1 19.1 The Transition Metals: A Survey 19.2 The First-Row Transition Metals 19.3 Coordination Compounds 19.4 Isomerism 19.5 Bonding in Complex Ions: The Localized Electron Model 19.6 The Crystal Field Model 19.7 The Molecular Orbital Model 19.8 The Biological Importance of Coordination Complexes Chapter 19. Transition Metals and Coordination Chemistry Vanadium metal (center) and in solution as V2+(aq), V3+(aq), VO2+(aq), and VO2+(aq), (left to right). Figure 19.1: Transition elements on the periodic table 2 19.5 Bonding in Complex Ions: The Localized Electron Model ? Recall formation of hybrid orbitals for bonding in molecules (chapt. 14). ? Formation of coordinate covalent bonds between a metal ion and a ligand (L) is through a Lewis acid-base interaction. M L = M L empty metal ion hybrid atomic orbital lone pair on the ligand in a hybrid atomic orbital Co3+ is d2sp3 hybridized in the complex ion Co(NH3)63+ Figure 19.19: Set of six d2sp3 hybrid orbitals on Co3+. Figure 19.20: Hybrid orbitals required for tetrahedral, square planar, and linear complex ions. ? Co2+ is sp3 hybridized in tetrahedral CoCl42- complex ion. ? Ni2+ is dsp2 hybridized in the square planar Ni(CN)42- complex ion. ? Ag+ is sp hybridized in the linear Ag(NH3)2+ ion. 3 Bonding in Complex metal ions ? Limitations of the LE Model of Bonding in complex metal ions: 1. The VSEPR model for predicting structure does not work for complex ions. 2. The LE model cannot predict the properties of the metal complex ion, e.g. magnetism and color. ? Magnetic properties and colors ? the crystal field model based on d-orbital splitting 19.6 The Crystal Field Model ? Accounts for the color and magnetic properties of complex ions ? Key assumptions: 1. ligands are approximated by negative point charges. 2. metal ? ligand bonding is entirely ionic. Octahedral Complexes (See Figure 19.21) ? Orientation of the 3d orbitals relative to the point-charge ligands leads to splitting of the 3d orbital energies (?) ? t2g set of orbitals (dxz, dyz, dxy) ? lower in energy ? eg set of orbitals (dz2 , dx2-y2 ) ? higher in energy Strong-field case eg t2g large ? E Weak-field case eg t2g small ? E 4 Figure 19.21: Octahedral arrangement of point-charge ligands and the orientation of the 3d orbitals. Figure 19.22: Energies of the 3d orbitals for a metal ion in an octahedral complex. 5 Figure 19.23: possible electron arrangements in the split 3d orbitals of an octahedral complex of Co3+ ? In a strong-field case (i.e. large ?), the electrons of the metal ion pair in the lower-energy t2g orbitals. ? In a weak-field case (i.e. small ?), the electrons will occupy all five orbitals before pairing occurs. Example 19.4 Fe(CN)63- experimentally known to have one unpaired electron. Does the CN- ligand produce a strong or weak field? Fe3+ has 3d5 electron configuration weak-field (high-spin) case eg t2g small ? strong-field (low-spin) case eg t2g large ? ? CN- is a strong-field ligand toward the Fe3+ ion. ? Spectrochemical series of ligands: CN- > NO2- > en > NH3 > H2O > OH- > F- > Cl- > Br- > I- strong-field ligands (large ?) weak-field ligands (small ?) decreasing ? values for a given metal ion 6 ? The magnitude of ? for a given ligand increases as the charge on the metal ion increases. Example Account for the magnetic properties of Co(NH3)63+ and CoF63- Co3+ with 3d6 electron configuration in both cases. Co(NH3)63+ diamagnetic eg t2g large ? CoF63- paramagnetic eg t2g low ? experimentally observed. ?Origin of Color in Complex metal ions ? The same d orbital energy splitting explains the colors of complex ions. eg t2g E ?E = hc/?= 1240 (eV)/ ?(nm) if ? = wavelength of light absorbed is in the visible ? The visible spectrum (Figure 19.24) 400 700 nm ?E = 3.1 eV 1.8 eV If ?E < 1.8 eV, the wavelength of the light absorbed will be in the infrared and not visible. eg t2g E ? Octahedral complexes of Cr3+ 3d3 7 Figure 19.24: Visible spectrum Figure 19.26: The complex ion Ti(H2O)63+ absorbs light and becomes excited. 8 Complexes with Other Coordination Geometries ? Octahedral complexes (already considered) ? Tetrahedral complexes ? Square planar complexes ? Linear complexes all tetrahedral complexes produce the weak-field case. e.g. CoCl42-, 3d7 (Example 19.6) For the crystal field diagrams of square planar and linear complexes See Figure 19.29 eg t2g ? E dxy dxz dyz dz2 dz2-y2 For a given ligand and metal ion: ?tet = 4/9 ?oct Figure 19.27: Tetrahedral and octahedral arrangements of ligands shown inscribed in cubes. 9 Figure 19.28: Crystal field diagrams for octahedral and tetrahedral complexes Figure 19.29: Crystal field diagram for a square planar complex oriented in the xy plane (b) crystal field diagram for a linear complex 10 Complex metal ions ? Bonding in complex metal ions ? the localized electron model is used but it has major limitations: 1. The VSEPR model for predicting structure does not work well for complex ions. 2. The LE model cannot predict the properties of the metal complex ion, e.g. magnetism and color. ? Magnetic properties and colors ? the crystal field model based on d-orbital splitting ?The Molecular Orbital Model of Complex Metal Ions ?Consider the general octahedral complex with formula ML6n+. ? dz2 and dx2-y2 orbitals point at the ligands and thus will form MOs with the ligand lone pair orbitals ? the ?MOs in the complex ion involve dz2 , dx2-y2 , 4s, 4px, 4py, and 4pz orbitals ? dxy, dxz, and dyz orbitals are not involved in ? bonding with the ligands Figure 19.30 Figure 19.31: The MO energy- level diagram for an octahedral complex ion (ML6n+). ? the dxz, dyz and dxy orbitals ( the t2g set) of the metal ion remain unchanged in the complex ion since they do not overlap with the ligand orbitals ? they are nonbonding orbitals. ? the eg* MOs are antibonding orbitals (since they are higher in energy than the atomic orbitals that mix to form them) ? the eg* MOs are composed of dz2 and dx2-y2 atomic orbitals primarily (little or no mixing with ligand atomic orbitals due to the large energy difference btw them) ? the MO model has the same d orbital splitting as the crystal field model. ML6n+ 11 Figure 19.32: MO energy-level diagram for CoF63-, which yields the high-spin (a) whereas Co(NH3)63+ yields low-spin (b). ? the MO model accounts for the magnetic and spectral properties of complex ions (just like the crystal field model). ? the MO model is a more realistic description of the metal-ligand bonding in complex ions. 19.8 The Biological Importance of Coordination Complexes of Metal Ions ? In biological systems, metal ion complexes find diverse applications ? transport and storage of oxygen and other essential elements ? electron transfer agents ? catalysts (enzymes) ? drugs ? etc. 12 Figure 19.33: The heme complex in which an Fe2+ ion is coordinated to four nitrogen atoms of a planar porphyrin ligand. Figure 19.34: Chlorophyll is a porphyrin complex of Mg2+, essential to photosynthesis. 13 Figure 19.35: Representation of the myoglobin molecule Figure 19.36: Representation of the hemoglobin structure. Each hemoglobin stores 4 oxygen (O2) molecules. Pei-Tzu 20.4 Isomerism
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