Chemistry of Transition Metals Part 1. General Considerations 1 •Filling of 3d, 4d, and 5d shells •In s- and p-block, electrons added to outer shell. •In d-block, electrons added to penultimate shell expanding it from 8-18 2 •Most elements have incompletely filled d-shell (interesting properties) Transition elements 1. Metals 2. Almost all: HARD, STRONG, High m.p., b.p. 3. Conduct heat & electricity 4. Form Alloys 5. Show variable oxidation states 6. At least one of the ions & compounds colored. 7. Form paramagnetic species because of partially filed shells 8. Form coordination compounds (complexes) and organometallic compounds. 3 Variable oxidation state 1. Increase in the number of oxidation states from Sc to Mn. All possible states exhibited by only Mn. Sc +3 2. Decrease in the number of Ti +1 +2 +3 +4 oxidation states from Mn to Zn, V +1 +2 +3 +4 +5 due to the pairing of d- electrons occurs after Mn Cr +1 +2 +3 +4 +5 +6 (Hund's rule). Mn +1 +2 +3 +4 +5 +6 +7 3. Stability of higher oxidation Fe +1 +2 +3 +4 +5 +6 states decreases along Sc to Zn. Mn(VII) and Fe(VI) are Co +1 +2 +3 +4 +5 powerful oxidizers. Ni +1 +2 +3 +4 4. Down the group, the stability of high oxidation states Cu +1 +2 +3 increases (easier availability of Zn +2 both d and s electrons for ionization). 4 Variable oxidation state 5 Color 1. The d-orbitals of the metal interact with the electron cloud of the ligands in such a manner that the d-orbitals become non-degenerate. When the d-level is not completely filled, it is possible to promote and electron from a lower energy d-orbital to a higher energy d- orbital by absorption of a photon of electromagnetic radiation having an appropriate energy (d-d transitions). 2. Metal to Ligand and Ligand to Metal Charge Transfer (LMCT and MLCT) transitions (KMnO4) 3. Electromagnetic radiations in the visible region of the spectrum often possess the appropriate energy for the above transitions. 6 Color 7 8 9 10 11 12 13 14 Self Study Isomerism Naming of complexes 15 Self Study 16 Part 2. Theories / Concepts Bonding in transition metal compounds -Werner Coordination Theory 18 electron rule -Valence Bond Theory -Crystal field theory -Molecular orbital approach High spin and low spin complexes Spectrochemical series CFSE 17 Jahn-Teller distortions Werner Coordination Theory Werner : 1893 (electron was discovered in 1896) Nobel prize in 1913. In complexes, metal ions show two different types of valency: Primary Valency: Non-directional, is the number of charges on the complex ion. 1. Secondary Valency: Directional, equals to the number of ligands coordinated to the metal. 18 18 electron rule (based on earlier EAN Rule: Sidgwick) Stable low oxidation state complexes are found to have a total of 18 bonding electrons metal electrons + lone pairs from ligands = 18 2 8 ! Ni(CO)4 -4s3d and 4 lone pairs 2 6 ! Fe(CO)5 -4s3d and 5 lone pairs 2 4 ! Cr(CO)6 -4s3d and 6 lone pairs The stability of these 18 electron species can be explained using MO theory. Corresponds to filling all the molecular bonding orbitals and none of the antibonding orbitals However, is rule only works for species with metals in a low oxidation state NOT FOR MOST COMPLEXES 19 Valence Bond Theory The idea that atoms form covalent bonds by sharing pairs of electrons was first proposed by G. N. Lewis in 1902. In 1927, Walter Heitler and Fritz London showed how the sharing of pairs of electrons holds a covalent molecule together. The Heitler-London model of covalent bonds was the basis of the VBT. The last major step in the evolution of this theory was the suggestion by Linus Pauling that atomic orbitals mix to form hybrid orbitals, such as the sp, sp2, sp3, 3 2 3 dsp , and d sp orbitals. 20 VBT – Assumptions / Features Coordination compounds contain metal ions, in which ligands form covalent-coordinate bonds to the metal. Ligands must have a lone pair of electrons. Metal should have an empty orbital of suitable energy available for bonding. Atomic orbitals are used for bonding (rather than molecular orbitals) This theory is useful to predict the shape and stability of the metal comples. Limitations: (1) Does not explain why some complexes are colored and others are not; (2) Does not explain 21 the temp. dependence of the magnetic properties. VBT – Example Co: 3d74s2 Co3+ : 3d64s0 3d 4s 4p 4d Outer sphere complex xx xxxx xx xx xx Reactive / labile High spin Paramagnetic sp3d2 3d 4s 4p 4d Inner sphere complex xx xx xx xxxx xx Stable Low spin Diamagnetic d2sp3 22 Crystal Field Theory (Text : JD Lee; pp.204-222) •This theory (CFT) largely replaced VB Theory for interpreting the chemistry of coordination compounds. •It was proposed by the physicist Hans Bethe in 1929. •Subsequent modifications were proposed by J. H. Van Vleck in 1935 to allow for some covalency in the interactions. These modifications are often referred to as Ligand Field Theory. •For a review on the evolution of bonding models see: C. J. Ballhausen, J. Chem. Ed. 1979 56 194-197, 215-218, 357-361. 23 CFT-Assumptions •The interactions between the metal ion and the ligands are purely electrostatic (ionic). •The ligands are regarded as point charges •If the ligand is negatively charged: ion-ion interaction. If the ligand is neutral : ion-dipole interaction •The electrons on the metal are under repulsive from those on the ligands •The electrons on metal occupy those d-orbitals farthest away from the direction of approach of ligands 24 Symmetric Field •The 5Xd orbitals in an isolated gaseous metal are degenerate. •If a spherically symmetric field of negative charges is placed around the metal, these orbitals remain degenerate, but all of them are raised in energy as a result of the repulsion between the negative charges on the ligands and in the d orbitals. E metal ion in a spherical negative field Mn+ metal ion spherical in free state negative field (vacuum) 25 Octahedral Field •If rather than a spherical field, discrete point charges (ligands) are allowed to interact with the metal, the degeneracy of the d orbitals is removed (or, better said, lifted). The splitting of d orbital energies and its consequences are at the heart of crystal field theory. eg E n+ M t2g •Not all d orbitals will interact to the same extent with the six point charges located on the +x, -x, +y, -y, +z and -z axes respectively. •The orbitals which lie along these axes (i.e. x2-y2, z2) will be destabilized more that 26 the orbitals which lie in-between the axes (i.e. xy, xz, yz). Octahedral Complexes 27 CFT-Octahedral Complexes 2 2 2 •For the Oh point group, the x -y , z orbitals belong to the Eg irreducible representation and xy, xz, yz belong to the T2g representation. ∆ •The extent to which these two sets of orbitals are split is denoted by 0 or alternatively 10Dq. As the baricenter must be conserved on going from a spherical field to an octahedral field, the t2g set must be stabilized as much as the eg set is destabilized. ∆ = + 0.6 o = − 0.4 ∆ o 28 Illustration of CFSE 3+ 1 − [Ti(H2O)6] : a d complex and the e occupies the lowest energy orbital, i.e. one of the three degenerate t2g orbitals. The purple colour is a result of the absorption of light which results in the 1 0 0 1 promotion of this t2g electron into the eg level. t2g eg –> t2g eg The UV-Vis absorption spectrum reveals that this transition occurs with a maximum at 20300 cm-1 which ∆ corresponds to o 243 kJ/mol. (1000 cm-1 = 11.96 kJ/mol or 2.86 kcal/mol or 0.124 eV.) 29 ∆ Typical 0 values are of the same order of magnitude as the energy of a chemical bond. •What happens for more than 1 electron in d orbitals? •The electron-electron interactions must be taken into account. •For d1-d3 systems: Hund's rule predicts that the electrons will not pair and occupy the t2gset. •For d4-d7 systems ( there are two possibilities): Either put the electrons in the t2g set and therefore pair the electrons (low spin case or strong field situation. Or put the electrons in the eg set, which lies higher in energy, but the electrons do not pair (high spin case or weak field situation). •Therefore, there are two important parameters to consider: The Pairing energy (P), and the eg -t2g Splitting (referred ∆ to as 0, 10Dq) •For both the high spin (h.s.) and low spin (l.s.) situations, it is possible to compute the CFSE. 30 ∆ o vs Pairing Energy ∆ 1 -1 Complex Config. o, cm- P, cm spin-state 2+ 6 [Fe(OH2)6] d 10,400 17,600 high-spin 4- 6 [Fe(CN)6] d 32,850 17,600 low-spin 3- 7 [CoF6] d 13,000 21,000 high-spin 3- 7 [Co(NH3)6] d 23,000 21,000 low-spin 31 ∆ o is dependent on: •Nature of the ligands •The charge on the metal ion •Whether the metal is a 3d, 4d, or 5d element Ligands which cause a small splitting are Weak field ligands -1 (∆ο in the range 7000 - 30000 cm ) and those cause a large splitting are Strong field ligands (CFSE typically > 30000 cm-1) Spectrochemical Series − − 2− − − − − − 2− I < Br < S < SCN < Cl < N3 , F < urea, OH < ox, O < − − − − H2O < NCS < py, NH3 < en < bpy, phen < NO2 < CH3 , C6H5 < CN− < CO.