Transition Metals: Complex Ions

Edexcel A level Chemistry (2017)

Topic 15A: Principles of transition metal chemistry

Here are the learning objectives relating to complex ions:

15/4. To know what is meant by the term ‘ligand’.

15/5. To understand that dative (coordinate) bonding is involved in the formation of complex ions.

15/6. To know that a complex ion is a central metal ion surrounded by ligands.

15/11. To understand the meaning of the term ‘coordination number’.

15/12. To understand why H2O, OH− and NH3 act as monodentate ligands.

15/13. To understand why complexes with six-fold coordination have an octahedral shape, such as those formed by metal ions with H2O, OH− and NH3 as ligands.

15/14. know that transition metal ions may form tetrahedral complexes with relatively large ligands such as Cl^—.     

Complex Ions and Transition metals.

First Row transition metals form complex ions.

Complex ions are ions that contain a central metal ion surrounded by particles: anions or neutral molecules that are dative covalently bonded to the metal ion as in the copper aqua complex below:

These particles: anions or neutral molecules are called ligands.

Examples of ligands include:  

water H₂O, ammonia NH₃, hydroxide ions OH^—.

For a particle to act as a ligand, it must be capable of donating a pair of electrons to the central metal cation to form a dative covalent bond. 

If the ion or neutral molecule has a lone pair available for donation it is a Lewis Base. 

If the ion or neutral molecule has only one lone pair available and makes only one dative covalent bond with the central metal ion, it is called a monodentate ligand. 

Water, ammonia and hydroxide ions are small ligands, all three are Lewis bases and usually form octahedral complexes with transition metal ions. 

An octahedral complex has six ligands bonded to the central metal cation as in the cobalt complex below:

Chloride ions and other halide ions are larger ligands and only four will fit round a transition metal ion.  These usually make tetrahedral or square planar complexes.  See the copper chloro complex below.

The number of ligands that fit round the central metal cation is called the coordination number of the complex ion. 

So here is the question: how can a central metal cation accommodate the extra 12 electrons of the six ligands surrounding it?

Those 12 electrons must be shared with the central metal ion and accommodated in its electron subshells/orbitals. 

In the case of the copper(II) ion with 9 electrons in its outer shell those 12 electrons must be accommodated in subshells from the fourth electron shell. 

This can be shown quite simply using the electrons in boxes model:

The situation is not as simple as it seems on the diagram since each dative covalent bond is identical to the other—they are all degenerate—so that will mean that the 4^th energy level subshells will hybridise to form the appropriate 6 molecular orbitals.  

At least the diagram suggest at this level: college level or A level, how the metal ion can accommodate the extra electrons without us going into and discussing the complexities of crystal field theory.