Coordination chemistry - MOT

51
V.SANTHANAM DEPARTMENT OF CHEMISTRY SCSVMV

Transcript of Coordination chemistry - MOT

Page 1: Coordination chemistry - MOT

V.SANTHANAMDEPARTMENT OF CHEMISTRY

SCSVMV

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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.

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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 dxz and dyz orbitals lower in energy.

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Tetragonal Complexes

The compression along the x and y axis causes orbitals with density along these axes to increase in energy.

.

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Tetragonal Complexes

For complexes with 1-3 electrons in the eg set of orbitals, this type of tetragonal distortion may lower the energy of the complex.

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Square Planar Complexes

For complexes with 2 electrons in the eg set of orbitals, a d8 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.

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Square Planar Complexes

Square planar complexes are quite common for the d8 metals in the 4th and 5th 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.

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Square Planar Complexes

Square planar complexes are rare for the 3rd period metals. Ni(II) generally forms tetrahedral complexes. Only with very strong ligands such as CN-, is square planar geometry seen with Ni(II).

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Square Planar Complexes

The value of ∆sp for a given metal, ligands and bond length is approximately 1.3(∆o).

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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.

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Tetrahedral Complexes

The dz2 and dx2-y2 orbitals point directly between the ligands in a tetrahedral arrangement.

As a result, these two orbitals, designated as e in the point group Td, are lower in energy.

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Tetrahedral Complexes The t2 set of orbitals,

consisting of the dxy, dyz, and dxz 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.

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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.

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Tetrahedral Complexes

In general, ∆T ≈ 4/9 ∆o. Since the splitting is smaller, all tetrahedral complexes are weak-field, high-spin cases.

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Octahedral, Tetrahedral & Square Planar

CF Splitting pattern for CF Splitting pattern for various molecular various molecular geometrygeometry

M

dz2dx2-y2

dxzdxy dyz

M

dx2-y2 dz2

dxzdxy dyz

M

dx

z

dz2

dx2-y2

dxy

dyz

OctahedralTetrahedralTetrahedral Square Square

planarplanar

Pairing energy Vs. Weak field < PeStrong field > Pe

Small High Spin

Mostly d8

(Majority Low spin)Strong field ligandsi.e., Pd2+, Pt2+, Ir+, Au3+

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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.

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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.

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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

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The Jahn-Teller Effect

The strength of the Jahn-Teller effect is tabulated below: (w=weak, s=strong)# e# e-- 11 22 33 44 55 66 77 88 99 11

00High High spinspin ** ** ** ss -- ww ww ** ** **

Low Low spinspin ww ww -- ww ww -- ss -- ss --

*There is only 1 possible ground state configuration.- No Jahn-Teller distortion is expected.

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SummaryCrystal Field Theory provides a basis for explaining many features of transition-metal complexes. Examples include why transition metal complexes are highly colored, and why some are paramagnetic while others are diamagnetic. The spectrochemical series for ligands explains nicely the origin of color and magnetism for these compounds. There is evidence to suggest that the metal-ligand bond has covalent character which explains why these complexes are very stable. Molecular Orbital Theory can also be used to describe the bonding scheme in these complexes.

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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.

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Ligand Field Theory

The A1g group orbitals have the same symmetry as an s orbital on the central metal.

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Ligand Field Theory

The T1u group orbitals have the same symmetry as the p orbitals on the central metal. (T representations are triply degenerate.)

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Ligand Field Theory

The Eg group orbitals have the same symmetry as the dz2 and dx2-

y2 orbitals on the central metal. (E representations are doubly degenerate.)

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Ligand Field Theory

Since the ligands don’t have a combination with t2g symmetry, the dxy, dyz and dxy orbitals on the metal will be non-bonding when considering σ bonding.

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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.

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Ligand Field Theory

The electrons from the ligands (12 electrons from 6 ligands in octahedral complexes) will fill the lower bonding orbitals.

{

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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.

{

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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.

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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.

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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 field. For the upper transition metals, the spin-only magnetic moment, μs, can be used to determine the number of unpaired electrons.

μs = [n(n+2)]1/2

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Magnetic Moments

The magnetic moment of a substance, in Bohr magnetons, can be related to the number of unpaired electrons in the compound.

μs = [n(n+2)]1/2

Where n is the number of unpaired electrons

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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.

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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.

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Considering π Bonding

OOhh EE 8C8C336C6C

22

6C6C44

3C3C2 2

(=C(=C4422))

ii 6S6S44

8S8S66

33σσhh

66σσdd

ΓΓππ1122 00 00 00 -4-4 00 00 00 00 00

y y

y

y y

y

x

x

x

x

x

x

zz

z

z

zz

Consider only the px and pz orbitals on each of the ligands to obtain Γπ.

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Considering π Bonding

OOhh EE 8C8C336C6C

22

6C6C44

3C3C2 2

(=C(=C4422))

ii 6S6S44

8S8S66

33σσhh

66σσdd

ΓΓππ1122 00 00 00 -4-4 00 00 00 00 00

This reduces to T1g + T2g + T1u + T2u. The T2g set has the same symmetry as the dxy, dyz and dxz orbitals on the metal. The T1u set has the same symmetry as the px, py and pz orbitals on the metal.

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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.

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π Bonding

The main source of π bonding is between the dxy, dyz and dxz orbitals on the metal and the d, p or π* orbitals on the ligand.

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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.

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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

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π Donor Ligands (LM)

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.

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π Donor Ligands (LM)

The bonding orbitals, which are lower in energy, are primarily filled with electrons from the ligand, the and antibonding molecular orbitals are primarily occupied by electrons from the metal.

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π Donor Ligands (LM)

The size of ∆o decreases, since it is now between an antibonding t2g orbital and the eg* orbital.This is confirmed by the spectrochemical series. Weak field ligands are also pi donor ligands.

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π Acceptor Ligands (ML)

Ligands such as CN, N2 and CO have empty π antibonding orbitals of the proper symmetry and energy to interact with filled d orbitals on the metal.

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π Acceptor Ligands (ML)

The metal uses the t2g set of orbitals (dxy, dyz and dxz) 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.

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π Acceptor Ligands (ML)

The metal uses the t2g set of orbitals (dxy, dyz and dxz) 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.

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π Acceptor Ligands (ML)

The interaction causes the energy of the t2g bonding orbitals to drop slightly, thus increasing the size of ∆o.

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Summary

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.

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Summary

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<COπ donor< weak π donor<σ only< π

acceptor