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CATHERINE DRENNAN:
That's today's handout.
We have valence bond
theory and hybridization.
So some people
ask, OK, now you're
going to tell me everything
you just learned.
It's not really
right and there's
something else that's better.
No.
All of these theories
are good theories.
They all do a very good job
predicting the properties
of molecules, but they all
have different strengths
and weaknesses.
And I think in terms
of what they're useful
for, molecular orbital theory is
very good in terms of thinking
about energy levels.
It's very good about
thinking about bond orders
or predicting whether
something's going to have
an unpaired electron or not.
Valence bond theory
and hybridization
are really good in
terms of thinking
about shapes of molecules.
So not so much about
energy levels, but shapes.
So all of these theories
are very, very useful
because we want to think
about how atoms come together
to form these molecules
and what are the properties
of the molecules.
So these theories
brought together really
give us a wonderful
picture of this.
And I really like valence
bond theory and hybridization
because I like shape.
I determine shapes of molecules,
complicated molecules,
for a living so I'm a big fan.
But I will say that when I
taught this the same lecture
last year, I
announced to the class
that I had had a dream where
all these atomic orbitals were
coming together and trying to
make other kinds of orbitals.
And I realized
that, perhaps, that
was a sign that
on Friday I should
start teaching
thermodynamics which is
what we're going to be doing.
We're going to start on
thermodynamics on Friday.
And last night, I had
another dream about orbitals.
So I think this is
some more orbitals
and then we go to
thermodynamics.
And I remembered my dream
because at that moment,
my giant dog jumped on top
of me as I was sleeping
to wake me up and realize
that thermodynamics
needs to come pretty soon.
OK.
But one more theory,
valence bond.
This is not so bad.
OK.
Bonds result from the
pairing of unpaired electrons
from the valence shell
of atomic orbitals.
That's it.
That's it.
So we have one, we
bring in another
so we can make
molecular hydrogen, H2,
because they each have
one unpaired electron
and they come together
to form a bond.
I like theories that
you can put on a magnet
on your refrigerator.
That's a good theory to me.
So also as part of
valence bond theory,
we have some names of bonds.
And we've been talking about
sigma molecular orbitals
and pi molecular orbitals.
And now, we're going to talk
about sigma bonds and pi bonds.
So we had orbitals in MO
theory, valence bond theory,
we now have bonds.
Sigma orbital is cylindrically
symmetric about the bonding
axis.
Thank goodness they didn't
define them differently.
That would have
been a nightmare.
So we have sigma orbitals
that are cylindrically
symmetrical about the
bond axis and sigma
bonds are cylindrically
symmetrical about the bond
axis.
So no nodal plane
along the bond axis.
Good.
We should be able
to remember that.
So with pi bonds, we have
electron density in two lobes
with a single nodal plane
along the bond axis.
So again, with pi
orbitals, we had
more-- it wasn't
cylindrically symmetric.
So this we should
be able to remember.
A couple other things about
sigma bonds and pi bonds,
a single bond is a sigma bond.
So when there's one
bond, it's a sigma bond.
So what's a double bond?
A double bond is a sigma
bond plus one pi bond.
So if it's double bond,
it's got two types
of bonds, sigma and pi.
And what do you think
a triple bond is?
Sigma bond and two pi bonds.
So you got a triple bond like
nitrogen, you got two pi's.
Hey, it's really a good life
when you have a triple bond.
All right.
Single bonds always
going to be sigma.
Double, sigma and pi.
Triple, sigma and two pi bonds.
OK.
So now we're going to
hybridize our orbitals.
And we're going to talk about
electron promotion, as well.
So start with carbon,
carbon based life.
Carbon is really important and
if you are an organic chemist,
and by organic, it means
studying things with carbon,
you care a lot
about hybridization.
And the stuff I'm
teaching you today,
you'll see a lot if you go on
to take Organic Chemistry 512.
So carbon, such as one in
methane, so we have our methane
molecule here.
The carbon has four
unpaired-- can form bonds
with four electrons, but to
do so we need to do something
with our electrons.
So carbon comes in, it has
two electrons in it's 2s
and it has two electrons
in it's 2p's, p orbitals,
but we want to form four bonds.
And in covalent bond
theory, every bond
you bring an electron
from one atom,
an electron from the
other, and they pair
and that forms a bond.
So we don't have four unpaired
electrons to make four bonds
with this configuration
of electrons,
so we can talk about promotion
of an electron from here
up there.
And if we do that , now we
have our four single unpaired
electrons ready to
make four bonds.
And carbon does like
to make four bonds.
It does it quite often.
So that's electron promotion.
To form those four
bonds, a 2s electron
is promoted to an
empty 2p orbital.
And then, we can
hybridized our orbitals
and that means that we want
to give all our orbitals
some s and some p character.
So here are our hybrid
orbitals and let
me show you the nomenclature.
So we're talking
about n equals two.
So we have a two.
We have s character
and we have p character
and we're using three p orbitals
to make our hybrid orbitals.
So we are going to make
a 2sp 3 hybrid orbital.
And we're going to
make four of them
because we've used four
atomic orbitals to make them.
So if we are using four,
we need to make four.
So let's kind of take a look
at what's going on here.
And we'll say that these
molecular orbitals differ only
in terms of their
orientation in space.
So they don't have
different shapes,
they're just
oriented differently.
So here we have our 2s,
remember it's symmetric,
and we have our
three p orbitals,
and they're all the same
except that they're all
oriented differently in space.
And when we bring
these together,
we form four hybrid orbitals and
they kind of look like turtles,
but they're turtles oriented
differently in space,
but otherwise they're the same.
So those are our sp 3
hybridized orbitals.
So carbon has this sp
3 hybridized orbital
and it has four
unpaired electrons
available to form bonds
with four hydrogens.
So let's bring our hydrogens
in to form our bonds.
And each hydrogen brings
with it it's one electron.
So now we have two
electrons in all four
of our hybrid orbitals.
And we can think about
where the energy came from.
I just moved that electron,
I didn't think about it.
I'm like, yeah, that
just goes up here.
So where did the energy
come from to do that?
And that is, it
came from bonding.
So this molecule now is more
stable because it's bonded.
Methane isn't quite
a stable molecule.
That's another problem
in and of itself.
So the bonding allows
you to do that.
You get back from this bonding.
So let's look at
those bonds then
that are formed that make that
electron promotion worthwhile.
And so you're forming a
bond between the carbon
and the hydrogen,
you're forming for them,
and you're forming single
bonds, they're sigma bonds,
and the bond is formed between
the carbon's 2sp 3 orbital
and the hydrogen's 1s orbital.
Hydrogen can't hybridize,
it's got one, 1s orbital.
That's all it's got,
can't do anything else.
And that gives you a
bond then, a sigma bond,
that you'll see this a lot
and you'll write this a lot.
This is how we're going
to name that sigma bond.
So we're going to say sigma.
We're going to
have a parentheses.
Identify the
element, it's carbon.
N is 2.
Type of orbital, sp
3 comma hydrogen,
the name of the other element,
and it's orbital, which is 1s.
So when it ask you to
name the type of bond,
this Is the complete answer
that we're looking for.
And we'll have more
practice on this.
Now we can also
think about the shape
that this molecule would have.
What is the angle here between
this hydrogen and that hydrogen
and frankly, between any of
the hydrogen carbon hydrogens?
Yup, 109.5.
And the name of that geometry?
Tetrahedral, right.
So sp 3 gives you a tetrahedral
based geometry here.
All right.
So now let's get
more complicated.
Let's bring two carbons in.
So we have ethane, two
carbon's, six hydrogens.
So this also has its
carbons are sp three,
and this is what we
saw before for methane,
but now I'm going to rotate this
around and that's one carbon,
but we need another
carbon, but first we
can think about this one carbon.
So one of the carbons of, ethane
it would have this 109.5 angle.
It has four unpaired electrons
available in it's four
hybrid orbitals to form
interactions, one with carbon
and three of them with
hydrogen. And then we
need another one of these
so we'll bring that in
and it comes in with its
set of hybrid orbitals
and it's set of electrons.
And we form a bond between them.
And the bond we're going to form
is a single bond, a sigma bond.
And now let's bring
in our hydrogens.
So we had six hydrogens,
three for each carbon.
And so there are now
two types of bonds.
We have the carbon-carbon
bond and we also
have the carbon-hydrogen bonds.
And so the carbon-carbon
bond, which is a sigma bond,
is sigma (C-- it
has carbon-- 2sp 3,
the other carbon is the same,
C2sp 3 and then the bracket.
So that's that sigma bond.
It's a single bond.
And here is our ethane molecule.
And then we have our
carbon hydrogen bonds,
they're also sigma.
Please don't give me pi
bonds to hydrogen. It only
has that one electron
tapping with two electrons.
It doesn't want do
anything complicated.
It doesn't have p
orbitals, just that one s.
So sigma C2sp 3, H1s.
And now we have
defined this molecule
so we brought together
two tetrahedral centers
and formed this molecule
with a single bond.
So let's talk about
nitrogen. Nitrogen,
also again, very important.
So here we have five
valence electrons.
What about electron promotion?
Should I do it?
No.
Because I mean, you
could put it up here,
but it can't make any more bonds
so it doesn't really matter.
So it doesn't occur
because it would not
increase the number of unpaired
electrons to form bonds,
but we can hybridize.
So we can still
hybridize our orbitals
and we can get four
hybrid orbitals,
because we're
going to use our 2s
and all three of
our 2p orbitals.
So we'll get the same
set of hybrid orbitals.
But this time, one of them
has two electrons in it.
So it's not ready
to bond, it's happy
according to
valence bond theory.
And these are our alone pairs.
But we can form three
bonds with these guys
so let's look at
an example, NH3.
So now we have our lone pair,
it's in this orbital up here,
and then we have three
orbitals available for bonding,
each with an unpaired
electron ready for the three
atoms of hydrogen to come in.
So we bring in our
three atoms of hydrogen,
each came with an
electron, and now
you can tell me with a
clicker about the angle
and the geometry
of this molecule.
Let's just do 10 more seconds.
All right.
So this is back to VSEPR again.
So we have an angle here.
It's based on an sn 4 system,
one lone pair, three bonded
atoms.
So it's based on our
109.5, but those lone pairs
make for bad
roommates and they're
pressing all of these
hydrogens together
and so the angle
is less than 109.5
and we name this structure
based on the atoms
we see, not the lone pairs,
so this is trigonal pyramidal.
And so here we have it
here so we're naming it
without thinking about the
position of those lone pairs
that are pressing down on the
bond so it looks trigonal,
like a triangle, but it's
also a little pyramid.
So VSEPR-- VSEPR and
hybridization, they just
go right together.
It's awesome.
OK.
So we can also name
the type of bond.
So our nitrogen had 2sp 3
hybridization and our hydrogen
just 1s, it's a sigma
bond, it's a single bond.
So we named that
sigma N2sp 3, H1s.
So nitrogen, now we're
going to go back to carbon--
sorry, to oxygen-- and think
about hybridization of oxygen.
Oxygen. Should I do
an electron promotion?
No.
It's not going to help me.
It's not going create any more
electrons available to form
bonds, but I can hybridize
and I can get the same four
hybrid orbitals, our four 2sp
3 orbitals, but now two of them
have two electrons
in them and two
are available to form bonds.
Oxygen loves to form bonds
with hydrogen and form water,
most of the planet is water.
There's a lot of water and water
is really important for life
so it's great that oxygen and
hydrogen get along so well.
So the oxygen, again, has two
lone pairs which are here.
You bring in our hydrogens and
they come with one electron.
And again, now, it's still a
steric number of four systems,
so it's less than 109.5,
and it's actually a lot less
than the nitrogen because you
have those two lone pairs that
are just taken up so much
room and squeezing together
these hydrogen atoms over here
creating this 104.5 angle.
So here we have
our oxygen molecule
with its two lone pairs
and its two hydrogens,
and what's the name
of that geometry?
Bent.
And again, we have
these polar bonds
that create a dipole so
it's a polar molecule, which
is very important in life.
And we can name that bond.
It's a sigma bond.
It's made up of
oxygen O2sp 2, H1s.
All right.
So that's sp 3 hybridization.
Now let's talk about
sp 2 hybridization.
So sp 2 hybridization.
So back to our atomic orbitals.
And now, we're not going
hybridized all of our orbitals.
We're just going hybridize our
2s and two of our p orbitals.
So we'll hybridized
these guys and we
will form three hybrid
orbitals and we will still
have one un-hybridized orbital.
We will have 2p y left alone.
So let's see what this does.
How is this
hybridization useful?
So let's talk about boron.
Boron has three
unpaired electrons,
but they are not all available
right now to form bonds,
according to
valence bond theory.
So here we do want to
do an electron promotion
to put one of them up
here so that now all three
are available to form bonds.
And we can again, hybridize
these three atomic orbitals
and form three hybrid orbitals.
So we have three 2sp
2 hybrid orbitals
and then we still
have our 2 py orbitals
so don't forget to mark it.
It seems lonely.
It's over here, but it's
going to be important later
so don't feel bad for it, yet.
All right.
So boron-- let's think
about these hybrid orbitals
and how this gives us
the structure that we
know occurs when we have boron.
So boron now has its three sp
2 orbitals and these are going
to lie in a plane
and they're going
to be as far apart
from each other
as they can to minimize
electron repulsion.
And if you're in
a plane, then you
need-- far apart as you
could be is 120 degrees.
And this is what gives us
our trigonal planar geometry.
So we saw that boron formed
these trigonal planar complexes
before and again, they're
trigonal planar because they're
like a triangle and they're
in a plane, trigonal planar.
And we can now bring
in our hydrogens.
The hydrogens come
with an electron
so we have an electron
for them and there we
have our structure.
We can also name that bond.
So again we have single bonds.
So sigma B, for boron,
2sp 2, H1s, and there
are three of those.
Carbon-- carbon
can also do this.
We talked about
carbon being sp 3.
Carbon can also be sp 2.
Hybridized carbon is
amazing that's why
life is based on carbon.
Carbon can do lots of things.
So again, we're going
hybridized two p orbitals,
one s orbital to give
three hybrid orbitals
and we have our 2py
over here in the corner,
but don't feel bad for it, it's
going to do something useful.
So we now have three electrons
in these hybrid orbitals
and now we have one electron in
our 2py un-hybridized orbital,
as well.
So let's see what carbon,
with this kind of arrangement
of orbitals, can do.
And again, we're going to
have trigonal planar geometry
for our 2sp 2 hybrid orbitals.
So we have carbon there and
these are all in a plane,
but now coming out of a plane
toward us is this 2py orbital.
So it's coming out 90 degrees
away from the trigonal planar
geometry.
So an example of
sp 2 hybridization
is in this molecule,
C2H4, and it
has a double bond, which
means if it's a double bond,
it has what kind of bonds in it?
Sigma and pi, right.
So one sigma, one pi bond.
So here now, and this is the
trigonal planar geometry.
It's supposed to be in a plane,
but you can't really see it
if it's really in a plane.
But 90 degrees away from that
plane is our 2py orbital.
We brought in our two hydrogens
so this carbon here, is carbon
is there, two
hydrogens are there.
This would be 120 degrees.
Now we're going to
bring in another one.
That's the one over here.
It comes in with it's carbon.
It comes in with it's
two hydrogens forming
these single bond, sigma
bonds, between the carbon
and those hydrogens.
And now we're going to
form a carbon-carbon bond.
This carbon-carbon bond is a
sigma bond and so it's C2sp 2,
C2sp 2, but we're not done.
We said this is a double bond,
so that's our sigma bond,
but we need our pi bond.
And now py, our
un-hybridized orbital,
is extremely excited because
it can form the pi bond.
So we form a pi bond
and that's formed
by our C2py, C2py
un-hybridized orbital.
And we also have four CH bonds
and those are single bonds,
those are sigma
bonds, and so they're
formed by our C2sp 2
carbon and hydrogen 1s.
And there are four of those.
So that's an example
of sp 2 hybridization.
And one thing that's
very important,
and here you can see
what that molecule looks
like-- doesn't all fall apart--
so this is a double bond,
these smaller kits don't
let me make double bonds,
so I have a sign double bond.
And you can see the angles and
the geometry of this molecule.
And another property
of something
with the double bond like
this is that it's not
really free to rotate.
So when you have these two
kind of points of attachment,
when we have these
orbitals forming
between your on
hybridized p orbitals,
that does not allow for
rotation around the double bond.
So if you're an
organic chemist wanting
to make a molecule
that's going to be rigid,
if you put a lot of
double bonds in it,
it can't twist and
turn very well.
It's often very rigid
which is useful.
So we'll stop here
and we'll finish up
on Friday sp 2 hybridization.
For the clicker question,
the bone over there on that
is also the same as
the one on the board.
The one on the board is
written with atoms in it
and it has squiggly lines
to abbreviate so make sure
that your answer is consistent
with the picture on the board,
as well.
How we doing?
OK.
All right.
Let's just take 10 more seconds.
Remember this is a
clicker competition
so we want to get the right
answer in for your recitation.
AUDIENCE: [SIDE CONVERSATIONS]
CATHERINE DRENNAN: All right.
That's pretty good because
that's the right answer.
OK.
So let's just take a look
at this for a minute.
First let me explain.
Let's settle down, quiet down.
Let me just explain
the diagram, too,
because you'll be
seeing these diagrams.
So when you just
have a bond, a line,
and there's no atom indicated,
that means it's carbon.
Organic chemists, I think,
came up with this rule.
Carbon, they just said if
nothing's indicated, of course,
it's carbon.
Carbon is such an
important element,
we don't really need to say
more about it than that.
So you could interpret
this diagram,
you have a carbon double
bonded to another carbon.
And then up here,
there's a carbon
in that ring so I just
put, in this diagram,
carbon with squigglies,
you'll see that sometimes.
That means that there's
more atoms there,
but I'm too lazy to draw them.
And on this side,
there's a carbon,
but there's more atoms there
and I'm too lazy to draw it,
another squiggly.
And then we have the double bond
so there's a carbon down here,
as well.
It wasn't indicated, just
the line in the drawing.
And you have to predict how many
hydrogens Hydrogens are often
not indicated.
This one is indicated.
There are other
hydrogens in this drawing
that are not indicated.
You need to figure
out where they go
and the material we're doing now
is going to help you do that.
And then I also drew
something and another squiggly
because I was too
lazy to draw the rest.
So these are different
kinds of diagrams
that you'll see that
all kind of mean
there's more than one
way to kind of write
the same structure.
So this particular molecule
was used to treat schizophrenia
in the 1950s and key to the
usefulness of the molecule
was that double bond.
As we talked about last time,
double bonds restrict movement.
You can't twist around
the double bond.
And so if you had exchange and
you had this group over there
and the hydrogen over here, it
wouldn't be an active molecule.
So this double bond
fixes the orientation
of those other atoms such
that it was an active molecule
and could be used
as a pharmaceutical
to treat schizophrenia.
So in terms of
the bonds then, we
have a double bond which means
we have one sigma and one pi
bond.
And so the sigma
bond down here, we
had to know what the
hybridization was.
And here, those carbons are
bonded to three other atoms
and so it would be sp 2 carbons
and also with the double bond
sp 2.
And then we also
have a pi bond and pi
bonds are made up of
non-hybridized orbitals, our py
or our px.
And so those are the ones
that make up the pi bond.
In all the other
variations, some
you had two sigmas, that's not
right, we have a sigma and a pi
so most people figure that out.
They picked the ones that had
those categories for the most
part.
And then you had
to pay attention
to whether it was sp 2 or sp 3.
And then here, this
one, the pi bond
is not made up of
hybridized orbitals,
it's made up of the
atomic orbital leftover.
So a lot to look at
that particular problem,
but this is really good
practice for the exam, which is
coming up a week from Monday.
There's going to be lots
of hybridization and today,
we're going to post extra
problems for the exam
so you have,
really, a whole week
to start getting
ready for this exam
and to keep up with
the new material.
And so extra problems
and an old exam
are also going to be
posted later today.
All right.
So let's just finish
with hybridization now
and this is good because
this is all stuff that's
going to be on the exam.
And also, the
instructions for the exam
are attached to today's
handouts and, of course,
remember no makeup exams
and clicker competition.
So I'm not really going
to go through anything
in the instructions.
It's very similar to last time
so you can take a look at that
and see on the
material, the material
starts with the periodic
table, periodic trends,
and goes through the
material that I'm
going to finish with partway
through today's lecture.
So at the end of
hybridization, that's
the end of exam two material.
So we're going to finish our
lecture notes from last time
so pull those out.
And then we're going to
move on to thermodynamics.
So once we hit thermodynamics,
that's exam three material,
so we're almost done with exam
two material, a week in advance
to get ready for the exam.
So that's great.
Lots of time to review
exam two material
and let's see if we can have
an A average on this exam.
That would make me really,
really really, happy.
I would wear my periodic
table leggings again
if we could get an A
average on the exam.
I'm just saying,
I'd be very excited.
OK.
So we better finish
up that material
so that you can get
started, get ready for this.
So we talking about valence
bond theory and hybridization
and forming these
hybrid orbitals.
And valence bond
theory is this idea
that if you have a single
electron in an orbital,
it's available to form a
bond and bonding happens when
two atoms bring together
single electrons
and those pair up
to form a bond.
So we talked about
electron promotion
before and let's just
review what that meant.
So if you have an
empty orbital, you
can promote one
of your electrons
to that empty orbital.
And so now we have
four valence electrons
after electron promotion so
that we have more possibility
of forming bonds.
Now if you don't have
an empty orbital,
you can't promote your electron.
If you do have an
empty orbital, you
can promote it more
single electrons
available for bonding.
If you don't have
an empty orbital,
there's nothing to do with that.
So that's the trick
to electron promotion.
All right.
So now we have one electron
in each of the four
valence atomic orbitals
that we have for carbon,
but we're going to only
hybridize two of them.
We saw already last time that we
can hybridize all four orbitals
and have sp 3.
We can hybridize just three of
those orbitals and have sp 2.
Now we're going
to see that we can
hybridize just two of
those orbitals and have sp.
So carbon is really amazing.
It can do all three of these
kinds of hybridization.
That's why carbon
based life forms
are able to exist
and do so much.
So we're going to now hybridize
our 2s and our two pz.
Z is just special and
so it gets to hybridize
with the 2s leaving two
of the other orbitals
just by themselves.
So we're going to form
two hybrid orbitals.
Again, if we hybridize
two atomic orbitals,
we're going to form
two hybrid orbitals.
And if we hybridized
2s and 2pz, we're
going to get hybrid
2sp orbitals.
And we'll have our 2px and our
2py just the same as always.
All right.
So we can think about this
in terms of shapes, as well.
So we have, again, our
spherically symmetric
s orbitals and our
p orbitals and we
have the three of
our p orbitals that
are the same shape they just
differ in orientation in space.
And so we're just going
hybridize our 2pz and our 2s
and so we'll have our
kind of funny looking
I think of them as turtle
shaped or hybrid orbitals
and then we also have
our 2px and 2py orbitals
the same as always.
So what are we going to do with
our two sp orbitals and our one
2px and one 2py?
Well, we can form a pretty
cool molecule with it.
So we're going to form
something that has
a carbon-carbon triple bond.
So this is C2H2.
So now in cyan is
the sp orbital,
the hybrid orbital, that is
formed on this carbon here.
And then we have
a 2px orbital here
in the plane of the screen.
And we have a 2py orbital
coming out toward us.
And, of course,
our 2pz orbital had
been hybridized with the 2s.
So here we have this structure.
We're going to bring
in our other carbon,
and the other carbon has
the same situation going on.
And we can form a bond
between the two carbons
with our sp orbitals.
And we can form, also, with
the sp hybrid orbitals, bonds
to hydrogen. So we have two
hydrogens, one over here
and one over there.
So what is the angle between
these hydrogens here?
Yeah.
So that's 180 and,
again, we have an example
here of the molecule we're
going to build that's going
to have a triple bond.
So now let's name
those types of bonds
or as sometimes
in problem set, it
will say describe the
symmetry of the bond.
And what it means by
that is the following.
It means that, that's either
say name the type of bond
or describe the symmetry.
There's multiple ways
to ask the question
and this is the answer
to those questions.
So the bond that's formed,
the first one that's formed,
between the two
carbons is a sigma bond
and it's formed
between the sp orbital.
So C2sp, C2sp.
That's the first one.
But this is a triple bond so
we have two more bonds to form.
And this is where our atomic px
and py orbitals will come in.
So we're going to form the
next bond, which is what?
A sigma or pi?
Pi bond and that can be
between our x, px, orbitals,
so pi C2px, C2px.
And now, we have the
2py orbitals, as well,
and that allows us to
form our triple bond.
So we're also going to
have a bond pi C2py, C2py.
So again, with the
triple bond, we're
going to have one
sigma and two pi bonds.
The sigma is formed
from the hybrid orbitals
and the pi bonds are formed
by the 2px and 2py orbitals.
So carbon, really impressive.
Carbon can form of these three
types of hybrid orbitals.
It can form molecules with
single bonds, double bonds,
and triple bonds.
So let's just have
a little cheat sheet
to think about that.
So again, this is for carbon
hydrocarbon molecules,
like we've looked at so far,
that have two carbons in them.
So let's look at carbon in
C2H6, so that's over here.
What is going to be
the hybridization when
you have a carbon that has
a bond to another carbon
and a bond to three
hydrogens here?
What hybridization?
sp 3.
That's right.
And it's going to have what
kind of a bond-- single,
double, or triple?
It's going to have a
single bond and it's
going to have
tetrahedral geometry
around both of the carbons.
So both of these
carbons are going
to have tetrahedral
geometry, which
is not a blank in your
note, but what's the angle?
109.5, right.
Thank you.
So we have carbons C2H2 are
going to be sp 2 hybridized.
And what kind of a
bond are they going
to have between the two carbons?
That will be a double bond.
And what is the
geometry of that?
Right.
Trigonal planar.
And here, you have to pretend
this is a double bond,
my model kit didn't come with
double bond possibilities
and I have to hold
it very carefully,
but if I hold it very y-- oh--
the bonds are still there.
You'll see that
the angles are 120
and so this is trigonal planar
geometry at each carbon.
We didn't tape that one.
You can see there's scotch
tape all over the others.
It was not a happy molecule.
OK.
So now, C2H2, what
kind of hybridization?
sp, that is our
friend the triple bond
and we're going to have
linear geometry and 180.
So both carbons have
linear geometry.
That works.
It's always triple bonds
are much more stable,
they don't fall apart as much.
OK.
So that's a cheat
sheet for carbon.
Now if you're thinking
about nitrogen or oxygen,
those often have
lone pairs on them.
Carbon likes to
form all bonds, it
doesn't care double,
triple, single, whatever,
but it doesn't really have
a lot of lone pairs on it.
But oxygen, nitrogen
have lone pairs
and whenever you
have lone pairs,
you have to worry about
what the geometry is
because the geometry gets
named based on the atoms
that you do see,
not the lone pairs.
So this cheat sheet works for
carbon without lone pairs.
If you have lone pairs, you've
got to go back to your vesper
and think about what the
names of the geometries are.
OK.
So rules, and I
posted this on Steller
for the problem set
that was due today,
and so very simple for
determining hybridization.
And this is the kind of
equation that will not
be on an equation
sheet for an exam,
you just need to know that.
So in determining
hybridization of an atom
in a complex
molecule, you're going
to be thinking about the number
of bonded atoms plus the number
of lone pairs is going
to be equal to the number
of hybrid orbitals.
So now, clicker
question, what is
the hybridization
of an atom that has
exactly two hybrid orbitals?
All right.
10 seconds.
Yes.
Right.
sp.
So we can take a look at that
two hybrid orbitals are formed
by one at 1s orbital
and 1p orbital
and if you have two things
bonded and no lone pairs,
that's what you would get.
Three hybrid orbitals would be
sp 2 and four would be sp 3.
So again, you're going
to just be thinking
in these problems
about how many atoms
are bonded to that central
atom and how many lone pairs do
you have.
And that's going to
then let you figure out
what your hybridization is.
And we have one
exception which is
that if an atom
has a single bond
and it's terminal on the
edge of the molecule, then
we're not going hybridize it.
So we can now take a look
at an example of this
and this is going to
be another-- yeah,
just keep your clickers out.
We've got a whole bunch of
clicker questions coming
at you kind of in a row here.
And if we have this molecule,
it has a central carbon
and three terminal atoms.
Now help me figure out what
kinds of bonds this will form.
So which one of these
has the correct bond
types for this molecule?
All right.
Make a decision.
Let's just take 10 more seconds.
Interesting.
I think some time I re-poll,
but I think that we'll just kind
of go over this one and then
we'll-- do you want to go ahead
and show the answer
and then-- this is--
AUDIENCE: [SIDE CONVERSATIONS]
CATHERINE DRENNAN: And if it
wasn't a clicker competition,
I might have you discuss
it more and re-poll,
but it's a competition.
So let's go over it.
So this one isn't in
your notes so if you
want to write it at
the bottom of the page,
we'll go over what
the answer is.
Hopefully there's
a typo in there,
but we'll see when
we go through.
All right.
So let's take a look
at this molecule.
Hydrogen is terminal
and single bonded,
but we've already
talked about hydrogen
so we kind of knew that.
Oxygen is terminal,
but it's double bonded
so we need to hybridize it.
Cl is terminal and single bonded
so we don't hybridize this
and we don't hybridize hydrogen,
never hybridized hydrogen. OK.
So let's look at the kind
of bonds that are formed.
So we have a sigma bond, single
bond, this carbon is C2sp 2.
It's bonded to three
different things
and it has no lone pairs
so that makes it three
that there's three
things so we have three
hybrid orbitals, which is sp 2.
Our hydrogen is just
1s, it's always just 1s.
That's all it is.
So let's look at this bond now.
So we have a single bond
between our carbon that is 2sp 2
and then, we also
have this oxygen.
We do hybridize it because
it has a double bond
and it has two
sets of lone pairs
and it's bonded
to one atom, so it
has three hybrid orbitals, so
it's sp 2 just like the carbon.
And then we have a pi
bond and the pi bond
is made up of atomic
orbitals, either 2px or 2py.
Chlorine is single
bonded so we're not
going to hybridize it
because it is single bonded
and it's terminal.
So it's a single bond, it's
from this carbon that's C2sp 2,
we already saw that,
and then the chlorine is
going to be terminal
and so it's Cl3pz
and so that's a
non-hybridized orbital.
So good practice
for the clicker.
I think that one
could help, but we're
going to have more practice now.
I threw in a bunch of extra
problems so that one was extra
and now, let's do the
one that is in the notes
from last time,
which is vitamin C.
So I'll give you another minute
if everyone has that one down.
OK.
So let's look at
vitamin C. So vitamin C
is needed to form
collagen in your body.
Without enough vitamin C in your
diet, you could be in trouble.
So it doesn't happen
too much anymore
because there's vitamin
supplements and all sorts
of things, but often, vitamin
C deficiency is associated
with sailors who went out to sea
and didn't have a healthy diet
and they became deficient
in vitamin C and got scurvy.
And so then they
had to figure out
they had to eat oranges
or other things that
were rich in vitamin C. In
terms of who should be concerned
about vitamin C deficiency, us,
primates, we don't make vitamin
C, so we have to get it in our
diet and also, Guinea pigs.
Most other animals make it.
I don't really know why-- maybe
this is why Guinea pigs are
called Guinea pigs, they're
good for scurvy experiments
because they don't make
vitamin C. All right.
So let's look at this
vitamin C molecule
and think about what
type of molecule
it is and this is
a quicker question.
So we have to remember
back more material that's
going to be on exam two.
Does that look like a
polar or non-polar molecule
and what's true about polar
and non-polar molecules?
All right 10 more seconds.
Great.
So it is polar and it,
therefore, water soluble
and so you know that because
if there's atoms in there that
have differences of
electronegativity
of greater than 0.4, carbon
and oxygen of a difference
in electronegativity
of greater than 0.4,
oxygen hydrogen also,
electronegativity
differences greater than 0.4.
So we have a lot of polar
bonds and they're not
canceling each other out.
It's not a symmetric
molecule so therefore, it
would be a polar molecule
and water soluble.
OK.
Great.
So you're good on your polar
covalent bonds which is also
going to be on exam two.
All right.
So let's go back
to hybridization
and have a little
more practice on that.
So don't put your clickers away.
Why don't you tell
me the hybridization
of carbon a labeled up
here and in your notes.
All right.
10 more seconds.
All right.
So we know what
clicker questions are
going to determine the winners.
OK.
So carbon a was sp 3 hybridized.
So if we look at it over here,
it has bonded to four things
so there's four
which makes it sp 3.
OK.
So let's just do the rest
and you can yell these out.
Carbon labeled b, what kind
of hybridization for carbon b?
AUDIENCE: Sp 3.
CATHERINE DRENNAN: Sp 3.
Carbon c?
AUDIENCE: Sp 3.
CATHERINE DRENNAN: Sp 3.
Again, you just want
to count how many bonds
you have going on or lone pairs,
but carbon doesn't usually
like to have lone pairs.
What about carbon d?
AUDIENCE: sp 2.
CATHERINE DRENNAN: Sp 2.
Right.
It only has-- if we look
at that one over here,
I'm supposed to point to this
one-- so carbon d over here,
it has three atoms
that it's bound to.
Carbon e?
sp 2 and carbon f?
AUDIENCE: sp 2.
CATHERINE DRENNAN: Sp 2.
Right.
So now that we did that,
we can use this information
when we think about
the bonds that
are formed between these
carbons and the other atoms.
So let's look at bonding now.
So if we look at
carbon b, two hydrogen,
that's going to be a sigma
bond, and you told me
that carbon b was sp
3 so we write that.
So describe the symmetry
around the bond,
name the bond, C2sp 3, H1s,
we do not hybridize hydrogen.
So now, for the bond between
b and a, again, a sigma bond.
We already looked at the
fact that carbon b is 2sp 3,
carbon a was the same.
Now if we look at the difference
between b and c, b was C2sp 3
and then c is also the same.
Remember to write
the twos, remember
to write the hybridization,
remember to write the element,
remember to write sigma
for the single bond.
Grading these questions
on the exam is not fun.
You've got to remember to
have all those things in there
so if you get them all in there,
it makes everyone very happy.
OK.
Now let's look at
carbon b to the oxygen.
It's also a single
bond, so sigma.
We know that carbon b is C2sp 3.
The oxygen here is
also going to be sp 3
because it has two bonded atoms
and two sets of lone pairs.
OK.
One more clicker.
All right.
10 more seconds.
Great.
Yup.
So that is correct and if we
take a look at that over here,
we have carbon d, it has bonded
to three things so it's sp 2
and the oxygen is bonded to
two atoms and two lone pairs
so it's sp 3.
We can keep going and finish
up between d and c now,
we have-- oops, sorry,
d and c up here,
we have d which is 2sp 2
bonded to three things,
c has bonded to four
things, it's C2sp 3.
And then finally, d to
e, we have two bonds.
We have a sigma bond so
that's between our two--
these two carbons here
are hybridized orbitals
and again, it's 2sp 2, 2sp 2.
And it's a double bond so we
have one sigma and one pi bond
and the pi bond is between
non-hybridized orbitals,
so it's C2py, C2py or
you could have used x,
I don't really care about that.
All right.
Good practice.
I think you're getting
the hang of this.
Again, there will be
more practice problems
on hybridization posted today
to get you ready for the exam
and also to figure
out these bonds.
Once you get the
hang of this, it's
really pretty trivial and
good points for an exam.
