CHE 425: ORGANOMETALLIC CHEMISTRY SOURCE: OPEN ACCESS FROM INTERNET; Striver and Atkins Inorganic Chemistry Lecturer: Prof. O. G. Adeyemi

ORGANOMETALLIC CHEMISTRY

Definitions:

Organometallic compounds are compounds that possess one or more metal-carbon bond.  The bond must be “ionic or covalent, localized or delocalized between one or more carbon atoms of an organic group or molecule and a transition, lanthanide, actinide, or main group metal atom.”

 Organometallic chemistry is often described as a bridge between organic and inorganic chemistry.

 Organometallic compounds are very important in the chemical industry, as a number of them are used as industrial catalysts and as a route to synthesizing drugs that would not have been possible using purely organic synthetic routes.

Coordinative unsaturation is a term used to describe a complex that has one or more open coordination sites where another ligand can be accommodated.

 Coordinative unsaturation is a very important concept in organotrasition metal chemistry.

Hapticity of a ligand is the number of atoms that are directly bonded to the metal centre.

 Hapticity is denoted with a Greek letter η (eta) and the number of bonds a ligand has with a metal centre is indicated as a superscript, thus η1, η2, η3, ηn for hapticity 1, 2, 3, and n respectively.

 Bridging ligands are normally preceded by μ, with a subscript to indicate the number of metal centres it bridges, e.g. μ2–CO for a CO that bridges two metal centres.

Ambidentate ligands are polydentate ligands that can coordinate to the metal centre through one or more atoms.

– – –  For example CN can coordinate via C or N; SCN via S or N; NO2 via N or N. The term can also be used to describe instances where a ligand can behave as monodentate or a chelating ligand.

Bite angle is the ligand–metal–ligand angle formed when a polydentate ligand coordinates to a metal centre.

1 A Chelate (Greek word for “claw” is a polydentate ligand that forms a ring that includes a metal. Examples are EDTA, acac–, en.

Heterobimetallic describes a complex in which there are two different metal centres.

Homobimetallic complexes have two metal centres that are the same elements. These need not have identical ligands or coordination number, but are usually found as symmetric dimers.

Homoleptic complexes are compounds in which all the ligands that bound to the metal centre are identical.

A ligand is a molecule or ion that is bonded directly to a metal centre, usually by a covalent or coordinate bond.

Monodentate ligands have only one point of attachment to the metal centre and occupy one coordination site only. Examples of monodentate ligands are NH3, CO, NMe3, H2O and PMe3.

Polydentate ligands have more than one point of attachment to the centre and occupy more than one coordination site.

 By definitions above bidentate ligands and ambidentate ligands are special cases of polydentate ligands.

The primary coordination sphere of a metal involves the set of ligands closest to the metal that are directly bonded. Mobile cations and anions are said to be in the outer or secondary sphere.

HISTORICAL OVERVIEW OF ORGANOMETALLIC CHEMISTRY

Many examples of organometallic compounds involving alkyls, alkenes, alkynes, carbonyls and aromatics are known. Organometallic chemistry is a relatively recent area of chemistry (40-50 years), even though organometallic compounds have been known since 1760; e.g. As2Me4O was made in 1760. Organometallic Chemistry has really been around for millions of years Naturally occurring Cobalimins contain Co—C bonds Vitamin B12:

Some other historical landmarks are:

1827 Zeise’s salt: Na[PtCl3(C2H4)]

1849 Frankland: [(C2H5)2Zn]

1890 Mond: [Ni(CO)4] – the first binary metal carbonyl, and useful in purifying Ni

1931 Hieber: [Fe(CO)4H2] – the first hydride to be made.

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1951 Ferrocene: [Fe(cp)2] – the first sandwich complex.

Nobel -Prize Winners in the area of Organometallic: Victor Grignard and Paul Sabatier (1912) Grignard reagent K. Ziegler, G. Natta (1963) Zieglar-Natta catalyst E. O. Fisher, G. Wilkinson (1973) Sandwich compounds K. B. Sharpless, R. Noyori (2001) Hydrogenation and oxidation Yves Chauvin, Robert H. Grubbs, Richard R. Schrock (2005) Metal-catalyzed alkene metathesis

BONDING

 Bonding is essentially covalent and can be Metal–Carbon σ (as in metal carbonyls and alkyls) or Metal–Carbon  bonds (as in metal alkenes, alkynes and arenes).

-– +  Ionic compounds are often formed with electropositive metals, e.g. (C6H5)3C Na and – 2+ (C5H5 )2Ca . These ionic compounds are insoluble in hydrocarbon solvents and are very sensitive to air and moisture.

ELECTRON COUNTING

The 18-electron rule

The 18-electron rule is like the octet rule to the organometallic chemists.

 It is useful in predicting the reactivity of organomettalic compounds.

 It is also often violated

 18-electron rule is also known as the effective atomic number rule (EAN rule).

 The rule is used to determine the number of valence electrons in a complex to determine whether that complex is likely to be stable or not.

 18 electrons arise from electrons from completely filled five d, three p, and one s orbitals.

 “A stable complex is obtained when the sum of metal d-electrons, the electron donated from the ligands, and the overall charge of the complex equals 18.”

3 Counting electrons in organometallic complexes

Two distinct methods are used to count electrons: the neutral or covalent method and the effective atomic number or ionic method.

 These are just two different methods of an accounting system that give the same final answer. All you need to do is to understand one and keep to that to avoid confusion.

What are d-electrons

The configuration of transition metal as [Ar]4s23dn.

 This representation is only true for isolated metal atoms.

 When a metal ion is put into an electronic field (surrounded with ligands), the d-orbitals drop in energy and fill first.

 The organometallic chemist therefore, considers the transition metal valence electrons to be all the d-electrons.

 Thus a transition metal such as Ti in the zero oxidation state attains a d4, and not d2 configuration.

 Thus for zero-valent metals, the electron count simply corresponds to the column it occupies in the periodic table. Fe, for example, in the 8th column is d8 and not d6; Re3+ is d4 (seventh column for Re, and then add 3 position charges or subtract three negative ones).

 That is, d-occupancy = group number – oxidation state

Electronic contribution of ligands

Method 1: The ionic (charged) model

The basic premise of this method is that we remove all of the ligands from the metal and, if necessary, add the proper number of electrons to each ligand to bring it to a closed valence shell state.

 For example, if we remove ammonia from a complex, NH3, has a completed octet and acts as a neutral molecule. When it bonds to the metal centre if does so through its lone pair (in a classic Lewis acid-base sense) and there is no need to change the oxidation state of the metal to balance charge.

 Ammonia is called a neutral 2-electron donor.

 In contrast, if we remove a methyl group from the metal and complete its octet, then we – formally have CH3 .

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 If we bond this methyl anion to the metal, the lone pair forms one metal-carbon bond and the methyl group acts as a 2-electron donor ligand.

 To keep charge neutrality we must oxidize the metal by one electron (i.e. assign a positive charge to the metal). This, in turn, reduces the d-electron count of the metal centre by one.

Method 2: the covalent (neutral) model

The major premise of this method is that we remove all of the ligand from the metal, but rather than take them to a closed shell state, we do whatever is necessary to make them neutral.  Consider ammonia again: when we remove it from the metal, it is a neutral molecule with one lone pair of electrons. Therefore, as with ionic, ammonia is a neutral 2-electron donor. For methyl, we diverge from the ionic model. When we remove it from the metal and make the fragment neutral we have a neutral radical.  Both the metal and methyl radical must donate one electron each to form a metal-carbon bond.

 The methyl group is, therefore, a one-electron donor, not a two-electron donor as it is under the ionic formation.

 Notice that this method does not give us any immediate information about the oxidation state of the metal, so we must go back and assign that later.

Some commonly encountered ligands donate the following numbers of valence electrons.

 1-electron donor: H• (in any bonding mode), and terminal Cl•, Br•, I•, R• (e.g. R = alkyl or Ph) or RO•; 2  2-electron donor: CO, PR3, P(OR)3, R2C=CR2 (η -alkene), R2C: (carbine); 3 • • • • •  3-electron donor: η -C3H5 (allyl radical), RC (carbine), μ-Cl , μ-Br , -I , μ-R2P ; 4 4  4-electron donor: η -diene, η -C4R4 (cyclobutadienes); 5 • • • • •  5-electron donor: η -C5H5 , μ3-Cl , μ3-Br , μ3-I , μ3-R2P ; 6 6 6  6-electron donor: η -C6H6 (and other η -arenes, e.g. η -C6H5Me);  1- or 3-electron donor: NO

Counting electrons provided by bridging ligands, metal-metal bonds and net charges requires care.

• •  When X (X = CI, Br, I) or R2P the ligand uses the unpaired electron and a lone pair to bridge two metal centres, i.e. one electron is donated to M, and two to the second metal, M´  In a doubly bridged species such as (CO)2Rh(μ-Cl)2Rh(CO)2 the μ-Cl atoms are equivalent as are the Rh atoms, and the two Cl bridges together contribute three electrons to each Rh.  A bridging H• provides only one electron in total, shared between the metal atoms it bridges, – e.g. [HFe2(μ-CO)2(CO)6] .  M—M single bond provides each M atom with one extra electron; an M=M double bond contributes two electrons to each metal.

Summary for 18-electron rule

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1. Assume that the metal has an oxidation state of zero

2. Add or subtract electrons according to the overall charge on the complex

3. Add electrons from the ligands

4. Consider a M―M to provide one electron to each metal; an M=M bond to provide two electrons to each metal.

5. Consider bridging two-electron ligands such as CO, CR2 to contribute one electron to each metal.

Exceptions to the rule

1. Electron-rich complexes of d8 transition metals.

Metals such as Rh(1), Ir(II), Pd(II) have a strong preference for square planar 16-electron 2– 2 – configuration, e.g., [Ni(CN)4] , [PtCI3(η -C2H4]

2. High steric hindrance: Ligands that are so bulky that not enough can be coordinated around a metal (there are few of such metals at the left hand side of the transition metal series) to obtain 18-electron count.

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CH3

H3C CH3

(Me3Si)2HC H3C CH3 Cr CH(SiMe3)2 Ti (Me3Si)2HC H3C CH3

CH3 H3C 9 VE

CH3

14 VE

3. Many organometallic compounds obey the 18-electron rule, but classical coordination compounds often have widely differing electron counts. Why?

 With 6 σ-donor ligands and the metal t2g and eg levels, three classes of complexes can be seen:

 Class I:  is small: valence electron (VE) is less than or greater than 18. T2g is non- bonding, eg is weakly antibonding. 12–22 VE can be accommodated.

2– IV 0 e.g.: [TiF6] Ti , d 12 VE 3– III 4 [Mn(CN)6] Mn , d 16 VE 2+ II 7 [Co(NH3)6] Co , d 19 VE 2+ 2+ 10 [Zn(en)3] Zn , d 22 VE

 Class II: M is 2nd/3rd row transition metal and Δ is larger. VE is less than or equal to 18. T2g is nonbonding and occupied with 0-6 electrons, eg is antibonding and unoccupied.

VI 0 e.g.: WCl6 W , d 12 VE 2– IV 2 WCl6 W , d 14 VE 2– IV 6 PtF6 Pt , d 18 VE 2– II 8 PtCl4 Pt , d 16 VE

 Class III: Δ large; VE = 18. These are organometallic complexes. T2g becomes bonding and is fully occupied, eg is antibonding and empty. The 18-electron rule is obeyed as far as steric constraints will allow.

– – 6 e.g.: V(CO)6 V , d 18 VE 0 8 Fe(CO)5 Fe , d 18 VE 1 8 cpCo(CO)2 Co , d 18 VE 0 8 Ni(COD)2 Ni , d 18 VE

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Examples:

Rationalize the following: Cr(CO)6, Ni(PF4)4, Fe(CO)4PPh3 and Mn2(CO)8

Answer: Fe 8e- 2Mn 14e- Cr 6e- Ni 10e- 4CO 8e- 10CO 20e- - - - - 6CO 12e 4PF3 8e Ph3P 2e Mn—Mn 2e 18e- 18e- 18e- 36e- = 18e-/Mn

-BONDING LIGANDS

1. Common -bonding ligands are:  Carbon monoxide  Isocyanides, and  Substituted phosphines, arsines, stibbines or sulfides,  Nitric oxide (nitrogen(II) oxide), and  Various molecules with delocalized -orbitals such as pyridines, 2,2-bipyridine, e.t.c.

2. Metals that could form complexes with the above ligands display low positive, zero, or negative formal oxidation states. 3. Properties of complexes:  These ligands then stabilize low oxidation states.  This property is associated with the fact that in addition to lone pairs of electrons, the ligands possess vacant -orbitals, which accept electrons density from filled metal orbitals fo form a type of -bonding that supplements the -bonding arising from lone- pair donation.  The ability of ligands to accept electron density is then called -acidity (from Lewis sense).

In this introductory lecture, we will limit ourselves to complexes of carbon monoxide.

METAL CARBONYL COMPLEXES

Ni(CO)4, the first metal carbonyl, was first discovered in 1888. Metal carbonyls belong to a very important class of organometallic compounds. Almost all transition metals form compounds with carbon monoxide as ligand.

Three important points about these compounds are:

 Carbon monoxide is not ordinarily considered a very strong Lewis base and yet it forms strong bonds with the metals it complexes with.

 The metals are always in low oxidation state, usually 0, but sometimes +1 or -1.

 The 18-electron rule is obeyed by at least 99% of all the organometallic complexes.

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Examples of stable carbonyl complexes of the first row transition metals are given in the Table below.

Stable metal carbonyls

Mononuclear V(CO)6 Cr(CO)6 Fe(CO)5 Ni(CO)4 Dinuclear Mn2(C0)10 Fe2(CO)9 Co2(CO)8 Trinuclear Fe3(CO)12 Tetranuclear Co4(CO)12 Hexanuclear Co6(CO)16

 The only exception to the 18-electron among the mononuclear carbonyl complexes is hexacarbonylvanadium(0), V(CO)6, which is paramagnetic and is a 17-electron molecule.

 V(CO)6 does not dimerise to give the 18-electron analogue like Mn2(CO)10 and Co2(CO)8 because each metal would attain a coordination number of 7, which may present too much steric hindrance to allow stability.

 This may also be due to ligand repulsion that can overcome weak M—M bond or present a kinetic barrier to dimerization.

 V(CO)6 is, however, less stable than other carbonyl complexes that obey the 18-electron – rule. It decomposes at 70 ℃ and readily forms [V(CO)6] to form the 18-electron anion.

 The tendency to form more complex carbonyls increases from left to right in the first row transition metal series.

 Second and third-row transition metals form parallel complexes to the ones given in the table above; e.g. Mo(CO)6, Tc2(CO)10 and Re2(CO)10

 Generally the carbonyls of the first row transition metals are more stable than their analogues in the second and third row series. For example, Fe(CO)5 > Ru(CO)5 > Os(CO)5; Fe2(CO)9 > Ru2(CO)9 > Os2(CO)9.

 Increasing strength of M—M interaction as one descends the transition metal series leads to enhanced stability for Os3(CO)12 and Ru3(CO)12 relative to the mononuclear, M(CO)5.

Synthesis of metal carbonyls

(a) From the metal: By direct action of CO on finely divided metal. Only [Ni(CO)4] and [Fe(CO)5] are so made.

9 o Ni + 4 CO 25 C 1 bar [Ni(CO)4]

150oC Fe + 5 CO [Fe(CO) ] 100 bar 5 o  [Ni(CO)4] is a highly toxic colourless liquid; b.p. 34 C. Purification of nickel by Monds process is attained by decomposing [Ni(CO)4]. o  [Fe(CO)5] is a yellow liquid; b.p. 103 C. [Fe(CO)5] is very sensitive to light and heat and decomposes in air.  Photolysis of [Fe(CO)5] yields an orange yellow solid [Fe2(CO)9].

h [Fe (CO) ] (m.p. 100 oC) [Fe(CO))5] 2 9

(b) From metal salts: by reduction of metal salts in the presence of CO.

benzene, AlCl CrCl + Al 3 [Cr(CO) ] + AlCl 3 300 bar CO 6 3

i AlEt3, Pr O [Mn (CO) ] Mn(OAc)2 + 10 CO 2 10

o  [Cr(CO)6] is colourless, diamagnetic, air-stable solid; m.p. is 154 C. It has18 VE.  [Mn2(CO)10] is a yellow crystalline, diamagnetic solid. It also obeys 18-electron rule.

300 bar CO 2 CoCO + 2 H + 8 CO [Co (CO) ] + 2CO + 2H O 3 2 130 oC 2 8 2 2

 [Co2(CO)8] is orange brown solid, diamagnetic; it obeys the 18-electron rule.

300 bar CO Co (CO) + 2 CO + 2 H O 2 CoCO3 + 2 H2 + 8 CO o 2 8 2 2 130 C orange-brown, diamagnetic, 18 VE

STRUCTURES OF METAL CARBONYLS

Metals will bind enough CO ligands to obey the 18-electron rule.

Complexes can be monomeric or dimeric or trimeric – and feature metal-metal bonds to attain the 18-electron configuration

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Binary Transition-metal Carbonyl complexes

CO CO CO OC OC CO Co CO M CO OC OC Co CO CO OC CO Solid state M = V, Cr, Mo, W OC OC CO Co CO OC CO OC CO Mn OC CO Co OC OC CO Mn CO CO CO OC CO Solution

Group 8

CO OC Group 9: M = Co, Rh M CO CO OC CO OC CO Ir OC CO M = Fe, Ru, Os OC Ir Ir CO CO Ir OC CO OC CO O OC C CO OC OC CO OC Fe CO CO OC CO Fe Fe Fe OC Fe CO CO CO OC OC OC C CO CO O OC CO M O CO OC C CO OC CO CO CO M M M CO OC OC M CO CO M M OC CO CO OC CO OC CO CO

M = Ru, Os

11 BONDING IN METAL CARBONYL

Two modes of bonding contribute to bonding of CO to metals: a) Overlap of filled carbon p-orbital with a d-orbital of the metal to form a -type bond. Thus, electron flows from carbon to metal. This increases the electron density on the metal atom. To reduce the electron density the metal by pushing electron back to the ligand (back-bonding). (b) Overlap of a filled d or hybrid dp—metal orbital with the empty p-orbital on CO, which can act as a receptor of electron density. This constitutes a contribution from -bond.

Note that this bonding mechanism is synergic, since back-bonding will tend to make the CO as a whole negative (increasing its basicity); at the same time the electron drift to the metal in the -bond tends to make the CO positive. In other word the two effects assist each other. The bond-order for M—CO bond is therefore greater than 1 but less than 2.

Physical evidence for bond-order/strength of M—CO bond

 Dipole moment of an M—C bond is very low (0.5 D)  Bond length: Shorter M—C and longer C—O bonds as compared to single and triple bonds; C—O bond is 1.128 Å, while in metal carbonyl complexes it is 1.15 Å.  Vibrational spectroscopy (studying CO stretching frequencies): see later.

A molecular orbital description of CO shows a carbon-centred lone pair (HOMO) and degenerate * levels (LUMO’s) p

s p p

s p s

s

C O

C O C O

T* (LUMO) sσ* (HOMO)

1. The carbon lone pair (sσ*) interacts with a vacant metal orbital in a -donor fashion.

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COORDINATION MODES

O O O C C C M M M M M M terminal doubly triply bridging bridging   

Doubly bridging is quite common.

CO CO CO CO CO CO

OC Cr CO OC Mn Mn CO OC OC OC CO CO CO octahedral octahedral  = 2000 cm-1 -1 CO CO = 2044, 2013, 1983 cm

CO OC CO Fe CO Ni OC CO OC CO CO tetrahedral trigonal bipyramid -1 -1 CO = 2057 cm CO = 2034, 2013 cn

INFRARED SPECTROSCOPY

As seen previously, the stretching frequency (VCO) of the carbonyl bond can be very characteristic. From ~ 2100 → 1650 cm-1

(i) Influence of charge -1 VCO (cm ) Ni(CO)4 2060

13 10 - d Co(CO)4 1890 2- F2(CO)4 1790 ______+ Mn(CO)6 2090 6 d Cr(CO)6 2000 - V (CO)6 1860

Charge influences extent of backbonding

→ More –ve charge → better back bonding ∴ strong M-C, weaker C-O

More +ve charge, → worse back bonding ∴ weaker M-C, stronger C-O

(ii) Properties of other ligands By replacing a CO group with a ligand which is a weaker ___ acceptor, then more electron deficient metal centre, less efficient back bonding (stronger C-O bond)

BONDING MODES OF CO O O O C C C M M M M M M terminal doubly triply bridging bridging   

Free Terminal ͵U2 ͵U2

VCO 2143 2120-1850 1850-1750 1730-1620 (CM-1)

4. SPECTROSCOPIC CHARACTERIZATION OF METAL CARBONYLS

(a) MASS SPECTROMETRY

As neutral, volatile molecules, mass spectrometry is useful to observe parent molecular ions.

Also, a sequential loss of carbonyl ligands can be seen and aids characterization.

(see figure 5)

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O O

C C

M M M M C C

O

CO bridges often occur in pairs and can be in equilibrium with the non-bridging mode e.g. O O C OC CO CO C CO OC

CO Co Co Co Co Co CO C C CO C O O C CO O O

D 3d C2V (in solution)

M C O Dative bond

M(σ) ← CO(σ) M-C bond is strengthened

2) The __ orbital accepts electron density from an occupied d-orbital. ∴CO is a strong ------acid

This process is known as ‘back-bonding’

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M-C Bond strengthened C-O bond weakened

So, both processes operate in opposite directions but are supportive i.e. Synergic bonding

(a) More electron rich metal centre, better back donation (weaker C-O bond) The M-CO bond in the trans position is strengthened, and the C-O bond weakened.

→ The higher the v(CO) frequency, the more labile the carbonyl ligand. e.g Ni(CO)3{PR3} v(CO) P(OPH)3 2085 PPh(OPh)2 2080 PPh2(OPh) 2075 PPh3 2069 t P B43 2056

Orders of ╥-acceptor strength can be established :-

NO > CO > RNC > PF3 > PCl3 > PCl2R > P(OR)3 > PR3 > RCN > NH3

e.g. [Ru(CO)2(PPh3)2 {L}] PPh3 OC Ru L 31P-{'H} NMR is a singlet OC

PPh3

L V(CO) (cm-1) CNC6H4Me 1899, 1865 Ch2CH2 1955, 1900 CS 1962, 1900 CF2 1982, 1912 SO2 2000, 1932 O2 2005, 1945 CS2 2012, 1945 NO (+) 2065, 2014 Increasing ╥-acidity of ligands Decreasing metal-based electron density

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(iii) SYMMETRY Number and intensities of CO bands depends on the symmetry around the central atom

The symmetry of a complex can often be determined by counting the number of infrared bands.

The expected number of IR- active bands can be derived from group theory.

(See figure 6)

5. REACTIVITY OF METAL CARBONYLS

(i) SUBSTITUTION REACTIONS CO ligands can be substituted by other ligands, thermally or photochemically.

-CO L' L M - CO L M L M - L' n +CO n n vacant coordination site

Second row complexes usually react faster than first- or third- row analogues.

(Some substitution reactions (typical) are shown in figure 7)

Photochemical substitution involves mild reaction conditions and allows synthesis of highly substituted or more labile products. hv . L . . Co2(CO)8 2 Co(CO)4 Co(CO)4 + Co(L) (CO)3

18VE 17VE

Co2(CO)7L 18VE product

(ii) REDUCTION (METAL CARBONYL ANIONS) Reduction of metal carbonyls gives anionic complexes. They are highly reactive and there is rich chemistry.

17 - Fe(CO)5 + NaOH Na[H Fe(CO)4]

NaH

2- Na2[Fe(CO)4] THF Co (CO) + 2 Na 2 Na+[Co(CO) ]- 2 8 4

Carbonyl metallates are strong bases/nucleophiles and basicity increases within a group.

First < second < third

Relative basicity of carbonyl anions - 1 Cp Co(CO)4 - 4 = Cp Cr (CO)3 - 67 C5H5 Cp Mo (CO)3

Co W (CO) - 550 O 3 - 5.5 x 106 Cp= Cp Ni (CO) - 7 ydo Cp Fe (CO)2 7 x 10 pentadienyl

METAL HYDRIDE COMPLEXES Hieber first discovered [FeH2(CO)4], but it was difficult to detect the hydrogen atoms (now use neutron diffraction).

Metal-hydride bond is covalent, directional and the majority of M-H complexes show protic and/or hydridic character.

1. SYNTHESIS (a) By protonation - + Co(CO)4 + H → HCo(CO)4 - + HFe (CO)4 + H → H2Fe(CO)4

(b) By reduction

NaBH4 FeI2(CO)4  H2Fe(CO)4

(c) From dihydrogen

Mn2(CO)10 + H2 → 2 HMn(CO)5

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Co2(CO)8 + H2 → 2 HCo(CO)4

The carbonyl hydrides are volatile liquids. They are stable under CO and at low temperatures. They can decompose to metal carbonyls and liberate H2.

Most hydrides are strong acids and this is affected by – (i) Substitution e.g. HCo(CO)3(PPh3) pKa = 7.0 Lower = more acidic

(ii) Decreases down a group H2M(CO)4 pKa 11.8 (Fe) 18.7 (Ru) 20.8 (Os)

(iii) Increases from LHS to RHS of periodic table

The hydrogen can be bound in several ways

SPECTROSCOPY

(i) Infrared – Terminal hydrides absord in the region 2100-1800 cm-1.

The polarity differences in the M-H bond give different intensities.

Bridging hydrides absord in the region 1550-1000 cm-1

(ii) ‘H NMR – very useful. Metal hydrides resonate between 0 and -20 pp bridging hydrides even higher (-10 to -30 ppm)

This negative shift is very diagnostic. The nature of the metal gives extremely strong shielding

Complexes v(M-H) (cm-1) (MH) (PPM) [HMn(CO)5] 1780 -7.5 [HRe(CO)5] 1832 -5.66 [H2Fe(CO)4] 1887 -11.1 [H2Ru(CO)4] 1980 -7.62 [Cr2(U-H)(CO)11] ~1250 -19.5

REACTIVITY

19 Metal hydrides are highly reactive. They can insert alkenes and alkynes, are protonated and deprotonated.

(see figure 8)

METAL ALKENE COMPLEXES Zeise’s salt was first discovered in 1827 but its significance not realized until later.

Cl Cl K Pt

CL

Alkene (Olefin) complexes are known of every metal. The first row and early metal complexes are not as stable as the heavier metals.

1. SYNTHESIS (i) By ligand substitution hv F2(CO)5 + Fe (CO)4 -CO

+ O BF AgBF4 Cp Fe (CO) I + H C = CH 2 2 2 -AgI Fe

C O CO

Alkene are weak donors, but electron-rich early transition metals in low oxidation states favour alkenes as better ╥-acceptor

(ii) By addition to coordinatively unsaturated complexes

[IrCL (CO) (PPh3)2] + R2C = CR2 Cl 16 VE Ph P square planar 3 CR2 Ir CR Ph3P 2 Cl 18 Ve

(iii) By reduction

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Reduction with alkali metals or megnesium reagents in the presence of alkenes is a common route. Me2 P Nal H + 3 TiCl4 + Me2 P P Me2 Ti

P Me2

K O Cp Co + C H 2 2 4 -KCp

Co 18 VE

reactive complex with labile ethebe ligands

2. BONDING OF ALKENES Alkenes bind to transition metals via their --- orbitals, by donting electron density into an empty metal d-orbital. There is also donation of electron density from the metal to the olefin, into its --- orbital.

(DEWAR-CHATT-DUNCANSON MODEL)

21 H H

C II-donation

M

C T

H H H H

C back donation % into alkene M II II* orbitals

C

H H

OF Co, back donation into C=C antibonding --- levels leads to a weakening (lengthening) of the C=C bond, (and a reduction in the C=C bond stretching frequency.

vc  c (cm-1 ) r (c  c) A o Free C2H4 1623 1.335

K[P+Cl3(C2H4)] 1516 1.375

Back-bonding has two structural consequences:- (i) lengthening of the C=C bond (ii) reduction of angles around C from CO. 120o to tetrahedral angles (for Sp3-C)

The extent of back-bonding depends on;- (i) the energy of the orbitals (ii) steric effects (iii) the presence or absence of electron-withdrawing substituents.

22

The alkene substituents bend back away from the metal – this may be beneficial.

3. REACTIVITY OF ALKENES This has important industrial relevance

(See figure 9)

(i) ligand substitution

As relatively weak - donors and ---acceptors, alkenes are easily displaced by ligands, such as phosphines.

Ni (COD)2 + 4R3P → Ni(PR3)4

(ii) reactions with electrophiles

With electrophiles e.g. HX, there is reduction loss of the alkene and formation of the LnMXy complex. C2H5 H2O (C5H5)2Ti

(C5H5)2Ti O Ti(C5H5)2

C2H2

(iii) reactions with necleophiles

The most important aspect of metal-alkene chemistry. Coordinated ethane reacts smoothly under ambient conditions with necleophiles, whereas free ethylene does not.

23 O

+ CH (COOMe)2 Fe OC C O O

Fe OC C CH (COOMe) O 2

Many catalytic reactions involve the transformation of the alkene by nucleophilic attack and it can occur in 2 ways;-

3. MULTIDECKER SANDWICH COMPLEXES

24

4. COMPLEXES WITH TILTED SANDWICH STRUCTURE

25

SANDWICH STRUCTURES:

Metallocenes (featuring cyclopentadienyl ligands)

In 1948, an orange compound with the formula FeC10H10 was formed. It was remarkably stable and the initial reaction was thought to be: H H

+ FeCl3 Fe MgBr

H

However, the product was correctly identified later as having a ‘sandwich’ or ‘double cone’ structure, with all five C-atoms of a cyclopentadienyl (CP) ligand interacting with the metal centre.

26

Fischer Wilkinson

(Noble prize in 1973)

The new compound, Bis-(cyclopentadienyl) iron was named “ferrocene”, due to its aromatic nature.

Many metal complexes are known and generally called ‘metallocenes’.

This discovery, bonding structure of ferrocene is a cornerstone of organometallic chemistry.

The ring ligands are huckel aromatic systems, and most transition metal sandwich complexes follow the 18-electron rule.

1. SYNTHESIS OF METALLOCENES

(A) METAL SALT + CYCLOPENTADIENYL REAGENT

Dicyclopentadiene has to be ‘cracked’ to give moniomeric C5H6. This is a weak acid (pKa 15) and it can be deprotonated by strong bases or by alkali metals.

THF MCL2 + 2 NaC3H5  (C3H5)2M M=V, Cr, Mn, Fe, Co

Ni(acac)2 + 2 C5H5MgBr → (C5H5)2M + MgCl2

2. METAL + CYCLOPENTADIENE

M + C5H6 → M(C5H5) + ½ H2 M=Li, Na, K 500o C M + 2 C5H6  (C5H5)2M + H2

3. METAL SALT + CYCLOPENTADIENE

Often an auxiliary base is required if the basicity of the salt anion is insufficient to deprotonate cyclopentadiene

- 2- Tl2SO4 + 2 C5H6 + 2 OH → 2 TLC5H5 + 2H2O + SO4

27

FeCl2 + 2C5H6 + 2 Et2NH → (C5H5)2Fe + 2 [Et2NH2]Cl

In other cases, a reducing agent is needed;-

3 EtOH 3 2+ RuCl3(H2O)x + 3 C5H6 + /2 Zn  (C5H5)2 Ru + C5H8 + /2 Zn

2. BONDING IN METALLOCENES The bonding can vary depending on where the central element is in the periodic table.

n=1 Alkali metals n+ - Ionic M (C5H5 )n n=2 Heavy alkanine Earth metals n=2,3 Lanthanides

Intermediate n=1, In, Tl N=2, Br, Mg, Sn, Pb, Mn, Zn, Cl, Hg

COVALENT molecular n=2 V, Cr, Fe, Co, Ni, Ru, Os Lattice n=3 Ti n=4 Ti, Zr, Nb, Ta, Mo, U, Th

See figure 10, for molecular orbital diagram of ferrocene (with a combination of two Cp ligands and iron)

As in Co, there is a similar bonding pattern

σ-Donor bonds [Cp)a1) → M (4s, 3pz)]

╥-Donor bonds [Cp(e1) → M (dxz, yz, Px, y)] 2 2 ╥-acceptor bonds [Cp(e2) ← M(dx -y , dxy)]

3. REACTIVITY OF FERROCENE

With 18 VE, ferrocene is the most stable member in the metallocene series. It sublimes readily, is not attacked by air or water but can be oxidized

SPECTROSCOPY

‘H NMR spectroscopy is very useful for characterizing metallocenes

(see figure 12)

Due to the different metals and different ligands there are characteristic shifts.

28

The η5 (pentahapto) ligands are firmly bound to the metal, and they rotate very rapidly around the metal-Cp axis, giving rise to typical sharp Cp singlets.

However other motions of the Cp ligand are possible;-

1. Cp ring rotation (ring whiz) 2. η1– Cp and σ- bond shift 3. η5- η1 – interchange 4. η3-η5 ring slippage

1 1 η -Cp binding is common with main group elements e.g. Hg (η -C5H5)2

1 5 5 1 1 both η and η ligands exist in [)η -C5H5) Fe(CO)2 (η -C5H5)]. The η -Cp ligand is fluxional and undergoes rapid 1,2 – sigmatropic shifts of the metal-carbon σ bond. So, above R. temp, all five carbons are in effect identical on the NMR time scale. At low temps, the singlet becomes a multiplet as the fluxionality is ‘frozen’ and the η1- bonding is asymmetric.

4 5 H 3 H 4 Fe 5 Fe 1 1 2 2 3 OC CO OC CO

3 H 2 Fe 4 5 1 OC CO

For Cp2W(CO)2, the NMR spectrum shift that both ligands are identical in solution (at room temperature) but the solid state structure shows that one Cp ligand is only three coordinate (η5 → η3 ring slippage).

29 CO CO W CO W -CO CO

5 5 n ,n n5,n3

Also common for indenyl ligands

M 5 M n 3 n

Reversibly, electrochemically or by oxidizing reagents such as iodine to give the blue ferrocenium + cation [Cp2Fe] .

The reactivity of ferrocene is vast.

See figure 11 for some reactions.

Ferrocene may be regarded as a fairly electron rich arene; in friedel-crafts acylations, it reacts Ca. 106 times faster than benzene. O O C C CH 3 CH3 CH3CoCL Fe AlCl3 Fe + Fe

Ch3 C O

Metallation Li Li

n=BuLi Fe TMEDA Fe + Fe

Li

30

(Very useful reactants)

31 METAL COMPLEXES OF CYCLIC POLYENES

Cyclic conjugated ligands CnHn+,o,- ar known to occur in four different classes of compounds;

1. SANDWICH COMPLEXES 5e 5e 5e 5e

7e 8e 9e 6e Cr Mn Fe Co

7e 6e 5e 4e

2. HALF-SANDWICH COMPLEXES

Mo Co Ni

OC CO OC CO C N O O

"piano chair" "Milking stool"

3. MULTIDECKER SANDWICH COMPLEXES

32

4. COMPLEXES WITH TILTED SANDWICH STRUCTURE

33 SANDWICH STRUCTURES:

METALLOCENES (FEATURING CYCLOPENTADIENYL LIGANDS)

In 1948, an orange compound with the formula FeC10H10 was formed. It was remarkably stable and the initial reaction was though to be:. H H

+ FeCl3 Fe MgBr

H

However, the product was correctly identified later as having a ‘sandwich’ or ‘double cone’ structure, with all five C-atoms of a cyclopentadienyl (CP) ligand interacting with the metal centre.

Fischer Wilkinson

(nobel prize in 1973)

The new compound, bis-(cyclopentadienyl iron was named “ferrocene”, due to its aromatic nature.

Many metal complexes are known and generally called ‘metallocenes’.

This discovery and bonding structure of ferrocene is a cornerstone of organometallic chemistry.

The ring ligands are Huckel aromatic systems, and most transition metal sandwich complexes follow the 18-electron rule.

1. SYNTHESIS OF METALLOCENES

A) METAL SALT + CYCLOPENTADIENYL REAGENT

Dicyclopentadiene has to be ‘cracked’ to give monomeric C5H6. This is a weak acid (pKa  15) and it can be deprotonated by strong bases or by alkali metals.

THF ML2 + 2 NaC5H5  (C5H5)2M M= V, Cr, Mn, Fe, Co

34

MBr2 + 2 C5H5MgBr → (C5H5)2M + 2MgBr2

2. METAL + CYCLOPENTADIENE

M + 2C5H6 → M(C5H5)2 + H2 M = Li, Na, K

500O C M + 2 C5H6  (C5H5)2 M + H2 M = Ru, Fe

3. METAL SALT + CYCLOPENTADIENE

Often an auxiliary base is required if the basicity of the salt anion is insufficient to deprotonate cyclopentadiene

- 2- TL2 SO4 + 2 C5H6 + 2OH → 2 TLC5H5 + 2 H2O + SO4

FeCL2 + 2C5H6 + 2 Et2NH → (C5H5)2Fe + 2 [Et2Nh2]CL

In other cases, a reducing agent is needed;-

3 EtOH 3 2+ RuCl3 (h2O)x + 3 C5H6 + /2 Zn  (C5H5)2Ru + C5H8 + /2 Zn

2. BONDING IN METALLOCENES

The bonding can vary depending on where the central element is in the periodic table.

n = 1 alkali metals n+ - Ionic M (C5H5 )n n = 2 heavy alkaline Earth metals n = 2,3 lanthanides

Intermediate n = 1, In, Tl n = 2, Br, Mg, Sn, Pb, Mn, Zn, Cd, Hg

Covalent molecular lattice n = 2 V, Cr, Fe, Cr, Ni, Ru, Os n =3 Tl n = 4 Ti, Zr, Nb, Ta, Mo, Cl, Th

See figure 10, for molecular orbital diagram of ferrocene (with a combination of two ligands and iron)

As in CO, there is a similar bonding pattern

σ- donor bonds [Cp(a1) → M (4s, 3p2)]

35 ╥- donor bonds [Cp(e1) → M (dxz, yz, Px, y)] 2 2 ╥ - acceptor bonds [Cp (e2) ← M (dx -y , dxy)]

3. REACTIVITY OF FERROCENE

With 18 VE, ferrocene is the most stable member in the metallocene series. It sublimes readily, is not attacked by air or water but can be oxidized.

Reversibly, electrochemically or by oxidizing reagents such as iodine to give the blue ferrocenium + cation [Cp2Fe] .

The reactivity of ferrocene is vast.

See figure 11 for some reactions.

Ferrocene may be regarded as a fairly electron rich arene; in Friedel-crafts acylations, it reacts Ca. 106 times faster than benzene. O O C C CH 3 CH3 CH3CoCL Fe AlCl3 Fe + Fe

Ch3 C O

Metallation Li Li

n=BuLi Fe TMEDA Fe + Fe

Li

Li Li

n=BuLi Fe TMEDA Fe + Fe

Li

Very useful reactants

36

SPECTROSCOPY ‘H NMR spectroscopy is very useful for characterizing metallocenes

(see figure 12)

Due to the different metals and different ligands there are characteristic shifts.

The η5 (pentahapto) ligands are firmly bounf to he metal, and they rotate very rapidly around the metal-Cp axis, giving rise to typical sharp Cp singlets.

However, other motions of the Cp ligand are possible;-

1. Cp ring rotation (ring whiz) 2. η1-Cp and σ- bond shift 3. η5-η1 interchange 4. η3-η5 ring slippage

1 1 η -Cp bonding is common with main group elements e.g. Hg(η -C5H5)2

1 5 5 1 1 Both η and η ligands exist in [(η -C5H5) Fe(CO)2 (η -C5H5)]. The η -Cp ligand is fluxional and undergoes rapid 1,2- sigmatropic shifts of the metal-carbon σ bond. So, above R temperature, all five carbons are in effect identical on the NMR time scale. At low temperatures, the singlet becomes a multiplet as the fluxionality is ‘frozen’ and the η1- bonding is asymmetric.

4 5 H 3 H 4 Fe 5 Fe 1 1 2 2 3 OC CO OC CO

3 H 2 Fe 4 5 1 OC CO

For Cp2W(CO)2, the NMR spectrum shows that both ligands are identical in solution (at room temperature) but the solid state structure shows that one Cp ligand is only three coordinate (η5→η3 ring slippage).

37 CO CO W CO W -CO CO

5 5 n ,n n5,n3

Also common for idenyl ligands

M 5 M n 3 n

38