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Today is a very big day for all of you.
Your first semester at MIT is half over.
And, in addition to that, you have a lot of challenges
remaining ahead of you for the rest of this semester.
I am going to help make the chemical challenges fun for you.
That is what I am here to do. I love chemistry.
I love teaching 5.112. This is the best part of my
week. So, here we are.
Do you guys know that chemistry can be incredibly fun?
Let me just show you how fun chemistry can be.
And that is just the gravy. The fun part is actually making
the molecules. Any of you see this picture in
U.S.A. Today recently?
Not reading that important journal, I see.
Well, this picture did appear very recently in U.S.A.
Today in a story that talked about the recent Nobel Prize in
chemistry that was awarded to Professor Richard R.
Schrock, of our department. And, in fact,
I don't know if you realize this, but Professor Schrock was
the first person ever awarded the Nobel Prize in Chemistry at
MIT for work conducted while at MIT.
That is a pretty amazing accomplishment because we have
so many tremendous chemists on the faculty here,
and he is the first one to break through in that manner.
And maybe you are here at a very special time,
and we will see a number more of these in the next few years.
May with you contributing, so much the better because you
are here. And let me just explain
something, too. Chemistry is fun.
And only on very rare occasions like this would Professor
Schrock be drinking champagne at 8:30 in the morning,
when this picture was taken. The Reuters reporter,
Brian, who took this photo, is actually a pretty good
friend of Professor Nocera here. And so he rode over from the
Schrock household. And Dick and Dan came in
together in the same car, as they often do,
from Winchester on that morning, which was a very
exciting day. If you come to see me for
office hours maybe I will tell you more about it in detail,
but it was an exciting day. And I am in here,
too. My name is Kit Cummins.
Kit is my nickname. My name is Christopher.
You can call me Christopher or Kit, or you can call me
Professor Cummins, or you can just say hey,
chemistry guy. That will be fine.
I am in this picture primarily because it is part of being in
the right place at the right time, just like I am here now in
front of you. But also I am in it because we
three are three in one of the areas in our department that
really focus our research on inorganic chemistry.
And what is inorganic chemistry?
I will say more about this during the semester.
But inorganic chemistry is really the chemistry of the
elements of the Periodic Table inclusively.
So, a pretty broad topic. And I like doing chemistry with
lots of elements. And so that is why these guys
and I share an association with each other.
And, in fact, I was a graduate student with
Professor Schrock. And I am very,
very proud to be able to say that.
Hopefully, my computer will be staying on, here.
And I want to point out, too, that people who love
science want to know something about its history.
And part of all of that is knowing the academic family from
which you derive, depending on the person you've
studied with in learning about chemistry.
And I studied with Schrock. And Schrock studied with
Osborn, a fantastic chemist in his own right,
no longer alive, sadly, to see his progeny,
Professor Schrock, win the Nobel Prize.
And Osborn studied with Wilkinson, and Wilkinson was
also a Nobel Prize winner. So, you do not have to go very
far back in that lineage to encounter that recognition,
the Nobel Prize, Schrock, and then it skipped
Osborne, and if you go back to Wilkinson, there is another one.
And you can keep going back. And I am not going to go
through that whole history today because I have another history
that I am going to need to go through with you today.
And I am going to bring this back over so it does not
distract you. And this other history that I
need to impart to you today has to do with some of the
contributions of Gilbert Newton Lewis.
This is one of the most famous of all chemists ever.
And if you like chemistry, I encourage you to think of the
word Newton, as it is found engraved in one of the pillars
in Killian Court at MIT as the middle name of Gilbert Newton
Lewis. It is conceivable that having
that middle name weighed heavily upon Gilbert during his
lifetime. But it remains a good moniker.
And in the future, when you have kids,
I encourage you to consider using it as a middle name.
It worked well in this particular case.
When you get the notes for today, which incidentally will
be available on the website after class today,
so you don't need to write down everything I am writing down.
You are actually going to get a little bit more in the online
version than what I am writing down here.
And if you find that you forgot something that I said during
class, just email me. I will tell you what I said.
It is not a problem. But if you want to write down
notes while I am speaking just to make notes that you can refer
to later to refresh your memory, that is fine.
I think your taking in information in lots of different
ways. We are going to use the
computer and we are going to use the blackboard today.
And you are going to have available stuff from the
Internet a little bit later, too.
And all of that is good. I think you should try to
reinforce and not worry if some of the information is a little
redundant. And never be afraid to ask
questions, especially, I like to take questions by
email or if you come knock on my office door.
In 1916, Gilbert Lewis published a landmark paper in
chemistry in the journal of the American Chemical Society which
remains the flagship journal of the American Chemical Society.
And this paper was entitled "The Atom and the Molecule."
Besides being an incredibly important and influential paper
in the history of chemistry, "The Atom and the Molecule" was
also the basis for --
Like I said, besides being a hugely
influential manuscript, "The Atom and the Molecule"
also laid the foundations for Lewis's thinking about
electronic structure theory that shaped the way we still today
talk about valence and chemical bonding.
So, an incredibly important paper from 1916.
In the lecture notes today, you will find the specific
volume and page number for that reference in the journal of the
American Chemical Society. And I encourage you in your
copious free time to download that paper.
It is really great. All the American Chemical
Society journal articles are available, from the very
beginning. And you can download it and
read it. And, having had this
introduction today, I think it will probably prove
to be a very fascinating exercise for you,
if you choose to do so. It also was the basis of a book
that Gilbert Newton Lewis later wrote on valance and on atoms
and molecules. That is also a classic in
chemistry. And, should you be shopping for
chemistry books on e-bay or something, you might find a rare
copy of it. And if you do and choose not to
buy it, please alert me and I will buy it.
But Gilbert Newton Lewis, "The Atom and the Molecule."
In this paper you are going to see, if you do choose to
download it, that he was thinking about polarity.
And he was also thinking about the relationship between the
polarity of a molecule relates to the polarity of a substance.
And I mean, when I say a polarity of a substance,
that that would be a sample, you could say a water molecule
down here and liquid water up here.
And you will see this kind of sigmoidal graph that Lewis made
in thinking about how polar condensed phase substances were
as compared with the structure of the molecules that comprise
the substance. Down here, in this region,
you will find that there are substances that are not terribly
polar. And what I have written here is
the organic shorthand for the hexane molecule.
Liquid hexane is an alkane, akin to methane,
which is natural gas. And so, if you are not familiar
with this way of writing organic molecules, I will just tell you
briefly that this terminus is CH three or a methyl group,
and then it is CH two, CH two, CH two,
CH two, and then CH three, another methyl group.
So, that is some organic shorthand.
And I will go through that kind of shorthand occasionally so
that you become familiar with it.
And that is the hexane molecule, a very non-polar
molecule composed of elements, carbon and hydrogen,
that have very similar electronegativity.
And then he goes on up the curve here, and you find this
familiar molecule, benzene, considered next,
as being somehow more polar than hexane, and liquid benzene
accordingly in its solvating properties in the condensed
phase, being a more polar medium, as it were.
And then, as we continue progress up this scale,
he writes next molecules like the diethyl ether molecule.
No doubt you are familiar with ether and its applications as an
anesthetic. The first use of ether as an
anesthesia in an operation in medicine was carried out just
across the river in Boston. There is this historical Ether
Dome there that you can go and look at.
But, nonetheless, we are just getting up a little
bit more polar than benzene. And then next you find
molecules like this one. That contains an organic
functional group called an ester.
That is ethyl acetate, which is a substance that you
might have smelled if you ever built models,
because some of the glues that you may have used would contain
it, or the paints. And so, it is a system that,
in terms of its molecular structure, is more polar.
And, when you use it in the condensed phase,
it behaves in a manner that you would describe as more polar.
And then, finally, you get up to some very polar
systems, like the water molecule.
Like the ammonia molecule. And, finally,
Lewis talks about a species like this, which is sodium twice
sulfur. That is actually called sodium
sulfide. And it is a salt analogous to
table salt, sodium chloride.
And being a salt, it is obviously very polar
because there is quite a large degree of charge separation
between the atoms in the molecule.
And, accordingly, if you take that salt and heat
it up until you have got it as hot as its melting point and
make it liquid, that molten salt has
individually charged ions moving around in solution.
And that provides a very polar medium, which would be excellent
at dissolving polar things. You will notice here that on
the non-polar side we have elements, carbon and hydrogen,
that are very close to each other in electronegativity.
And then, as we examine this relationship,
we start putting in elements like oxygen or nitrogen,
which are very electronegative close to the upper right-hand
side of the Periodic Table. That introduces polarity both
in the molecule and in substances comprised of the
molecule. And then all the way over to
the involved atoms, sodium and sulfur are very
different in electronegativity and, therefore,
have very strongly separated charges.
And, in thinking about this, Lewis proposed the notion of a
continuum. And his musing on this are
really quite lucid in that paper, but the idea was that a
molecule, or a part of a molecule even,
you can analyze molecules in terms of smaller pieces of them
that can be identified within the larger molecule,
can lie on either side of a continuum in reference to a
number of different properties. And the section of class that
we are launching into today deals with acid-base properties.
Let's mention that first, acid and base.
These are two ends of a continuum.
But he also recognized that you could have a molecule that would
be an oxidant. And the counterpart to that is
a reductant. And then, here is a term that
maybe those of you who have studied organic chemistry in
more detail than most would recognize.
This is electrophile.
Electrophile has a counterpart called a nucleophile.
And then, the fourth one that I would like to mention is
acceptor.
And, of course, the counterpart to that is
donor.
Please remember those four pairs because when you can
understand and make predictions about those four properties as a
function of molecular structure, then you will be a chemist.
I mean that is a large part of what it is all about,
understanding all the different chemical properties that will
make it interested in terms of acid-base, oxidant-reductant,
electrophile-nucleophile, and acceptor and donor.
And the currency of chemistry can be seen, with reference to
this, to be the electron. That is what chemistry is all
about. And so, we will be talking a
lot about atoms and molecules with reference to their
properties and how the properties stem from knowing
something about the electrons in the molecule.
And I think you are going to find this to be quite fun.
With respect to the properties of the medium,
let's talk about a substance like hydrogen chloride.
This is just to illustrate what I was saying over there.
If you take a molecule like hydrogen chloride,
and it is a gas, of course.
And if you make a solution of it in hexane,
hexane is going to be our solvent, and then what you find
is that in this low polarity medium hexane solution,
liquid hexane, the HCl molecule remains intact
and floats around in solution, interacting weakly with hexane
molecules that surround it. But if you take this HCl
molecule and put it into a solvent like water,
then something on the other end of the continuum transpires
because we get ionization to produce hydrogen ions.
And they go off and are separated in solution from
chloride ions. And we will talk more about
just what is going on when you separate H plus from Cl
minus and let them diffuse apart in aqueous
solution. But you can see,
I think, the dichotomy. Either the proton shares in the
electrons of the chloride, or it just pops off and ionizes
and goes out into solution. And Lewis wanted to understand
this. And, to understand it,
he developed a theory. And his theory was the "cube
theory." In this part we get to draw some cubes.
And he decided that if you thought of the nucleus,
which he calls the kernel, of an atom as being located at
the center of the system, then if you had an atom like
neon, which is a noble gas with eight valance electrons,
that these valance electrons would want to get away from each
other as far as possible because they repel each other.
They are all negatively charged.
This could be, here, the neon atom.
Here, on that cube, is how he would represent the
electronic structure of the neon atom, with the eight electrons
here represented by peach-colored circles,
each the vertex of a cube. So, they got far away from each
other. And you can see that in the
historical archives, some of his original sketch
books containing these ideas are there, and now you will know
what they mean. You will see that he progressed
beyond understanding the atom to understanding the molecule.
And here is one of the molecules that he talks about.
Here we have a molecule containing 14 valance electrons,
all arranged in that manner. And this could be,
for example, the I two molecule.
And let me draw it in terms of a dot structure,
--
-- like that. And there was known at the
time, partly because of the contributions of Lewis himself,
something called the Rule of Eight.
And that was a reference to this notion that most atoms want
to have eight electrons in their valance shell,
and it is also something that we call today the Octet Rule.
So, when you see the Rule of Eight mentioned in his paper,
it is the familiar Octet Rule. And you can see here that one
of the key features is two electrons shared equally.
And when electrons are shared equally like that,
you have a molecule with symmetrical charge distribution,
charge referring to the electrons where they are in
space. And then, he showed that you
could think about this another way.
You could consider some kind of a geometric distortion of the I
two molecule, like this.
Distorted, not only because of my drawing, but because of the
way the atoms themselves are thought to be rearranging in
this process.
Notice that what has happened here is that the I two
molecule is doing something electronically such that one
iodine is starting to pull away from the other.
This one over here is carrying with it the electron that came
from over here. And, in that case,
what we have is a molecule in which we are developing partial
negative charge on the right-hand side and partial
positive charge on the left-hand side.
As I was talking up here, I was thinking about what I was
going to say here, because the ideal is that
molecules, even if they are inherently non-polar like the I
two molecule, in their equilibrium structure
they can have electron fluctuations that lead to
partial charges in the manner shown here.
And Lewis was working beautifully toward the
development of a theory like this that would allow the
description of electronic structure to be made graphically
apparent. And this is something that
chemists really like to do, is we like to be able to see
stuff. I hope I am going to be able to
show you some really cool stuff this semester.
And, in fact, even some really cool stuff
before we finish here today. So, that is the development of
charge as represented using the Lewis cube theory in an I two
molecule. Next, one thing you could do is
say, let's just continue this process, and the right-hand
iodine with its eight electrons pops off.
That would give you an ionized situation in which you had an I
minus floating away from an I plus molecule.
And that situation is drawn in the notes, and because I am
running out of space on this board I am not going to draw
that one up. Instead I am going to go onto
another triumph of the Lewis cube theory, which was its
ability to explain different chemical bond orders.
By the way, Lewis never got the Nobel Prize.
He certainly should have. A dramatic oversight.
He was born in Massachusetts. He was a Harvard professor for
a while and did not like it, and came to MIT to be a
professor for a while. And, although he liked that,
that was not far enough away from Harvard,
so he went to Berkeley, where he became the department
head and was very successful over a long period of time in
learning about chemistry and in teaching chemistry.
I will invite you, and you will see this in the
notes, too, to go ahead and see if you can find a good biography
of Lewis because it is very instructive.
Anyway, here is another molecule considered by Lewis's
cube theory.
How many electrons are being shared by the two nuclei or
atoms involved? Four electrons shared.
We can circle them here. That is four electrons shared.
Therefore, that is not a single bond.
We call that a double bond. And so, that might be a
representation of, for example,
the O two molecule. Here we have I two,
here we have distorted I two with charges developing,
and here we have O two. O two is a very
interesting diatomic molecule. We are going to talk about that
later in the semester, but for now let's just leave it
at that. But then, even though this was
quite a successful theory, the cube theory,
it has certain problems, certain shortcomings,
such as, we can do single bonds with it quite well,
and now we can do double bonds with the cube theory,
but what do you do if you have to make a triple bond?
That was a problem for Lewis. That should be a problem for
you, too. Cubes just don't fit together
that way. They have edges and faces,
and that is it. And they have vertices.
You can do one vertex. So, triple bond is a problem.
And we know that there are molecules that have triple
bonds. And we cannot handle them with
the cube theory, so we are going to have to
something a little different, for that reason.
But there is also another reason that had to do with just
considering the chemistry of light elements like hydrogen and
helium. Those elements tend to
associate themselves with only two electrons.
And so, Lewis decided, well, maybe the rule of two is
more fundamental than the rule of eight, and that we should be
thinking about pairs of electrons somehow.
Because he looked at what was known about chemistry and
realized that in most molecules, the number of electrons is
always even, is always divisible by two.
That is something interesting about two electrons at a time
being important. You take that together with the
triple bond problem, and Lewis started thinking in
terms of electron pairs, and this is where this
"electron pair theory" comes from.
Although he spent a lot of time on the cube theory,
he was perfectly willing to cast that theory aside in favor
of something that maybe was going to work a little better,
the electron pair theory. And how does that look?
Well, of course, his initial pictures still had
the cube, which just shows that once the human mind fixes on
something, it fixes on it with great tenacity and it can be
hard to shake your way of thinking about things.
But hopefully you will be able to do that.
He still decided, let's arrange these electron
pairs on a cube with the nucleus of the atom supposed to be at
the center, and let's keep these electrons, now,
nicely collected in pairs. And let's get these pairs as
far away from each other as possible, for the same reasons
we used to be putting eight electrons as far away from each
other as possible, namely, electron repulsion.
Now, the problem of electron repulsion between the pairs was
not really solved. And Lewis decided that the pair
theory worked so well that he was not going to worry about
that, and so we won't worry about it yet,
either. You have your four pairs of
electrons arranged at the center of four of the edges of the
cube, and that geometry is what we call a tetrahedron,
which I am drawing the edges of here in pink.
You have a tetrahedral disposition of four pairs of
electrons, now, for an atom.
This might again be the neon atom.
But the nice thing about this choice is that if you want to
you can make single bonds with this theory, and you can make
double bonds --
-- with this theory, like that.
And then, due to the properties of the tetrahedron,
you can also make triple bonds. This is marvelous.
You can see that the famous chemist Gilbert Newton Lewis was
talking about electronic structure in terms of
symmetrical things, like tetrahedra and cubes,
as a way of arranging electrons reasonably in space.
This could, again, be our iodine molecule.
This could be, now, our O two molecule.
And, if we continue that progression in the Periodic
Table, what is this molecule? Dinitrogen, N two,
the major component of our atmosphere, another interesting
molecule that we will be talking about as the semester progresses
forward. But chemical language and
electronic structure theory are very much still permeated with
this affection for symmetry and the problems that you can solve
using symmetry to analyze electronic structure theory.
Let's talk next about another molecule because we want to talk
about Lewis's development of the electron pair theory,
but also we want to understand how that relates to acid-base
properties. And so, let me draw here
another molecule.
And when you look at this molecule, which is the aluminum
trichloride molecule, in which I am drawing out
explicitly all of the electrons, not forgetting these two.
And having drawn the aluminum trichloride molecule
and looked at it like that, you can ask yourself is the
number of electrons divisible by eight?
Yes it is. We have a 24 valance electron
system here. It is divisible by eight.
Of course, it is also divisible by two.
You can also ask yourself is this molecule,
as written, one that satisfies the rule of eight for all four
of the atoms? No.
We have a problem there because this thing is electron
deficient.
And if you were Lewis, you would want to know,
well, if it is electron deficient, this molecule must be
doing something about that to get more electrons.
Which atom is missing electrons, here?
The aluminum. How can we give aluminum more
electrons, in a structure like this?
Can we get it up to the rule of eight?
Double bonds, awesome.
Beautiful. We can make a double bond.
And, if we do that, we can allow aluminum to
satisfy the rule of eight, or allow the rule of eight to
satisfy aluminum. Whichever way you want to look
at it. And see what we have done,
here? We have taken one of the
electron pairs, say, this one,
and we have moved it in between the aluminum and the chlorine,
so it joins the other one in there.
And we get a double bond. And now, everybody has an
octet. But what is the problem with
that? The problem with that is that
chlorine is very electronegative,
and it does not want its electrons stolen away by the
very electropositive aluminum center.
There are other ways that aluminum can ultimately satisfy
its octet, and this is how we get into the realm of "Lewis
acid-base theory." And I want to show you one of
the other ways that aluminum can satisfy its octet,
and then, we are going to move to the computer so we can
visualize some of these properties.
If I put my aluminum here -- and do you remember this
molecule up here, the ammonia molecule?
I am going to say that if I draw that ammonia molecule in
place, here --
The ammonia molecule comes in, and it has its eight electrons.
It is sharing six of them with three hydrogens.
The idea is that it has an extra pair of electrons on that
nitrogen in the ammonia molecule, that it can share with
the aluminum. And now we go ahead and
complete the picture. And we can have eight electrons
around everybody.
And so the rule of eight is now satisfied for the three
chlorines, for the aluminum, and for the nitrogen of the
ammonia molecule. The ammonia molecule is what we
call a "Lewis base."
And the aluminum trichloride molecule is what we
call a "Lewis acid," for reasons that should now be
quite clear. And I want to show you how we
can look at these things. And if you will give me just a
second here, I am going to load the aluminum trichloride
molecule up onto the screen.
What we are going to do is remember that there are lots of
different ways in which you can visualize molecules.
And I want to show you some of the more interesting ways.
All that is is a picture of the aluminum trichloride
molecule with spheres drawn at kind of arbitrary radii.
The size of these spheres is not representative of any
particular physical quantity, but we can change that.
And so now, what we are going to do is draw the thing a
slightly different way. And what you can see is that
the positions of those nuclei are in this flat box.
I made the box kind of flat. It is kind of a slab because
this AlCl three molecule is conveniently kind of
flat in its structure. And what are we going to do
now? We are going to draw a contour
surface. And I like this.
We are going to use the solid surface method,
first. And can you see that now?
These curves that enclose the four nuclei of the aluminum
trichloride molecule do have some physical
significance, because what I have done here
is taken results, in this case from the quantum
chemistry calculation, but you could also take it from
an X-ray crystallographic structure determination study.
And I have drawn a three-dimensional contour map.
The value of the electron density at every point on the
surface of one of those ovaloid shapes is the same.
Just like when you have a contour map to describe in two
dimensions an interesting complicated surface,
we can do this in three dimensions as well.
And one of the things you notice, when you look at a
surface like this, is that the value of the
electron density actually goes down to some low value in
between the nuclei, for this.
When you see molecules represented as balls and sticks,
or if you represent molecules this way, just showing a pair of
dots between the aluminum and chloride, you may not be getting
a physically complete picture of the situation.
You can get a much more complete picture if you have
access to the actual electron density.
And then, you can even go a little further,
and I want to show you something very cool.
First I am going to need to load in a second molecule.
And that will be the product of the first "Lewis acid-base
reaction" that we have discussed.
Because, I didn't say it over there, but that was your first
exposure to a Lewis acid-base reaction, in which the ammonia
molecule comes in and binds to the aluminum trichloride
molecule, in order to satisfy the valence
requirements of the respective atom types.
And so, here we are going to look at that molecule,
and we are going to look at its electron density.
And I think you are going to like this very much.
Here it comes. And then I am going to
represent this as a solid surface.
The value of the electron density everywhere on this
surface is 0.1 electron density units.
I want you to start thinking about how that shape appears
there because I am going to underscore that a little bit
more in just a second. But one thing you should notice
is that at the NH three molecule, which is pictured at
the top of that representation, what you are seeing is that the
value of the electron density does not get very small as you
go between the nitrogen and hydrogen nuclei because both
nitrogen and hydrogen are quite similar in electronegativity as
compared with the aluminum atom in this system.
The fascinating thing about illustrating a molecule in this
way is that we can actually associate another property with
the color on the surface. Now what we are doing is
looking at the aluminum trichloride ammonia system.
And let me see if I can make this thing gradually spin,
so that I can talk about it without being pinned to my
computer. What we have on the surface on
each of these electron density isosurface enclosed regions is a
color that corresponds with a value of a function that tells
us about the propensity of electrons to be paired at that
point in space. The Lewis theory is now colored
onto the electron density isosurface of this Lewis acid,
Lewis base complex molecule. And the values that correspond
to those regions in space where electrons are most likely to be
found paired up are colored blue.
Where they are least likely to be found paired up are colored
kind of red-orange, here.
And then an intermediate color is this green.
And so, you can see that you have electron pairs associated
with each of the three NH bonds, up at the top.
You have another electron pair here in blue,
which is very proximate to this aluminum center,
that is very electron deficient as we had found right over here
because it lacks an octet. And then, each of these
chlorines is able to attract electron pairs.
Because chlorine is very electronegative,
it attracts the electron pairs to itself and sort of withholds
them from the aluminum center, but yet, that development of
charge separation keeps everything together in what we
can call an ionic interaction, an electrostatic interaction.
You can see here covalent bonding, donor-acceptor complex
formation, and electrostatic contributions to bonding.
And all of this is just an illustration of the brilliant
ideas of Gilbert Newton Lewis.