Overview of Structure and Reactivity of Transition Metals

Transition metals are a group of special metals that are located in the periodic table between the alkaline earth metals and the nonmetals.  They have unfilled d-shells with electrons.  By definition transition metals are those metals in the periodic table with incompletely filled d-shells with electrons.  They are located above the lanthanides and the actinides, which have f-shell that are occupied with electrons. 

Transition metals are governed by their metal complexes, chemistry, and their organometallic chemistry.  They differ in their bonding to ligands than to other atoms by a special type of bonding which is called dativ bonding.  Normally transition metals can form a coordination number of 6 but also a coordination number of 4 is observed. 

Transition metals in the periodic table begin with the scandium column and end in the zinc metal column.  The zinc metal column has filled d-shell with electrons.  Therefore some authorities consider the zinc group to be not part of the transition metals due to this fact.  The zinc metal column behaves chemically like the magnesium column in that both are governed by the +2 oxidation state. 

The ionic radius of transition metals decreases as we go across a row in the periodic table.  The ionic radius also increases as we go down a column in the periodic table.  This is so due to the increase in nuclear charge as we go from left to right in the periodic table. 

The bonding in transition metals complexes is governed by the d-orbitals on the metals.  The stability of transition metal complexes is determined or governed by the 18 electron rule.  Complexes which have 19 electrons are easily oxidized and they are a good reducing agents.  17 Electrons complexes are on the other hand easily reduced and are therefore good oxidizing agents. 

D orbitals electron transitions govern also the spectroscopic  properties of the metal complexes.  D-d electron transitions confer a color to the solution that contains the complex.  Also the magnetic properties of the metals are confered by the d electrons configuration.

Metal complexes that have 5-d orbitals are split into two sets of orbitals in octahedral field.  tThese two levels contain three and two d orbitals respectively.  The set of three d orbitals lies lower in energy than the set of 2 d orbitals.  The set of three d orbitals can accept up to 6 electrons while the set of 2-d orbitals can accept up to 4 electrons. 

D10 electron configuration thus has 10 electrons in the 5 d orbitals in which all the d orbitals are filled.  This configuration is diamagnetic and is a relatively stable configuration.  D-5 configuration is also a stable configuration.  It can be high spin or low spin. 

The ligand that is attached to the metal can confer extra stability to the complex by further splitting the d 0rbitals energy levels.  Ligands that have high splitting of these two energy levels such as CO and  -CN  give the complex a diamagnetic property.  On the other hand ligands that do not participate in backbonding have high spin complexes such as water complexes.  This is so due to the low energy splitting of the 2-d energy levels.

In coordination number of 4 the complexes of transition metals there is two possible arrangements for the ligands around the metal.  These are tetrahedral arrangement and square planar arrangement.  The d-orbitals splitting in tetrahedral arrangement is similar to that of the octahedral arrangement of the d orbitals except that the d levels are reversed.  The diamagnetic and paramagnetic properties of the complexes also depend on the ligand attached to the metal.  In square planar configuration diamagnetism is predominant in all complexes.  D8 is typical for these complexes.  Nickel groups possess this type of coordination.