Coordination Chemistry II Bonding including crystal field theory

Coordination 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: d 2 sp 3 (using the 3 d, 4 s and 4 p orbitals), or sp 3 d 2 (using the 4 s, 4 p and 4 d orbitals). The valence bond approach isn’t used because it fails to explain the electronic spectra and magnetic moments of most complexes.

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 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.

d Orbital Splitting __ __ e g dz 2 dx 2 -y 2 __ __ __ 0. 6∆o Spherical field 0. 4∆o __ dxy ∆o __ __t 2 g dxz dyz Octahedral field

d Orbital Splitting In some texts and articles, the gap in the d orbitals is assigned a value of 10 Dq. The upper (eg) set goes up by 6 Dq, and the lower set (t 2 g) goes down by 4 Dq. 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 t 2 g 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 Based on measurements for a given metal ion, the following series has been developed: I-<Br-<S 2 -<Cl-<NO 3 -<N 3 -<F-<OH-<C 2 O 42 -<H 2 O <NCS-<CH 3 CN<pyridine<NH 3<en<bipy<phen <NO 2 -<PPh 3<CN-<CO

The Spectrochemical Series The complexes of cobalt (III) show the shift in color due to the ligand. (a) CN–, (b) NO 2–, (c) phen, (d) en, (e) NH 3, (f) gly, (g) H 2 O, (h) ox 2 –, (i) CO 2–. 3

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/sy mmetry/octahedral. html

Octahedral Symmetry

Ligand Field Theory Consider the group orbitals of all six ligands in octahedral geometry. Oh Γσ 3 C 2 E 8 C 3 6 C 2 6 C 4 i 6 S 4 8 S 6 3σh 6σd 2 (=C 4 ) 6 0 0 2 2 0 0 0 4 2 This reduces to A 1 g + Eg + T 1 u

Ligand Field Theory The A 1 g group orbitals have the same symmetry as an s orbital on the central metal.

Ligand Field Theory The T 1 u group orbitals have the same symmetry as the p orbitals on the central metal. (T representations are triply degenerate. )

Ligand Field Theory The Eg group orbitals have the same symmetry as the dz 2 and dx 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 2 g symmetry, the dxy, dyz and dxy orbitals on the metal will be non-bonding when considering σ bonding.

Ligand Field Theory The molecular orbital diagram is consistent with the crystal field approach. Note that the t 2 g 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 4 s and 3 d 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 t 2 g 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 field. For the upper transition metals, the spin -only magnetic moment, μs, can be used to determine the number of unpaired electrons. μ = [n(n+2)]1/2

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

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 t 2 g 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 2 g) set of orbitals before occupying the eg 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 eg set of orbitals before they pair up and fill the lower (t 2 g) 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 d 1 d 2 LFS 0. 4Δo. 8 E d 3 1. 2 d 4 d 5 d 6 d 7 d 8 d 9 d 10. 6 0. 4. 8 1. 2. 6 0

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. M 2+ + 6 H 2 O(l) M(H 2 O)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 o 1 2 3 4 5 6 7 8 9 d d d d LFSE 0. 4Δo. 8 1. . 6 0. 4. 8 2 d d 1. . 6 2 d 10 0

High Spin vs. Low Spin 3 d metals are generally high spin complexes except with very strong ligands. CN- forms low spin complexes, especially with M 3+ ions. 4 d & 4 d metals generally have a larger value of ∆o than for 3 d metals. As a result, complexes are typically low spin.

Effect of Metal Charge on Splitting

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.

z x x z z y x x y x z y y x y z Considering π Bonding Consider only the px and pz orbitals on each of the ligands to obtain Γπ. Oh 3 C 2 E 8 C 3 6 C 2 6 C 4 i 6 S 4 8 S 6 3σh 6σd 2 (=C 4 ) Γπ 1 2 0 0 0 -4 0 0 0

Considering π Bonding Oh 3 C 2 E 8 C 3 6 C 2 6 C 4 i 6 S 4 8 S 6 3σh 6σd 2 (=C 4 ) Γπ 1 2 0 0 0 -4 0 0 0 This reduces to T 1 g + T 2 g + T 1 u + T 2 u. The T 2 g set has the same symmetry as the dxy, dyz and dxz orbitals on the metal. The T 1 u set has the same symmetry as the px, py and pz orbitals on the metal.

Considering π Bonding π reduces to: T 1 g + T 2 g + T 1 u + T 2 u. n The T 1 g and T 2 u group orbitals for the ligands don’t match the symmetry of any of the metal orbitals. n The T 1 u 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. n The T 2 g 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 dxy, dyz and dxz 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* 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 S 2 -. The filled p or d orbitals on these ions interact with the t 2 g set of orbitals (dxy, dyz and dxz) on the metal to form bonding and antibonding molecular orbitals.

π Donor Ligands (L M) 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.

σ Bonding vs. Bonding

π Donor Ligands (L M) The size of ∆o decreases, since it is now between an antibonding t 2 g orbital and the eg* 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 2 g 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.

π Acceptor Ligands (M L) The metal uses the t 2 g 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.

π Acceptor Ligands (M L) The interaction causes the energy of the t 2 g bonding orbitals to drop slightly, thus increasing the size of ∆ o.

Effect of Ligand on Splitting

Summary 1. All ligands are σ donors. In general, ligands that engage solely in σ bonding are in the middle of the spectrochemical series. Some very strong σ donors, such as CH 3 - 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.

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-<H 2 O<NH 3<PPh 3<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 dz 2 and dx 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 Td, are lower in energy.

Tetrahedral Complexes The t 2 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.

Tetrahedral Complexes The size of the splitting, ∆T, is considerably 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.

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

Tetragonal Complexes Six coordinate complexes, notably those of Cu 2+, 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 dxz and dyz 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. .

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.

Square Planar Complexes For complexes with 2 electrons in the eg 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 stabilization energies, favoring fourcoordinate 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 d 4, low spin d 7 and d 9 configurations, as the degeneracy occurs in the eg set of orbitals. d 4 d 7 d 9 eg t 2 g

The Jahn-Teller Effect The strength of the Jahn-Teller effect is tabulated below: (w=weak, s=strong) # e- 1 2 3 4 5 6 7 8 9 10 High * * * s - w w * * * spin Low w w - s - s spin *There is only 1 possible ground state configuration. - No Jahn-Teller distortion is expected.