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A very simply little demonstration.
And I would like you to take a look at the two materials that
are in the two vials I have here.
One of these that I am taping is a beige-colored solid,
and it is iron dichloride, ferrous chloride.
And then over here, I have an organic molecule that
is colorless, just white.
And what I am going to do is first make a solution of the
ferrous chloride in water.
I am just using a very small amount.
And what you will see is that the FeCl two dissolves
up pretty nicely in water, and it gives a solution that
has very little color to it. You might be able to see that
it looks maybe pale yellow. Can you see that?
Okay. To that solution of ferrous
chloride in water, I am now going to add this
organic molecule. I will draw the molecule for
you before we finish today. And I am just going to add a
small amount of this colorless organic molecule --
-- to get a very nice intense red color in the aqueous
solution. And the person I am going to be
talking about today is the one who really figured out what is
going on in a reaction like that.
And this was a great mystery for a long period of time,
until Alfred Werner came along.
I would encourage you while you are at home, perhaps,
or visiting friends over this coming weekend,
to go onto the Internet and go to the Nobel Prize website,
where you can read a very nice short bibliography of Alfred
Werner, because he won the Nobel Prize in 1913.
And he won this prize based on a theory that he developed
stemming from observations he made regarding reactions of
metal salts with various substances.
And I am going to point out initially, here,
that he studied the reaction of cobalt three chloride.
I just used iron dichloride. This is cobalt trichloride
reacting with six equivalents of ammonia.
And he observed that if to aqueous solution of cobalt
trichloride was added six equivalence of NH three,
ammonia, followed by silver nitrate, that that resulted in
no AgCl precipitate. And that is rather astounding.
Sorry. We will get to "no,"
but this one is "all." How does he do this experiment?
He just puts these things in solution, adds silver nitrate,
and either there is or is not a precipitative silver chloride,
which is very insoluble. And that precipitate can be
collected by filtration, dried, and then weighed.
And then, in comparison with the mass of the added
substances, you would know how much of the chloride that was
put into the reaction actually came out as insoluble silver
chloride precipitate. And he did a series of
experiments. And so, if he uses cobalt
trichloride and less ammonia, namely five equivalents of
ammonia, then he finds that that leads instead to two-thirds of
the possible AgCl precipitate. And continuing down.
If he uses now only four equivalence of ammonia,
then the addition of silver nitrate provides one-third of
the possible of the total precipitated chloride.
And then, finally, if he drops down the number of
equivalents of ammonia to three, then we get none,
zero of the AgCl precipitating. And a further observation,
to add to these four observations,
was that in no case, here, did the solution give a
reaction with hydrogen chloride. What is the significance of
that? Hydrogen chloride,
of course, is a strong Bronsted acid.
And, if you have a base in solution, it should react with
that Bronsted acid. And what are we adding here?
We are adding ammonia. And ammonia is a base,
isn't it? But when you do the experiment
like this and then test for any reactivity with hydrogen
chloride, there is no reactivity with hydrogen chloride.
So what is going on? And, secondly,
normally, this reaction with silver nitrate is used to
quantitatively precipitate chloride from solution.
So it is a quantitative analytical test for chloride in
solution. And the less ammonia we add,
the less silver chloride we are getting as a precipitate.
How are these facts related? Well, Alfred Werner put it all
together, and he correctly formulated these complexes.
Let me write this as follows.
In that first instance, when we have added six
ammonias, Werner decided that the reason that ammonia is not
in solution in a form that is reactive with hydrogen chloride
is because the ammonia is coordinated to the metal.
And so he wrote the formula this way.
Cobalt NH three six times. And this species
is a trication. And to balance those three
positive charges, we find that there must be
three chloride ions outside of what we are going to call the
inner coordination sphere of the cobalt complex.
Now, there were lots of different preparations that had
been reported in the literature back at this time of materials
that seem to be composed of metal ions and mixtures of
chloride or ammonia or other types of molecules.
And this kind of a formulation of them was completely unique.
And I really think that in the history of chemistry,
you can compare Alfred Werner's leap, his development of the
coordination theory as very much analogous to Kekule's
description of planar benzene with all equivalent C-C bond
distances. This is really a tremendous
leap in our thinking about molecules.
And then, in the case where he is adding five equivalents of
ammonia, those five equivalents all go onto the cobalt,
and so does one chloride ion. So he writes it that way.
And that chloride ion is balancing one of the three
positive charges on the cobalt plus three ion.
So this overall, now, has a two plus charge,
and there are two chlorides external to balance the charge
on that. And then in case three,
we have added four NH three per cobalt to solution.
Four of them go on the metal and two chlorides remain and
interact with the metal in a way that we will discuss shortly.
And that system now has two of the plus three charges on cobalt
three balanced by chlorides that are in the inner coordination
sphere. And only a single chloride,
now, is needed externally to balance that charge to give
overall a neutral system. And then, finally,
when only three equivalents of ammonia are added to solution,
those three equivalents per cobalt bind to the metal.
And all three of the original chlorides can be included in the
primary coordination sphere, balancing the three positive
charges on the cobalt three ion and giving overall a neutral
coordination complex. This is coordination theory.
And the dominance of organic chemistry at that point in time
was very great. Most of the people who were
thinking about these unusual substances were thinking that
they might have structures analogous to those that organic
molecules have. And typical hydrocarbon
molecules like n-pentane or n-hexane have sequential joined
CH two groups, repeating CH two groups
in a line. And so the type of formula that
you were seeing people write for these molecules at that point in
time was, for example, a cobalt.
And then NH three, NH three, NH three,
NH three, somehow all stuck
together in a way that does not seem very intuitive to us today
because we know so much more now, partly due to the
accomplishments of Alfred Werner.
This systematic set of observations,
the use of silver chloride's insolubility as a means of
precipitating it out so that you could distinguish between
external chloride from chloride that is actually in the
coordination complex. And let me define coordination
complex.
The coordination complex is a metal ion.
Plus its ligands.
So there is another word that you need to learn in this
context. Here, I would like to define
the term ligand as an atom, or a molecule,
or an ion, that can bind directly to a metal like cobalt
in its primary coordination sphere.
And that means that they are directly connected to the metal.
And the amazing thing here and what was so different from
organic chemistry at this time was the idea that a single metal
ion can have a fairly large number of ligands.
In this particular case, Werner analyzed his experiments
with the assumption of coordination number being equal
to six.
So here, it is six. But coordination number is a
variable that depends on the metal itself and depends on the
specific choice of the ligands. Some molecules are known in
which there are very low coordination numbers.
A coordination number can be as small as two or one in some very
special instances for insoluble molecules.
And for very large metal ions, sometimes the coordination
number can be as great as about
So 12 atoms or ions or molecules directly connected to
a central metal atom. And ligands don't have to be as
simple as chloride or ammonia. Ligands can have some pretty
interesting architectures. And you can even dream up new
ligands with which to decorate a metal ion and with which to
imbue it with special properties for purposes like catalysis.
We will be talking soon about metaloenzymes.
These are proteins as ligands to metal complexes.
And very many important enzymes are metaloenzymes that have
these elements from the 3D part of the periodic table bonded.
Here is what we call the d-block, --
-- or transition elements.
And, in the case of the 3D series, you will know that we
have metals like titanium, vanadium, chromium,
manganese, iron, cobalt, nickel.
These are called transition elements because oftentimes in
ions stemming from these elements, as you go from left to
right across the periodic table, you are adding more electrons
to an incompletely filled d-shell.
And at the end today, we are going to talk a little
bit about the bonding properties of transition elements.
And that will hearken back to what I said with my discussion
of carbon monoxide and why it is a poison.
And it is the interaction, actually, with certain
d-orbitals on the iron in hemoglobin that makes CO a toxic
substance. And so how does this work?
How do ligands coordinate two metals?
Well, one simple way is if you have a ligand like ammonia that
is a base, it can also be a nucleophile, and the metal can
be the corresponding electrophile.
I can draw that to represent a lone pair of electrons on the
nitrogen. Now that we have studied
molecular orbital theory, you will know that I can also
call this the highest occupied molecular orbital of the NH
three molecule. And it is the one responsible
for the basicity of the ammonia molecule and the one responsible
for its ability to serve as a ligand in coordination
complexes, like these. And you might also suspect that
we might have some contributions to this highest occupied
molecular orbital from hydrogen 1s linear combinations.
I will just draw that in to make it a little bit more
accurate. And so, you can think of this
as a big fat lone pair that will coordinate to Lewis acids.
And the metal ion is a Lewis acid.
But it is a very interesting Lewis acid because,
unlike the BH three molecule that has a single empty
orbital, this metal seems to be able to act as a Lewis acid six
times and coordinate six bases to it in forming this
coordination complex. And if we go ahead and
crystallize molecules of this sort and use X-ray diffraction
studies to determine the bond angles and bond distances in
systems like this, what we would find is that
these nitrogens are located at the vertices of a nice,
regular octahedron. So, in the case of our first
one, we can draw it out this way.
This first one is what would result if, to that aqueous
solution of cobalt trichloride, we were to add six equivalents
of ammonia. This is Werner's first system.
It is a molecule oriented like this.
That lone pair that comes from the highest occupied molecular
orbital of ammonia is directed right at the metal from each of
the six ammonia ligands. And this system does have a
three plus charge that is balanced by three chloride ions
in solution. This locating of six nitrogens
in an array in space that approximates a regular
octahedron is what makes the octahedron such a central aspect
of the theory of transition element chemistry.
And, if you are going to design molecules that do include these
transition metal ions, --
-- whether you are going to do it for their color,
like the red color there of the iron complex that we made a few
moments ago, or whether you are going to do it to take advantage
of the properties associated with unpaired electrons like
magnetism, for example, you would begin any such
approach with the octahedron as your starting point.
Let's go ahead and consider some of the other examples
provided to us by Alfred Werner.
If instead of six, we are adding only five NH
three molecules for every cobalt,
then what happens --
-- is indeed we do get an octahedron, but one of the
chlorides is not ionized. It is bound directly to the
metal, and it is serving as a ligand.
And this species, therefore, has a two plus
charge. The cobalt ion is still
considered here to be in the plus three oxidation state.
And this system is balanced by two chloride ions that are
floating around externally in solution and that are not in the
primary coordination sphere. These atoms here that are part
of ammonia molecules that are bonded directly to the metal are
in the inner coordination sphere.
That is, the inner or primary coordination sphere.
What do you think happens if you take a metal salt and
dissolve it in water? I did that a moment ago with
ferrous chloride. I dissolved it in water.
Water is a very polar solvent. It promotes the formation of
ions in solution because of its great polarity.
It is good at solvating ions, water is, as a medium.
If I take FeCl two and add it to water,
as I did a moment ago, and it ionizes,
what is happening to the iron?
The iron is going to take up water molecules into its inner
coordination sphere. When you dissolve FeCl two
in solution, which might often be written
quite simply as FeCl two aqueous,
what you really have in solution is the system in which
six water molecules are bonded to that iron.
And, because I used FeCl two, this system had a two
plus charged balanced by two of the chloride ions that
dissociate from the iron and ionize and go out into solution
to be solvated separately from the cation by water.
That initial weakly-colored solution contained iron in this
form, hexaaquairon two.
And I will tell you a little bit about what made the color
change in a moment.
But first I would like to discuss an issue that arises in
the Werner system. And this is the problem of
isomerism.
Werner found that you could make different cobalt complexes
that would have the same chemical formula,
but, for example, one would be red and one would
be green, or one would be yellow, for example,
even though they have the same chemical formula.
And that was because, as he correctly reasoned,
they were forming isomers.
And this comes into play, for example,
when you add only four equivalents of ammonia to
solution. And here is why.
If I put the first chloride up on top, as I have done here,
there are two choices of where to put the second one that are
not the same. I can either put a chloride
here, such that we have a bond angle of 90 degrees between the
two chlorides. And I will draw in our
remaining ammonia molecules that are coordinating to the cobalt.
This is an isomer that we would call "cis." Cis denotes a
proximal arrangement of the two chlorides with a 90 degree bond
angle between them. And then the alternative here
would be to put the other chloride 180 degrees away from
the first one. And that gives us what we call
the trans iosomer.
And note that both of these would have a single plus charge.
Trans means across. So the two chloride ligands are
located in a mutually trans disposition here.
Isomerism is very important. I will discuss a couple other
types of isomerism that you can get and that Werner contributed
to our understanding of very greatly.
And let me do that by completing consideration of
this. You might ask yourself in the
case where we added only three ammonias to the solution is
there a possibility for the formation of isomers?
And the answer again would be yes, we can have two
possibilities. And this is for a neutral
system that contains three ammonias and three chloride
ligands. And let's say I put the first
one here, the second one here, the third one here.
That is one of our possible isomers of this neutral
coordination complex. And then the other possibility,
the only other possibility is with one there,
one there, -- -- and then the third one here.
And so you can try to draw different structures.
And you will see that these are the only two possible structures
that you can draw for a combination of three ammonia
ligands and three chloride ligands surround a central
cobalt three plus ion in an octahedral
array.
And these have names, too.
This one is the so-called fac isomer.
And that fac is an abbreviation of the word facial,
because if you remember that the octahedron is composed of a
set of eight equilateral triangles, then the polyhedron
that we call the octahedron has both vertices and faces.
And these chlorides, in this particular case,
can be seen to define one of the eight faces of the
octahedron. And so that is the facial
isomer. And then, the other type of
isomer for this type of structure is called mer.
And that is an abbreviation of the word meridional,
which would be like the meridians of longitude that you
see on the globe. They start at the top and run
down through the equator and all the way down to the South Pole.
That is your meridional isomer. These isomers,
here, are called geometric isomers.
There are different types of isomerism.
Because the complexes that differ only with regard to the
spatial arrangement of the ligands, but not with respect to
the formula of the system, these would be types of isomers
known as geometric. We have the possibility of CIS,
TRANS, FAC, MER geometric isomers for molecules that have
the same formula. And then, there is a further
type of isomerism. And here, again,
the contributions of Alfred Werner were exceedingly important, because it was thought that this next type of
isomerism was restricted to organic molecules.
And this is stereoisomerism.
Stereoisomerism is a little more subtle than geometric
isomerism. And it is a little more subtle
because two molecules that are stereoisomers of each other are
related in the same way that your left hand and your right
hand are related. They are non-superimposable
mirror images. If you can find a way to
separate molecules that are chiral, then you can have a
sample that can do interesting things, like rotate the plane of
polarized light. This happens when you have
chiral molecules. And if a molecule is chiral,
that is to say it is non-superimposable on its mirror
image.
And in order to see whether a molecule is or is not
superimposable on its mirror image, you really need to get
good at visualizing things in three-dimensions and at rotating
molecules around in your mind. You can also do it on the
computer. And doing it on the computer
will help you prepare for doing it on the exam,
where you have to do it in your mind.
But if you like architecture, and you like visualizing things
in three-dimensions, you should know that that is a
lot of what we do in chemistry. You should think about these
molecules, these 3D structures, in ways that allow you to test
for a property like stereoisomerism.
And I mentioned that it was thought that stereoisomerism was
a property associated with organic molecules.
And organic molecules were compounds of carbon that were
thought to be associated very fundamentally with life and
living things. And so the fact that Werner in
one of his most amazing accomplishments was ultimately
able to synthesize a coordination complex that
contained no carbon at all but exhibited stereoisomerism just
shattered that theory and really helped to bring science onto a
much more firm footing. And that parallelism between
organic and inorganic chemistry, I think, has stemmed from this
aspect of its history. And so let's look at an example
of a molecule that is chiral --
-- that could be made from cobalt.
And if you imagine carrying out a reaction like we were talking
about up above but not even giving it enough ammonia to
displace all the water molecules then you could have an
intermediate like this.
And in this type of species what I've got are two water
molecules, two ammonia molecules, two chlorides.
And so if this is cobalt three, we would have a single positive
charge on that ion. And what I can represent here
by a dashed line would be a mirror plane.
That is our mirror. And we are going to reflect
this molecule through that mirror plane to see what its
mirror image would look like. And then, if you can rotate it
around in your mind, we can determine whether it is
or is not superimposable on that mirror image.
I am generating the mirror image by reflecting this water
to this position. This ammonia back here reflects
over to here. The top ammonia reflects still
into the top position. This water behind the board
reflects behind the board. And over here,
this chloride coming out in front of the board reflects over
to here. And we have one more chloride,
down on the bottom. That molecule is now our mirror
image. And let's go ahead and rotate
it, like this. Because it is a little hard,
I am going to highlight the position of the two ammonia
ligands. And to see if this mirror image
is superimposable on the structure we started with,
I am going to rotate this around so that we can put the
two ammonia ligands coincident with the two shown here
underlined in green on the left. We are going to do a rotation.
And I need to rotate this. I am going to rotate here,
around the cobalt-chlorine bond access.
And I am actually going to go in the negative direction to
generate the following structure.
This puts this ammonia up top, and it will put this one down
below. We have NH three and NH three
here. Let me underline them.
So those are in positions, coincident.
And this rotation also will carry that chloride from the
bottom up here, into what I may call an
equatorial position. And it puts a water molecule
down. And that rotation about this
cobalt chlorine bond left the cobalt and chlorine on that bond
axis unrotated. And then in the back,
we have this OH two molecule.
And what you can see is, if you now bring this over,
what we have, in fact, is a situation where
we are not currently superimposable with that choice.
I generated the mirror image. I have rotated it by 90 degrees
around the cobalt-chlorine bond axis to bring these two ammonias
coincident with these two. So you can see that,
whereas we have a water molecule on the bottom here,
we have a chloride over here. So that is not superimposable.
But we can do one more rotation to check the other possibility,
and that rotation will be a rotation by 180 degrees around
an axis, here, that bisects the
nitrogen-cobalt-nitrogen bond angle of 90 degrees.
We will rotate 180 degrees around that axis,
and that will bring our ammonias, again,
into a position so as to be coincident.
And rotating around that axis brings a water around front here
and puts a chloride in back, rotating around there,
and it swaps this chloride with that water molecule.
So we now have chloride down and OH two over here.
And so if we take this, we identify our ammonia
positions by green underlining, they're coincident,
here. And now where we have a
chloride coming out, we have a water coming out,
so our mirror image is not superimposable on the structure
that we generated it from through the process of
reflection through that mirror plane.
And so, what we can say is that this molecule and this one
constitute a pair of stereoisomers.
And because this condition was satisfied that the mirror image
was not superimposable on the structure we generated it from,
the molecule is chiral. And you will see that I have
chosen a molecule that contains no carbon, and yet it is chiral
and it has stereoisomers. And that was thought impossible
prior to the time of Werner.
Let me show you another example of a molecule that is chiral.
And I am going to use this example, also,
to illustrate another important feature that ligands can have.
And that is that they can have more than one atom that can bond
to the metal at the same time.
I am drawing a cobalt ion three plus complex that
has six nitrogens directly bonded to the cobalt.
But now, look what I am going to do.
I am going to put some organic material in here and link these
nitrogens by a CH two CH two unit,
a CH2-CH2 chain here. So it is CH2-CH2.
These carbons that I am representing as vertices here
each have two additional hydrogens that I am not showing.
And that is typical organic shorthand.
And I am going to suggest that this molecule would be generated
by adding three of these ligands to the metal center.
And for each nitrogen, if you consider it as being
derived from ammonia, one of the hydrogens of the
ammonia is replaced with a nitrogen-carbon bond.
And we have used this organic moiety here to tether two
nitrogens together. This is a very popular and
ancient ligand in coordination chemistry.
And, by drawing in simplistic form the two lone pairs on the
two nitrogens, you can see that this set of
four atoms is able to organize itself so as to simultaneously
point two lone pairs at the same metal center.
That is permitted by this bridge.
This particular ligand is called ethylenediamine.
And it is called en for short, ethylenediamine.
And it is an example of a bidentate ligand.
And that means that it has two teeth with which to bite down on
the metal center. It is a double Lewis base.
And when it binds to the metal center, we call that the process
of chelation. When a bidentate or a
multidentate, which would be maybe a
tridentate or a tetradentate ligand, binds to a metal through
multiple points, we call that a ligand chelate.
And we call the process one of chelation that forms these ring
structures with the metal as part of the ring produced
through multipoint binding of the ligand to the metal center.
And you are going to see that it is possible to have all kinds
of different architectures for ligands in proteins or in
synthetic systems. And the reason that I carried
out over here earlier is one in which I added three equivalents
of a bidentate ligand to this solution of iron two plus.
And, when that occurred,
this bidentate ligand displaced the water molecules from the
inner coordination sphere of the metal.
And the bidentate ligand that I used was this one.
This is a very common chelating ligand, a planar aromatic
ligand. And you can see that,
like ethylenediamine, its architecture promotes the
pointing of a pair of electrons toward the same point in space.
So that this ligand can bind itself to a metal center through
two nitrogen lone pairs simultaneously.
And it is the interaction of the d-electrons on the iron
center with the unsaturated pi system of this organic ligand
that produces the red color in ways that we are going to
explore in more detail in one of our next lectures.
But, before we do that, we are going to need to
understand something about d-orbitals.
And, as you have learned, when you are forming molecular
orbitals in systems that consist of either s or p orbitals,
you needed to know something about the nodal properties of
those atomic orbitals in order to build proper molecular
orbitals. And that will certainly be the
case for these more interesting elements that have d orbitals.
Not just s and p valance orbitals, but also a set of
d-orbitals. And I call those the 3d
elements because their principle quantum number for those
elements is three. And what we need to now know is
what do these orbitals have, as far as nodal properties,
depending on the other quantum numbers?
And I will draw a set of coordinate axes,
here, on which to map these orbitals.
It should be pretty straightforward for you to keep
straight the nodal properties of the d orbitals of which there is
a set of five. We had one s orbital for a
given valance shell, and we had a set of three p
orbitals, and there is a set of five d orbitals for the d-block
elements. And they can have different
values for the quantum number m. One is zero.
One is plus one. One is plus two.
And m can be minus one, and m can equal minus two.
And this quantum number determines the angular nodal
properties of the d orbital in question.
Here, let's draw a fairly simple one.
Let's say that we have x, y, and z.
Then what we might have is a d orbital that looks like this.
d orbitals often have four lobes.
In fact, you will see that we represent four of the d orbitals
this way and not the fifth. And let me use this pink to
represent the negative phase. And so this orbital here is
(d)xz. And that means that it has
nodes. You have two planes that are
nodes for a (d)xz orbital. And one of these is the
x,y-plane. And then the other one is the
y,z-plane. Those are planes when you go
from one side through one of those planes to the other side.
The wave function changes sign. And, just like each p orbital
has a single nodal plane, each d orbital has two.
And this is (d)xz. And we can also have one that
we call d x squared minus y squared.
And that one lies right along the coordinate axes like this,
with the four lobes being skewered by the x-axis and the
y-axis. And we have negative phase
located along y for the d x squared minus y squared.
And you can see that the nodal
surfaces here both contain z. The nodes contain the z-axis
and bisect the x and y coordinate axes.
There is one plane up here that contains the z-axis and one over
there located at 90 degrees to the first one.
Those are the nodal planes for d x squared minus y squared.
In addition to that (d)xz
orbital, I have a (d)yz orbital, which is located with its lobes
lying between the y and z coordinate axes.
And it will have phasing as indicated here in pink.
That is (d)yz. And it looks exactly like
(d)xz. And it is just rotated by 90
degrees around the z-axis relative to (d)xz.
And then, finally, we have one that looks just
like d x squared minus y squared.
And this one is (d)xy. And, like d x squared minus y
squared, the (d)xy orbital lies in the x,y-plane.
And its lobes point between the axes, as shown here with that
phasing. And then, finally --
And we will return to this point next week.
Our m equals zero orbital is our d z squared **d(z^2)**.
And d z squared lies along and is skewered by the z-axis.
It looks like a p orbital, except the sign is the same on
top and on bottom. And then it has this beautiful
torus here that is in the x,y-plane like that,
so that its nodal surfaces are actually conical rather than
planes. That is our set of five d
orbitals with which we are going to do a lot more to understand
the chemistry and coordination complexes.
Have a great break, and please don't forget to read
about Alfred Werner.