Professor Dave here, let’s look at coordination compounds.
Normally, when you see stock images of chemistry laboratory equipment, it’s filled with a
variety of brightly-colored liquids.
This is very misleading, as most chemistry does not look like this, it’s just to make
the image pop.
But there is one exception.
The ions of certain transition metals and the complexes they form can have very vivid colors.
What are these complexes?
They are called coordination compounds, and they make some fascinating structures, so
let’s learn about them now.
Coordination compounds are compounds in which a central metal ion acts as a Lewis acid,
and a certain number of ligands, each with some electron excess, act as Lewis bases and
coordinate to the central atom to form a complex.
These ligands can be atoms, molecules, or ions, but they usually have a lone pair that
can coordinate to the electron-deficient metal ion to form what is called a coordinate covalent
bond, sometimes also called a dative bond.
These are different from the ionic and covalent bonds that we are used to.
With coordinate covalent bonds, both of the electrons in the bond come from a Lewis base,
and are being donated to a metal center, like the lone pairs on each of these water molecules,
which interact with the central metal ion to form this octahedral complex.
If we want to show the coordination sphere, or the metal ion plus its ligands, using our
standard notation, we will do this by placing brackets around the complex, as is shown here.
Outside of these brackets will be listed any formal charge exhibited by the complex.
The coordination number is the number of donor atoms that are bound to the central atom in
the complex, so that would be two, four, and six, for these three complexes respectively.
With these situations, all the ligands are monodentate ligands, meaning they interact
with the central atom through one atom only.
Ligands can also be bidentate, such as ethylenediamine, whereby both the nitrogen atoms on each ligand
will interact with the metal atom.
So in this case, even though there are three ligands, the complex has a coordination number of six.
There can even be polydentate ligands, where interaction with the metal atom occurs in
many places.
These can sometimes be called chelating ligands.
Hemoglobin is such a ligand, because all four nitrogen atoms in the heme are coordinated
to the iron atom.
So technically this is a tetradentate ligand.
Coordination compounds are fascinating structures that display a wide variety of geometries.
What determines the type of geometry that a coordination compound will exhibit?
More than anything else, that would be the coordination number of the compound.
Once again, this is the number of donor atoms that are interacting with the central metal
atom, or in other words, the number of coordinate covalent bonds that are being made to the
central atom in the complex.
There are a few types of geometry that we will see over and over again, as they are
the most common.
These would be octahedral, for complexes with a coordination number of six, as well as tetrahedral
and square planar, both for complexes with a coordination number of four.
But we can have coordination numbers well above six, and these will generate some interesting
geometries, like pentagonal bipyramid, square antiprism, and dodecahedral.
Many of these do not follow the geometries expected from the VSEPR model, due to details
regarding nonbonding d electrons.
As we mentioned, a few coordination numbers offer more than one option for geometry, and
the geometry adopted by the complex will depend on the precise set of ligands involved.
Octahedral complexes, with a coordination number of six, are very common.
In this geometry, all the bond angles are ninety degrees.
Going back to compounds with a coordination number of four, once again, both tetrahedral
and square planar geometries are possible.
We can’t predict which will occur by VSEPR, we will have to use crystal field theory in
order to predict this, and the factor that is largely responsible is the number of d
electrons in the valence shell for the central metal atom.
But we will get to that in a moment.
Beyond understanding the geometry of coordination compounds, we also have to be able to name
them, so that we can communicate concepts about these molecules.
The rules for the nomenclature of coordination compounds are as follows.
First, if a coordination compound is ionic, we will name the cation first, and then the
anion, in accordance with typical nomenclature rules.
Second, we will always name the ligands first, and we will do so in alphabetical order.
Once these have all been named, then we will name the central metal atom.
So what are the names of these ligands?
Well for neutral ligands, it’s easy.
These will just be the normal name of the molecule that is acting as a ligand.
However, there are some common exceptions to this rule, including water, ammonia, carbon
monoxide, and nitrogen monoxide, which as ligands will be called aqua, ammine, carbonyl,
and nitrosyl.
For anionic ligands, these will end in the letter O, and the names will be quite intuitive,
like fluoro, bromo, and the rest of the halogens, as well as others like cyano, or hydroxo.
The third and final rule is that if a particular ligand appears more than once in a coordination
compound, we must specify this using a special prefix indicating the number of times it appears.
Sometimes these prefixes are modified if the ligand itself begins with one of those prefixes,
or a vowel, so as to remove ambiguity, resulting in the prefixes bis, tris, and so forth.
Now to put things together, when the complex as a whole is a cation, we will list the metal
by the name of its element, followed by a roman numeral in parentheses to indicate its
oxidation state.
So as we can see in these examples, the ligands are listed first, with prefixes as necessary
to indicate the number of them, then the metal ion with its oxidation state, and then any
counterions to the complex if applicable.
When the complex as a whole is neutral, we will follow the same rules as for cations,
and in this case there necessarily will not be any counterions to list.
When the complex is an anion, however, we will have to add the suffix “ate” to the
metal, giving us things like platinate instead of platinum, and so forth.
With stannate, we will see that sometimes the latin name is used for the metal atom
instead of the name of the element itself.
Let’s try an example, Na2PtCl6.
First let’s assess the charges.
Sodium ions are plus one, so that’s a total of plus two, meaning the complex itself has
a minus two charge as a whole.
Each chloro group is minus one, so six minus ones will need a plus four on the platinum
to get to minus two, so platinum must have an oxidation state of plus four.
Now we just list everything.
Six chloro ligands gives us hexachloro, then platinate with the roman numeral four, remembering
that because it’s an anion we use the suffix “ate”, and then sodium at the beginning
to list the counterions.
So that’s a little bit of introductory information regarding coordinate compounds.