Module 2 for students.pptx

Engineering Chemistry Module 2 Metal Complexes and Organometallics By Dr. S. Mohana Roopan Email: mohanaroopan.s@v...

Engineering Chemistry Module 2 Metal Complexes and Organometallics By Dr. S. Mohana Roopan Email: [email protected] Ph. 9865610356 Inorganic complexes - Structure, bonding and application Organometallics – Introduction, stability, structure and applications of metal carbonyls and ferrocene. Metals in biology: Haemoglobin and chlorophyll- structure and property  Inorganic Complexes: Structure, Bonding and Applications When Two or more stable compounds are mixed in stoichiometric amounts, two types of addition compounds are formed. i.e. Double salts & Complex salts. Double Salts and Coordination Compounds  A double salt is a salt that contains more than one different cation or anion. ❖ Ex: Ferric alum (NH4)2SO4.Fe2(SO)4.24H2O In water: NH4+, SO42-, Fe3+ Complexes or coordination compounds are molecules that possess a metal center that is bound to ligands (atoms, ions, or molecules that donate electrons to the metal). These complexes can be neutral or charged. When the complex is charged, it is stabilized by neighboring counter-ions. Note that double salts should not be confused with a complex. 2  Double Salt Complex Salt Double salts are ionic compounds that are Complex salts are ionic compounds that contain formed by the combination of two different salt a central metal atom which is surrounded by compounds. ligands that are linked to it via covalent bonds. Completely dissociate in aqueous medium Complex salts do not completely dissociate in aqueous medium. When double salts are added to water, they give When complex salts are added to water, they do simple ions. not give simple ions. Analysis can be done by determining the ions Can not be done by identifying the ions in that are present in the aqueous solution. aqueous solution. Ex: Potassium sodium tartrate, bromlite, Tutton’s Ex.: Potassium ferrocyanide, Argento cyanide, salt tetra amino cupric sulphate. 3 Ligand s  A ligand is an ion or molecule, which donates a pair of electrons to the central metal atom or ion to form a coordination complex. ▪ Coordinate covalent bond : metal-ligand bond ▪ Monodentate : one bond to metal ion ▪ Bidentate : two bond to metal ion ▪ Polydentate : more than two bonds to a metal ion possible 4 Chelating Agents  Chelation is a type of bonding of ions and molecules to metal ions.  It involves the formation of Coordinate bonds between a polydentate ligand and a Chelation single central atom. These ligands are called chelants, chelators, chelating agents, or sequestering agents.  Chelating agents are organic compounds capable of linking together metal ions to form complex ring-like structures called chelates. Porphyrin  Ex.: (i) Phosphates are used to tie up Ca2+ and Mg2+ in hard water to prevent them from interfering with detergents. 5  Werner Coordination Albert Werner, 1913 Noble prize Theory Werner's Theory: Alfred Werner, Swiss chemist put forward a theory to explain the formation of complex compounds. Warner studied the following metal complexes: CoCl3 forms four different compounds with NH3. Werner Coordination Theory Original Color Lone per formula “Free” Cl- Ions Modern Formulation Formulation Unit Per Formula Unit CoCl5.6 NH3 Orange 1 3 [Co(NH3)6]Cl3 CoCl3. 5 NH3 Purple 3 2 [Co(NH3)5Cl]Cl2 CoCl3.4 NH3 Green 2 1 trans – [Co(NH3)4Cl2]Cl CoCl3. 4 NH3 Violet 2 1 Cis - [Co(NH3)4Cl2]Cl Limitation of Werner Coordination Theory  It failed to explain why all elements don’t form coordination compounds.  It failed to explain the directional properties of bonds in coordination compounds.  It does not explain the colour, and the magnetic and optical properties shown by coordination compounds. 8  Lewis Acid Base Gilbert N. Lewis, 1920s Theory ❖ Lewis Acid/Base reactions: Base : electron pair donor; Acid : electron pair acceptor. ❖ Ligands : Lewis bases ❖ Metals : Lewis acids ❖ Coordinate covalent bonds ❖ Metal Complexes - Formation of a complex was described as an acid - base reaction according to Lewis.  Sidgwick’s Effective atomic number (EAN) rule is based on the octet Sidgwick’s Rule theory of Lewis and this is the first attempt to account for the bonding in complexes. Bonding in coordination compounds ❖ There has been much work done in attempting to formulate theories to describe the bonding coordination compounds. ❖ The first theory adopted to describe the bonding coordination compounds is Valence Bond Theory (VBT). ❖ Followed that Crystal Field theory (CFT) has been adopted. ❖ Ligand Field theory (LFT) & Molecular Orbital (MO) theory are considered as sophisticated model when compare to CFT. These theories continue to contribute to current discussions of coordination compounds. 10 Valence Bond Theory Linus Pauling, 1931 (VBT) Valence bond theory predicts that the bonding in a metal complex arises from overlap of filled ligand orbitals and vacant metal orbitals. Covalent bond is formed by the overlapping of two half filled atomic orbitals of two atoms having electrons with opposite spin. 11 Valence Bond Theory 12  Tetrahedral Geometry Tetrahedral nickel complex [NiCl4]2- 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Since Cl- is a weak field ligand, pairing of electrons don’t not take place (3d orbitals) Two unpaired electron - paramagnetic and attracted by magnets  Square Planar Geometry Square planar nickel complex [Ni(CN)4]2- 3d 4s 4p Ni (3d84s2) Ni2+ [Ni(CN)4]2- dsp2 All paired electrons – diamagnetic - weakly repelled by magnets Since CN- is a strong field ligand, pairing of electrons take place (3d orbitals) All paired electron - diamagnetic – Weakly repelled by magnets  Octahedral sp3d2 Outer orbital octahedral complex Geometry Octahedral cobalt complex Sp3d2 hybridization 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Since F- is a weak field ligand, pairing of electrons don’t not take place (3d orbitals) Four unpaired electron - paramagnetic and attracted by magnets 15  Octahedral d2sp3 Inner orbital octahedral complex Geometry Octahedral iron complex [Fe(CN)6]3- 3d 4s 4p Fe 3d64s2 Fe3+ [Fe(CN)6]3- 1s2 2s2 2p6 3s2 3p6 4s2 3d6 CN– Strong ligand d2sp3 Since CN- is a strong field ligand, pairing of electrons take place (3d orbitals) One unpaired electron - paramagnetic – attracted by magnets 16  Limitation of Valence Bond Theory ❖ The valence bond approach could not explain the following ▪ It involves a number of assumptions. ▪ It does not explain the colour exhibited by the coordination compounds. ▪ It does not provide a quantitative interpretation of the thermodynamic or kinetic stability of the coordination compounds. ▪ It does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes. ▪ It does not distinguish between weak and strong ligands.  Crystal field theory Hans Bethe in 1929 (CFT) ❖ The Crystal Field Theory (CFT) is a model for the bonding interaction between transition metals and ligands. ❖ The repulsion between ligand lone pairs and the d-orbitals on the metal results in a splitting of the energy of the “d” orbitals. Orbital occupancy for high- and low-spin complexes of d4 to d7 metal ions high high high low spin: high low spin: spin: low spin: spin: low spin: spin: strong- spin: strong- weak- strong- weak- strong- weak- field ligand weak- field ligand field field ligand field field ligand field field ligand ligand ligand ligand  CFT Assumptions ⮚ Interaction between the metal ion and the ligands are purely electrostatic (ionic). ⮚ Ligands are considered as point charges. ⮚ Ion-ion interaction, if the ligand is negatively charged and ion-dipole ⮚ Interaction between electrons of the cation and those of ligands interaction, if the ligand is neutral are entirely repulsive. This is ⮚ Electrons on the metal are under responsible for splitting of d repulsive from those on the orbitals. ligands. ⮚ CFT does not consider the ⮚ Electrons on metal occupy those overlapping between metal and ligand orbitals. d-orbitals farthest away from the direction of approach of ligands. ⮚ The d-orbitals lose their degeneracy due to the approach of ligands during the formation of complex. 19  Octahedral Complex and d-Orbital Energies ⮚ For the Oh point group, the dx2-y2, dz2 orbitals belong to the eg irreducible representation and xy, xz, yz belong to the T2g representation. ⮚ The splitting extent of these two sets of orbitals is denoted by ∆0 or 10 Dq. As the barycenter must be conserved on going from a spherical field to an octahedral field, the t2g set must be stabilized as much as the 20 eg set is destabilized. The magnitude of crystal field splitting energy (Δ0 or Δt) depends on different factors are given below: 1. Oxidation state of the central metal ion: Larger the charge on the central metal ion, greater is the extent of Splitting. Ex. [Fe(H2O)6]+3 has Fe+3 ion and its Δ0 13700 cm-1; [Fe(H2O)6]+2 has Fe+2 ion and its Δ0 10400 cm-1 [Co(H2O)6]+3 has Co+3 ion and its Δ0 18600 cm-1; [Co(H2O)6]+2 has Co+2 ion and its Δ0 9300 cm-1 2. Type of d-orbital: Generally the magnitude of crystal field splitting is in the order 3d