coordination chemistry

CRYSTAL FIELD THEORY In crystal field theory, the electron pairs on the ligands are viewed as point negative charges that interact with the d orbitals on the central metal.. CRYSTAL FIE

COORDINAT ION

CHEMISTRY II

BONDING, INCLUDING CRYSTAL FIELD THEORY

AND LIGAND FIELD THEORY

BASIS FOR BONDING

THEORIES

Models for the bonding in transition metal complexes must be

consistent with observed behavior

Specific data used include stability (or

formation) constants, magnetic

susceptibility, and the electronic (UV/Vis) spectra of the complexes.

BONDING

APPROACHES

Valence Bond theory provides the hybridization for octahedral complexes For the first row transition metals, the

hybridization can be: d2sp3 (using the 3d, 4s and 4p orbitals), or sp3d2 (using the 4s, 4p

CRYSTAL FIELD

THEORY

In crystal field theory, the electron pairs on the ligands are viewed as point negative charges that interact with

the d orbitals on the central metal The

nature of the ligand and the tendency

toward covalent bonding is ignored.

D ORBITALS

CRYSTAL FIELD THEORY

Ligands, viewed as point charges, at the corners of

an octahedron affect the various d orbitals differently.

CRYSTAL FIELD THEORY

CRYSTAL FIELD THEORY

The repulsion between ligand lone pairs and the d orbitals on the metal

results in a splitting of the energy of the d orbitals.

Octahedral field

eg

t2g

D ORBITAL

SPLITTING

In some texts and articles, the gap in the d orbitals is assigned a value of 10Dq The upper (eg) set goes up by 6Dq, and the lower set (t2g) goes down by 4Dq.

The actual size of the gap varies with the metal and the ligands.

D ORBITAL

SPLITTING

The colors exhibited by most transition metal complexes arises from the

splitting of the d orbitals As electrons

transition from the lower t2g set to the eg

set, light in the visible range is absorbed.

D ORBITAL SPLITTING

The splitting due to the nature of the ligand can be observed and measured using a spectrophotometer Smaller values of ∆ o result in colors in the green range

Larger gaps shift the color to yellow

THE SPECTROCHEMICAL

SERIES

The complexes of cobalt (III) show the shift in color due to the

ligand

(a) CN–, (b) NO2–, (c) phen, (d) en, (e)

NH3, (f) gly, (g) H2O, (h)

ox2–, (i) CO3 2–

LIGAND FIELD

STRENGTH

OBSERVATIONS

1 ∆o increases with increasing oxidation

number on the metal.

Mn+2<Ni+2<Co+2<Fe+2<V+2<Fe+3<Co+3

<Mn+4<Mo+3<Rh+3<Ru+3<Pd+4<Ir+3<Pt+4

2 ∆o increases with increases going down a group of metals.

LIGAND FIELD

THEORY

Crystal Field Theory completely ignores the nature of the ligand As a result, it

cannot explain the spectrochemical series.

Ligand Field Theory uses a molecular orbital approach Initially, the ligands can be

viewed as having a hybrid orbital or a p

orbital pointing toward the metal to make σ bonds.

OCTAHEDRAL

SYMMETRY

http://www.iumsc.indiana.edu/morphology/ symmetry/octahedral.html

LIGAND FIELD THEORY

LIGAND FIELD THEORY

The A 1g group orbitals have the same

symmetry as an s orbital on the

central metal.

LIGAND FIELD THEORY

The T 1u group orbitals have the same

symmetry as the p orbitals on

the central metal

(T representations are triply degenerate.)

LIGAND FIELD THEORY

The E g group orbitals have the same

symmetry as the d z 2 and d x 2 -y 2

orbitals on the central metal

(E representations are doubly degenerate.)

LIGAND FIELD THEORY

Since the ligands don’t have a combination with t 2g symmetry, the d xy , d yz and d xy

orbitals on the metal will be bonding when considering σ bonding.

non-LIGAND FIELD THEORY

The molecular orbital diagram is consistent with the crystal field

approach

Note that the

t2g set of orbitals is non-bonding, and the eg set of orbitals

is antibonding.

LIGAND FIELD THEORY

The electrons from the ligands (12 electrons from 6 ligands in

octahedral complexes) will fill the lower bonding orbitals.

{

LIGAND FIELD THEORY

The electrons from the 4s and 3d orbitals of the metal (in the first transition row) will occupy the middle

portion of the diagram.

{

EXPERIMENTAL

EVIDENCE FOR

SPLITTING Several tools are used to confirm the

splitting of the t2g and eg molecular orbitals

The broad range in colors of transition metal complexes arises from electronic

transitions as seen in the UV/visible spectra

of complexes.

Additional information is gained from measuring the magnetic moments of the

complexes.

EXPERIMENTAL EVIDENCE FOR SPLITTING Magnetic

susceptibility measurements can be

used to calculate the number of unpaired electrons in a compound.

Paramagnetic substances are attracted to a magnetic field.

MAGNETIC MOMENTS

A magnetic balance can be used to

determine the magnetic moment of a

substance If a substance has unpaired

electrons, it is paramagnetic, and attracted to a

MAGNETIC MOMENTS

Complexes with 4-7 electrons in

the d orbitals have two possibilities for the

distribution of electrons The complexes

can be low spin, in which the electrons

occupy the lower t2g set and pair up, or they

can be high spin In these complexes, the

electrons will fill the upper eg set before

pairing.

HIGH AND LOW SPIN

COMPLEXES

If the gap between

the d orbitals is large, electrons

will pair up and fill the lower (t 2g ) set of orbitals before occupying the e g set of orbitals The

complexes are called low spin.

HIGH AND LOW SPIN

COMPLEXES

In low spin complexes, the size of ∆ o is greater than the pairing energy

of the electrons.

HIGH AND LOW SPIN

COMPLEXES

If the gap between

the d orbitals is small, electrons

will occupy the e g set of orbitals before they pair up and fill the lower (t 2g ) set of orbitals before.

The complexes are called

high spin.

HIGH AND LOW SPIN

COMPLEXES

In high spin complexes, the size of ∆ o is less than the pairing energy of the electrons.

LIGAND FIELD STABILIZATION ENERGY

The first row transition metals in water are all weak field, high spin cases.

do d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

EXPERIMENTAL

EVIDENCE FOR LFSE

The hydration energies of the first row

transition metals should increase across the period as the size of the metal ion gets smaller.

M2+ + 6 H2O(l)  M(H2O)62+

EXPERIMENTAL EVIDENCE FOR LFSE

The heats of hydration show two “humps”

consistent with the expected LFSE for the metal ions The values for d 5 and d 10 are the same

as expected with a LFSE equal to 0.

EXPERIMENTAL

EVIDENCE OF LFSE

do d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

HIGH SPIN VS LOW

SPIN

3d metals are generally high spin

complexes except with very strong ligands

CN- forms low spin complexes, especially with M3+ ions.

4d & 4d metals generally have a larger

value of ∆o than for 3d metals As a result, complexes are typically low spin

NATURE OF THE

LIGANDS

Crystal field theory and ligand field theory differ in that LFT considers the nature of the ligands Thus far, we have

only viewed the ligands as electron pairs used for making σ bonds with the metal Many ligands can also form π bonds with the metal Group theory greatly simplifies the construction of molecular orbital

diagrams.

CONSIDERING Π

BONDING

To obtain Γred for π bonding, a set of cartesian coordinates is established for each of the ligands The direction of the

σ bonds is arbitrarily set as the y axis (or

the py orbitals) The px and pz orbitals are used in π bonding.

x x x

z

z

z z

z

z

Consider only the px and

pz orbitals on each of the ligands to obtain Γπ.

CONSIDERING Π BONDING

Τπ reduces to: T1g + T2g + T1u + T2u

The T1g and T2u group orbitals for the ligands don’t

match the symmetry of any of the metal orbitals.

The T1u set has the same symmetry as the px, py and

pz orbitals on the metal These orbitals are used

primarily to make the σ bonds to the ligands.

The T2g set has the same symmetry as the dxy, dyz and

dxz orbitals on the metal

Π BONDING

The main source of π bonding is between the d xy ,

d yz and d xz orbitals on the metal and the d, p or π* orbitals on the ligand.

Π BONDING

The ligand may have empty d or π* orbitals and

serve as a π acceptor ligand, or full p or d orbitals and serve as a

π donor ligand.

Π BONDING

The empty π antibonding orbital on CO can accept

electron density from a filled d orbital on the metal CO is a pi

acceptor ligand.

empty π* orbital

filled d

orbital

Π DONOR LIGANDS

(LM)

All ligands are σ donors

Ligands with filled p or d orbitals may also

serve as pi donor ligands Examples of π donor ligands are I-, Cl-, and S2- The filled p

or d orbitals on these ions interact with the

t2g set of orbitals (dxy, dyz and dxz) on the

metal to form bonding and antibonding

molecular orbitals.

Π DONOR LIGANDS (LM)

The size of ∆ o decreases, since it is now between an antibonding t 2g orbital and the e g *

orbital.

This is confirmed by the spectrochemical series Weak field ligands are also pi donor ligands.

Π ACCEPTOR LIGANDS (ML)

Ligands such as

CN, N 2 and CO have empty π

antibonding orbitals of the proper

symmetry and energy to interact

with filled d orbitals on the metal.

Π ACCEPTOR LIGANDS (ML)

The metal uses the

t 2g set of orbitals (d xy , d yz and d xz )

to engage in pi bonding with the

ligand The π* orbitals on the

ligand are usually higher in

energy than the d orbitals on the

metal.

Π ACCEPTOR LIGANDS (ML)

The metal uses the

t 2g set of orbitals (d xy , d yz and d xz )

to engage in pi bonding with the

ligand The π* orbitals on the

ligand are usually higher in

energy than the d orbitals on the

metal.

Π ACCEPTOR LIGANDS (ML)

The interaction causes the energy of the t 2g

bonding orbitals to drop slightly,

thus increasing the size of ∆ o

1 All ligands are σ donors In general,

ligand that engage solely in σ bonding are

in the middle of the spectrochemical series Some very strong σ donors, such as CH3-

and H- are found high in the series.

2 Ligands with filled p or d orbitals can

also serve as π donors This results in a

smaller value of ∆o.

3 Ligands with empty p, d or π* orbitals

can also serve as π acceptors This results

in a larger value of ∆o.

I-<Br-<Cl-<F-<H2O<NH3<PPh3<CO

π donor< weak π donor<σ only< π

acceptor

4 – COORDINATE

COMPLEXES

Square planar and tetrahedral complexes are quite common for certain

transition metals The splitting patterns of

the d orbitals on the metal will differ

depending on the geometry of the complex.

TETRAHEDRAL COMPLEXES

The d z 2 and d x 2 -y 2 orbitals point directly between the ligands in a tetrahedral arrangement As a result, these

two orbitals, designated as e in the point group T d, are lower in energy.

TETRAHEDRAL COMPLEXES

The t 2 set of orbitals,

consisting of the d xy , d yz , and d xz orbitals, are directed more in the direction of the ligands

These orbitals will be higher in energy in a tetrahedral field due

to repulsion with the electrons on the ligands

TETRAHEDRAL COMPLEXES

The size of the splitting, ∆ T ,

is considerable smaller than with comparable octahedral complexes This

is because only 4 bonds are formed, and the metal orbitals used in bonding don’t point right at the ligands as they do in octahedral complexes.

TETRAGONAL

COMPLEXES

Six coordinate complexes, notably those of Cu2+, distort from

octahedral geometry One such distortion

is called tetragonal distortion, in which the

bonds along one axis elongate, with

compression of the bond distances along the other two axes.

TETRAGONAL COMPLEXES

The elongation

along the z axis causes the d

orbitals with density along the axis to drop in energy As a result, the d xz and d yz orbitals lower in

energy.

TETRAGONAL COMPLEXES

The compression

along the x and y axis causes

orbitals with density along these axes to increase in energy

.

SQUARE PLANAR COMPLEXES

For complexes with 2 electrons in the e g set of orbitals, a d 8

configuration, a severe distortion may occur, resulting in a 4-coordinate

square planar shape, with the ligands

along the z axis no longer bonded to

the metal.

SQUARE PLANAR COMPLEXES

Square planar complexes are quite common for the d 8

metals in the 4 th and 5 th periods: Rh(I), IR(I), Pt(II), Pd(II) and Au(III) The lower transition metals have large ligand

field stabalization energies, favoring four-coordinate complexes.

SQUARE PLANAR COMPLEXES

Square planar complexes are rare for the 3 rd period metals Ni(II) generally forms

tetrahedral complexes Only with very strong ligands such as CN - , is square planar geometry seen with Ni(II)

SQUARE PLANAR COMPLEXES

The value of ∆ sp for a given metal, ligands and bond length

is approximately 1.3(∆ o )

THE JAHN-TELLER

EFFECT

If the ground electronic configuration

of a non-linear complex is orbitally

degenerate, the complex will distort so as

to remove the degeneracy and achieve a lower energy.

THE JAHN-TELLER

EFFECT

The Jahn-Teller effect predicts which structures will distort It does not

predict the nature or extent of the

distortion The effect is most often seen

when the orbital degneracy is in the orbitals that point directly towards the ligands.

THE JAHN-TELLER

EFFECT

In octahedral complexes, the effect is most pronounced in high spin d4, low spin d7 and d9 configurations, as the degeneracy occurs in the eg set of orbitals.

d4 d7 d9

eg

t2g

-*There is only 1 possible ground state configuration.

- No Jahn-Teller distortion is expected.

EXPERIMENTAL

EVIDENCE OF LFSE

do d1 d2 d3 d4 d5 d6 d7 d8 d9 d10