ln05-1 electronic structure / bonding in d-block...

LN05-1 Electronic structure / bonding in d-block complexesMany, many properties of transition metal complexes (coordination number, structure, colour, magnetism, reactivity) are very sensitive to the number of d-electrons and how they are arranged in the d-orbitals .

For the transition element valence orbitals the energy-ordering is: ns < (n-1)d < np

Why? Recall: orbital energies affected by principal quantum number (n), effective nuclear charge experienced by electrons (Zeff) and e–-e– repulsions as subshells are filled.

Removal of one or more electrons (oxidation) reduces overall e– repulsion and lowers energy; this effect is most pronounced for d-orbital energies, relative to s or p.

H&S 19.2

BUT For higher oxidation states, Mn+, the energies of (n-1)d orbitals tend to be lower in energy than the ns orbitals.

e.g. Element free atom configuration configuration in complexesScandium 1s22s22p63s23p64s23d1 1s22s22p63s23p63d3

Iron 1s22s22p63s23p64s23d6 1s22s22p63s23p63d8

LN05-2 Oxidation states and d-electron countsTransition metal ions (mostly) have no s-e–, only d-e–, in their valence shell……so, by convention, we discuss the electron configurations of metal ions as “dn” where “n” is the number of d-electrons in the valence shell of the T.M. ion.

Determining the d-electron count in a transition metal complex

1. Determine the oxidation state:

oxidation state = - (Σ ligand charge) + (charge on inner sphere complex)

Once the formal charge of the metal is determined, the d-electron count can be determined based on the group that the metal belongs to:

for dn, n = (group number) – (oxidation state of metal)

H&S 20.

LN05-3 Oxidation states and d-electron counts

Examples:Rh(PPh3)3Cl:

K3[Fe(CN)6]

H&S 20.

LN05-4 Oxidation states

(Figure 19.3, H&S) Blue = common oxidation state. [ ] = rare oxidation state.Blank = non-existent

Group 3 4 5 6 7 8 9 10 11 12

LN05-5 Oxidation states – general comments

1. Most metals can adopt more than 1 oxidation state. Exceptions are groups 3 and 12

2. Within a group, the heavier metals tend to (slightly) favour higher oxidation states.

3. In the early to mid part of the d-block (up to Group 8), oxidation states range up to and including the maximum possible oxidation state.

4. From Group 9 onwards, very high oxidation states become less favoured. For most metals in this part the preferred oxidation states are maximum +3 (and occasionally +4 for heavier metals)

LN05-6 The Electroneutrality principle H&S 19.6

Formal oxidation states are mainly useful for “bookkeeping” purposes and should not be interpreted as a representation of the ‘real’ charge on a metal ion.

Purely covalentmodel

Purely ionic model

Partly ionic and covalent (polar

covalent):electroneutrality

principle

old-fashioned way of representing Lewis acid-

base interaction

LN05-7 Bonding in transition metal complexesVarious models have been used to discuss structure and bonding in relation to experimentallydetermined properties (structure, colour, magnetism).

• Valence bond theory (H&S 20.2) - not really used anymore; included for “historical reasons”

•Crystal field theory (H&S 20.3) - Conceptually simple and has some predictive power; but,based on dubious premise that there are NO covalent interactions between a metal and itsligands

•Molecular orbital theory (H&S 20.4) - The most complex, but most complete model. We willsimply (but thoroughly) analyze them and focus on particular part of them (i.e. MOs that arebased on the metal d orbitals)

•Ligand field theory (H&S 20.5) - Related to crystal field theory but with additionalparameters (“fudge factors”) included to represent metal-ligand covalent bonding

H&S 20.4

LN05-8 Atomic orbitals

3D representations of common atomic orbitals. Know these!

H&S 1.6

3dyz 3dxy 3dxz

3dx2-y23dz2

LN05-9 Molecular orbital diagram for octahedral complexesH&S 20.4

LN05-10 Octahedral MOs: a closer look

t2g

eg*

Non-bonding orbitals

M-L antibonding (σ*) orbitals

H&S 20.4

LN05-11 MO theory: tetrahedral complexes H&S 20.4

x

y

z

LN05-12 MO theory: square planar complexes

x

y

z

H&S 20.4

LN05-13 Qualitative comparison of MO diagrams

LN05-14 The 18 electron “rule”

LN05-15 The 18 electron “rule”

Stable transition metal complexes tend to have a total of 18 valence electrons associated with the transition metal.

H&S 20.4 & 24.3

Exceptions:

1. complexes of early transition metals (not enough d-electrons)

2. Late transition metals (too many d-electrons), particularly with weak sigma donors (more on this in a bit)

3. Square planar complexes – these obey the 16 electron rule

Note that we use the “ionic” model of M-L bonding, where formally anionic ligands are two-electron donors (this is the model where we emphasize polar nature of the M-L bond). There is an alternative model for counting (mentioned in 24.3) in which M-L bonds are broken homolytically. We don’t use this (but it’s valid).

LN05-16 Pi bonding in Octahedral complexes H&S 20.4

a1g + eg + t1u a1g + eg + t1u

t2g

As before: L is a σ donor L is a π donor(and still a σ donor)

LN05-17 Pi bonding in Octahedral complexes

Recall that, in the absence of π orbitals on the ligands, t2g molecular orbitals are non-bonding and are just the dxy, dxz, and dyz orbitals from M.

When the L have π donor orbitals, the t2g become t2g* - i.e. M-L (π) antibonding:

H&S 20.4

t2g (no ligand π orbitals) t2g* (ligand π donor orbitals)


a1g + eg + t1u

a1g + eg + t1u

t2g

As before: L is a σ donor L is a π acceptor (and still a σ donor)


Recall that, in the absence of π orbitals on the ligands, t2g orbitals are non-bonding and are just the dxy, dxz, and dyz orbitals from M.

When the L have π donor orbitals, the t2g become t2g* - i.e. M-L (π) antibonding:

t2g (no ligand π orbitals) t2g (ligand π acceptor orbitals)

A ligand group π-acceptor orbital:

LN05-20 Bonding in organometallic pi complexes

So far we have discussed ligands which can be pi donors, pi acceptors, or neither – but all (so far) are sigma donors in which the donor orbital is a lone pair on the ligand:

BUT we have also seen a few ligands that don’t appear to have a sigma donor orbital:

How do we understand metal-ligand bonding in “pi complexes”?What is the correct way to think about these ligands in terms of

the metal’s coordination number?the metal’s valence electron count?

2-ethylene 6-benzene 5-cyclopentadienide


Metal-ligand bonding in ethylene complexes: the Dewar/Chatt/Duncanson model

HOMO

LUMO

1. sigma-type interaction involving ethylene HOMO. Ethylene is using is π bond (HOMO) as a σdonor


2. π-type interaction involving backbonding from a metal d-orbital to the ethylene LUMO.

Metal-ligand bonding in ethylene complexes: the Dewar/Chatt/Duncanson model

HOMO

LUMO

H&S Fig 24.5


There is a sigma donating part and a pi backbonding part to the bonding in metal-ethylene complexes.

For the purposes of electron counting, we focus mainly on the sigma donating part: the pair of electrons in the “sigma bond” come from ethylene.

Therefore ethylene contributes two electrons as a neutral ligand.

In general, a two-carbon unit bonded to a metal in a pi complex is described as occupying one coordination site and contributes two electrons to the complex

Ligand number of e- charge number of coord. sites2-ethylene 2 0 14-butadiene 4 0 26-benzene 6 0 34-benzene 4 0 25-cyclopentadienide 6 -1 3


M-L bonding in benzene and Cp complexes is more complicated – these rings have more π molecular orbitals.

For these ligands we can focus on the electronic and structural similarities to traditional σ donor ligands: both benzene and Cp- are best described as occupying three mutually cis L type positions:


Bis(6-benzene)chromium(0)

Bis(5-cyclopentiadienyl)iron(II)“ferrocene”

LN05-26 Representing pi complexes

Pseudo-6-coordinate

(NOT 2-coordinate)

12-coordinateNO

6-coordinateOK but harder to draw

Pseudo-6-coordinate

(NOT 2-coordinate)

10-coordinateNO

6-coordinateOK but harder to draw

LN05-27 Spectrochemical series

The spectrochemical series is an empirical ordering of ligands in terms of their effect on ΔO:

I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < bipy < phen < CN- < CO

The parameter ΔO is an important determinant of the overall electronic structure, and hence properties (colour, magnetism, reactivity) of transition metal complexes.

H&S 20.4

LN05-28 High spin vs low spin: octahedral H&S 20.1

Orbitals are populated based on Hunds rule

d0-d3

d4-d7

d8-d10

t2g

eg*

t2g

eg*

t2g

eg*

t2g

eg*

LOW SPIN HIGH SPIN

LN05-29 High spin vs low spin: octahedral

The spetrochemical series allows for general predictions of preferences for HS vs LS configurations, particularly when comparing two complexes with different ligands.

Generally the spectrochemical series cannot be used to absolutely determine the HS/LS preference of one compound, except for ligands at either end of the series.

LN05-30 High spin vs low spin: other geometries H&S 20.1

In general Δt tends to be about half the magnitude of Δo (assuming equivalent ligands etc.)

As a result, tetrahedral complexes are nearly always high spin when there is a choice (d3-d6)

Within these for orbitals, use “high spin” confiugurations……..

Only populate this orbital if necessary (> 8 d electrons)

LN05-31 High spin vs low spin: other geometries H&S 20.1

Within these for orbitals, use “high spin” confiugurations……..

Only populate this orbital if necessary (> 8 d electrons)

LN05-32 Jahn-Teller distortions H&S 20.3

Sometimes the structures of “octahedral” complexes deviate in subtle but important ways:

regular octahedron:All Ni-O bonds 2.07 Å

distorted octahedron:Red Ni-O bonds 1.95 ÅBlue Ni-O bonds 2.38 Å

Ni(II) = d8 Cu(II) = d9


The Jahn-Teller theorem: Electronically degenerate states are susceptible to structural changes which remove the degeneracy

or

Cu(II) = d9

Regular octahedron Distorted octahedron


Distortions can be elongation of z-axis, or compression:


Distortions can be predicted for the following d-electron counts in octahedral complexes:

d1 d2 d4 (HS and LS) d5 (LS) d6 (HS) d7 (HS and LS) d9

Consequences of Jahn-Teller distortions:1. Structures2. spectroscopy (see later)3. Reactivity (see later)

ln05-1 electronic structure / bonding in d-block...

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