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Transition Metals and Coordination Chemistry



Transition metals consist of the elements in groups 3–12 of the periodic table (Sc through Zn). In their elemental states, they display metallic luster (shiny appearance) and are good conductors of heat and electricity. Transition metals generally act as Lewis acids, accepting electron pairs from ions or polyatomic compounds (molecules or ions) that often function as Lewis bases called ligands. These Lewis acid/base donor-acceptor interactions lead to the formation of covalent bonds between transition metals and ligands. These are called coordinate covalent bonds or dative bonds. The entire molecule (or ion) is called a coordination complex. The simplest transition metal complexes contain a central metal center, usually a cation, that has one or more covalent bonds to several ligands. The three-dimensional shape of a transition metal complex can often be predicted using valence shell electron repulsion theory concepts. The metal ion with attached ligands is a coordination complex, and its shape depends on how ligands bond. The same ligands can bond with the central metal ion in different geometries, forming isomers, with distinct properties. Hemoglobin and chlorophyll are examples of two coordination complexes found in biological systems.

The crystal field theory description of bonding is a simple electrostatic model that is used to predict numbers of unpaired electrons and in some cases the colors of transition metal complexes. Electrons occupy nnd orbitals according to Aufbau filling principles. Most complexes maximize the number of unpaired electrons due to electron-electron repulsion; pairing of electrons is generally avoided unless all other nnd orbitals contain one (or more) electrons, according to Hund’s rule (electrons first singly occupy empty orbitals before pairing in an orbital).

At A Glance