Hi. It's Mr. Andersen and this is chemistry essentials video 22. It's on Lewis Diagrams
and VSEPR Models. And it's a powerful model if we can use these two together to make predictions
about the structure of something like methane over here. It works great if we're looking
at atoms that are bonded covalently and also in ions. And so we could take methane for
example, look at its constituent parts, so carbon and hydrogen. We could figure out their
Lewis dot diagram. If you don't know how to do that I'll put a little video up here so
you can see that. Then you could create a Lewis structure. And finally we can create
a VSEPR model. And once we have that we can learn all of these things about a molecule
such as methane. And so it works on molecules. We start with Lewis diagrams. And I'll give
you a method for drawing Lewis diagrams. What we can do is compare different diagrams that
we come up with by looking at their formal charges. We can look at alternatives. And
lots of times we'll find multiple alternatives. And then we have to show resonance. Another
model is to use a VSEPR model. VSEPR model is simply looking at where the pairs of electrons
are. And they're going to repel each other. This is simply going to be Coulomb's Law.
And so based on that repulsion we can create these three dimensional models of what those
Lewis diagrams look like. So these two together is a pretty powerful model that you can use
in chemistry. What we can use is use that to predict geometry, bond angles, bond energy,
bond length. And we can even put it polarity based on that as well. Lot's of times we're
going to have to extend it. It's just a model. It's not actually what the atoms look like.
And so we can do extensions on this based on overlap of bonds. And also sometimes it's
confusing when we get odd valence electrons. And so I'll present another theory. It's called
molecular orbital theory. Won't go into detail, but I'll kind of explain where that would
be useful. And so first thing we have to do is to know how to draw Lewis structures. There's
lots of different methods on how to do this out there. I couldn't find any good mnemonic
devices. And so we're going to use VSEPR twice. VSEPR is going to be our model of repulsion.
But it's also an easy way to remember how you draw Lewis structures. And so if we're
drawing a Lewis structure, the first thing you want to do is add up all of the valence
electrons. Next thing you want to do is just sketch out the skeleton of where the atoms
might be connected. Next you're going to add the electrons. In general you should add it
to the most electronegative atoms first and kind of work from the outside to the inside.
After you've done that we can look for pairs of bonds. And then finally we're going to
review the formal charges. Now I know this seems confusing, but we'll work through a
number of different molecules and I think it will make sense. So let's say we're starting
with this. This is hydrogen cyanide. If we're going to draw the Lewis structure of it, the
first thing we want to do is draw and add up the total number of valence electrons.
So to figure out valence electrons we're ignoring the metals. You can see I've removed that
D block right here. But what we've got is if you're in this first column, then you're
going to have one valence electron. So hydrogen. If you're beryllium you're going to have two.
If you're carbon you're going to have four. And so we go up to this molecule and we're
going to add up all the valence electrons. And so that's 1 plus 4 plus 5 is 10. So we
start by V, by adding the valence electrons. Next thing we do is simply sketch out the
skeleton. And easy way to do that is to just sketch it out by just drawing bonds between
each of that atoms as they are in the molecule itself. And so we've done that. Now since
these are pairs of electrons, we had to remove 4 of those electrons so now we're down to
6. Next thing we do is we're going to add the electrons. And so we would show add that
to the most electronegative or generally the one on the outside of the molecule. We'll
get to carbon in a second. So I'm going to add that to satisfy the octet rule. Remember
nitrogen wants 8 electrons. So it's going to have 2, 4, 6 and then it's sharing the
one, so it's going to have 2 electrons here. So nitrogen at this point is totally happy.
But you should see that carbon is not so happy. And so now we go to P. We're going to look
for pairs of bonds to try to make carbon a little bit happier. So what we could do is
we could share 1 of those electrons from nitrogen, or one of those pairs. And now nitrogen still
satisfies the octet rule, but carbon is closer. Now it has 6 around it. And so we could try
that again. And now we've got a Lewis structure for hydrogen cyanide. What's the last thing
we're going to do is review the formal charges. And a lot of people don't know what a formal
charge is. So we're going to start and go across from the left to the right with hydrogen.
How many electrons does hydrogen have? How many valence electrons does it have? Well
it has 1. How many have I assigned it in this model? Well I've assigned it 1 electron. Remember
it's sharing one of its electrons with carbon. But one of the electrons is really it. And
so we're going to have the 1 valence electron it has minus the one assigned. And so it's
going to have a formal charge of 0. Let's go to carbon. How many valence electrons?
It's going to have 4. How many is it sharing? Well it's sharing 1 here, 1 here, 1 here,
1 here. So it's sharing 4. So it also has a formal charge of 0. Let's go to nitrogen.
Nitrogen has 5 valence electrons. It's got these 2 just to itself. But it's also sharing
these 3 over. So it's got a formal charge of 0. If I got formal charges of 0 all the
way across I am happy. That means that I've got a pretty good model. So let's go on to
another one. Let's look at carbon dioxide. What do I do first? I'm going to add up the
valence electrons. So it's 4 for carbon plus 6 for each of the oxygen. So I get 16. Now
I'm going to sketch out the skeleton. So this is a quick tip that will help you. If you
have an atom written by itself, and then a number of other atoms written after that,
generally this one is going to be in the middle. And so when I sketch out the skeleton, it's
going to look like that. With the carbon in the middle. Remember I've used 4 of my electrons.
So now I'm back to 12 valence electrons. So I'm going to add electrons. I'm going to add
them to the electronegative oxygen on the outside. So how many have I added? I've added
all of my electrons. I'm back down to 0. But you can see that carbon's not happy. So what
I can do is I can do that pair switching right here. And now I've got carbon. It look happier
at this point. Now if we look at review the formal charges, let's go from left to the
right. If we look at oxygen, oxygen is going to have 6. How many does it have? Well you
can see that it has 6 minus 1. So it's actually going to have a formal charge of negative
1. So that's kind of not a great sign on my model so far. If we look at carbon, it's 0.
And if we look at oxygen it's going to have a formal charge of plus 1 because 6 valence
electrons minus 1, 2, 3, 4, 5. And so it's going to have plus 1. And so could you put
together this model in a different way so that I would have lower formal charges? For
sure. What I could is I could switch and get a pair on one side. I could switch and get
a pair on one side. And now I've got formal charges of 0, 0 and 0. And so this is what
it looks like. Now I should probably jimmy this a little bit. I should switch these
pair of bonds over here. And these over here. But this is pretty close. Now let's go to
ozone. O3. So I'm going to add the valence electrons. Sketch out the skeleton. Let me
add those electrons. You can see I've got an extra one compared to when I did carbon
dioxide. And so now if we draw that, I've got 1 structure like this. I could draw my
formal charges. So that's okay but can you visualize this? That I could kind of invert
it the other way? And I could also get those formal charges? Remember formal charges, the
closer they are to 0, the more stable that kind of structure is. But these one's have
the exact same formal charges. And so what am I going to draw? I'm going to draw resonance.
I'm going to draw this arrow in the middle, this two way arrow. And that means that this
structure or this structure are properly. So we have to show both of those structures.
Now let's get to the VSEPR model. So the VSEPR model is the valence shell electron-pair repulsion.
What does that mean? It's very simple. It simply means that when you have a pair of
electrons, they have this negative charge. And they're going to push on other pairs of
electrons. And they're going to push on themselves. And so think of it like a balloon. If you
were to hold one balloon it would look like this. But if I were to give you two balloons,
they're going to move apart from each if we hold it in the middle. And so that tells us
what the electrons are going to do and therefore what the structure is going to be. So imagine
if we have carbon dioxide like this, it's going to be the same. Where these pair of
electrons, these pair of electrons, you're hand is kind of holding it in the middle,
and so what shape, 3-dimensional shape do we think it's going to have? It's going to
have this kind of a shape. We call that linear. Linear is in a line. I know it doesn't look
3-dimensional, but it will get there in just a second. What's going to be the bond angle
between the two? We call that 180 degrees. And this is a linear model. So you can see
now how VSEPR allows us to start kind of visualizing these in 3-dimensions. One quick note. If
you're doing organic chemistry, lots of times they'll call this sp hybridized at this point.
I'm not going to go into what that means. In AP Chem it's not important. But it is important
that you understand the terminology. Let's go to three electrons now. Excuse me, 3 pair
of electrons. So we've got three balloons. It's going to organize itself like this. What's
something that looks like that? Maybe nitrate. And so if we're looking at nitrate like this,
it's going to have the nitrogen in the middle. Those electrons around the outside. And so
this is going to be that VSEPR model. It's going to be trigonal planar. So it's going
to have these 3 electrons coming out. Or these three atoms coming out. It's going to be 120
degree angle between the two. But it's also going to be flat at this point. Again, quick
organic chemistry note, we call the sp2 hybridized. So again, we haven't got to this third structure
yet. We haven't got to this 3rd dimension, but we're quickly going to get there. Let's
say we're looking at something like ozone. Ozone, you might think is going to be linear.
But if you look at this pair of electrons out here, what kind of a structure are we
going to get from ozone? We're going to get a bent kind of a structure. And now let's
get to 4 pairs of electrons. With 4 pairs of electrons, methane is that example I gave
you at the beginning. What's that going to look like? Well if you hold four balloons
in your hand it's going to have this tetrahedral shape. The bond angle is going to be 109.5.
And it's going to really rise out of the paper itself. You call this sp3 hybridized. And
that's kind of where it ends in AP Chem. You're knowledge of all these bond angles and geometry.
Know that we can also have more pairs of electrons. And as we do we get to what's called, let's
say we're looking at PCl5, that's going going to be trigenal bipyramidal. An so they're
going to branch off like this. What you really have is a trigonal planar here in the middle
and then you're going to have a linear structure right down here through the middle. Why do
we call it pyramidal? If I were to connect those with lines you can see how we kind of
have two pyramids on top of each other. Octahedral is when we're going to have six pairs of electrons
coming off. Again bond angles. You don't have to know it, but you do have to know that we
call this an octahedral at this point. And so what are some extensions? What are going
to be some things that we have to build on this VSEPR model? Well one thing we're going
to have to do is understand that not all of the bonds are going to be the same. Single
bonds like this are going to be different than double bonds like this. And that's because
the orbitals are going to start to overlap. And so in a single bond what we have is called
a sigma bond. Sigma bond is where they're sharing those electrons between the nuclei.
But as we start to add more electrons that are being shared, what we're using is an overlap
of orbitals. And so we get what are called pi bonds. And so pi bonds, so this would be
a sigma bond, and then a pi bond, or I could say this is a sigma and this is a pi bond.
Those pi bonds as we add them are going to be less powerful. But they're also going to
kind of lock the molecule in itself so you get formation of isomers. Also sometimes the
math doesn't add up. In other word's we're going to have a odd number of electrons. So
how do we add these pairs if we're adding one single electron? So nitrogen dioxide is
actually going to have a Lewis structure like this. But it's going to have one single electron
out here. So we call that a free radical. And so there are extensions. And so we can build
a better model, it's called the molecular orbital model, or orbital theory. It's based
on quantum theory mathematics. It's just a different model and we can answer a lot of
these questions. But it's probably a little too intense and a little to sophisticated
for AP Chem. And so again, did you learn the following? To use Lewis geometry and VSPER
models to predict geometry, hybridization, polarity? If you did, you did great. Try one
on your own. Let's say I give you water. Could you draw the Lewis structure? Could you draw
the VSEPR model? Watch out for the those electrons. And I hope that was helpful.