8) Organometallic Chemistry-1 (1).pdf

ORGANOMETALLIC CHEMISTRY Organometallic chemistry ⏤ the chemistry of compounds that contain metal–carbon bonds 476 Chapter containing 13 | Organometallic Chemistry In practice, complexes several other ligands similar to CO in their bonding, such as NO and N2, are often included. FIGURE 13.1 Examples...

ORGANOMETALLIC CHEMISTRY Organometallic chemistry ⏤ the chemistry of compounds that contain metal–carbon bonds 476 Chapter containing 13 | Organometallic Chemistry In practice, complexes several other ligands similar to CO in their bonding, such as NO and N2, are often included. FIGURE 13.1 Examples of Sandwich Compounds. O O O C C C Mn Fe Examples: B Rh Sandwich compounds ⏤ metal atoms with cyclic organic ligands containing delocalized π systems H3C B B Rh CH3 B Fe B S B Mn C C C O O O Cr FIGURE 13.2 Examples of Cluster Compounds. S O 1CO23 Fe C C C O O O ORGANOMETALLIC CHEMISTRY Cluster compounds FIGURE 13.2 Examples of Cluster Compounds. • may contain only two or three metal atoms or many dozens O C • may contain single, double, triple, quadruple, or quintuple bonds O OC C O C Cr • in some cases have ligands that bridge two or more metal atoms. Carbide clusters Co Co P 1OC23 CO C O CO 1CO22 Ru Sn Sn Sn Sn Sn Sn • contains C bonded to 5, 6, or more surrounding metals Fe 1OC22Ru 1OC23Ru C Fe1CO23 C 1CO23Fe C O O C Sn Sn Sn C O O C P OC Co O C OC Co 1CO23 Fe Fe 1CO23 Ru 1CO23 Ru 1CO23 Ru 1CO23 Many other types of organometallic compounds have interesting structures and ch cal properties. Figure 13.3 shows additional examples of the variety of structures that o in this field. ORGANOMETALLIC CHEMISTRY Many other types of organometallic compounds have interesting structures chemical 13.1 Historicaland Background | 477 properties. FIGURE 13.3 More Examples of Organometallic Compounds. Ta O S O Li Hg F F F F F S S Li Li S S Li O S Mn C C C1CMe322 OC C O O Ta Hg Hg Hg Hg I Hg Zr Zr I PR3 PR3 Pt Pt PR3 PR3 F F F F until experiments performed by Birnbaum in 1868. The structure of the compound was not determined until more than 100 years later!2 Zeise’s salt was the first compound identified F Cl H H - nearly so. More structures recent X-ray studies of solid ferrocene have as Grignard reagents. These complexes often have complicated and diffraction contain magnesium–carbon sigma bonds. Their synthetic utility was recognized early; by 1905, crystalline phases, with an eclipsed conformation at 98 K and with conform more than 200 research papers had appeared on the topic. Grignard other rings slightly twistedreagents (D5) in and higher-temperature crystalline modification reagents containing metal–alkyl s bonds, such as organozinc and organolithium reagents,sandwich compound ferrocene led The discovery of the prototype have been of immense importance in the development of synthetic organic chemistry. other sandwich compounds, of other compounds containing metal ato F F Historical Background Organometallic chemistry developed slowly from the discovery of Zeise’s salt in 1827 PR3 C H ring in a similar fashion, and to a vast array of other compounds 5 5 until around 1950. Some organometallic compounds, such as Grignard reagents, found utility F organic synthesis, ThePt first organometallic compound (Zeise’s salt) organic ligands. Therefore, it ismetal– often stated that the discovery of ferroc in but there was little systematic study of compounds containing ** Fulvalene of modern organometallic chemistry. to be was synthesized in 1827 and fulvalene from cyclopentadienyl bromide, carbon bonds. In 1951, in an attempt to synthesize PR3 reported 5 This of F F Kealy and PausonRreacted H the Grignard H reagent R An introductory history organometallic chemistry would be in cyclo@C H MgBr with FeCl . reaction 5 5 3 contained an ethylene group. 1951 Ferrocene, Fe(C , was 5H5)2vitamin mentioning the oldest organometallic compound known, B12 coe H ) Fe, ferrocene: did not yield fulvalene but an orange solid having the formula (C 5 5 2 C C C C ORGANOMETALLIC CHEMISTRY H formed from FeCl cyclo3 and rally occurring cobalt complex ( Figure 13.6 ) contains a cobalt–carbon s cyclo@C5H5MgBr + FeCl h (C H ) Fe H H H 3 5 5 2 cofactor in ferrocene a number of enzymes that catalyze 1,2attempt shifts in biochemical sy C5H5MgBr in an to e compound was not H ompound identified H stable; it could be sublimed in air without decomposing synthesize fulvalene The product wasClsurprisingly ctrons of the organic and was resistant to catalytic C hydrogenation and Diels–Alder reactions. In 1956, X-ray Pt FIGURE 13.5 Conformations of the structure of the Cl C Ferrocene. s occupying corners Cl H Fe Fe H dicular to the plane. onoxide as a ligand FIGURE 13.4ofAnion of Zeise’s Anion Zeise’s salt Mond reported the Salt. D5d D5h l for the purification Staggered rings Eclipsed rings 1890 – Mond prepared Ni(CO)4 . Barbier 1898 in 1898 and – Grignard reagents (alkyl magnesium complexes) FIGURE 13.6 Vitamin B12 CH2OH mplexes now known H Coenzyme. O uctures and contain O H H H O P ized early; by 1905, O - HO Fe D5 Skew rings D5d D5h Staggered rings Eclipsed rings ORGANOMETALLIC CHEMISTRY FIGURE 13.6 Vitamin B12 Coenzyme. oldest organometallic What is the compound known? H3C CH2 NH CO Vitamin B12 coenzyme O- CH • naturally occurring cobalt complex CH2 • it is a cofactor in a number of enzymes that catalyze 1,2 shifts in biochemical systems N CH3 N CH3 CH2CH2CONH2 CH3 H CH3 N CH2 CH2 H2NOC HO CH3 H • contains a cobalt–carbon sigma bond H H H O O P O H3C CH2 CONH2 CH3 N Co N N H CH3 CH2 CONH2 CH2CH2CONH2 CH3 CH3 H CH2CONH2 CH2 CH2 *For Skew rings CH2OH H O D5 O H H OH H OH H N N NH2 N N an interesting article on the discovery of ferrocene’s structure, see P. Laszlo, R. Hoffmann, Angew. Chem., Int. Ed., 2000, 39, 123. **A special issue of the Journal of Organometallic Chemistry (2002, 637, 1) has been devoted to ferrocene, For example, because the cyclopentadienyl ligands in ferrocene bond through all five atoms, they are designated h5@C5H5. The formula of ferrocene may therefore be written (h5@C5H5)2Fe. The h5@C5H5 ligand is designated the pentahaptocyclopentadienyl ligand. ORGANOMETALLIC CHEMISTRY Hapto comes from the Greek word for fasten; therefore, pentahapto means “fastened in five places.” C5H5, probably the second most frequently encountered ligand in organomeOrganic Ligands and Nomenclature tallic chemistry (after CO), most commonly bonds to metals through five positions, but certain circumstances, it may bond through only one or three positions. As a ligand, The number ofunder atoms through which a ligand bonds is indicated by the Greek letter η (eta) C H is commonly abbreviated Cp. 5 5 followed by a superscript indicating the number of ligand atoms attached to the metal. * The corresponding formulas and names are designated as follows: Number of Bonding Positions Formula 1 h1@C5H5 Monohaptocyclopentadienyl M 3 h3@C5H5 Trihaptocyclopentadienyl M 5 h5@C5H5 Pentahaptocyclopentadienyl M Name • The η5–C5H5 ligand (or Cp) is designated the pentahaptocyclopentadienyl ligand. • Hapto (Greek) means “fastened in five places.” Ligandmeans Namefasten; therefore, pentahapto Ligand Name CO C Carbonyl Carbene (alkylidene) Benzene h5@C5H5 5 M Pentahaptocyclopentadienyl ORGANOMETALLIC CHEMISTRY Several of the ligands may bond through different numbers of atoms. Ligand CO C C Name Ligand Name Carbonyl Benzene Carbene (alkylidene) 1,5-cyclooctadiene (1,5-COD) (1,3-cyclooctadiene complexes are also known) Carbyne (alkylidyne) Cyclopropenyl (cyclo-C3H3) Cyclobutadiene (cyclo-C4H4) H2C CH2 Ethylene HC CH Acetylene p-Allyl (C3H5) CR3 Alkyl O Cyclopentadienyl (cyclo-C5H5) (Cp) Acyl C R FIGURE 13.7 Classic Organic Ligands. *For Classic Organic Ligands ligands having all carbons bonded to a metal, sometimes the superscript is omitted. Ferrocene may therefore be written (h@C5H5)2Fe and dibenzenechromium (h@C6H6)2Cr. Similarly, p with no superscript may occasionally Organometallic Chemistry ORGANOMETALLIC CHEMISTRY As in theare case of other coordination compounds, bridging are designated the Bridging ligands designated by the prefix µ, followed by aligands subscript indicatingbythe prefix m, followed bybridged. a subscript indicating the number of metal atoms bridged. Bridging number of metal atoms carbonyl ligands, for example, are designated as follows: Number of Atoms Bridged None (terminal) Formula CO 2 m2@CO 3 m3@CO 13.3 The 18-Electron Rule THE 18-ELECTRON RULE In organometallic chemistry,wethe electronic structures many compounds based on a In main group chemistry, have encountered the octetof rule, in which electronicare structures totalcan valence electron on count of 18of ona the central atom. of 8 electrons. Similarly, be rationalized the basis valence shellmetal requirement in organometallic chemistry, the electronic structures of many compounds are based on a total valence electron count of 18 on the central metal atom. As with the octet rule, there are many exceptions to the 18-electron rule, but the rule nevertheless provides useful guidelines to the chemistry of many organometallic complexes, especially those contain- ORGANOMETALLIC CHEMISTRY Counting Electrons Cr(CO)6 • A Cr atom has 6 electrons outside its noble gas core. • Each CO is considered to act as a donor of 2 electrons. • The total electron count is therefore: Cr 6(CO) 6 electrons 6 x 2 electrons = 12 electrons Total = 18 electrons • Cr(CO)6 is therefore considered an 18–electron complex. • It is thermally stable and can be sublimed without decomposition. • On the other hand, Cr(CO)5, a 16–electron species, and Cr(CO)7, a 20–electron species, are much less stable and are known only as transient species. + - • Likewise, the 17–electron [Cr(CO)6] and 19–electron [Cr(CO)6] are far less stable than the 18– electron Cr(CO)6. ORGANOMETALLIC CHEMISTRY (η5-C5H5)Fe(CO)2Cl Method A: Donor Pair Method • This method considers ligands to donate electron pairs to the metal. • To determine the total electron count, take into account the charge on each ligand and determine the formal oxidation state of the metal. e.g. η5–C5H5 is considered by this method as C5H5-, a donor of 3 electron pairs; it is a 6–electron donor. CO is counted as a 2–electron donor. Cl- is counted as a 2–electron donor. Therefore, (η5–C5H5)Fe(CO)2Cl is formally an iron(II) complex. Iron(II) has 6 electrons beyond its noble gas core. ORGANOMETALLIC CHEMISTRY Therefore, the electron count is Fe(II) 6 electrons η5-C5H5 6 electrons 2(CO) 4 electrons Cl - 2 electrons Total = 18 electrons Method B: Neutral-Ligand Method • This method uses the number of electrons that would be donated by ligands if they were neutral. • For simple inorganic ligands, ligands are considered to donate the number of electrons equal to their negative charge as free ions. ORGANOMETALLIC CHEMISTRY For example, Cl is a 1–electron donor (charge on free ion = –1) O is a 2–electron donor (charge on free ion = –2) N is a 3–electron donor (charge on free ion = –3) The oxidation state of the metal does not need to be specified to determine the total electron count by this method. Fe has 8 electrons beyond its noble gas core. (η5–C5H5)Fe(CO)2Cl Fe atom 8 electrons η5-C5H5 5 electrons 2(CO) 4 electrons Cl 1 electrons Total = 18 electrons ORGANOMETALLIC CHEMISTRY The charge must be included in determining the total electron count. • In the neutral–ligand method, for anions the charge of the complex is added as a number of electrons to the total, and for cations the magnitude of the charge of the complex is subtracted as a number of electrons from the total. [(η5–C5H5)Fe(CO)2]– [Mn(CO)6]+ Fe atom 8 electrons Mn atom η5–C5H5 5 electrons 6(CO) 12 electrons 2(CO) 4 electrons Charge –1 electrons Charge 1 electrons Total = 18 electrons 7 electrons Total = 18 electrons ORGANOMETALLIC CHEMISTRY A metal–metal single bond counts as one electron per metal, a double bond counts as two electrons per metal, and so forth. (CO)5Mn⏤Mn(CO)5 Mn atom 7 electrons 5(CO) 10 electrons Mn⏤Mn bond 1 electrons Total = 18 electrons In ligands such as CO that can interact with metal atoms in several ways, the number of electrons counted is usually based on σ donation. • e.g. although CO is a π acceptor and (weak) π donor, its electron-donating count of 2 is based on its σ donor ability alone. 6 6 6 (C7H7 + ) h7@C7H7 (Cycloheptatrienyl) 7 ORGANOMETALLIC CHEMISTRY E X A M P L E 13 . 2 rganometallic Chemistry Electron Counting Schemes for Common Ligands TABLE 13.1 Electron Counting Schemes for Common Ligands Ligand Method A H 2 (H - ) 1 Cl, Br, I 2 (X - ) 1 OH, OR 2 (OH - , OR - ) 1 CN 2 (CN - ) 1 - Both methods of electron counting are illustrated for the following complexes. Method B - CH3, CR3 2 (CH3 , CR3 ) 1 NO (bent M i N i O) 2 (NO - ) 1 NO (linear M i N i O) 2 (NO + ) 3 CO, PR3 2 NH3, H2O Method A ClMn(CO)5 Method B Mn(I) 6 e- Mn 7 e- Cl - 2 e- Cl 1 e- 10 e - 5 CO 18 e 6 e- 2 (h5@C5H5)2Fe Fe(II) 2 2 “ CRR! (Carbene) 2 2 (Ferrocene) 2 h5@C5H5 - H2C “ CH2 (Ethylene) 2 2 CNR 2 2 “ O, “ S 4 (O2 - , S2 - ) h3@C3H5 (p@allyl) ‚ CR (Carbyne) 12 e - [Re(CO)5(PF3)] + 6 e- 2 4 (C3H5 -) 5 CO 10 e - 3 3 3 PF3 ‚N 6 (N ) 3 Ethylenediamine (en) 4 (2 per nitrogen) 4 Bipyridine (bipy) 4 (2 per nitrogen) 4 Butadiene 4 4 h5@C5H5 (Cyclopentadienyl) 6 (C5H5 - ) h6@C6H6 (Benzene) h7@C7H7 (Cycloheptatrienyl) 18 e 8 e- Fe 2 h5@C5H5 18 e Re(I) 3- 10 e - 5 CO (contd.) 2 e- + charge * 18 e - Ti(2–) 6 e- 5 6 CO 12 e - 6 6 6 (C7H7 + ) 2– charge 7 * 18 e - 7 e- Re 10 e - 5 CO 2 e- PF3 + charge Method A 18 e [Ti(CO)6]2- 10 e - -1 e - Method B 18 e Ti (continues) 4 e- 6 CO 12 e - 2– charge 2 e18 e - * Charge on ion is accounted for in assignment of oxidation state to metal. E X A M P L E 13 . 2 Both methods of electron counting are illustrated for the following complexes. Method A Method B The electron-counting method of choice is a matter of individual preference. Metho A includes the formal oxidation state of the metal; Method B does not. Method B may b ORGANOMETALLIC CHEMISTRY Assignment: 1) Determine the valence electron counts for the transition metals in the following complexes: a) [Fe(CO)4]2– b) [(η5–C5H5)2Co]+ c) (η3–C5H5)(η5–C5H5)Fe(CO) d) Co2(CO)8 (has a single Co⏤Co bond) 2) Identify the first-row transition metal for the following 18-electron species: a) [M(CO)3(PPh3)]– b) HM(CO)5 c) (η4–C8H8)M(CO)5 d) [(η5–C5H5)M(CO)5]2 (assume single M⏤M bond) ORGANOMETALLIC CHEMISTRY Why 18 Electrons? If the octet represents a complete valence electron shell configuration (s2p6), then the number 18 represents a filled valence shell for a transition metal (s2p6d10). • But this analogy does not provide an explanation for why so many complexes violate the 18electron rule. • The valence–shell rationale does not distinguish between types of ligands (e.g., σ donors, π acceptors) • This distinction is important in determining which complexes obey and which violate the rule. A good example of a complex that adheres to the 18–electron rule is Cr(CO)6. • The MOs of interest in this molecule are those that result primarily from interactions between the d orbitals of Cr and the σ-donor (HOMO) and π-acceptor orbitals (LUMO) of the six CO ligands. ORGANOMETALLIC CHEMISTRY MO Levels of Cr(CO)6 Chapter 13 | Organometallic Chemistry RE 13.8 Molecular Orbital y Levels of Cr(CO)6. • Chromium(0) has 6 electrons outside its noble gas core. cular Orbital Energy Levels CO)6 by Gary O. Spessard ary L. Miessler. Reprinted rmission. • Each CO contributes a pair of electrons to give a total electron count of 18. •In the MO diagram, these 18 electrons appear as the 12 σ electrons—the σ electrons of the CO ligands, stabilized by their interaction with the metal orbitals—and the 6 t2g electrons. antibonding •Addition of one or more electrons to Cr(CO)6 would populate the eg orbitals, which are antibonding; the consequence would be destabilization of the molecule. •Removal of electrons from Cr(CO)6 would depopulate the t2g orbitals, which are bonding as a consequence of the strong π-acceptor ability of the CO ligands •A decrease in electron density in these orbitals would also tend to destabilize the complex. 12 σ electrons Chromium(0) has 6 electrons outside its noble gas core. Each CO contributes a pair of electrons to give a total electron count of 18. In the molecular orbital diagram, these 18 electrons appear as the 12 s electrons—the s electrons of the CO ligands, stabilized by their interaction with the metal orbitals—and the 6 t2g electrons. Addition of one or more electrons to Cr(CO)6 would populate the eg orbitals, which are antibonding; the consequence •The result is that the 18-electron configuration for this molecule is the most stable. ORGANOMETALLIC CHEMISTRY • Ligands that are both strong σ donors and π acceptors are the most effective at forcing adherence to the 18–electron rule. Examples of exceptions [Zn(en)3]2+ • is a 22–electron species • it has both the t2g and eg* orbitals filled. • although en (ethylenediamine) is a good σ donor, it is not as strong a donor as CO. • As a result, electrons in the eg orbitals are not sufficiently antibonding to cause significant destabilization of the complex, and the 22-electron species, with 4 electrons in eg orbitals, is stable. ORGANOMETALLIC CHEMISTRY [TiF6]2– • is a 12-electron species eg* • the fluoride ligand is a π donor as well as a σ donor. • The π-donor ability of F– destabilizes the t2g orbitals of the complex, making them slightly antibonding. ¢o t2g d eg* weakly antibonding; more than 18 electrons possible t2g nonbonding; fewer than 18 electrons possible s \ • has 12 electrons in the bonding σ orbitals and no electrons in the antibonding t2g or eg* orbitals Minimum of 12 electrons M ML6 1Oh2 6L Exceptions to the 18-Electron 2Rule. + Examples of exceptions may be noted. [Zn(en)3] is a 22-electron species; * both the t2g and eg orbitals filled. Although en (ethylenediamine) is a good s do is not as strong a donor as CO. As a result, electrons in the eg orbitals are not suffi antibonding to cause significant destabilization of the complex, and the 22-electron s with 4 electrons in eg orbitals, is stable. An example of a 12-electron species is TiF this case, the fluoride ligand is a p donor as well as a s donor. The p@donor ability ngly CO)5 486 Chapter 13 | Organometallic Chemistry ORGANOMETALLIC CHEMISTRY anar FIGURE 13.11 Molecular y forSquare-Planar Complexes Orbital Energy Levels for a Square-Planar Complex. Examples of square-planar complexes include the d8 16-electron complexes. n in ially diais a ailed ntermetal four arily m of N N C C Ph3P C Rh N p Cl Pt Cl PPh3 Ph3P PPh3 Cl Ir H2 N dz 2 N H2 C b1g antibonding a1g •Putting electron in b1g to make an 18-electron complex would destabilize it. dx 2-y 2 s 2- Ni C Cl N p* e dxz dyz g d b dxy 2g s O PPh3 σ-donor orbitals M Wilkinson’s complex Vaska’s complex FIGURE 13.10 Examples of 8 Square-Planar d Complexes. ML4 1D4h2 4L p@acceptor characteristics, a 16-electron configuration is more stable than an 18-electron Orbital Energy Levels for may a Square-Planar configuration. Molecular Sixteen-electron square-planar complexes be able to acceptComplex. one or two ligands at the vacant coordination sites (along the z axis) to achieve an 18-electron configuration. This is a common reaction of 16-electron square-planar complexes (Chapter 14). E X E R C I S E 13 . 3 ORGANOMETALLIC CHEMISTRY 13.4 Ligand Ligands in Organometallic Chemistry s* Carbonyl (CO) Complexes s* • CO is the most common ligand in organometallic chemistry. p* p* • CO serves as the only ligand in binary carbonyls such as Ni(CO)4, LUMO W(CO)6, and Fe2(CO)9 LUMO • CO may bond to a single metal, or it may serve assa bridge s between two or more metals. HOMO HOMO • The HOMO of CO has its largest lobe on carbon. It is through this orbital, occupied by an electron pair, that CO exerts its σ-donor p function p C O C O • CO also has two empty π orbitals (or LUMO); these also have N2 CO larger lobes on carbon than on oxygen; a metal atom can Frontier Orbitals of CO donate electron density to these πtheorbitals. bonding between metals and CO, the synthesis and reactions of CO complexes * ** examples of various types of CO complexes. metal and CO; however, the strength of this bonding depends on several factors, including the charge on the complex and the ligand environment of the metal. ORGANOMETALLIC CHEMISTRY E X E R C I S E 13 . 4 N2 has molecular orbitals rather similar to those of CO, as shown in Figure 13.12. Bonding with CO Would you expect N2PitoInteractions be a stronger or weaker pCO acceptor than CO?atom Sigma and between and a metal Sigma donation M C Pi acceptance O M FIGURE 13.13 Sigma and Pi Interactions between CO and a Metal Atom. Overall interaction p C O M s C O Lobe of acceptor orbital M C O • The overall effect is synergistic. • CO can donate electron density via a σ orbital to a metal atom; • the greater the electron density on the metal, the more effectively it can return electron density to the π* orbitals of CO. • The net effect can be strong bonding between the metal and CO. Experimental evidences of CO bonding are IR spectroscopy and X-ray Crystallography. • The C⏤O stretch in organometallic complexes is often very intense. -1 Free CO has a C⏤O stretch at 2143 cm • -1 Cr(CO) has a C⏤O stretch at 2000 cm 6 • • The lower energy for the stretching mode means that the C⏤O bond is weaker in Cr(CO)6 • Both σ donation (which donates electron density from a bonding orbital on CO) and π acceptance (which places electron density in C⏤O antibonding orbitals) would be expected to weaken the C⏤O bond • In carbon monoxide, the C⏤O distance has been measured at 112.8 pm. • Weakening of the C⏤O bond would be expected to cause this distance to increase. • Complexes containing CO have C⏤O distances approximately 115 pm for many carbonyls. • dence is provided by X-ray crystallography. In carbon monoxide, the s been measured at 112.8 pm. Weakening of the C i O bond would e this distance to increase. Such anORGANOMETALLIC increase in bond length isCHEMISTRY found in ng CO, with C i O distances approximately 115 pm for many carbonyls. The charge on a CO complex is also reflected in its IR spectrum. surements provide definitive measures of bond distances, in practice it is t to use infrared spectra to obtain data on the strength of C iC⏤O O bonds. 5 isoelectronic hexacarbonyls have the following stretching bands (compare with a carbonyl complex is-1also reflected in its infrared spectrum. Five 𝒗(CO) = 2143 cm for free CO):* carbonyls have the following C i O stretching bands (compare with • The formal charges on the metals increase from -2 for -1 * for free CO): v(CO), cm−1 Complex [Ti(CO)6] [V(CO)6] 2- - 1859 2000 Cr(CO)6 + 2100 2+ 2204 [Mn(CO)6] [Fe(CO)6] 2- 1748 [Ti(CO)6]2- to +2 for [Fe(CO)6]2+. • The titanium in [Ti(CO)6]2-, with the most negative formal charge, has the strongest tendency to donate to CO. • The consequence is strong population of the π* orbitals of CO in [Ti(CO)6]2- and reduction of the strength of the C⏤O bond. • In general, the more negative the charge on the organometallic species, the greater the tendency of the metal to donate electrons to the π* orbitals of CO, and the metal, formally containing lower the energy of the C⏤O stretching vibrations. CO)6] contains the most highly reduced that titanium has the weakest ability to attract electrons and the greatest onate electron density to CO. The formal charges on the metals increase 22+ 2)6] to +2 for [Fe(CO)6] . The titanium in [Ti(CO)6] , with the most shared by the carbon and the oxygen, giving rise to a stronger bond and a higher-energy C i O stretch. ORGANOMETALLIC CHEMISTRY Bridging Modes of CO Many cases are known in which CO forms bridges between two or more metals. Many Bridging Modes of CO bridging modes are known (Table 13.2). Many cases are known in which CO forms bridges between two or more metals. TABLE 13.2 Bridging Modes of CO Approximate Range for v(CO) in Neutral Complexes (cm−1) Type of CO Free CO 2143 Terminal M i CO 1850–2120 Symmetrica m2 i CO M Symmetrica m3 i CO M O C O C M M M 1700–1860 1600–1700 O C m4 i CO < 1700 (few examples) M M M M *Asymmetrically bridging m2- and m3@CO are also known. *For reviews of metal carbonyl anions and other complexes containing metals in negative oxidation states, see J. E. Ellis, Organometallics, 2003, 22, 3322 and Inorg. Chem., 2006, 45, 3167. • The bridging mode is strongly correlated with the position of the C⏤O stretching band. • In cases in which CO bridges two metal atoms, both metals can contribute electron density into π* orbitals of CO to weaken the C⏤O bond and lower the energy of the stretch. • Consequently, the C⏤O stretch for doubly bridging CO is at a much lower energy than for terminal COs. ORGANOMETALLIC CHEMISTRY Interaction of three metal atoms with a triply bridging CO further weakens the C⏤O bond; IR band for the C⏤O stretch is still lower than in the doubly bridging case. 490the Chapter 13 | Organometallic Chemistry O O C C M M s Donor M M p Acceptor O O O C C C OC CO Fe Fe -1 n = 2082, 2019 cm CO C C O C O O C O nCO = 1829 cm-1 FIGURE 13.14 Bridging CO. For comparison, the carbonyl stretches in organic molecules are typically in the range 1700 to 1850 cm-1. The bridging mode is strongly correlated with the position of the C i O stretching band. In cases in which CO bridges two metal atoms, both metals can contribute electron * density into p orbitals of CO to weaken the C i O bond and lower the energy of the stretch. Consequently, the C i O stretch for doubly bridging CO is at a much lower energy –1bridging The mode is strongly correlated with the position of the C i O stretching alkylThe ketones near 1700 cm .) bridging mode is strongly correlated with the position of the C i O stretching band.twoIn cases inboth which CO bridges two metal atoms, both metals2-electron can contribute electron band. In cases in which terminal CO bridges metal atoms, metals can contribute electron Ordinarily, and bridging carbonyl ligands can be considered O * * O bond and lower the energy of the density into p orbitals of CO to weaken the C i density into p orbitals of CO to weaken the C i O bond and lower the energy of the C CO donors, with thethe donated electrons shared atoms in the bridging cases. ORGANOMETALLIC stretch. Consequently, C i O stretch for doubly bridging COCHEMISTRY is at a by muchthe lower metal energy stretch. Consequently, the C i ofOthree stretch for doubly bridging CO is at a much lower energy than for terminal COs. An example is shown in Figure 13.14 . Interaction metal Foratoms example, in the complex to the left, the bridging CO is a 2-electron donor overall, with a triply bridging CO further the C COs. i O bond; infrared band Re Re than forweakens terminal Antheexample is for shown in Figure 13.14. Interaction of three metal with donated to each metal. The electron for each Re atom accordthe Caisingle O stretch electron is still lower than in the doubly bridging case. (For comparison, carbonyl count Terminal and bridging CO ligands can be considered 2-electron donors, with the donated withinathetriply bridging C C C stretches in organic moleculesatoms are typically range 1700 to 1850 CO cm –1, further with many weakens the C i O bond; the infrared band for ingalkyl to method B cm is–1.)theinC the O ketones nearatoms 1700 electrons shared by the metal bridging cases. i O stretch is still lower than in the doubly bridging case. (For comparison, carbonyl O O Ordinarily, terminal and bridging carbonyl ligands can be considered 2-electron –1 O stretches in organic molecules are typically in the range 1700 to 1850 cm , with many C CO Re donors, with the donated electrons shared by the metal atoms in the bridging cases. – 1 ketones nearCO1700 cm .) donor overall, For example, in the complex alkyl to the left, the bridging is a 2-electron with a single electron donated to each metal. The electron count for each Re atom accord5 Ordinarily, terminal and bridging carbonyl O ing to method B is 5 5 Re C C O C O O C CO Re Re C C C O O O Re 7e ligands can be considered 2-electron h @C H 5e donors, with the donated electrons shared by the metal atoms in the bridging cases. 2 (CO) in (terminal) 4 ebridging CO is a 2-electron donor overall, Re For example, 7 e the complex to the left, the h5@C5H5 5 e1 - electron count for each Re atom accordwith a2single electron donated to each metal. The (m @CO) 1 e 2 (CO) (terminal) 2 4 eing to method B is 1 e 1 (m @CO) 2 2 1e M i M bond 1e M i M bond Total = 18 e Total = 18 e Re 7e - 5 Nearly linear bridging carbonyls, such as those in [(h @C5H5)Mo(CO) particuh @C52]H 2 are 5 5 5 larly interesting. When [(h @C5H5)Mo(CO)3]2 is heated, some carbon monoxide is released; 10 2 this (CO) (terminal) the product, [(h5@C5H5)Mo(CO)2]2, reacts reaction: 5 readily with CO to reverse 5 5 3 12 5 ! 2 (m2@CO) 5 5 [(h @C5H5)Mo(CO) 5 35]2 m [(h @C 2 5H 2 5)Mo(CO)2]2 + 2 CO 5e Nearly linear bridging carbonyls, such as those in [(h @C5H5)Mo(CO)2-]2 are particu4e larly interesting. When [(h @C H )Mo(CO) ] is heated, some carbon monoxide is released; The Mo⏤Mo bond distance shortens by approximately 79 pm, • the product, [(h @C H )Mo(CO) ] , reacts readily with CO to reverse this1 ereaction:10 consistent with an increase in the bond1 order from 1 to 3. 1889, 1859 cm 1960, 1915 cm eMi M metal–metal bond Total = 18 e important interaction donation a metal d orbital to the π • An !is in This reaction is accompanied by changes in the infrared spectrum the CO region, as from listed 5 O C O C Mo Mo -1 Mo Mo C C Mo C Mo * O 5 5 [(h @C H )Mo(CO) ] m [(h H )Mo(CO) ] + 2 CO distance also shortens by approximately 79 @C pm, 5consistent above. The Mo i Mo bond 5 5 3 2 5 2 2 orbital of CO; itorder weakens the C⏤O bond; 1shifts bands 5 to lower with an increase in the metal–metal bond from 1 to 3. Calculations have indicated 1 1889, 1859 1960, 1915Nearly cm linear bridging carbonyls, suchcmas those in [(h @C5H5)Mo(CO)2]2 are particu* that an important interaction is donation from a metal d orbital to 5the p orbital of CO O energies larly interesting. When [(h @Cligand (Figure 13.15).11 Such donation weakens the carbon–oxygen bond in the and results 3]2 is heated, some carbon monoxide is released; 5H5)Mo(CO) Mo O C C O O C O O - C - C FIGURE 13.15 Bridging CO in [(h5-C5H5)Mo(CO)2]2. Mo -1 in the observed shift of the C i stretching bands lower energies. theO product, [(hto5@C 5H5)Mo(CO)2]2, reacts Infrared spectra of carbonyl complexes are discussed further in Section 13.8. 10 readily with CO to reverse this reaction: This reaction is accompanied by changes in the infrared spectrum in the CO region, as listed above. The Mo i Mo bond distance also shortens by! approximately 79 pm, consistent Mo Mo 5 5 [(h @C H )Mo(CO) ] m [(h @C5H5)Mo(CO) + 2 CO with an increase in the metal–metal bond order 3from 1 to 3. Calculations have 5 5 2 2]2 indicated is possible. However, V(CO)6 is easily reduced to [V(CO)6] - , an 18-electron complex. E X E R C I S E 13 . 6 ORGANOMETALLIC CHEMISTRY Verify the 18-electron rule for five of the binary carbonyls—other than V(CO)6, Co6(CO)16, and Rh6(CO)16 i shown in Figure 13.16. An interesting feature of the structures of binary carbonyl complexes is that the tendency of CO to bridge transition metals decreases going down the periodic table. For example, in Fe2(CO)9 there are three bridging carbonyls; but in Ru2(CO)9 and Os2(CO)9, there is a single bridging CO. A possible explanation is that the orbitals of bridging CO are less able to interact effectively with transition-metal atoms as the size of the metals increases, along with the metal–metal bond lengths. Binary Carbonyl Complexes Binary carbonyls, containing only metal atoms and CO, are numerous. Mononuclear 3M1CO2x4 O C OC M C O OC CO OC M = Ni, Pd Polynuclear (CO represented by M OC Co C O Fe C O C O Fe Fe21CO29 Fe Fe O C O C CO Co Co OC C O C Fe31CO212 C O CO C O Co21CO28 1solid2 O C O C CO CO C O OC O OC C M31CO212 M = Ru, Os M C O M C O M21CO210 M = Mn, Tc, Re FIGURE 13.16 Binary Carbonyl Complexes. M Ir M M41CO212 M = Co, Rh O C Ir Ir Ir41CO212 M M M M M • Most of these complexes obey the 18-electron rule. • The cluster compounds Co6(CO)16, Rh6(CO)16 and V(CO)6 do not obey the rule. M CO C O M Ir O M M M Fe C O M = V, Cr, Mo, W Co21CO28 1solution2 O C M CO CO M C O C OC O O C OC CO OC M = Fe, Ru, Os O O C C Co CO C O Binuclear 3M21CO2x4 O C O C O C for clarity) M61CO216 M = Co, Rh • The tendency of CO to bridge transition metals decreases going down the periodic table. • Orbitals of bridging CO are less able to interact effectively with transition-metal atoms as the size of the metals increases, along with the metal– metal bond lengths. ORGANOMETALLIC CHEMISTRY Oxygen-Bonded Carbonyls • CO can sometimes bond through oxygen as well as carbon. • The O of CO has the ability to act as donor toward Lewis acids such as AlCl3 with the overall function of CO serving as a bridge between the two metals. 13.4 Ligands in Organometallic Chemistry | FIGURE 13.17 Oxygen-B Carbonyls. H3C Zr O C Mo C C O OC O (a) W C C O O (b) Al1C6H523 • Attachment of a Lewis acid to the O results in significant weakening and lengthening 13.4.2 Ligands Similar to CO of the C⏤O bond and a shift of the C⏤O stretching vibration to lower energy in the IR. Several diatomic ligands similar to CO are worth mention. Three—CS (thiocarbonyl), CSe (selenocarbonyl), and CTe (tellurocarbonyl)—are of interest in part for purposes of comparison with CO. The realm of CS complexes has been explored extensively since the first thiocarbonyl complex was reported in 1966,14 but the chemistry of complexes containing 15 the decrease in C—E stretching frequency, the major contributor to this phenomenon is the increasing mass of the heteroatom E. t structural and Other ligands are isoelectronic with CO and, not surprisingly, exhibit structural and ORGANOMETALLIC CHEMISTRY CN have been chemical parallels with CO. Two examples are CN - and N2. Complexes of CN - have been and Turnbull’s known longer than carbonyl complexes. Blue complexes (Prussian blue and Turnbull’s Ligands Similar to CO nts and inks for blue) containing the ion [Fe(CN)6]3 - have been used as pigments in paints and inks for antiallyCS, weaker CSe, and CTe are similar approximately to CO in their modesis in that they behave as both σweaker threebonding centuries. Cyanide a stronger s donor and a substantially * Unlike most s.* Unlike most p acceptor than CO; overall, it is close to CO in the spectrochemical series. donors and π acceptors and can bond to metals in terminal or bridging modes. e bonds readily organic ligands, which bond to metals in low formal oxidation states, cyanide bonds readily N m Me2N N TABLE 13.3 Complexes of CO, CS, CSe, and CTe N Cl Ru CE Cl N , 29, 519. v(C—E), cm−1 Ru—C Distance, nm CO 1934 1.829 CS 1238 1.793 CSe 1129 1.766 CTe 1024 1.748 Data from Y. Mutoh, N. Kozono, M. Araki, N. Tsuchida, K. Takano, Y. Ishii, Organometallics, 2010, 29, 519. NMe2 king, Inorg. Chem., CE *A comparison of these ligands in mixed ligand complexes of Fe is in C. Loschen, G. Frenking, Inorg. Chem., 2004, in 43, going 778. distance down this series is consistent with increasing π-acceptor • The decrease in Ru—C activity of the ligands, populating orbitals that are bonding with respect to the Ru—C bond. - scheme A, therefore, bent NO is considered the 2-electron donor NO ; by the neutral ligand model, it is considered a 1-electron donor. ORGANOMETALLIC CHEMISTRY Although these electron-counting methods in NO complexes are useful, they do not describe how NO actually bonds to metals. The use of NO + , NO, or NO - does not necNO Complexes essarily imply degrees of ionic or covalent character in coordinated NO; these labels are simply convenient means of counting electrons. • The NO (nitrosyl) ligand shares many similarities with CO. Useful information about the linear and bent bonding modes of NO is summarized in Figure 13.19. Many complexes containing each mode are known, and examples are also • Like CO, it is a σ donor and π acceptor and can serve as a terminal or bridging ligand E 13.18 Examples of NO Complexes. O Ni N O Linear Cl N Ir Ph3P O N Cr PPh3 CO Bent + O N O N Cr N O CC O O Bridging Cr NS NS complex • NO+ is isoelectronic with CO; therefore, in its bonding to metals, linear NO is considered by scheme A as NO+, a 2-electron donor; by method (B), linear NO is counted as a 3-electron donor. • By method A, bent NO is considered the 2-electron donor NO ; by the neutral ligand model, it is considered a 1-electron donor. - ORGANOMETALLIC CHEMISTRY Hydrides and Dihydrogen Complexes Hydride Complexes r 13 | Organometallic Chemistry As a ligand, H may be considered a 2-electron donor as hydride (:H-, method A) or a 1-electron neutral donor (H atom, method B). H in combination with other ligands. Such complexes may be made in a variety of ways. Probably the most common synthesis is by reaction of a transition metal complex with H . 2 The most common synthesis of hydrides is by reaction of a transition metal complex with H2. For example, For example, Co2(CO)8 + H2 h 2 HCo(CO)4 trans@Ir(CO)Cl(PEt3)2 + H2 h Ir(CO)Cl(H)2(PEt3)2 M H H Carbonyl hydride complexes can also be formed by the reduction of carbonyl com-s donation Dihydrogen Complexes plexes, followed by the addition of acid. For example, R = • In 1984, Kubas synthesized M(CO)3(PR3)2(H2), where M =+Mo or W and Co2(CO)8 + 2 Na h 2 Na [Co(CO)4] cyclohexyl or isopropyl - + H M H [Co(CO) h orbital HCo(CO) donated toHan empty on4 the metal, • The σ electrons in H2 can be 4] + p acceptance and the empty σ* orbital of the ligand can accept electron density from One of the most interesting aspects of transition-metal hydride chemistry is the rela- in H2 complexes Bonding an occupied d orbital of the metal. FIGURE 13.20 Bonding in * tionship between this ligand and the chemistry of the dihydrogen ligand, H2. Dihydrogen Complexes Dihydrogen Complexes. 13.4 Ligands in Organometallic Chemistry | 497 cyclic p systems, after which we will consider how p e will first describe linear and then olecules containing such systems can bond to metals. p r and then cyclic p systems, after which we will consider how The antibonding interaction a nodal plane perpendicular to the internuclear axis, but ORGANOMETALLIC CHEMISTRY 498 Chapter 13 | has Organometallic Chemistry systems can bond to metals. * near Pi Systems the bonding interaction has no such nodal plane. 498interaction ChapterPi 13 | aOrganometallic Chemistry The antibonding has nodal plane perpendicular to the internuclear axis, but Ligands with Extended Systems e simplest case of an organic molecule having linear p system ispethylene, which has Nexta is the three-atom system,interact the p@allyl radical, C3H5the . Inlowest this case, there are in four ways, with energy p molec the bonding interaction has no such nodal plane. single p bond resulting from the interactions of two 2 p orbitals on its carbon atoms. three 2 p orbitals to be considered, one from each of the carbon atoms participating in the interactions between neighboring p orbitals, and the en anic molecule having a linear p system is ethylene, which has interact in four ways, with the lowest energy p molecular orbital having all constructive Next is the three-atom p system, the p@allyl radical, C H . In this case, there are 3 5 Linear Pi Systems eractions ofthree these2pp orbitals result in one bonding and one antibonding p orbital, as shown: p system. The possible interactions are as follows: ing with the number of nodes between atoms. from the interactions of two 2 p orbitals on its carbon atoms. interactions between neighboring p orbitals, and the energy of the other pthe orbitals increas orbitals to be considered, one from each of the carbon atoms participating in the ing with the number of nodes between atoms. ls result in one and one antibonding p orbital, shown: p bonding system. possible interactions are asasfollows: p orbitals interacting Relative energy H CThe CH HC p orbitalsthe interacting HC CH CH CH2 2 p orbitals interacting 2 HRelative CH energyCH2 2C 2 2 CHenergy CH p orbitals interacting H2C Relative p* p* p* p* p p CH2 CH CHenergy CH2 p orbitals inter Relative 2 p orbitals interacting Relative energy pn pn e antibonding interaction has a nodal plane perpendicular to the internuclear axis, but n has a nodal plane perpendicular to the internuclear axis, but eno bonding interaction has no such nodal plane. p such nodal plane. p is the p system, radical, pNext system, the three-atom p@allyl radical, C3H5. Inthe thisp@allyl case, there are C3H5. In this case, there are ee 2p orbitals to be considered, one from each of the carbon atoms participating in the idered, one from each of the carbon atoms participating in the The lowest energy p molecular orbital for this system has all threeorbital p orbitals The lowest energy p molecular for interacting this system has all three p orbitals interacting system. as follows: ractions The are constructively, aspossible follows: interactions to give a are bonding molecular orbital. Higher in energy is Similar the nonbonding constructively, to give a bonding molecular orbital. energy is the nonbonding patterns can beintwo obtained longer p included systems Similar patterns can be obtained for longer p Higher systems; more for examples are (p in which a nodalenergy plane bisects theinmolecule, cutting through the central carbon cutting porbital orbitals interacting Relative CH2 The number of π molecular orbitals is equal to the number of carbons in number π system. n),CH p orbitals interacting Relative energy H C CH orbital (p ), which a nodal plane bisects the molecule, through central carbonin 13.21 . The ofthe p the molecular orbitals i molecular orbitals isthe equal to number of carbons n in Figure 13.21. The number ofinpFigure 2 2 atom. In this case, the p orbital on the central carbon does not participate in the molecular atom. In thisthe case, the p orbital on the central carbon does not participate in the molecular p system. the p system. p* orbital; a nodal plane passes through the center of this p orbital and thereby cancels it p* orbital; a nodal plane passes through the center of this p orbital and thereby cancels it * from participation. Highest in energy is theCyclic antibonding p orbital, in which there is an * Pi Systems Cyclic Pi Systems from participation. Highest in energy is the antibonding p orbital, in which there is an molecular orbitals have the same energy; p molecular orbitals having the same numnodes in cyclic p systems of hydrocarbons are degenerate (have the same energy). molecular orbitals orbital have the same energy; p molecular the same numotal p molecular diagram for cyclo@C thereforehaving be summarized as 3H3 can orbitals odes in cyclic p systems of hydrocarbons are degenerate (have the same energy). s: al p molecular orbital diagram 500 for cyclo@C can| therefore be summarized as 3H3 13 Chapter Organometallic Chemistry Cyclic Pi Systems p orbitals interacting Relative energy cyclo C3H3 : ORGANOMETALLIC CHEMISTRY cyclo-C3H3 p orbitals interacting Relative energy FIGURE 13.22 Molecular Orbitals for Cyclic Pi Systems. p orbitals interacting cyclo-C5H5 Relative energy p orbitals interacting Relative energy cyclo-C6H6 (Benzene) simple way to determine the p orbital interactions and the relative energies of the p systems that are regular polygons is to draw the polygon with one vertex pointed Each vertex corresponds the relative energy of a molecular orbital. Furthersimple way tothen determine the ptoorbital interactions and the relative energies of the number ofare nodal planes perpendicular to the of the increases pthe systems that regular polygons is to draw theplane polygon withmolecule one vertex pointed goes to higher energy, with the bottom orbital having zero nodes, the next pair of Each vertex cyclothenCcorresponds to the relative energy of a molecular orbital. Further4H4 s a single node, and so on. For example, this scheme predicts that the next cyclic p he number (cyclobutadiene) of nodal planes perpendicular to the plane of the molecule increases * m, cyclo@C4H4 (cyclobutadiene), would have molecular orbitals as follows: goes to higher energy, with the bottom orbital having zero nodes, the next pair of Relative energythat the next cyclic p a single node, and so on. For example, this scheme predicts 2-node p orbital cyclo@C4H4 (cyclobutadiene),One would have molecular orbitals as follows:* Two 1-node p orbitals One 2-node p orbital Relative energy One 0-node p orbital Two 1-node p orbitals r results are obtained for other cyclic p systems; two of these are in Figure 13.22. se diagrams, nodal planes are disposed For example, in cyclo@C4H4, One 0-nodesymmetrically. p orbital ngle-node molecular orbitals bisect the molecule through opposite sides; the nodal results are obtained for other p systems; two orbital of these Figure 13.22 are oriented perpendicularly to cyclic each other. The 2-node forare thisinmolecule also . diagrams, nodal planes. planes are disposed symmetrically. For example, in cyclo@C4H4, rpendicular We arethe now gle-node molecular bisect the sides; nodal his method may seemorbitals oversimplified, butmolecule the nodal through behavioropposite and relative energies areready to consider metal–ligand interactions involving such systems. We will 13.5 Bonding between Metal Atoms and Organic Pi Systems the following geom o consider metal–ligand interactions involving such systems. We will lest of the linear systems, ethylene, and conclude with ferrocene. ORGANOMETALLIC CHEMISTRY Systems Bonding Between Metal Atoms and Organic Pi Systems lexesLinear Pi Systems volve ethylene, C2H4, as a ligand, including the anion of Zeise’s salt, Pi–Ethylene Complexes In such complexes, ethylene commonly acts as a sidebound ligand with Ethylene commonly acts as a side bound ligand with the following geometry with respect to the metal: etry with respect to the metal: H H C C The hydrogen C M M Pt as shown. Ethylen C C C p@bonding electro H s donation p acceptance H can be donated ba Side bound ligand binding Bonding in Ethylene Complexes in ethylene complexes are typically bent back away from the metal, FIGURE 13.23 Bonding in orbital of the ligan density the metal in a sigma using its π-bonding electron pair. • Ethylene donates electrondonates density electron to the metal in atosigma fashion, using fashion, its Ethylene Complexes. p acceptance enco pair, as shown in Figure 13.23 . At the same time, electron density • Electron density can be donated back to the ligand in a π fashion from a metal d orbital to the empty π* k to the orbital ligand of in the a piligand. fashion from a metal d orbital to the empty p* If this picture . This is another example of the synergistic effect of s donation and with the measured ntered earlier with the CO ligand. tor, depending on the electron distribution between the metal and the ligand. The highest energy p orbital acts as an acceptor; thus, there can be synergistic sigma and pi interactions between allyl and the metal. The C i C i C angle within the ligand is generally near 120°, consistent with sp2 hybridization. ComplexesAllyl complexes (or complexes of substituted allyls) are intermediates in many reactions, some of which take advantage of the capability of this ligand to function in both a ORGANOMETALLIC CHEMISTRY Pi–Allyl The allyl group most commonly functions as a trihapto ligand, using delocalized π orbitals, or as a monohapto ligand, primarily σ bonded to a metal. h3-C h1-C 3H5: 02 Chapter 13 | Cl Organometallic Chemistry Ni Pd C6H5 3H5: H O C CO C OC Mn C H C C H O O C6H5 Pd Cl FIGURE 13.24 Examples of Allyl Complexes. CH2 h3 and h1 fashion. Loss of CO from carbonyl complexes containing h1@allyl ligands often 1 3 results in conversion of h @ to h For example, from carbonyl complexes@allyl. containing η1-allyl ligands often results Loss of CO η1- to η3-allyl. 3Mn1CO254- 1h1-C + C3H5Cl 3H52Mn1CO25 + ClH ¢ or hn Other metal orbitals of suitable symmetry H2C C C C M C C Mn1CO25 1h3-C3H52Mn1CO24 + CO Mn1CO24 in conversion of FIGURE 13.25 Bonding in h3-Allyl Complexes. px H H The [Mn(CO)5] - ion displacesdxzCl - from allyl chloride to give an 18-electron product containing h1@C3H5. The allyl ligand switches to trihapto when a CO y is lost, preserving Identify the transition metal in the following 18-electron complexes: a. (h5@C5H5)(cis@h4@C4H6)M(PMe3)2(H) (M) = second@row transition metal) ORGANOMETALLIC CHEMISTRY 5 Other Linear Pi Systems b. (h @ C5H5)M(C2H4)2 (M = first-row transition metal) 13.5.2 Cyclic Pi Systems π systems have the possibility of isomeric ligand forms (cis and trans). • Butadiene and longer conjugated Cyclopentadienyl (Cp) Complexes Larger cyclic ligands may have a π systemgroup, extending through part in ofa the ring. The cyclopentadienyl C5H5, may bond to metals variety of ways, with many examples known of the h1@, h3@, and h5@bonding modes. The discovery of the first cyclo1,3-isomer hasferrocene, a 4-atom systemincomparable toorganometallic butadiene • Cyclooctadiene (COD); the pentadienyl complex, was aπlandmark the development of chemistry and double stimulated bonds, the search one for other containing p@bonded organic •1,5-cyclooctadiene has two isolated or compounds both of which may interact with ligands. Substituted cyclopentadienyl ligands are also known, such as C5(CH3)5, often manner similar to ethylene. abbreviated Cp*, and C5(benzyl)5. H Ni H H3C Mo 1CO23 CH3 H3C Fe H 3C FIGURE 13.26 Examples of Molecules Containing Linear Pi Systems. Ti Ti Examples of Molecules Containing Linear Pi Systems 1CH322 P P 1CH322 Zr H a metal in a ORGANOMETALLIC CHEMISTRY 504 Chapter 13 | Organometallic Chemistry Cyclic Pi Systems The ten group orbitals arising from the C5H5 ligands are shown in Figure 13.27. Cyclopentadienyl Complexes The process of developing the molecular orbital picture of ferrocene now becomes one of matching the group orbitals with the s, p, and d orbitals of appropriate symmetry on Fe. Ferrocene, (η5-C5H5)2FeWe will illustrate one of these interactions, between the dyz orbital of Fe and its appropriate group orbital (the 1-node group orbitals are shown in Figure 13.27). This interaction results in a bonding and an antibonding orbital: Ferrocene is the prototype of a series of sandwich compounds, the metallocenes, with the Selected Atomic Orbitals by Gary O. Spessard and L. Miessler. 5 Gary 5 2 Reprinted by permission. formula C H ) M. FIGURE 13.27 Group Orbitals for C5H5 Ligands of Ferrocene by Gary O. Spessard and Gary L. Miessler. Reprinted by permission. Interactions, between the dyz orbital of Fe and the 1-node group orbitals. in which one cyclopentadienyl ligand has been modified into h @C5H6 (Figure 13.30). + HFerrocene, however, is by no means chemically inert. It undergoes a variety of reacCo Co alticinium ion tions, including many on the cyclopentadienyl rings. A good example is that of electrophilic acyl substitution (Figure 13.31), a reaction paralleling that of benzene and its derivatives. FIGURE 13.30 Reaction of PF3)4 + organic productsIn general, electrophilic aromatic substitution reactions are much more rapid for ferrocene Cobalticinium with Hydride. than for benzene, an indication of greater concentration of electron density in the rings of - Ferrocene undergoes a variety of reactions, including many on the cyclopentadienyl rings. 8e the sandwich compound. Among the most interesting ferrocene-containing compounds is a molecule that was + soughtcompound for many years before its synthesis in 2006, hexaferrocenylbenzene, a type of moleceutral, 18-electron sandwich H ular “Ferris wheel” (Figure 13.32). This compound, originally obtained from the reaction 4 H dified into h @C5H6 (Figure 13.30 ). of hexaiodobenzene and diferrocenylzinc, has six ferrocenyl groups as substituents on a -of hexaferrocenylbenzene benzene ring. The highly crowded nature is illustrated by the + H ly inert. It undergoes a variety of 29reacCo Co alternating up/down arrangement of ferrocenes around the benzene ring, with the benzene . A good example is that of electrophilic itself adopting a chair conformation with alternating C i C distances of 142.7 and 141.1 pm. metallocenes—with two atoms, rather than one in the center of a sandwich ling that of benzene and its Binuclear derivatives. Reaction of Cobalticinium FIGURE 13.30 Reaction ofHydride. structure—are also known. Perhaps the best known ofwith these metallocenes is decamethyldiztions are much more rapidincocene, for ferrocene (h5@C5Me5)2Zn2Cobalticinium (Figure 13.33), which prepared from decamethylzincocene, with was Hydride. 18 e ORGANOMETALLIC CHEMISTRY ation of electron density in the rings of O C CH3 + ning compounds is a molecule that was H hexaferrocenylbenzene, a type of molec- + - H+ + 3CH3CO4 Fe Fe d, originally obtained from the reaction ferrocenyl groups as substituents on a aferrocenylbenzene is illustrated by the Electrophilic Acyl Substitution in Ferrocene. und the benzene ring, with the benzene g C i C distances of 142.7 and 141.1 pm. O C Fe CH3 FIGURE 13.31 Electrophilic Acyl Substitution in Ferrocene. ORGANOMETALLIC CHEMISTRY Molecular “Ferris wheel" Binuclear metallocenes—with two atoms, rather than one in the center of a sandwich structure—are also known. (hexaferrocenylbenzene) was sought for 508 many Chapter 13 | Organometallic Chemistry years before its synthesis in 2006. 508 Chapter 13 | FIGURE 13.32 Hexaferrocenylbenzene. Organometallic Chemistry FIGURE 13.32 Hexaferrocenylbenzene. Fe Zn Fe Fe Fe Fe Fe Fe Fe O Zn Zn Fe Fe O Zn Zn FIGURE 13.33 Decamethylzincocene and Decamethyldizincocene. Fe Zn FIGURE 13.33 Decamethylzincocene and Decamethyldizincocene. Fe30 Particularly notable is (h5@C5Me5)2Zn2, the first exam(h5@C5Me5)2Zn, and diethylzinc. ple of a stable molecule with a zinc–zinc bond; moreover, its zinc atoms are in the exceptionally rare +1 oxidation state. Among Group 12 elements, mercury by far exhibits this oxidation state most often, with the best known example the Hg22 + ion*; cadmium and zinc ORGANOMETALLIC CHEMISTRY Complexes Containing Cyclopentadienyl and CO Ligands “Inverse” sandwich 5 (h @C5 Many complexes are known containing both Cp and CO ligands. ple of FIGURE 13.35 Complexes Group 4 Group 5 Group 6 Group 7 Group 8 tionall Containing C H and CO. O C O CO oxidat O Mo CO Ti V1CO2 Mn CO CO O CO compo Ca Similar structures Similar structures Similar structure Similar structures 5 (h @C5 for Zr, Hf for Nb, Ta for W for Tc, Re with a CO Fe V CO contai Ca CO A O O OO C O CC O by the C O O O C CO C Re Re Fe Fe Cr Cr cyclic C C OC O C C CC C O O OO O O prepar Similar structures Similar structure FIGURE 13.34 An Inverse • Ca(I) ions on the outside and the cataly for Mo, W for Ru Sandwich Compound, O O cyclic π ligand 1,3,5-triphenylbenzene C tion is C CO [(thf) Ca{m-C H -1,3,5-Ph } 3 6 3 3 CO in between. Os Os Mo Mo of a + Ca(thf)3]. (Molecular structure 13.5 Bonding between Metal Atoms and Organic Pi Systems | 509 5 5 4 OC OC Similar structure for Cr OC C O drawing created from CIF data, with hydrogen atoms omitted for clarity.) + FIGURE 13.36 Examples of Molecules Containing Cyclic Pi Systems. Comp Many sandw ORGANOMETALLIC CHEMISTRY Fullerene Complexes •As immense π systems, fullerenes were recognized early as ligands to transition metals. •Fullerene-metal compounds have been prepared for a variety of metals. 510 Chapter 13 | Organometallic Ch •These compounds fall into several structural types: Adducts to the oxygens of osmium tetroxide. Example: I i N O O N Os O O C60(OsO4)(4-t-butylpyridine)2 Complexes in which the fullerene itself behaves as a ligand. Examples: Fe(CO)4(η2-C60), Mo(η5-C5H5)2(η2-C60), [(C6H5)3P]2Pt(η2-C60) FIGURE 13.37 Structure of C60(OsO4)(4-t-butylpyridine)2. T Rb3C The i Addu The The X struct the do OsO4 1:1 ad in Fig Fulle As a to me (Figu 60 4 2 in Figure 13.37. Fullerenes as Ligands39 As a ligand, C60 behaves primarily as an electron-deficient alkene (or arene), and it bonds to metals in a dihapto fashion through a C i C bond at the fusion of two 6-membered rings (Figure 13.38). There are also instances in which C60 bonds in a pentahapto or hexahapto fashion. Dihapto bonding was observed in the first complex to be synthesized, in which C60 acts as a ligand toward a metal, [(C6H5)3P]2Pt(h2@C60),40 also shown in Figure 13.38. A common route to the synthesis of complexes with fullerene ligands is by displacement of other ligands, typically those weakly coordinated to metals. For example, the platinum complex in Figure 13.38 can be formed by the displacement of ethylene: ORGANOMETALLIC CHEMISTRY Fullerenes containing encapsulated (incarcerated) atoms, called incarfullerenes. These may contain one, two, three, or four atoms, sometimes as small molecules, inside the fullerene structure. Examples: UC60, LaC82, Sc2C74, Sc3C82 Intercalation Compounds of Alkali Metals [(C6H5)3P]2Pt(h2@C2H4) + C60 h [(C6H5)3P]2Pt(h2@C60) The d electron density of the metal can donate to an empty antibonding orbital of a fullerene. This pulls the two carbons involved slightly away from the C60 surface. In addition, the distance between these carbons is elongated slightly as a consequence of this interaction, which populates an orbital that is antibonding with respect to the C i C bond. This increase in C i C bond distance is analogous to the elongation that occurs when ethylene and other alkenes bond to metals (Section 13.5.1). In some cases, more than one metal can become attached to a fullerene surface. A spectacular example is [(Et3P)2Pt]6C6041 in Figure 13.39. In this structure, the six (Et3P)2Pt units are arranged octahedrally around the C60. These contain alkali metal ions occupying interstitial sites between fullerene clusters. Examples: NaC60, RbC60, KC70, K C Fullerene as Ligand FIGURE 13.38 Bonding of C60 3 to Metal. 60 Reprinted by permission. Bonding of C60 to Metal by Gary O. Spessard and Gary L. Miessler. Reprinted by permission. As a ligand, C60 behaves primarily as an electrondeficient alkene (or arene), and it bonds to metals in a dihapto fashion through a C—C bond at the fusion of two 6-membered rings Bonding of C60 to Metal. Center of C60 Pt Pt ORGANOMETALLIC CHEMISTRY Pt Complexes of other fullerenes have Detail also been prepared. 512 Chapter 13 | Organometallic Chemistry 2-C70) Ir(CO)Cl(PPh3)2. Stereoscopic View of (η omplexes of other fullerenes have also been prepared. An example is 70)Ir(CO)Cl(PPh3)2 (Figure 13.40). As in the case of th