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visit MIT OpenCourseWare at ocw.mit.edu PROFESSOR: OK. All right, let us
take ten more seconds. All right, 92%. I like it. That's great, I told you there'd
be a formal charge question, and there was, and
you learned it. Awesome, that's what I like. All right, that was carbon
had four, no lone pairs. So, 4 minus 0 minus 1/2 of
eight bonding electrons, which is four. So, 4 minus 4 is 0. OK, so today we're having
a clicker competition. And problem set 4 will
be posted later today. And today's goal is for all
recitations, but recitation 12 to try to unseat recitation 12. And the goal of recitation
12 is, of course, to win an unprecedented
third week in a row. So, that's the goal
for everyone today. Now, remember recitations
that win a multiple weeks will be in the playoffs for
the clicker competition at the end of the semester for
a specially designed t-shirt. OK, more on clicker competitions
actually later today. So, if we can get settled in. I know it's exciting. And you are allowed to
talk to your recitations, other recitation members during
the clicker questions, that is allowed. But I'll need to cover
a little material first. So today, we're going to talk
about shapes of molecules, and I brought some molecules
with me today to help me out. And we're going to be
talking about VSEPR theory. So, why is shapes of
molecules important? So shape, which we can
also call geometry. So, you're going to ask about
the geometry of the molecule, you are asking
what is it's shape. It's particularly
important in chemistry because shape can
dictate properties of a particular molecule. It can tell you
about-- or dictate melting points, or boiling
points, or reactivity. And I'm a biological
chemist, so I care a lot about shapes of
molecules because in biology shape is really important. So, you have enzymes in your
body catalyzing reactions. And for those
enzymes to work, they are specially designed
to react with one particular kind of molecule and
not any molecule in the cell. So, they're designed
to recognize the shape of that molecule. So, biochemistry really
works by shape recognition, so shape is very important. So, there's a lot of ways to
get information about shape. But there's one
very simple theory that does exceedingly
well in predicting the shapes of small molecules. And that is called the Valence
Shell Electron Pair Repulsion theory, which is known as VSEPR. And is also known as
the V-S-E-P-R theory. I will call it VSEPR because it
is really hard to say V-S-E-P-R theory. So, VSEPR is the
topic today and this is based on Lewis structures. So, this is highly
exciting because you just finished a problem
set that had you draw lots of Lewis structures. And now on the next
problem set, you can draw more Lewis
structures and then tell us about his shape. So, that's very exciting. So, you're going to retain
all of the knowledge that you've gained in
the last problem set and continue on
with that problems. All right, so this is
a very simple theory. And it's based on the
idea that valence shell-- valence electron pairs
repel each other. Electrons repel each other. They're negatively charged,
they repel each other. That is the heart of
this theory, very simple. And as many of you
may have gathered, I love it one very
simple theories explain a lot of stuff. I enjoy that. Gets it, it doesn't get
it right 100% of the time, I'm OK with that, I'm
good with about 90%. So, this works
really pretty well. So, again we're talking
about the geometry around a central atom. And the idea is
that the atoms are lone pairs around that
central atom are going to move in such a way,
attain such a shape, such that the repulsion is minimized. So, this is all about minimizing
repulsion, minimizing stress. Again, I'm a big fan
of minimizing stress, so I also like VSEPR
for that reason as well. AUDIENCE: [INAUDIBLE] PROFESSOR: Lewis structures? What causes stress? Problem sets. Oh, but they build character. All right, VSEPR Very
simple nomenclature, again I like simple
nomenclature. A is the central atom. Maybe should be C,
but atom is there, so central atom is A.
X is the bonding atom, and E is the lone
pair of electrons. So, E lone pair electrons. X is whatever bonding atom,
and A, A is in the center. A is the central atom. One more term you need to know
for VSEPR and that's steric number, . And that's used to
predict geometries. So, what is steric number equal? Steric number is
the number of atoms bonded to the central atom,
plus the number of lone pair electrons. And you count one pair as one. And so, you want to
note, when considering VSEPR double triple
bonds are all-- and single bonds all the same. So, you don't have to worry
about double triple bonds right now. They're all counts the same. It's only the number
of bonded atoms and the number of loan pairs. So, let's look at
some examples of this. So, we have our central atom
A, bonded to two bonding atoms X, with one lone
pair of electrons. So, if you were asked what
is the formula for this, the VSEPR formula, it would
be AX2E, as shown there. And then you might be
asked to steric number. And the steric number in
this case would be what? Three. We have two bonding atoms. One lone pair of electrons. Again, the lone pair
just counts as one. So, I could also draw this
structure with a double bond to one of the Xs. And this would have the equation
AX2E, exactly the same formula. Because we don't care about
the double bond in the formula, only the number
of bonding atoms. Only the number of lone pairs. And this would have
an SN number of what? Three, right. Because it only matters about
the lone pairs and the atoms. That's all we care about here. We don't care about the double
bonds, triple bonds, whatever, we'll care about them later. But for right now
for geometry, we're only caring about bonded
atoms and lone pairs. Some of you will be very happy
after doing the last problem set to know, that you can
apply VSEPR theory to all the resonance structures that
you may come up with when you're doing Lewis structures. So, if a molecule has one
or more resonance structures that's OK. VSEPR can be applied
to any one of them. Also, if there's more
than one central atom, you need to consider
those atoms separately. So, in this case,
you would be asked about the geometry around
the carbon and the geometry around the oxygen. But a lot
of the examples we have today just have one central atom. All right, so that's just a
little introduction to VSEPR Now, there's two cases that
we're going to consider today, one are molecules
without lone pairs and one are molecules
with lone pairs. Without lone pairs is
a little bit easier than with lone pairs. So, let's start there. And we have a nice
table for you. I'll tell you that these
lecture notes are of high value. I've had people who've taken
this course want to come back, it's a nice summary
of all these shapes. So, you want to keep this
in a nice, secure location. A highly desirable notes. OK, so let's look
at a formula type. The simplest we have AX2 here. So, that has a SN
number of 2, because it has two bonded atoms. And it has this molecular shape. And I brought an example here. And this is, of course, a
linear molecule and therefore the angle-- and here we're
talking about the angle around the central atom. So, from this black atom to
that black atom, we'd have 180. So, that's the simplest
linear molecule. So, it's got a little
more complicated and add three bonding atoms. So, here we have a case of AX3. It has a SN number of 3. There are three things
bonded to the central atom. And its shape is
trigonal planar. Trigonal should be remembered,
it looks like a triangle. And you should also
remember that it's planar. You can hold it this way
and see all of the atoms are in one plane. And I'm making
some of this point because later on the
exam, when you're asked to name geometries of things. People come up with all
sorts of crazy things. So, if I spend a little time
doing my demonstration here of the shapes of molecules, it
will pay off later in the exam. It'll sear in your brain. It will be hard to forget. Trigonal planar. What are the angles? 120. Right. OK, let's move on. We'll have four atoms. SN number of 4, AX4. So, this is our
tetrahedral geometry. And I'm going to just
note, as it says down here, that when you have a
thick arrow coming out at you, like this bond
and a thin one going back, that means the one coming
out is coming straight out towards you. The dashed line is going
back into the screen. And the two that are
not thickened or dashed are in the plane. So, if I hold it like this, we
have two atoms in the plane, one coming out, one going back. And so, that's how-- if you
see that drawn that way, and you will, you can think
about it in three dimensions. All right, so what's the
angle of tetrahedral? 109.5. So, a lot of people have already
familiar with this is great, if you're not, you will
learn it quickly when you do the problems. All right, we'll keep going. We have five. And you move over here. So, a SN number of 5. We have AX5. And here is our shape. So, this is trigonal
bipyramidal. So, if you think about the
shape of your trigonal planar, you see along here, we
have our trigonal planar. So, we have our trigonal shape,
but we also now have an atom above and an atom below. And that forms a kind
of a pyramid on top and a pyramid on the bottom. So, it's trigonal bipyramidal. And there are two
sets of angles now. What are the angles
in the equatorial? 120. I love when people yell
things out, it's awesome. Angle from the axial
to the equatorial? 90. Awesome. All right. And one more basic shape group,
and that's this one here. We have six atoms bonded
to the central atom, AX6. And we have this shape. So, now I'm holding
it, two axial atoms. Two atoms coming out
towards you, two going back. Octahedral geometry. And what are all
the angles here? 90, awesome. All right. So, let's look at some examples. Hopefully, everyone got the
90 and can write it down. If not, I'm sure someone will
yell it out again for you. All right, so let's
look at some examples. We were talking about
CO2, and you calculated the formal charge on it for me. So, we have AX2. SN number of 2. This is linear and
our angle is 180. And now we can explain
why this is non-polar. We learned before that
we had polar bonds where we had difference
of electronegativity between the carbon
and the oxygen, which would make a polar bonds. You have two polar bonds, so why
is this not a polar molecule? And it's not a polar molecule
because this oxygen is pulling this way, this one's
pulling that way, and that makes it non-planar
because there's no net dipole. So, shape is important for this. All right, let's look
at the next example. We've seen a lot of
these examples before. We're talking about
Lewis structure. What was boron an example of? Incomplete octet. That's right. So, here's an example, it
has three things bound to it. And it has the shape
of trigonal planar. And the middle atom
is incomplete octet. But it's OK. Boron and which other one are
OK being incomplete octets? Aluminum, right. So, the angles are 120 here. Now, we move on. We have this molecule CH4. Can someone tell
me what this is? Methane. So, we have AX4, SN number of 4. And we have our tetrahedral
shape and our angles of? 109.5. Right. So, methane we should
all care about. Greenhouse gas, but
also some people believe will be the salivation
to our biofuels and energy problem. We will see if
that's true or not. OK, another example over here. So we have phosphorus in the
middle and five chlorines. So, what is this an example of,
in terms of Lewis structures? Yep, so this is trigonal
bipyramidal or bipyramidal. I think both are
right, I don't know. And so we have our
120 and our 90. And this was an example
we had in Lewis structures or something very similar
of an expanded octet. So, we have more. We have five bonds
around the phosphorus and that's OK because it's
N equals three or greater. So, we have an example
again of an expanded octet. And if we keep going, we'll
have another expanded octet. So, here we have AX6. So, we have six bonds
around the sulfur. Sulfur is OK with that. And what is the geometry here? Octahedral angles? 90. Yep, all right, so we
have several examples here of things with expanded octets
and also deficient octets. All right, so this is
pretty straightforward. People get a lot
of points on exams, except when they come up
with all sorts of weird kind of shapes that don't exist. But most of the time you can
learn this and it's great. With lone pairs it's a little
more complicated, but also much more fun. All right, so let's talk
now about what happens when we have lone pairs. So, electrons in bonds are
hanging out in their bond and they're not really doing
much, but being in their bond. So, they have less spatial
distribution than loan pairs. Meaning that, electrons in
bonds take up less space. And electrons in lone
pairs, they can be anywhere. They're not restricted
to their bond, so they take up more
space and therefore, cause more repulsion. So, the whole idea
of when you're talking about VSEPR
with lone pairs is that you're thinking about
electron pair repulsion. And if you have a
lone pair, that's going to give you more
repulsion than bonded electrons because lone pair electrons
can take up more space. So, it's a very simple
idea, but it actually works to explain a lot of
stuff and again, the geometry re-arranges to minimize
that repulsion. So, when we're talking
about repulsive forces then, we go from the most repulsion,
lone pair, lone pair. That's like two messy roommates
living together, that's very repulsive situation. Lone pair, bonding pair. And then bonding pair, bonding
pair is the least repulsive. To neat roommates that
usually works out quite well. So, if we keep this
in mind, we can now predict shapes of molecules
based on this repulsion. So, we can rationalize
shapes based on VSEPR theory. So, now we can think about
AX4E has a seesaw shape. Which of these two
shapes is seesaw? So, we have one case--
I'll make this one-- where we have the lone pair
in an axial position. And we have another case
where we have the lone pair in an equatorial position. So, when you have it
in the axial position, here, you have three bonding
pairs of bonding electrons pretty close. They're 90 degrees
away from each other. So, there are three sets
of bonding electrons that will be repelled
pretty strongly. Now, clicker question. Think about what's
going to be true here. How many things with
an equatorial lone pair will be repelled strongly? Very repulsive,
those lone pairs. OK. So, that was the answer
we're going for two. Because there are two sets
of bonding electrons here that are 90 degrees away. There are two that are 120,
but 120 is bigger than 90. So, there are two that
are close whereas, with this geometry there three
sets of bonding electrons that are 90 degrees away. So, it turns out that
the equatorial lone pair is more favorable. It has a little bit more
room in the equatorial area to spread out. And so, this is the
shape that we find. We never see this
shape, does not exist. We'll just take that
away, it doesn't exist. And this is called seesaw. And let me demonstrate to
you why that is the case. May I'll do it this way. So, how many of you had
seesaws in a playground? Quite a number. Seesaws are not
considered that safe. You know, the heavy
kid sits on it and keeps you up in the
air for days or whatever, until your mom comes and
gets you from the playground. Or gets off suddenly,
and the seesaw flips, and you go flying in the air. But studies show actually, that
dangerous playground equipment builds neural networks and is
good for cognitive development. So, I think we could do
a survey here and see if there's a correlation
between how many of you played with seesaws
when you were a kid. And whether you ended up
as an MIT student or not. That could be interesting,
but now none of you will ever forget that this
shape is seesaw, right? You will remember this forever? OK. All right. So, this idea also
is responsible for a T-shaped molecule. So, we'll add another one. And I think this
is a little harder to rationalize why it wouldn't
go in a different place, but it does. This is how it goes. Both the lone pairs are
an equatorial positions, and we get something that
looks like a T. So, if you have AX3E2 SN 5 number, it's
a T-shaped molecule. We can also think about our
core octahedral geometry. And if you have two
lone pairs, so here's our octahedral with
six bonding atoms. And if you have two lone
pairs and four bonding atoms. If you have AX4E2,
it has this shape. The lone pairs go
on opposite sides. And now they have their
repelling the bonding electrons here, but they're far away from
each other, which is favorable. And this is called
square planar. So, it's square and it's planar. Square planar. OK, the two lone
pairs are far apart on the opposite
sides of the bond. So, in addition to predicting
these shapes, which work pretty well, we can also
think about the geometries. And when you get lone
pairs, you find deviations from your standard geometries. So, let's look at some examples. So, with molecules with a
lone pairs, such as NH3, the angles tend to be smaller. So, we saw methane
before, it was 109.5. You yelled that
nicely out for me. These are both SN 4. This is AX4. This is AX3E, but it's
still an SN 4 system. So, it looks like this, we
have-- here we have methane, here we have NH3. And now, we want to think about
which would be more repulsive, the bonded electrons
or the lone pair electrons in terms of
the geometries around. So, if we have hydrogen carbon
hydrogen angle of 109.5 here, then here with
nitrogen we're thinking about hydrogen
nitrogen hydrogen. And this lone pair is pushing
on those bonded electrons, and it's taken up
a lot of space. This is the messy roommate. The messy roommate
has a lot of stuff that's spreading out
all over your room. And so instead of having
109.5 amount of space, you now have 106.7 amount
of space in your room. Because it's just-- the
messy roommates just spreading out all
over the place, pushing down on those bonds
and the bond's contract. All right. Now, let's go back to
trends in the periodic table for a minute. We learned that
atomic size increases as we go down the periodic
table because what increases? N. Principal quantum number
N increases, so atomic size increases as we go down. Think of this as a messy
roommate having more stuff. So, the lone pairs now
occupy a larger volume. Still messy, but now messy
with a lot more stuff. And this stuff
impedes in your space. So, the angles tend to be even
smaller between the bonded atoms. So, we had NH3 before at 106.7. But nitrogen is up here. Phosphorus is below it, so
it has a bigger atomic size and it has lone pairs that
occupy larger volumes. So, this is what happens.
[EXPLOSION SOUND EFFECT] Some of you may have
experienced that. That is a messy roommate
with a lot of stuff. Now, if you have a
roommate that's very neat, they're putting their
clothes in drawers. When the clothes are in drawers,
just like electrons and bonds, they don't really go anyplace. But when the messy roommate
does not put their clothes-- when they're lone
pair clothes-- then they go and this
is what happens. OK, so we can predict now by
looking at periodic trends, and thinking about how much
room those lone pair electrons are taking up, we can think
about and predict the angle between the bonded atoms. All right. So, this is pretty cool,
but can it save the world? What do you think, can
VSEPR save the world? Do something important? Of course it can, it's
part of chemistry. So let's think about a pressing
problem in the world right now and how VSEPR can address it. And if you listen to NPR or
open a newspaper these days, you're hearing about a lot
of car bombs, explosions, things are not good
in the Middle East. We're hearing about
all sorts of-- I heard on the news about
bombs going off at a school and then when parents
rushed toward the school more bombs went off. I mean just really
horrible stories. And even when the conflict
is over, a lot of times those improvised
explosives are still there. They still exist
in the countries. And it's estimated by the
UN, by the United Nations, that landmines kill 15,000 to
20,000 people, mostly children, women, elderly who are
out in the farm fields trying to grow food
for their family and they step on
something and blow up. So, how do you find
these explosive devices? How do you find
explosive devices that are actively being
used now in dangerous parts of the world? And how do you find the
explosive devices left behind when the war is over? So, if you are
Stephanie, you use VSEPR That's what you do
to find those explosives. So, in her own words now
I'm going to tell you-- or she's going to tell you, why
VSEPR what she calls V-S-E-P-R, because she can say that better
than I can, to find explosives. STEPHANIE SYDLIK: My
name is Stephanie Sydlik and I am a graduate student
in Tim Swager's research group at MIT. The research that were perhaps
the most well-known for is for sensing explosives,
such as TNT. [EXPLOSION SOUND EFFECT] New, bigger, better
explosives have been developed. And two of these
are RDX and PETN. And these kind of have more
bang for your buck, if you will. Unfortunately, they have an even
lower vapor pressure than TNT. Which means there's
even less molecules of the explosive in the air
and it makes them even harder to detect. The dogs that they
would send out are actually sensing
cyclohexanone and acetone, which are molecules
that are used in the purification of
these two explosives. Both cyclohexanone and acetone
have a carbonyl in them. And this carbonyl
then can interact with a group known as the urea. We have two nitrogens
connected to a carbon that has a double bond to
an oxygen between them. And these nitrogens
have hydrogens on it, that can hydrogen bond with that
acetone or the cyclohexanone that we're looking for. What happens is that the
lone pair from the carbonyl reaches out and grabs those
hydrogens and pulls them away from the nitrogens. And this makes the
nitrogen hydrogen bond more lone pair-like. As it becomes more
lone pair-like, we see more repulsion between
the lone pair-like bond of electrons and the
neighboring bonds. And by the V-S-E-P-R theory,
we know that this is going to cause the bond angle
to become larger between the nitrogen hydrogen bond
and the accompanying bonds and smaller between the other
bonds around the nitrogen. And this causes large scale
changes in the polymer. So, we can see differences
in the way in the large scale the polymer interacts with
light as fluorescents, so it will start to glow. Or, a refractive index change,
which is also a different way the polymer
interacts with light. We have instruments that
can very easily measure both fluorescents
and refractive index. And with these
very easy signals, we now know that our
acetone or cyclohexanone and therefore, the
explosive is there. For soldiers, this
is a really big deal. In Iraq and Afghanistan,
there are minefields and improvised explosive
devices almost everywhere. And the soldiers
over there really have to watch where they step. So, if we can come up
with a handheld device, and we have in the
past come up with some. And I'm hoping that my
technology might in the future also go towards these
types of devices that will be attached
to a robot and sent out to sniff out the area before
the soldiers go there. You can really save a lot of
the soldiers lives as well. It's very cool to do the
hands on work in the chemistry laboratory, and then know that
what you've done at your bench will then one day be
actually used by someone and potentially save their life. PROFESSOR: OK. So, back to VSEPR
and lone pairs. Let's look at some
examples and think about the shapes we've
seen some of these already, but let's look at some more. So, now we can have
AX2E, SN number of 3. So, two bonded
atoms one lone pair. This has bent geometry,
but again think about this lone pair. Now, we're going to
talk about the angles. And we're talking about
angles in this class, we're talking about the angles
between the bonded atoms. So, the angle from
the lone pair down is going to be bigger because
the lone pair is really repulsive, but we're
going to be thinking about the angle between
the atoms you see. So, when we ask
about angle, we're asking about this angle between
one bonded atom, central atom, and the other bonded atom. So, this lone pair doesn't
look that repulsive. So, I just want you to sort
of think more about this. So keep this in mind. This is really more what a
messy roommate is all about. So, if this is your lone pair
pressing down on those bonds, what do you expect the
angle is going to be? What's the angle in
the normal case first? [BALLOON POP] So, we have 120. And now think--
and actually it's a clicker question, think about
what the answer is going to be. We talked about the normal. All right, 10 more seconds. All right, yes, less than 120. So, you don't know
exactly, you can't say "Oh, that's going to
be 118.5 or something." But you can say less than 120. That's how you would express it. And I'll just remind
myself to say, if you like these model
kits and want your own, some toothpicks and
gum drops can create some awesome VSEPR model kits. And we'll try to
bring some of these into recitation
for people who want to have gum drops and
toothpicks for making models. OK, let's keep going. So, now we have our
tetrahedral based system AX3E and an SN number of 4
based on tetrahedral. And so here, now, we
have trigonal pyramidal. So, we have a bunch, this
is why it's confusing. It's not bi-pyramidal,
there's only one pyramid here and it looks like a triangle. So, trigonal pyramidal. And now what are the
angles going to be? And you can just
yell this one out. Yeah, 109.5. And now let's keep going. And we have another
clicker question. All right, 10 more seconds. OK. So the trick here was to
think about the parent geometry of the system. And so this is the
parent geometry is the tetrahedral system. And we know that because
it has a SN number of 4. And so, when you have
a SN number of 4, then it's going to
be less than 109.5. And this is called
a bent geometry. So, again you can think about
it within terms of those-- if you have some whole
cans, those lone pairs are pressing down on the
bonds and compressing them. So, it's less than 109.5. I just teach chemistry because
I like to buy whole cans and have a justification for
it, is really the bottom line. OK, so if we keep going now,
we have our friend seesaw, which you're never
going to forget. And I'll rebuild
my-- oops wrong one, rebuild my seesaw over here. So, now what are
the angles here? There's two of them. Think about the
equatorial angles. Yep, I'm hearing
it, less than 120 and the axial would
be less than 90. So, the lone pair is
pressing down both on the 90 and on that 120. Probably more repulsive for
the 90, but all you have to do is say less than
for both of them. So, the trick is just think
about the parent geometry. What are the angles in
the parent geometry? And then, it's less than. And if we keep going with this,
we had our T-shaped molecule, as well. So, when we added another
lone pair to our SN 5 case. And what's the angle
now going to be with those two lone pairs? Less than 90. Right. All right. So, there so many
possibilities for lone pairs. If we add yet another
lone pair to the system, what's my geometry? So, this is now going
to be linear geometry. And what's the
angle going to be? Yeah, it's just going to be 180. So, we don't have
a less then here because whatever
way it would bend it would be just moving toward
more repulsive loan pairs. So, there's no way
you can minimize the repulsion in this case. So, it's just going to
be a linear molecule. So, now we're going to move
into our SN six category. And we're going to
talk about a shape that we haven't
talked about yet. So, based on this what happens? We have our parent
geometry of octahedral. And so this has six
bonded atoms, SN 6. But now we're going to take
off one of the bonds-- put it the same as the
figure-- we're going to take off one of the bonds
and put on the lone pair. And this is called
square pyramidal because you have this
square here in your axial. It looks like a square. But when you consider you have
an atom on top coming down to these four atoms on
the side, that again looks like a pyramid. So, this is square pyramidal. And what angle do
you think you're going to have here for
these bonded atoms? Yeah, that will be less than 90. All right, so we can keep going
and we saw this one before. If we take off another
bonded electron and put it in a
second lone pair. As we saw before
those lone parents want to be a far apart from
each other as they can. So, one goes on top,
one goes on the bottom. And this was our
square planar geometry because it is so
square and it's planar. Now, what do you
think the angles are? 90. Right. There's no where
to escape when you have a messy roommate on top and
a messy roommate on the bottom. If you're in a triple
between messy roommates, there's just nothing you can do. And you can't minimize
the repulsion at all. You just have to live with it. OK, so if we keep
going again and now we're going to add
another lone pair. And it really doesn't
matter where we put that, it's all equivalent. And it comes up with the
shape, that's T-shaped again that makes sense when you
look at the structure. And now what do you
think the angles are? Less than 90, and you
would be correct in that. All right, so if we put
on yet one more lone pair and take off a bonded electron
or take off one more bonded atom and put on an electron. What's our geometry? All right, so we
have this structure, it's linear 180 no place to go. So, we keep going far enough
to come back to linear a lot. Now, let's just look at a
couple more-- some real life examples of molecules and
think about their geometries and shapes. I have a couple more
clicker questions, so let's start with
our friend water. And I had water here. So, now let's think
about what the formula type is going to be for water. And what's our formula type? AX Yup. AX2E2, two bonded atoms, two
lone pairs SN number of 4. And do you remember what
this geometry is called? Bent, yeah. And this explains then
what we talked about before that this is a polar molecule. So, these are polar
bonds between the oxygen and the hydrogen.
But in this case, it's not a linear molecule
it's a bent molecule, so it creates a net dipole. And so that makes
it a polar molecule. And if water had any other shape
and was not a polar molecule, then life would be
entirely different because water is
the solvent of life. So, this shape pretty
important for anything. Actually, they researched on
what medical doctors thought was the most important
topic to learn as an undergraduate
in premed education. And the number one topic that
was most important was water. There, and you
just learned about. It OK, so we keep
going now and you have to answer a lot of
things about this on a clicker question. All right, just 10 more seconds. You got to finish your hand out. All right, so let's go
take a look at this one and fill it in. So, we have our
AX4, which is a SN 5 system because we
have one E. And then we have our seesaw geometry. All right, so let's
just fill in the rest here and see who's won
the clicker competition. So, for the next
system we have Br. And now it's expanded,
so we have AX3E2 SN 5. We've added another lone pair. And now this makes
a T-shape, but it's kind of a little benty shape
because of the repulsion. We come down. We have AX2E5 SN 5. We've added three
more with xenon. Xenon is expanded here. We have our linear shape. We also have xenon
with four things bound. And if we did the Lewis
structure of that, we'd realize there's lone pairs
on the top and the bottom. AX4E2 SN 6. And this is our square
planar geometry. So, you can predict a lot about
just doing a Lewis structure or thinking about
where the lone pairs. You can predict geometries. And let's see who has won today. We have an upset. All right, Sam. All right, see
everybody on Monday.
