolysis via charge transfer excitation resulted in seven products [569]. The primary products were the structural isomers H3PO and H2POH (which could not be interconverted by further irradiation), and (HO)2HPO. HPO was detected as an intermediate species in the formation of HOOPO and HP(O2)O. The concentration of (HOPO2) increased throughout the irradiation process [569]. A variety of other P/O/H species were observed when PH3 was reacted with oxygen atoms [570]. The 64 Page 64 of 213

reaction of P2 and P4 with O3 was spontaneous with no evidence for an intermediate ozone complex [571, 572] and the reactivity of As4 towards ozone was greater than P4 [573]. In the case of the reaction of AsH3 with O3, the initial complex that formed was very similar to that observed for H3P•O3 with red-shifted O3 modes, a red-shifted AsH3 symmetric bending mode and a blue shifted AsH3 stretching mode, together small amounts of cis- and

ip t

trans-H2AsOH. [568]. Red visible photolysis yielded H3AsO, H2AsOH and HAsO (tentatively). Higher energy photolysis destroyed the HAsO intermediate with the formation

cr

of HOAsO2. This behaviour parallels that of PH3, but AsH3 was slightly more reactive. With SbH3 there was much greater evidence of product formation (HSbO3 and H2SbOH) on

us

deposition in addition to a H3Sb•O3 complex [574]. Red photolysis generated H3SbO. The HSbO3 was destroyed by UV photolysis, which also decreased the amount of H3SbO but

an

increased the proportion of H2SbOH, together with bands probably from HOSbO2. The reactivity of the Group 15 hydrides towards O3 increased down the group as there were

M

no codeposition reactions products observed for NH3 and PH3 until after photolysis of the initial 1:1 complex. However, H2AsOH and H2SbOH were observed with AsH3 and SbH3,

d

respectively, but photolysis was required for the deuterated analogues. Whilst the HSbO3 ozonide compound was formed on deposition, no such species was observed for AsH3. On

Ac ce pt e

the basis of the red shifts in the ν3 asymmetric stretching mode of O3, (O3 1039.9 cm−1; H3N•O3 1047 cm−1; H3P•O3 1037.2 cm−1; H3As•O3 1035.5 [568]; H3Sb•O3 1032.0 cm−1 [574]) the interaction appears slightly larger in the H3Sb•O3 and H3As•O3 complexes than in the H3P•O3 complex.

The co-condensation of PPh3 in O2 matrices yielded Ph3P=O in 90% yield and Ph2(Ph-O)P=O in 10% yield, but only after photolysis [575].

1.26.5.2. PX3 ozone complexes In analogous work on the reaction of phosphorus halides with ozone, X3P•O3 complexes have been observed, but the detailed vibrational data does not appear to have been reported/tabulated for them [576-578]. In contrast to the hydrides, there was significant wavelength dependence in the photolysis and it was noted that the highest yields of the red photolysis products were for PH3, AsH3, P4 and PCl3, all of which form strong bonds to 65 Page 65 of 213

oxygen [578]. It was also noted that stronger ozone complexes can activate the ozone towards red photochemistry, and this was suggested to arise because the complex lowers the barrier to dissociation by providing a strongly exothermic dissociation-recombination process [578]. The reaction of AsCl3 and O3 in matrices led to the first characterisation of OAsCl3 [576].

Ozone group 16 complexes

ip t

1.26.6.

When ozone was allowed to react with Me2S in argon matrices, Me2SO was formed on deposition with no sign of a complex [579], but a more recent article indicates that a Me2S•O3

cr

complex is formed, which converts to Me2SO on red light rather than visible photolysis

us

[580]. A weak complex between MeOH and O3 was identified when ozone was used as a source of O atoms for the oxidation of CH4 and MeOH, but it was not involved in the

1.26.7.

Ozone group 17 complexes 1.26.7.1.

M

H2S before photolytic formation of HSOH [535].

an

reaction pathways [581]. There was no evidence for an intermediate complex between O3 and

Ozone halogen complexes

d

Reactions between F atoms (from photolysis of F2) and ozone in argon matrices resulted in

Ac ce pt e

the first observation of the FO•O2 complex, with νOO at 1522 cm−1 and νFO at 968 cm−1, redshifted by 34 and 61 cm−1, respectively, from their “free” values [582]. On 532 nm photolysis, the complex decayed and was replaced by FO. The data indicate that the complex was not a weak van der Waals complex, but was an asymmetric association complex, where an F atom has added to a terminal atom of the O3 molecule. The reaction of Cl atoms and O3 has been studied by two groups, with slightly different products reported, of either ClOO• and OClO as the primary products [583, 584] or ClOO•, ClClO and ClO• as the primary products [585]. A study of complexes between ozone and Br2, Cl2 or BrCl [586] was a reinvestigation of a previous report [587] where the perturbed ν3 mode of ozone was red-shifted compared to free O3. The octet of equal intensity in the ozone ν3 region observed when scrambled 16O/18O was used confirmed an asymmetric structure with coordination via a terminal oxygen to give complexes of the form O3•Cl2, O3•Br2, and O3•BrCl. The latter was confirmed from the 66 Page 66 of 213

photolysis which only gave ClBrO. An asymmetric O3•Br2 complex had been identified previously [588]. The formation of a 1:1 complex between O3 and ICl has been reported, most probably with electron donation from a terminal oxygen of ozone into a σ* orbital on the I of ICN, i.e.

ip t

O3•ICN or OOO•ICN, with the structure confirmed by 18O enrichment experiments [589]. The reaction of ozone and ICN or BrCN led to the observation of an O3•ICN rather than

cr

ICN•O3 complex. Coordination was claimed via the central oxygen of the ozone, but although scrambled ozone was used there was no discussion of this with respect to the

us

bonding configuration of the O3•ICN complex. There was no evidence for complex formation

1.26.7.2.

an

between O3 and BrCN [590].

Ozone organohalogen complexes

M

Ozone and F3CI form a molecular complex where the ozone ν3 mode was shifted down by 7 cm−1, and the ν1 mode up by 2 cm−1 [548]. The presence of an octet of absorptions in the O3 ν3 region in the scrambled

16,18

O3 experiments indicated that the complex was asymmetric

d

with respect to the interchange of the two terminal oxygen atoms in the O3 sub-unit, and that

Ac ce pt e

therefore it was bound via a terminal O. Ozone and MeI also form a molecular complex [591], and although no isotopic data is discussed the structure was believed to be similar to that proposed for F3CI complex [548]. Clark has carried out a series of experiments investigating the reaction of ozone with a selection of organohalogen reagents [592-597]. Codeposition of ozone with iodoethane [592], 2-iodopropane, pentafluoroiodoethane, 1,1,1trifluoroiodoethane,

1,1,2,2-tetrafluoroiodoethane,

1,1,1,2-tetrafluoriodoethane

[594],

chloroiodomethane [595], led to the formation of 1:1 complexes which were very photolabile using near infrared radiation. 16O/18O enrichment of the ozone resulted in the observation of characteristic six line patterns for these fluoroalkane complexes, clearly indicating coordination via the central oxygen. This is in variance to the coordination via the terminal O proposed by Andrews for iodomethane, but little evidence was presented for this other than the observation of small shifts from the parent ozone bands [591] and a similarity to the F3CI complex [548]. For bromochloromethane there was no evidence for complex formation, and the photolysis conditions required to form new compounds were much more forcing [595]. No bands that could be assigned to a complex were observed for ozone with 67 Page 67 of 213

tribromofluoromethane [593], dibromochloromethane and bromodichloromethane in solid argon [596]. Analogous experiments involving ozone and 1,2-dibromoethene or 1,2dichloroethene also indicated the formation of photolabile complexes where the double bond was acting as an electron donor, presumably to the central oxygen of ozone [597].

ip t

1.27. Me2O, Me2S and Me2Se complexes

These have been considered in other sections and the symmetric C-O-C, C-S-C and C-Se-C

cr

stretching mode data are presented in Table 16. There are also reports of hydrogen bonded complexes between Me2S and Me2S2 and nitrous [598] and nitric acid [599].

us

1.28. SO2 complexes

SO2 can formally act as both a Lewis base with coordination via O as in the SO2•BF3

an

complex [178, 179] or in hydrogen bonding [553, 600, 601]; or as a relatively weak Lewis acid where the S atom can act as a π* acceptor. In the H2O2•SO2 complex there are examples

M

of both Lewis acidity and hydrogen bonding [536]. The complexes where SO2 is acting as a Lewis acid are those considered in this section, Lewis base behaviour being considered elsewhere, together with some reference to hydrogen bonding complexes [600]. SO2

Ac ce pt e

d

complexes (with CH3I) are one of a few which have been investigated in p-H2matrices [15]. A recent computational study has reviewed some of the SO2 complexes both in the gas phase and in matrices [602]. This concluded that the complexes between SO2 and HCl or HF were hydrogen bonded, but the complexes with H2O, H2S, NH3 and PH3 were characterised by an interaction between the lone pair orbital on the electron donor (e.g. O, S, N) and the π* orbital on the SO2 so that the dative bond is orthogonal to the SO2 plane [602]. A representation of the structure of the H3N•SO2 complex is given in Figure 19. The binding energies were: −17.1 kJ mol−1 for H3N•SO2; −8.9 kJ mol−1 for H2O•SO2; −3.0 kJ mol−1 for H3P•SO2; and −1.8 kJ mol−1 for H2S•SO2 and these are shown graphically in Figure 20. The binding energy for the SO2•BF3 complex is reported as −9.5 kJ mol−1 [178]. The structure of the SO2 dimer in matrices has been a source of uncertainty for a number of years. The initial work of Nord in N2 matrices indicated a dimer structure of non-equivalent sub-units with overall Cs symmetry [603]. This was prior to its observation in the gas phase via microwave [604, 605] and cavity ringdown spectroscopy [606]. Although the N2 matrix 68 Page 68 of 213

observations were confirmed by later studies [607, 608], in an argon matrix there appeared to be a strengthening of the interaction between the SO2 units [609] with overall Ci symmetry. This change was put down to a stronger anisotropic solute-matrix potential distorting the highly symmetric dimers in the molecular matrix (N2), whereas the rare gas matrices generally preserve the gas phase structures [609]. However, more recent studies using Kr and Xe matrices are consistent with a Cs dimer structure [610] as found in N2 matrices [603] and

ip t

the gas phase [604-606]. This work [610] also assigned spectral features in Ar to the same Cs geometry as the other hosts. There was no evidence for a Ci structure, and the peaks originally

cr

assigned to this [609] were assigned to SO2 clusters in meta-stable (hcp) sites [610]. Calculations reveal that the lowest energy region is relatively flat, and that dispersion forces

us

are likely to be responsible for the observed Cs global minimum [611].

an

Nord carried out a series of experiments in 1982 and reported 1:1 complexes between SO2 and NH3, H2O, Cl2, HBr and C2H4 [612], Fredin having previously investigated the C2H4•SO2 complex [613]. The matrix spectrum of the H3N•SO2 complex was very different to that of

M

the solid state indicating as noted above that solid state studies of complex formation can sometimes provide very little information about the free or isolated complexes [612]. The

d

shifts in the νOH modes in the H2O•SO2 complex were large compared to other non-hydrogen

Ac ce pt e

bonded complexes, but indicated that the water was acting as a lone pair donor towards the SO2. For the Cl2•SO2 and C2H4•SO2 complexes the ν1 mode of SO2 was red-shifted whereas in the H2O and NH3 complexes it was blue-shifted. It was not observed in BrH•SO2, but the other modes clearly indicated this contained a hydrogen bonding interaction. Ault had previously reviewed some of his work on SO2 complex formation in matrices in 1988 [10]. However, as there is considerable work published both before and after this it is worth considering the data as a whole and this is collected together in Table 17.

1.28.1.1. SO2 complexes with group 15 donor atoms The Me3N•SO2 complex is one of the best characterised charge transfer (donor acceptor) complexes. The solid state crystal structure reveals a N-S distance of 2.06 Å with the SO2 unit lying normal to the N-S bond [614], and there are also many computational reports [567, 615-617]. The crystal structure of Me2HN•SO2 has also been reported with N-S bond lengths of 2.015(3) and 1.991(3) for the two crystallographically independent adducts [618]. The gas 69 Page 69 of 213

phase microwave spectra revealed N-S bond lengths of 2.26(3) Å for Me3N•SO2 [619] and 2.34(3) Å for Me2HN•SO2 [620]. Therefore, as for the BF3 nitrile complexes, the length of the donor-acceptor bond is very dependent on the medium in which it is determined. The selfconsistent reaction field method (based on Onsager’s reaction field theory) has been used to study the effect of polar media on the partially bound Me3N•SO2and Me3N•SO3 complexes

ip t

[621]. Ault built on the previous work [612, 622] in his study of the H3N•SO2 and related amine

cr

complexes [623]. Both argon and nitrogen matrices were used, and whilst multiplet structure was observed in argon, single bands were observed in nitrogen matrices. Therefore, the data

us

given in Table 17 for these complexes is that from both argon and nitrogen matrices. 1:1 SO2 complexes were observed for NH3, NH2Me, NHMe2 and NMe3 in both merged and twin-jet

an

mode, except for NH2Me where a dehydration reaction resulted in the formation of MeNSO and H2O in single-jet mode. These observations were taken to imply evidence for the formation of 1:1 complexes in the gas phase. The only evidence for the 2:1 complex was

M

(H3N)2•SO2. The vibrational data from a subsequent computational study [616] for NH3, NH2Me, NHMe2 and NMe3 1:1 complexes were in good agreement with the experimental

d

data. The previous calculations [567] had indicated that the plane of the SO2 molecule was

Ac ce pt e

approximately perpendicular to the S-N bond, which was co-linear with the amine C3 axis as found in the solid state [614]. The computational study also indicated the thermodynamic stability of all of the 1:2 complexes [616]. Subsequent experiments using a pulsed molecular beam Fourier transform microwave spectrometer gave N-S bond lengths of 2.34(3) Å for Me2HN•SO2 [620] and 2.26(3) Å for Me3N•SO2 [619] which shorten to 2.015/1.991 and 2.06 Å, respectively, in the solid state [614, 618]. A representation of the structure is given in Figure 19. Crystal structures also gave N-S bond lengths of 2.006 and 2.017 Å for the Me2HN•SO2 complex and 2.036 and 2.207 Å for the C5H11N•SO2 complex [624]. Building on their earlier work involving nitrile•BF3 complexes (section 1.4.1 [151, 152]) and the ammonia and amine complexes [612, 614, 618-623] where it had been shown that the medium affected the extent of interaction, Phillips et al. studied nitrile•SO2 complex formation [625] as it was anticipated that there would be more pronounced affects due to the weaker nitrile base, compared to the previous ammonia studies. There was no evidence of product formation between MeCN and SO2 down to 193 K and product formation was only observed in thin films deposited at 80 K, which was 20 K lower than that required for the 70 Page 70 of 213

nitrile BF3 complexes [625]. The MeCN•SO2 complex bands were similar in argon and nitrogen matrices as well as the thin films of MeCN•SO2 [625]. The calculations predicted a N-S distance of 3.02 Å. Therefore, they concluded that MeCN•SO2 was a rather weak donoracceptor complex, with little or no propensity for medium induced structural change [625]. Microwave spectroscopy showed that the C5H5N•SO2 complex has Cs symmetry with the

ip t

pyridine C2 axis being approximately perpendicular to the SO2 plane with a N-S distance of

cr

2.61(3) Å, and total dipole moment of 4.552(5) D [626].

When TTP (TTP = 1, 4, 8, 11-tetrathiacyclotetradecane) and SO2 were codeposited in argon

us

matrices, the νSO band at 1322 cm−1 was shifted by some 28 cm−1 from the monomer which

an

was interpreted as being due to the formation of a 1:1 complex [627].

1.28.1.2. SO2 complexes with group 16 donor atoms

M

The SO2 monomer was found to be photochemically inactive in oxygen matrices, with no evidence for [SO4]* in an oxygen matrix [628]. However, the dimer in an oxygen matrix was readily photolysed to give SO3. An investigation of the oxidation of SO2 to SO3 using ozone D O atoms or by oxygen transfer from within the O3•SO2 complex [629, 630].

Ac ce pt e

1

d

and O atoms in argon matrices showed that this could occur via either reaction of the SO2 and 16

O/18O

isotopic substitution in the O3•SO2 complex indicated the inequivalence of the two oxygen atoms in SO2. Whilst a six line pattern with the same intensity distribution as ozone was observed, the broadness of the mixed isotope features indicated that it was more likely to be indicative of an asymmetric ozone unit. Small shifts in the vibrational modes indicated a weak interaction. The relative intensity of the ν3 and ν1 modes of the SO2 sub-unit were very similar to those in the SO2 monomer and this was taken as indicating negligible charge transfer between the SO2 and O3, and that the interaction was mainly due to dipolequadrupole or dipole-induced-dipole interactions [629, 630]. The H2O•SO2 1:1 complex is important in modelling atmospheric reactions and was first studied in matrices by Nord in 1982 [612], and subsequently in 1988 by Schriver et al. [631]. They concluded that the oxygen atom acts as an electron donor in a charge transfer complex, rather than being hydrogen bonded. The 1:2 complex involves a second SO2 unit hydrogen bonding to the 1:1 complex, and the 2:1 complex is best thought of as a water dimer involved 71 Page 71 of 213

in a cyclic structure with one hydrogen bond and one charge transfer interaction to the SO2 [631]. Gas phase microwave data [632] indicated a double-decker (stacked/eclipsed) structure with a S-O distance of 2.824(2) Å but with the two planes of the sub-units tilted ca. 45° from parallel, although subsequent calculations indicate this might be nearer 14° [633]. Hirabayashi reported the first spectroscopic observation of the 3:1 complex in addition to 1:1 and 2:1 complexes in argon matrices and calculated the binding energy of the H2O•SO2

ip t

complex to be −13.8 kJ mol−1 in a double-decker style structure [634]. In the 2:1 and 3:1 complexes, there is still a H2O•SO2 interaction, but the additional water molecules are held in

cr

place by hydrogen bonds [634]. H2O•SO2 and H2O•SO3 matrix isolated complexes can also both be formed from photo-oxidation of H2S in solid oxygen matrices at 15 K [534]. More

us

recent calculations gave CCSD(T) dissociation energies of 10.9 (aug-cc-pVDZ) and 12.6 (aug-cc-pVTZ) kJ mol−1 for the double decker structure and much smaller dissociation

an

energies (4.4 and 5.3 kJ mol−1) for a hydrogen bonded structure involving the terminal O of SO2 [635]. Other computational reports give values of −9.2 kJ mol−1 [531] and −11.6 kJ

M

mol−1 [633].

In the HOO•SO2 complex (HOO was prepared via microwave discharge of H2 and trapped

d

with O2), the HOO radical acted as a lone pair donor to the sulfur atom of SO2 [636]. A

Ac ce pt e

significant blue-shift in the HOO bending mode indicated an interaction between the H and one of the oxygen atoms of SO2 to give a structure similar to that observed for the H2O2 complex [536].

Argon matrices containing SO2, O2 and OCS, and SO2 and OCS were photolysed and studied using excitation and thermoluminescence as well as IR spectroscopy. The fluorescence behaviour was explained in terms of an energetically coupled S•SO2 complex, and a weak red shift (8 cm−1) of the SO2 modules in the IR spectrum was also indicative of a weak S•SO2 complex [637].

Li used pre-mixed SO2 and S atom donor bases (Me2S, Et2S, or MeSEt) with a single-jet approach to study the weak complexes formed. Two different complexes in both Ar and N2 matrices were observed for the interaction of SO2 with Me2S, Et2S, or MeSEt. One of these was a 1:1 complex between the thioether and SO2 monomer, e.g. Me2S•SO2, and the other was a 1:1 complex between the thioether and the SO2 dimer e.g. Me2S•(SO2)2 [638]. In the case of thiols, 1:1 complexes were observed between SO2 and EtSH, 2-C3H7SH and t72 Page 72 of 213

Me3CSH, with the interaction most likely via the formation of a disulfide linkage rather than a hydrogen bond, based on the vibrational shifts observed [639]. A subsequent report investigated the interaction of SO2 with the bases, HSCH2CH2SH, MeSCH2SH and MeSSMe, which are isomers all containing two S donor atoms. For HSCH2CH2SH, and MeSSMe only, HSCH2CH2(H)S•SO2 and MeS(Me)S•SO2 were observed but for MeSCH2SH both (Me)SCH2(H)S•SO2 and HSCH2(Me)S•SO2 were observed [640, 641]. For MeSH a 1:1

ip t

complex was observed and supporting calculations yielded a S-S separation of 3.45 Å [642]. This is significantly larger than the sum of the covalent radii, but less than the van der Waals

cr

radii indicating a weak interaction.

us

The salt/molecule reaction technique in conjunction with matrix isolation has been used to synthesize Cs[SO2F]-, Cs[SOF3]- and Cs[SO2F3]- from CsF and SO2, SOF2 and SO2F2,

an

respectively [643].

M

1.28.1.3. SO2 complexes with unsaturated organic ligands Matrix isolated C4H6•SO2 [644] and C2H4•SO2 complexes [612] complexes were indicated by the presence of broadened monomer SO2 bands, whereas for C6H6•SO2 and C6H5Me•SO2

d

there were distinct bands red-shifted from those of the parent SO2 indicating formation of 1:1

Ac ce pt e

complexes [644]. The subsequent gas phase microwave spectrum indicated that the C2H4•SO2 complex had a stacked structure with Cs symmetry with 3.504(1) Å separating the centres of mass of the two sub-units [645]. The gas phase microwave spectrum of the C6H6•SO2 complex indicated that the SO2 was parallel with the benzene ring with the C6H6 π electron density interacting with the π* orbitals of SO2, whereas in the C5H5N•SO2 complex the pyridine rotated by 70° so that it became nearly perpendicular to that of SO2 [646, 647].

1.28.2.

Summary of SO2 complex behaviour

Figure 21 collects together all the data for which there are both matrix vibrational and proton affinity data in Table 17. From this there appear to be three different correlations between the proton affinity and the position of the ν3mode in SO2. There is a good correlation between the proton affinity and asymmetric stretching mode of SO2 (shown by the black line) for most of the bases studied. However, for the amines, there is a much stronger correlation as shown by the blue line. The thiols and the thioethers could have a similar correlation to that observed

73 Page 73 of 213

for the amines, or alternatively, the thioethers could stand alone. There is a similar pattern observed for the ν1 mode as well, but as the shifts are smaller, there is a greater scatter.

1.29. SO3 complexes SO3 is a strong Lewis acid and a potent oxidizer, it also has a pivotal role in atmospheric

ip t

chemistry. The spectrum of SO3 in a variety of rare gas matrices has been reported [648]. In argon the ν3

cr

asymmetric stretching mode was assigned to a peak at 1385.9 cm−1. In a Ne matrix with a small proportion of Xe the spectrum was almost identical to the pure Ne spectra, except for

us

some evidence of the ν2 out of plane bending mode of SO3 at 480.3 cm−1 being shifted from 497.6 cm−1 in pure Ne (no such were effects observed on the other modes) [648]. This indicated that the SO3 was perturbed by a Xe atom located along the C3 axis. This value of

an

the ν2 mode was almost identical to that of SO3 in neat Xe. With O2 doped Ne matrices, a similar shift to 490.1 cm−1 was observed, together with activation of the O-O stretching

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mode. SO3 can also be prepared in matrices from the oxidation of SO2 by 1D oxygen atoms following photolysis of SO2 and O3containing argon matrices [630].

d

Matrix experiments and calculations indicated that SO3 dimers may either have a molecular

Ac ce pt e

“Al2Cl6” D2h structure or a more loosely bound complex type structure with the SO3 units stacking with S...O interactions [649]. The calculations indicated that the binding enthalpy for the complex was −6.3 kJ mol−1 using B3LYP and −20.5 with MP2; for the D2h dimer the values were +8.0 and −20.3 kJ mol−1, respectively. The difference in the DFT and MP2 values for the binding energies of the complex were put down to the magnitude of the dispersion forces within the complex. The complex is disfavoured by entropy and only becomes exogenic below 110 K, even so it was commented that the formation of the dimeric complex from the monomer was close to being thermo neutral. The calculated spectra indicated the presence of the loosely bound dimer (and trimer) in the experimental data and there was no evidence for the molecular S2O6 (or S3O9) species [649]. The argon matrix values for the ν3 mode in SO3 were in good agreement with the previous ones [648], but it was suggested that in an intermediate study with four bands reported for the ν3 mode in SO3 [629], that some of these were due to the presence of N2 in the argon matrix [649].

74 Page 74 of 213

1.29.1. SO3complexes with Group 15 donors Although SO3 and NMe3 were known to form a stable complex including an X-ray crystal structure which indicated that it exists as +H3N•SO3− zwitterions in the solid state [650]. However, there was limited spectroscopic data available for the gas phase complexes between SO3 and ammonia, methylamines and pyridine prior to Ault’s twin-jet matrix studies in 1986 where 1:1 complexes were characterised by shifts in the vibrational modes compared

ip t

to their parent compounds [70]. All of the complexes displayed a red-shifted SO3 asymmetric stretching mode from 44 cm−1 for H3N•SO3, 50 cm−1 for MeH2N N•SO3, 52 cm−1 for Me2HN N•SO3, and 68 cm−1 for Me3N N•SO3, in the case of the C5H5N•SO3 complex the shift was 60

cr

cm−1. In addition the SO3 symmetric stretching mode at ca. 1070 cm−1, was activated by

us

complex formation. The gas phase value for this infrared inactive mode was given as 1065 cm−1. The ν2 mode of NH3 in H3N•SO3 was blue-shifted by over 300 cm−1. All of these observations indicated a very substantial interaction between the amines and the sulfur

an

trioxide, and were described as being characteristic of the SO3 acting as a π* acceptor. Although hydrogen bonding could be invoked to explain the structure, the presence of a νS-N

M

mode in Me3N•SO3 clearly indicated a Lewis acid/base complex and the formation of a S-N coordinate bond. It was noted that although the degenerate asymmetric mode should be split in the substituted amines, this was not observed. However, in the study of a range of oxygen

d

bases, this splitting was observed, confirming the formation of the 1:1 complexes [651]. See

Ac ce pt e

Figure 22 for a representation of the structure.

Subsequent work on phosphorus and sulfur bases reported the first characterisation of 1:1 complexes between SO3 and PH3, PMe3, PF3, H2S, and Me2S [652]. The red-shift in the asymmetric stretching mode of SO3 ranged from only 8 cm−1 for F3P•SO3, through 75 cm−1 for H3P•SO3 to 117 cm−1 for Me3P•SO3 as expected for the variation in Lewis basicity. In particular, the shift for PH3 is much greater than that for NH3 indicating a stronger interaction. For the H3P•SO3 and Me3P•SO3 complexes a weak band near to 1060 cm−1 was assigned to the symmetric SO3 stretch, activated upon complex formation [652]. Following on from the matrix work there have been more recent reports describing kinetic results as well as the use of microwave spectroscopy to structurally characterise some SO3 group 15 complexes.

75 Page 75 of 213

The rate of reaction between SO3 and NH3 at 298 K in flow tubes has been determined as 6.9 (±1.5) x 10−11 cm3 molecule−1 s−1 at He pressures of 1 – 2 torr [653], that of SO3 and MeH2N was 1.1 (±0.2) x 10−11 cm3 molecule−1 s−1 [653] and that between SO3 and MeOH was 1.17 (±0.16) x 10−13 cm3 molecule−1 s−1 [654]. Leopold has used microwave spectroscopy and Stark effect measurements to study a range of

ip t

gas phase SO3 complexes, many of which were initially characterised using matrix techniques. For H3N•SO3 the structure was found to be C3v with a N-S bond length of

cr

1.957(23) Å [655] and the nitrogen quadrupole coupling constant implied that ca. 0.36 e- was transferred when the complex was formed [655]. The N-S bond length is 0.186(23) Å longer

us

than in the solid where the structure consists of +H3N•SO3− zwitterions [650]. Stark effect measurements yielded a dipole moment of 6.204(11) D for H3N•SO3 [656]. The N-S bond

an

length in Me3N•SO3 was 1.912(2) Å with a dipole moment of 7.1110(69) D which is over four times that of the individual components [657]. Analysis of the hyperfine structure indicated that about 0.6 e- was transferred from the nitrogen to SO3 during complex

M

formation. The calculated binding energies for H3N•SO3 and Me3N•SO3 were −81.9 and −151 kJ mol−1, respectively [658]. The difference in the solid state and gas phase N-S

d

distances for Me3N•SO3 has been modelled using self-consistent reaction field methods, as

Ac ce pt e

for the Me3N•SO2 complex [621]. The N-S distance in N2•SO3 was 2.934(12) Å and the C-S distance in OC•SO3 was 2.854(12) Å [659]. In the OC•SO3•Ar complex the C-S distance increased to 2.849(4) Å and the Ar-S distance (on the opposite side) was 3.411(11) Å, which is 0.061 Å longer than that of Ar•SO3. The dipole moment of N2•SO3 was 0.46(1) D [660]. The Stark effect derived dipole moments were 0.8488(13) D for OC•SO3 and 0.602(15) D for OC•SO3•Ar [659]. The calculated binding energies were: Ar•SO3, −4.51 kJ mol−1; N2•SO3, −10.9 kJ mol−1; OC•SO3, −16.4 kJ mol−1; Ar•SO3•CO, −20.5 kJ mol−1 [659]. In HCN•SO3 the N-S distance was 2.577(6) Å and in MeCN•SO3 the N-S distance was 2.466(16) Å [661]. Stark measurements gave dipole moments of 4.4172(31) D for HC14N•SO3, and 6.065(18) D for MeC14N•SO3 [143]. The calculated binding energies for HCN•SO3, MeCN•SO3 and HCCCN•SO3 were −30.5, −37.6 and −30.9 kJ mol−1, respectively [658]. In C5H5N•SO3 the N-S distance was 1.915(1) Å, and this short bond length indicated that the formation of the dative bond was nearly complete, and the quadrupole coupling constants indicated a transfer of ca. 0.54 e- from pyridine to SO3 [658].The binding energy was calculated to be −107 kJ mol−1 [658]. In addition to the HCN•SO3 complex, microwave spectroscopy has also been used to study the HCN•SO3•Ar and HCN•SO3•CO complexes, where the Ar or CO are 76 Page 76 of 213

weakly bound on the opposite side of SO3 as the HCN [662]. The N-S distance in HCN•SO3•CO increased to 2.6563(14) Å and the S-C increased to 3.0109(48) Å [662] compared to 2.577(6) Å [661] and 2.854(12) Å [659] in HCN•SO3 and OC•SO3, respectively. Leopold has used this microwave data to show that there is a very good correlation between N-S distances, the proton affinity and the vertical ionisation energies for the nitrogen lone

ip t

pair of electrons in the free base [663].

cr

1.29.2. SO3 complexes with Group 16 donors In an initial study there was evidence for a H2O•SO3 complex, but the data were insufficient other than to assume that the oxygen of the water lies on the C3axis and donates its electron

us

density to the SO3 Lewis acid [648]. It also appeared that the H2O•SO3 complex does not rearrange to H2SO4 in Ne at 5 K. Higher concentrations of water led to “solvated” sulfuric acid [648]. There was also evidence for H2O•SO3 complex following photolysis of an

an

SO2/H2O/O3 mixture in a nitrogen matrix, and the observed isotopic shifts confirmed coordination via an O...S interaction [629]. Likewise, the photo-oxidation of H2S in solid O2

M

also yielded the H2O•SO3 complex [534]. Subsequent experimental work and calculations showed that in addition to the H2O•SO3 1:1 complex there was evidence for 2:1 and even

d

higher complexes, but there was no evidence for a 1:2 complex. The H2O•SO3 complex has a double decker eclipsed structure with Cs symmetry with the oxygen of the water lying above

Ac ce pt e

the S and the two water O-H bonds eclipsed with two of the SO3 S-O bonds [649]. The energies of reaction of complex formation for H2O•SO3 were −32.7 (B3LYP) or −39.2 (MP2) kJ mol−1 [649], and these are in good agreement with a value of −30.9 kJ mol−1 published the previous year [531].

In addition to group 15 complexes [70, 652], Ault also studied a range of group 16 bases and observed the formation of 1:1 complexes with SO3 in nitrogen matrices [651, 652]. The redshifts of the asymmetric stretching mode in SO3 complexes with O donors (Me2O, MeOH, Me2CO, H2CO, (CH2)2O, (CH)4O, (CH2)4O) were smaller than those of the N donor complexes, indicating a weaker interaction, but in contrast to the N donor complexes there was evidence of the splitting of this mode confirming the formation of the 1:1 complex via interaction of the O and S [651]. However, the SO3 symmetric stretching mode was not activated to any detectable degree. As well as O donor complexes, 1:1 H2S•SO3 and Me2S•SO3 complexes have also been reported [652]. The red-shift in the asymmetric stretching mode in SO3 was greater in Me2S•SO3 than H2S•SO3 and was also accompanied by 77 Page 77 of 213

activation of the SO3 symmetric stretch at 1042 cm−1. As for the phosphine complexes, this was red-shifted [652], in contrast to the blue shift observed for the amines but was blueshifted for the amines [70]. The reaction of SO3 and O atoms in argon matrix formed SO4, but there was no sign of a

ip t

precursor O3•SO3 complex formed with the O3 [664]. Due to the atmospheric importance of SO3, H2O and H2SO4, Loewenschuss et al. have

cr

published a series of experimental papers on the complexes of SO3 with H2O [649] CO2 [665], CO [666] and N2 [667]. The H2O complex has been considered above. The

us

calculations for the OCO•SO3 complex indicate that lowest energy structure has one of the oxygen atoms of the CO2 is bound to the sulphur in SO3 at a distance of 2.94 Å, with the

an

carbon lying 4.01 Å above one of the oxygens in SO3 [665]. The bonding energy was calculated as –6.2 kJ mol−1 with a bonding enthalpy of −2.3 kJ mol−1. For the potential N2•SO3 complex, the small shifts were interpreted in terms of cage effects rather than well a

M

defined complex [667]. In addition they also studied H2SO4 [668, 669] and its complexes with H2O [670-672], NH3 [673], N2 and NO [667], CO [666], HCl [674] CO2 and CO2SO3

d

complexes [665]; and the H2O complexes of dimethylsulfone [675, 676] and methanesulfonic

Ac ce pt e

acid [677], the majority of which involve hydrogen bonding. There are also supporting calculations in these papers, as well as others [678-680]

1.29.3. Summary of SO3 complexes The experimental matrix isolation data for SO3 complexes is collected together in Table 18. For the H2O•SO3 complex in Ne [648], Ar [649] and O2 [534] matrices the asymmetric SO3 stretching mode was blue-shifted, whereas in all the other complexes it was red-shifted. Ault pointed out that these red-shifts were not consistent with the proton affinity of the base, but were much better described by the HSAB approach, and that in the case of PF3 it was the whole molecule that needed considering rather than just the donor atom [652]. There is a good correlation between the proton affinity and asymmetric stretching mode in the SO3 sub-unit for both N-donor and O-donor bases, and these data have been used to generate the fit in Figure 23. Whilst the values for the F3P•SO3 and H2S•SO3complexes 78 Page 78 of 213

follow this trend, those for H3P•SO3, Me2S•SO3 and especially Me3P•SO3 complexes lie a long way from this relationship and do not really correlate with the proton affinities. As has been pointed out previously [652], the proton affinity of the base is not the only parameter that should be invoked, as the hardness and softness (HSAB) of both the acid and base need to be considered. In this case, there is a reasonable relationship between all the third row donor atoms as shown in Figure 23, but it would be even stronger if it only included the data

ip t

for the H3P•SO3, Me2S•SO3 and Me3P•SO3 complexes.

cr

Recent calculations have indicated that there are significant differences between the gas and condensed phase properties of both MenH3-nN•SO3 (n = 0-3) and MemH2-nO•SO3 (m = 0-2)

us

due to the ability of the condensed phase medium to support a larger charge separation between the donor and acceptor. There is also a strong correlation between the extent of

M

oxygen-containing complexes of SO3 [681].

an

charge transfer and binding energy in gas and condensed phases for both the nitrogen and

1.30. Complexes of SF4, SOF2 and SO2F2

1:1 complexes were observed between SF4 and NH3, pyridine and acetone; SOF2 and NH3,

d

MeNH2, and pyridine; and SO2F2 and NH3. Although complexes of SF4 and MeNH2 or

Ac ce pt e

Me3N, and SOF2 and Me3N with were expected, they could not be identified due to masking by parent bands [682]. Red shifts were observed in the νS-F and νS-O modes for all of the complexes, together with a significant blue shift in the NH3 ν2 mode for the ammonia complexes. The shift for H3N•SOF2 was slightly smaller than that of H3N•SF4 as expected from the differences in Lewis acidity. The perturbation of the acid sub-unit modes was greater for SF4 than SOF2 indicating greater acceptor properties of SF4. Whilst a 1:1 complex was observed for SO2F2 and NH3, this was only weakly bound. No bands attributable to a Me3N•SF4 or Me3N•SOF2 complex were observed [682]. These data are collected in Tables 19 and 20.

1.31. OCS complexes OCS O3 reactions have atmospheric implications as OCS is thought to be the dominant nonvolcanic source of stratospheric sulfuric acid [683]. Although the formation of an SCO•O3 1:1 complex was reported, even with isotopic data it was not possible to distinguish between 79 Page 79 of 213

a symmetric complex where the central oxygen atom of O3 is linked with SCO, or an asymmetric complex, where one of the ozone terminal oxygen atoms forms a dimer with SCO. Photolysis of the O3 yielded 1D oxygen atoms, leading to CO2 and S, with the S atoms being further oxidized to SO2, SO3 and SO4 by O2 and diffusing O atoms. [683]. As noted in section 1.20 there was no complex formation between PF3 or PCl3 and OCS but

ip t

photolysis yielded CO and SPF3 or SPCl3 [498]. Although CO was observed no sulfur atom transfer products were observed for NF3, PH3, AsF3 or HCl [498]. The implication is that

cr

both 1D and 3P S atoms result from photolysis of matrix isolated OCS using 230-260 nm radiation. Whether sulfur atoms add to a given substrate X depends on the efficiency with

us

which the initial “hot” adduct [XS]* is quenched, as well as the photolability of the product

an

XS [498].

OCS complexes with halogens and interhalogens are covered in section 1.36.10.

M

Calculations indicated that the H2O•CS2 and H2O•OCS complexes would be better described as H2O•SCS and H2O•SCO, as the interaction was between the water oxygen atom and the

d

sulfur in each case. The binding energies were calculated as −4.5 kJ mol−1 and −5.4 kJ mol−1

Ac ce pt e

for H2O•SCS and H2O•SCO, respectively [531]. The structure of the HC≡CH•OCS complex has been determined by gas phase microwave spectroscopy [684] and also been the subject of computational work [685]. The experimental data [684] indicated that the two sub-units were nearly parallel with the triple bond above the central carbon, whilst the computational work [685] explored the possibility of T-shaped geometries as well as the parallel one, and unsurprisingly found that the parallel geometry was the lowest in energy, as was the depth of its potential well.

1.32. Complexes of S2 and Se2

The S2•O2 and Se2•O2 complexes formed from the reaction of S2 or Se2 and O2, have cyclic parallelogram symmetric structures, and are chemically intermediate between the unstable O4 and stable S4 [686, 687], Se4 [688] and Te4 [689] molecules. The O-O fundamentals were red-shifted whilst the S-S (and Se-Se) fundamentals were blue shifted, implying a transfer of electron density from the π* orbitals on the S2 or (Se2) into a similar orbital on O2 [690, 691]. Photolysis of these complexes resulted in a variety of products, including Se2O2 which adopts 80 Page 80 of 213

a three membered ring structure with an exocyclic Se=O unit [689]. Analogous tellurium oxygen complexes were not observed when laser ablated Te atoms were the source in the presence of oxygen [692]. Calculations indicated that the energetically favourable “face to face” binding of the diatomic chalcogenide species leads to delocalisation of the electron

ip t

density in the high lying occupied π* anti-bonding orbitals [693].

1.33. Oxygen complexes

cr

A significant number of investigations employing O2 are concerned with its reaction with

us

atomic species which lie outside the scope of this review.

Complexes between O2 and N2O [694], O2 and NO [695, 696] and O2 and cis-N2O4 [696] in matrices have been reported. A detailed matrix isolation experimental and computational

an

study of the bonding of O2 to the FSO3 radical has shown that the O2 is end on bonded to one of the terminal O atoms of the FSO3 sub-molecule [697]. A cyclic structure involving side on

M

bonded O2 had similar thermochemical properties, and it was the use of detailed isotopic

Ac ce pt e

Group 17

d

labelling that allowed for the identification of the end-on complex.

1.34. Introduction

Base HF complexes were reviewed by Andrews [549], and Legon has reviewed the “halogen” bond in a number of gas phase complexes studied by rotational spectroscopy [698700]. Often the matrix work would predate the microwave spectroscopic study, but not in all cases.

1.35. Hydrogen halide complexes

As these are essentially all hydrogen bonded [701], the illustrative examples have been dealt with elsewhere.

1.36. Halogen and interhalogen complexes The halogens and interhalogens can act as Lewis acid σ* acceptors, and form complexes that are often described as charge transfer. Of these ClF is attractive to the matrix isolation community as it has readily observable IR active stretching modes at 768 and 762 cm−1 (due to 35Cl and 37Cl isotopomers) that can be used to probe the Lewis base-acid behaviour [69]. 81 Page 81 of 213

For the halogens themselves, the IR inactive stretching mode is activated upon complexation, but depending on the extent of the interaction it is not always observable. New absorption features close to the parent modes indicate complex formation rather than elimination, as does the activation of the IR inactive stretching modes in F2 and Cl2. The halogen vibrational mode data is collected together in Table 21. ClF complexes with O and S and Se donor Lewis bases.

ip t

1.36.1.

Twin-jet deposition of ClF and a range of oxygen bases diluted in argon at 10 K revealed new vibrational modes shifted from those of the parent. (All the ClF experiments were carried out

cr

using a twin-jet approach due to reactivity of ClF.) As the shifts in the base vibrational modes

us

were those involving oxygen, this indicated interaction of O with ClF, rather than hydrogen bonding, and probably with the Cl part of the molecule. The following 1:1 complexes were characterised: H2CO•ClF; H3C(H)O•ClF; (CH2)2O•ClF; Me2O•ClF; Me2CO•ClF; Et2O•ClF

an

and (CH)4O•ClF [69]. The shift in the νClF mode correlates well with the proton affinity of the base. Similar 1:1 complexes were also observed between ClF or Cl2 and H2S, MeSH and

M

Me2S [702]. The νClF and νClCl modes were red-shifted, increasing with proton affinity of the base. This work was in agreement with earlier work on H2S•Cl2 but in substantial

Ac ce pt e

complex [703].

d

disagreement with the previous assignment of the νClCl mode at 525 cm−1 for the Me2S•Cl2

The reaction of ClF with H2Se and Me2Se bases in solid argon led to the straight-forward assignment of the spectral features to a 1:1 complexes H2Se•ClF. In the case of Me2Se, the data was not sufficient to discriminate between a Me2Se•ClF complex, and a four coordinate addition product of the probably of the form Me2SeClF [704]. Subsequent microwave experiments have yielded 2.523(7) Å [705] and 2.437(2) [706] for OCl distances in (CH2)2O•ClF and H2CO•ClF, respectively with the base approximately perpendicular to the Cl-F bond. The S-Cl distances in H2S•ClF and H2S•Cl2 were 2.857(3) [707] and 3.249(2) Å [708], again with the H2S approximately normal to the Cl-F. Recent calculations give interatomic distances of 2.875 and 3.211 Å, respectively for these complexes, as well a distances for 16 other interhalogen or halogen complexes of H2S, Me2S and H2CS [709].

82 Page 82 of 213

The shifts in the νClF (or νClCl) modes for the O-donors (Table 21 and Fig. 12) were approximately 50 - 100 cm−1, whereas shifts of up to 300 cm−1 were observed for S-donors, despite very similar proton affinities, indicating that HSAB factors may be at work as well.

1.36.2.

Br2, Cl2, I2 and ICl complexes with Me2S and Me2Se

ip t

Me2S•Br2 is a yellow-orange solid that has an appreciable vapour pressure. Matrix isolation experiments involving both the co-condensation of Me2S and Br2 as well as sampling the vapour above the solid complex revealed essentially identical IR spectra characteristic of a

cr

1:1 Me2S•Br2 van der Waals complex with the Br2 end-on bonded to the S of Me2S [710].

us

The gas phase PES experiments did not yield any observable products, presumably requiring a third body for stabilisation [710]. Analogous experiments involving Cl2 yielded a 1:1 complex in the co-condensation reactions as observed previously [702], but in the gas phase

an

this converted to the covalently bound Me2SCl2 product, which went on to decompose to monochlorodimethylsulfide (MeSCH2Cl) and HCl [710]. Similar matrix experiments were

M

carried out on Me2S complexes with I2 and ICl. Identical IR spectra characteristic of 1:1 van der Waals complexes were observed, irrespective of whether the reagents were co-condensed in the matrix, or the vapour above the adduct (prepared by standard synthetic methods) was

d

trapped in the matrix [711]. A purely computational report on the complexes between Me2Se

Ac ce pt e

and Cl2, Br2, I2, by the same group predicted that the Cl2 and Br2 complexes would be present as covalent intermediates in matrix or PES spectra, whilst the I2 complex would be a van der Waals adduct [712].

1.36.3.

ClF complexes with N donor Lewis bases

The reaction of ClF with a selection of nitrogen containing bases in solid argon and nitrogen matrices led to the formation of the following 1:1 complexes: Me3N•ClF, H3N•ClF, CH2CHCN•ClF and MeCN•ClF [713]. Whilst the amines will clearly be electron lone pair donors (Figure 24), the nitriles could conceivably interact with ClF via the N lone pair, or the π system of the C≡N bond. The blue-shifts of the νCN modes were consistent with an axial interaction via the N lone pair. The red-shifts in the νClF modes of 263, 168, 142 and 88 cm−1, for Me3N•ClF, H3N•ClF, CH2CHCN•ClF and MeCN•ClF, respectively, indicate a reduction in the extent of interaction from Me3N to MeCN [713]. Large shifts between Ar and N2 matrices were observed, which were interpreted as suggesting considerable ionic character in the complexes [713]. Subsequently, gas phase microwave spectroscopy has been used to 83 Page 83 of 213

determine the structures of Me3N•ClF [714, 715], H3N•ClF [716, 717] and MeCN•ClF [718]. Me3N•ClF has a C3 axis with either C3 or C3v symmetry and the N-Cl distance of 2.090 Å indicates that the main contributor to the valence bond description is [Me3NCl]+F− and is hence a Mulliken inner type complex [714, 715], consistent with the matrix observations [713]. Both H3N•ClF and MeCN•ClF have C3v symmetry and N-Cl distances of 2.37(1) Å [716, 717] and 2.561(2) Å [718], respectively, with smaller ionic contributions to the bonding

ip t

than in Me3N•ClF. In HCN•ClF the N-Cl distance is 2.639(3) Å [719]. Calculations examining the MenH3-nN•ClF complexes indicated that there is a greater ionic contribution in

cr

the Me3N•ClF complex than the H3N•ClF complex, and that as n increases the ClF distance

1.36.4.

us

increases, the N-Cl distance decreases, and the stabilisation energy increases [720].

F2, Cl2 complexes with NH3

an

As part of a larger study on the hydrogen bonding interactions involving NH3 within matrices, Andrews investigated the reaction of NH3 and F2 with the aim of providing the first

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spectroscopic characterisation of the highly reactive NH2F molecule [721, 722]. On deposition of NH3/Ar and F2/Ar samples at 12 K a weak H3N•F2 complex was observed, characterised by a perturbed NH3 ν2 mode at 966.2 cm−1 (equidistant between the inversion

d

doublet components of NH3). The lack of the inversion doublets indicated that the F2 adduct

Ac ce pt e

prevents inversion of the NH3 submolecule, as also found for the H3N•HF complex [459, 723]. The presence of a weak F-F mode at 781 cm−1 (compared to 892 cm−1 for F2 observed in matrix Raman experiments [724]) indicated axial attachment of the F2 to the nitrogen lone pair as this arrangement would result in higher IR activity and provide a more favourable overlap with the fluorine σ* (2p) MO. When this complex was photolysed a number of FH2N•HF complexes were formed, representing the first spectroscopic evidence for NH2F. Although attempts were made to prepare uncomplexed NH2F using both photolysis during deposition and discharge, there was only marginal evidence for its formation [721]. The assignment of the H3N•F2 complex was confirmed by other, later matrix measurements [725, 726]. A subsequent fast-mixing pulsed-nozzle FT microwave study (which did not reference the matrix work [721]) confirmed a C3v geometry for H3N•F2 with a N-F distance of 2.708(7) Å [727]. EPR spectroscopy of F atom NH3 interactions in argon matrices has yielded data on the radical-molecular complex •NH2-HF [725, 728-730] as well as •NH2 radicals [730]. Due to the computational complexity of dealing with F2, the H3N•F2 complex and those of related amines has kept the computational chemists occupied [731-735]. 84 Page 84 of 213

For H3N•Cl2 a NH3 umbrella mode at 988 cm−1 and the perturbed Cl2 stretching mode at 479 cm−1 indicated a stronger interaction for the H3N•Cl2 complex than for H3N•F2 [736]. After the matrix work, Legon, published a series of papers where pulsed nozzle Fourier transform microwave spectroscopy was used to determine the geometries of the complexes,

ip t

and this confirmed Ault’s description of the structure as having C3v symmetry with the N coordinated to the halogen. The experimental values for the r(N-X) interaction were: H3N•F2,

cr

2.708 Å [727]; H3N•Cl2, 2.73(3) Å [737, 738]; H3N•Br2, 2.72(2) Å [739]; H3N•ClF, 2.37(1) Å [716];. H3N•BrCl, 2.672 Å [740]; H3N•ICl, 2.711(2) Å [741]. The calculated

us

“intermolecular” modes were ca 100 cm−1 for H3N•F2, ca. 140 cm−1 for H3N•Cl2 and ca. 240 cm−1 for H3N•ClF. The calculated stabilisation energies were ca. −10 kJ mol−1 for H3N•F2,

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ca. −20 kJ mol−1 for H3N•Cl2, and ca. −40 kJ mol−1 for H3N•ClF [733]. These indicate an increase in intermolecular interaction from the F2 through the Cl2 to ClF complex with NH3 which correlates very well with the ν2 mode of NH3 for these complexes (Table 11). The

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r(N•Cl) distance of 2.561(2) in MeCN•ClF has been subsequently determined using microwave spectroscopy [718]. The HCN•XY (XY = ClF, BrCl, Cl2) linear complexes have been

Ac ce pt e

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described as Mulliken, weak outer type complexes [719].

1.36.5.

Cl2 and ClF complexes with PH3, PMe3, PF3 and PCl3

For PH3 and PMe3 very strongly bound molecular complexes with ClF and Cl2 were observed in matrices [742]. The Cl-F stretching mode shifting from 768 cm−1 to 406 cm−1 in the H3P•ClF complex and 311 cm−1 in the Me3P•ClF complex. These are similar values to those of the ClF− anion (ca. 340 cm−1), indicating substantial rearrangement of the charge in the complexes. In the case of Cl2 a product band at 512 cm−1was assigned to the νCl-Cl mode in the H3P•Cl2 complex. Whilst no νCl-Cl mode was observed for the Me3P•Cl2 complex this was thought most likely to be because it lay beyond the range of the spectrometer, rather than being absent. The observation of an “intermolecular” vibration between the acid and base sub-units at 265 cm−1 was also indicative of a strong interaction as these are rarely observed above 250 cm−1 [743]. When PF3 was employed using twin-jet deposition with ClF numerous intense product bands including those from PF4Cl and PF5 were observed, indicating direct reaction rather than 85 Page 85 of 213

complex formation [742]. With Cl2 and PCl3 or PF3 there was no evidence of product bands until after Hg arc photolysis which yielded intense features due to PF3Cl2 and PCl5. The twinjet experiment involving PCl3 and ClF only gave two weak absorptions, with no effect of photolysis. These do not correspond to the expected PFCl4 product, and due to their weak intensity were not assigned [742].

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The order of reactivity is ClF > Cl2 and PF3 > PCl3. Comparison with other studies involving ClF indicates the strength of interaction is R3P > R2S > R2O. This is consistent with ClF

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hydrogen by a methyl group increased the extent of interaction.

cr

acting as a soft Lewis acid via its more polarisable chlorine atom. Also the substitution of

As for other examples, Legon has employed rotational spectroscopy after the matrix isolation

an

studies to obtain detailed structural information, however, the prior matrix work does not appear to be have been referenced. The structures of H3P•Cl2, H3P•Br2, H3P•BrCl and H3P•ICl all had C3v symmetry and P-X distances of 3.240(15) Å [744]; 3.0440(4) Å [745],

1.36.1.

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2.869(1) Å [746] and 2.963(1) Å [747], respectively.

F2 complexes with O, S and P donor Lewis bases.

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Codeposition of H2O and F2 in solid argon resulted in the formation of the H2O•F2 complex

Ac ce pt e

with a νFF mode at 877.5 cm−1. The initial calculations supporting the experimental work indicated that the oxygen lone pair, σ bonded Lewis acid-base complex (Figure 25(a)) was more stable than the π bonded sideways geometry (Figure 25(b)) by 0.8 kJ mol−1, and the hydrogen bonded isomer by 2 kJ mol−1 [748]. However, the subsequent microwave spectrum was consistent with a Cs structure with a Lewis acid-base interaction between the F2 and H2O with a O-F distance of 2.7480(27) Å and the H2O plane 48.5° from the F-F axis [749, 750]. Our calculations using B3LYP, 6-311G+dp indicate that the structure in Figure 25(b) is about 3 kJ mol-1 more stable than that in Figure 25(a). The calculated equilibrium dissociation energy, De, was 5.3 kJ mol−1 compared to 21.2 kJ mol−1 for the analogous ClF complex [749, 750]. Calculations carried out by other workers at around the same time reported a binding energy of −6.1 kJ mol−1 for the H2O•F2 [751]. When PH3 was co-condensed with F2 the only products observed were PH3F2, PHF2 and PH2F [722]. No intermediate complex analogous to H3N•F2 was observed [459, 723]. It was postulated that the reaction proceeded through a penta-coordinated complex with one axial and one equatorial fluorine, before relaxing to the more stable D3h structure with two axial fluorines which then eliminated H2 to yield PHF2, or 86 Page 86 of 213

HF to give PH2F. Hence PH3F2 with D3h symmetry obeys the “Muetterties rule” [362] for substituted fluorophosphoranes [752] rather than that expected from simple VSEPR considerations. Whilst the reaction between ClF or Cl2 and sulfur [702] or phosphorus [742] bases resulted in the formation of complexes, in the analogous reactions between F2 and alkyl sulphides and trimethylphosphine there was no evidence for the formation of complexes or intermediates, instead direct reaction was observed, leading to hydrogen displacement or

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direct addition products [563]. For PMe3 the reaction product was PMe3F2 [722, 743] and for Me2S the product was tentatively identified as Me2SF2 [563]. For MeSH the products were

cr

not those observed in F atom reactions and were most likely to be MeSF and CH2FSH as well as HF [563]. The reaction between F2 and H2S led to an initial H2SF product which then

us

underwent HF elimination when photolysed with Hg arc radiation [563]. Subsequent work by Andrews identified a H2S•F2 complex (although (H2S)2•F2 could not be entirely ruled out), in

an

which the F2 was dissociated by 590-100 nm radiation to yield several other products [753]. It should be noted that for the analogous H2O•F2 complex 290 nm radiation was required [748].

distance in H2S•F2 was 3.20(1) Å [754].

ClF and F2 complexes with AsH3

d

1.36.2.

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The microwave derived O-F distance in H2O•F2 was 2.7480(27) Å [749, 750] and the S-F

Ac ce pt e

The reaction of ClF with AsH3 in solid argon gave H3As•ClF [704], however, AsH3 and F2 react to give AsH2F as the major product, apparently through an AsH3F2 intermediate [755].

1.36.3.

ClF and Cl2 complexes with O, S and N macrocyclic ligands

The red-shifts in the νClF and νClCl modes for 1:1 complexes between ClF or Cl2 and crown ethers were similar to those observed for other single atom bases, especially oxygen donor atom bases [756]. The spectral data indicated that there was little or no reorganisation of the crown ether, and little cooperativity between the oxygen atoms in the ring. However, the νClF and νClCl modes were broader than those compared to simpler bases. Analogous experiments with HF and HCl led to the formation of hydrogen bonded complexes [757]. With sulfur and nitrogen containing macrocycles such as cyclam, tetramethylcyclam (TMC) and 1,4,8,11tetrathiacyclotetradecane (TTP) 1:1 complexes were observed for ClF and hydrogen bonded complexes between HCl and cyclam or TTP [627].

87 Page 87 of 213

1.36.4.

ClF, Cl2 and Br2 complexes with t-butyl chloride and t-butyl bromide

Twin-jet reactions between ClF, Cl2 or Br2 and t-butyl chloride or t-butyl bromide led to formation of complexes in matrices [758]. For the 1:1 complexes of ClF and t-butyl chloride a shifted νClF mode as well as shifts in the base modes was observed, but with Cl2 and Br2 complexes, no activation of the νClCl or νBrBr modes was observed due to the relative weakness of the interaction in the complex. In the case of t-butyl bromide, the major

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products, even with twin-jet deposition were t-butyl chloride and its complexes with ClF and Cl2. This indicates a very rapid halogen exchange, and that as the reaction was likely to occur

cr

on the cryogenic surface, the activation barrier must be low. There was some evidence for the formation of Me3CBr•ClF and Me3CBr•ClF complexes indicating that they can be isolated if

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sufficient energy is removed from the reagents prior to reactive collision. When single-jet reactions were carried out there was no evidence of ClF remaining in the spectrum, and the

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observed bands were assigned to Me3CF indicating extensive reaction had taken place [758].

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1.36.5. ClF, Cl2 and Br2 complexes with cycloalkanes Single-jet experiments were used to investigate the reaction of cyclopropane and its derivatives (methylcyclopropane, 1,1-dimethylcyclopropane and bromocyclopropane) with

d

Cl2, whereas for the experiments involving ClF twin-jet approach was used, and both twin

Ac ce pt e

and single-jet experiments were used for Br2 [759]. The shifts in the halogen stretching modes (and partial activation in the case of Cl2 and Br2), together with the shifts in the cyclopropane vibrational modes were taken to indicate the first identification of 1:1 complexes of cyclopropane and its derivatives with ClF, Cl2 and Br2. It was postulated that the interaction was between the halogen and one of the carbon-carbon bonds in the ring (except for bromocyclopropane). Subsequent microwave studies (which did not reference the matrix work) confirmed the mode of coordination as being between the Cl and one of the C-C bonds with the three carbon nuclei as well as the chlorine and fluorine nuclei being coplanar in cyclopropane [760], but in methylenecyclopropane the ClF was perpendicular to the molecular plane with the Cl interacting with the double bond [761]. 1.36.6.

ClF and Cl2 complexes with alkenes and alkynes

In all of the above complexes it has been assumed that the ClF (or Cl2) is coordinated end-on to the donor atom of the base. However, for the complexes of halogens such as ClF with the π electron density of alkynes or alkenes a T-shaped structure was inferred from twin-jet matrix experiments for the 1:1 complexes between ClF and C2H2, C3H4, C4H6, C2H4, C3H6, cis-C4H8 88 Page 88 of 213

and iso-C4H8, and the 1:1 complexes between Cl2 and C2H2, C3H4, C2H4 and C3H6 [762]. The position of the νClF mode in the complexes was found to be dependent on whether the halogen was interacting with a double or triple bond, and also the number of methyl substituents associated with the multiple bond. For example for ethylene the shift in the νClF mode was 106 cm−1, whereas for acetylene it was 53 cm−1, indicating greater interaction for the former. Increasing the number of methyl groups resulted in a further decrease in the νClF mode,

ip t

indicating greater interaction. Subsequent microwave spectroscopic studies [763-766] confirmed the T shaped geometry proposed from the matrix experiments [762]. For the C2H2

cr

complex the distance from the midpoint of the triple bond of C2H2 to Cl was 2.87 Å [764], whereas the distance from the centre of the double bond of C2H4 to Cl in the C2H4 complex

us

was 2.768(3) Å [765]. This correlates well with the observed shifts in the νClF modes in the matrix data [762]. The T shaped geometry was also the minimum energy structure from the

an

calculations that indicated the interaction energy for the C2H2 ClF complex was −8.4 kJ mol−1

1.36.7.

M

[767].

ClF complexes with benzene and its derivatives

1:1 complexes between ClF and benzene, as well a series of substituted benzenes have been

d

stabilised in argon matrices [768]. For the C6H6•ClF complex the νClF mode red-shifted by 60

Ac ce pt e

cm−1 from that in free ClF. The perturbations of the benzene vibrational modes implied that the ClF was sat axially above the benzene ring, interacting with the π-electron density. For C6H5Br, there appeared to be two 1:1 complexes formed, one as for C6H6with the ClF above the ring, and a second one with the ClF interacting with the Br [768]. Calculations confirm that the Cl interacts with the π electrons of the benzene ring, but there was significant deviation from C6v symmetry [769]. A subsequent microwave study confirmed that the ClF sits atop the benzene at 3.765 Å, but is tilted by 14.4°, and it executes an approximately circular path sampling the π electron density on one face of the aromatic molecule governed by a Mexican hat potential [770].

1.36.8. Interhalogen and halogen complexes summary In the above sections it is clear that Legon has carried out a detailed and in-depth investigation of halogen and inter-halogen complexes using rotational spectroscopy, and in particular pulsed-nozzle or fast-mixing nozzle, Fourier transform microwave spectroscopy which enables the study of very reactive species. This has provided a wealth of additional 89 Page 89 of 213

information on the geometries and properties of the complexes observed in the matrix experiments, which surprisingly were not always referenced. There are some general reviews [698-700, 771-773], as well as reviews of F2 [774], Cl2 [775], ClF [766, 776], BrCl [777] and ICl complexes [778] and DFT calculations [779]. A complete listing is given in Table 22. The extensive matrix study by Ault led to the conclusion that CIF is a soft Lewis interacting

ip t

through its polarisable chlorine atom with the lone pair on the Lewis base, and this was supported by all the subsequent microwave studies. A graphical summary of the matrix data

cr

(based on that given in [713]) is presented in Figure 26. It is clear that there is a reasonable correlation between the νClF stretching mode and base proton affinity for similar bases, but

us

that between types of bases there are clear differences. This leads to stronger complexes with soft bases (R3P and R2S) than with hard bases (such as R2O) and thus the order of Lewis

an

basicity for ClF is R3P > R2S > R2O, which does not necessarily reflect the proton affinities of the base [713, 742]. In particular, the softer third row donor atoms, cause a much larger shift in νClF indicating a more substantial interaction, which can be understood on the basis of

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the hard-soft (HSAB) approach. Therefore, whilst proton affinities give some indication of

d

the strength of interaction, other factors also need to be considered.

Ac ce pt e

The plot of νClF versus the Cl-donor atom distance (derived from gas phase microwave studies) in Figure 27 shows a very good correlation between these two parameters, except for H2S. However, this is the only example of a third period donor atom, and as the covalent and van der Waals radii of S are of the order of 0.3 to 0.4 Å greater than those of O, this is not unexpected.

Recent calculational efforts by Li et al. have been directed to ternary complexes such as H3N•XY•HF [780] and H3N•BF3•XY [781] (XY = ClF, Cl2, BrF, BrCl, Br2), XCl-FH2P-NH3 (X = F, OH, CN, NC and FCC) [782] and whilst the authors indicate these are ripe for investigation by matrix isolation techniques, this does not seem to have been followed up.

1.36.9.

Halogen and interhalogen CO complexes

In contrast to the work presented above where the matrix experiments preceded the microwave work, for the CO halogen complexes the gas phase IR spectrum indicated that OC•Cl2 was linear [783] and microwave spectroscopy gave r(C•Cl) of 3.092 Å [784] which was 90 Page 90 of 213

subsequently recalculated as 3.134 Å [785]. The other experimental microwave r(C•X) distances are: 3.105 for OC•Br2 [786]; 2.770 Å for OC•ClF [787]; 3.011 Å for OC•ICl [785]; and 3.004 Å for OC•BrCl [788]. The subsequent matrix work was carried out to search for the less stable complexes, with weaker interactions, and in particular those where the O of CO was bound to the halogen. These experiments have resulted in the observation in argon matrices of the 1:1 CO complexes of Cl2, Br2, BrCl [789]; Cl2, BrCl, ICl, IBr [790], IF [791],

ip t

ClF and BrF [792]. The νCO modes, in conjunction with calculations were used to identify the bonding modes of both the CO and the interhalogens, and these are given in Table 23. The

cr

work of Schriver [789] and Romano [790] used the reaction of the halogen (or interhalogen) with CO in argon matrices to form 1:1 complexes, where νCO modes higher than νCO in “free”

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CO were used to identify the mode of bonding (OC•XY vs. OC•YX) in the interhalogens. In addition isocarbonyl complexes were observed for: CO•ICl; CO•ClI; CO•IBr; and CO•BrI

an

[790], and these were characterised by a νCO mode lower than that of “free” CO. For the IF, ClF and BrF complexes, the CO complexes were prepared by photodissociation of FC(O)Br and FC(O)Cl in argon matrices [791, 792]. In the case of FC(O)Br, three isomers were

M

identified, OC•BrF, OC•FBr and CO•BrF [792]. For FC(O)Cl, only one complex, OC•ClF was observed. Romano et al. used photolysis of the CO and ICl or IBr complexes to prepare

d

IC(O)Cl and IC(O)Br for the first time, as well as the known compounds OCCl2, OCBr2 and

Ac ce pt e

BrC(O)Cl from Cl2, Br2 and BrCl complexes with CO [793]. 1.36.10.

Complexes between the halogens and interhalogens and OCS and OCSe

Della Védova et al. have carried out a series of experiments on the interaction of OCS with dihalogens such as Cl2, ICl, IBr, BrCl in argon matrices. Whilst it was reported that weak complexes were formed between OCS and Cl2, IBr and ICl, no vibrational data was reported other than a comment that there were some changes in the patterns in the OCS fundamental at 2049.8 cm−1 [794]. A subsequent paper from the same group has reported the 1:1 adducts between OCS and ClF, Cl2 and BrCl by photolysis of FC(O)SCl, ClC(O)SCl and BrC(O)SCl, respectively [795]. Whilst condensation of OCS and BrCl gave evidence of complexes, assignment of the overlapping bands from Cl2, ClBr and Br2 components made assignment difficult. However, the OCS•ClF and OCS•Br2 1:1 complexes were formed on condensation of the parent species, and OCS and Cl2 gave predominantly 1:2 complexes formulated as Cl2•OCS•Cl2 [795]. Calculations revealed that the binding energies were small, and even for the largest (OCS•ClF) this was less than −10 kJ mol−1, with many of the remainder being less than −5 kJ mol−1. The calculations also indicated that there were two minima on the potential 91 Page 91 of 213

energy surface for many of the complexes, one with an angular structure close to 90° between OCS and the halogen, and another one where they were collinear. The most tightly bound appear to be those with an S•X linkage and an angular structure [795]. In analogous reactions involving OCSe and F2 [796], Cl2 or Br2 [797], the OCSe•F2 [796], OCSe•Cl2, OCSe•Br2, and SeCO•Br [797] complexes were used as precursors for the

ip t

photochemical formation of FC(O)SeF, ClC(O)SeCl and BrC(O)SeBr molecules, respectively.

Complexes between the halogens and interhalogens and CS2

cr

1.36.11.

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Formation of what are described as van der Waals complexes between CS2 and Cl2, Br2, Brl, ICl or IBr, have been reported involving the interaction of the lone pair on one of the sulfur atoms in CS2 with the σ* orbital on the dihalogen, resulting in planar L shaped complexes

an

[798]. The calculated S···X distances were 3.3 to 3.5 Å., which are shorter than the sum of the van der Waals radii by 0.3 to 0.5 Å, thus fulfilling the dp criterion. These complexes have

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then been used to prepare a variety of new species [799]. The complexes between CS2 and

d

HCl have the same structural motif [800] as those of the halogens [798].

Ac ce pt e

Group 18 As expected the coordination chemistry of the group 18 elements is limited, but matrix techniques have allowed for some significant breakthroughs. These include the first observation of KrF2 [801], XeH2 [802], HXeCl, HXeBr, HXeI, HKrCl [803] and a wide selection of compounds which have recently been reviewed [804]. Although we have detected XeH2 using thermal hydrogen atoms produced from the photolysis of H2Te in xenon matrices, we could find no evidence for the formation of KrH2 in krypton matrices [805].

Perutz and Turner showed that photolysis of metal hexacarbonyls yields metal pentacarbonyls with IR and UV-vis spectra that were both dependent on the matrix material, indicating the vacant site was occupied by a matrix atom with a weak metal-matrix bond [806]. This is often overlooked as one of the first reports of metal-noble gas bonding interaction. Many others now claim that there is some noble gas (Ng) bonding in a variety of compounds isolated in cryogenic solids [807-812]. 92 Page 92 of 213

Probably the most significant development is the identification of the first stable argon compound, HArF, formed from the photolysis of HF in argon matrices which was reported in 2000 by Räsänen’s group in Helsinki [813]. There has been considerable experimental and theoretical work since then which has been reviewed in 2009 [804] and 2011 [814]. Whilst species may be stable above the cryogenic range [815, 816].

cr

Conclusions

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this was only stable at very low temperatures, there is indication that some of the matrix

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Cryogenic techniques, and in particular matrix isolation have allowed for the identification and characterisation of a wide range of main group coordination compounds, ranging from

an

those with fairly strong interactions through to those with very small interactions, including intermediates not accessible under ambient conditions. Many of these were observed for the first, and sometimes only time, in matrices, but in a few cases detailed matrix experiments

M

allowed for the filling in of existent data, especially for the weakly bound complexes. Whilst much has been done, it is clear that there is ample opportunity for the experimentalist to

Ac ce pt e

complexes.

d

continue to use matrix isolation methods to prepare and stabilise a wide variety of main group

Although there are a plethora of calculations available, these use a wide variety of methods so are not always comparable. However, those by Ford are probably the easiest to make use of as they are specifically related to matrix isolation results (his and others), and have been carried out and reported in a consistent manner. A summary of the calculated binding energies calculated by Ford et al. for a range of bases with BF3, SiF4 and SO2 Lewis acids are shown in Figure 28 (based on a figure in [393]). It is clear that the largest binding energies are for NH3 complexes of the three acids, and that BF3 is the strongest of the Lewis acids for each base. Unsurprisingly there is a reasonable correlation between the interaction energies and the shifts of the modes in the acid and bases and Figure 29 collects together all the proton affinity data in the earlier graphs in the form of % shift of the relevant vibrational mode vs proton affinity in order to provide some normalisation. Whilst it may be naive as all vibrational modes will not be affected to the same extent, it is clear that the BF3 shifts are the largest. This a busy figure, but in addition to reinforcing the general trend that the vibrational 93 Page 93 of 213

shifts increase with base proton affinity, it allows for comparison of the shift induced by each base by reading the data vertically, in conjunction with the proton affinities in Table 1. For example the NH3 data is at 853.6 kJ mol-1. The base proton affinity gives a good indication of the strength of interaction (as measured by the shift in the acid and base vibrational modes) for Me2Zn, Me2CD, Me3Ga, Me3In, with a wide range of donors atom types. Whilst the correlation is good for SiF4 and GeF4 complexes, data is only available for O and N donors so

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this may be a false sense of security. For BF3, SO2, SO3 and ClF complexes there is a less straight-forward relationship between proton affinity and strength of interaction, indicating

cr

that other considerations such as HSAB need to be invoked as well.

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It has also been shown that there is a good correlation between the matrix derived data for either acid or base vibrational modes and the distance between the donor and acceptor atom

an

derived from the gas phase rotational spectroscopic data. With caution this can be used to estimate the distances in the complexes for which gas phase data is not available. Whilst it is tempting to attempt a similar exercise with the calculated bond energies, the weakness of the

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interaction means that the values are very susceptible to methodology and basis set chosen. In addition, there is a selection of parameters reported, which have been subject to a variety of

Ac ce pt e

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corrections. Therefore, this task lies outside the scope of this review.

Acknowledgements This review is dedicated to the memory of Steve Ogden, who introduced me to, and taught me so much about the art of matrix isolation and science of vibrational spectroscopy. The University of Hull is gratefully thanked for supporting the preparation of this review.

94 Page 94 of 213

Figure Captions Figure 1 Schematic representation of the formation of the dative bond in the H3N•BH3 complex.

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Figure 2. Calculated structure of the H3N•ZnMe2 complex. Figure 3. Plots of the ZnC2 asymmetric stretching mode of Me2Zn complexes versus the

cr

proton affinity of the base. (a) Argon matrix data, (b) Cold thin film data. Uncomplexed

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Me2Zn is at 616.6 cm−1 in argon matrices and 602 cm−1 in thin films.

Figure 4. Plots of the CdC2 asymmetric stretching mode of Me2Cd complexes versus proton

an

affinity of the base. (a) Argon matrix data, (b) Cold thin film data. Uncomplexed Me2Cd is at

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544.3 cm−1 in argon matrices and 525/520 cm−1 in thin films.

Figure 5. Representation of the structure of the H3N•BF3complex.

Ac ce pt e

(See text for references)

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Figure 6. Plot of calculated reaction energies for a variety of BF3 Lewis acid-base complexes.

Figure 7. Plots of the 11BF3 asymmetric stretching mode (a) and symmetric deformation mode (b) for some BF3 complexes isolated in cryogenic matrices versus the proton affinity of the bases. The modes for uncomplexed 11BF3 in solid argon are at 1445 and 676 cm−1. Figure 8. Plot of the 11BF3 asymmetric stretching mode in BF3 complexes isolated in cryogenic matrices versus the experimental boron-donor atom distance in the gas phase. The mode for uncomplexed 11BF3 in solid argon is at 1445 cm-1. Figure 9. Plots of the (a) AlC3, (b) GaC3 and (c) InC3 asymmetric stretching modes for Me3Al, Me3Ga and Me3In complexes in cryogenic matrices versus proton affinity of the base. The modes for uncomplexed Me3Al, Me3Ga and Me3In in argon matrices are at 688, 575 and 505 cm−1.

95 Page 95 of 213

Figure 10. Representation of the structure of the H3N•SiF4 complex. Figure 11. Plot of calculated interaction energies for SiF4 Lewis acid-base complexes (data from [393]). Figure 12. Plots of the SiF and GeF stretching modes of (a) SiF4 and (b) GeF4 complexes in

ip t

cryogenic matrices versus the proton affinity of the base. The modes for uncomplexed SiF4

cr

and GeF4 in argon matrices are 1025 and 800 cm−1, respectively.

Figure 13. Calculated structure of ECl2(N2)n complexes (E =Si, Ge, Sn, n = 1 or 2) (Redrawn

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from [405])

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Figure 14. Plot of the NH3 symmetric deformation mode for NH3 complexes isolated in cryogenic matrices versus the experimental nitrogen-acceptor atom distance in the gas phase.

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Figure 15. Representation of the structure of the H3N•CO2 complex.

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Figure 16. Representation of the H3N•N(O)Cl complex. Figure 17. Representation of the H2O•CO2 complex. Figure 18. Representation of the H3N•O3 complex. Figure 19. Representation of the H3N•SO2 complex. Figure 20. Calculated interaction energies for SO2 Lewis acid base complexes. Figure 21. Plot of SO2 asymmetric stretching mode of SO2 complexes in cryogenic matrices versus proton affinity of the base. Uncomplexed SO2 is at 1347 cm−1 in argon matrices. Figure 22. Representation of H3N•SO3 complex. Figure 23. Plot of asymmetric SO3 stretching mode of SO3 complexes in cryogenic matrices versus proton affinity of the base. Uncomplexed SO3 is at 1385 cm−1 in solid argon. 96 Page 96 of 213

Figure 24. Representation of H3N•ClF complex. Figure 25. Representation of H2O•F2 complex. Figure 26. Plot of the ClF stretching mode in ClF complexes in cryogenic matrices versus ClF is at 768 cm−1 in argon matrices.

cr

35

ip t

proton affinity of the base. Straight lines are best fits through the relevant data. Uncomplexed

Figure 27. Plot of the ClF stretching mode for ClF complexes isolated in cryogenic matrices

us

versus the experimental chlorine-donor atom distance in the gas phase.

an

Figure 28. Calculated interaction energies of a selection of Lewis acid base complexes. (Redrawn from [393])

Ac ce pt e

d

versus proton affinity of the base.

M

Figure 29. Plot of the % shift in vibrational modes of main group Lewis acid base complexes

97 Page 97 of 213

us

cr

ip t

Figure 1

Ac ce pt e

d

M

an

Figure 2

98 Page 98 of 213

Figure 3

(b) Thin film

610

605

605

600

600

595

595

590

590

585

585

580

580

575

575

570

ip t

615

cr

(a) Ar matrix

610

570

N donors P donors As donors O donors S donors Se donors

565 560 555 550 650

700

565

N donors P donor As donor O donors S donors Se donors

560 555 550 750

800

850

900

950

650

-1

700

750

800

850

900

950

-1

Proton affinity / kJ mol

Ac ce pt e

d

M

an

Proton affinity / kJ mol

us

ZnC2 asymmetric stretching mode / cm

-1

615

99 Page 99 of 213

Figure 4

(a) Ar matrix

540

(b) Cold solid film

530

530

525

525

520

520

515

515

510

ip t

535

510

N donors O donors S donors Se donors

505

N donors O donors S donors

505

500

500 725

750

775

800

825

850

700

725

750

an

700

-1

775

800

825

850

-1

Proton affinity / kJ mol

Ac ce pt e

d

M

Proton affinity / kJ mol

cr

535

us

CdC2 asymmetric stretching mode / cm

-1

540

100 Page 100 of 213

Ac ce pt e

d

M

an

us

cr

ip t

Figure 5

101 Page 101 of 213

d

Ac ce pt e

N2

-100

cr

ip t

-150

N2O

-1

-175

H2S

-125

H3P

C2N2

FCN

-50

us

an

-75

HCN

0 HC3N

-25

H2O

MeH2P

M

Me2HP

NH3

Me3P

MeH2N

Me2HN

Me3N

Lewis acid base complexes /kJ mol

Calculated interaction energies of BF3

Figure 6

-200

Base

102

Page 102 of 213

Figure 7

1450

670

(a)

(b)

660 -1

1400

1325 1300 1275 1250

N donors O donors S donors C donors

1225 1200

640

ip t

1350

650

630 620

610

600

N donors O donors S donors C donors

590

580 500

550

600

650

700

750

800

cr

1375

500

850

-1

us

BF3 symmetric deformation mode / cm

BF3 asymmetric stretching mode / cm

-1

1425

550

600

650

700

750

800

850

-1

Proton affiniy / kJ mol

Ac ce pt e

d

M

an

Proton affiniy / kJ mol

103 Page 103 of 213

N2

1450

SO2 CO

1400

ip t

1350

1250

NH3

cr

1300

MeCN

1.6

1.8

2.0

2.2

2.4

2.6

2.8

Ac ce pt e

d

M

an

B - donor atom distance / Å

3.0

us

1200

11

BF3 asymmetric stretching mode / cm

-1

Figure 8

104 Page 104 of 213

Figure 9

690

ip t

680

670

cr

660

650

640

us

AlC3 asymmetric stretching mode / cm

-1

(a) Me3Al complexes

N donors P donors O donors

630

620 600

650

700

750

800

850

900

950

an

-1

Proton affiniy / kJ mol (b) Me3Ga complexes

M

575

570

565

d

560

555

N donors P donors As donors Sb donors O donors S donors

Ac ce pt e

GaC3 asymmetric stretching mode / cm-1

580

550

545

540

600

650

700

750

800

850

900

950

900

950

-1

Proton affiniy / kJ mol

InC3 asymmetric stretching mode / cm-1

515

(c) Me3In complexes

510

505

500

495

490

485

N donors P donors As donors Sb donors

480

600

650

700

750

800

850 -1

Proton affiniy / kJ mol

105 Page 105 of 213

Ac ce pt e

d

M

an

us

cr

ip t

Figure 10.

106 Page 106 of 213

d

Ac ce pt e Base

cr

-1

ip t

-80

0 HCl

us

-60

H2S

-20

an

-40

PH3

HF

M

H2O

H3N

Lewis acid base complexes /kJ mol

Calculated interaction energies of SiF4

Figure 11

-100

107

Page 107 of 213

Figure 12

(a) SiF4 complexes

(b) GeF4 complexes

800 790

1000

780

980 970 960 950 940

760 750 740 730 720

N donors O donors 500

600

710 700

800

900

1000

-1

500

600

700

800

900

1000

-1

Proton affinity / kJ mol

Ac ce pt e

d

M

Proton affinity / kJ mol

N donors O donors

an

930

770

cr

990

us

Ge-F stretching mode /cm

Si-F stretching mode /cm

-1

-1

1010

ip t

1020

108 Page 108 of 213

Figure 13.

E

N

N Cl

N

N

Cl

Cl

d

M

an

us

cr

ip t

Cl

E

N

Ac ce pt e

N

109 Page 109 of 213

Figure 14 BH3

SO3

1300

BF3

1250

SiF4

ip t

1200 1150 1100

ClF

cr

1050

F2

950 1.5

1.6

1.7

1.8

1.9

2.0

2.1

2.2

2.3

2.4

2.5

2.6

2.7

2.8

d

M

an

Nitrogen - acceptor atom distance / Å

us

Cl2

1000

Ac ce pt e

NH3 symmetric deformation mode / cm

-1

1350

110 Page 110 of 213

an

us

cr

ip t

Figure 15

Ac ce pt e

d

M

Figure 16

111 Page 111 of 213

an

us

cr

ip t

Figure 17

Ac ce pt e

d

M

Figure 18

112 Page 112 of 213

Ac ce pt e

d

M

an

us

cr

ip t

Figure 19

113 Page 113 of 213

d

Ac ce pt e -10

cr

-12

-1

ip t

-14

-2

Base H2S

-6

us

-8

H3P

-4

an

M

0 H2O

H3N

Lewis acid base complexes /kJ mol

Calculated interaction energies of SO2

Figure 20

-18

-16

114

Page 114 of 213

Figure 21

1350

ip t

1340 1330

cr

1320 1310 1300

1280 1270 1260 550

N-donors (amines) O-donors S-donors (thiols) S-donors (thioethers) pi-donors HBr, S, MeCN 600

650

700

us

1290

750

800

an

SO2 asymmetric stretching mode / cm

-1

1360

900

950

Ac ce pt e

d

M

Proton affinity /kJ mol

850

-1

115 Page 115 of 213

Ac ce pt e

d

M

an

us

cr

ip t

Figure 22

116 Page 116 of 213

Figure 23

ip t

1380

1360

cr

1340

1280

N donors P donors O donors S donors 700

750

800

850

900

950

d

M

Proton affinity / kJ mol

-1

an

1300

us

1320

Ac ce pt e

SO3 asymmetric stretching mode / cm

-1

1400

117 Page 117 of 213

an

us

cr

ip t

Figure 24

Figure 25 (a)

Ac ce pt e

d

M

(b)

118 Page 118 of 213

750

O donors

700

alkenes

600

H2S

N donors

500

cr

550

us

S donors AsH3

an

450 400

P donors

350

M

ClF stretching mode / cm

-1

650

O donors N donors P donors S donors alkene alkyne

alkynes

ip t

Figure 26

300 700

750

800

d

650

850

900

950

-1

Ac ce pt e

Proton affinity / kJ mol

119 Page 119 of 213

Figure 27

750

H2CO (CH2)2O

MeCN

ip t

H2S

650

600

NH3

cr

550

Me3N 450 2.0

2.1

2.2

2.3

2.4

2.5

2.6

2.7

2.8

2.9

d

M

an

Cl-donor atom distance /Å

us

500

Ac ce pt e

ClF stretching mode / cm

-1

700

120 Page 120 of 213

Figure 28

-140

BF3 SiF4

ip t

-120

SO2

cr

-100 -80

us

-60 -40

an

Calculated interaction energy / kJ mol

-1

-160

-20 0 H2O

HF

PH3

M

NH3

H2S

HCl

Ac ce pt e

d

Base

121 Page 121 of 213

Figure 29

20

Me2Zn Me2Cd

18

BF3

ip t

Me3Al Me3Ga

14

Me3In

12

cr

SiF4 GeF4

10

SO2

8

us

SO3 ClF

6 4

an

% shift in vibrational mode

16

2

550

600

650

700

M

0 750

800

900

950

-1

Ac ce pt e

d

Proton affinity / kJ mol

850

122 Page 122 of 213

H2O2

643.8

O2 O3 SO2 SO3 SF4

396.3 595.9

550.7

421 625.5 676 577 575.3

793.6 782.1 764.5 683.3 710.3

825.3 812.0 792.0 712.9 742.0

MeOH EtOH MeSH EtSH n-C3H8SH i-C3H8SH t-C4H9SH

724.5 746 742 758.4 763.6 772.3 785.1

754.3 776.4 773.4 789.6 794.9 803.6 816.4

764.5 801 801.2 827.0 815.3 770.0

792 828.4 830.9 856.7 846.5 797.4

870.9 820.2

918.8 858.0

745.3 773.9 794.7

774.2 801.3 822.1

C2H4 C3H6 C4H8

651.5 722.7 719.9

680.5 751.6 747

C2H2 C3H4 C3H6

616.7 723 722.7

641.4 748 751.6

C6H6 C6H5CH3

725.4 756.3

750.4 784.0

Group 15 H3N MeH2N Me2HN Me3N H5C2NH2

819.0 864.5 896.5 918.1 878.0

853.6 899.0 929.5 948.9 912.0

HCN MeCN MeNC H5C5N PhCN Me3CCN CH2CHCN

681.6 748.0 806.6 898.1 780.9 806.6 753.7

712.9 779.2 839.1 930.0 811.5 839.1 784.7

Me2O Et2O Me2S Et2S MeEtS 1,4-dioxane 1,4-thioxane 1,4-dithiane diglyme MeOCH2CH2OMe

HN3 N2 N2O

723.5 464.5 523.3

756.0 493.8 549.8

(CH2)2O (CH2)3O (CH2)4O

d

M

H2CCO Me2CO Me2O H2CO HCOOH

Ac ce pt e

674.5

cr

681.9 725 594.0 426.3 540.5 772.3

us

657.7 700.9 562.8 402.2 515.8 739.8

an

CS2 CSe2 COd CO CO2 HNC

ip t

Table 1. Gas phase basicities and proton affinitiesa of common bases used in matrix isolation studies. Base Gas phase Proton Base Gas phase Proton affinity basicity affinity basicity −1 −1 −1 /kJ mol /kJ mol /kJ mol−1 /kJ mol Group 14 Group 16 SiH4 613.4 639.7 H2O 660.0 691.0 GeH4 687.1 713.4 H2S 673.8 705 CF4 527 H2Se 676.4 707.8 SiF4 476.6 502.9 H2Te 704.5 735.9

123 Page 123 of 213

575.2 531.8 591.0

(CH)4O

770.9

803.4

853.6 785 747.9 731

(CH2)2S propylene sulfide OCS S Group 17 HF

777.6 801.5 602.6 640.2

807.4 833.3 628.5 664.3

NH3 PH3 AsH3 SbH3

819.0 750.9 712.0

456.7

484

Me3N Me3P Me3As

918.1 926.3 864.9

948.9 958.8 897.3

HBr HI F2

601.3 305.5

569 627.5 332

NF3 PF3 AsF3

662.8

ip t

548.7 505.3 560.3

cr

N2O NO NO2

CH3F 571.5 598.9 CH3Cl 621.1 647.3 CH3Br 638.0 664.2 CH3I 665.5 691.7 Data taken from [72], if only proton affinities are given then data from [71], bold indicates atom involved where base is ambibasic

Ac ce pt e

d

M

an

us

604 695.3 648

124 Page 124 of 213

Table 2. ZnC2 asymmetric stretching mode in ZnMe2 and its complexes in argon matrices and thin cryogenic films. Ar matrix /cm−1 (a) Thin film Proton affinity/ Ref /cm−1 kJ mol−1 (b) Me2Zn 619.5 616.6 614.0 602 [88]

H3P•ZnMe2 (H3P)2•ZnMe2

611.3

607.9

H3As•ZnMe2

611.5

608.7

605.5

599

H2O•ZnMe2 Me2O•ZnMe2 Me2CO•ZnMe2 (Me2CO)2•ZnMe2 H3C(H)O•ZnMe2 (H3C(H)O)2•ZnMe2

599.9 599.5 591.9 596.1 584.3

605 596.9 595.5 589.1 593.5 581.9

593.5 593.8 586.1 591.5 579.5 607.5 607.3

596.7

552 501 557 518

853.6 853.6 948.9 948.9

594 577

785 785

590.9

[88] [88] [88] [88]

[88] [88] [88]

585.6 587.5 573.3 575 556

691.0 792.0 812.0 812.0 754.3 754.3

[89] [88] [88] [88] [90] [90]

596

705

595 584.9 551.8 569 569

707.8 830.9 830.9 773.4 773.4

[88] [88] [88] [88] [88] [90] [90]

an

747.9

M

609.5 609.3 609.4 594.4

d

610.9 610.9

Ac ce pt e

H2S•ZnMe2 D2S•ZnMe2 H2Se•ZnMe2 Me2S•ZnMe2 (Me2S)2•ZnMe2 H3C(H)S•ZnMe2 (H3C(H)S)2•ZnMe2

580.5

ip t

582.4 569 577 527

cr

584.8

us

H3N•ZnMe2 (H3N)2•ZnMe2 Me3N•ZnMe2 (Me3N)2•ZnMe2

Me2Se•ZnMe2 596.7 594.2 591.7 588.3 [88] (Me2Se)2•ZnMe2 552.0 [88] (a) Where three values are given these refer to the 64Zn, 66Zn and 68Zn isotopomers. If one value given then unresolved peak. (b) Proton affinities taken from [71, 72]

125 Page 125 of 213

544.3 527.6 521.6

Me2O•CdMe2 Me2CO•CdMe2 Me(H)O•CdMe2 (Me(H)O)2•CdMe2

535 523 532 527

H2S•CdMe2 Me2S•CdMe2 (Me2S)2•CdMe2 Me(H)S•CdMe2

537.9 533 521 536

M

Me2Cd H3N•CdMe2 (H3N)2•CdMe2

525 504 475

d

542 498 482.5

Ac ce pt e

Me2Cd Et2Cd (Me2Te)1,2•CdEt2

an

us

cr

ip t

Table 3. CdC2 asymmetric stretching mode in Me2Cd and its complexes in argon matrices and thin cryogenic films. Ar matrix thin film Proton affinity(a) Ref /cm−1 /cm−1 Me2Cd 520 [100] TMEDA•CdMe2 479 989 [100] TMDPE•CdMe2 518 [100] diglyme•CdMe2 505 918.8 [100] MeOCH2CH2OMe•CdMe2 512 858.0 [100] 1,4-dioxane•CdMe2 518 797.4 [100] 1,4-thioxane•CdMe2 520 [100] 1,4-dithiane•CdMe2 520 [100] Et2S•CdMe2 503 856.7 [100] (CH2)2S•CdMe2 517 807.4 [100] propylene sulfide•CdMe2 508 833.3 [100] OCS•CdMe2 520 628.5 [100]

H2Se•CdMe2 536.9 Me2Se•CdMe2 532 (Me2Se)2•CdMe2 521 (a) Proton affinities taken from [71, 72]

521 506 513

853.6

[85] [85] [85] [103] [103] [103]

792.0 812.0 754.3

[104] [104] [104] [104]

705 830.9

[103] [104] [104] [104]

773.4 707.8

[103] [104] [104]

126 Page 126 of 213

Table 4. Table of boron fluoride Lewis base complexes in cryogenic matrices . Band position /cm−1

Complex/m atrix

Proton affinity of base / kJ mol−1

Ref

11

10

BF3 sym str 11 B B

BF3 sym def 10 11 B B

BF3 gas phase BF3/Ne BF3/Ar BF3/Ar BF3/Ar

1505.78

1453.98

885.64 5

719.28

691.21

1499.2 1498.4

713.5 703.5

685.6 676.2 676.0 664

BF3/Ar BF3/Ar

1493.5 (1494.8 1490.2)(b) (1495.4 1491.9 1488.3)(b)

1447.5 1446.9 1447.0 1442g (1447 1444 1441 1437)(b) 1442.3 (1443.4 1439.2)(b) (1447.1 1443.7 1440.5 1436.6)(b) 1442.3

BF3/Ar

1493.5

(BF3)2/Ar

1514 (1475/147 3)(b) (1494 1489)(b)

BF3/N2 BF3/N2 BF3/N2 BF3/N2 BF3/N2

(1495 1490)(b) (1495.8 1492.0) (b)

(1442 1438)(b) 1440 (1443 1438)(b) (1445.5 1440.9)

(1494.8 1490.2)(b) (1494.6 1490.2)(b) 1505 1476(c)

(1443.5 1439.1)(b) (1443.3 1439.1)(b) 1456 1424(c)

(1489.8 1488.3)(b) 1488.3 (1490.1 1487.6)(b)

(1438.7 1437.3)(b) (1435.9 1432.2)(b) (1436.7 1434.2)(b)

1494.5 1268

1443.4 1216

(BF3)2/N2

OC•BF3/Ar OC•BF3/Ar OC•BF3/N2

N2•BF3/Ar H3N•BF3/Ar

[817819] [168] [123] [195] [122]

us 673.5 664.4

[184] [183]

654.1

[175]

673.5

[173, 174] [120, 122]

an

702.9 684.1 676.0

702.9

Ac ce pt e

BF3/N2

1456 1424(b)

676

M

BF3/Ar

885.84 3

d

(1495 1492 1488)(b)

B

cr

BF3 asym str 10 B

ip t

(a)

(e)

650 680

653

[120]

684

653 658

[128] [122]

684.0

662.3

684.0

664.4

[173, 174, 184] [183]

879.6

683.8

657.6

[175]

820

678

651

[120] [122]

659.6

636.2

594.0

[123]

658.2

636.2

594.0

[125]

678.4 672.1

649.8 646.8

594.0

[125]

680.3

654.3

493.8 853.6

[123] [137]

(b)

879.6

853.2

(b)

888/87

888/87

(b)

127 Page 127 of 213

4d 1021.3 (856)a

4d 932.1 1013.6 (856)a

617

655.9 667.9 601

853.6 853.6 779.2

[184] [184] [151]

1281

1277

1226

1293.3

1248.7

833/83 8(b) 844/83 6 813.5

833/83 8 (b) 844/83 6 811.7

600/60 3 (b) 640/63 3 (b) 618.0

584/58 6 (b) 625/61 8 (b) 601.2

779.2

[152]

779.2

[152]

779.2

[156]

1277.8

1235.0

843.5

835.8

640.6

625.3

779.2

[156]

1284.8

1239.0

810.8

809.0

616.9

601.7

779.2

[156]

1295

1253

848

848

602

590

811.5

[160]

1290

1247

840

840

[160]

1203.8

852.2

840.4

603/60 7 638.0

810.9

1244.1

617/62 1 651.7

839.1

[195]

OCO•BF3/A r OCO•BF3/N

1486.4

1435.5 1453.9(c) 1435.5 1453.9(c) 1285 1241(c) 1240 1225(c)

681.2

676.7

540.5

[183]

676.8

540.5

[183]

691.0

[166]

691.0

[128]

691.0

[167]

875.9

691.0

[168]

812

792

[166]

575.2

[173, 174] [173] [174] [167, 180] [167, 180] [175]

1313

H2O•BF3/N2

1314.6 1283.9(c)

1264.7 1251.9(c)

863

874.5

Ac ce pt e

H2O•BF3/Ar

M

H2O•BF3/Ar

d

1486.4

2

H2O•BF3/Ne Me2O•BF3/ Ar Me2O•BF3/ Ar Me2O•BF3/ N2 Et2O•BF3/A r (CF3CH2)2O •BF3/Ar N2O•BF3/Ar N2O•BF3/N2 OSO•BF3/A r OSO•BF3/N 2

Me2S•BF3/A r

1251.6 1222.7(c) 1240.1 and 1207.8(c)

1323.0 1303.8(c) 1221, 1260(c) 1219.2 1106.6(c) 1166.9 1094.3(c) 1241.1 1207.6(c)

(1484.2 not obs)(c)

not observed

1486.7 1453.9(c) 1470.3 1454.5(c)c 1485.2 1454.4(c)

1436.0 1400.5(c) 1431.5 1422.1(c) 1433.3 1424.3(c)

1395.1

1443.1 1363.9(c)

ip t

1318

685.3

cr

1216.2 1169.4 1248

us

1266.7 1222.9 1293

an

H3N•BF3/Ar H3N•BF3/N2 MeCN•BF3/ Ar MeCN•BF3/ Ne MeCN•BF3/ N2 MeCN•BF3/ Ar MeCN•BF3/ N2 MeCN•BF3/ Xe C6H5CN•BF 3/Ar Me3CCN•B F3/Ar MeNC•BF3/ Ar

874.5

691.5

460 (385 in origina l) 664.0

815.1

811.6

637.8

624.1

792

817.1

812.0

652.5

637.1

792

796.8

762.0

828.4 694.9

668.0

865.1

865.1

667.8

642.2

575.2

[175]

875.9

875.9

649.2

634.6

672.3

[179]

863.6

863.6

651.8

636.7

672.3

[179]

837.7

837.7

688.2

668.2

830.9

[173] [174]

128 Page 128 of 213

2

H2S•BF3/Ar

1424.1 1410.1(c)

1459.1 (1380.4 1378.4)(c) 1270 1321(c)

Me2CHF•B F3/Ne

1434 1415(c)

BF4

1197

(a)

853.5

853.5

690.5

670.2

830.9

[173] [174]

705

[166]

610

[181] [820]

Proton affinities taken from [71, 72]. Doublets due to trapping/matrix sites. (c) Doublets due to lifting of degeneracy of parent E mode. (d) Mixed B-N str and BF3 symmetric stretch (e) Formally IR inactive, activated by matrix effects or complexation.

Ac ce pt e

d

M

an

us

cr

(b)

ip t

Me2S•BF3/N

129 Page 129 of 213

us

853.6 792

[145] [166] [166]

an

399

Ac ce pt e

d

M

BBr3/Ar 849.0 812.8 H3N•BBr3/Ar 839 804 Me2O•BBr3/Ar 722/680 (a) Proton affinities taken from [71, 72].

cr

ip t

Table 5 Table of boron chloride and bromide Lewis base complexes in cryogenic matrices Complex/matrix Band position /cm−1 Proton Ref affinity of base / kJ mol−1 (a) BX3 asym str BX3 sym BX3 sym str def 10 11 10 B B B 11B 10B 11B BCl3/Ar 995 956 [166] H3N•BCl3/Ar 853 836 506 506 853.6 [166] Me2O•BCl3/Ar 815/777 415 792 [166]

130 Page 130 of 213

Table 6. νCN modes in BF3 nitrile complexes in cryogenic matrices Ne Com plex F311B•NCMe F311B•NCMe F310B•NCMe F311B•NCCD3 F310B•NCCD3 F311B•15NCMe F310B•15NCMe F311B•NCMe

Mo nom er

Ar Complex

Monom er

Xe Comple x

Monom er

N2 Comple x

2258.3

2357.6

2262.3

F310B•NCMe

2365.7

2258.3

2357.6

2262.3

2367.0

F311B•NCCD3

2362.0

2267.0

2356.1

2269.9

2365.3

F310B•NCCD3

2361.9

2267.0

2355.4

2269.9

2365.2

2320

[156] [156] [156] [152] [160] [160]

M

2323/2327

2257. 2 2257. 2 2266. 4 2266. 4

cr

us

2368

an

2352 2356

2367.0

Ac ce pt e

d

F311B•NCMe F310B•NCMe F311B•NCC6H5 F310B•NCC6H5 F311B•NCCMe3 F310B•NCCMe3

[150] [151] [151] [151] [151] [151] [151] [156]

ip t

2380 2365 2365 2363 2363 2339 2339 2365.6

Ref Mono mer

131 Page 131 of 213

Table 7 BF3 complexes in cryosolutions Base Band position /cm−1 BF3 asym str BF3 sym str BF3 sym def 10 11 10 10 11 B B B 11B B B 679.5 0

[127, 198]

1493. 1

1441.9

697.9

671.0

[198]

690.1

664.0

[198]

691.2

664.8

[198]

656.2 7

[127]

642.6 0

[127]

1437.9 5

1486. 11

1435.0 5

1481

cr

1488. 97

us

1442.8

an

1493. 4

ip t

707.0 2

[200]

Ac ce pt e

(liquid argon) OCS•BF3 (liquid Ar)

1443.7 5

M

3

1495. 00

d

BF3 (liquid Ar) N2•BF3 (liquid Ar) (N2)2•BF3 (liquid Ar) N2•BF3 (liquid N2) OC•BF3 (liquid Ar) (OC)2•BF

Ref

132 Page 132 of 213

853.6 948.9

[320] [320]

H3P•AlMe3 Me3P•AlMe3

662 668

785 958.8

[325] [325]

Me2O•AlMe3 Me2CO•AlMe3

682 697, 661

792 812.0

[325] [325]

Me2S•AlMe3

not reported

GaMe3

575

F3N•GaMe3 H3N•GaMe3 Me3N•GaMe3

570 559 543

an

[325]

cr

624 623

us

D3N•AlMe3 Me3N•AlMe3

ip t

Table 8. AlC3, GaC3, InC3 asymmetric stretching mode in complexes in argon matrices. Ar matrix Proton Ref affinity /kJ /cm−1 mol−1 (a) AlMe3 688 [325]

Ac ce pt e

d

604 853.6 948.9

M

[342, 821]

[342] [333] [85]

F3P•GaMe3 H3P•GaMe3 Me3P•GaMe3

569 561 549

695.3 785 958.8

[342] [342] [342]

H3As•GaMe3 H3Sb•GaMe3

565 566

747.9 731

[335] [336] [342]

Me(H)O•GaMe3 Me2O•GaMe3

566 561

754.3 792

[343] [344]

H2S•GaMe3 Me(H)S•GaMe3 Me2S•GaMe3

566 562 561

705 773.4 830.9

[343] [343] [344]

Me2Se•GaMe3

561

[344]

InMe3 H3N•InMe3

505 497

[349] [353] [350]

853.6

133 Page 133 of 213

[349] [350] [349, 350] [349] [350]

Ac ce pt e

d

M

an

us

cr

ip t

H3P•InMe3 497 785 H3P•InMe3 498 785 H3As•InMe3 499 747.9 H3Sb•InMe3 499 731 H3Sb•InMe3 498 731 (a) Proton affinities taken from [71, 72].

134 Page 134 of 213

Table 9

Me3N•SiF4/Ar C5H5N•SiF4/Ar

957 954 950 942 934 940 931 938 932 945 955

HCN•SiF4/Ar MeCN•SiF4/Ar

994 986

MeH2N•SiF4/Ar MeH2N•SiF4/N2

493.8 853.6 853.6

us

854 838 835 832 845

[408] [361] [361]

899.0

[363]

899.0

[363]

836

929.5

[363]

847 837

948.9 930.0

[363] [376]

970 974

712.9 779.2

[377] [377]

841

Ac ce pt e

Me2HN•SiF4/Ar

νSi-F (ax) /cm−1

an

(b)

[358]

M

N2•SiF4/N2 H3N•SiF4/Ar H3N•SiF4/N2

1029.4 1005.2 νSi-F (eq) /cm−1 1022

d

(SiF4)2/Ar

cr

ip t

Table 9. νSi-F, modes of SiF4 halide Lewis acid base complexes in argon matrices. Complex/matrix νSi-F /cm−1 Proton affinity Ref of base /kJ mol−1(a) SiF4/Ar 1025 [389] SiF4/Ar 1023.0 [358] SiF4/N2 1022.3 [358] SiF4/Xe 1022.7 [358]

H2O•SiF4 984 691.0 Me(H)O•SiF4 973 852 754.3 Me2O•SiF4 969 962 792 Et2O•SiF4 965 (CH2)2O•SiF4 973 774.2 (CH2)3O•SiF4 964 856 801.3 (CH2)4O•SiF4 964 822.1 (a) Proton affinities taken from [71, 72]. (b) not observed, masked by rotation inversion of NH3

[68] [68] [68] [68] [389] [68, 389] [389]

135 Page 135 of 213

[375]

899.0

[375]

948.9

[375]

HCN•GeF4/Ar

712.9

[377]

C5H5N•GeF4/Ar H2O•GeF4

Me(H)O•GeF4 Me2O•GeF4

Me2CO•GeF4

(CH2)2O•GeF4 (CH2)3O•GeF4 (CH2)4O•GeF4 (CH)4O•GeF4 [GeF5]−

an [377]

930.0

[376]

691.0

[388]

712.9

[388]

754.3

[388]

792.0

[388]

812.2

[388]

774.2

[389]

801.3

[389]

822.1

[389]

803.4

[389]

Ac ce pt e

D2CO•GeF4

779 737 768 750 760 722 759 718 747 715 751 722 755 714 755 712 768

779.2

M

MeCN•GeF4/Ar

775 698 673 770 719 679 729 661

730 696 (a) Proton affinities taken from [71, 72].

cr

853.6

us

738 718 MeH2N•GeF4/Ar 712 727 697 Me3N•GeF4/Ar 715 700

d

H3N•GeF4/Ar

ip t

Table 10. νGe-F modes of GeF4 Lewis acid base complexes in argon matrices. νGe-F /cm−1 Proton affinity Ref of base /kJ mol−1 (a) GeF4/Ar 795 [375, 392] GeF4/Ar 800 [389]

[392]

136 Page 136 of 213

Table 11

1352 Not identified with any certainty

H3N•AlMe3 H3N•GaMe3 H3N•InMe3

1212 1203 1119

[255] [137]

Ac ce pt e

H3N•BCl3 H3N•BBr3

us

1343 1309

an

Group 13 H3N•BH3 H3N•BF3

[88] [103]

M

1086.8 1044.4

d

Group 12 H3N•ZnMe2 H3N•CdMe2

cr

ip t

Table 11. ν2 mode (umbrella mode) of NH3 in Lewis acid base complexes in cryogenic matrices. Ar /cm−1 N2 /cm−1 Ref NH3 956 Q(1+1) [456] + 962 P(1 0) 970 R(0-0) NH3 970 [375] NH3 974.3 [466, 479]

[137] [137]

[320] [333] [350]

Group 14 H3N•CO H3N•OC H3N•CO2 (H3N)2•CO2

997 1028 1003

[481] [481] [479] [479]

NH3•SiF4 H3N•GeF4 H3N•SiCl4 H3N•GeCl4

1253 1245 1289 1210

[361] [375] [361] [375]

H3N•SiH4

982 Q(1+1) 984 P(1+0) 1004 R(00)

H3N•Si H3N•Ge H3N•SiO2 H3N•SiCl3H

1185.5 1128.7 1277.3 1003

994

[399]

[430] [430] [429] [396] 137 Page 137 of 213

Group 16 H3N•SO3 H3N•SO2

1317 1038

H3N•SF4 H3N•SOF2 H3N•SOF2 Group 17 H3N•HF H3N•F2 H3N•Cl2 H3N•ClF

1093 966.2 988 1054 1047

1050 1044 not reported

[682] [682] [682]

1022

[459] [721] [736] [713]

ip t

[12] [566]

cr

985.7

us

H3N•O3

[70] [623]

NH3 NH3

M

NH3

N2 /cm−1 Ref [456] 970

Ac ce pt e

974.3

[375] [466] [479]

d

Ar /cm−1 956 Q(1+1) 962 P(1+0) 970 R(0-0)

an

Table 12. ν2 mode (umbrella mode) of NH3 in Lewis acid base complexes in cryogenic matrices in rank order.

H3N•F2 H3N•SiH4

966.2 982 Q(1+1) 994 984 P(1+0) 1004 R(0-0) H3N•O3 985.7 H3N•Cl2 988 H3N•CO2 997 H3N•SiCl3H 1003 (H3N)2•CO2 1028 1003 H3N•SO2 1038 H3N•SOF2 1044 H3N•CdMe2 1044.4 H3N•SF4 1050 H3N•ClF 1054 1022 1047 H3N•ZnMe2 1086.8 H3N•HF 1093 H3N•InMe3 1119

[721] [399]

[12] [566] [736] [479] [396] [479] [612, 623] [682] [103] [682] [713] [88] [459] [350] 138 Page 138 of 213

1343 1352

ip t

1317

[430] [430] [333] [375] [320] [375] [361] [429] [361] [137] [70] [255] [137]

cr

1128.7 1185.5 1203 1210 1212 1245 1253 1277.3 1289 1309

Ac ce pt e

d

M

an

us

H3N•Ge H3N•Si H3N•GaMe3 H3N•GeCl4 H3N•AlMe3 H3N•GeF4 NH3•SiF4 H3N•SiO2 H3N•SiCl4 H3N•11BF3 H3N•SO3 H3N•11BH3 H3N•11BCl3

139 Page 139 of 213

[88] [103]

H3P•AlMe3 H3P•GaMe3 H3P•InMe3

979 976 972

[325] [103] [349]

H3P•O3

988.5 986.3

[12] [652]

an

H3P•SO3

cr

983.0, 981.5 989, 983*

us

H3P•ZnMe2 H3P•CdMe2

ip t

Table 13. Umbrella mode (ν2) of PH3, AsH3, SbH3 and in their complexes in argon matrices. Ar /cm−1 Ref PH3 994 [12, 103, 348] PH3 988/994 [742] PH3 992 [325]

986 981 981

[348] [742] [742]

AsH3 AsH3

912 912.0, 907.5, 905.1

[348] [88]

H3As•ZnMe2 H3As•CdMe2

903.3, 900.9 903

H3As•AlMe3 H3As•GaMe3 H3As•InMe3

893 890, 893

[335] [336] [349] [568]

H3As•HF H3As•ClF H3As•Cl2

905.7 902.7 902 898,890 898, 896, 892

[348] [704] [704]

SbH3 SbH3 SbH3

797/780 779.2 779

[348] [574] [349]

H3Sb•AlMe3 H3Sb•GaMe3 H3Sb•InMe3 H3Sb•InMe3 H3Sb•O3

not reported 769 769 786.7

[342] [349] [349] [574]

d

M

H3P•HF H3P•Cl2 H3P•ClF

Ac ce pt e

[88] [103]

H3As•O3

140 Page 140 of 213

[348]

Ac ce pt e

d

M

an

us

cr

ip t

H3Sb•HF 777 *ambiguity in data presented in paper.

141 Page 141 of 213

Table 14

1816 1814 1816 1820 1827 1823

576 575 570 566 569

[482] [482] [482] [482]

us

C6Me6 C6F6 C6H5Cl C6H5Br C6H5Me C6H6

312 312

[482] [482] [482] [482] [482] [482]

an

565 548 549

315 315 315 312 312

M

1840 1850 1822 1821

Ac ce pt e

d

Et2O Me2CO Me2S Me2Se

cr

ip t

Table 14 Vibrational modes observed for ClNO and its complexes in argon matrices νNO /cm−1 νNCl /cm−1 δClNO /cm−1 Ref ClNO/Ar 1804.95 585.35 [329] 584.38 (ClNO)2/Ar 1842.5 [329] ClNO/N2 1830.7 575.8-575.0 316.7-313.0 [335] (ClNO)2/N2 1865.5 568.2 312.5 [335] ClNO/Ar 1805 585 321 [482]

142 Page 142 of 213

Table 15

594.0

[565]

853.6

[566]

853.6

[552]

H3N•O3/N2

1047.6 1049.1 1048.4

H3P•O3/Ar F3P•O3/Ar Cl3P•O3/Ar Br3P•O3/Ar

1037.2 not reported 1032 1032

785 695.3

[12] [577, 578] [578] [578]

H3As•O3/Ar F3As•O3/Ar

1035.5 not reported

747.9

[568] [577]

H3Sb•O3/Ar

1032.0

731

[574]

H2O•O3/N2 H2CO•O3/N2 NO2•O3/Ar

1047.4 1047.2 1040.4 1038.6 1036.4 1034.0 1040

691.0 712.9 591.0

[552] [552] [564]

591.0

[564]

628.5

[683]

Ac ce pt e

H3N•O3/Ar

1042.56 1043.37

d

Ozone as an acid OC•O3/Ar

M

an

us

cr

ip t

Table 15 Ozone ν3 and ν1 modes in Lewis acid-base complexes in argon matrices. Complex/matrix Ar matrix Proton Ref gas affinity /kJ mol−1 (a) ν3 /cm−1 ν1 /cm−1 ν2 /cm−1 O3/Ar 1041.28, [565] 1039.61 O3/Ar 1040.0 1105 [822] O3/Ar 1041.2 1105.1 704.2 [823] 1039.7 703.6 O3/N2 1043.0 1108.4 704.3 [823] O3/N2 1042.8 1108 704.3 [552] O3/Ar 1040.9 [564] 1039.4 O3/Kr 1035.8 [564] 1034.6 O3/N2 1042.9 [564]

NO2•O3/Ar SCO•O3/Ar

706.6

706.4 706.2

143 Page 143 of 213

H2CCH2•O3/N2

1175 ?

O3•SO2/Ar

1039.8

O3•Cl2/Ar O3•Br2/Ar O3•BrCl/Ar O3ICl/Ar O3ICl/Kr

1032.1 1030.0 1028.0 1019.5 1011.3

O3ICl/N2

1023.0

O3•ICN/Ar

1033.9, 1028.6

O3•IMe/Ar O3•ICF3/Ar O3•IH2Me/Ar

1032 1040.8, 1039.5, 1037.2, 1035.8, 1034.4, 1033.7, 1032.5, 1031.7

[552]

680.5

[564]

680.5 680.5

[564] [564]

641.4

[564] [592]

M

an

us

cr

ip t

1100.8

Ac ce pt e

Ozone as a base O3•ZnMe2/Ar

680.5

1110.8 1107.0 1104.2 1110.0

1102.8 1106 1100.8

[595] [594] [597] [597] [52]

d

1041.5 1040.2 H2CCH2•O3/N2 1040.6 1040.2 H2CCH2•O3/Ar 1038.5 H2CCH2•O3/Kr 1033.9 1032.4 HCCH•O3/Ar 1037.9 C2H5I•O3/Ar 1040.8 1039.5 1037.2 1035.8 1034.4 1033.7 1032.5 1031.7 CH2ClI•O3/Ar 1035.7 1033.3 C2F5I•O3/Ar 1039.5 1033.8 1032.0 BrHCCHBr•O3/Ar 1033.5 1031.8 ClHCCHCl•O3/Ar 1034.8

676

[630] [586] [586] [586] [589] [589] [589] [590] [591] [548] [592]

144 Page 144 of 213

709.1

[552]

Ac ce pt e

d

M

an

us

cr

ip t

O3•HBr/N2 1014.0 (a) Proton affinities taken from [71, 72].

145 Page 145 of 213

ip t

Table 16 Values of the O-E-O symmetric stretch in Me2E (E =O, S, Se) and their complexes in argon matrices. Complex O-C-O Ref symmetric stretch /cm−1 Me2O 926 [68] Me2O 925 [344] Me2O 924 [89] 912.3 912 895 901

[89] [104] [325] [344]

Me2O•11BF3 Me2O•11BF3 Me2O•11BCl3 Me2O•11BBr3

≈950(a) 950.2 903 887

[166] [173, 174] [166] [166]

Me2O•SiF4 Me2O•SiHCl3 Me2O•GeF4

917 900 892

[68] [396] [388]

Me2O•ClF

904

[69]

S-C-S symmetric stretch /cm−1 694 695 693

[344] [710] [173, 174] [325]

Me2S•ZnMe2 Me2S•CdMe2 Me2S•AlMe3 Me2S•GaMe3 Me2S•11BF3

not reported not reported 678 687 not observed

[89] [104] [325] [344] [173, 174]

Me2S•ONCl Me2S•SO2 Me2S•SO3

not reported 691 not reported

[482] [638] [652]

Me2S•Cl2 Me2S•Cl2 Me2S•ClF Me2S•Br2

not reported 694 not reported 694

[702] [710] [702] [710]

us an

M

d

Ac ce pt e

Me2S Me2S Me2S

cr

Me2O•ZnMe2 Me2O•CdMe2 Me2O•AlMe3 Me2O•GaMe3

146 Page 146 of 213

Ac ce pt e

d

M

an

us

cr

ip t

Se-C-Se symmetric stretch /cm−1 Me2Se 652 [344] Me2S•ZnMe2 not reported [89] Me2S•CdMe2 not reported [104] Me2Se•GaMe3 644 [344] Me2Se•ONCl not reported [482] Me2Se•ClF not reported [702] Me2Se•Cl2 not reported [702] (a) Obscured by BF3bands, estimated from deuterium enrichment experiments.

147 Page 147 of 213

Table 17 Table 17. SO2 asymmetric (ν3) and symmetric (ν1) stretching modes in SO2 and its complexes in cryogenic matrices. Complex Ar matrix /cm−1 N2 matrix / cm−1 Proton Ref affinity / kJ mol

SO2 SO2 SO2

1342

H2O•SO2

1152

cr

ν1

1153.8

[644] [623] [637] [634] [628] [630] [638] [609]

1155.8

853.6

[612]

1149

853.6

1121 1120 1156

899.0 929.5 948.9 779.2

[623] [623] [623] [623] [623] [625]

1155.8

691.0

[612]

d

Ac ce pt e

H3N•SO2 (H3N)2•SO2 MeH2N•SO2 Me2HN•SO2 Me3N•SO2 MeCN•SO2

1350 1146 1152.2 1153.2(b) 1352.5 1148.5(c) 1152.3(b) 1147.2(c) 1152 1152.3 1151.7 1337.5 1334.5 1338 1320 1309(d) 1277(d) 1270(e) 1148 1345

M

H3N•SO2

ν3

us

1356/1352 1355.0 1356(b) 1352.2(c) 1355.0(b) 1351.3(c) 1356 1355.5 1355.0

ν1 1146

an

ν3 1347

SO2 SO2 SO2 SO2 SO2

ip t

−1 (a)

1346.6 1344.7

H2O•SO2

1343

1150

691.0

[631, 639]

H2O•SO2 (H2O)2•SO2 (H2O)3•SO2

1343.5 1342.4 1337.6

1150.0 1157.2 1160.2

691.0

[634] [634] [634]

O3•SO2

1352.4 1349.5 1330.1

1150.2 1149.6 1140.4

625.5

[630]

1346 1308 1308 1309 1340 1337

1146 1129 1128 1130 1145 1142(e)

HOO•SO2 S•SO2 Me2S•SO2 Et2S•SO2 MeEtS•SO2 Et(H)S•SO2

[636] 830.9 856.7 789.6

[637] [638] [638] [638] [639]

148 Page 148 of 213

1322

HSCH2CH2(H)S•SO2 MeS(Me)S•SO2

1142

[639]

1144

[640]

1144 1142 1135(e) 1134 1126 1121(e) 1145 1143(e)

[640]

Cl2•SO2

ip t

TTP(f)•SO2

C4H9(H)S•SO2

[639]

cr

1338

t

1144 1140(e)

1345.9 1347.7 1346.7

HBr•SO2

1347.9

1150.4 1149.0

1150.3

[640] [640] [642] [627] [612] [612]

[612] [644] [644]

Ac ce pt e

d

M

C2H4•SO2 C6H6•SO2 1337 1142 C6H5Me•SO2 1335 1142 (a) Proton affinities taken from [71, 72]. (e) Predicted from 18O studies. (b) stable crystal site (ccp) (c) unstable crystal site (hcp) (d) average value from matrix split sites. (e) matrix site splittings (f) TTP = 1,4,8,11-tetrathiacyclotetradecane

us

Me(H)S•SO2

2-C3H7(H)S•SO2

an

(Me)SCH2(H)S•SO2 HSCH2(Me)S•SO2

1323(e) 1342 1338 1325(e) 1342 1336(e) 1339 1337(e) 1338 1335 1329(e) 1341 1314

149 Page 149 of 213

Table 18 Table 18. SO3 asymmetric stretching mode in SO3 and its complexes in cryogenic matrices Complex Ne Ar Xe N2 O2 Proton Ref /cm−1 /cm−1 /cm−1 /cm−1 /cm−1 affinity /kJ mol−1 SO3 SO3 SO3 SO3 SO3

1390.8

1385.9

1383.1 1398

cr

1385.8

us

1385.0 1385.1

H3N•SO3 MeH2N•SO3 Me2HN•SO3 H5C5N•SO3 Me3N•SO3 N2•SO3

1393

Ac ce pt e

d

F3P•SO3 H3P•SO3 Me3P•SO3 H2O•SO3 H2O•SO3 H2O•SO3 H2O•SO3 H2O•SO3

Me2O•SO3

H3C(H)O•SO3 Me2CO•SO3 H2CO•SO3 (CH2)2O•SO3 (CH)4O•SO3 (CH2)4O•SO3

[648] [70] [534] [649] [670] [666]

an

1396.7 1394.6 1354 1348 1346 1338 1330

853.6 899.0 929.5 930.0 948.9

[70] [70] [70] [70] [70] [667]

1390 1323 1281

695.3 785 958.8

[652] [652] [652]

691.0 691.0 691.0

[648] [534] [649] [629]

M

OC•SO3

ip t

(a)

1401.4

1397.0

1390.6

1399 1390

1396.5 1392.4

[670] 1362, 1378(b) 1369, 1387(b)

(c)

1362, 1372(b) 1359, 1383(b) 1370 1356, 1377

792

[651]

754.3

[651]

812.0 712.9

[651] [651]

774.2

[651]

803.4 822.1

[651] [651]

150 Page 150 of 213

OCO•SO3

1398.2 1395.3

[665]

H2S•SO3

1377, 1380 (b) 1301, 1335 (b)

Me2S•SO3 (a)

Proton affinities taken from [71, 72]. splitting due to loss of degeneracy, (c) bands obscured

705

[652]

830.9

[652]

Ac ce pt e

d

M

an

us

cr

ip t

(b)

151 Page 151 of 213

Table 19 Table 19. νS-F for SF4 and its complexes in nitrogen matrices N2 /cm−1 864 830

699 675 662 662

841

853.6 899.0 930.0

[682] [682] [682] [682]

680

812.0

[682]

us

Me2CO•SF4 (a) Proton affinities taken from [71, 72].

549

ip t

888 841

Ref

cr

SF4 H3N•SF4 MeH2N•SF4 H5C5N•SF4

Proton affinity /kJ mol−1(a)

Table 20

Ac ce pt e

d

M

an

Table 20. νS-F and νS-O for SOF2 and SO2F2 and their complexes in nitrogen matrices. N2 /cm−1 Proton Ref affinity /kJ mol−1 (a) νS-O νS-O νS-F νS-F SOF2 1331 796 731 [682] H3N•SOF2 1277 780 710 853.6 [682] MeH2N•SOF2 1265 770 696 899.0 [682] H5C5N•SOF2 773 930.0 [682] SO2F2 1493 1267 H3N•SO2F2 (a) Proton affinities taken from [71, 72].

883 879

849 841

853.6

[682] [682]

152 Page 152 of 213

Table 21

948.9 853.6 779.2

[713] [713] [713] [713]

H3P•ClF Me3P•ClF

406 311

785 958.8

[742] [742]

H3As•ClF

465

747.9

[704]

H2CO•ClF H3C(H)O•ClF (CH2)2O•ClF Me2O•ClF Me2CO•ClF Et2O•ClF (CH)4O•ClF

716, 710 697, 693 692, 686 682 687, 682 673 657, 652

712.9 754.3 774.2 792 812.0 828.4 803.4

an

M

d

Ac ce pt e

cr

505 600 626 680

us

Me3N•ClF H3N•ClF H2CCHCN•ClF MeCN•ClF

ip t

Table 21 . νClF and νClCl modes in ClF and Cl2 Lewis acid base complexes in argon matrices. Proton affinity / Ref Complex νClF /cm−1 (b) kJ mol−1 (a) ClF 768, 762 [69]

[69] [69] [69] [69] [69] [69] [69]

18-crown-6•ClF TTP•ClF cyclam•ClF TMC•ClF

673 468 458 476

[756] [627] [627] [627]

H2S•ClF H3C(H)S•ClF Me2S•ClF

631 536 471

705 773.4 830.9

[702] [702] [702]

H2Se•ClF Me2Se•ClF

580 448

707.8

[704] [704]

Me3CCl•ClF

713, 706

[758]

c-C3H6•ClF c-MeC3H6•ClF c-Me2C3H6•ClF c-BrC3H5•ClF

739, 733 735, 729 699, 693 690,684

[759] [759] [759] [759]

2-butyne•ClF

636

[762] 153 Page 153 of 213

641.4 748 775.8 680.5 751.6 747 747 750.4

Complex

νClCl /cm−1

H3P•Cl2

512

785

[742]

H3As•Cl2

519

747.9

[704]

H2S•Cl2 H3C(H)S•Cl2 Me2S•Cl2 Me2S•Cl2

516 472 360 525

705 773.4 830.9 830.9

H2Se•Cl2 Me2Se•Cl2

512 295

707.8

428 434 394

d

M

an

us

Ref

Ac ce pt e

c-C3H6•Cl2 c-MeC3H6•Cl2 c-Me2C3H6•Cl2

[762] [762] [762] [762] [762] [762] [762] [768] [768] [768] [768] [768] [768]

ip t

715 688 636 662 609 571 586 711, 705 705 720 710 652 700

cr

C2H2•ClF C3H4•ClF C4H6•ClF C2H4•ClF C3H6•ClF cis-C4H8•ClF iso-C4H8•ClF C6H6•ClF C6H5CH3•ClF C6H5Br•ClF C6H5Br•ClF C6Me6•ClF C3H5C6H5•ClF

[702] [702] [702] [703] [704] [704] [759] [759] [759]

Complex c-MeC3H6•Br2 285 [759] c-Me2C3H6•Br2 260 [759] (a) Proton affinities taken from [71, 72]. (b) Where two values are quoted, the lower one is due to the 37Cl component. TTP = 1,4,8,11-tetrathiacyclotetradecane, TMC = tetramethylcyclam

154 Page 154 of 213

ICl [778] [741]

C2H4 C2H3F

thiophene C6H6 allene buta-1,3-diene Ar

[828]

[831] [747]

[835] [838]

[836] [839]

[787] [760] [761] [763, 764, 766] [763, 765, 766] [857]

[788]

[785]

[850]

[851]

[855]

[856]

[861]

[862]

us

[750, 834] [707] [840] [841] [705] [706] [843] [844] [845]

an

[833] [837]

M

[842]

[785, 846]

Ac ce pt e

CO cyclopropane, methylenecyclopropane C2H2

[749, 750] [832] [754] [708]

[786]

d

H2O H2S furan 2,5-dihydrofuran H2CO (CH2)2O CO2 SO2 (CH2)2S

cr

ip t

Table 22 Halogen and interhalogen complexes studied by microwave spectroscopy F2 Cl2 Br2 ClF BrCl Review [774] [775] [776] [777] [766] Lewis base NH3 [727] [737, 738] [739] [716, 717] [740] Me3N [715, 824] [714, 715] NF3 [717] HCN [825] [826] [719] [827] MeCN [829] [718] N2 [830] PH3 [744] [745] [746]

[847, 848]

[849]

[852, 853]

[854]

[858] [770] [859] [860]

155 Page 155 of 213

Table 23

2154.1 and 2150.8 2145.5 and 2143.2 2045.7

CO•BrI CO•ClI CO•IBr CO•ICl CO•BrF

2136.5 2136.5 2129.7 2128.4 2127.7

an

OC•IBr OC•BrI BrCO•

[792] [791] [791] [790, 793] [790, 793]

M

2158.4 2156.3 2138.2 2157.0 and 2154.0 2145.6 and 2144.2

Ac ce pt e

d

OC•ClF OC•IF OC•FI OC•ICl OC•ClI

us

cr

ip t

Table 23 νCO modes of complexes between CO and the halogens and interhalogens in argon matrices. Ref Complex νCO /cm−1 CO 2138.5 [789] OC•Cl2 2140.6 [789] OC•Cl2 2140.7 [790, 793] OC•Br2 2143.9 [789] OC•Br2 2144.1 [793] OC•ClBr 2138.0 [789] OC•BrCl 2147.9 [789] OC•BrCl 2148.3 [790, 793] OC•BrF 2163.3 [792] OC•FBr 2139.6 and 2135.6 [792]

[790, 793] [790, 793] [793]

[790] [790, 793] [790] [790, 793] [792]

156 Page 156 of 213

References [1] A. Haaland, Angew. Chem. Int. Ed. Engl., 28 (1989) 992-1007.

ip t

[2] R.S. Mulliken, J. Phys. Chem., 56 (1952) 801-822.

cr

[3] R.S. Mulliken, J. Am. Chem. Soc., 74 (1952) 811-824.

us

[4] A.E. Reed, L.A. Curtiss, F. Weinhold, Chem. Rev., 88 (1988) 899-926.

[5] V. Jonas, G. Frenking, J. Chem. Soc., Chem. Commun., (1994) 1489-1490.

an

[6] V. Jonas, G. Frenking, M.T. Reetz, J. Am. Chem. Soc., 116 (1994) 8741-8753.

254 (2010) 2031-2077.

M

[7] E.I. Davydova, T.N. Sevastianova, A.V. Suvorov, A.Y. Timoshkin, Coord. Chem. Rev.,

d

[8] A.Y. Timoshkin, H.F. Schaefer, Journal of Physical Chemistry C, 112 (2008) 13816-

Ac ce pt e

13836.

[9] E. Whittle, D.A. Dows, G.C. Pimentel, J. Chem. Phys., 22 (1954) 1943-1943. [10] B.S. Ault, Rev. Chem. Intermed., 9 (1988) 233-269. [11] M. Moskovits, G.A. Ozin, J. Chem. Phys., 58 (1973) 1251-1252. [12] R. Withnall, M. Hawkins, L. Andrews, J. Phys. Chem., 90 (1986) 575-579. [13] S. Tam, M.E. Fajardo, Rev. Sci. Instrum., 70 (1999) 1926-1932. [14] M. Bahou, C.W. Huang, Y.L. Huang, J. Glatthaar, Y.P. Lee, J. Chin. Chem. Soc., 57 (2010) 771-782. [15] Y.F. Lee, Y.P. Lee, J. Chem. Phys., 134 (2011) 124314. 157 Page 157 of 213

[16] M.E. Fajardo, Matrix isolation spectroscopy in solid parahydrogen: a primer, in: L. Khriachtchev (Ed.) Physics and Chemistry at Low Temperatures, Pan Stanford Publishing, Singapore, 2011. [17] B. Mile, P.D. Sillman, A.R. Yacob, J.A. Howard, J. Chem. Soc., Dalton Trans., (1996)

ip t

653-663. [18] H.A. Joly, T. Newton, M. Myre, Phys. Chem. Chem. Phys., 14 (2012) 367-374.

us

cr

[19] F.D. Brunet, H.A. Joly, J. Phys. Chem. A, 116 (2012) 4267-4273.

[20] F.D. Brunet, J.C. Feola, H.A. Joly, J. Phys. Chem. A, 116 (2012) 2439-2452.

an

[21] H.A. Joly, L. Beaudet, M. Levesque, M. Myre, J. Phys. Chem. A, 115 (2011) 1184111851.

M

[22] W.A. Herrebout, B.J. van der Veken, Cryogenic solutions as a tool to characterize redand blue-shifting C-H...X hydrogen bonding, in: L. Khriachtchev (Ed.) Physics and

Ac ce pt e

d

Chemistry at Low Temperatures, Pan Stanford Publishing, Singapore, 2011. [23] K.S. Rutkowski, W.A. Herrebout, S.M. Melikova, P. Rodziewicz, B.J. van der Veken, A. Koll, Spectrochim. Acta A, 61 (2005) 1595-1602. [24] J.J.J. Dom, B. Michielsen, B.U.W. Maes, W.A. Herrebout, B.J. van der Veken, Chem. Phys. Lett., 469 (2009) 85-89.

[25] W.A. Herrebout, A. Gatin, G.P. Everaert, A.I. Fishman, B.J. van der Veken, Spectrochim. Acta A, 61A (2005) 1431-1444. [26] F. Stienkemeier, A.F. Vilesov, J. Chem. Phys., 115 (2001) 10119-10137. [27] J.P. Toennies, A.F. Vilesov, Angew. Chem. Int. Ed., 43 (2004) 2622-2648. [28] S. Kuma, M.N. Slipchenko, K.E. Kuyanov, T. Momose, A.F. Vilesov, J. Phys. Chem. A, 110 (2006) 10046-10052. 158 Page 158 of 213

[29] S. Kuma, M.N. Slipchenko, T. Momose, A.F. Vilesov, Chem. Phys. Lett., 439 (2007) 265-269. [30] K. Kuyanov-Prozument, D. Skvortsov, M.N. Slipchenko, B.G. Sartakov, A. Vilesov, Matrix isolation spectroscopy in helium droplets, in: L. Khriachtchev (Ed.) Chemistry and

ip t

Physics at Low Temperatures, Pan Stanford Publishing, Singapore, 2011. [31] M. Ito, T. Ebata, N. Mikami, Annu. Rev. Phys. Chem., 39 (1988) 123-147.

cr

[32] M. Herman, R. Georges, M. Hepp, D. Hurtmans, International Reviews In Physical

an

[33] D.J. Nesbitt, Chem. Rev., 112 (2012) 5062-5072.

us

Chemistry, 19 (2000) 277-325.

[34] Y. Futami, S. Kudoh, M. Takayanagi, M. Nakata, Chem. Phys. Lett., 357 (2002) 209-

M

216.

d

[35] F. Ito, J. Chem. Phys., 137 (2012) 014505.

Ac ce pt e

[36] P. Gilli, L. Pretto, V. Bertolasi, G. Gilli, Acc. Chem. Res., 42 (2009) 33-44. [37] R.N. Perutz, Chem. Rev., 85 (1985) 77-96. [38] H.M. Himmel, A.J. Downs, T.M. Greene, Chem. Rev., 102 (2002) 4191-4241. [39] L. Andrews, Metal atom reactions to form novel small molecules, in: L. Khriachtchev (Ed.) Physics and Chemistry at Low Temperatures, Pan Stanford Publishing, Singapore, 2011.

[40] I.R. Dunkin, Matrix-Isolation Techniques: A Practical Approach, Oxford University Press, Oxford, 1998. [41] S.F. Parker, C.H.F. Peden, P.H. Barrett, R.G. Pearson, J. Am. Chem. Soc., 106 (1984) 1304-1308.

159 Page 159 of 213

[42] M. van der Heyden, M. Pasternak, G. Langouche, J. Phys. Chem. Solids, 46 (1985) 1221-1226. [43] Y. Yamada, Hyperfine Interact., 139 (2002) 77-85.

ip t

[44] N.A. Young, J. Chem. Soc., Dalton Trans., (1996) 249-251. [45] P. Roubin, S. Varin, C. Crepin, B. GauthierRoy, A.M. Flank, R. Delaunay, M. Pompa,

us

[46] O.M. Wilkin, N.A. Young, J. Synch. Rad., 6 (1999) 204-206.

cr

B. Tremblay, J. Chem. Phys., 109 (1998) 7945-7948.

[47] P. Roubin, S. Varin, C. Crepin, B. Gauthier-Roy, A.M. Flank, P. Lagarde, F. Tenegal,

an

Low Temp. Phys., 26 (2000) 691-698.

[48] A.J. Bridgeman, G. Cavigliasso, N. Harris, N.A. Young, Chem. Phys. Lett., 351 (2002)

M

319-326.

d

[49] I.J. Blackmore, A.J. Bridgeman, N. Harris, M.A. Holdaway, J.F. Rooms, E.L.

Ac ce pt e

Thompson, N.A. Young, Angew. Chem. Int. Ed., 44 (2005) 6746-6750. [50] J.F. Rooms, A.V. Wilson, I. Harvey, A.J. Bridgeman, N.A. Young, Phys. Chem. Chem. Phys., 10 (2008) 4594-4605.

[51] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 95 (1991) 3502-3506. [52] P. Varma, B.S. Ault, J. Phys. Chem. A, 112 (2008) 5613-5620. [53] A.J. Downs, S.C. Peake, Matrix Isolation, in: Molecular Spectroscopy, The Chemical Society, London, 1973, pp. 523-607. [54] B.M. Chadwick, Matrix isolation, in: Molecular Spectroscopy, The Chemical Society, London, 1975, pp. 281-383.

160 Page 160 of 213

[55] B.M. Chadwick, Matrix Isolation, in: Molecular Spectroscopy, The Chemical Society, London, 1979, pp. 72-135. [56] D.W. Ochsner, D.W. Ball, Z.A. Kafafi, A Bibliography of Matrix Isolation Spectroscopy:1985-1997, Naval Research Laboratory, Washington, DC, 1998.

ip t

[57] D.W. Ball, Z.H. Kafafi, L. Fredin, R.H. Hauge, J.L. Margrave, A Bibliography of Matrix Isolation Spectroscopy: 1954 - 1985, Rice University Press, Houston, 1988.

cr

[58] M.J. Almond, A.J. Downs, Spectroscopy of matrix isolated species, John Wiley,

us

Chichester, 1989.

[59] M.J. Almond, R.H. Orrin, Annu. Rep. Prog. Chem., Sect. C, Phys. Chem., 88 (1991) 3-

an

44.

M

[60] M.J. Almond, Annu. Rep. Prog. Chem., Sect. C, Phys. Chem., 93 (1997) 3-55. [61] M.J. Almond, K.S. Wiltshire, Annu. Rep. Prog. Chem., Sect. C, Phys. Chem., 97 (2001)

Ac ce pt e

d

3-60.

[62] M.J. Almond, N. Goldberg, Annu. Rep. Prog. Chem., Sect. C, Phys. Chem., 103 (2007) 79-133.

[63] H.E. Hallam, Vibrational Spectroscopy of Trapped Species, John Wiley & Sons, London, 1973.

[64] S. Craddock, A.J. Hinchcliffe, Matrix Isolation, Cambridge University Press, Cambridge, 1975.

[65] M. Moskovits, G.A. Ozin, Cryochemistry, Wiley-Interscience, New York, 1976. [66] M. Moskovits, L. Andrews, Chemistry and Physics of Matrix Isolated Species, Elsevier, Amsterdam, 1989.

161 Page 161 of 213

[67] L. Khriachtchev, Physics and Chemistry at Low Temperatures, Pan Stanford Publishing, Singapore, 2011. [68] B.S. Ault, J. Am. Chem. Soc., 105 (1983) 5742-5746.

ip t

[69] N.P. Machara, B.S. Ault, Inorg. Chem., 24 (1985) 4251-4254. [70] C.S. Sass, B.S. Ault, J. Phys. Chem., 90 (1986) 1547-1551.

cr

[71] S.G. Lias, J.F. Liebman, R.D. Levin, J. Phys. Chem. Ref. Data, 13 (1984) 695-808.

us

[72] E.P.L. Hunter, S.G. Lias, J. Phys. Chem. Ref. Data, 27 (1998) 413-656.

an

[73] M. Kaupp, H.G. von Schnering, Angew. Chem. Int. Ed. Engl., 32 (1993) 861-863.

M

[74] M. Kaupp, M. Dolg, H. Stoll, H.G. von Schnering, Inorg. Chem., 33 (1994) 2122-2131.

d

[75] W. Liu, R. Franke, M. Dolg, Chem. Phys. Lett., 302 (1999) 231-239.

Ac ce pt e

[76] P. Pyykkö, M. Straka, M. Patzschke, Chem. Comm., (2002) 1728-1729. [77] S. Riedel, M. Straka, M. Kaupp, Phys. Chem. Chem. Phys., 6 (2004) 1122-1127. [78] S. Riedel, M. Straka, M. Kaupp, Chem. Eur. J., 11 (2005) 2743-2755. [79] S. Riedel, M. Kaupp, P. Pyykkö, Inorg. Chem., 47 (2008) 3379-3383. [80] J. Kim, H. Ihee, Y.S. Lee, J. Chem. Phys., 133 (2010) 144309. [81] R.L. Deming, A.L. Allred, A.R. Dahl, A.W. Herlinger, M.O. Kestner, J. Am. Chem. Soc., 98 (1976) 4132-4137. [82] X. Wang, L. Andrews, S. Riedel, M. Kaupp, Angew. Chem. Int. Ed., 46 (2007) 83718375.

162 Page 162 of 213

[83] L. Andrews, X. Wang, Y. Gong, T. Schlöder, S. Riedel, M.J. Franger, Angewandte Chemie International Edition, 51 (2012) 8235-8238. [84] M.J. Almond, M.P. Beer, S.A. Cooke, D.A. Rice, H.M. Yates, J. Mater. Chem., 5 (1995) 853-854.

ip t

[85] M.J. Almond, C.E. Jenkins, D.A. Rice, C.A. Yates, J. Mol. Struct., 222 (1990) 219-233. [86] M. Bochmann, M.A. Chesters, A.P. Coleman, R. Grinter, D.R. Linder, Spectrochim.

us

cr

Acta A, 48 (1992) 1173-1178.

[87] V. Boiadjiev, W.T. Tysoe, Chem. Mater., 10 (1998) 1141-1152.

an

[88] H.B. Bai, B.S. Ault, J. Phys. Chem., 98 (1994) 6082-6088.

M

[89] H.B. Bai, B.S. Ault, J. Phys. Chem., 98 (1994) 10001-10007.

d

[90] H. Bai, B.S. Ault, J. Phys. Chem., 99 (1995) 10492-10497. [91] A. Haaland, J.C. Green, G.S. McGrady, A.J. Downs, E. Gullo, M.J. Lyall, J. Timberlake,

Ac ce pt e

A.V. Tutukin, H.V. Volden, K.-A. Ostby, Dalton Trans., (2003) 4356-4366. [92] N. Maung, J. Mol. Struct. (Theochem), 434 (1998) 255-264. [93] M.J. Almond, M.P. Beer, M.G.B. Drew, D.A. Rice, J. Organomet. Chem., 421 (1991) 129-136.

[94] M.J. Almond, M.P. Beer, M.G.B. Drew, D.A. Rice, Organometallics, 10 (1991) 20722076.

[95] M.J. Almond, M.P. Beer, K. Hagen, D.A. Rice, P.J. Wright, J. Mater. Chem., 1 (1991) 1065-1070. [96] M.J. Almond, D.A. Rice, C.A. Yates, J. Mol. Struct., 268 (1992) 51-60.

163 Page 163 of 213

[97] M.J. Almond, M.P. Beer, S.A. Cooke, D.A. Rice, L.A. Sheridan, Spectrochim. Acta A, 50 (1994) 1919-1925. [98] M.J. Almond, M.P. Beer, P. Heath, C.A. Heyburn, D.A. Rice, L.A. Sheridan, J. Organomet. Chem., 469 (1994) 11-13.

ip t

[99] M.J. Almond, D.A. Rice, L.A. Sheridan, P.J. Craig, G. Stojak, M.C.R. Symons, U.S. Rai, J. Chem. Soc., Farad. Trans., 90 (1994) 3153-3157.

cr

[100] M.J. Almond, S.A. Cooke, D.A. Rice, L.A. Sheridan, J. Phys. Chem., 99 (1995) 14641-

us

14646.

[101] M.J. Almond, B. Cockayne, S.A. Cooke, D.A. Rice, P.C. Smith, P.J. Wright, J. Mater.

an

Chem., 5 (1995) 1351-1355.

Chem., 6 (1996) 1639-1642.

M

[102] M.J. Almond, B. Cockayne, S.A. Cooke, D.A. Rice, P.C. Smith, P.J. Wright, J. Mater.

Ac ce pt e

d

[103] H.B. Bai, B.S. Ault, J. Mol. Struct., 377 (1996) 235-246. [104] H.B. Bai, B.S. Ault, J. Mol. Struct., 384 (1996) 191-202. [105] D. McNally, B.S. Ault, J. Phys. Chem. A, 116 (2012) 1914-1922. [106] A.J. Downs, M.J. Goode, C.R. Pulham, J. Am. Chem. Soc., 111 (1989) 1936-1937. [107] C.R. Pulham, A.J. Downs, M.J. Goode, D.W.H. Rankin, H.E. Robertson, J. Am. Chem. Soc., 113 (1991) 5149-5162.

[108] L. Andrews, X. Wang, Science, 299 (2003) 2049-2052. [109] X. Wang, L. Andrews, S. Tam, M.E. DeRose, M.E. Fajardo, J. Am. Chem. Soc., 125 (2003) 9218-9228. [110] L. Andrews, X. Wang, Angew. Chem. Int. Ed., 43 (2004) 1706-1709. 164 Page 164 of 213

[111] X. Wang, L. Andrews, J. Phys. Chem. A, 108 (2004) 3396-3402. [112] T.J. Tague, L. Andrews, J. Am. Chem. Soc., 116 (1994) 4970-4976. [113] A.J. Downs, H.J. Himmel, L. Manceron, Polyhedron, 21 (2002) 473-488.

ip t

[114] A.J. Downs, C.R. Pulham, Chem. Soc. Rev., 23 (1994) 175-184.

us

[116] J.A.J. Pardoe, A.J. Downs, Chem. Rev., 107 (2007) 2-45.

cr

[115] A.J. Downs, Coord. Chem. Rev., 189 (1999) 59-100.

[117] S. Aldridge, A.J. Downs, The Group 13 Metals Aluminium, Gallium, Indium and

an

Thallium: Chemical Patterns and Peculiarities, Wiley, 2011.

M

[118] P.D. Ownby, J. Solid State Chem., 177 (2004) 466-470.

d

[119] J.M. Bassler, P.L. Timms, J.L. Margrave, J. Chem. Phys., 45 (1966) 2704-2706.

Ac ce pt e

[120] F.M.M. O'Neill, G.A. Yeo, T.A. Ford, J. Mol. Struct., 173 (1988) 337-348. [121] L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 300 (1993) 325-338. [122] L.M. Nxumalo, T.A. Ford, Vib. Spectrosc., 6 (1994) 333-343. [123] J. Gebicki, J. Liang, J. Mol. Struct., 117 (1984) 283-286. [124] K.C. Janda, L.S. Bernstein, J.M. Steed, S.E. Novick, W. Klemperer, J. Am. Chem. Soc., 100 (1978) 8074-8079.

[125] L.M. Nxumalo, T.A. Ford, S. Afr. J. Chem., 48 (1995) 30-38. [126] V.M. Rayon, J.A. Sordo, J. Phys. Chem. A, 101 (1997) 7414-7419. [127] E.J. Sluyts, B.J. van der Veken, J. Am. Chem. Soc., 118 (1996) 440-445.

165 Page 165 of 213

[128] D.G. Evans, G.A. Yeo, T.A. Ford, Farad. Discus. Chem. Soc., (1988) 55-64. [129] C.E. Truscott, B.S. Ault, J. Mol. Struct., 157 (1987) 67-71. [130] S. Geller, R.E. Hughes, J.L. Hoard, Acta Crystallogr., 4 (1951) 380-380.

ip t

[131] S. Geller, J.L. Hoard, Acta Crystallogr., 3 (1950) 121-129.

cr

[132] J.L. Hoard, T.B. Owen, A. Buzzell, O.N. Salmon, Acta Crystallogr., 3 (1950) 130-137.

us

[133] J.L. Hoard, S. Geller, W.M. Cashin, Acta Crystallogr., 4 (1951) 396-398.

an

[134] S. Geller, J.L. Hoard, Acta Crystallogr., 4 (1951) 399-405.

[135] J.L. Hoard, S. Geller, T.B. Owen, Acta Crystallogr., 4 (1951) 405-407.

M

[136] S. Geller, M.E. Milberg, Acta Crystallogr., 4 (1951) 381-381.

d

[137] R.L. Hunt, B.S. Ault, Spectros. Int. J., 1 (1982) 31-44.

Ac ce pt e

[138] F. Gaffoor, T.A. Ford, Spectrochim. Acta A, 71 (2008) 550-558. [139] R.L. Kuczkowski, P.S. Bryan, Inorg. Chem., 10 (1971) 200-201. [140] P. Cassoux, R.L. Kuczkowski, A. Serafini, Inorg. Chem., 16 (1977) 3005-3008. [141] A.C. Legon, H.E. Warner, J. Chem. Soc., Chem. Commun., (1991) 1397-1399. [142] D. Fujiang, P.W. Fowler, A.C. Legon, J. Chem. Soc., Chem. Comm., (1995) 113-114. [143] D.L. Fiacco, Y. Mo, S.W. Hunt, M.E. Ott, A. Roberts, K.R. Leopold, J. Phys. Chem. A, 105 (2001) 484-493. [144] C.T. Kwon, H.A. McGee, Inorg. Chem., 12 (1973) 696-697. [145] A. Moroz, R.L. Sweany, Inorg. Chem., 31 (1992) 5236-5242. 166 Page 166 of 213

[146] A.G. Avent, P.B. Hitchcock, M.F. Lappert, D.S. Liu, G. Mignani, C. Richard, E. Roche, J. Chem. Soc., Chem. Comm., (1995) 855-856. [147] A.C. Testa, Spectrochim. Acta A, 55 (1999) 299-309.

ip t

[148] C. Matthaus, R.A. Wheeler, Spectrochim. Acta A, 57 (2001) 521-534. [149] K.R. Leopold, M. Canagaratna, J.A. Phillips, Acc. Chem. Res., 30 (1997) 57-64.

cr

[150] I.R. Beattie, P.J. Jones, Angew. Chem. Int. Ed. Engl., 35 (1996) 1527-1529.

us

[151] N.P. Wells, J.A. Phillips, J. Phys. Chem. A, 106 (2002) 1518-1523.

an

[152] A.A. Eigner, J.A. Rohde, C.C. Knutson, J.A. Phillips, J. Phys. Chem. B, 111 (2007) 1402-1407.

d

(2005) 722-731.

M

[153] J.A. Phillips, J.A. Halfen, J.P. Wrass, C.C. Knutson, C.J. Cramer, Inorg. Chem., 45

833.

Ac ce pt e

[154] J.A. Phillips, C.J. Cramer, Journal of Chemical Theory and Computation, 1 (2005) 827-

[155] L.M. Nxumalo, M. Andrzejak, T.A. Ford, J. Mol. Struct., 509 (1999) 287-295. [156] R. Hattori, E. Suzuki, K. Shimizu, J. Mol. Struct., 750 (2005) 123-134. [157] D.J. Giesen, J.A. Phillips, J. Phys. Chem. A, 107 (2003) 4009-4018. [158] J.A. Phillips, C.J. Cramer, J. Phys. Chem. B, 111 (2007) 1408-1415. [159] M.A. Dvorak, R.S. Ford, R.D. Suenram, F.J. Lovas, K.R. Leopold, J. Am. Chem. Soc., 114 (1992) 108-115. [160] J.A. Phillips, D.J. Giesen, N.P. Wells, J.A. Halfen, C.C. Knutson, J.P. Wrass, J. Phys. Chem. A, 109 (2005) 8199-8208. 167 Page 167 of 213

[161] L.M. Nxumalo, M. Andrzejak, T.A. Ford, J. Chem. Inf. Comput. Sci., 36 (1996) 377384. [162] A.E. Shirk, J.S. Shirk, Inorg. Chem., 22 (1983) 72-77.

Inorg. Chem., (2000) 2245-2253.

cr

[164] D. Mootz, M. Steffen, Z. Anorg. Allg. Chem., 483 (1981) 171-180.

ip t

[163] C. Aubauer, G. Engelhardt, T.M. Klapotke, H. Noth, A. Schulz, M. Warchhold, Eur. J.

us

[165] D. Mootz, M. Steffen, Acta Crystallogr. B, 37 (1981) 1110-1112.

an

[166] R.L. Hunt, B.S. Ault, Spectros. Int. J., 1 (1982) 45-61.

[167] D.W. Ball, M.J. Zehe, NASA Technical Memorandum, (1993) 106422.

M

[168] M.E. Jacox, K.K. Irikura, W.E. Thompson, J. Chem. Phys., 113 (2000) 5705-5715.

d

[169] G.A. Yeo, T.A. Ford, S. Afr. J. Chem., 59 (2006) 129-134.

Ac ce pt e

[170] D.W. Ball, J. Mol. Struct. (Theochem), 331 (1995) 223-228. [171] G.M. Begun, W.H. Fletcher, A.A. Palko, Spectrochimica Acta, 18 (1962) 655-665. [172] M. Weinmann, J. Nuss, M. Jansen, Acta Cryst. E, 62 (2006) O5590-O5591. [173] L.M. Nxumalo, T.A. Ford, Mikrochim. Acta, (1997) 383-385. [174] L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 656 (2003) 303-319. [175] L.M. Nxumalo, T.A. Ford, Spectrochim. Acta A, 53 (1997) 2511-2524. [176] K.R. Leopold, G.T. Fraser, W. Klemperer, J. Am. Chem. Soc., 106 (1984) 897-899. [177] L.M. Nxumalo, G.A. Yeo, T.A. Ford, S. Afr. J. Chem., 51 (1998) 25-34.

168 Page 168 of 213

[178] S.A. Peebles, L. Sun, R.L. Kuczkowski, L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 471 (1998) 235-242. [179] L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 661 (2003) 153-159.

ip t

[180] D.W. Ball, M.J. Zehe, W. Morales, Tribol. Trans., 43 (2000) 767-773. [181] R.R. Knauf, H.M. Helminiak, J.P. Wrass, T.M. Gallert, J.A. Phillips, J. Phys. Org.

us

[182] S.J. David, B.S. Ault, Inorg. Chem., 24 (1985) 1238-1241.

cr

Chem., 25 (2012) 493-501.

an

[183] L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 437 (1997) 69-80.

[184] L.M. Nxumalo, M. Andrzejak, T.A. Ford, Vib. Spectrosc., 12 (1996) 221-235.

M

[185] L.M. Nxumalo, T.A. Ford, J. Mol. Struct. (Theochem), 82 (1991) 135-159.

d

[186] L.M. Nxumalo, T.A. Ford, J. Mol. Struct. (Theochem), 369 (1996) 115-126.

Ac ce pt e

[187] L.M. Nxumalo, G.A. Yeo, T.A. Ford, Theor. Chem. Acc., 96 (1997) 157-165. [188] T.A. Ford, J. Mol. Struct., 834 (2007) 30-41. [189] T.A. Ford, J. Phys. Chem. A, 112 (2008) 7296-7302. [190] T.A. Ford, J. Mol. Struct. (Theochem), 897 (2009) 145-148. [191] T.A. Ford, Int. J. Quantum Chem., 112 (2012) 478-488. [192] G.A. Yeo, T.A. Ford, J. Mol. Struct. (Theochem), 771 (2006) 157-164. [193] L.M. Nxumalo, T.A. Ford, J. Mol. Struct. (Theochem), 357 (1995) 59-65. [194] T.A. Ford, D. Steele, J. Phys. Chem., 100 (1996) 19336-19343.

169 Page 169 of 213

[195] R. Hattori, E. Suzuki, K. Shimizu, J. Mol. Struct., 738 (2005) 165-170. [196] R.C. Taylor, H.S. Gabelnick, K. Aida, R.L. Amster, Inorg. Chem., 8 (1969) 605-612. [197] W.A. Herrebout, G.P. Everaert, B.J. van der Veken, M.O. Bulanin, J. Chem. Phys., 107

ip t

(1997) 8886-8898. [198] W.A. Herrebout, B.J. van der Veken, J. Am. Chem. Soc., 120 (1998) 9921-9929.

cr

[199] B.J. van der Veken, E.J. Sluyts, J. Mol. Struct., 349 (1995) 461-464.

us

[200] B.J. van der Veken, E.J. Sluyts, W.A. Herrebout, J. Mol. Struct., 449 (1998) 219-229.

an

[201] B.J. van der Veken, E.J. Sluyts, J. Phys. Chem. A, 101 (1997) 9070-9076. [202] W.A. Herrebout, A.A. Stolov, B.J. van der Veken, J. Mol. Struct., 563-564 (2001) 221-

M

226.

d

[203] B.J. van der Veken, E.J. Sluyts, J. Am. Chem. Soc., 119 (1997) 11516-11522.

Ac ce pt e

[204] W.A. Herrebout, B.J. van der Veken, J. Am. Chem. Soc., 119 (1997) 10446-10454. [205] G.P. Everaert, W.A. Herrebout, B.J. van der Veken, J. Lundell, M. Räsänen, Chem. Eur. J., 4 (1998) 321-327.

[206] A.A. Stolov, W.A. Herrebout, B.J. van der Veken, J. Am. Chem. Soc., 120 (1998) 7310-7319.

[207] W.A. Herrebout, B.J. van der Veken, Phys. Chem. Chem. Phys., 1 (1999) 3445-3452. [208] W.A. Herrebout, J. Lundell, B.J. van der Veken, J. Phys. Chem. A, 102 (1998) 1017310181. [209] W.A. Herrebout, J. Lundell, B.J. van der Veken, J. Mol. Struct., 480-481 (1999) 489493. 170 Page 170 of 213

[210] W.A. Herrebout, J. Lundell, B.J. van der Veken, J. Phys. Chem. A, 103 (1999) 76397645. [211] W.A. Herrebout, R. Szostak, B.J. van der Veken, J. Phys. Chem. A, 104 (2000) 84808488.

ip t

[212] W.A. Herrebout, B.J. van der Veken, J. Mol. Struct., 550-551 (2000) 389-398.

[213] G.P. Everaert, W.A. Herrebout, B.J. van der Veken, J. Phys. Chem. A, 105 (2001)

us

cr

9058-9067.

[214] G.P. Everaert, W.A. Herrebout, B.J. van der Veken, Spectrochim. Acta A, 61 (2005)

an

1375-1387.

[215] M.L. Lesiecki, J.S. Shirk, J. Chem. Phys., 56 (1972) 4171-4177.

d

Trans., (1976) 666-676.

M

[216] I.R. Beattie, H.E. Blayden, Hall, S.M., S.N. Jenny, J.S. Ogden, J. Chem. Soc., Dalton

Ac ce pt e

[217] I.R. Beattie, H.E. Blayden, J.S. Ogden, J. Chem. Phys., 64 (1976) 909-910. [218] J.S. Shirk, A.E. Shirk, J. Chem. Phys., 64 (1976) 910-911. [219] Y.S. Yang, J.S. Shirk, J. Mol. Spectrosc., 54 (1975) 39-42. [220] R.G.S. Pong, A.E. Shirk, J.S. Shirk, Ber. Bunsenges. Phys. Chem., 82 (1978) 79-80. [221] R.G.S. Pong, A.E. Shirk, J.S. Shirk, J. Chem. Phys., 70 (1979) 525-531. [222] D.W. Ball, J. Phys. Chem., 99 (1995) 12786-12789. [223] M. Krossner, G. Scholz, R. Stosser, J. Phys. Chem. A, 101 (1997) 1555-1560. [224] A. Boutalib, J. Mol. Struct. (Theochem), 623 (2003) 121-126. [225] I. Jimenez-Fabian, A.F. Jalbout, A. Boutalib, Cent. Eur. J. Chem., 5 (2007) 1007-1018. 171 Page 171 of 213

[226] J. Bahlo, H.J. Himmel, H. Schnöckel, Angew. Chem. Int. Ed., 40 (2001) 4696-4700. [227] J. Bahlo, H.J. Himmel, H. Schnöckel, Inorg. Chem., 41 (2002) 2678-2689. [228] J. Bahlo, H.J. Himmel, H. Schnöckel, Inorg. Chem., 41 (2002) 4488-4495.

ip t

[229] K.A. Grencewicz, D.W. Ball, J. Phys. Chem., 100 (1996) 5672-5675.

us

[231] M.A.A. Beg, H.C. Clark, Can. J. Chem., 40 (1962) 393-398.

cr

[230] R. Ahlrichs, M.R. Bär, M. Häser, E. Sattler, Chem. Phys. Lett., 184 (1991) 353-358.

an

[232] D.C. Mente, J.L. Mills, R.E. Mitchell, Inorg. Chem., 14 (1975) 123-126. [233] D.C. Mente, J.L. Mills, Inorg. Chem., 14 (1975) 1862-1865.

M

[234] E. Wiberg, U. Heubaum, Z. Anorg. Allg. Chem., 225 (1935) 270-272.

Ac ce pt e

2735.

d

[235] J.R. Durig, S. Riethmiller, V.F. Kalasinsky, J.D. Odom, Inorg. Chem., 13 (1974) 2729-

[236] D.R. Martin, R.E. Dial, J. Am. Chem. Soc., 72 (1950) 852-856. [237] B. Rapp, J.E. Drake, Inorg. Chem., 12 (1973) 2868-2873. [238] J.D. Odom, V.F. Kalasinsky, J.R. Durig, Inorg. Chem., 14 (1975) 2837-2839. [239] C. Aubauer, T.M. Klapotke, A. Schulz, J. Mol. Struct. (Theochem), 543 (2001) 285297.

[240] C. Loschen, K. Voigt, J. Frunzke, A. Diefenbach, M. Diedenhofen, G. Frenking, Z. Anorg. Allg. Chem., 628 (2002) 1294-1304. [241] G. Frenking, K. Wichmann, N. Frohlich, C. Loschen, M. Lein, J. Frunzke, V.M. Rayon, Coord. Chem. Rev., 238 (2003) 55-82.

172 Page 172 of 213

[242] F. Bessac, G. Frenking, Inorg. Chem., 42 (2003) 7990-7994. [243] J.A. Plumley, J.D. Evanseck, J. Phys. Chem. A, 113 (2009) 5985-5992. [244] I. Alkorta, J. Elguero, J.E. Del Bene, O. Mó, M. Yáñez, Chem. Eur. J., 16 (2010)

ip t

11897-11905. [245] S. Reinhardt, M. Gastreich, C.M. Marian, Phys. Chem. Chem. Phys., 2 (2000) 955-963.

cr

[246] S. Reinhardt, M. Gastreich, C.M. Marian, Z. Anorg. Allg. Chem., 626 (2000) 1871-

us

1880.

[247] A.Y. Timoshkin, A.V. Suvorov, H.F. Bettinger, H.F. Schaefer, J. Am. Chem. Soc., 121

an

(1999) 5687-5699.

M

[248] A.Y. Timoshkin, H.F. Schaefer, J. Am. Chem. Soc., 125 (2003) 9998-10011.

d

[249] J.C. Dobrowolski, R. Kawecki, J. Mol. Struct., 734 (2005) 235-239.

Ac ce pt e

[250] A. Ogawa, H. Fujimoto, Inorg. Chem., 41 (2002) 4888-4894. [251] L.C. Wang, K. Bao, X. Meng, X.L. Wang, T.T. Jiang, T.A. Cui, B.B. Liu, G.T. Zou, J. Chem. Phys., 134 (2011) 024517.

[252] P. Wang, Dalton Trans., 41 (2012) 4296-4302. [253] L.R. Thorne, R.D. Suenram, F.J. Lovas, J. Chem. Phys., 78 (1983) 167-171. [254] W.T. Klooster, T.F. Koetzle, P.E.M. Siegbahn, T.B. Richardson, R.H. Crabtree, J. Am. Chem. Soc., 121 (1999) 6337-6343. [255] J. Smith, K.S. Seshadri, D. White, J. Mol. Spectrosc., 45 (1973) 327-337. [256] D.G. Allis, M.E. Kosmowski, B.S. Hudson, J. Am. Chem. Soc., 126 (2004) 7756-7757. [257] J. Dillen, P. Verhoeven, J. Phys. Chem. A, 107 (2003) 2570-2577. 173 Page 173 of 213

[258] G.A. Venter, J. Dillen, J. Mol. Struct.: THEOCHEM, 915 (2009) 112-121. [259] S.B. Hyder, T.O. Yep, J. Electrochem. Soc., 123 (1976) 1721-1724. [260] M. Hirayama, K. Shohno, J. Electrochem. Soc., 122 (1975) 1671-1676.

ip t

[261] A.C. Adams, C.D. Capio, J. Electrochem. Soc., 127 (1980) 399-405.

cr

[262] D. Franz, M. Hollenstein, C. Hollenstein, Thin Solid Films, 379 (2000) 37-44.

[263] M.C.L. Gerry, W. Lewis-Bevan, A.J. Merer, N.P.C. Westwood, J. Mol. Spectrosc., 110

us

(1985) 153-163.

an

[264] J.D. Carpenter, B.S. Ault, Chem. Phys. Lett., 197 (1992) 171-174.

M

[265] S. Sakai, Chem. Phys. Lett., 217 (1994) 288-292.

d

[266] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 95 (1991) 3507-3511. [267] S. Aldridge, A.J. Downs, C.Y. Tang, S. Parsons, M.C. Clarke, R.D.L. Johnstone, H.E.

Ac ce pt e

Robertson, D.W.H. Rankin, D.A. Wann, J. Am. Chem. Soc., 131 (2009) 2231-2243. [268] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 97 (1993) 11397-11401. [269] E.L. Smith, D. Sadowsky, J.A. Phillips, C.J. Cramer, D.J. Giesen, J. Phys. Chem. A, 114 (2010) 2628-2636.

[270] E.L. Smith, D. Sadowsky, C.J. Cramer, J.A. Phillips, J. Phys. Chem. A, 115 (2011) 1955-1963.

[271] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 96 (1992) 4288-4294. [272] Y. Kawashima, H. Takeo, C. Matsumura, J. Mol. Spectrosc., 116 (1986) 23-32. [273] J.D. Carpenter, B.S. Ault, J. Mol. Struct., 319 (1994) 139-143.

174 Page 174 of 213

[274] Y. Kawashima, H. Takeo, C. Matsumura, J. Chem. Phys., 74 (1981) 5430-5435. [275] B.S. Ault, Chem. Phys. Lett., 157 (1989) 547-551. [276] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 96 (1992) 7913-7916.

ip t

[277] J.D. Carpenter, B.S. Ault, J. Phys. Chem., 97 (1993) 3697-3700.

us

[279] J.D. Carpenter, B.S. Ault, J. Mol. Struct., 298 (1993) 17-21.

cr

[278] A.M. Mebel, D.G. Musaev, K. Morokuma, J. Phys. Chem., 97 (1993) 7543-7552.

an

[280] E.L. Gamble, P. Gilmont, J. Am. Chem. Soc., 62 (1940) 717-721.

[281] J.R. Durig, Y.S. Li, I.A. Carreira, J.D. Odom, J. Am. Chem. Soc., 95 (1973) 2491-

M

2496.

d

[282] R.W. Parry, T.C. Bissot, J. Am. Chem. Soc., 78 (1956) 1524-1527.

Ac ce pt e

[283] J.D. Carpenter, B.S. Ault, Inorg. Chim. Acta, 286 (1999) 1-6. [284] H.H. Lindner, T. Onak, J. Am. Chem. Soc., 88 (1966) 1886-1889. [285] J.R. Creighton, G.T. Wang, M.E. Coltrin, J. Cryst. Growth, 298 (2007) 2-7. [286] C. Jones, G.A. Koutsantonis, C.L. Raston, Polyhedron, 12 (1993) 1829-1848. [287] S. Aldridge, A.J. Downs, Chem. Rev., 101 (2001) 3305-3365. [288] P.T. Brain, H.E. Brown, A.J. Downs, T.M. Greene, E. Johnsen, S. Parsons, D.W.H. Rankin, B.A. Smart, C.Y. Tang, J. Chem. Soc., Dalton Trans., (1998) 3685-3691. [289] H.J. Himmel, A.J. Downs, T.M. Greene, Chem. Comm., (2000) 871-872. [290] C.Y. Tang, R.A. Coxall, A.J. Downs, T.M. Greene, L. Kettle, S. Parsons, D.W.H. Rankin, H.E. Robertson, A.R. Turner, Dalton Trans., (2003) 3526-3533. 175 Page 175 of 213

[291] M. Chaillet, A. Dargelos, C.J. Marsden, New J. Chem., 18 (1994) 693-700. [292] F.A. Kurth, R.A. Eberlein, H. Schnöckel, A.J. Downs, C.R. Pulham, J. Chem. Soc., Chem. Comm., (1993) 1302-1304.

[294] J.M. Parnis, G.A. Ozin, J. Phys. Chem., 93 (1989) 1215-1220.

ip t

[293] G.V. Chertihin, L. Andrews, J. Phys. Chem., 97 (1993) 10295-10300.

cr

[295] P. Pullumbi, C. Mijoule, L. Manceron, Y. Bouteiller, Chem. Phys., 185 (1994) 13-24.

us

[296] P. Pullumbi, Y. Bouteiller, L. Manceron, C. Mijoule, Chem. Phys., 185 (1994) 25-37.

an

[297] A. Almenningen, G. Gundersen, T. Haugen, A. Haaland, Acta Chem. Scand., 26 (1972) 3928-3934.

d

12215-12222.

M

[298] H.E. Warner, Y. Wang, C. Ward, C.W. Gillies, L. Interrante, J. Phys. Chem., 98 (1994)

Ac ce pt e

[299] H. Anane, A. Jarid, A. Boutalib, J. Phys. Chem. A, 103 (1999) 9847-9852. [300] A. Boutalib, A. Jarid, I. Nebot-Gil, F. Tomas, J. Phys. Chem. A, 105 (2001) 65266529.

[301] A. Jarid, A. Boutalib, I. Nebot-Gil, F. Tomas, J. Mol. Struct. (Theochem), 572 (2001) 161-167.

[302] A. Boutalib, J. Phys. Chem. A, 106 (2002) 8933-8936. [303] A. Boutalib, J. Phys. Chem. A, 107 (2003) 2106-2109. [304] A.F. Jalbout, A. Boutalib, J. Phys. Chem. A, 110 (2006) 12524-12527. [305] M. Cherkaoui, A. Boutalib, Ann. Chim. Sci.Mater., 34 (2009) 203-209.

176 Page 176 of 213

[306] A. Martín-Sómer, A. Lamsabhi, O. Mó, M. Yáñez, J. Phys. Chem. A, 116 (2012) 69506954. [307] J. Moc, K. Bober, K. Mierzwicki, Chem. Phys., 327 (2006) 247-260.

Dalton Trans., 39 (2010) 5637-5642.

cr

[309] X. Wang, L. Andrews, J. Phys. Chem. A, 107 (2003) 11371-11379.

ip t

[308] A.J. Downs, T.M. Greene, E. Johnsen, C.R. Pulham, H.E. Robertson, D.A. Wann,

us

[310] X. Wang, L. Andrews, J. Phys. Chem. A, 108 (2004) 4440-4448.

an

[311] A. Kohn, H.H. Himmel, B. Gaertner, Chem. Eur. J., 9 (2003) 3909-3919. [312] H.J. Himmel, L. Manceron, A.J. Downs, P. Pullumbi, J. Am. Chem. Soc., 124 (2002)

M

4448-4457.

[313] H.J. Himmel, L. Manceron, A.J. Downs, P. Pullumbi, Angew. Chem. Int. Ed., 41

d

(2002) 796-799.

Ac ce pt e

[314] K.D. Dobbs, M. Trachtman, C.W. Bock, A.H. Cowley, J. Phys. Chem., 94 (1990) 5210-5211.

[315] R.M. Graves, G.E. Scuseria, J. Chem. Phys., 96 (1992) 3723-3731. [316] A. El Guerraze, A.M. El-Nahas, A. Jarid, C. Serrar, H. Anane, M. Esseffar, Chem. Phys., 313 (2005) 159-168.

[317] G.E. Coates, R.A. Whitcombe, J. Chem. Soc., (1956) 3351-3354. [318] R. Sanchez, C. Arrington, C.A. Arrington, J. Am. Chem. Soc., 111 (1989) 9110-9111. [319] G.W. Bethke, M.K. Wilson, J. Chem. Phys., 26 (1957) 1118-1130. [320] B.S. Ault, J. Phys. Chem., 96 (1992) 7908-7912. 177 Page 177 of 213

[321] F. Watari, S. Shimizu, K. Aida, E. Takayama, Bull. Chem. Soc. Jpn., 51 (1978) 16021608. [322] G. Schomburg, E.G. Hoffmann, Z. Elektrochem. Angew. Phys. Chem., 61 (1957) 1110-1117.

ip t

[323] J. Müller, J. Am. Chem. Soc., 118 (1996) 6370-6376.

us

[325] J.L. Laboy, B.S. Ault, J. Mol. Struct., 300 (1993) 351-362.

cr

[324] J.R. Creighton, G.T. Wang, J. Phys. Chem. A, 109 (2005) 10554-10562.

an

[326] R. Tarao, Bull. Chem. Soc. Jpn., 39 (1966) 2126-2132. [327] R. Tarao, Bull. Chem. Soc. Jpn., 39 (1966) 725-728.

M

[328] R. Tarao, Bull. Chem. Soc. Jpn., 39 (1966) 2132-2134.

d

[329] S. Takeda, R. Tarao, Bull. Chem. Soc. Jpn., 38 (1965) 1567-1574.

Ac ce pt e

[330] B.S. Ault, J.L. Laboy, J. Mol. Struct., 475 (1999) 193-202. [331] B.S. Ault, J. Organomet. Chem., 572 (1999) 169-175. [332] B.S. Ault, J. Phys. Chem., 98 (1994) 77-80. [333] B.S. Sywe, J.R. Schlup, J.H. Edgar, Chem. Mater., 3 (1991) 737-742. [334] M.J. Almond, C.E. Jenkins, D.A. Rice, K. Hagen, J. Organomet. Chem., 439 (1992) 251-261. [335] E.A. Piocos, B.S. Ault, J. Am. Chem. Soc., 111 (1989) 8978-8979. [336] E.A. Piocos, B.S. Ault, J. Phys. Chem., 95 (1991) 6827-6830. [337] C.W. Bock, M. Trachtman, Structural Chemistry, 4 (1993) 15-18. 178 Page 178 of 213

[338] M.J. Almond, C.E. Beer, D.A. Rice, K. Hagen, J. Mol. Struct., 346 (1995) 83-94. [339] G.T. Wang, J.R. Creighton, J. Phys. Chem. A, 110 (2006) 1094-1099. [340] J. Müller, S. Bendix, Chem. Comm., (2001) 911-912.

us

[343] E.A. Piocos, B.S. Ault, J. Mol. Struct., 317 (1994) 223-232.

cr

[342] E.A. Piocos, B.S. Ault, J. Phys. Chem., 96 (1992) 7589-7593.

ip t

[341] J. Müller, B. Wittig, H. Sternkicker, S. Bendix, J. Phys. IV, 11 (2001) 17-22.

an

[344] E.A. Piocos, B.S. Ault, Inorg. Chem., 32 (1993) 5246-5250.

[345] N. Maung, G.H. Fan, T.L. Ng, J.O. Williams, A.C. Wright, J. Mater. Chem., 9 (1999)

M

2489-2494.

d

[346] N. Maung, J. Mol. Struct. (Theochem), 432 (1998) 129-137.

Ac ce pt e

[347] J.R. Creighton, G.T. Wang, J. Phys. Chem. A, 109 (2005) 133-137. [348] R.T. Arlinghaus, L. Andrews, J. Chem. Phys., 81 (1984) 4341-4351. [349] E.A. Piocos, B.S. Ault, J. Phys. Chem., 97 (1993) 3492-3496. [350] E.A. Piocos, B.S. Ault, J. Mol. Struct., 476 (1999) 283-288. [351] A. Locy, B.S. Ault, Chem. Phys., 392 (2012) 192-197. [352] B.H. Cardelino, C.E. Moore, C.A. Cardelino, D.O. Frazier, K.J. Bachmann, J. Phys. Chem. A, 105 (2001) 849-868. [353] A.P. Kurbakova, S.S. Bukalov, L.A. Leites, L.M. Golubinskaya, V.I. Bregadze, J. Organomet. Chem., 536-537 (1997) 519-529.

179 Page 179 of 213

[354] V.O. Gel’mbol’dt, A.A. Ennan, A.V. Sakharov, V.F. Sukhoverkhov, Zh. Neorg. Khim., 34 (1989) 2847-2856. [355] V.O. Gel'mbol'dt, L.V. Koroeva, A.A. Ennan, Russ. J. Coord. Chem., 33 (2007) 160167.

ip t

[356] W. Levason, G. Reid, W.J. Zhang, Coord. Chem. Rev., 255 (2011) 1319-1341.

[357] E.I. Davydova, A.Y. Timoshkin, T.N. Sevastianova, A.V. Suvorov, G. Frenking, J.

us

cr

Mol. Struct. (Theochem), 767 (2006) 103-111.

[358] S.K. Ignatov, T.D. Kolomiitsova, Z. Mielke, A.G. Razuvaev, D.N. Shchepkin, K.G.

an

Tokhadze, Chem. Phys., 324 (2006) 753-766.

[359] W.Y. Lee, J.R. Strife, R.D. Veltri, J. Am. Ceram. Soc., 75 (1992) 2803-2808.

M

[360] W.Y. Lee, J.R. Strife, R.D. Veltri, J. Am. Ceram. Soc., 75 (1992) 2200-2206.

d

[361] B.S. Ault, Inorg. Chem., 20 (1981) 2817-2822.

Ac ce pt e

[362] E.L. Muetterties, W. Mahler, R. Schmutzler, Inorg. Chem., 2 (1963) 613-618. [363] T.J. Lorenz, B.S. Ault, Inorg. Chem., 21 (1982) 1758-1761. [364] C.J. Marsden, Inorg. Chem., 22 (1983) 3177-3178. [365] A.R. Rossi, J.M. Jasinski, Chem. Phys. Lett., 169 (1990) 399-404. [366] J.N. Hu, L.J. Schaad, B.A. Hess, J. Am. Chem. Soc., 113 (1991) 1463-1464. [367] R.S. Ruoff, T. Emilsson, A.I. Jaman, T.C. Germann, H.S. Gutowsky, J. Chem. Phys., 96 (1992) 3441-3446. [368] A.V. Sakharov, V.F. Sukhoverkhov, A.A. Ennan, V.O. Gel'mbol'dt, Zh. Neorg. Khim., 34 (1989) 1914-1916. 180 Page 180 of 213

[369] I.R. Beattie, G.A. Ozin, J. Chem. Soc. A, (1969) 2267-2269. [370] A.Y. Timoshkin, T.N. Sevast'yanova, E.I. Davydova, A.V. Suvorov, H.F. Schaefer, Russ. J. Gen. Chem., 72 (2002) 1576-1585. [371] A.Y. Timoshkin, T.N. Sevast'yanova, E.I. Davydova, A.V. Suvorov, H.F. Schaefer,

ip t

Russ. J. Gen. Chem., 72 (2002) 1911-1915.

[372] T.N. Sevast'yanova, A.Y. Timoshkin, E.I. Davydova, O.V. Dubov, A.V. Suvorov, H.F.

us

cr

Schaefer, Russ. J. Gen. Chem., 73 (2003) 48-53.

[373] A.Y. Timoshkin, T.N. Sevast'yanova, E.I. Davydova, A.V. Suvorov, H.F. Schaefer,

an

Russ. J. Gen. Chem., 73 (2003) 765-775.

[374] E.I. Davydova, A.Y. Timoshkin, T.N. Sevast'yanova, A.V. Suvorov, H.F. Schaefer,

M

Russ. J. Gen. Chem., 73 (2003) 1742-1750.

d

[375] A.M. McNair, B.S. Ault, Inorg. Chem., 21 (1982) 1762-1765.

Ac ce pt e

[376] B.S. Ault, J. Mol. Struct., 129 (1985) 287-298. [377] B.S. Ault, J. Mol. Struct., 130 (1985) 215-226. [378] M.F. Davis, W. Levason, G. Reid, M. Webster, Dalton Trans., (2008) 2261-2269. [379] I.R. Beattie, G.A. Ozin, J. Chem. Soc. A, (1970) 370-377. [380] D.K. Frieson, G.A. Ozin, Can. J. Chem., 51 (1973) 2685-2696. [381] J.L. Margrave, K.G. Sharp, P.W. Wilson, Inorg. Nucl. Chem. Lett., 5 (1969) 995-997. [382] J.L. Margrave, K.G. Sharp, P.W. Wilson, J. Am. Chem. Soc., 92 (1970) 1530-1532. [383] P.G. Sennikov, S.K. Ignatov, O. Schrems, Russ. J. Inorg. Chem., 55 (2010) 413-420.

181 Page 181 of 213

[384] P. Sennikov, M. Ikrin, S. Ignatov, A. Bagatur'yants, E. Klimov, Russ. Chem. Bull., 48 (1999) 93-97. [385] S.K. Ignatov, P.G. Sennikov, L.A. Chuprov, A.G. Razuvaev, Russ. Chem. Bull., 52 (2003) 837-845.

ip t

[386] S.K. Ignatov, P.G. Sennikov, B.S. Ault, A.A. Bagatur'yants, I.V. Simdyanov, A.G. Razuvaev, E.J. Klimov, O. Gropen, J. Phys. Chem. A, 103 (1999) 8328-8336.

cr

[387] S.K. Ignatov, P.G. Sennikov, A.G. Razuvaev, L.A. Chuprov, Russ. Chem. Bull., 50

us

(2001) 2316-2324.

an

[388] A.M. Walther, B.S. Ault, Inorg. Chem., 23 (1984) 3892-3897. [389] B.S. Ault, J. Mol. Struct., 127 (1985) 357-367.

M

[390] D.K. Padma, A. Iyengar, Ind. J. Chem. A, 23 (1984) 63-64.

Ac ce pt e

1787-1790.

d

[391] B.S. Suresh, V. Chandrasekhar, D.K. Padma, J. Chem. Soc., Dalton Trans., (1984)

[392] A.M. McNair, B.S. Ault, Inorg. Chem., 21 (1982) 2603-2605. [393] P. Ramasami, T.A. Ford, J. Mol. Struct. (Theochem), 940 (2010) 50-55. [394] R.W. Chorley, P.W. Lednor, Adv. Mater., 3 (1991) 474-485. [395] H.J. Himmel, N. Schiefenhövel, M. Binnewies, Chem. Eur. J., 9 (2003) 1387-1393. [396] M.-L.H. Jeng, B.S. Ault, Inorg. Chem., 29 (1990) 837-842. [397] M. Brehm, B.S. Ault, J. Mol. Struct., 649 (2003) 95-103. [398] N. Goldberg, M.J. Almond, J.S. Ogden, J.P. Cannady, R. Walsh, R. Becerra, Phys. Chem. Chem. Phys., 6 (2004) 3264-3270. 182 Page 182 of 213

[399] D.R. Meininger, B.S. Ault, J. Phys. Chem. A, 104 (2000) 3481-3486. [400] S.R. Davis, L. Andrews, J. Chem. Phys., 86 (1987) 3765-3772. [401] S.W. Hu, Y. Wang, X.Y. Wang, T.W. Chu, X.Q. Liu, J. Phys. Chem. A, 107 (2003)

[402] T.C. McInnis, L. Andrews, J. Phys. Chem., 96 (1992) 5276-5284.

ip t

9189-9196.

cr

[403] Z. Mielke, K.G. Tokhadze, Chem. Phys. Lett., 316 (2000) 108-114.

us

[404] N.N. Zemlyanskii, I.V. Borisova, M.S. Nechaev, V.N. Khrustalev, V.V. Lunin, M.Y.

an

Antipin, Y.A. Ustynyuk, Russ. Chem. Bull., 53 (2004) 980-1006.

[405] S.E. Boganov, V.I. Faustov, M.P. Egorov, O.M. Nefedov, Russ. Chem. Bull., 53 (2004)

M

960-979.

d

[406] A. Patyk, W. Sander, J. Gauss, D. Cremer, Chem. Ber., 123 (1990) 89-90. [407] A.V. Lalov, S.E. Boganov, V.I. Faustov, M.P. Egorov, O.M. Nefedov, Russ. Chem.

Ac ce pt e

Bull., 52 (2003) 526-538.

[408] S.E. Boganov, V.I. Faustov, M.P. Egorov, O.M. Nefedov, Russ. Chem. Bull., 47 (1998) 1054-1060.

[409] D. Tevault, K. Nakamoto, Inorg. Chem., 15 (1976) 1282-1287. [410] M. Tacke, C. Klein, D.J. Stufkens, A. Oskam, P. Jutzi, E.A. Bunte, Z. Anorg. Allg. Chem., 619 (1993) 865-868.

[411] N. Goldberg, J.S. Ogden, M.J. Almond, R. Walsh, J.P. Cannady, R. Becerra, J.A. Lee, Phys. Chem. Chem. Phys., 5 (2003) 5371-5377. [412] S.E. Boganov, V.I. Faustov, M.P. Egorov, O.M. Nefedov, Russ. Chem. Bull., 43 (1994) 47-49. 183 Page 183 of 213

[413] S.E. Boganov, V.I. Faustov, S.G. Rudyak, M.P. Egorov, O.M. Nefedov, Russ. Chem. Bull., 45 (1996) 1061-1067. [414] S.E. Boganov, M.P. Egorov, O.M. Nefedov, Russ. Chem. Bull., 48 (1999) 98-103. [415] G. Maier, H.P. Reisenauer, J. Glatthaar, R. Zetzmann, Chem. Asian J., 1 (2006) 195-

ip t

202.

cr

[416] G. Maier, H.P. Reisenauer, H. Egenolf, Organometallics, 18 (1999) 2155-2161.

us

[417] J.C. Amicangelo, C.T. Dine, D.G. Irwin, C.J. Lee, N.C. Romano, N.L. Saxton, J. Phys. Chem. A, 112 (2008) 3020-3030.

an

[418] R. Withnall, L. Andrews, J. Phys. Chem., 94 (1990) 2351-2357.

M

[419] W.P. Neumann, Chem. Rev., 91 (1991) 311-334.

d

[420] T.J. Drahnak, J. Michl, R. West, J. Am. Chem. Soc., 101 (1979) 5427-5428.

Ac ce pt e

[421] M. Haaf, T.A. Schmedake, R. West, Acc. Chem. Res., 33 (2000) 704-714. [422] C.A. Arrington, J.C. Poutsma, N. Stone, Abs. Pap.Am. Chem. Soc., 201 (1991) 128PHYS.

[423] M. Veith, E. Werle, R. Lisowsky, R. Koppe, H. Schnöckel, Chem. Ber., 125 (1992) 1375-1377.

[424] O.M. Nefedov, M.P. Egorov, A.I. Ioffe, L.G. Menchikov, P.S. Zuev, V.I. Minkin, B.Y. Simkin, M.N. Glukhovtsev, Pure Appl. Chem., 64 (1992) 265-314. [425] W. Sander, G. Bucher, S. Wierlacher, Chem. Rev., 93 (1993) 1583-1621. [426] V.N. Khabashesku, S.E. Boganov, K.N. Kudin, J.L. Margrave, J. Michl, O.M. Nefedov, Russ. Chem. Bull., 48 (1999) 2003-2015.

184 Page 184 of 213

[427] R.S. Sheridan, J. Phys. Org. Chem., 23 (2010) 326-332. [428] C. Wentrup, Acc. Chem. Res., 44 (2011) 393-404. [429] M.F. Zhou, M.H. Chen, Chem. Phys. Lett., 349 (2001) 64-70.

ip t

[430] M.H. Chen, A.H. Zheng, H. Lu, M.F. Zhou, J. Phys. Chem. A, 106 (2002) 3077-3083.

cr

[431] M. Zhou, L. Zhang, H. Lu, L. Shao, M. Chen, J. Mol. Struct., 605 (2002) 249-254.

us

[432] J. Dong, L. Miao, M.F. Zhou, Chem. Phys. Lett., 355 (2002) 31-36.

an

[433] M.E. Palumbo, P. Leto, C. Siringo, C. Trigilio, Astrophys. J., 685 (2008) 1033-1038. [434] P.A. Gerakines, M.H. Moore, Icarus, 154 (2001) 372-380.

M

[435] N. Piétri, T. Chiavassa, A. Allouche, J.P. Aycard, J. Phys. Chem. A, 101 (1997) 1093-

d

1098.

Ac ce pt e

[436] N. Piétri, I. Tamburelli, J.P. Aycard, T. Chiavassa, J. Mol. Struct., 416 (1997) 187-195. [437] I. Tamburelli, T. Chiavassa, F. Borget, J. Pourcin, J. Phys. Chem. A, 102 (1998) 422425.

[438] I. Couturier-Tamburelli, T. Chiavassa, J.P. Aycard, J. Am. Chem. Soc., 121 (1999) 3756-3761.

[439] I. Couturier-Tamburelli, J.P. Aycard, M.W. Wong, C. Wentrup, J. Phys. Chem. A, 104 (2000) 3466-3471.

[440] I. Couturier-Tamburelli, B. Sessouma, J.P. Aycard, J. Mol. Struct., 560 (2001) 197-203. [441] B. Sessouma, I. Couturier-Tamburelli, M. Monnier, M.W. Wong, C. Wentrup, J.P. Aycard, J. Phys. Chem. A, 106 (2002) 4489-4497.

185 Page 185 of 213

[442] B. Sessouma, I. Couturier-Tamburelli, A. Allouche, J.P. Aycard, Chem. Phys., 278 (2002) 9-19. [443] B.J. Ortman, R.H. Hauge, J.L. Margrave, Z.H. Kafafi, J. Phys. Chem., 94 (1990) 79737977.

cr

[445] S. Ekern, M. Vala, J. Phys. Chem. A, 101 (1997) 3601-3606.

ip t

[444] J. Szczepanski, S. Ekern, M. Vala, J. Phys. Chem., 99 (1995) 8002-8012.

us

[446] Y. Yamada, K. Onaka, C. Obayashi, H. Sato, T. Tominaga, J. Radioanal. Nucl. Chem. Lett., 199 (1995) 477-492.

an

[447] Y. Yamada, T. Kumagawa, Y.T. Yamada, T. Tominaga, J. Radioanal. Nucl. Chem. Lett., 201 (1995) 417-429.

d

1017-1020.

M

[448] S. Bukshpan, P. Hendrickx, K. Milants, H. Pattyn, Solid State Commun., 80 (1991)

Ac ce pt e

[449] S. Bukshpan, Chem. Phys. Lett., 177 (1991) 269-271. [450] E.B. Saitovitch, E.O. Rodrigues, V. Drago, H. Micklitz, Hyperfine Interact., 56 (1990) 1679-1682.

[451] C. Obayashi, H. Sato, T. Tominaga, J. Radioanal. Nucl. Chem. Lett., 127 (1988) 75-85. [452] F.J. Litterst, A. Schichl, E. Baggio-Saitovitch, H. Micklitz, J.M. Friedt, Ber. Bunsenges. Phys. Chem., 82 (1978) 73-75.

[453] A. Bos, A.T. Howe, B.W. Dale, L.W. Becker, J. Chem. Soc., Farad. Trans. II, 70 (1974) 440-450. [454] A. Bos, A.T. Howe, J. Chem. Soc., Farad. Trans. II, 70 (1974) 451-464.

186 Page 186 of 213

[455] A. Bos, A.T. Howe, B.W. Dale, L.W. Becker, J. Chem. Soc., Chem. Commun., (1972) 730-731. [456] L. Abouaf-Marguin, M.E. Jacox, D.E. Milligan, J. Mol. Spectrosc., 67 (1977) 34-61.

[458] W. Levason, G. Reid, Coord. Chem. Rev., 250 (2006) 2565-2594.

ip t

[457] J.A. Altmann, M.G. Govender, T.A. Ford, Mol. Phys., 103 (2005) 949-961.

cr

[459] G.L. Johnson, L. Andrews, J. Am. Chem. Soc., 104 (1982) 3043-3047.

us

[460] L. Andrews, X. Wang, Z. Mielke, J. Phys. Chem. A, 105 (2001) 6054-6064.

an

[461] L. Andrews, X. Wang, J. Phys. Chem. A, 105 (2001) 6420-6429.

M

[462] L. Andrews, X. Wang, J. Phys. Chem. A, 105 (2001) 7541-7550.

d

[463] G.Q. Liu, Y.M. Zhao, J. Mol. Struct. (Theochem), 821 (2007) 173-173.

Ac ce pt e

[464] L. Andrews, S.R. Davis, G.L. Johnson, J. Phys. Chem., 90 (1986) 4273-4282. [465] R.B. Bohn, L. Andrews, J. Phys. Chem., 95 (1991) 9707-9712. [466] R.B. Bohn, L. Andrews, J. Phys. Chem., 93 (1989) 3974-3979. [467] Z. Mielke, K.G. Tokhadze, Z. Latajka, E. Ratajczak, J. Phys. Chem., 100 (1996) 539545.

[468] A. Bil, Z. Latajka, Chem. Phys. Lett., 406 (2005) 366-370. [469] R. Lascola, R. Withnall, L. Andrews, Inorg. Chem., 27 (1988) 642-648. [470] F. Duvernay, T. Chiavassa, F. Borget, J.P. Aycard, Chem. Phys., 298 (2004) 241-250. [471] N. Piétri, B. Sessouma, F. Borget, T. Chiavassa, I. Couturier-Tamburelli, Chem. Phys., 400 (2012) 98-102. 187 Page 187 of 213

[472] K. Hiraoka, A. Yamashita, Y. Yachi, K. Aruga, T. Sato, H. Muto, Astrophys. J., 443 (1995) 363-370. [473] H. Hidaka, M. Watanabe, A. Kouchi, N. Watanabe, Phys. Chem. Chem. Phys., 13 (2011) 15798-15802.

cr

[475] P.G. Sennikov, J. Phys. Chem., 98 (1994) 4973-4981.

ip t

[474] X. Wang, P.F. Souter, L. Andrews, J. Phys. Chem. A, 107 (2003) 4244-4249.

us

[476] Y.J. Chen, M. Nuevo, J.M. Hsieh, T.S. Yih, W.H. Sun, W.H. Ip, H.S. Fung, S.Y. Chiang, Y.Y. Lee, J.M. Chen, C.Y.R. Wu, Astronomy & Astrophysics, 464 (2007) 253-257.

an

[477] R.K. Khanna, M.H. Moore, Spectrochim. Acta A, 55 (1999) 961-967.

(2005) 940-952.

M

[478] P.D. Holtom, C.J. Bennett, Y. Osamura, N.J. Mason, R.I. Kaiser, Astrophys. J., 626

Ac ce pt e

211-217.

d

[479] J.B. Bossa, F. Duvernay, P. Theule, F. Borget, T. Chiavassa, Chem. Phys., 354 (2008)

[480] L. Fredin, B. Nelander, Chem. Phys., 15 (1976) 473-484. [481] J. Lundell, M. Krajewska, M. Räsänen, J. Mol. Struct., 448 (1998) 221-230. [482] H. Bai, B.S. Ault, J. Phys. Chem., 94 (1990) 606-610. [483] L.H. Jones, B.I. Swanson, J. Phys. Chem., 95 (1991) 86-90. [484] D. Lucas, L.J. Allamandola, G.C. Pimentel, Croat. Chem. Acta, 55 (1982) 121-128. [485] N. Sanna, L. Schriver-Mazzuoli, A. Hallou, A. de Saxce, A. Schriver, J. Chem. Phys., 103 (1995) 6930-6940.

188 Page 188 of 213

[486] A. Hallou, L. Schriver-Mazzuoli, A. Schriver, N. Sanna, A. Pieretti, Asian J. Spectrosc., 1 (1997) 189-200. [487] A. Pieretti, N. Sanna, A. Hallou, L. Schriver-Mazzuoli, A. Schriver, J. Mol. Struct., 447 (1998) 223-233.

ip t

[488] L. Schriver-Mazzuoli, A. Hallou, A. Schriver, J. Phys. Chem. A, 102 (1998) 97729778.

cr

[489] G. Cazzoli, C.D. Esposti, P. Palmieri, S. Simeone, J. Mol. Spectrosc., 97 (1983) 165-

us

185.

an

[490] P. Antoniotti, S. Borocci, F. Grandinetti, Eur. J. Inorg. Chem., (2004) 1125-1130. [491] R. Lascola, R. Withnall, L. Andrews, J. Phys. Chem., 92 (1988) 2145-2149.

d

322-327.

M

[492] M.E. Jacox, D.E. Milligan, W.A. Guillory, J.J. Smith, J. Mol. Spectrosc., 52 (1974)

Ac ce pt e

[493] M.E. Jacox, J. Mol. Spectrosc., 84 (1980) 74-88. [494] J.V. Gilbert, L.J. Smith, J. Phys. Chem., 95 (1991) 7278-7281. [495] J.V. Gilbert, B.D. Hofsetz, J. Phys. Chem., 96 (1992) 4321-4324. [496] A.M. Zarubaiko, J.V. Gilbert, J. Phys. Chem., 97 (1993) 4331-4334. [497] P. Wermer, B.S. Ault, Inorg. Chem., 20 (1981) 970-973. [498] M. Hawkins, M.J. Almond, A.J. Downs, J. Phys. Chem., 89 (1985) 3326-3334. [499] E.G. Robertson, D. McNaughton, J. Phys. Chem. A, 107 (2003) 642-650. [500] W. Storzer, D. Schomburg, G.V. Röschenthaler, R. Schmutzler, Chem. Ber., 116 (1983) 367-374. 189 Page 189 of 213

[501] A.L. Khidir Aljibury, R.L. Redington, J. Chem. Phys., 52 (1970) 453-459. [502] L.E. Alexander, I.R. Beattie, J. Chem. Phys., 56 (1972) 5829-5831. [503] I.R. Beattie, R. Crocombe, A. German, C.J. Marsden, G.J. van Schalwyk, A.

ip t

Bukovszky, J. Chem. Soc., Dalton Trans., (1976) 1380-1387. [504] X.F. Wang, L. Andrews, M. Knitter, P.A. Malmqvist, B.O. Roos, J. Phys. Chem. A,

cr

113 (2009) 6064-6069.

us

[505] X.F. Wang, L. Andrews, R. Lindh, V. Veryazov, B.O. Roos, J. Phys. Chem. A, 112 (2008) 8030-8037.

an

[506] X.F. Wang, L. Andrews, Dalton Trans., (2009) 9260-9265.

M

[507] X. Wang, J.T. Lyon, L. Andrews, Inorg. Chem., 48 (2009) 6297-6302.

d

[508] L. Andrews, X.F. Wang, B.O. Roos, Inorg. Chem., 48 (2009) 6594-6598.

Ac ce pt e

[509] E.G. Hope, W. Levason, Coord. Chem. Rev., 122 (1993) 109-170. [510] W. Levason, S.D. Orchard, G. Reid, Coord. Chem. Rev., 225 (2002) 159-199. [511] G. Maes, J. Smets, J. Mol. Struct., 270 (1992) 141-160. [512] T.A. Ford, Can. J. Anal. Sci. Spectrosc., 43 (1998) 113-121. [513] J. Ceponkus, P. Uvdal, B. Nelander, J. Chem. Phys., 134 (2011) 064309. [514] J. Ceponkus, P. Uvdal, B. Nelander, J. Phys. Chem. A, 116 (2012) 4842-4850. [515] A. Schriver, L. Schriver-Mazzuoli, P. Chaquin, E. Dumont, J. Phys. Chem. A, 110 (2006) 51-56. [516] T. Svensson, B. Nelander, G. Karlstrom, Chem. Phys., 265 (2001) 323-333.

190 Page 190 of 213

[517] X. Michaut, A.M. Vasserot, L. Abouaf-Marguin, Low Temp. Phys., 29 (2003) 852-857. [518] J. Lundell, Z. Latajka, Chem. Phys., 263 (2001) 221-230. [519] M.E. Jacox, W.E. Thompson, J. Chem. Phys., 123 (2005) 064501.

ip t

[520] L. Andrews, S.R. Davis, J. Chem. Phys., 83 (1985) 4983-4989.

cr

[521] T.D. Mokomela, I. Rencken, G.A. Yeo, T.A. Ford, J. Mol. Struct., 275 (1992) 33-54.

[522] S. Coussan, A. Loutellier, J.P. Perchard, S. Racine, Y. Bouteiller, J. Mol. Struct., 471

us

(1998) 37-47.

an

[523] S. Hirabayashi, K. Ohno, H. Abe, K.M.T. Yamada, J. Chem. Phys., 122 (2005) 194506. [524] P.D. Cooper, H.G. Kjaergaard, V.S. Langford, A.J. McKinley, T.I. Quickenden, T.W.

M

Robinson, D.P. Schofield, J. Phys. Chem. A, 109 (2005) 4274-4279.

d

[525] J.A.G. Gomes, J.L. Gossage, H. Balu, M. Kesmez, F. Bowen, R.S. Lumpkin, D.L.

Ac ce pt e

Cocke, Spectrochim. Acta A, 61 (2005) 3082-3086. [526] S. Hirabayashi, K.M.T. Yamada, Chem. Phys. Lett., 418 (2006) 323-327. [527] K. Mierzwicki, Z. Mielke, M. Sałdyka, S. Coussan, P. Roubin, Phys. Chem. Chem. Phys., 10 (2008) 1292-1297.

[528] D. Forney, M.E. Jacox, W.E. Thompson, J. Chem. Phys., 121 (2004) 5977-5984. [529] M.E. Fajardo, S. Tam, M.E. DeRose, J. Mol. Struct., 695 (2004) 111-127. [530] J. Ceponkus, B. Nelander, J. Chem. Phys., 124 (2006) 024504. [531] A.L. Garden, J.R. Lane, H.G. Kjaergaard, J. Chem. Phys., 125 (2006) 144317. [532] E. Isoniemi, L. Khriachtchev, J. Lundell, M. Räsänen, J. Mol. Struct., 563 (2001) 261265. 191 Page 191 of 213

[533] J. Lundell, E. Nordquist, M. Räsänen, J. Mol. Struct., 416 (1997) 235-242. [534] T.L. Tso, E.K.C. Lee, J. Phys. Chem., 88 (1984) 2776-2781. [535] R.R. Smardzewski, M.C. Lin, J. Chem. Phys., 66 (1977) 3197-3205.

ip t

[536] S. Pehkonen, J. Lundell, L. Khriachtchev, M. Pettersson, M. Räsänen, Phys. Chem. Chem. Phys., 6 (2004) 4607-4613.

us

[538] A. Engdahl, B. Nelander, Science, 295 (2002) 482-483.

cr

[537] A. Engdahl, B. Nelander, G. Karlstrom, J. Phys. Chem. A, 105 (2001) 8393-8398.

an

[539] A. Engdahl, B. Nelander, Chem. Phys., 293 (2003) 203-209.

M

[540] A. Engdahl, B. Nelander, Phys. Chem. Chem. Phys., 2 (2000) 3967-3970.

d

[541] J. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A, 104 (2000) 2033-2037.

Ac ce pt e

[542] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A, 105 (2001) 6430-6435. [543] J.R. Goebel, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A, 105 (2001) 11365-11370. [544] J.R. Goebel, K.A. Antle, B.S. Ault, J.E. Del Bene, J. Phys. Chem. A, 106 (2002) 64066414.

[545] J. Lundell, S. Jolkkonen, L. Khriachtchev, M. Pettersson, M. Räsänen, Chem. Eur. J., 7 (2001) 1670-1678.

[546] S. Pehkonen, M. Pettersson, J. Lundell, L. Khriachtchev, M. Räsänen, J. Phys. Chem. A, 102 (1998) 7643-7648. [547] B. Makiabadi, H. Roohi, Chem. Phys. Lett., 460 (2008) 72-78. [548] L. Andrews, M. Hawkins, R. Withnall, Inorg. Chem., 24 (1985) 4234-4239.

192 Page 192 of 213

[549] L. Andrews, Absorption spectroscopy of molecular ions and complexes in noble-gas matrices, in: L. Andrews, M. Moskovits (Eds.) Chemistry and Physics of Matrix Isolated Species, Elsevier, Amsterdam, 1989, pp. 15-46. [550] R. Withnall, L. Andrews, J. Phys. Chem., 89 (1985) 3261-3268.

ip t

[551] M. Bahou, L. Schriver-Mazzuoli, A. Schriver, J. Chem. Phys., 114 (2001) 4045-4052.

cr

[552] L. Nord, J. Mol. Struct., 96 (1982) 37-44.

us

[553] L. Andrews, R. Withnall, R.D. Hunt, J. Phys. Chem., 92 (1988) 78-81.

an

[554] Z. Mielke, L. Andrews, J. Phys. Chem., 94 (1990) 3519-3525.

[555] L. Schriver, C. Barreau, A. Schriver, Chem. Phys., 140 (1990) 429-438.

M

[556] A. Engdahl, B. Nelander, J. Chem. Phys., 122 (2005) 126101.

d

[557] P.R. Jones, H. Taube, J. Phys. Chem., 77 (1973) 1007-1011.

Ac ce pt e

[558] M.D. Hoops, B.S. Ault, J. Mol. Struct., 929 (2009) 22-31. [559] M.D. Hoops, B.S. Ault, J. Am. Chem. Soc., 131 (2009) 2853-2863. [560] M. Clay, B.S. Ault, J. Phys. Chem. A, 114 (2010) 2799-2805. [561] B.E. Coleman, B.S. Ault, J. Phys. Chem. A, 114 (2010) 12667-12674. [562] B.E. Coleman, B.S. Ault, J. Mol. Struct., 976 (2010) 249-254. [563] J.G. Deng, J.H. Chen, C.M. Geng, H.J. Liu, W. Wang, Z.P. Bai, Y.S. Xu, J. Phys. Chem. A, 116 (2012) 1710-1716. [564] H. Frei, L. Fredin, G.C. Pimentel, J. Chem. Phys., 74 (1981) 397-411.

193 Page 193 of 213

[565] V. Raducu, D. Jasmin, R. Dahoo, P. Brosset, B. Gauthierroy, L. Abouaf-Marguin, J. Chem. Phys., 101 (1994) 1878-1884. [566] R. Withnall, L. Andrews, J. Phys. Chem., 92 (1988) 2155-2161.

ip t

[567] R.R. Lucchese, K. Haber, H.F. Schaefer, J. Am. Chem. Soc., 98 (1976) 7617-7620. [568] L. Andrews, R. Withnall, B.W. Moores, J. Phys. Chem., 93 (1989) 1279-1285.

cr

[569] R. Withnall, L. Andrews, J. Phys. Chem., 91 (1987) 784-797.

us

[570] R. Withnall, L. Andrews, J. Phys. Chem., 92 (1988) 4610-4619.

an

[571] L. Andrews, M. McCluskey, Z. Mielke, R. Withnall, J. Mol. Struct., 222 (1990) 95108.

M

[572] M. McCluskey, L. Andrews, J. Phys. Chem., 95 (1991) 2988-2994.

d

[573] L. Andrews, Z. Mielke, Inorg. Chem., 28 (1989) 4001-4006.

Ac ce pt e

[574] L. Andrews, B.W. Moores, K.K. Fonda, Inorg. Chem., 28 (1989) 290-297. [575] I. Reva, L. Lapinski, M.J. Nowak, Chem. Phys. Lett., 467 (2008) 97-100. [576] F.W.S. Benfield, A.J. Downs, G.P. Gaskill, G.E. Staniforth, J. Chem. Soc., Chem. Commun., (1976) 856-858.

[577] A.J. Downs, G.P. Gaskill, S.B. Saville, Inorg. Chem., 21 (1982) 3385-3393. [578] B.W. Moores, L. Andrews, J. Phys. Chem., 93 (1989) 1902-1907. [579] D.E. Tevault, R.L. Mowery, R.R. Smardzewski, J. Chem. Phys., 74 (1981) 4480-4487. [580] D. Wakamatsu, N. Akai, A. Kawai, K. Shibuya, Chem. Lett., 41 (2012) 252-253.

194 Page 194 of 213

[581] C. Lugez, A. Schriver, R. Levant, L. Schriver-Mazzuoli, Chem. Phys., 181 (1994) 129146. [582] E.Y. Misochko, A.V. Akimov, C.A. Wight, J. Phys. Chem. A, 103 (1999) 7972-7977.

[584] S.C. Bhatia, J.H. Hall, J. Phys. Chem., 85 (1981) 2055-2060.

cr

[585] R.O. Carter, L. Andrews, J. Phys. Chem., 85 (1981) 2351-2354.

ip t

[583] S.C. Bhatia, J.H. Hall, Inorg. Chem., 20 (1981) 629-630.

us

[586] M. Bahou, L. Schriver-Mazzuoli, A. Schriver, P. Chaquin, Chem. Phys., 216 (1997)

an

105-118.

[587] L. Schriver-Mazzuoli, O. Abdelaoui, C. Lugez, A. Schriver, Chem. Phys. Lett., 214

M

(1993) 519-526.

d

[588] S.D. Allen, M. Poliakoff, J.J. Turner, J. Mol. Struct., 157 (1987) 1-15. [589] M. Hawkins, L. Andrews, A.J. Downs, D.J. Drury, J. Am. Chem. Soc., 106 (1984)

Ac ce pt e

3076-3082.

[590] R.J.H. Clark, L.J. Foley, S.D. Price, J. Phys. Chem. A, 104 (2000) 10675-10682. [591] M. Hawkins, L. Andrews, Inorg. Chem., 24 (1985) 3285-3290. [592] R.J.H. Clark, J.R. Dann, J. Phys. Chem., 100 (1996) 532-538. [593] R.J.H. Clark, J.R. Dann, J. Phys. Chem., 100 (1996) 9271-9275. [594] R.J.H. Clark, J.R. Dann, L.J. Foley, J. Phys. Chem. A, 101 (1997) 9260-9271. [595] R.J.H. Clark, J.R. Dann, J. Phys. Chem. A, 101 (1997) 2074-2082. [596] R.J.H. Clark, J.R. Dann, L.J. Foley, J. Chem. Soc., Dalton Trans., (1999) 73-78.

195 Page 195 of 213

[597] R.J.H. Clark, L.J. Foley, J. Phys. Chem. A, 106 (2002) 3356-3364. [598] M. Wierzejewska, J. Mol. Struct., 520 (2000) 199-214. [599] M. Wierzejewska, Vib. Spectrosc., 23 (2000) 253-262.

ip t

[600] M. Wierzejewska, Z. Mielke, R. Wieczorek, Z. Latajka, Chem. Phys., 228 (1998) 1729.

us

[603] L. Nord, J. Mol. Struct., 96 (1982) 19-25.

an

[602] T.A. Ford, J. Mol. Struct., 924-26 (2009) 466-472.

cr

[601] M. Wierzejewska, A. Olbert-Majkut, J. Phys. Chem. A, 107 (2003) 10944-10952.

M

[604] J.D.D. Nelson, G.T. Fraser, W. Klemperer, J. Chem. Phys., 83 (1985) 945-949. [605] A. Taleb-Bendiab, K.W. Hillig, R.L. Kuczkowski, J. Chem. Phys., 94 (1991) 6956-

d

6963.

Ac ce pt e

[606] F. Ito, Chem. Phys. Lett., 436 (2007) 335-340. [607] Wierzejewska-Hnat, M., A. Schriver, L. Schriver-Mazzuoli, Chem. Phys., 183 (1994) 117-126.

[608] L.M. Nxumalo, T.A. Ford, J. Mol. Struct., 347 (1995) 495-507. [609] L. Schriver-Mazzuoli, A. Schriver, M. Wierzejewska-Hnat, Chem. Phys., 199 (1995) 227-243.

[610] F. Ito, S. Hirabayashi, Chem. Phys., 358 (2009) 209-218. [611] R.G.A. Bone, C.R. Lesueur, R.D. Amos, A.J. Stone, J. Chem. Phys., 96 (1992) 83908410. [612] L. Nord, J. Mol. Struct., 96 (1982) 27-35. 196 Page 196 of 213

[613] L. Fredin, Chem. Scr., 4 (1973) 97-102. [614] D. van der Helm, J.D. Childs, S.D. Christian, J. Chem. Soc., D, Chem. Commun., (1969) 887-888.

ip t

[615] J.E. Douglas, P.A. Kollman, J. Am. Chem. Soc., 100 (1978) 5226-5227. [616] M.W. Wong, K.B. Wiberg, J. Am. Chem. Soc., 114 (1992) 7527-7535.

cr

[617] R. Steudel, Y. Steudel, Eur. J. Inorg. Chem., (2007) 4385-4392.

us

[618] J.A. Phillips, D. Britton, K.R. Leopold, Journal of Chemical Crystallography, 26 (1996)

an

533-538.

[619] J.J. Oh, M.S. Labarge, J. Matos, J.W. Kampf, K.W. Hillig, R.L. Kuczkowski, J. Am.

M

Chem. Soc., 113 (1991) 4732-4738.

d

[620] J.J. Oh, K.W. Hillig, R.L. Kuczkowski, J. Phys. Chem., 95 (1991) 7211-7216.

Ac ce pt e

[621] J. Choo, S. Kim, Y. Kwon, J. Mol. Struct. (Theochem), 594 (2002) 147-156. [622] I.C. Hisatsune, J. Heicklen, Can. J. Chem., 53 (1975) 2646-2656. [623] C.S. Sass, B.S. Ault, J. Phys. Chem., 88 (1984) 432-440. [624] N. Maier, J. Schiewe, H. Matschiner, C.P. Maschmeier, R. Boese, Phosphorus, Sulfur Silicon Relat. Elem., 91 (1994) 179-188.

[625] A.A. Eigner, J.P. Wrass, E.L. Smith, C.C. Knutson, J.A. Phillips, J. Mol. Struct., 919 (2009) 312-320. [626] J.J. Oh, K.W. Hillig, R.L. Kuczkowski, J. Am. Chem. Soc., 113 (1991) 7480-7484. [627] H. Bai, B.S. Ault, J. Mol. Struct., 238 (1990) 223-230. [628] J.R. Sodeau, E.K.C. Lee, J. Phys. Chem., 84 (1980) 3358-3362. 197 Page 197 of 213

[629] L. Schriver, D. Carrere, A. Schriver, K. Jaeger, Chem. Phys. Lett., 181 (1991) 505-511. [630] H. Chaabouni, L. Schriver-Mazzuoli, A. Schriver, J. Phys. Chem. A, 104 (2000) 34983507.

ip t

[631] A. Schriver, L. Schriver, J.P. Perchard, J. Mol. Spectrosc., 127 (1988) 125-142. [632] K. Matsumura, F.J. Lovas, R.D. Suenram, J. Chem. Phys., 91 (1989) 5887-5894.

cr

[633] R. Steudel, Y. Steudel, Eur. J. Inorg. Chem., (2009) 1393-1405.

us

[634] S. Hirabayashi, F. Ito, K.M.T. Yamada, J. Chem. Phys., 125 (2006) 034508.

an

[635] J. Cukras, J. Sadlej, J. Mol. Struct. (Theochem), 819 (2007) 41-51.

M

[636] T. Svensson, B. Nelander, Chem. Phys., 286 (2003) 347-352.

d

[637] G. Cook, O.D. Krogh, J. Chem. Phys., 74 (1981) 841-848.

Ac ce pt e

[638] S. Li, Y.S. Li, Spectrochim. Acta A, 47 (1991) 201-209. [639] S. Li, Y.S. Li, J. Mol. Struct., 248 (1991) 79-88. [640] S. Li, Y.S. Li, J. Mol. Struct., 301 (1993) 21-28. [641] Y.S. Li, S. Li, Spectrochim. Acta A, 50 (1994) 509-519. [642] S. Li, H. Kurtz, P. Korambath, Y.S. Li, J. Mol. Struct., 550 (2000) 235-244. [643] K. Garber, B.S. Ault, Inorg. Chem., 22 (1983) 2509-2513. [644] B.S. Ault, J. Mol. Struct., 238 (1990) 111-117. [645] A.M. Andrews, A. Taleb-Bendiab, M.S. Labarge, K.W. Hillig, R.L. Kuczkowski, J. Chem. Phys., 93 (1990) 7030-7040.

198 Page 198 of 213

[646] M.S. Labarge, J.-J. Oh, K.W. Hillig, R.L. Kuczkowski, Chem. Phys. Lett., 159 (1989) 559-562. [647] A. Talebbendiab, K.W. Hillig, R.L. Kuczkowski, J. Chem. Phys., 97 (1992) 2996-3006.

ip t

[648] V.E. Bondybey, J.H. English, J. Mol. Spectrosc., 109 (1985) 221-228. [649] A. Givan, A. Loewenschuss, C.J. Nielsen, M. Rozenberg, J. Mol. Struct., 830 (2007)

cr

21-34.

us

[650] F.A. Kanda, A.J. King, J. Am. Chem. Soc., 73 (1951) 2315-2319.

an

[651] C.S. Sass, B.S. Ault, J. Phys. Chem., 90 (1986) 4533-4536. [652] C.S. Sass, B.S. Ault, J. Phys. Chem., 91 (1987) 551-554.

M

[653] G.L. Shen, M. Suto, L.C. Lee, J. Geophys. Res., 95 (1990) 13981-13984.

d

[654] G.L. Shen, M. Suto, L.C. Lee, Int. J. Chem. Kinet., 22 (1990) 633-639.

Ac ce pt e

[655] M. Canagaratna, J.A. Phillips, H. Goodfriend, K.R. Leopold, J. Am. Chem. Soc., 118 (1996) 5290-5295.

[656] M. Canagaratna, M.E. Ott, K.R. Leopold, Chem. Phys. Lett., 281 (1997) 63-68. [657] D.L. Fiacco, A. Toro, K.R. Leopold, Inorg. Chem., 39 (2000) 37-43. [658] S.W. Hunt, K.R. Leopold, The Journal of Physical Chemistry A, 105 (2001) 54985506.

[659] M.B. Craddock, C.S. Brauer, K.J. Higgins, K.R. Leopold, J. Mol. Spectrosc., 222 (2003) 63-73. [660] K.H. Bowen, K.R. Leopold, K.V. Chance, W. Klemperer, J. Chem. Phys., 73 (1980) 137-141. 199 Page 199 of 213

[661] W.A. Burns, J.A. Phillips, M. Canagaratna, H. Goodfriend, K.R. Leopold, J. Phys. Chem. A, 103 (1999) 7445-7453. [662] C.S. Brauer, M.B. Craddock, K.J. Higgins, K.R. Leopold, Mol. Phys., 105 (2007) 613625.

us

[664] R. Kugel, H. Taube, J. Phys. Chem., 79 (1975) 2130-2135.

cr

213-218.

ip t

[663] S.W. Hunt, D.L. Fiacco, M. Craddock, K.R. Leopold, J. Mol. Spectrosc., 212 (2002)

[665] A. Givan, A. Loewenschuss, C.J. Nielsen, J. Mol. Struct., 604 (2002) 147-157.

an

[666] A. Givan, L.A. Larsen, A. Loewenschuss, C.J. Nielsen, J. Chem. Soc., Farad. Trans., 94 (1998) 2277-2286.

M

[667] A. Givan, A. Loewenschuss, C.J. Nielsen, Phys. Chem. Chem. Phys., 1 (1999) 37-43.

d

[668] A. Givan, L.A. Larsen, A. Loewenschuss, C.J. Nielsen, J. Mol. Struct., 509 (1999) 35-

Ac ce pt e

47.

[669] Y.K. Choe, E. Tsuchida, T. Ikeshoji, Int. J. Quantum Chem., 109 (2009) 1984-1990. [670] A. Givan, L.A. Larsen, A. Loewenschuss, C.J. Nielsen, J. Chem. Soc., Farad. Trans., 94 (1998) 827-835.

[671] M. Rozenberg, A. Loewenschuss, J. Phys. Chem. A, 113 (2009) 4963-4971. [672] M. Rozenberg, A. Loewenschuss, C.J. Nielsen, Phys. Chem. Chem. Phys., 12 (2010) 4024-4031. [673] M. Rozenberg, A. Loewenschuss, C.J. Nielsen, J. Phys. Chem. A, 115 (2011) 57595766. [674] A. Givan, A. Loewenschuss, C.J. Nielsen, J. Phys. Chem. A, 104 (2000) 3441-3445. 200 Page 200 of 213

[675] A. Givan, H. Grothe, A. Loewenschuss, C.J. Nielsen, Phys. Chem. Chem. Phys., 4 (2002) 255-263. [676] A. Givan, H. Grothe, A. Loewenschuss, J. Mol. Struct., 648 (2003) 159-169.

ip t

[677] A. Givan, A. Loewenschuss, C.J. Nielsen, J. Mol. Struct., 748 (2005) 77-90. [678] Y. Miller, G.M. Chaban, R.B. Gerber, J. Phys. Chem. A, 109 (2005) 6565-6574.

cr

[679] A.J. Bienko, Z. Latajka, Chem. Phys. Lett., 374 (2003) 577-582.

us

[680] S. Brutti, L. Bencivenni, V. Barbarossa, S. Sau, G. De Maria, J. Chem. Thermodyn., 38

an

(2006) 1292-1300.

[681] B.N. Ida, P.S. Fudacz, D.H. Pulsifer, J.M. Standard, J. Phys. Chem. A, 110 (2006)

M

5831-5838.

d

[682] C.S. Sass, B.S. Ault, J. Phys. Chem., 89 (1985) 1002-1006.

Ac ce pt e

[683] K. Jaeger, R. Weller, O. Schrems, Ber. Bunsenges. Phys. Chem., 96 (1992) 485-488. [684] S.A. Peebles, R.L. Kuczkowski, The Journal of Physical Chemistry A, 103 (1999) 3884-3889.

[685] S. Pakhira, C. Sahu, K. Sen, A.K. Das, Chem. Phys. Lett., 549 (2012) 6-11. [686] G.D. Brabson, Z. Mielke, L. Andrews, J. Phys. Chem., 95 (1991) 79-86. [687] P. Hassanzadeh, L. Andrews, J. Phys. Chem., 96 (1992) 6579-6585. [688] G.D. Brabson, L. Andrews, J. Phys. Chem., 96 (1992) 9172-9177. [689] P. Hassanzadeh, C. Thompson, L. Andrews, J. Phys. Chem., 96 (1992) 8246-8249. [690] G.D. Brabson, L. Andrews, C.J. Marsden, J. Phys. Chem., 100 (1996) 16487-16494.

201 Page 201 of 213

[691] G.D. Brabson, A. Citra, L. Andrews, R.D. Davy, M. Neurock, J. Am. Chem. Soc., 118 (1996) 5469-5473. [692] K. Sundararajan, K. Sankaran, V. Kavitha, J. Mol. Struct., 876 (2008) 240-249.

ip t

[693] G. Orlova, J.D. Goddard, J. Phys. Chem. A, 103 (1999) 6825-6834. [694] M. Bahou, L. Schriver-Mazzuoli, C. Camy-Peyret, A. Schriver, T. Chiavassa, J.P.

cr

Aycard, Chem. Phys. Lett., 265 (1997) 145-153.

us

[695] H. Beckers, H. Willner, M.E. Jacox, ChemPhysChem, 10 (2009) 706-710.

an

[696] H. Beckers, X.Q. Zeng, H. Willner, Chem. Eur. J., 16 (2010) 1506-1520. [697] H. Beckers, P. Garcia, H. Willner, G.A. Arguello, C.J. Cobos, J.S. Francisco, Angew.

M

Chem. Int. Ed., 46 (2007) 3754-3757.

d

[698] A.C. Legon, Angew. Chem. Int. Ed., 38 (1999) 2687-2714. [699] A.C. Legon, The interaction of dihalogens and hydrogen halides with Lewis bases in

Ac ce pt e

the gas phase: An experimental comparison of the halogen bond and the hydrogen bond, in: Halogen Bonding: Fundamentals and Applications, 2008, pp. 17-64. [700] A.C. Legon, Phys. Chem. Chem. Phys., 12 (2010) 7736-7747. [701] B.S. Ault, J. Mol. Struct., 127 (1985) 343-356. [702] N.P. Machara, B.S. Ault, J. Phys. Chem., 91 (1987) 2046-2050. [703] U.P. Agarwal, A.J. Barnes, W.J. Orville-Thomas, Can. J. Chem., 63 (1985) 1705-1707. [704] N.P. Machara, B.S. Ault, Inorg. Chem., 27 (1988) 2383-2385. [705] S.A. Cooke, G.K. Corlett, C.M. Evans, A.C. Legon, J.H. Holloway, J. Chem. Phys., 108 (1998) 39-45. 202 Page 202 of 213

[706] H.I. Bloemink, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 251 (1996) 275-286. [707] H.I. Bloemink, K. Hinds, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 242 (1995) 113-120.

ip t

[708] H.I. Bloemink, S.J. Dolling, K. Hinds, A.C. Legon, J. Chem. Soc., Farad. Trans., 91 (1995) 2059-2066.

cr

[709] X.Y. Zhang, Y.L. Zeng, X.Y. Li, L.P. Meng, S.J. Zheng, Structural Chemistry, 22

us

(2011) 567-576.

an

[710] S. Beccaceci, J.S. Ogden, J.M. Dyke, Phys. Chem. Chem. Phys., 12 (2010) 2075-2082. [711] S. Beccaceci, N. Armata, J.S. Ogden, J.M. Dyke, L. Rhyman, P. Ramasami, Phys.

M

Chem. Chem. Phys., 14 (2012) 2399-2407.

[712] L. Rhyman, N. Armata, P. Ramasami, J.M. Dyke, J. Phys. Chem. A, 116 (2012) 5595-

Ac ce pt e

d

5603.

[713] N.P. Machara, B.S. Ault, J. Phys. Chem., 92 (1988) 2439-2442. [714] H.I. Bloemink, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 254 (1996) 59-68. [715] C. Domene, P.W. Fowler, A.C. Legon, Chem. Phys. Lett., 309 (1999) 463-470. [716] H.I. Bloemink, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 248 (1996) 260-268.

[717] E.R. Waclawik, A.C. Legon, J.H. Holloway, Chem. Phys. Lett., 295 (1998) 289-297. [718] M.D. Page, E.R. Waclawik, J.H. Holloway, A.C. Legon, J. Mol. Struct., 509 (1999) 5565. [719] K. Hinds, A.C. Legon, J.H. Holloway, Mol. Phys., 88 (1996) 673-682. 203 Page 203 of 213

[720] A. Karpfen, J. Phys. Chem. A, 105 (2001) 2064-2072. [721] L. Andrews, R. Lascola, J. Am. Chem. Soc., 109 (1987) 6243-6247. [722] L. Andrews, R. Withnall, Inorg. Chem., 28 (1989) 494-499.

ip t

[723] L. Andrews, J. Phys. Chem., 88 (1984) 2940-2949.

cr

[724] W.F. Howard, Jr., L. Andrews, J. Am. Chem. Soc., 95 (1973) 3045-3046.

[725] E.Y. Misochko, I.U. Goldschleger, A.V. Akimov, C.A. Wight, Low Temp. Phys., 26

us

(2000) 727-735.

an

[726] A.V. Akimov, I.U. Goldschleger, E.Y. Misochko, C.A. Wight, J. Phys. Chem. A, 106 (2002) 9756-9760.

d

598-604.

M

[727] H.I. Bloemink, K. Hinds, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 245 (1995)

134.

Ac ce pt e

[728] I.U. Goldschleger, A.V. Akimov, E.Y. Misochko, Mendeleev Commun., (1999) 132-

[729] I.U. Goldschleger, A.V. Akimov, E.Y. Misochko, J. Mol. Struct., 519 (2000) 191-198. [730] E.Y. Misochko, I.U. Gol'dshleger, A.V. Akimov, Russ. Chem. Bull., 49 (2000) 829835.

[731] Z. Latajka, S. Scheiner, Y. Bouteiller, H. Ratajczak, J. Mol. Struct., 376 (1996) 343351.

[732] A. Karpfen, Chem. Phys. Lett., 299 (1999) 493-502. [733] A. Karpfen, J. Phys. Chem. A, 104 (2000) 6871-6879. [734] A. Karpfen, Chem. Phys. Lett., 316 (2000) 483-488. 204 Page 204 of 213

[735] H. Feng, W.G. Sun, Y.M. Xie, H.F. Schaefer, Chem. Asian J., 6 (2011) 3152-3156. [736] L. Fredin, B. Nelander, G. Ribbergard, Chem. Phys., 12 (1976) 153-159. [737] A.C. Legon, D.G. Lister, J.C. Thorn, J. Chem. Soc., Chem. Comm., (1994) 757-758.

3212.

cr

[739] H.I. Bloemink, A.C. Legon, J. Chem. Phys., 103 (1995) 876-882.

ip t

[738] A.C. Legon, D.G. Lister, J.C. Thorn, J. Chem. Soc., Farad. Trans., 90 (1994) 3205-

us

[740] H.I. Bloemink, A.C. Legon, J.C. Thorn, J. Chem. Soc., Farad. Trans., 91 (1995) 781-

an

787.

[741] E.R. Waclawik, A.C. Legon, Phys. Chem. Chem. Phys., 1 (1999) 4695-4700.

M

[742] N.P. Machara, B.S. Ault, J. Phys. Chem., 92 (1988) 73-77.

d

[743] N.P. Machara, B.S. Ault, J. Mol. Struct., 172 (1988) 129-138.

Ac ce pt e

[744] A.C. Legon, H.E. Warner, J. Chem. Phys., 98 (1993) 3827-3832. [745] E.R. Waclawik, A.C. Legon, Chem. Eur. J., 6 (2000) 3968-3975. [746] A.C. Legon, J.M.A. Thumwood, E.R. Waclawik, J. Chem. Phys., 113 (2000) 52785286.

[747] J.B. Davey, A.C. Legon, E.R. Waclawik, Phys. Chem. Chem. Phys., 2 (2000) 22652269.

[748] T.C. McInnis, L. Andrews, J. Phys. Chem., 96 (1992) 2051-2059. [749] S.A. Cooke, G. Cotti, J.H. Holloway, A.C. Legon, Angew. Chem. Int. Ed. Engl., 36 (1997) 129-130.

205 Page 205 of 213

[750] S.A. Cooke, G. Cotti, C.M. Evans, J.H. Holloway, Z. Kisiel, A.C. Legon, J.M.A. Thumwood, Chem. Eur. J., 7 (2001) 2295-2305. [751] W.F. Wang, J.M. Sirota, D.C. Reuter, J. Mol. Struct. (Theochem), 541 (2001) 31-37. [752] H. Beckers, J. Breidung, H. Burger, R. Kuna, A. Rahner, W. Schneider, W. Thiel, J.

ip t

Chem. Phys., 93 (1990) 4603-4614.

cr

[753] L. Andrews, T.C. McInnis, Y. Hannachi, J. Phys. Chem., 96 (1992) 4248-4254.

us

[754] G. Cotti, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 264 (1997) 513521.

an

[755] L. Andrews, T.C. McInnis, Inorg. Chem., 30 (1991) 2990-2993.

M

[756] H. Bai, B.S. Ault, J. Mol. Struct., 196 (1989) 47-56.

d

[757] B.S. Ault, J. Phys. Chem., 93 (1989) 279-282.

Ac ce pt e

[758] H.B. Bai, B.S. Ault, J. Phys. Chem., 95 (1991) 3080-3084. [759] B.S. Ault, J. Phys. Chem., 90 (1986) 2825-2829. [760] K. Hinds, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 93 (1997) 373-378. [761] S.A. Cooke, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 93 (1997) 42534258.

[762] B.S. Ault, J. Phys. Chem., 91 (1987) 4723-4727. [763] H.I. Bloemink, K. Hinds, A.C. Legon, J.H. Holloway, J. Chem. Soc., Chem. Comm., (1995) 1833-1834. [764] K. Hinds, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 92 (1996) 12911296. 206 Page 206 of 213

[765] H.I. Bloemink, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 250 (1996) 567-575. [766] A.C. Legon, Chem. Eur. J., 4 (1998) 1890-1897. [767] S.M. Resende, W.B. DeAlmeida, Mol. Phys., 91 (1997) 635-641.

cr

[769] W.B. Dealmeida, J.S. Craw, Mol. Phys., 78 (1993) 1351-1364.

ip t

[768] H. Bai, B.S. Ault, J. Phys. Chem., 94 (1990) 199-203.

(1998) 2295-2302.

an

[771] A.C. Legon, Chem. Soc. Rev., 22 (1993) 153-163.

us

[770] S.A. Cooke, C.M. Evans, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 94

M

[772] A.C. Legon, Chem. Comm., (1996) 109-116.

d

[773] A.C. Legon, Chem. Comm., (1998) 2585-2586.

Ac ce pt e

[774] A.C. Legon, Chem. Comm., (1998) 2737-2738. [775] A.C. Legon, Chem. Phys. Lett., 237 (1995) 291-298. [776] A.C. Legon, Chem. Phys. Lett., 279 (1997) 55-64. [777] A.C. Legon, J. Chem. Soc., Farad. Trans., 91 (1995) 1881-1883. [778] A.C. Legon, Chem. Phys. Lett., 314 (1999) 472-480. [779] O.K. Poleshchuk, V. Branchadell, B. Brycki, A.V. Fateev, A.C. Legon, J. Mol. Struct. (Theochem), 760 (2006) 175-182. [780] Q.Z. Li, Q.Q. Lin, W.Z. Li, J.B. Cheng, B. Gong, J.Z. Suo, ChemPhysChem, 9 (2008) 2265-2269. [781] X.L. An, B. Jing, Q.Z. Li, J. Phys. Chem. A, 114 (2010) 6438-6443. 207 Page 207 of 213

[782] Q.Z. Li, R. Li, X.F. Liu, W.Z. Li, J.B. Cheng, ChemPhysChem, 13 (2012) 1205-1212. [783] S.W. Bunte, J.B. Miller, Z.S. Huang, J.E. Verdasco, C. Wittig, R.A. Beaudet, J. Phys. Chem., 96 (1992) 4140-4143.

ip t

[784] W. Jäger, Y. Xu, M.C.L. Gerry, J. Phys. Chem., 97 (1993) 3685-3689. [785] J.B. Davey, A.C. Legon, E.R. Waclawik, Phys. Chem. Chem. Phys., 1 (1999) 3097-

cr

3101.

us

[786] E.R. Waclawik, J.M.A. Thumwood, D.G. Lister, P.W. Fowler, A.C. Legon, Mol. Phys., 97 (1999) 159-166.

an

[787] K. Hinds, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 242 (1995) 407-414.

M

[788] S. Blanco, A.C. Legon, J.C. Thorn, J. Chem. Soc., Farad. Trans., 90 (1994) 1365-1371. [789] A. Schriver, L. Schriver-Mazzuoli, P. Chaquin, M. Bahou, J. Phys. Chem. A, 103

d

(1999) 2624-2631.

Ac ce pt e

[790] R.M. Romano, A.J. Downs, J. Phys. Chem. A, 107 (2003) 5298-5305. [791] M.S. Chiappero, G.A. Argüello, P. García, H. Pernice, H. Willner, H. Oberhammer, K.A. Peterson, J.S. Francisco, Chem. Eur. J., 10 (2004) 917-924. [792] P. García, H. Willner, H. Oberhammer, J.S. Francisco, J. Chem. Phys., 121 (2004) 11900-11906.

[793] R.M. Romano, C.O. Della Védova, A.J. Downs, Y.A. Tobón, H. Willner, Inorg. Chem., 44 (2005) 3241-3248. [794] Y.A. Tobón, L.I. Nieto, R.M. Romano, C.O. Della Védova, A.J. Downs, J. Phys. Chem. A, 110 (2006) 2674-2681.

208 Page 208 of 213

[795] A.L. Picone, C.O. Della Védova, H. Willner, A.J. Downs, R.M. Romano, Phys. Chem. Chem. Phys., 12 (2010) 563-571. [796] J.A. Gómez Castaño, A.L. Picone, R.M. Romano, H. Willner, C.O. Della Védova, Chem. Eur. J., 13 (2007) 9355-9361.

ip t

[797] J.A. Gómez Castaño, R.M. Romano, H. Willner, C.O. Della Védova, Inorg. Chim. Acta, 361 (2008) 540-550.

us

cr

[798] R.M. Romano, A.L. Picone, A.J. Downs, J. Phys. Chem. A, 110 (2006) 12129-12135. [799] Y.A. Tobón, R.M. Romano, C.O. Della Védova, A.J. Downs, Inorg. Chem., 46 (2007)

an

4692-4703.

[800] A.L. Picone, R.M. Romano, J. Mol. Struct., 978 (2010) 187-190.

M

[801] J.J. Turner, G.C. Pimentel, Science, 140 (1963) 974-975.

d

[802] M. Pettersson, J. Lundell, M. Räsänen, J. Chem. Phys., 103 (1995) 205-210.

Ac ce pt e

[803] M. Pettersson, J. Lundell, M. Räsänen, J. Chem. Phys., 102 (1995) 6423-6431. [804] L. Khriachtchev, M. Räsänen, R.B. Gerber, Acc. Chem. Res., 42 (2009) 183-191. [805] J.F. Rooms, PhD thesis, The University of Hull, 2005. [806] R.N. Perutz, J.J. Turner, J. Am. Chem. Soc., 97 (1975) 4791-4800. [807] Y.Y. Zhao, Y. Gong, M.H. Chen, C.F. Ding, M.F. Zhou, J. Phys. Chem. A, 109 (2005) 11765-11770. [808] Y.Y. Zhao, G.J. Wang, M.H. Chen, M.F. Zhou, J. Phys. Chem. A, 109 (2005) 66216623. [809] R. Yang, Y. Gong, M.F. Zhou, Chem. Phys., 340 (2007) 134-140. 209 Page 209 of 213

[810] R. Yang, Y. Gong, H. Zhou, M.F. Zhou, J. Phys. Chem. A, 111 (2007) 64-70. [811] Y.Y. Zhao, M.F. Zhou, Science China-Chemistry, 53 (2010) 327-336. [812] I. Infante, L. Andrews, X.F. Wang, L. Gagliardi, Chem. Eur. J., 16 (2010) 12804-

ip t

12807. [813] L. Khriachtchev, M. Pettersson, N. Runeberg, J. Lundell, M. Räsänen, Nature, 406

cr

(2000) 874-876.

us

[814] W. Grochala, L. Khriachtchev, M. Räsänen, Noble gas chemistry, in: L. Khriachtchev (Ed.) Physics and Chemistry at Low Temperatures, Pan Stanford Publishing, Singapore,

an

2011.

[815] E. Tsivion, S. Zilberg, R.B. Gerber, Chem. Phys. Lett., 460 (2008) 23-26.

d

Chem. A, 116 (2012) 4510-4517.

M

[816] M. Tsuge, S. Berski, R. Stachowski, M. Rasanen, Z. Latajka, L. Khriachtchev, J. Phys.

Ac ce pt e

[817] A. Maki, T. Masiello, T.A. Blake, J. Mol. Struct., 742 (2005) 3-20. [818] T. Masiello, A. Maki, T.A. Blake, J. Mol. Spectrosc., 234 (2005) 122-136. [819] R. Kirkpatrick, T. Masiello, A. Weber, J.W. Nibler, J. Mol. Spectrosc., 237 (2006) 97103.

[820] R.L. Hunt, B.S. Ault, Spectrochim. Acta A, 37 (1981) 63-69. [821] S. Kvisle, E. Rytter, Spectrochim. Acta A, 40 (1984) 939-951. [822] L. Andrews, R.C. Spiker, JR. , J. Phys. Chem., 76 (1972) 3208-3213. [823] L. Schriver-Mazzuoli, A. Schriver, C. Lugez, A. Perrin, C. Camy-Peyret, J.M. Flaud, J. Mol. Spectrosc., 176 (1996) 85-94.

210 Page 210 of 213

[824] H.I. Bloemink, S.A. Cooke, J.H. Holloway, A.C. Legon, Angew. Chem. Int. Ed. Engl., 36 (1997) 1340-1342. [825] S.A. Cooke, G. Cotti, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 262 (1996) 308-314.

cr

[827] K. Hinds, A.C. Legon, Chem. Phys. Lett., 240 (1995) 467-473.

ip t

[826] A.C. Legon, J.C. Thorn, J. Chem. Soc., Farad. Trans., 89 (1993) 4157-4162.

us

[828] W.A. Herrebout, A.C. Legon, E.R. Waclawik, Phys. Chem. Chem. Phys., 1 (1999) 4961-4966.

an

[829] G. Cotti, S.A. Cooke, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 260 (1996) 388-394.

d

Soc., Farad. Trans., 92 (1996) 2671-2676.

M

[830] S.A. Cooke, G. Cotti, K. Hinds, J.H. Holloway, A.C. Legon, D.G. Lister, J. Chem.

411.

Ac ce pt e

[831] J.B. Davey, A.C. Legon, E.R. Waclawik, J. Mol. Struct. (Theochem), 500 (2000) 403-

[832] J.B. Davey, A.C. Legon, J.M.A. Thumwood, J. Chem. Phys., 114 (2001) 6190-6202. [833] A.C. Legon, J.M.A. Thumwood, E.R. Waclawik, Chem. Eur. J., 8 (2002) 940-950. [834] S.A. Cooke, G. Cotti, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Comm., (1996) 2327-2328.

[835] J.B. Davey, A.C. Legon, Phys. Chem. Chem. Phys., 3 (2001) 3006-3011. [836] J.B. Davey, A.C. Legon, E.R. Waclawik, Phys. Chem. Chem. Phys., 2 (2000) 16591665. [837] A.C. Legon, J.M.A. Thumwood, Phys. Chem. Chem. Phys., 3 (2001) 2758-2764. 211 Page 211 of 213

[838] H.I. Bloemink, A.C. Legon, Chem. Eur. J., 2 (1996) 265-270. [839] A.C. Legon, E.R. Waclawik, Chem. Phys. Lett., 312 (1999) 385-393. [840] S.A. Cooke, G.K. Corlett, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 94

ip t

(1998) 2675-2680. [841] S.A. Cooke, G.K. Corlett, C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett.,

cr

275 (1997) 269-277.

us

[842] C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 267 (1997) 281-287.

an

[843] S.A. Cooke, A.C. Legon, J.H. Holloway, J. Mol. Struct., 406 (1997) 15-21. [844] G. Cotti, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 255 (1996) 401-409.

M

[845] C.M. Evans, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 255 (1996) 119-128.

d

[846] S.D. Springer, L.A. Rivera-Rivera, B.A. McElmurry, Z.C. Wang, Leonov, II, R.R.

Ac ce pt e

Lucchese, A.C. Legon, J.W. Bevan, J. Phys. Chem. A, 116 (2012) 1213-1223. [847] H.I. Bloemink, K. Hinds, A.C. Legon, J.C. Thorn, Chem. Phys. Lett., 223 (1994) 162166.

[848] H.I. Bloemink, S.A. Cooke, K. Hinds, A.C. Legon, J.C. Thorn, J. Chem. Soc., Farad. Trans., 91 (1995) 1891-1900.

[849] J.B. Davey, A.C. Legon, Chem. Phys. Lett., 350 (2001) 39-50. [850] H.I. Bloemink, K. Hinds, A.C. Legon, J.C. Thorn, J. Chem. Soc., Chem. Comm., (1994) 1229-1230. [851] J.B. Davey, A.C. Legon, Phys. Chem. Chem. Phys., 1 (1999) 3721-3726.

212 Page 212 of 213

[852] H.I. Bloemink, K. Hinds, A.C. Legon, J.C. Thorn, J. Chem. Soc., Chem. Comm., (1994) 1321-1322. [853] H.I. Bloemink, K. Hinds, A.C. Legon, J.C. Thorn, Chem. Eur. J., 1 (1995) 17-25.

ip t

[854] A.C. Legon, J.M.A. Thumwood, Phys. Chem. Chem. Phys., 3 (2001) 1397-1402. [855] H.I. Bloemink, K. Hinds, A.C. Legon, J.C. Thorn, Angew. Chem. Int. Ed. Engl., 33

cr

(1994) 1512-1513.

us

[856] J.M.A. Thumwood, A.C. Legon, Chem. Phys. Lett., 310 (1999) 88-96.

[857] J.B. Davey, J.H. Holloway, A.C. Legon, E.R. Waclawik, Phys. Chem. Chem. Phys., 1

an

(1999) 2415-2420.

M

[858] S.A. Cooke, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 298 (1998) 151-160.

d

[859] S.A. Cooke, J.H. Holloway, A.C. Legon, Chem. Phys. Lett., 266 (1997) 61-69.

2365.

Ac ce pt e

[860] S.A. Cooke, J.H. Holloway, A.C. Legon, J. Chem. Soc., Farad. Trans., 93 (1997) 2361-

[861] J.B. Davey, A.C. Legon, E.R. Waclawik, Chem. Phys. Lett., 346 (2001) 103-111. [862] J.B. Davey, A.C. Legon, E.R. Waclawik, Chem. Phys. Lett., 306 (1999) 133-144.

213 Page 213 of 213