Download Electron Counting

Electron Counting
This, for some reason I’ve never worked out, isn’t as easy as it should be. It is, however,
important. It allows us to say things about the stability of organometallic complexes and
make predictions about the structures of complexes and deduce what oxidation state a metal
centre is in. All useful information.
Effectively, electron counting is a method for determining, by visual inspection, how many
electrons a given complex has in its molecular orbitals (its covalent bonding system). There
are certain magic numbers (which I’ll go through below) that tell you your complex is happy.
There are two schools as to how to go about determining this, the ionic method and the
covalent method. Both must give the same answer, so although you only really need to ever
use one method, using the other is a good way to double-check your answer.
It’s easier to show than to explain, I think.
Firstly, using both ways, you need to assess which ligands are dative and which aren’t. This
partly comes with experience, but if you’re not confident you can work it out by looking at the
ligand by itself and seeing if the binding atom conforms to the octet rule (remember that?). If
it’s got eight electrons in its outer shell, contributed from itself and from any other atoms
bound to it, it’ll act as a dative ligand; if it’s got seven or fewer, it’ll be negatively charged.
Then, using both methods in parallel:
ionic:
covalent:
So, in essence, in the ionic method, you determine which bonds naturally break into ionic
pairs (in this, the Pd–Aryl and Pd–I bonds) and which are purely dative (donation from the
phosphorous lone pair). This gives you the charge on the metal centre (positive, 2+). From
this, looking at your periodic table, you can determine how many outer shell electrons this
metal has (eight), and add this number to the number of electrons the ligands each give (in
this case, each gives two electrons).
In the covalent method you treat all the bonds as covalent and determine which bonds are
dative and which have electrons contributed from both the ligand and the metal. The dative
ligands give 2 electrons, whilst the other ligands give one electron. You count how many
electrons the metal has assuming it to be uncharged (i.e. exactly as the periodic table has it,
in this case 5s24d8, ten electrons).
You do also need to take account of the overall charge of the complex - if we look at Zeise’s
salt again, you can see how the complex itself is negatively charged as it needs to be
balanced by a potassium cation.
You only need to count the complex itself, without the counterion. Hence, as another
example:
ionic:
covalent:
If you don’t quite understand this, there’s a very good page online that goes through this
again, has a few more worked examples and some opportunities for you to try yourself. It
also gives a very useful table which shows many of the most common ligands and how
many
electrons
they
donate
in
the
ionic
See: http://www.ilpi.com/organomet/electroncount.html
and
covalent
counting
method.
Task 1
Provide electron-counts (using both the ionic and covalent methods) for the following
complexes:
[Hint - use the ligand table in the above link to help with any unfamiliar ligands]
Eighteen/sixteen electron rule
Right, you now know how to count electrons. This is good, though it probably seems pretty
useless. It isn’t. It’s very important in combination with the eighteen and sixteen electron
rules.
These rules come from the molecular orbital analysis I briefly talked about earlier but didn’t
actually show. They are what is left when you boil down all the group theory stuff and get to
the actually useful information. They’re used to say, in essence, that a given complex is
stable based on how many electrons it has in its bonding system. The (very straightforward)
rules are as follows:
The sixteen electron rule: if a complex has a square-planar geometry and has 16 electrons,
it is stable.
The eighteen electron rule: if a complex has any geometry other than a square-planar
geometry and has 18 electrons, it is stable.
Easy. Chemistry’s at its best when the maths it requires can be done with only fingers and
toes. Some common inorganic complex geometries are shown in figure 3.
Figure 3 - Common inorganic complex geometries - left to right: square-planar, tetrahedral,
octahedral, trigonal bipyramidal, square-based pyramidal
Think of the sixteen and eighteen electron rules a bit like the octet rule that you learn at
GCSE, except applied to organometallic systems. They’re pretty much that simple, but
they’re very powerful. Just one quick example of their use is in rationalising why some
complexes exist in square-planar configurations (e.g. Zeise’s salt) when a tetrahedral
configuration would allow them to spread their ligands out more (more stable in a purely
ionic compound).
These rules also allow you to determine what oxidation state the metal at the centre of an
organometallic complex is in. Just assume the complex to have eighteen electrons (or
sixteen if it’s square planar), use the ionic method to strip away the electrons contributed by
the ligands, account for the overall charge, then work out how many electrons the metal has
to have and how many fewer (or even more) that number is than the periodic table gives.
Easy. (Disclaimer - this only works for truly covalent organometallic compounds, the maths
doesn’t work on some complexes like the titanium complex shown in Figure 1
Task 2
o
Why might it be that the above method only works for truly covalent organometallic
compounds?
o
Comment on the stability of all the complexes you’ve electron-counted so far in this
work sheet.