Posted by: Mark Foreman | February 25, 2012

Nitrosyl complexes

Dear Reader,

Since the time I started at University a great deal of stuff has been written on NO, it was the subject of a noble prize in the early 1990s. One of the key things about NO is its ability to bind to transition metal centres. Back in 1993 I had a go at making a couple of nitrosyl complexes. This was in the inorganic chemistry lab class at Imperial College where I made iron and nickel nitrosyl complexes. One of the important things is that both bent and linear nitrosyls exist.

Now if we consider a classic ligand which binds via a sp carbon (cyanide), then when I did a survey of all the cyanides which bind via the carbon to a single transition metal then I found the metal-carbon-nitrogen angles were all very close to 180 degrees. Here is a histogram below showing the results.

Metal-carbon-nitrogen bond angle for transition metal cyanides

I then repeated the same search for transition metal complexes of pyridine, you should be able to see that the metal-nitrogen-carbon angle for a typical pyridine complex is very close to 120 degrees. I choose pyridine as it has a well defined sp2 nitrogen which binds to metals.

Metal nitrogen carbon bind angles in transition metal complexes of pyridine

Now if we try the same thing with nitrosyl complexes then we will see how while the majority of nitrosyls have a metal-nitrogen-oxygen angle of close to 180 degrees, a sizeable minority of them have metal-nitrogen-oxygen angles of about 120 degrees. Here is the graph below.

Metal nitrogen oxygen angles in nitrosyls

So it is clear that the nitrosyl ligands have more than one coordination mode. If we assume that the nitrosyl nitrogen is a sp2 nitrogen then it will have three sp2 orbitals which make sigma bonds. We will have the sigma bond to the oxygen, the lone pair which points out into space and a sp2 orbital containing one electron. This last orbital can form a sigma bond to a metal. As this last orbital has one electron then it will contribute one electron to the valence count of the metal.

Below I am showing the sp2 orbitals of a nitric oxide molecule if we draw it as having a sp2 nitrogen, in the second picture the singlely occupied molecular orbital of the nitric oxide is forming a sigma bond with the metal which is in the rather fetching purple. Note that as a good Royal College of Science man I like black, white and purple.

Bent nitrosyl ligand

Now if we start with the assumption that the nitrosyl nitrogen in our nitric oxide is a sp nitrogen then we will get a molecule which has a lone pair on nitrogen pointing along the same axis as the N-O bond. The molecule will also have two sets of pi systems at 90 degrees to each other.

nitric oxide for a linear nitrosyl

Now if we make a diagram of the molecular orbitals of the nitric oxide then it will look like this. It is important to note that the pi antibonding orbitial now has a single electron in it.

MO diagram for nitric oxide

Now when the nitric oxide bonds to the metal for form a linear nitrosyl the lone pair on the nitrogen will bind to the metal to form the sigma bond, this gives the metal two electrons. When the occupied d orbitals on the metal interact with the pi* orbitals of the nitric oxide then the electron in the pi* orbitals will be shared with the metal in the new pi bonds between the metal and the nitrosyl, thus in this way one extra electron is shared with the metal.

As a result the metal gets a total of 2+1 electrons to add to its valence count, which makes three.


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