chapter 10 coordination chemistry ii:...
TRANSCRIPT
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)
Experimental Evidence for Electronic Structures;Thermodynamic Data
What is the stability constants?
The equilibrium constants for formation of coordination complex.
Experimental Evidence for Electronic Structures;Thermodynamic Data
Stability constants
Thermodynamic values →Prediction of properties, structures
HSAB concepts
Experimental Evidence for Electronic Structures;Thermodynamic Data
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).
Experimental Evidence for Electronic Structures;Thermodynamic Data
HSAB concepts
Experimental Evidence for Electronic Structures;Thermodynamic Data
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
Experimental Evidence for Electronic Structures;Magnetic Susceptibility
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
Experimental Evidence for Electronic Structures;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.
Experimental Evidence for Electronic Structures;Magnetic Susceptibility
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.
Theories of Electronic Structure;Crystal field theory
Repulsion between d-orbital electrons and ligand electrons→ Splitting of energy levels of d-orbitals
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
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.
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
Theories of Electronic Structure;Crystal field theory
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
Theories of Electronic Structure;Crystal field theory
∆E = strong field – weak field∆E > 0 weak field∆E < 0 strong field
Theories of Electronic Structure;Crystal field theory
What determine ?Depends on the relative energies of the metal ions and ligandorbitals and on the degree of overlap.
Theories of Electronic Structure;Crystal field theory
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
Ligand field theory; Orbital Splitting and Electron Spin
What determine ?Depends on the relative energies of the metal ions and ligand orbitals and on the degree of overlap.
Ligand field theory; Orbital Splitting and Electron Spin
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
Ligand field theory; Orbital Splitting and Electron Spin
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
Ligand field theory; Orbital Splitting and Electron Spin
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
Ligand field theory; Pi-Bonding
LUMO orbitals:can be used for π bonding with metal
HOMO
Ligand field theory; Pi-Bonding
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.
Angular Overlap:Sigma-Donor Interactions
Angular Overlap:Sigma-Donor Interactions
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-Acceptor Interactions
Angular Overlap:Pi-Acceptor Interactions
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:Pi-Acceptor 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:Magnitudes of eσ eπ and ∆
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
Angular Overlap:Magnitudes of eσ eπ and ∆
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
Angular Overlap:Four- and Six-Coordinate Preference
Angular Overlap:Four- and Six-Coordinate Preference
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