Electronic structure Crystal-field theory Ligand-field theory. Electronic-spectra electronic spectra of atoms
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1 Chapter 19 d-metal complexes: electronic structure and spectra Electronic structure 19.1 Crystal-field theory 19.2 Ligand-field theory Electronic-spectra 19.3 electronic spectra of atoms 19.4 electronic spectra of complexes 19.5 Charge-transfer bands 19.6 Selection rules and intensities 19.7 Luminescence
2 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
3 Experimental Evidence for Electronic Structures Thermodynamic Data Magnetic Susceptibility Electronic Spectra Coordination Numbers and Molecular Shapes
4 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)
5 Experimental Evidence for Electronic Structures; Thermodynamic Data What is the stability constants? The equilibrium constants for formation of coordination complex.
6 Experimental Evidence for Electronic Structures; Thermodynamic Data Stability constants HSAB concepts Thermodynamic values Prediction of properties, structures
7 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).
8 Experimental Evidence for Electronic Structures; Thermodynamic Data HSAB concepts
9 Experimental Evidence for Electronic Structures; Thermodynamic Data Chelating Ligands Entropy Effect Chelate Effect Five or six membered ring en vs methyl amine Figure in head. Stability.
10 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
11 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.
12 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 (m s ) Total spin magnetic moment Spin quantum # S (sum of m s ) Isolated oxygen atom 1s 2 2s 2 p 4 S = +1/2 +1/2 +1/2-1/2 = 1 Electron spin Orbital magnetic moment (m l ) Total orbital magnetic moment Orbital quantum # L (sum of m l ) Max. L for the p 4 L = = 1
13 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
14 Experimental Evidence for Electronic Structures; Electronic Spectra Give a direct evidence of orbital energy level Give an information for geometry of complexes
15 Theories of Electronic Structure Valence bond theory Crystal field theory Ligand field theory Angular overlap method
16 Theories of Electronic Structure; Valence bond theory Hybridization ideas Octahedral: d 2 sp 3 d orbitals could be 3d or 4d for the first-row transition metals. (hyperligated, hypoligated)
17 Theories of Electronic Structure; Valence bond theory Fe(III) Isolated ion; 5 unpaired e - In O h compound; 1 or 5 unpaired e - Co(II) Low spin Low spin High spin High spin
18 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.
19 Theories of Electronic Structure; Crystal field theory Repulsion between d-orbital electrons and ligand electrons Splitting of energy levels of d-orbitals
20 Theories of Electronic Structure; Crystal field theory
21 Theories of Electronic Structure; Crystal field theory
22 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); 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.
23 Theories of Electronic Structure; Crystal field theory
24 Theories of Electronic Structure; Crystal field theory
25 Theories of Electronic Structure; Crystal field theory
26 Theories of Electronic Structure; Crystal field theory
27 Theories of Electronic Structure; Crystal field theory
28 Theories of Electronic Structure; Crystal field theory
29 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
30 Theories of Electronic Structure; Crystal field theory E = strong field weak field E > 0 weak field E < 0 strong field
31 Theories of Electronic Structure; Crystal field theory What determine? Depends on the relative energies of the metal ions and ligand orbitals and on the degree of overlap.
32 Theories of Electronic Structure; Crystal field theory Spectrochemical Series for Metal Ions Oxidation # Small size & higher charge Pt 4+ > Ir 3+ > Pd 4+ > Ru 3+ > Rh 3+ >Mo 3+ > Mn 4+ > Co 3+ > Fe 3+ > V 2+ > Fe 2+ Co 2+ > Ni 2+ > Mn 2+ Only low spin aqua complex
33 Ligand field theory; Molecular orbitals for Octahedral complexes CFT & MO were combined The d x2-y2 and d z2 orbitals can form bonding orbitals with the ligand orbitals, but d xy, d xz, and d yz orbitals cannot form bonding orbitals
34 Ligand field theory; Molecular orbitals for Octahedral complexes The combination of the ligand and metal orbitals (4s, 4p x, 4p y, 4p z, 3d z2, and 3d x2-y2 ) form six bonding and six antibonding with a 1g, e g, t 1u symmetries. The metal T 2g orbitals do Electron not have in bonding appropriate orbitals provide symmetry the - nonbonding potential energy that holds molecules together
35 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. d 0 ~d 3 and d 8 ~d 10 only one electron configuration possible no difference in the net spin Strong fields lead to low-spin complexes Weak fields lead to high-spin complexes
36 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.
37 Ligand field theory; Orbital Splitting and Electron Spin Spectrochemical Series for Metal Ions Oxidation # Small size & higher charge Pt 4+ > Ir 3+ > Pd 4+ > Ru 3+ > Rh 3+ >Mo 3+ > Mn 4+ > Co 3+ > Fe 3+ > V 2+ > Fe 2+ Co 2+ > Ni 2+ > Mn 2+
38 Ligand field theory; Ligand field Stabilization Energy
39 Ligand field theory; Orbital Splitting and Electron Spin Orbital configuration of the complex is determined by o, c, and e In general o for 3 + ions is larger than o for 2 + ions with the same # of e -. o > low-spin o < high-spin For low-spin configuration Require a strong field ligand
40 Ligand field theory; Ligand field Stabilization Energy
41 Ligand field theory; Orbital Splitting and Electron Spin The position of the metal in the periodic table Second and third transition series form lowspin 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
42 Ligand field theory; Pi-Bonding The reducible representation is
43 Ligand field theory; Pi-Bonding LUMO orbitals:can be used for bonding with metal HOMO
44 Ligand field theory; Pi-Bonding metal-to-ligand bonding or back-bonding -Increase stability -Low-spin configuration -Result of transfer of negative charge away from the metal ion Ligand-to metal bonding -decrease stability -high-spin configuration
45 Ligand field theory; Square planar Complexes; Sigma bonding
46 Ligand field theory; Square planar Complexes; Sigma bonding ll e - from metal 16 e - 8 e -
47 Ligand field theory; Tetrahedral Complexes; Sigma bonding The reducible representation is A 1 and T 2
48 Ligand field theory; Tetrahedral Complexes; Pi bonding The reducible representation is E, T 1 and T 2
49 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.
50 Angular Overlap: Sigma-Donor Interactions The strongest 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.
51 Angular Overlap: Sigma-Donor Interactions
52 Angular Overlap: Sigma-Donor Interactions Stabilization is 12e
53 Angular Overlap: Pi-Acceptor Interactions The strongest interaction is considered to be between a metal d xy orbitals and a ligand * orbital. Because of the overlap for these orbitals is smaller than the overlap, e < e.
54 Angular Overlap: Pi-Acceptor Interactions
55 Angular Overlap: Pi-Acceptor Interactions
56 Angular Overlap: Pi-Donor Interactions In general, in situations involving ligands that can behave as both acceptors and donors (such as CO, CN - ), the acceptor nature predominates.
57 Angular Overlap: Pi-Donor Interactions
58 Angular Overlap: Pi-Acceptor Interactions
59 Angular Overlap: Types of the ligands and the spectrochemical series Spectrochemical Series for Ligands CO > CN - > PPh 3 > NO 2- > phen > bipy > en NH 3 > py > CH 3 CN > NCS - > H 2 O > C 2 O 4 2- OH - > RCO 2- > F - > N 3- > NO 3- > Cl - > SCN - S 2- > Br - > I - donor only acceptor (strong field ligand) donor(weak field ligand)
60 Angular Overlap: Magnitudes of e e and Metal and ligand
61 Angular Overlap: Magnitudes of e e and Angular overlap parameters derived from electronic spectra e is always larger than e. overlap isoelectronic The magnitudes of both the and parameters with size and electronegativity of the halide ions. overlap
62 Angular Overlap: Magnitudes of e e 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.
63 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.
64 Angular Overlap: The Jahn-Teller Effect
65 Angular Overlap: Four- and Six-Coordinate Preference Only bonding is considered. Angular overlap calculations Low-spin square planar Large # of bonds formed in the octahedral complexes.
66 Angular Overlap: Four- and Six-Coordinate Preference
67 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.
68 Angular Overlap: Other shapes Strength of interaction /4 9/8 9/8 0 0
69 Angular Overlap: Other shapes Trigonal-bipyramidal ML 5 (D 3h ) -donor only
70 Homework Exercise 10-1~10-11 Problem 2, 6a, 6c,11, 13, 16.
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