Chapter 10 Coordination Chemistry II: Bonding
10-1 Experimental Evidence for Electronic Structures
10-2 Theories of Electronic Structure
10-3 Ligand Field Theory
10-4 Angular Overlap
10-5 The Jahn-Teller Effect
10-6 Four- and Six-Coordinate Preferences
10-7 Other Shapes
“Inorganic Chemistry” Third Ed. Gary L. Miessler, Donald A. Tarr, 2004, Pearson Prentice Hall http://en.wikipedia.org/wiki/Expedia Experimental Evidence for Electronic Structures
Thermodynamic Data Magnetic Susceptibility Electronic Spectra Coordination Numbers and Molecular Shapes 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.
상대적 세기 Experimental Evidence for Electronic Structures; Thermodynamic Data
Stability constants HSAB concepts
Thermodynamic values → Prediction of properties, structures 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 state low polarizability high electronegativity energy 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
Chelating Ligands
Entropy Effect en vs methyl amine
Chelate Effect Stability…. Five or six membered ring
Check values !! 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- paired → repelled by a magnetic field
Paramagnetic: all e- not paired → attracted into a magnetic field
Magnetic Susceptibility: Measuring Magnetism Experimental Evidence for Electronic Structures; Magnetic Susceptibility
Magnetic Susceptibility
Gouy method A 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. Experimental Evidence for Electronic Structures; Magnetic Susceptibility
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.
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 Experimental Evidence for Electronic Structures; Magnetic Susceptibility
Two sources of magnetic moment – spin (S) and Angular (L) motions of electrons
Spin quantum number Angular momentum quantum number The equation for the magnetic moment
Contribution from L is small in first transition series
2.00023 ≈ 2 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 ideas Octahedral: 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)
Low spin Low spin
High spin High 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
Understand !! Theories of Electronic Structure; Crystal field theory
Electrostatic approach In an Octahedral field of ligand e- pairs; any e- in them are repelled by the field.
Crystal field stabilization energy (CFSE); Understand !! 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
Can you draw this ? 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
Can you draw this ? 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
Which way ? Theories of Electronic Structure; Crystal field theory
What determine ? Depends on the relative energies of the metal ions and ligand Spectrochemical Series for Metal Ions orbitals and on the degree of overlap. Oxidation # ↑→ ∆↑ Small size & higher charge Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+ Co2+ > Ni2+ > Mn2+ Theories of Electronic Structure; Crystal field theory
Spectrochemical Series for Metal Ions
Oxidation # ↑→ ∆↑ Small size & higher charge Only low spin aqua complex Pt4+ > Ir3+ > Pd4+ > Ru3+ > Rh3+ >Mo3+ > Mn4+ > Co3+ > Fe3+ > V2+ > Fe2+ Co2+ > Ni2+ > Mn2+ Ligand field theory; Molecular orbitals for Octahedral complexes
CFT & MO were combined
The dx2-y2 and dz2 orbitals can form bonding orbitals with the ligand orbitals, but dxy, dxz, and dyz orbitals cannot form bonding orbitals Are you agree ? Ligand field theory; Molecular orbitals for Octahedral complexes Concept !! 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 orbitals do not have appropriate symmetry - nonbonding Check this !! 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 orbitals interact 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 Understand it !! Strong fields lead to low-spin complexes Weak 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-. Understand ?
∆o > π low-spin ∆o < π high-spin
For low-spin configuration Require a strong field ligand Ligand field theory; Orbital Splitting and Electron Spin
The position of the metal in the periodic table Rationalize !! 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 Can you -Low-spin configuration draw it ? -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 ⊥
e- from metal
16 e-
8 e- Ligand field theory; Tetrahedral Complexes; Sigma bonding
The reducible representation is
A1 and T2 Ligand field theory; Tetrahedral Complexes; Pi bonding
The reducible representation is
E, T1 and T2 Angular Overlap
LFT →
No explicit use of the energy change that results Difficult 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 ligand orbitals 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
- - σ donor only CO > CN > PPh3 > NO2 > phen > bipy > en - 2- NH3 > py > CH3CN > NCS > H2O > C2O4 ------OH > RCO2 > F > N3 > NO3 > Cl > SCN S2- > Br- > I-
π acceptor (strong field ligand) π donor(weak field ligand) Angular Overlap: Magnitudes of eσ eπ and ∆
Metal and ligand
Rationalize !! Angular Overlap: Magnitudes of eσ eπ and ∆
Angular overlap parameters derived from electronic spectra
eσ is always larger than eπ. overlap Check trend !! The magnitudes of both the σ and π parameters ↓ with ↑ size and ↓ electronegativity of the halide ions. overlap 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
Angular overlap calculations Only σ bonding is considered.
Large # of bonds formed in Low-spin square planar the octahedral complexes. 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
1 1
1
Strength of σ–interaction 1 1
2+3/4 9/8 9/8 0 0 Angular Overlap: Other shapes
Trigonal-bipyramidal ML5 (D3h) σ-donor only