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Transition metal complexes u u u A transition metal complex is species consisting of a transition metal coordinated (bonded to) one or more ligands (neutral or anionic non-metal species) Transition metal complexes are important in catalysis, materials synthesis, photochemistry, and biological systems Display diverse chemical, optical and magnetic properties Transition metals Coordination numbers u u Transition metal ions usually form complexes with a well defined number of ligands Complexes with coordination numbers four and six are the most common, although five coordination is also well established 1 Stereochemistry u Four coordinate transition metal ions adopt either tetrahedral or square planar coordination – FeCl42- (tetrahedral), AuCl4- (square planar) u u Six coordinate species are nearly always octahedral Five coordinate species are either trigonal bipyramidal or square base pyramidal Example geometries Ligands u u u Ligands are species (neutral or anionic) bonded to the metal ion They may be attached to the metal through a single atom (monodentate) or bound to the metal through two or more atoms (bidentate, tridentate etc.) Polydentate ligands are called chelating ligands 2 Example chelating ligands Ligands and oxidation state u u u Low oxidation state complexes can be stabilized by using ligands such as cyanide and carbon monoxide Intermediate oxidation state complexes often have ligands such as chloride, ammonia or water High oxidation state complexes usually have fluoride or oxide as ligands Isomerism in metal complexes 3 Geometrical isomers Optical isomerism Linkage isomerism u Some ligand can link to a metal through one of two or more different atoms – NO2- can link through oxygen or nitrogen » [Co(ONO)(NH 3)5]+ and [Co(NO 2)(NH 3)5]+ – NCS- can link through S or N – SO32- can link through O or S 4 Ionization and hydration isomerism u Ionization isomerism – [Co(NH3)5(Br)]SO4 versus [Co(NH3)5(SO4)]Br – or [Pt(NH3)3(Cl)]Br versus [Pt(NH3)3(Br)]Cl u Hydration isomerism – [Cr(H2O)6]Cl 3 versus [Cr(H 2O)5Cl]Cl 2.H 2O versus [Cr(H2O)4Cl 2]Cl.2H 2O Naming of complexes u u Your own responsibility Should be able to name a variety of simple coordination complexes Bonding in metal complexes u u u A number of theories are commonly used 18 electron rule - only applies to low oxidation state complexes, particularly organometallics (things with carbon bonded to the metal center) Crystal field theory - an electrostatic model – pretty good for many complexes u MO theory 5 18 electron rule u Stable low oxidation state complexes are found to have a total of 18 bonding electrons – metal electrons plus lone pairs from ligands u u u u Ni(CO)4 - 4s23d8 and 4 lone pairs Fe(CO)5 - 4s23d6 and 5 lone pairs Cr(CO)6 - 4s23d4 and 6 lone pairs Can be explained using MO theory Crystal Field Theory u Consider the ligands are point negative charges or as dipoles. How do these charges interact with the electrons in the d-orbitals? Octahedral complexes 6 Ligand field splitting u u u The ligand field splitting depends upon the oxidation state of the metal and the ligand type High oxidation state favors large D Effect of ligand is given by the spectrochemical series – I-< Br-< Cl -< F-< OH-< OH 2< NH3< en< CN-< CO High spin and low spin complexes u HS versus LS is determined by the relative size of the ligand field splitting and the pairing energy u If Dis bigger than the pairing energy the complex will be low spin Tetrahedral complexes 7 Square planar complexes Magnetic properties u The loss of degeneracy of the d-orbitals due to crystal field splitting explains why some complexes are diamagnetic and others are paramagnetic – e.g. Ni(CN)42- (square planar) is diamagnetic – but NiCl42- (tetrahedral) is paramagnetic Colors u The colors of most transition metal complexes arises as a consequence of the dorbital splitting 8 Hydration energies Electronic spectra u TM complexes with more than one delectron often show absorption bands at multiple wavelengths u The appearance of multiple bands is due to electron-electron repulsion – not just one transition corresponding to D – may have more than one state for a given electron configuration d2 complexes 9 MO theory u u Crystal field theory explains many phenomena but not for instance the spectrochemical series MO theory offers a better qualitative picture Octahedral complex and sigma donor ligands u Dis the energy gap between nonbonding orbitals and antibonding orbitals p -donor ligands u p -donor ligands(e.g. halide)give smaller Ds 10 p -acceptor ligands u p -acceptor ligands (e.g. CO) give large Ds Carbonyl complexes u Carbonyl ligands are strong p -acceptors – backbonding weakens the C=O bond as seen by vibrational spectroscopy Coordination equilibria u For many complexes in solution there is an equilibrium between different species – Ni(OH2)62+ + 6NH3 = Ni(NH3)62+ + 6H2O u Equilibrium constant depends upon metal and ligands 11 Chelate effect u The equilibrium constants for chelate complexes are usually larger than for similar non-chelate complexes – Ni(OH2)62+ + 6NH3 = Ni(NH3)62+ + 6H2O K1 – Ni(OH2)62+ + 3en = Ni(en)32+ + 6H2O K 2 – K2 > K1 u The effect is due to entropic differences between the two reactions Kinetics u For many complexes the exchange of ligands in the coordination sphere occurs rapidly u However, for complexes of Co(III) and Cr(III) ligand exchange kinetics are slow – labile – said to be kinetically inert 12