Structure Bonding

Prepared by Dr. Kevin Shaughnessy, Spring 2011 The University of Alabama, Tuscaloosa, AL

Chapter 1: Introduction—Structure and Bonding

Types of Ligand Coordination

Three generic classes of ligands: L: a neutral electron pair donor (i.e., CO, PR3) X: an anionic electron pair donor (i.e., X-, H-) M-M: neutral 1 electron donor

Distinguishing L-type ligands from X-type ligands: To distinguish between X and L type ligands, consider removing ligand from metal center with the ligand taking the electron pair in the M-L bond. If the ligand in it's free state is neutral, it is an L-type ligand. If it would be anionic, it is X-type.

Ligands that can donate more than one pair of electrons can be classified using L and X designations:

Types of ligand coordination:

Terminal: Ligand is bound to only one metal center (L-M or X-M)

Bridging (µ): Ligand is attached to two different metal centers (M-L-Mʼ or M-X-Mʼ). For L-type ligands, the lone pair is usually shared between the two metals (count 1 electron for each metal). For X ligands, the lone pair can also be shared. If the X ligand has additional lone pairs (i.e., halide, alkoxide), then the additional lone pairs can be used to coordinate to the second metal center.

Hapticity (η): Ligand attached to a metal center through more than one atom. Generally used to describe ligands with conjugated π-systems

N Pd PMe3

CH3

ClPMe3 L-typeNL-type

CH3 X-type

Cl— X-type

L24 e donor

O

NRLX

4 e donor

H3C

CH3

CH3

CH3H3CL2X

6 e donor

MM M!4 !5 !6

Chapter 1: Structure and Bonding: 2


Hapticity for ligands such as Cp can be variable depending on how it is coordinated. Usually Cp is a η5-L2X ligand, but it can also coordinate as an η3-LX, or even an η1-X. The process of going from η5 to η3 is often called ring slippage.

Chelation: Ligand attached through more than one atom usually separated by one or more atoms. Chelating ligands are sometimes classified as being bidentate (2 points of attachment), tridentate (three points of attachment), or tetradentate (4 points of attachment).

Kappa convention (κ): The kappa convention is sometimes used to indicate the coordinating atoms of a polydentate ligand. If only some of the possible coordination sites are bonded to the metal, the coordinating atoms are indicated with a kappa. Some authors use kappa to indicate how many of the coordinating atoms are attached in a polydentate ligand.

The 18-Electron Rule

Recall: First row elements have 4 valence orbitals (1 s + 3 p) so they can accommodate up to 8 valence electrons--the octet rule.

Transition metals have 9 valence orbitals (1 s + 3 p + 5 d). Upon bonding to a ligand set, there will be a total of 9 low lying orbitals (see MO theory discussion below). Therefore, we can expect that the low lying MOs can accommodate up to 18 valence electrons--The 18-Electron Rule. Organometallic complexes with 18 electrons are predicted to be particularly stable because they will have a closed shell of electrons. Complexes with 18 electrons are often referred to as being coordinatively saturated.

Counting electrons: There are two models for counting electrons. Both give the same answer, but offer different advantages and disadvantages.

Example: CH4

Covalent model: Since C-H bonds are covalent, assume that the electrons are shared equally between carbon and hydrogen. To count the electrons, we dissect the molecule giving each atom 1 electron of the bonding pair.

Pt

Me2N

ClNMe2

Cl

!2N-coordinationbidentate ligand

Fe

P

SS

S

Cl

!P, !3S-coordinationtetradentate ligand

Pt

NH

NH2

NHCl NH2

!3-[N,N'-di(2-aminoethyl)ethane 1,2-diamine]chloroplatinum ion

or[N-(2-amino-!N-ethyl)-N'-(2-

aminoethyl)ethane 1,2-diamine-!2N,N']chloroplatinum ion



Ionic model: Alternatively, we can treat the bonds as being ionic. This allows us to assign a formal oxidation state to the carbon atom. This can be useful to determine whether a particular transformation is an oxidation or a reduction. In this model, both electrons are given to the atom with the higher electronegativity. For a C-H bond, this is the carbon.

Similarly for a transition metal complex, the electron count is the sum of the metal valence electrons + the ligand centered electrons.

Covalent Model: # e = # metal electrons + # ligand electrons - complex charge

Metal: The number of metal electrons equals it's row number (i.e., Ti = 4e, Cr = 6 e, Ni = 10 e) Ligands: In general L donates 2 electrons, X donates 1 electron. See the table below.

Ionic Model: # e = # metal electrons (dn) + # ligand electrons

Metal: You must first determine the formal oxidation state of the metal. The number of electrons is the row number minus the charge on the metal (i.e., Ti(4+): 4 - 4 = 0 e, d0; Pd (2+): 10 - 2 = 8 e, d8). The formal oxidation state is simply the charge on the complex minus the charges of the ligands. Ligands: In general and L and X are both 2 e donors. See the table below.

The covalent model is used by Crabtree in his book. It may be more realistic in most cases, since M-L bonds are usually covalent, but does not indicate the formal oxidation state.

In my opinion, the ionic model is easier and gives a clearer picture of the actual chemistry, since the formal oxidation state is part of the calculation. All discussions in this class will use the ionic model, so I would encourage you to learn that one. You should also be aware of the covalent method, since you will encounter it from time to time.

Examples:

HMn(CO)5

Covalent Model Ionic Model

•H = 1 e 5 X CO = 10 e

Mn = 7 e Total = 18 e

H- = 2 e 5 X CO = 10 e Mn(I) d6 = 6 e

Total = 18 e

HCH HH

H

CH H

H

H: 4 X 1 e = 4C: 4eTotal = 8 electrons

HCH HH

H

CH H

H

4-H+: 4 X 0 e = 0C(-4): 8 eTotal = 8 electrons

MnOC

OC CO

CO

H

CO



[CpFe(CO)]-


Cp• = 2 X CO =

Fe = Charge (-1) =

Total =

Cp- = 2 X CO =

Fe(0) d8 = Total =

[CpNi(µ-CO)]2


Cp• = 2 X 1/2 CO =

Ni-Ni = Ni =

Total =

Cp- = 2 X 1/2 CO =

Ni-Ni = Ni(I) d9 =

Total =

Ligand Name Covalent Model lonic Model Electron Count Charge Electron

Count

Nitrosyl (linear, NO+), terminal or bridging) 3 +1 2

Carbene (CYR, where Y = substituent with π interaction with carbene, i.e., OR, NR2, Ph, X

2 0 2

CO, CNR (bridging or terminal) 2 0 2

PR3, AsR3, SbR3, NR3, imines, nitriles, ethers, sulfides, etc.

2 0 2

Terminal dinitrogren (N2) 2 0 2

η2(π)-Alkene

2 0 2

η2(π)-Alkyne 2 0 2

Fe

COOC

Cp Ni

OC

NiCO

Cp

MY

R

M

M



η2(π)-Carbonyl

2 0 2

η2(σ)-Dihydrogen 2 0 2

µ-η2-Alkyne 4 0 4

η4-acyclic diene

4 0 4

η6-Arene

6 0 6

Hydride (H-) 1 -1 2

Terminal or bridging Alkyl (-CR3) 1 -1 2

η1-Acyl

1 -1 2

η1-Aryl, alkenyl, or alkynyl 1 -1 2

ER3 or EX3 (E = Si, Ge, Sn) 1 -1 2

Alkoxide (-OR), thiolate (-SR), amide (-NR2), or phosphide (-PR2)

1 -1 2

Superoxide 1 -1 2

Halide (F-, Cl-, Br-, I-) or pseudohalide (-CN, -OTs, etc) 1 -1 2

Bridging alkoxide, thiolate, amide, or phosphide 3 -1 4

Bridging Halide (µ-X) M-X-M

3 -1 4

η2-Acyl 3 -1 4

η2-Alkenyl (terminal or bridging)

3 -1 4

η3-Allyl 3 -1 4

η5-Cyclopentadienyl (Cp-)

5 -1 6

Carbene (CR2 where R = substituent with no π 2 -2 4

OM

H HM

M M

M

M

M

O

M O O

MO

M MM

Terminal Bridging

M

M



interactions with the carbene carbon)

µ-CYR or CR2

2 -2 4

Imide (M=NR) 2 -2 4

Oxide (M=O) 2 -2 4

Peroxide (terminal or bridging) 2 -2 4

Alkylidine or carbyne, terminal 3 -3 6

µ-Alkylidine

3 -3 6

Nitride 3 -3 6

Application of the 18-electron rule:

The 18-electron rule can be used as predictor for the number of ligands a particular metal will coordinate.

V(CO)6 Cr(CO)6 (CO)5Mn-Mn(CO5 Fe(CO)5 (CO)3Co(µ-CO)2Co(CO)3 Ni(CO)4

d5

17 e d6

18 e d7

18 e d8

18 e d9

18 e d10

18 e The 18 electron rule works best for low-valent metals with small ligands that are strong σ donors and/or π acceptors (i.e., H- and CO). These ligands give large Δ, thus there is a strong preference for filling the dπ orbitals (requiring 18 electrons), and are small enough to allow the metal to be coordinatively saturated.

For complexes that follow the 18 electron rule, it can be used to predict reactivity as we will see throughout the semester.

Note, that just like the octet rule, the 18-electron rule is not an absolute requirement. There are many exceptions.

Common exceptions to the 18 electron rule:

d8 metals: The d8 metals (groups 8 - 11) have a tendency to form square-planar 16 electron complexes. This tendency is weakest for group 8 (Fe(0), Ru(0), and Os(0)) and is very strong for groups 10 and 11 (Pd(II), Au(III)). Square planar, 16 electron complexes of of d8 metals results in completely filled orbitals except the high energy dx 2 − y 2 (see MO discussion below).

MR

RR

CR M

M

O OM

M C RR

CM M

M

M N

Me Pd Me

PMe3

PMe3

Me3P Rh Cl

CO

CO

Me3P Au CH3

CH3

CH3



d0 metals: The high-valent d0 complexes often have lower electron counts than 18.

Bulky Ligands: Sterically demanding ligands will often result in lower than expected electron counts.

> 18 electron complexes: Complexes with formally 19 or 20 electrons are known, but they are usually unstable, or adopt alternate configurations.

Properties of Transition Metals

Ionization potential:

Row trends: General trend that IP increases (harder to oxidize) moving to right across transition series.

Column trends: Trends across the series are not consistent. For middle and late transition metals generally third row elements best able to form higher oxidation states (≥+3) than second, which is greater than first row elements. High oxidation states are rare for first row elements, but more accessable for 2nd or 3rd row elements.

Bond Stengths: Metal-ligand bond strengths tend to increase moving down a given column

Formal Oxidation State and d electron configuration

Oxidation states in organometallic complexes are merely formalisms that may bear little resemblance to the actual positive charge on the metal.

"Formalisms are convenient fictions which contain a piece of the truth--and it is so sad that people spend a lot of time arguing about the deductions they draw, often ingeniously and artfully, from formalisms, without worrying about their underlying assumptions." (Roald Hoffman, JACS, 1984, 106, 2006)

Consider the range of possible formal charges on metal centers: [Fe(CO)4]2- (Collman's reagent): Fe(-II), although the metal likely has little if any negative charge [ReH9]2-: Re(VII), although it is made by reduction of ReO4

- (also Re(VII)) with sodium in ethanol

Ta3 X Np- = 6carbene = 4Ta(V) d0 = 0Total = 10 e

ZrCH3

CH3

2 X Cp- = 12 e2 X CH3

- = 4 eZr(IV) d4 = 0 eTotal = 16 e

t-BuPPdt-But-Bu

t-BuP

t-But-Bu

Co Ni

20 e19 e

- eCo

18 e

Ni

18 e

E˚ = -.94 V vs SCE



Therefore, formal oxidations and reductions do not necessarily result in a decrease or increase in electron density at the metal center. Conversely, reactions where the oxidation state does not change can greatly affect the electron density of the metal center.

For π-coordinated ligands (i.e., alkenes, dienes, etc.) a number of different formal oxidation states can be determined. In the butadiene complex below, the ligand can either be considered an L2 ligand donating 4 electrons, an LX2 donating 6 electrons, or an X4 ligand donating 8 electrons. Therefore, the iron center can be Fe(0), Fe(II), or Fe(IV). In each case the complex has 18 electrons.

The d-electron count is a closely related formalism. The d-electron count is obtained by subtracting the charge on the metal from the row number. Although a formalism, the d-electron count can be used to predict the structure of organometallic complexes in some cases.

Coordination Number

Geometry Preferred dn Example

2 Linear 3

T-Shaped

d8

3

Trigonal

d0, d5, d10

4

Tetrahedral

d0, d5, d10

Fe(CO)42- H+HFe(CO)4-

Fe(-II), d1018 e

Fe(0), d818 e

Oxidation by protonation

(Cy3P)2IrH5

Ir(V), d418 e

H+Ir

H

Cy3P H

PCy3

H H

H H

2 H2 = 4 e2 H- = 4 e2 PCy3 = 4 eIr(III) d6 = 6 eTotal = 18 e

Reduction by protonation

FeOC CO

CO

FeOC CO

CO

FeOC CO

COFe(0), d8 Fe(II), d6 Fe(IV), d4

Ph3P Au Cl

Pt-Bu3

PdPh

Br

PPh3

PdPh3P PPh3

PPh3

PdPh3P PPh3

PPh3



4

Square Planar

d8

5

Trigonal Bipyramidyl

d8, d6 (distorted)

6

Octahedral

d0, d3, d5, d6

8

Dodecahedral

d1

9

Tricapped Trigonal Prism

d0

To explain the bonding of ligands to metals, we have to consider the molecular orbitals involved metal-ligand bonding. Before the development of molecular orbital theory, the interaction of ligands with metals was described using crystal field theory. We'll discuss this briefly and then consider the true orbital interactions.

Crystal Field Theory

Assumption: Ligands act as points of negative charge surrounding the metal center

For an isolated metal ion, the 5 d orbitals are degenerate. As an octahedral set of ligands approach, the d orbitals pointing along the x, y, and z axes (dz 2 , dx 2 − y 2 ) are destabilized. The other orbitals (dxy, dyz, and dxz) are less destabilized.

PPh3

Pd

PPh3

H3C Cl

OC FeCOCO

CO

CO

MoOCOC CO

CO

CO

CO

Mo

PMe3H

H PMe3

Me3P H

HMe3P

H Re HH

H

H

H

HH

H2-



eg orbitals (dσ) can form σ bonds with the ligands t2g orbitals (dπ) can form π bonds with the ligands

Crystal field splitting of other common geometries:

dz2 dx2-y2

dxy dxz dyz

eg

t2g

!

eg

t2g

Mn+

ML6n+



High spin and Low spin complexes

Co(III) has the electron configuration [Ar]3d44s2 as a free atom. Upon coordination, the 4d orbitals are stabilized relative to the 3s and the metal takes on a [Ar]3d64s0 electronic configuration. This configuration is usually shortened to d6.

A d6 electronic configuration would be expected to strongly favor an octahedral ligand arrangement, since this would result in the a completely filled set of the low-lying dπ (t2g) orbitals. This electronic configuration is called the low-spin form, since all electron spins are paired.

If Δ is small enough, though, a high-spin form with a t2g4eg

2 configuration is possible. In this configuration there are 4 unpaired electrons.

dxy dyz dxz

dx2 dx2-y2

dxz dxz

dz2

dxy

dx2-y2

TetrahedralSquarePlanar

Δ

dxy, dyz, dxz

dx2-y2 and dz2 dyz and dxz

dxy

dz2

dx2-y2

Tetrahedral Square Planar

Δ



The size of the ligand field splitting (Δ) is dependent on the ligands as well as the metal center. – 2nd and 3rd row transition metals tend to have higher Δ – Higher oxidation state metals have higher Δ – π-acceptor and strong σ-donor ligands give high Δ, while π-donor ligands give low Δ

Ligand Field Theory

Crystal Field Theory is qualitative. To get a better understanding we can turn to a MO picture of bonding in coordination complexes.

For an octahedral complex of pure σ donor ligands, the s, three p, and five d orbitals of the metal interact with the lone pair orbitals of the ligand. – 6 of the metal orbitals (s, three p, and 2 dσ) are of the appropriate symmetry to interact with the ligand orbital set. – Therefore, a set of 6 bonding M-L σ MOʼs and 6 anti-bonding M-L σ * MOʼs are formed. The three dπ do not find a symmetry match and remain as non-bonding orbitals. – As a result, an octahedral complex can accommodate 18 electrons in the bonding and non-bonding orbitals. We will see that this maximum is similar to the octet rule in the first row elements.

An octahedral d6 metal complex will have 18 valence electrons. The 6 ligands will each contribute 2 electrons (12 total) and the metal will contribute 6. These 18 electrons will fill all of the bonding orbitals (6 bonds) plus the three non-bonding dP orbitals (assuming a low spin complex). Note that in a high-spin complex the filled e2g orbitals are actually M-L antibonding orbitals.

eg

t2g

!

t2g

eg!

Low-spin High-spin

I- < Br- < Cl- < F- < H2O < NH3 < PPh3 < CO, H < SnCl3-

π-donor π-acceptor/strong !-donor

Increasing "



Types of Ligands

σ-donors: Typical M-L bonding arrangement. Lone pair orbital of ligand interacts with metal orbitals.

π-acceptors:

Ligands such as CO have empty MOʼs of proper symmetry to overlap with the filled dπ orbitals of the metal center. CO is an example of a π-acceptor ligand (often called a π-acid). In the case of CO, the π* MO of CO is the orbital that interacts with the metal dπ.

In the MO description above, the CO π* orbitals will interact with the Mdπ orbitals. This interaction stabilizes the filled Mdπ orbitals. As a result of this bonding interaction, the metal is donating electron density back to the ligands in a process called back bonding. – π acceptors form very strong M-L bonds due to this back bonding – π-acceptors increase the ligand splitting by stabilizing the t2g set of orbitals increasing Δ - back bonding allows electron density to be donated back to the ligand. This allows low-valent metals with filled d orbitals to form complexes with ligands

!

t2g

eg

dπ

M-L "*

M-L "

Mn+

6 ligandlone pairs

ML6n+

Deg

Deg



Other π-acceptor ligands: NO, N2, CN-R, PF3

π-donors:

Ligands with lone pairs on the coordinating atom (i.e., -OR, -X, -NR2) can act as π-donor ligands. The atom can rehybridize to place the lone pairs in orbitals with π symmetry that can overlap with the metal dπ orbitals.

As a result, the number of electrons that can be donated by these ligands will depend on how electron deficient the metal is. For example, oxygen could donate 2, 4, or 6 electrons. This concept will be further developed in the next chapter of notes.

C OM

dππ*

M-L !

M-L "

!

dπ

π*

!



For a metal with filled dπ orbitals, π-donation will result in a destabilization of these orbitals due to repulsion between the lone pair electrons and the dπ electrons. This results in a smaller Δ and a weaker M-L bond.

For d0 metals, such as W6+, π-donation is a favorable interaction and leads to stronger M-L bonds.

O RM

d! p

M-L "

M-L !

!

dπ

Ligandlone pairs

!

π



π-complexes:

Unsaturated molecules can donate π electrons to form M-L bonds rather than non-bonding lone pairs. Ethylene is a classic example.

The π orbital of ethylene acts as a σ donor. The π* orbital of ethylene is of the proper symmetry to act as a π-acceptor (back bonding).

Any compound with a π bond can potentially form a π complex with a metal center (alkynes, arenes, ketones, etc.)

σ-complexes:

Dihydrogen can bind to metal centers as an intact molecule. H2 has neither lone pairs nor π electrons. The electron pair that interacts with the metal center is the σ-bonding electrons of the H-H bond. The σ∗-orbital can act as a π-acceptor allowing back bonding to occur.

Other σ-complexes that have been identified include: C-H, Si-H, Sn-H, P-H, S-H, B-H, M-H

HH

HH

Cl PtCl

Cl

Zeise's Salt--1825Organometallics, 2001, 20, 2-6

M

dπ

C

C

H

HH

H

π

π*M-L !

M-L "

M

dπ

H

H

!

!"

M-L #

M-L !

Structure Bonding

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