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Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 57 4. Main-Group Metals and Organometallic Compounds The information we will be considering is somewhat spread-out in Shriver-Atkins. The s and p-block metals are discussed in Chapter 9 (sections 9.1 – 9.5; 9.12 – 9.18) and Chapter 15 (all of this chapter). Be sure to read these sections of the text and study the example problems that are liberally spread through these sections. This presentation of the periodic table serves to remind us that 80% of all elements are metals. We dedicate a separate course (Chem3810) to consider primarily the chemistry of the transition elements, i.e. those that have partly filled d orbitals in at least one common oxidation state. This excludes the elements of Group 12. Also, silver in group 11 is dominated by chemistry of the +1 oxidation state, so it too does not behave as a typical transition metal. The chemistry of copper in the +1 oxidation state, of Ag(I), Zn(II), Cd(II) and Hg(II) will therefore be briefly considered here by lumping them with the p-block metals. The chemistry of the lanthanides (the first f-series) is closely related to that of the Group 3 elements, and like these elements the dominant oxidation state is always 3+. Thus these elements are pseudo-noble gas ions, and their chemistry is predominantly that of ionic and weakly coordinated species analogous to the Group 2 elements. On the other hand, the actinides behave more as the d-block elements do, and there is an extensive chemistry particularly of uranium, since the 238 U isotope is long-lived and the element is not too radioactive to preclude conventional chemical studies. In this course we will focus on the Group 1 and 2 metals and the p-block elements, with an emphasis on their covalent derivatives. 4.1 Some basic properties of the metallic elements 4.1.1 Enthalpies of Vaporization of the metallic elements (kJ/mol) The graphic at right is a bar graph of the enthalpies of vaporization of the metallic elements in kJ/mol. For metallic lattices, this data represents the strength of the metal-metal interactions in the solid-state metallic lattice. We recognize that this distribution has the general trend of being low at the left and the right-hand side. It peaks in the middle, which corresponds to d-block elements with a half-filled or close to half-filled d-valence orbital. Thus the resistance of tungsten towards evaporation has been attributed to covalent d-d bonding. Consistent with this interpretation is that tungsten is one of the most brittle of the metallic elements. It is a conductor, and it is correct to call it a metal, but it is a metal with strong covalent W–W bonding in addition to the metallic bond that allows for its metallic properties. This aspect of tungsten chemistry accounts for its ubiquitous role as the filament material of incandescent light bulbs. The s and p block metals, on the other hand, tend to be the most volatile. The larger enthalpies of vaporization for the pblock metals reflects the presence of partly filled p-orbitals which can lead to stable E–E bonding. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 58 4.1.2 Electron-Deficient Metal Compounds: Clusters Metals are electron deficient species, and the metallic bond that leads to the classic properties of metals (high thermal and electrical conductivity, malleability, metallic sheen, typical silver colours) are a consequence of the delocalized bonding that occurs as they attempt to form the best bonding patterns that are available with the limited number of electrons that are present. It should therefore not be a surprise that when these elements are reacted with fewer ligands than required by the octet rule or the 18-electron rules they tend to form cluster compounds. These can be thought of naively as small chunks of the metallic element, with the structure usually adopted at room temperature for bulk metal, surrounded on the outside by a suitable set of ligands. The distribution of such cluster compounds over the periodic table is shown in the following graphic: There are several element regions where cluster formation is common. We have already mentioned simple boron derivatives that are electron deficient (e.g. the as seen in self-dimerization of BH3 ). There are also cage structures for the hydrocarbons, such as tetrahedrane, C4 H4 , and cubane, C8 H8 , but these are exceptions rather than the rule. Also, they are not truly electron deficient. Clusters are also common for organometallic derivatives of lithium and sodium (see below). With Rb and Cs the formation of suboxides is common, and there is an extensive fragment population among the transition elements. The heavy p-block metals tend to form aggregated structures. We will likely not have time to discuss this important class of elements. 4.2 s-Block Metals The crustal abundances of the s-block elements are shown in the graphic at right (note that the numbers in this bar graph are the logarithms of the abundances.) Sodium and potassium are high abundance minerals, and both are mined in large quantities. Magnesium and calcium are largely present as a part of rocks and minerals (limestone, dolomite, etc.) Cesium and beryllium are expensive both because of their lower abundance and because of difficulties in handling. Cesium must be stored in an inert atmosphere or under vacuum. Beryllium is highly toxic. The latter element finds an extremely important application as the most X-ray transparent element with reasonable structural strength. It is universally employed as the “window” on X-ray tubes (for both medical and scientific applications) that allows the X-rays flux to emanate at high intensity from the vacuum chambers in which they are generated. In Chemistry 2810 we focused extensively on the reactivity of the metallic elements of Group 1. In aqueous solution, lithium has the most negative standard reduction potential, but cesium has the lowest ionization potential. How can this apparent contradiction be squared? It is of course due to the difference in solvation between these elements in aqueous solution. The thermodynamic cycles shown below compare lithium and cesium in the elemental form with both the gas phase and aqueous phase ions. Although lithium has almost 150% the ionization energy of cesium, it also has a correspondingly higher enthalpy of hydration. The net balance between these two terms is that the solution ionization energy of these elements are remarkably similar, and this is true for the intervening Group 1 elements as well. However, on balance the solution ionization energy (which correlates closely to the redox potentials) is greater for lithium. The large value of the lithium hydration energy is attributed to its very small size (r(Li+) = 0.90 Å for CN6). This is just one more example of the general principle that the second-row elements have distinctly different properties from those of the subsequent elements within any chemical group. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 59 Thermochemical cycles for oxidation of Li and Cs 4.2.1 Complex Formation: Crown ethers and cryptands Modern developments in the chemistry of the s-block elements have largely dealt with non-aqueous environments. In order to dissolve these reactive elements in suitable solvents, it is essential to supply appropriate ligands that stabilize their charges and facilitate their dissolution. Typical solvents for this work have included ethers, amines and aromatic hydrocarbons (benzene, toluene). By far the most important complexes are those formed between the metal ions and crown ethers and cryptand ligands. The crowns, developed by Peterson and later by Cram, are cyclic polyethers that have the maximum chelate effect, and whose size can be adjusted by the number of donor atoms and the spacing of the oxygen atoms by ajustment of the number of CH2 groups in the rings. They can thus be tuned to select out different ions based on their sizes. The crown effect can be optimized even further by the cryptand ligands, which combine crown ethers with trialkylamine functionality. The nitrogen centers are three coordinate, turning rings into cages. The cryptand ligands have been extensively developed into selective chelating agents for the s-block elements. The structures of come examples are provided below, and the graph at right maps the affinity for 2.2.1 and 2.2.2 crypts for the Group 1 metal cations in terms of the formation constants for complex formation (again expressed as logarithms). The former is optimally sized to coordinate sodium ions, and the latter is best for potassium ions. 4.2.2 Ammoniacal solutions of alkali metals: Electrides All the Group 1 metals, as well as the more reactive Group 2 metals Ca, Sr, Ba dissolve in anhydrous liquid NH3 . (NOTE: not the common aqueous ammonia that is widely used in the chemistry laboratory.) Ammonia is a gas at room temperature, but it is liquefied relatively easily by compression, and the liquid can be dispensed if kept at or below the boiling point of –33.4 °C. Ammonia, like water, has a high enthalpy of vaporization, and open flasks of the liquid can be Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 60 handled for reasonable periods of time. For work with electrides, it is essential to keep the solutions dry and a typical approach to handling them is to use a dry-ice condenser in an apparatus protected by nitrogen or with a drying tube. Note that metallic sodium is the standard reagent for drying ammonia, the main purification method that is required. Indeed, this is a self-indicating drying agent, for so long as the blue colour of the electride persists, the ammonia must be dry. Typical apparatus for purifying and handling liquid ammonia are shown in the figure at right. The cooling agent most commonly used is solid dry ice suspended in a suitable heattransfer fluid (methanol is effective). The solutions of metals in liquid ammonia develop an intense blue colour. From such solutions, the Group 1 metals can be recovered unchanged (contrast this with their behaviour in water!) The Group 2 metals are recovered as [M(NH3 )6 ] complexes, a form of expanded metal. From ammoniacal solutions of lithium and sodium, the complexes [Li(NH3 )6 ] (yellow) and [Na(NH3 )6 ] (blue) can be isolated at low temperature, but on warming these release ammonia and revert to finely divided metal (CAUTION: highly reactive state!) The blue colour of such solutions is attributed to the formation of electrides. That is, a solvated free electron: liquid NH 3 , bp M → M +( solvate ) + e−( solvate) The solvated electron exist in cavities of 3-4 Å radius as determined by the volume expansion on additon of the metals to the liquid. The origin of the colour apparently is due to quantization of the solvated electron. The most popular model has been to solve Schrödinger’s equation for an electron in a spherical box (related to the 3-D cubic box developed in Chemistry 3730.) The results are a set of wavefunctions closely related to those of the hydrogen atom, but with different quantum number rules. The solution can be graphed as shown in the picture at right. The wavefunctions are labeled the same as the H-atom orbitals for convenience, but note that the pattern of energies is different. It is assumed that the blue colour is due to a 1s→2p transition (absorption) leading to the observed λmax at ~1500 nm in the near infra-red. The blue colour is that of the reflected wavelengths. This wavelength is not a property of the alkali element used to generate the electride, but is a property of the medium, as shown by the following spectra: Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 61 As expected for solutions containing free electrons, they have distinctive EPR spectra. The spectra from liquid ammonia solutions are hard to measure, and somewhat uninformative. However, using organic amines it is possible to obtain more meaningful spectra, and the spectrum shown at right is an example of this. Note that the observed spectrum seems to be a result of several overlapping sets of signals, implying that in this solvent the electron has more than just one environment. The observed hyper-fine coupling is to the 14 N nuclei of the amine solvent. Electride solutions are powerful reducing agents. In the Chemistry 3810 laboratory they are used to reductively cleave a phosphorus carbon bond. There are many practical applications for electrides in synthetic chemistry. The solutions are metastable, and they can be decomposed, especially when catalyzed with iron salts. The presence of rust in most chemical laboratories requires special care that vessels for handling electrides are kept extremely clean. The decomposition reaction is analogous to the reaction of the elements with water, producing the corresponding alkali or alkaline earth amides. The most commonly used amide is NaNH2 . It is a colourless ionic compound, but is very hygroscopic and reacts instantly with water to produce sodium hydroxide (solvent leveling). In non-aqueous solvents it acts as a stronger base than sodium hydroxide. 2 NH 3 + 2e− → 2NH 2− + H2 4.2.3 Ether, alkylamine and cryptand solutions of alkali metals: alkalides Another way to generate electrides is to add an excess of crown ether or cryptand in a suitable solvent to which is added an alkali metal. This produces a solution of the type [M(crown)n ]+e– . An example of such an adduct of cesium, [Cs(18-crown-6)2 ]+e– , is shown in a figure below. However, when only half an equivalent of a crown or cryptand is added, a different reaction takes place with elemental alkali metal. For example: 2 Na solvent + crypt − [222] → [ Na (crypt − [222]] + Na − In this case, the free electron is trapped by another sodium atom, and the result is the formation of the sodide ion. Such anionic alkali elements are called alkalides. Recall from Chem2810 that the alkali elements have exothermic electron gain enthalpies: H –72.77 Li –59.63 Electron-Gain Enthalpy Values for Some Elements (kJ/mol)* Be +48† B –26.7 C –121.85 N +7 Na –52.87 Mg +39 Al –42.6 Si –133.6 P –72.07 K –48.39 Rb –46.89 Cs –45.51 Ca +30 Sr +30 Ba 0 Ga –30 In –30 Tl –20 Ge –120 Sn –120 Pb –35.1 As –78 Sb –103 Bi –91.3 (the first entry refers to the formation of X – from X; the first entry refers to the formation of X 2– from X – ) O –140.98 +844 S –200.41 +532 Se –194.97 Te –190.16 Po –180 F –328.0 Cl –349.0 Br –324.7 I –295.16 At –270 *Data taken from H. Hotop and W. C. Lineberger: Journal of Physical Chemistry, Reference Data, Vol. 14, p. 731, 1985. (This paper also includes data for the transition metals.) Some values are known to more than two decimal places. † Elements with a positive electron-gain enthalpy indicate that a stable anion A of the element does not exist in the gas phase. Chemistry 3810 Lecture Notes This means that the reaction: Dr. R. T. Boeré Na( g ) Page 62 + e − → Na(−g ) is favoured. Usually this is not a favourable reaction in solution, but the cryptand ligand reverses normal solvation trends. Adducts of this type have even been crystallized, as shown in the following figures, which depict structures from X-ray diffraction studies of [Na(crypt-[222])]+Na – . Note the large size of the Na – ion, as expected based on shielding consideration. Note that alkalides are the formal analogues to the hydride ion in Group 1. + – [Na(crypt-[222])] Na 4.3 + – [Cs(18-crown-6)2] e p-Block Metals The Group 11 and Group 12 metals are formally analogous to the Group 1 and 2 elements, respectively, but differ due to the presence of a poorly shielding d 10 filled shell of electrons. Thus, they have uniformly higher ionization energy, and while cesium is one of the most reactive metals, gold is among the least reactive! We have already considered many consequences of this phenomenon in Chem2810, so we will not revisit this material now. For the p-block elements, the important thing to recognize is that they can exist in more than one oxidation state. In Chem2810 we discussed this important distinction using Frost diagrams. Such diagrams are reproduced below for the Group 13 and Group 14 elements. The lower oxidation states, e.g. Tl+ and Sn 2+, are a consequence of not oxidizing the ns2 electron pair. The difficulty of oxidizing these electrons increases down the periodic table, so that Tl3+ and Pb 4+ are both strong oxidizing agents. But of course for boron or aluminum, the 3+ oxidation state are not only very stable, but the only accessible state. There are many practical consequences of these changes in the chemistry relevant to the chemical properties of the compounds of these elements. Finally we note that the heavier p-block elements are only very mild reducing agents. 4.4 Handling alkali element and organometallic compounds Most cluster compounds of Groups 1 and 2, and of virtually all organometallic compounds that have been prepared, have very high reactivity. Specialized methods have had to be developed in order to handle them safely and without Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 63 compromising their integrity. Typically, reactive species are susceptible to water, and indeed to any protic source. Thus they are typically handled only in dried, aprotic solvents, and their manipulations are carried out under an atmosphere of inert nitrogen gas (or if available, argon, which is superior as an inert gas medium to nitrogen for several reasons, but is quite a bit more expensive.) Most organometallics are also quite sensitive to oxygen, with some like methyl lithium spontaneously burning in air. 4.4.1 The Glove Box There is a variety of isolation chambers that are marketed (or locally built) depending on their application. Thus, nuclear chemistry tends to use “hot boxes” where highly radioactive samples can be handled, often using remote control of mechanical arms and levers. Biohazards are also contained in special chambers. However, the concept that is usually understood by “glove box” is an inertatmosphere box with an automated catalytic purification system, and an interlock chamber for loading and emptying the box without contaminating the purified air. The goal of such glove boxes is to eliminate all water and oxygen from an atmosphere where otherwise fairly standard chemistry can be undertaken, the operator reaching in to manipulate the reagents, flasks, beakers etc. through full-length rubber gloves. A typical glove box is shown in the photo at right. Several modifications are now commercially available, such as double boxes allowing two workers to work on a procedure either facing each other or side-by-side in otherwise conventional boxes that have been joined together. Some boxes have integrated freezers, and others have microscopes for examining products and/or mounting crystals in a protected environment. 4.4.2 The Double Manifold Vacuum Line and Schlenk Apparatus Many labs involved in this area of research do the majority of their work in so-called Schlenk apparatus. Some pieces of Schlenk apparatus are shown below. They must be used in conjunction with a specialized gas distribution system known as a double manifold vacuum line. One possible design for such a line is shown in the figure at right. The taps can alternately remove all gases from the apparatus which is attached to the bottom connectors, either directly or, more commonly, via a vacuum hose. Or it can introduce a purified inert gas such as dinitrogen or argon into the evacuated vessel. Schlenk apparatus was specifically designed to allow for all the usual operations that chemists are used to performing in air within a controlled atmosphere. The arrangment shown in the figure at left shows how a vacuum filtration can be performed under nitrogen or argon. Once the mother liquor has been removed, the pieces are separated and the glass joints sealed with caps and stoppers. Then the filtrate and/or precipitate can be used further. For example, precipitates need to be thoroughly dried under vacuum. The filtrates are often evaporated under vacuum to recover the solute. Many modifications to the basic Schlenk concept have been proposed. The following illustration shows how side-arm flasks can be used to peform a filtration under inert gas. This procedure would be used to remove unwanted solids from the reaction mixture. Chemistry 3810 Lecture Notes 4.5 Dr. R. T. Boeré Page 64 Introduction to Main Group Organometallic Compounds There is a strong conceptual link between the element hydrides and organometallic derivatives. In the language introduced by Nobel winner Roald Hoffman, a singly-bonded hydrocarbon anion group is “isolobal” with a hydride anion: C CH3 H R "isolobal" symbol = same electron count and same orbital type (σ) The table at right places the methyl derivatives of the main group elements into the periodic system (compare to the similar table in section 3.1 for the hydrides.) They fall into the same classification schemes of ionic, electron poor, electron precise and electron rich. The series is most extensive for methyl; some elements do not form stable derivatives of primary longer chain alkyl groups because of instability towards β-elimination. For most methyl derivatives, the phenyl derivative is known, however, and these are often quite stable and considerably less reactive. There is an extensive modern chemistry with exotic hydrocarbon derivatives that we will see examples of later in this section. β-elimination is the reaction: MCH 2CH 2 R → MH + CH 2 CHR In which the H atom transferred to the metal comes from the second, or β-carbon. 4.5.1 Examples of E–CH3 Compounds Consider these representative examples: Compound Category Bonding Structure KMe CsMe ionic ionic ionic ionic K+ CH3 – ionic lattice Cs + CH3 – ionic lattice [Be(CH3 )2 ]x [Mg(CH3 )2 ]x* poylmeric poylmeric (3c,2e) (3c,2e) doubly-bridged polymeric chain doubly-bridged polymeric chain Li4 Me4 Zn(CH3 )3 HgMe2 Al2 (CH3 )2 electron poor electron poor electron precise electron poor (4c,2e) (2c,2e) 2c,2e normal bond (3c,2e and 2c,2e) (tetrahedron structure) linear linear SiMe4 GeMe4 electron precise electron precise 2c,2e normal bond 2c,2e normal bond tetrahedral tetrahedral As(CH3 )3 electron rich 2c,2e bond pyramidal with lone pair *Note: the more common magnesium methyl compound is CH3 Mg +Br– , i.e. a Grignard reagent! Dimethyl magnesium is made from the Grigard by a redistribution reaction, see below. The thermal stability of element methyl compounds also mirrors that of the hydrides. The graph below presents the M – C bond enthalpies (for methyl derivatives , in kJ/mol) at 298 K. (Data from M.E. O’Neill and K. Wade, Comprehensive organometallic chemistry, ed. G. Wilkinson, F.G.A. Stone and E.W.Abel, Vol. 1, Pergamon Press, Oxford, 1982.) The table presents the standard enthalpies of formation. Note that some are exothermic and some endothermic. Since the C–H bonds in the methyl portion are strongly exothermic w.r.t. their elements, even mildly exothermic compounds tend to be easily cleaved at the M–C bonds. Especially the heaviest members of each group have easily cleaved bonds, and for most of the last century tetraethyl lead (cheaper to manufacture than tetramethyl lead, but very similar in properties) was added to inferior gasoline grades to improve the smoothness of combustion in gasoline engines (a so called “anti-knock” agent.) This effect is based on the thermal cracking of PbEt 4 into lead and ethyl radicals which act as initiators for the radical combustion of the Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 65 gasoline hydrocarbons. This agent is now banned in North America only because of feared health hazards associated with the production of lead oxides released from the exhaust gases of the engines, especially from vehicles. 4.5.2 Synthesis of main group organometallics There are many synthetic routes for the synthesis of main group organometallics. Some are only used in a laboratory setting because they are too costly for large-scale industrial use. On the other hand, many of the industrial methods are difficult to implement in the laboratory because of the need for highly specialized equipment and working conditions. The following reactions will serve as model equations: • Direct preparation from a halocarbon (usually chloride or bromide) and the metallic element Preparation of organolithium compounds: 2 Li + RX → LiR + LiX Preparation of Grignard reagents: Mg + RX ether → RMgX Rochow method for the production of methyl silyl chlorides Cu Si + CH3Cl → (CH3 ) 2CCl 2 + (CH3) 3SiCl + CH3SiCl3 Fluidizedbedreactor This last reaction, uses a copper catalyst, and is an essential industrial process; without it the production of silicones, an important class of synthetic hybrid inorganic/organic polyemr not economical. • Transmetallation 2 Ga + 3 Hg(CH 3 ) 2 → Ga(CH3 )3 + 3 Hg This is the displacement of one metal from a hydrocarbon group by another. The displacing metal must be less electronegative than the displaced metal. Hence Hg, with the very high value of χ = 2.0 is often used. This particular reaction is done at about 60°C in a sealed, shaken, heavy wall glass reactor, sometimes called a Carius tube). • Metathesis (which can, however, best be thought of as nucleophilic substitution.) Li4 Me4 + SiCl4 Al 2Me 6 + 2 BF3 • → → 4LiCl + SiMe 4 2 AlF3 + 2 BMe 3 Redistribution reaction type (this is another form of metathesis, and here the label is highly appropriate SiCl 4 + SiMe 4 → SiClMe 3 + SiCl 2Me 2 + SiCl 3Me etc This reaction has considerable importance for the silicone industry, which is economically very important. • Addition to a multiple bond E-H + CH 2 =CH 2 → E-CH 2 -CH3 Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 66 Here E can be, e.g. BH2 , SiH3 , etc. This is used extensively in organic chemistry in hydroboration for antiMarkovnikov addition to alkenes. [Organikers oxidize off the BH2 with H2 O2 to make alcohols, and with peroxybenzoic acid to make carboxylic acids.] A variation on this reaction is used industrially to produce ethyl compounds by direct reaction of the metal and ethene in the presence of hydrogen gas. This is a very important process for cheap large scale production of main group organometallic compounds. 4.5.3. Properties of organometallics • • • • 4.6 They tend to oxidize easily. Many are flammable in air! Consequently they are reducing agents. They have nucleophilic character. The more electropositive the metal, the more carbanionic the organic group is. Thus alkyl lithium reagents and Grignard reagents are usually reacting as nucleophiles. In the Chemistry 3810 laboratory a methyl grignard (CH3 MgBr) is used with S=PCl3 in a reaction where it acts both nucleophile and reducing agent! Electron deficient organometallics are powerful Lewis acids; they form complexes with Lewis bases, and for this reason, basic solvents such as ethers are extensively used to stabilize reactive organometallic reagents: case in point, Grignard reagents are always made in ethers. An exception is n-butyl lithium, which is prepared and reacted in hexane solution. Amines, pyridine and Me2 S are all used as stabilizing agents for reactive organometallics. The extremely varied reactions of methyl lithium are depicted in the radial diagram at right. Structure and Bonding in Main Group Organometallic Compounds The structure of, and bonding in, main group organometallics can best be understood using a delocalized molecular orbital approach. We will consider several representative examples; it is impossible to cover all the many varied examples that have been identified by chemists to date. 4.6.1 Structure and bonding of the tetrameric cluster methyl lithium Methyl lithium, readily available from chemical suppliers as a solution in ether, has been shown to exist in the solid state using X-ray diffraction and in less-coordinating solvents as the tetrahedral Me4 Li4 cage shown below. There is a central Li4 tetrahedron, and an outer C4 tetrahedron. These two clusers interpenetrate one another, but importantly they both have the same set of point group symmetries. Another way to think of this is as that of each methyl group pointing out from the center of one face of the Li4 tetrahedral cluster. There have been several attempts to describe the bonding in this highly unusual carbon compound. First, we recognize the classification of methyl lithium as an extremely electron deficient compound. Just as diborane forms from the electron deficient BH3 , so H3 C–Li is not stable. Rather than dimerized, it aggregates into a tetramer in order to obtain sufficient bonding character. We note first that the carbon-lithium bond in H3 C–Li is expected to be highly polar. The MO’s of this fragment are very similar to those of the heteronuclear diatomics we treated earlier. At the AM1 level of theory, we get for this fragment: Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 67 1a1 1e 2a1 The occupied orbitals shown at the right are 1a 1 , 1e and 2a 1 , all of which are quite polarized towards the C with the exception of 2a 1 . The empty orbitals are localized strongly on the lithium – thus we have a very polar carbon lithium bond, as we might expect from the electronegativity differences. What happens in the tetramer? There have been numerous discussion in the literature about the bonding in this very important cluster compound. A definitive answer to the problem is provided, however, by Bickelhaupt FM, Hommes NJRV, Gu erra CF, Baerends EJ Organometallics 15 2923-2931 (1996). Their calculations, using density functional theory lead to the energy-level diagram and fragment orbital topologies shown in the next figure. The calculations were done using a neutral metal cluster and a tetrahedral array of ligand sp hybrid basis functions. These results from a highlevel calculation can be mimicked by using two fragments in HyperChem at the AM1 level of theory. Be sure to set each as having a spin multiplicity of 5, or the energy levels will not look right. The picture developed at right is essentially displayed in the four highest lying MO’s in AM1. The remaning electrons are involved in fairly standard C–H bonding orbitals. As a result of the extensive sp mixing that occurs for the element lithium, the energy of the 1a1 SOMO fragment orbital is remarkably similar to that of the 2a 1 SOMO of the lithium cluster. These therefore undergo a remarkably covalent, low polarity 8c,2e bond. The remaining interactions between the t2 sets of orbitals are distinctly polarized, and this accounts for the high negative charge on the C atoms and high positive charge on the lithium atoms of the tetramer. Using AM1, set up as described above, the in-phase overlaps of these eight fragment orbitals correspond the four highest-lying filled orbitals. The lower orbitals are more extensively involved in C–H σ bonding. Note that these MO’s correspond to linear combinations of the 2a 1 orbitals of the monomer. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 68 This description of the bonding in methyl lithium explains the driving force for tetramer formation to be the creation of the lo wer a1 MeLi monomer bonding MO. The bonds are not activation barrier for reaction of MeLi strong, and methyl lithium remains a highly reactive compound. But in Me4Li 4 the absence of other reagents, they E effectively make do as best as n possible by forming this cluster e structure. This idea can be graphed r g using a typical reaction profile for a y generic reaction of methyl lithium as shown at right. Thus the products of reactions monomer corresponds to a with electrophiles maximum, and Me4 Li4 to a local minimum. If a suitable electrophile is added (see diagram above of the choices available) then the reaction will exceed very exothermically to the final products, which tend to be electron precise, or at the very least less electron deficient than methyl lithium is itself. 4.6.2 Structure and bonding in alkyl and aryl aluminums The next example that we will treat in detail is the hydrocarbon derivatives of aluminum. So long as the hydrocarbon groups are not too bulky, these tend to exist in a diborane-like dimerized structure. This is true, for example, in Al2 Me6 , Al2 Ph 2 R4 and in the mixed compounds Al2 Cl2 Me4 and Al2 (OR)2 Me4 . The structures of examples of such compounds are shown below: Al2Me6 Al2Cl2Me4 Al2Ph2R4 Note that just as in diborane, the angles at the bridging groups is unusually small (e.g. 75°at the methyl groups and similar values at the bridging aryl. However the chlorine bridged dimer has a more “normal” bond angle of 91°. This again suggests substantial Al–Al bonding in these dimerized clusters. We note that the preference for occupation of the bridging position is given by the series: (Cl, OR > aryl > alkyl). The dimers are weak and are readily broken if stronger Lewis bases are present. At the PM3 level of theory, the orbitals of Al2 Me6 are highly reminiscent of those found in diborane. The energy level diagram at the right corresponds to the highest 6 filled levels in the methyl bridged dimer. The orbital at –11.81 eV is the sixth down from the top, but is the major contributor to fragment bonding in the dimer. This MO uses a sp σ type orbital to form the bonds in the bridge, where H had to use a 1s orbital. However, otherwise this MO strongly resembles the 2a1g MO of B2 H6 . Also note that this molecule has effective D2h symmetry, so long as the methyl groups are freely rotating in the molecule. The text tries to approach bonding in closely related Al2 Ph 2 R4 using a VB 3c,2e bond model. The delocalized MO method is obviously superior here. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 69 Of course, there are now many additional orbitals present, due to the need to have bonding between the C and H atoms of the methyl groups. This considerably complicates the MO diagram, and necessitates a very thorough consideration of the bonding theory and a willingness to reject those orbitals that make only a small contribution to the inter-dimer bonds. For this reason theoreticians will often replace hydrocarbon groups in calculations for hydrogen atoms (isolobal analogy). However, it is fairly easy to divide those orbitals that deal with the cluster from those that deal with terminal atom bonding. The inter-dimer bond is weak and can be broken by typical Lewis bases (ethers and amines) Another indication for this weak dimerization is what happens with large R-groups. For example, the structure of the tris(2,6-dimethylphenyl) analogue is shown at right. The large 2,6-substituted aromatic rings keep the molecule from dimerizing, and hence it forms a discrete AlR3 fragment, and is planar at the central Al atom. 4.6.3 Structure and bonding in alkyldislanes and polysilanes Consider the molecule Me3 SiSiMe3 which is expected to have an element-element single bond. Using the PM3 method, we calculate the indicated energy levels for the frontier orbitals of this compound. The HOMO and LUMO are also shown. Noteworthy here is the small size of the HOMO-LUMO gap, less than 8 eV. This is much smaller that the value calculated by the same method for the all-carbon analog 2,2,3,3-tetramethylbutane at 14.1 eV). A consequence of this as well as the fact that there is only a single bond between these atoms is that the Si–Si bond is readily cleaved by UV light. It is possible to prepare catenated silicon organometallics, e.g.: Si Si Si Si Si Si Si Si Si Si Si These are made by the reductive elimination of halides e.g. from R2 SiCl2 . A suitable reagent is for example Li[C10 H8 ]. How do these polymers compare with polyethylene? First, it is possible to make them with R side groups rather than just H, because they are not as sterically crowded as the smaller organic polymer. What is much more fundamentally interesting is that they show chemical and spectroscopic behaviour more like polyalkynes! For example: • low-lying empty orbitals (UV absorptions, which for C-C bonds do not occur at all, are found to be similar to those for C=C bonds, with the band maxima at ca. 200 nm • evidence of electron delocalization • electrochemically easier to reduce, etc.) The origin of this effect is related to the basic postulate of MO theory: σ stabilization and σ* destabilization is a function of the extent of overlap. Si-Si overlap is poorer than C–C, thus there is a smaller HOMO-LUMO gap for the Si-Si bond, as Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 70 shown in the comparative energy-level diagram at the right. Moreover, for longer polysilane chains, an effect is seen on the orbitals because of cross-overlap resulting in delocalized MO's over the whole chain, as shown in the schematic figure at the right. E N E R G Y σ* σ* 2sp3 2sp3 3 3sp3 3sp σ σ C Si The net result is that polysilanes approach semi-conductor properties. This is exactly analogous to the difference in the properties of diamond and elemental silicon, which has the diamond structure. Some data: Band gap, kJ mol-1 Resistivity, ; C 580 106 Si 105 6 104 Ge 58 50 Sn 7 1 This property is exploited in the use of poly(alkylsilanes) as photoresists in the making of circuit paterns on solid silica. A representation of a straight-chain poly(dimethylsilane) is shown in the line diagram at right. The rubbery material is applied, and then a mask is placed over it. Intense UV light then cleaves the bonds between the silicon atoms, so that these are easily rinsed off with solvent, and the remaining space can be sputtered with elemental metal (Al or Cu!) After this the residual photoresist is rinsed off as well, and additional silicon is laid down from SiH4 . This tendency towards weaker E–E bonds O O O O O increases down the periodic table. The most extensive O Si Si Si series of heavier analogues to the hydrocarbons is in Si Si this very family of poly(dimethylsilanes). In order to obtain thermally stable and non-light-sensitive silicon polymers, it is necessary to replace half the silicon units with a suitable heteroatoms . The best known examples are the poly(dimethylsiloxanes), commonly known by their trade name of silicones. Such heterochain polymers are common for the 3rd and heavier period elements. They are typically extremely temperature resistant, and remain flexible over a wide range of temperature. 4.7 Stereochemistry at silicon The redistribution reaction of SiCl4 and SiR4 was mentioned above. This is a quite general process, and mechanistic studies show these to be associative in nature: rate = k [ SiR3 X ][ X ] ] This is what is called an SN 2 reaction for carbon (substitution with inversion): This reaction is drawn to show inversion Y of configuration, which is normally observed for the carbon analog. With silicon, either retention or inversion is seen. Retention occurs when Y– is a poor leaving group, e.g. H– or OH– . This is rationalized by a rearrangement of the intermediate. When the Y elimination of Y– is a slow process, this has time to occur, otherwise it does not occur and inversion occurs. The alternate mechanism is drawn at right. The increased importance for this over C is explained by the greater stability of fivecoordinate Si over C. This is primarily an effect of larger size. + Si B - A A X - Y C Si A X Y - + C B B C B X A A Si Si C + - X Y Si X B A Y Si B C A X Y C Si B X + C- Chemistry 3810 Lecture Notes 4.8 Dr. R. T. Boeré Page 71 p(π )-p(π ) Multiple Bonds Between Elements from Period 3 and Beyond Text books of inorganic chemistry not so long ago contained statements such as the following: "Si, Ge, Sn and Pb do not form p multiple bonds under any circumstances. Thus numerous types of carbon compounds, such as alkenes, alkynes, ketones, nitriles, etc. have no analogs." Cotton and Wilkinson, "Advanced Inorganic Chemistry, 1972 Statements such as these have now been completely disproved! The major breakthroughs came in the early years of the 1980's. Continuous new discoveries are being made in this whole area of the chemistry of multiple bonds to "heavy" (3rd period and beyond) main-group elements. Most of these advances are based on organometallic derivatives of these elements - this is where the big progress has been made. We will cover aspects of multiple bonds for the Group 14 and 15 elements at this time. This material belongs just as much to the chapters on those elements, but the textbook has chosen to place them in the organometallic chapter. So we will deal with them now. This is just one example on how difficult it is to categorize chemical compounds! 4.8.1. Historical survey There had been many unsuccessful attempts to prepare analogues of 2nd period double-bonded compounds, many being made in the early years of the previous century. These claims are demonstrated in the names of these supposed compounds, e.g.: H R Ph H N O Known compounds: H alkenes H Si Supposed new analogues: H R H R ketones Si N azobenzene Ph R Si P O P H R R H disilenes silicones diphosphenes But by the mid-thirties, all these claims had been refuted. Instead, those compounds that had been prepared were all shown to be cyclic oligomers or high polymers, in which there were only single bonds. R H H R Si disilenes H R H R R Si R Si R Si Si R Si Si R silicones P diphosphenes O R P R Si R R Si O R R P R R n n P n n Si R P R P O R R P R R nn Si O R R Formation of small rings (above), and high polymers (below) R R R Si O R Si R Si nn Si R O * R n n Si R R R * Presumed polymers not well characterized Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 72 4.8.2. The double bond rule Persistent failure over many decades led chemists of the 40's and 50's to formulate the "classical double bond rule", which stated that: Elements having a principal quantum number greater than 2 are not able to form (p-π) π bonds with themselves or with other elements. This truism was rationalized by various arguments, such as: • Overlap It was argued that there was very poor sideways overlap between p-p orbitals due to the large size or the 3rd period and heavier elements • Diffuseness of orbitals It was suggested that the higher-quantum-number p orbitals were more diffuse than the 2p orbitals. Hence the electron density in an overlapping orbital would be lower, leading to weak π bonds. • • Energy of orbitals It was suggested that 3p, 4p and 5p orbital had too high energies to participate readily in bonding. Thermochemistry It was suggested that p-bond energies in heavy double bonds would be too small to prevent polymerization. This was largely an argument from silence, but went something like this: energy(kJ mol–1 ) 347 619 bond C-C C=C 2n R H H Si R H H H R n Si R R R n* H H R R 619 R n n R R 2 ∆ H = large -ve R * 2n R Si Si H R Si * Si energy(kJ mol–1 ) 226 ??? bond Si-Si Si=Si R ∆ H = -75 kJmol -1 * R × 347 Notice that even C=C double bonds are thermodynamically less stable than two corresponding single bonds, although the difference is small. In fact, C=C double bonds are kinetically stabilized, and the preparation of polymers, e.g. polyethylene (plastic bags, etc) involves using a catalyst to overcome the kinetic barrier and induce an exothermic reaction. Nevertheless, if the Si=Si bond were much weaker, the driving force for polymerization could be so large that it would be hard to prevent, or so the reasoning went... 4.8.3 Transient species indicating double bonds might exist During the 1960’s and 1970’s evidence for the transient existence of species with heavy mutiple bonds was accumulated. The general philosophy was to produce the species by a suitable reaction in solution or in the gas phase, and then to look either at trapped adducts, or at the decomposition, and argue for the existence of the unsaturated compound. Thus Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 73 tetramethyldisilene was generated by a retro-Diels Alder reaction, and then trapped again by a different diene with a higher stability constant. Such results left open the question, however, as to whether a genuine π-bond had been formed, or whether perhaps a transient diradical species had been formed. In the case of phosphorus both double and triple bonded species were postulated, but again the evidence was indirect, since only the decomposition products were actually isolated. All of this work served to whet the appetite for the actual isolation of multiply bonded species, but none of it was definite proof. 4.8.4 Silicon – the breakthrough The key synthetic principle that was applied in the first successful isolation of Si=Si double-bonded compounds was to provide sufficient steric bulk that the spontaneous polymerization was inhibited, but not so bulky that the reactivity of the new functional group was completely hindered. The early reports were for almost identical compounds produced by two US labs (Wisconsin and MIT). Robert West’s group at Wisconsin prepared starting materials with Si–Si single bonds, and then used a photochemical means to cleave the single bond, as we discussed above. The resulting silenes apparently dimerized to provide the beautiful yellow disilene. His choice of very bulky group was the mesityl group (2,4,6trimethylphenyl) – we saw this group previously as stabilizing planar AlR3 species. The competing group at MIT used an active metal reduction of Si–Cl bonds. This resulted in trimers (i.e. trisilacylopropanes). However, these trimers convered to disilenes under the action of either heat or UV-light. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 74 Are the properties of these compounds consistent with a genuine π-type double bond? The evidence in favour was obtained first of all from a structure determined single-crystal X-ray diffraction study. The evidence was: • The co-planarity of the central C2 SiSiC2 core is clearly visible. • The substituents are eclipsed, despite their very large steric bulk. • Careful measurement of the Si–Si bond distances gave 2.160 Å for the Wisconsin compound, versus a typical Si–Si bond of 2.32 Å. This represents an 8% shortening. For comparison, tetraphenylethene, a suitable C=C model system shows 12% shortening from the hydrogenated analogue. Thus the Si=Si is significantly shortened for a nonpolar bond. • The final evidence was obtained from a measurement of the Z → E isomerization energy, using a closely-related unsymmetrical disilene, (Mes)(t Bu)Si=Si(t Bu)(Mes), for which Eactivation = 131 kJ mol–1 . For comparative purposes, the barrier to inversion of Z to E isomerism of trans-stilbene is only 179 kJ mol–1 . Thus the barrier in the disilene is significantly higher than expected for such Si–Si bonds. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 75 Note that the absence of the π-component to a double bond will allow for free-rotation about the E–E bond. This was elegantly proven several year’s ago by the following crystal structures. The first one is of the dianion of Ph 2 C=CPh 2 , in which occupation of the π* orbital reduced the C–C bond order to 1. It is found to be staggered in the solid state. Thus the dianion [Ph 2 C–CPh 2 ]2– is found in a crystal structure of one of its salts to be almost completely twisted to the “eclipsed” form, as might be expected based on steric hindrance between the large phenyl groups (note that the hydrogen’s are not shown in this diagram because they are thought to obscure the salient details of the structure. The second example is that of a dication of a different alkene, {Me2 N}2 C=C{Me2 N}2 . Here the electrons have been removed from the π orbital, again allowing for essentially free rotation about the now single C–C bond in [{Me2 N}2 C=C{Me2 N}2 ]2+. Other properties of the disilenes that are consistent with the claim that they have a genuine σ + π double bond are as follows: • Their colour is typically yellow or orange, consistent with a λmax ~ 400 nm. • Their electrochemical reduction and oxidation are more facile than similarly substituted alkenes. Thus Ered = –2.12 V, while Eox = +0.38 V. Both of these values are considerably smaller than for the alkenes. • Detailed calculations suggest that both the HOMO and the LUMO of this compound should be considerably more reactive than that of alkenes. This suggests that the compounds will be reactive to both electrophiles and nucleophiles. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 76 A direct output from the Wisconsin group’s structure is shown in the graphic at right. The circles represent the volume of space within which most of the electron density around the nucleus of these atoms vibrates in the crystal lattice. Thus, although Si is the largest atom in the structure, it has some of the smallest atom circles, because it is a heavy element and these often vibrate the least. The distinctive yellow colour of these disilenes imply that they absorb visible light in the blue region of the spectrum. Molecular orbital studies of this type of molecule has shown the following results: The HOMO-LUMO gap in a disilene is indeed much smaller than that in a normal alkene. The band gap is half as high as in the C=C bond case! But with this kinetic stabilization the product lasted for some time. The following orbitals are generated from a PM3 calculation. They are the HOMO and the LUMO of a model disilene. ←HOMO - LUMO→ 4.8.5 Stable disilenes are still reactive molecules Alkenes are among the most important hydrocarbons, because the double bond is classified as a functional group in organic chemistry, i.e. a region of heightened reactivity compared to the very low reactivity of saturated hydrocarbon molecules. This reactivity is due to the high-lying HOMO and low-lying empty LUMO, resulting in both addition and donation reactions occurring from these two orbitals. An alkene becomes less reactive with increased subsitution of the sp 2 carbon atoms – eventually the double bond gets too greatly protected – shielded – so that it cannot donate or receive electrons (especially if these come attached to an atom!) What then for diselenes? Have we “stericaly protected” them to the point where they cannot react at all? Fortunately, this is not necessarily the case. Indeed, dimerization and polymerization is a reaction with oneself. If instead these disilenes react cleanly with reagents more reactive than itself, then a proper reactivity may be expected for this compound. Indeed, the latter is the case! The following diagram includes a reaction pattern for West’s original tetra(mesityl)disilene. As can be seen at a gla nce, most of the common reactions of alkenes, such as addition of halogen or HX apply to the disilene. In addition, there are some reactions that do not occur with alkenes, but do with the more reactive disilenes. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 77 CH3 Mes Cl H3C CH3 Si Mes Mes H Mes Si Si Cl Mes Mes Mes Si Cl Mes R = Mes (for mesityl) Mes Mes Si Si O O HCl Cl2 Mes Mes O O Ph Mes Ph Mes Si Si Mes Ph Ph Mes Mes Si O Mes Mes Mes CMe2 Mes Si Si Ph ROH Mes O Si Si Ph Mes Mes Mes O2 Me2CO Ph Mes Mes R = H, alkyl H Mes Mes Si O Si Mes Ph Mes OR Si Mes 4.8.6 Extensions of kinetically stabilized double bonds to germanium and tin Since the first reports by West and Masumune, similar ideas have been used to try and make even heavier elements with stable double bonds. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 78 The following data has been obtained from solid-state crystallographic data. The terms are defined in the sketch above. Note that all these compounds have considerable bond shortening. MO calculations have been peformed to try and understand this unusual bond pyramidalizaton. There have been two main attempts to try and rationalize these large deviations from co-planarity. The first is a delocalized MO approach (ie. What we are familiar with). The second is that of a VB model put forward by some British scientists as kl a possible explanation both for the pyramidalization of, especially, the tin compounds. This model considers two “carbenelike” fragments, each undergoing internal Lewis Acid-Base interactions between the filled “sp 2 ” and the empty p orbitals. This model is most useful in that it explains the existence of free tin-based R2 Sn “carbene-like” fragments. The MO calculations show a much more systematic approach to these systems. Note how the data is presented: the energies of bending about the θ-angle are ploted for all the Group 14 elements. Silicon is a lot more flexible than the carbon analogue. This trend continues until for both Ge and Sn, the flat form is no longer the global minimum, and the structure will be bent. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 79 The origin of this greater “floppiness” of the heavier Group 14 element compounds can be found from molecular orbital calculations. Consider the following energy level diagram, which compares planar and pyramidalized forms of both compounds. At root, it is the much higher-lying nature of the weak π-bond in the Ge =Ge compound that accounts for the much greater ease of second-order mixing upon symmetry. Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 80 Chemistry 3810 Lecture Notes Dr. R. T. Boeré 4.8.7 Extension of the principle to the Group 15 elements Stable Group 15 p(π)–p(π) bonds have been known since the early 1980’s. Page 81 Chemistry 3810 Lecture Notes Dr. R. T. Boeré Some crystal structures of diphosphenes and diarsenes Page 82 Chemistry 3810 Lecture Notes The Electronic Structure of R–P=P–R Dr. R. T. Boeré Page 83 Chemistry 3810 Lecture Notes Dr. R. T. Boeré Page 84 4.8.8 Reactivity of diphosphenes Again the reactivity of these compounds has not been blocked – just their propensity to oligomerize and polymerize. The following chart allows us to see some typical reactions, almost all of which are in some ways unique! 4.8.9 Conclusions 1. 2. 3. 4. 5. Thermodynaically stable p(π) - p(π) bonds exist both between the heavy elements and between heavy and light elements. Such bonds must be kinetically stabilized against polymerization with bulky ligands. Despite the bulky groups, the bonds are still reactive, and hence interesting. There are now many derivatives of such systems due to a wide variety of reactions at the multiple-bond centers. The homologous compounds of elements down a group show a continuum in behaviour and properties. Molecular orbital theory is well suited to showing such continuous relationships and why observed changes in properties within groups occur.