LN7_stkk2322 Metal Carbonyl Revised

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Metal CO Complexes (18 Electrons Rule)

Metal Carbonyl !

Carbonyl – carbon monoxide , CO

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Metal carbonyls are frequent starting materials

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Metal carbonyls are known for most transition metals

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Metal carbonyls are reactive — labile — CO is easily substituted

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CO lost as gas is swept from system, pushes reaction forward

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Metal carbonyls are easily studied by IR and provide information on structure, symmetry, electron density changes at metal

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Carbonyl complexes The carbonyl carbonyl ligand forms a huge number of complexes with metal ions, most commonly in low oxidation states, where it binds to the metal through its C-donor, C-donor, as in the complexes below, below, where all a ll the metal ions are zero-valent:

[Ni(CO)4] Td

[Fe(CO)5] TBP (D3h)

[Cr(CO)6] Oh

Metal-Carbon bonding in carbonyl complexes The carbonyl ligand is a !-acid. -acid. This  This is an acid in the Lewis sense, where it receives receives electrons from the metal ion, and it is a ! –acid because this involves involves !-bonding. The !-bonding involves involv es overlap of the !* orbitals of the CO with d orbitals from the t2g set of the metal, and so is d!-p! bonding. The canonical structures involved in the !-acid nature of CO are:

M!C≡O A

-

+

M=C=O B

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Metal Carbonyls Bonding in Metal carbonyls understood with MOT

"-bonding: donation from M to LUMO on C

!-bonding: donation from HOMO on C

C atomic orbitals

CO molecular orbitals

O atomic orbitals

Carbonyl Complexes (CO) !

Bonding !

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Review of CO Molecular Orbitals !

HOMO resides mostly on C, !-donation

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LUMO resides mostly on C, #-acceptance

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Reinforce each other and provide strong bonding

Bonding of CO to a Metal

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Characteristics of CO complexes !

Infrared Spectroscopy ! !

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Free CO stretch $ = 2143 cm-1 Cr(CO)6 CO stretch $ = 2000 cm-1 because #-back donation from metal weakens the C%O bond by adding e - to antibonding #* orbital  Negative charge charge on complex complex further further weakens weakens CO bond: bond: [V(CO)6]- $ = 1858 cm-1 

[Mn(CO)6]+ $ = 2095 cm-1

d) Bridging CO further further weakened by by extra #-back donation (e- count = 1/M)

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X-Ray Crystallography !

Free CO bond length = 112.8 pm

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M—CO carbonyl bond length = 115 pm

Metal-Carbon bonding in carbonyl complexes The carbonyl ligand is a !-acid. This is an acid in the Lewis sense, where it receives electrons from the metal ion, and it is a ! –acid because this involves !-bonding. The !-bonding involves overlap of the !* orbitals of the CO with d orbitals from the t2g set of the metal, and and so is d!-p! bonding. The canonical structures involved in the !-acid nature nature of CO are: +

M!C≡O A

-

M= C = O B

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Metal-Carbon bonding in carbonyl complexes !

What stabilizes CO complexes is M→C ! –bonding. The The lower the formal charge on the metal ion, the more willing it i s to donate electrons to the ! –orbitals of the CO. CO. Thus, metal ions with higher formal charges, e.g. Fe(II) form CO complexes with much greater difficulty than do zero-valent metal ions such as Cr(O) and Ni(O), or negatively ne gatively charged charged metal ions such as V(I). The ! –ov  –overlap erlap is envisaged envisaged as involving involving d-orbitals of the metal and the !* –orbitals of the CO:

d"-p"* overlap d-orbital of metal metal

* orbitals orbitals of CO

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 orbitals  orbitals of CO

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M

IR spectra and Metal-Carbon bonding in carbonyl complexes The  "CO stretching frequency of the coordinated CO is very informative as to the nature of the bonding. Recall Recall that the stronger a bond gets, the higher its stretching frequency.. Thus, the more important frequency impor tant the M=C=O (C=O is a double bond) canonical structure, the lower the  "CO stretching frequency as compared to the M-C≡O structure (C≡O is a triple bond): (Note:  "CO for free CO is 2041 cm-1) [Ti(CO)6]2-  [V(CO)6]-  [Cr(CO)6] [Mn(CO)6]+  [Fe(CO)6]2+  "

CO

1748

1858

increasing M=C double bonding

1984

2094

2204 cm -1

decreasing M=C double bonding

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IR spectra and bridging versus terminal carbonyls Bridging CO groups can be regarded as having a double bond C=O group, as compared to a terminal C≡O, which is more like a triple bond: ~ triple bond

M-C≡O terminal carbonyl (~ 1850-2125 cm -1)

the C=O group in a bridging M ~ double bond carbonyl is more like the C=O in C=O a ketone, which typically has M υC=O = 1750 cm -1

bridging carbonyl (~1700-1860 cm-1)

One can thus use the CO stretching frequencies around 1700-2200 cm-1 to detect the presence presence of bridging CO groups.

IR spectrum and bridging versus terminal carbonyls in [Fe2(CO)9] O CO OC

OC

Fe

OC

C Fe

CO

C C

O

CO

O

terminal carbonyls

bridging carbonyls

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Organometallic Chemistry. Concepts.1 Predicting structure

18 electron Rule or Effective Atomic Number (EAN) Rule “a stable organometallic compound has 18 valence electrons at the metal” Examples

The 18-electron Rule !

Counting Electrons !

The octet rule governs organic and simple ionic compounds: s + 3p orbital

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The 18-electron rule governs organometallics (with many exceptions) !

s + 3p + 3d orbitals

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Donor ligands provide the electrons other than the delectrons

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Carbonyl complexes and the 18-electron rule One might wonder why in the above complexes Ni(0) has four C≡O groups attached atta ched to it, Fe(0) five C≡O, and Cr(0) six C≡O. A very simple rule allows us to predict the numbers of donor groups attached to metal ions in organometallic complexes, called the eighteen electron rule. The latter rule states that the sum of the d-electrons possessed by the metal plus those donated by the ligands (2 per C ≡O) must total eighteen: [Ni(CO)4] Ni(0)

=

4 x CO =

[Fe(CO [F e(CO))5]

[Cr(CO)6]

d10

Fe(0) =

d8

8

5 x CO

10

18 e

18e

Cr(0) =

d6

6 x CO = 12 18e

Formal oxidation states are all zero.

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The “Donor Pair” method of electron counting counting (Method A) !

Common organometallic ligands are assigned an electron count and charge

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The charge on ligands helps determine d-electron count of metal

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Add up all electrons from Metal d orbitals and ligands to find total ecount

(isonitrile or isocyanide) (oxo, sulfido)

(nitrido)

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Carbonyl complexes and the 18-electron rule To obey the 18-electron rule, many carbonyl complexes are anions or cations, as in:

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[Mn(CO)6]+  

[Fe(CO)4]2-

V(0 V( 0) = d5

Mn(0 Mn( 0) = d7

Fe( e(0 0) = d8

6 CO = 12e

6 CO = 12e

4 CO = 8e

[V(CO)6]

1-

= +1e

1+

= 18e Formal oxidation state = V(-I)

 = -1e

2-

= 18e Formal oxidation state = Mn(I)

 = 2e = 18e

Formal oxidation state = Fe(-II)

NOTE: In applying the 18-electron rule, r ule, metal ions are always considered to be zero-valent zero-valent,, not the formal oxidantion state

Metal-Metal bonding in Carbonyl complexes complexes A species such as [Mn(CO)5] would have only 17 e. The 18e rule can be obeyed by two such entities forming a Mn-Mn bond, where each Mn contributes one electron to the valence valence shell of the other Mn, giving the metal-metal bonded species [(CO)5MnMn(CO)5]. To To check on the t he 18e rule, we look at a t one metal at a time: Mn(0) = d7 5 C!O = 10 Mn-Mn = 1

Mn-Mn bond Mn

Mn

18 e [Mn2(CO)10]

+

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Bridging Carbonyls in carbonyl complexes bridging carbonyls

Co

Co

Co-Co bond

[Co2(CO)8]

Carbonyls may form bridges between two metals, where they donate one electron to each metal in working out the 18 electron rule. In [Co2(CO)8] at left each Co has three terminal CO’s, two bridging CO’s, and a Co-Co bond:

Co(0)

=

d9

3 CO’s

=

6

2 bridge CO’s

=

2

Co-Co bond

=

1 18 e

Bridging Carbonyls in Fe 2(CO)9 bridging carbonyls

Fe-Fe bond

[Fe2(CO)9]

Fe2(CO)9 has each Fe with three terminal CO’s, three bridging CO’s, and an Fe!Fe bond. The 18 electron rule holds for each Fe atom as:

Fe(0)

=

d8

3 CO’s

=

6

3 bridge CO’s

=

3

Fe-Fe bond

=

1 18 e

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Charged ligands and the eighteen-electron rule The formally charged ligands that are important in organometallic chemistry are mainly soft ligands such as Cl-, Br-, and I-, with CN- also occurring. Hydride (H-) is also very important. These mono-anionic ligands all contribute one electron for the 18-electron rule, as in the following examples:

[Mn(CO)5Cl]

[H2Fe(CO)4]

Mn(0)

=

d7

Fe(0)

5 CO’s

=

10

1 Cl

=

1

[HCo(CO)4] d8

Co((0) Co

4 CO’s =

8

4 CO’s = 8

2H

2

1H

=

=

18 e

18 e

= d9

= 1 18 e

Examples of Electron Counting !

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Cr(CO)6 !

Total charge on ligands = 0, so charge on Cr = 0, so Cr = d 6

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6 CO ligands x 2 electrons each = 12 electrons

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Total of 18 electrons

(&5-C5H5)Fe(CO)2Cl !

Total charge on ligands = 2 -, so Fe2+ = d6

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(&5-C5H5- = 6) + (2CO x 2 = 4) + (Cl- = 2) = 12 electrons Total of 18 electrons

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Charged complex: [Mn(CO) 6]+ ! !

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Total ligand charge = 0, so Mn + = d6 12 electrons from 6 CO ligands gives a total of 18 electrons

M—M Bond: (CO)5Mn—Mn(CO)5 !

Each bond between metals counts 1 electron per metal: Mn—Mn = 1 e-

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Total ligand charge = 0, so Mn 0 = d7 5 CO ligands per metal = 10 electrons for a total of 18 electrons per Mn

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Justification for and exceptions to the 18-electron Rule !

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MO Theory  predicts that 18 electrons fill  bonding orbitals

'o

This number is more stable than more (filling antibonding orbitals) or less

• When is the 18-electron rule most valid? a) With octahedral complexes of large large 'o.

 b) Ligands are good !-donors and good #-acceptors (CO)

• Exceptions to the 18-electron rule are common ! !

Weak Wea k field ligands with small 'o make filling eg* possible ( > 18e-) #-donor ligands can make t 2g antibonding ( < 18 e-)

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