Figure 23.2 Radii of transition metals as a function of group number.

First Ionization Energies of Transition Metals The first ionization energy increases gradually from left to right on the periodic table. 5d 3d 4d Figure 19-4. The first ionization energies of transition metals show gradual upward trends across each row of the periodic table.

Oxidation States of 3d Transition Metals Element group Sc 3 Ti 4 V 5 Cr 6 Mn 7 Fe 8 Co 9 Ni 10 Cu 11 Zn 12 Oxidation state Valance configuration +1   d 4 d 5 d 6 d 7 d 8 d 9 d 10 +2 d 1 d 2 d 3 +3 d 0 +4 +5 +6 +7 Key +3 +4 *Table lists the configuration of the ion corresponding to each observed oxidation state. The most important oxidation states of each element are color screened. Table 19-1. Oxidation states displayed by the 3d transitions metals*

Formation of Coordinate Covalent Bonds A ligand donates a lone pair of electrons to form a bond to a metal. Ex. The Ni–N bonds in [Ni(NH3)6]2+ form by overlap of the lone pair sp3 orbital on the nitrogen atom with an empty valence orbital on the metal. Donor Metal

Two Ni2+-Ligand Complexes Both water and ammonia form six covalent bonds with Ni2+, resulting in octahedral geometry. Figure 19-5. The Ni2+ cations in [Ni(H2O)6]2+ and [Ni(NH3)6]2+ bond to six ligands, one at each vertex of an octahedron.

Colors of Two Ni2+-Ligand Complexes Figure 19-6. Solid nickel(II) sulfate, nearly colorless, dissolves in water to give a green solution containing [Ni(H2O)6]2+ cations (left). The addition of ammonia produces a blue solution of [Ni(NH3)6]2+ cations (right). Solvent evaporation gives a blue-violet precipitate of [Ni(NH3)6]SO4. [Ni(H2O)6]2+ [Ni(NH3)6]2+

Figure: 24-01

Coordination Number - Two Complexes with coordination number two always adopt linear geometry about the metal cation. Figure 19-7. Complexes with coordination number two always adopt linear geometry about the metal cation.

Coordination Number – Four Tetrahedral and Square planar Figure: 24-03a,b

Coordination Number – Six Octahedral Figure: 24-04a,b

Figure: 24-T02

Figure: 24-05

Bidentate Ligands Figure: 24-06

Figure 23.12 The EDTA4– ligand (top) and the complex ion [Co(EDTA)]– (bottom).

Figure 23.17 Ferrichrome.

Figure 23.18 The iron-transport system of a bacterial cell.

Figure: 24-10

Heme Iron Oxygen-carrying component of blood. Planar structure. Multi-ring structure of C and N atoms. Extensive delocalized π system. Binds one Fe2+ cation at its center. Iron Figure 19-1. Heme is a planar, multi-ring structure of C and N atoms, binding an Fe2+ cation at its center.

Figure: 24-11

Figure: 24-12

Figure: 24-13

Figure: 24-14

Repulsion of Ligand Electrons and Metal Electrons is Greatest with Overlap dx2–y2 points directly toward the ligands. Overlap results in increased repulsion. dxy points between the ligands. Lack of overlap results in less repulsion. y  y  Figure 19-12. The dx2–y2 orbital is concentrated along the bond axes, pointing directly at the ligand lone pairs. The dxy orbital points between the bond axes. There is less electron-electron repulsion for dxy than there is for dx2–y2. x  x  dxy orbital  dx2–y2 orbital

The Five d Orbitals Interacting with an Octahedral Set of Ligands Figure 19-13. The five d orbitals and their relationship to an octahedral set of ligands. Whereas two orbitals point directly at the ligands, the other three orbitals point between the ligands.

The Crystal Field Level Diagram for Octahedral Coordination Complexes Electron-cation attraction stabilized all five d orbitals. Electron-electron repulsion destabilizes the five d orbitals by different amounts. Figure 19-14. The crystal field energy level diagram for octahedral coordination complexes.

Crystal Field Splitting Energy The difference in energy between the eg and t2g sets. Symbolized by the Greek letter, Δ. eg dx2–y2 dz2 Δ t2g dxy dxz dyz

The Spectrochemical Series The spectrochemical series lists the common ligands in order of increasing ability to split the energies of the t2g and eg subsets of orbitals. The ranking of ligands is influenced most strongly by the donor atom: Generally decreases across Row 2 of the periodic table. Generally decreases down the halogen column. Molecular orbital theory is best used to explain the trend. Figure 19–17. The spectroscopic series ranks ligands in order of their ability to increase the energy difference between the t2g and eg subsets of the valence d orbitals.

Relationships Among Wavelength, Color, and Crystal Field Splitting Energy (Δ) Wavelength (nm) Color absorbed Complementary color Δ (kJ/mol) >720 Infrared Colorless <165 720 Red Green 166 680 Red-orange Blue-green 176 610 Orange Blue 196 580 Yellow Indigo 206 560 Yellow-green Violet 214 530 Purple 226 500 239 480 249 430 279 410 Lemon-yellow 292 <400 Ultraviolet >299 Table 19-6. Relationships Among Wavelength, Color, and Crystal Field Splitting Energy (Δ)

Colors of Cr3+ Coordination Complexes The colors of Cr3+ coordination complexes depend on the magnitude of the crystal field splitting energy. Higher Δ, shorter λ. The spectrochemical series indicates the relative magnitude of Δ. Figure 19-22. The colors of Cr3+ coordination complexes depend on the magnitude of the crystal field splitting energy.

Figure: 24-18

Figure: 24-20

Figure: 24-21

Figure: 24-22