Presentation is loading. Please wait.

Presentation is loading. Please wait.

2-1 Orbitals and energetics Bonding and structure Ligand field theory Charge Transfer Molecular orbital theory Provide fundamental understanding of chemistry.

Similar presentations

Presentation on theme: "2-1 Orbitals and energetics Bonding and structure Ligand field theory Charge Transfer Molecular orbital theory Provide fundamental understanding of chemistry."— Presentation transcript:

1 2-1 Orbitals and energetics Bonding and structure Ligand field theory Charge Transfer Molecular orbital theory Provide fundamental understanding of chemistry dictating radionuclide complexes Structure based on bonding §Coordination important in defining structure àStructure related to spectroscopic behavior àElectron configuration important in structure *d 8 are square planar *d 0 and d 10 tetrahedral

2 2-2

3 2-3 Coordination Number Coordination number 2 §Two geometric possibilities àLinear (D ∞h ) àBent (C 2v ) §Common with divalent elements àHigher coordination possible through bridging Pu O O O HH

4 2-4 Coordination Number Coordination number 3 §Planar (D 3h ) §Pyramidal (C 3v ) §Some T-shaped forms (C 2v ) Found with trivalent central elements §For metal ions, not very common N H H H BH H H

5 2-5 Coordination Number Coordination number 4 §Formed by C 3 basic geometries §Tetrahedral (T d ) §Square geometry (C 4h ) §One lone pair (C 2v ) C H H H H Xe F F F F

6 2-6 Coordination Number Coordination number 5 §Trigonal bipyramid (D 3h ) §Square pyramid (C 4v ) Interconvertibility between geometries §Compounds can vary between shapes §Trigonal bipyramid seems to be more common àCommon with metal pentachloride species Cu Cl In Cl

7 2-7 Coordination Number Coordination number 6 Very common coordination number §Ligands at vertices of octahedron or distorted octahedron àOctahedron (O h ) àTetragonal octahedron (D 4h ) *Elongated or contracted long z axis àRhombic (D 2h ) *Changes along 2 axis àTrigonal distortion (D 3d )

8 2-8 Coordination Number O h ->D 4h O h ->D 2h O h ->D 3d or

9 2-9 Higher coordination 7 coordination 8 coordination 9 coordination

10 2-10 Hard and soft metals and ligands Based on Lewis acid definition §Ligand acts as base, donates electron pair to metal ion Hard metal ion interact with hard bases §Hard ligands N, O, F §Soft ligands P, S, Cl àLigand hardness decreases down a group Hard metals §High positive charges §Small radii §Closed shells or half filled configurations

11 2-11 Soft or hard metals and ligands Soft metals §Low positive charges §Large ionic radius §Non-closed shell configurations àTend to be on right side of transition series Lanthanides and actinides are hard §Actinides are softer than lanthanides àLigands with soft groups can be used for actinide/lanthanide separations

12 2-12 Hard Intermediate Soft

13 2-13 Chelation and stability Ligands with more than 1 complexing functional group §Carbonate, ethylenediamine §Enhanced stability through chelation effect §ethylenediamine binding stronger than 2 ammonia groups àBidentate àTridentate §Ligands can wrap around metal ion forming stronger complex

14 2-14 EDTA complex

15 2-15 Effective atomic number Metal bonding can be described with effective atomic number §Number of electrons surrounding metal is effective atomic number àTransitions metal have 9 possible bonds *5 d, 3p, 1 s Ë18 electrons §Possible to have effective atomic number different than 18 àFew d electrons àElectronegative ligands

16 2-16 Effective atomic number 16 electron §Square planar §d 8 configuration (Au, Pt) Greater than 18 electron §8-10 d electrons

17 2-17 Ligand Field Theory O h complexes §Six similar bonds, nine valence orbitals §To make six similar bonds, mixing must occur àHybridization of s, p, and 2 d orbitals *d x2-y2, d z2 Ëd 2 sp 3 hybrid Interaction of d orbitals with bonding ligand results in observed properties (magnetic, color) §Ligand field theory

18 2-18 d orbital hybrid

19 2-19 Ligand Field Theory Symmetry adapted linear combination §Combination of orbitals with symmetry considerations

20 2-20 Ligand Field Theory nd, n+1s, and n+1p orbitals on the metal overlap with one orbital on each of the six ligands §forms 15 molecular orbitals Six are bonding §energies are lower than original atomic orbitals Six are antibonding with higher energy Three are nonbonding Ligand-field theory describes how s,p, and d orbitals on the metal to overlap with orbitals on the ligand

21 2-21 d orbital splitting

22 2-22 Charge transfer Allowed transitions in UV-Visible §Ligand to metal §Metal to ligand Related to redox of metals and ligands §MnO 4 - àO ligands to Mn metal Absorption of radiation involves the transfer of an electron from the donor to an orbital associated with the acceptor. Molar absorptivities from charge-transfer absorption are large (greater that 10,000 L mol-1 cm-1).

23 2-23 MO theory The number of molecular orbitals = the number of atomic orbitals combined Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-bonding orbital (higher energy) Electrons enter the lowest orbital available The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle) Electrons spread out before pairing up (Hund's Rule)

24 2-24 MO theory

25 2-25

26 2-26 Ligand Field Theory Treats overlaps of ligand and metal orbitals Stems for SALC §Sigma §Combine sigma orbitals for each set àt 2g has no sigma §For molecular orbital combine  C M    C L  La1g §Pi bonding  Donor decrease   àAcceptor increases *Related to electrochemical series

27 2-27

28 2-28 Bonding and electronic structure Crystal field theory §Lone pair modeled as point àRepels electrons in d orbital àd orbitals have energy differences due to point *Results in ligand field splitting ËAbout 10 % of metal- ligand interaction Ëe and t orbitals *Ignores covalent contribution §Energy difference is ligand field splitting parameter (Δ o ) àCan be determined from absorption spectrum *e g  t 2g transition

29 2-29 Crystal Field Theory Ti(OH 2 ) 6 3+ §Absorbance at 500 nm, cm -1 §1000 cm -1 = kJ/mol   0 =243 kJ/m   0 found to vary with ligand àFor metal ion increases with oxidation state and increases down a group I - < Br - < SCN - ~C l- < F - < OH- ~ ONO-

30 2-30 Crystal field theory Ligand field stabilization energies  t 2g stabilized (40 % of  o )  e g increase energy (60 % of  o )  LFSE=(-0.4 t 2g e g )  o àLFSE few % of energy

31 2-31 Crystal Field Theory Weak and strong field limits §Related to location of 4 th d 4 electron àt 2g 4 or t 2g 3 e g 1 *All in t 2g has coulombic repulsion (P) but promotion to e g need  o energy   o

32 2-32 Crystal Field Theory Magnetic properties §Determination of spin state àDiamagnetic *Move out of a magnetic field àParamagnetic *Move into a magnetic field §Dipole moment àSpin only paramagnetism due to quenching of orbital angular momentum with ligand *μ=[N(N+2)] 1/2 μ B ; with μ B = 9.274E-24 JT -1 and N number of unpaired electrons ËFor d 6 N= 4 or 0, depending on spin

33 2-33 Crystal Field Theory Accounts for observations on trends §Ionic radius

34 2-34 Crystal Field Theory T d §Weak field splitting §e lower energy than t àBased on orbital spatial distributions Tetragonal complex §Splitting into 4 levels §Can distort into square planar à4d 8 and 5d 8 Jahn-Teller effect §Distortion of geometry to achieve energy stabilization (see previous) àEnergy of distorted complex lower

Download ppt "2-1 Orbitals and energetics Bonding and structure Ligand field theory Charge Transfer Molecular orbital theory Provide fundamental understanding of chemistry."

Similar presentations

Ads by Google