With our survey of the periodic table complete, we are now going to settle into the d-block and
learn a lot more about transition metals. In the general chemistry series we got an introduction
to coordination compounds, so some of what we will discuss here may be familiar, such as the
concept of a metal atom bonding with ligands. As we know, these ligands can be atoms or molecules,
either neutral or ions. But we are going to go a lot deeper now, in learning about precisely how
metal atoms interact with ligands, and what kinds of ligands we can expect to see.
The first thing we need to discuss is the types of bonds that can occur between a metal and a ligand.
One would be tempted to look at the bond between a metal and a ligand as a normal covalent bond,
which as we know according to valence bond theory is a situation where each of
them contributes one electron to the bond, and the orbital overlap is what creates the bond.
From the perspective of molecular orbital theory, that would look like this,
where each contributes one electron, so both of them sit down here in this sigma bonding orbital.
However, in most situations, a more accurate description of metal-ligand
bonding is the coordinate covalent bond, or dative bond. In a dative bond, the metal acts
as a Lewis acid and the ligand acts as a Lewis base, contributing both electrons to the bond.
If using molecular orbital theory that would look like this, with the ligand usually sitting at a
lower energy than the metal, and contributing both of these electrons to the sigma bonding orbital.
This is not really a new concept, as we have seen electron-rich species
acting as Lewis bases before. The one thing that will take a bit of getting used to
is that when we have ligands contributing pairs of electrons to form dative bonds with a metal
cation, we do not end up with formal charges in the same way we would expect in organic chemistry.
In other words, if a nitrogen atom uses its lone pair to pick up a proton,
it ends up with a formal positive charge, but that will not be the case when neutral ligands
use their electron density to coordinate to a metal atom. So for example, if this cobalt center
was coordinated to six water molecules acting as ligands, these oxygen atoms would not have formal
positive charges, and the whole complex would have an overall charge of 3+, which by convention
corresponds to the oxidation state of the cobalt atom. These dative bonds are of varying strength.
For example, these cobalt-oxygen bonds are particularly long and weak. Many metal-ligand
bonds are quite weak, in fact, and in such a case we would refer to these ligands as labile,
which means that they are easily exchanged for other ligands that make stronger bonds.
Indeed, new ligands displacing existing ligands in a transition metal complex
is a very important reaction that we will be discussing in great detail later.
The geometries of these complexes should already be very familiar from general chemistry. We just
saw an example of an octahedral complex. Other geometries include trigonal bipyramidal, square
pyramidal, tetrahedral, square planar, trigonal planar, linear, essentially all the ones we
already know, just now with metals in the center. It will not always be obvious which geometry a
complex will take on, as this may depend on the identity of the metal atom as well as the ligands.
For example, with four ligands, will a complex be tetrahedral or square planar? Answering this
question will require a tutorial all to itself. But one thing that we should look at briefly right
now is a molecular orbital diagram for transition metal complexes. On the most basic level, we
simply have to understand that on the left we have the orbitals from the metal, the lowest energy
often being the d orbitals, then higher up is the s orbital, and then the p orbitals. On the right
we have the orbitals from the ligands, which tend to be lower energy than those of the metal. When
these combine we get the molecular orbitals in the center. Down low we have the M-L bonding orbitals.
These get filled first, with the electrons donated by the ligands. Then we have the d pi
orbitals which will hold the metal’s d electrons, so these constitute the highest occupied molecular
orbitals, or HOMO. Then above that we have the unoccupied molecular orbitals, starting with
the remaining d orbitals. This dz2 orbital tends to be the lowest unoccupied molecular orbital,
or LUMO, so we will frequently see incoming ligands interacting with this orbital.
Speaking of ligands, let’s now become familiar with some common ligands and describe how they
bond to the metal atom. First up, the hydride ion, with its pair of electrons.
Its lone s orbital will overlap directly with some orbital on the metal, as in this case with
the dz2 orbital, to produce a sigma bond. The halide ions will all behave this way as well,
though these have four pairs of valence electrons and are sp3 hybridized. One of these lone pairs
will be donated to form the dative bond. Now from the halogens we move to the chalcogen ligands.
Oxide and sulfide ions, which each have four lone pairs, can both act as ligands, and they
can do so from either a bridging position, where they make two bonds to two different metal atoms,
or a terminal position, where they make two bonds to the same metal atom. Oxygens
in these positions would be sp3 and sp2 hybridized respectively, as we would expect.
In the bridging scenario, the electron pairs are donated into the dz2 orbitals of two different
metal atoms, whereas in the terminal scenario, the first bond, a sigma bond, is oriented this way
as well, while the second bond, a pi bond, occurs due to overlap of the unhybridized p
orbital on the oxide with the dyz orbital on the metal. This is essentially a double dative bond.
Alkoxides and sulfoxides, which each have three lone pairs, behave in the same manner, although
they tend to only donate one pair of electrons to make the sigma bond with one metal atom.
The same can be said for ethers and sulfides, as well as water, which each have two lone pairs.
These can donate one pair of electrons to a metal atom to form a sigma bond.
There are some more interesting ligands involving oxygen, such as the peroxide ion.
This has a 2- charge, and the highest occupied molecular orbital is a pi star orbital. So this
ligand can bond with two metals atoms, one per oxygen, or both oxygens can bond with the same
metal atom. A similar ligand is the superoxide ion. This has just a 1- charge because one of
the oxygens is a radical. The one with the full octet will do the coordinating,
again from the pi star orbital, so this ligand will typically bond to just one metal atom.
Moving on to pnictogens, there are a number of nitrogenous ligands. The nitride ion has
a 3- charge and thus four lone pairs. Three bonds can be made to a metal atom,
one of which, a sigma bond, will occur between the sp hybrid
orbital on nitrogen and the dz2 orbital on the metal, and the other two are pi bonds between
the two unhybridized p orbitals on nitrogen and the dxz and dyz orbitals on the metal.
Imides, or NR2-, can do something similar, just with one sigma bond and one pi bond.
Amide and phosphide ions, or NR2- and PR2-, can make one sigma bond to a metal atom,
as can neutral amines and phosphines. Finally the nitrosyl ion, or NO+, can also act as a ligand.
This is an oxygen and nitrogen with a triple bond between them, each having a lone pair,
and a formal positive charge on the oxygen. With this we introduce a new concept, that of
backbonding from the metal. With this ligand, a sigma bond forms due to the overlap of the
2p sigma orbital from this nitrogen atom and the dz2 orbital on the metal. And then an additional
interaction occurs where the metal donates a pair of electrons from one of its d orbitals
into the p pi star orbital of the nitrogen. So the nitrosyl ion will be a good ligand for
metals that have electrons in these d orbitals with which to provide this backbonding effect.
We should also note that apart from this linear bonding for the nitrosyl cation, the nitrosyl
anion can also coordinate, but in bent fashion, as is shown here. The number of electrons contributed
will not change, but we must be aware of this discrepancy in charge, as well as geometry.
Continuing to ligands from the carbon group, first we have alkylidynes, or CR3-.
This will be similar to the nitride, where we get a sigma bond as usual,
and then two pi bonds between the p orbitals on carbon and the dxz and dyz orbitals on the metal.
Then alkylidenes, or CR22-, will form a sigma bond and one pi bond.
Carbanions will donate the one lone pair to form a sigma bond, as will analogous ligands replacing
carbon with silicon, germanium, or tin. Then finally, an extremely common ligand will be the
carbonyl, or CO ligand. When encountered in this form it will look like this, with a triple bond
between the atoms and a lone pair on each. This will work very similarly to the nitrosyl ion.
The carbon donates its lone pair in the 2p sigma orbital to generate the sigma bond, and then
the 2p pi star orbital on the carbon accepts backbonding from one of the d orbitals in the
metal. When this type of backbonding is present, the ligand is bound quite strongly to the metal.
So that covers the general concept of bonding in transition metal complexes.
We have learned about dative bonds and how they differ from the covalent bonds we are used to,
the types of orbital overlap that occur between metals and ligands,
including the new concept of backbonding, as well as a brief survey of simple and common ligands
based on elements from Groups 14 to 17. Now let’s move forward and check out some trickier ligands.