Chemical Bonding in Coordination Compounds

Well hello, again.

Today I'm going to talk about the models used to describe the different type of bondings that occur in coordination compounds. There are various theories that historically have been used to try to describe the interaction that appear:

• Valence Bond Theory

• Crystal Field Theory

• Ligand Field Theory

• Angular Overlap Theory

Before we start I will define some words that will be going to appear along the text: 

Complex: Usually a complex is defined by a number of ligands that are bonded to a central atom (usually a metal) that have a defined geometry. 

Ligand: An atom or molecule bonded to a metal center. 

Coordination number: In complexes the coordination number is simply the number of species that are bonded to the central cation (the metal). For example [Fe(CN)₆]^3-, the iron has a coordination number of 6.

I'm inserting here a table with the common geometries from the VSEPR model, so that you can see them if they are not familiar to you: 

So.

We've already talked about the VB theory, but now I'm going to talk about its applications to coordination compounds. So, Pauling stated the geometry of the coordination complex should depend on the number of hybridization processes that occur. This will depend on the metal's nature, the number of ligands and the coordination number of the metal. For example: 

For tetrahedral geometries, for example a  [CoCl4]2- will (according to Co electronic configuration) 3 unpaired electrons, this unpaired electrons make the complex paramagnetic, that is, susceptible to external magnetic fields. The bonds with the metal will be formed in the 4s and 4p orbitals. This orbitals are, according to Pauling, hybridized. 

The red box shows the hybridized orbitals.

Now for square planar geometries:

You can observe how the hybrid orbitals include one of the 3d orbitals. 

Now octahedral geometries can have two type of compounds according to Pauling. High spin and low spin. You can see the analogy in the next electronic configurations of two different complexes: 

The difference is that the CN complex takes 3d orbitals and the F complex takes the 4d orbitals.

• High spin: takes 4d
• Low spin: takes 3d

How can you predict which complexes will be high or low spin? It is according to the nature of the ligand. If you can say that the ligand has a strong field nature, then it is low spin. Thus if the ligand is of weak field nature you can say it will be high spin. 

But how do you know that?

Answer: The spectrochemical series

I− < Br− < S2− < SCN- < Cl− < NO3− < N3− < F− < OH− < C2O42− ≈ H2O < NCS− < CH3CN < py (pyridine) < NH3 < en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-phenanthroline) < NO2− < PPh3 < CN− ≈ CO

Where do strong ligands start? This is a really good question, it is not really well stablished and basically after ammonia you can consider them strong.

And so, according to the hybridization you can predict its magnetic properties. If there are unpaired electrons the complex is paramagnetic and if there are no unpaired electrons it is called diamagnetic.

Also, something important:

• Tetrahedral geometries are always high spin, its independent of the nature of the ligand.
• Square planar are always low spin (same reason).

So, what have we learned?

We now know the concepts of low and high spin, we've presented the spectrochemical series for the determination of the nature of a ligand and with all this we can predict how the electrons will distribute in a complex giving arise to paramagnetism or diamagnetism.

Well, that's all, folks!
See ya, in the next entries I'm going to talk about the other theories I mentioned.

References:

Atkins, P. . (2014). Química Inorgánica. México: Reverté.

Cotton, F. . (1999). Advanced Inorganic Chemistry. USA: Wiley.