chapter 10 coordination chemistry ii: bonding · 2017-11-13 · chapter 10 coordination chemistry...

10-4 Angular Overlap

10-5 The Jahn-Teller Effect

10-6 Four- and Six-Coordinate Preferences

10-7 Other Shapes

10-3 Ligand Field Theory

10-2 Theories of Electronic Structure

10-1 Experimental Evidence for Electronic Structures

Chapter 10 Coordination Chemistry II: Bonding

“Inorganic Chemistry” Third Ed. Gary L. Miessler, Donald A. Tarr, 2004, Pearson Prentice Hallhttp://en.wikipedia.org/wiki/Expedia

Experimental Evidence for Electronic Structures

Coordination Numbers and Molecular Shapes

Electronic SpectraMagnetic SusceptibilityThermodynamic Data

Experimental Evidence for Electronic Structures;Thermodynamic Data

One of the primary goal of a bonding theory is to explain the energy of compound.

The energy is openly not determined directly by experiment.

Thermodynamic measurements of enthalpies and free energies of reaction are used to compare.

Bonding strength → Stability constants(formation constants)


What is the stability constants?

The equilibrium constants for formation of coordination complex.


Stability constants

Thermodynamic values →Prediction of properties, structures

HSAB concepts


HSAB concepts

The gist of this theory is that soft acids react faster and form stronger bonds with soft bases, whereas hard acids react faster and form stronger bonds with hard bases, all other factors being equal.

The classification in the original work was mostly based on equilibrium constants for reaction of two Lewis bases competing for a Lewis acid.

Hard acids and hard bases tend to have:small size high oxidation statelow polarizabilityhigh electronegativityenergy low-lying HOMO (bases) or energy high-lying LUMO(acids).


HSAB concepts


Entropy Effect

Chelating Ligands

en vs methyl amine

Figure in head….

Stability….

Chelate EffectFive or six membered ring

Experimental Evidence for Electronic Structures;Magnetic Susceptibility

The magnetic properties of a coordination compound can provide indirect evidence of the orbital energy level.

Hund’s rule → the max. # of unpaired e-.

Diamagnetic: all e- paried→ repelled by a magnetic field

Paramagnetic: all e- paried→ attracted into a magnetic field

Magnetic Susceptibility: Measuring Magnetism


Gouy methodA sample that is to be tested is suspended from a balance between the poles of a magnet. The balance measures the apparent change in the mass of the sample as it is repelled or attracted by the magnetic field.

Magnetic Susceptibility


Electron spin → Spin magnetic moment (ms)

Total spin magnetic moment → Spin quantum # S (sum of ms)

Isolated oxygen atom 1s22s2p4

S = +1/2 +1/2 +1/2 -1/2 = 1

Electron spin → Orbital magnetic moment (ml)Total orbital magnetic moment → Orbital quantum # L (sum of ml)

Max. L for the p4

L = +1 +0 -1 +1 = 1

In physics and applied disciplines such as electrical engineering, the magnetic susceptibility is the degree of magnetization of a material in response to an applied magnetic field.


The equation for the magnetic moment

Contribution from L is small in first transition series

2.00023 ≈ 2

Two sources of magnetic moment – spin (S) and Angular (L) motions of electrons

Spin quantum number

Angular momentum quantum number

Experimental Evidence for Electronic Structures;Electronic Spectra

Give a direct evidence of orbital energy level

Give an information for geometry of complexes

Theories of Electronic Structure

Valence bond theory

Crystal field theory

Ligand field theory

Angular overlap method

Theories of Electronic Structure;Valence bond theory

Hybridization ideasOctahedral: d2sp3

d orbitals could be 3d or 4d for the first-row transition metals. (hyperligated, hypoligated)

Theories of Electronic Structure;Valence bond theory

Fe(III)Isolated ion; 5 unpaired e-

In Oh compound; 1 or 5 unpaired e-

Co(II)

High spin

Low spin

High spin

Low spin

Theories of Electronic Structure;Crystal field theory

Crystal field theory (CFT) is a model that describes the electronic structure of transition metal compounds, all of which can be considered coordination complexes.

CFT successfully accounts for some magnetic properties, colours, hydration enthalpies, and spinel structures of transition metal complexes, but it does not attempt to describe bonding.

CFT was developed by physicists Hans Bethe and John Hasbrouck van Vleck in the 1930s.

CFT was subsequently combined with molecular orbital theory to form the more realistic and complex ligand field theory (LFT), which delivers insight into the process of chemical bonding in transition metal complexes.


Repulsion between d-orbital electrons and ligand electrons→ Splitting of energy levels of d-orbitals


Electrostatic approachIn an Octahedral field of ligand e- pairs; any e-

in them are repelled by the field.Crystal field stabilization energy (CFSE);the actual distribution vs the uniform field.Good for the concept of the repulsion of orbitals by the ligands but no explanation for bonding in coordination complexes.


Why are complexes formed in crystal field theory?Crystal Field Stabilization Energy (CFSE)Or Ligand Field Stabilization Energy (LFSE)→ the stabilization of the d orbitals because of metal-ligand environments


∆E = strong field – weak field∆E > 0 weak field∆E < 0 strong field


What determine ?Depends on the relative energies of the metal ions and ligandorbitals and on the degree of overlap.


Spectrochemical Series for Metal Ions

Oxidation # ↑→ ∆↑Small size & higher charge

Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+

Co2+ > Ni2+ > Mn2+

Only low spin aqua complex

Ligand field theory; Molecular orbitals for Octahedral complexes

CFT & MO were combined

The dx2-y2 and dz2 orbitals can form bonding orbitalswith the ligand orbitals, but dxy, dxz, and dyz orbitalscannot form bonding orbitals

Ligand field theory; Molecular orbitals for Octahedral complexes

The combination of the ligand and metal orbitals (4s, 4px, 4py, 4pz, 3dz2, and 3dx2-y2) form six bonding and six antibonding with a1g, eg, t1u symmetries.

The metal T2g orbitalsdo not have appropriate symmetry - nonbonding

Electron in bonding orbitals provide the potential energy that holds molecules together

Ligand field theory; Orbital Splitting and Electron Spin

Strong-field ligand – Ligands whose orbitalsinteract strongly with the metal orbitals→large ∆o

Weak-field ligand.

d0~d3 and d8 ~d10 – only one electron configuration possible → no difference in the net spin

Strong fields lead to low-spin complexesWeak fields lead to high-spin complexes


What determine ?Depends on the relative energies of the metal ions and ligand orbitals and on the degree of overlap.


Spectrochemical Series for Metal Ions

Oxidation # ↑→ ∆↑Small size & higher charge

Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+

Co2+ > Ni2+ > Mn2+

Ligand field theory; Ligand field Stabilization Energy


Orbital configuration of the complex is determined by ∆o, πc, and πe

In general ∆o for 3+ ions is larger than ∆o for 2+ ions with the same # of e-.

∆o > π low-spin∆o < π high-spin

For low-spin configurationRequire a strong field ligand

Ligand field theory; Ligand field Stabilization Energy


The position of the metal in the periodic table

Second and third transition series form low-spin more easily than metals form the first transition series-The greater overlap between the larger 4d and 5d orbitals and the ligand orbitals-A decreased pairing energy due to the larger volume available for electrons

Ligand field theory; Pi-Bonding

The reducible representation is


LUMO orbitals:can be used for π bonding with metal

HOMO


metal-to-ligand π bonding or π back-bonding-Increase stability-Low-spin configuration-Result of transfer of negative charge away from the metal ion

Ligand-to metal π bonding-decrease stability-high-spin configuration

Ligand field theory; Square planar Complexes; Sigma bonding

Ligand field theory; Square planar Complexes; Sigma bonding

ll ⊥

8 e-

16 e-

e- from metal

Ligand field theory; Tetrahedral Complexes; Sigma bonding

The reducible representation isA1 and T2

Ligand field theory; Tetrahedral Complexes; Pi bonding

The reducible representation isE, T1 and T2

Angular Overlap

LFT →

No explicit use of the energy change that resultsDifficult to use other than octahedral, square planar, tetrahedral.

Deal with a variety of possible geometries and with a mixture of ligand. → Angular Overlap Model

The strength of interaction between individual ligandorbitals and metal d orbitals based on the overlap between them.

Angular Overlap:Sigma-Donor Interactions

The strongest σ interaction

There are no examples of complexes with e- in the antibonding orbitals from s and p orbitals, and these high-energy antibonding orbitals are not significant in describing the spectra of complexes. → we will not consider them further.


Stabilization is 12eσ

Angular Overlap:Pi-Acceptor Interactions

The strongest π interaction is considered to be between a metal dxy orbitals and a ligand π* orbital.

Because of the overlap for these orbitals is smaller than the σ overlap, eπ < eσ.

Angular Overlap:Pi-Donor Interactions

In general, in situations involving ligands that can behave as both π acceptors and π donors (such as CO, CN-), the π acceptor nature predominates.

Angular Overlap:Pi-Donor Interactions

Angular Overlap:Types of the ligands and the spectrochemical series

Spectrochemical Series for Ligands

CO > CN- > PPh3 > NO2- > phen > bipy > en

NH3 > py > CH3CN > NCS- > H2O > C2O42-

OH- > RCO2- > F- > N3

- > NO3- > Cl- > SCN-

S2- > Br- > I-

π acceptor (strong field ligand) π donor(weak field ligand)

σ donor only

Angular Overlap:Magnitudes of eσ eπ and ∆

Metal and ligand


Angular overlap parameters derived from electronic spectra

eσ is always larger than eπ. overlap

The magnitudes of both the σ and πparameters ↓ with ↑ size and ↓electronegativity of the halide ions. overlap

isoelectronic


Can describe the electronic energy of complexes with different shapes or with combinations of different liagnds.

The magnitude of ∆o→ Magnetic properties and visible spectrum.

Angular Overlap:The Jahn-Teller Effect

There cannot be unequal occupation of orbitals with identical orbitals.To avoid such unequal occupation, the molecule distorts so that these orbitals no longer degenerate.In other words, if the ground electron configuration of a nonlinear complex is orbitally degenerate, the complex will distort to remove the degeneracy and achieve a lower energy.

Angular Overlap:The Jahn-Teller Effect

Angular Overlap:Four- and Six-Coordinate Preference

Only σ bonding is considered.

Large # of bonds formed in the octahedral complexes.

Angular overlap calculations

Low-spin square planar


How accurate are these predictions?

Their success is variable, because of there are other differences between metals and between ligands. In addition, bond lengths for the same ligand-metal pair depend on the geometry of the complex.

The interactions of the s and p orbitals.

The formation enthalpy for complexes also becomes more negative with increasing atomic number and increasing ionization energy.

By careful selection of ligands, many of the transition metal ions can form compounds with geometries other than octahedral.

Angular Overlap:Other shapes

Strength of σ–interaction

11

1

11

2+3/4 9/8 9/8 0 0

Angular Overlap:Other shapes

Trigonal-bipyramidal ML5 (D3h) σ-donor only

chapter 10 coordination chemistry ii: bonding · 2017-11-13 · chapter 10 coordination chemistry...

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