Now we're going to talk about one of the most important concepts in all of chemistry and
that's electronegativity.
So as you guys already know, chemical bonds are formed by the sharing of valence electrons
between two atoms.
Alright, so when two atoms share their electrons with each other, that forms a bond.
But the extent of that sharing will determine the identity and strength of that bond.
What that means is that all bonds are not created equal.
Some of them are very, very strong because they have intense sharing and some of them
are very, very weak because they barely have any sharing at all.
The unequal sharing of electrons in one direction or another is called a dipole moment.
And that can be symbolized using the mu symbol, the mu Greek letter.
Just so you know, the dipole moment is calculated based on two variables.
It's calculated based on the charge and it's also calculated – the charge difference
between the two atoms – and it's also calculated based on the distance between two atoms.
So the charge between any two atoms is going to be related to their difference in electronegativity.
I just wanted to point out that even though we use these two variables to figure out what
the dipole moment is, the one that we're going to deal with the most is actually going to
be the charge.
The reason is because the distances are going to be very similar for a lot of these bonds,
so that means the biggest difference is going to be the electronegativity.
It's going to be is it very, very charged or is it not very charged at all.
So that's why we're going to look at the charge mostly.
And to figure out what the charge is, we're going to use the following scale.
Now this scale is called the Pauling Electronegativity Scale.
You're going to see slightly different versions of it in your book.
For example, these are to only two significant figures, so in your book, it might be to three
significant figures.
Or there's possibilities that the numbers will be slightly, slightly different.
So for example, carbon might be like 2.44 or something like that.
But in the end, these relatively round to each 0.5, which makes it really convenient
when we're talking about what types of bonds we're dealing with.
So just to remind you guys, fluorine is the most electronegative.
It goes all downhill from there.
Another thing to point out is that hydrogen is actually unusually electronegative for
where it is in the periodic table.
Check it out.
Look what it's next to.
It's next to a 1.0, a 0.9 a 0.8 and then all of the sudden you get 2.1 out of nowhere.
It's like, what?
What's going on?
That's actually going to make hydrogen able to make covalent bonds and I'll show that
in a second.
So I want to couple this diagram, that you guys should, don't necessarily memorize it,
but be familiar with it, and I want to link that to my difference in electronegativity.
Now when you were in gen chem, your professor might have told you there's two types of bonds.
There's three types of bonds.
There's polar and then there's ionic.
Then there's like just totally covalent.
I remember when I was in gen chem, my professor made very clear distinctions like there's
one and there's the other and they're two different things.
But it turns out that it's really a spectrum.
It's really not just like one is completely polar or one is completely ionic.
Actually, it's a spectrum.
And the way you figure out what the identity of that bond is, is by taking the difference
in the two electronegativities of the two atoms.
If that difference is less than 0.5, we're going to call that covalent.
That means that there's a lot of sharing.
You can imagine that, for example, let's say that I had an F and an F. Are you guys cool
with that?
Both of these have the same electronegativities, so this one has an electronegativity of 4.0,
this one has an electronegativity of 4.0.
They both want electrons really, really bad, but they're both going to pull equally.
Think of this as a game of tug-of-war, where you have these two heavyweights pulling on
a rope.
They're both pulling as hard as they can, but it doesn't budge.
The reason is because they're both pulling exactly as hard, so the electrons get perfectly
shared in between the two atoms.
Does that make sense?
That's a covalent bond meaning that the distribution of electrons is very similar amongst the two
different atoms.
Then we have the range between 0.5 and 2.0, these are going to be polar or what's also
known as polar covalent, so write that in just so you guys know that.
That means that they are covalent in the sense that there is still some sharing, but they're
polar, polar means that it's a dipole.
Remember that polar comes from the word dipole.
What that means is that there is an unequal sharing.
That means that one of these guys is going to get a little bit more than the other.
A good example of that would be let's just say if I made this carbon and I made this
fluorine.
The electronegativity of carbon is 2.5.
The electronegativity of fluorine is 4.0.
Which one wants the electrons more?
The fluorine.
The way that I write a dipole is kind of like a vector back in physics.
You write an arrow going in towards the direction of the greatest electronegativity, but in
order to be fancy, chemists had to invent their own thing, so they make a little line.
That little line just means it's a dipole.
If I were to draw the electron cloud of this, remember that these electrons are just kind
of going around in these clouds of areas that they're able to circulate, what it would look
a lot like is like this, with a fluorine getting a ton of electrons and the carbon getting
very few.
Why?
Because this is a tug-of-war where you've got a Sumo wrestler on one side and like I'm
on the other side, so I'm just screwed.
So I am not keeping up with the Sumo wrestler and he's just taking almost all the electrons
from me.
It's very selfish of him, but whatever.
That's a polar covalent bond.
Getting that so far?
I also want to point out – we're going to talk about this a lot more in further chapters,
but this is going to result in partial charges.
What that means is that the fluorine on average is going to have way more electrons, so the
fluorine is going to have a partial negative.
Now the partial sign is a lower case delta, so the squiggly, if that were a lower case
delta.
And what it means is that it's more negative than other atoms in the molecule.
There's no number associated with it.
It's just a qualitative thing, like this is more negative, not quantitative.
Then the carbon, since it's missing electrons, that's me, right?
The guy just destroyed my ass, like I have no electrons left, so that would be a partial
positive because I don't have any electrons left, so he took them all.
I just want to show you guys that when you have polar bonds that leads to a dipole moment
and that also leads to partial charges.
That's going to be a huge part of organic chemistry as we move forward.
We're going to be analyzing lots of partial charges.
Then finally, you have ionic.
Ionic bonds are similar to what you learned in gen chem about like table salt.
You have Na and you have Cl.
Remember NaCl.
Let's go ahead and look at those differences.
Na is down at 0.9.
Cl is at 3.0.
So what is that difference?
That difference is actually going to be if you subtract the two, 3.0 minus 0.9.
I'm not going to do the carryovers or whatever.
That's going to give you 2.1. 2.1 puts it in this range which is ionic.
What that means is that that bond is actually kind of misleading.
There's almost no sharing at all.
The electrons almost fully rest on the one that's more electronegative.
In this case, if I had to write the negative charge somewhere, where would I write it?
The sodium or the chlorine?
The chlorine.
So actually it turns out that this can be either written as a bond or it can be written
as Na positive Cl negative.
Why?
Because there really is no sharing going on.
Really what's happening is that the Cl is really negative and the Na is positive, so
it's just attracted to it, but they're not sharing electrons between them.
It turns out – I'm just going to put one of those little triple arrow signs – that
means that by definition they're the same thing.
When you have an ionic bond you can either write it as a bond and you just know that
there's a dipole there or you can write it as the individual charges.
That's what happens with NaCl.
So hopefully what you guys can see now is that there are some polar bonds that are more
ionic and some polar bonds that are more covalent.
It just depends where you sit on that spectrum.
Does that make sense?
Cool.
So now you're probably thinking, “Okay, Johnny, well, I'm not going to have these
numbers memorized all the time or do I have to?
How do I figure out these questions of polarity?”
What I would say is there are a few general rules that are going to work for 99% of the
time.
Instead of memorizing each number we can just remember these rules.
So the first thing is that bonds to carbon and hydrogen are always going to be covalent.
If you figure out the numbers, you're going to find out that the difference is 0.4 and
that difference is so small that that puts it right into the covalent area.
Cool?
So bonds between carbon and hydrogen are covalent.
Bonds between two identical atoms are always going to be covalent.
What I'm doing here is just giving you rules that you can just use as a generalization.
Adjacent atoms on the periodic table, if they're attached to each other, bonded, are going
to be polar.
That's an example where, let's say, I give you nitrogen and oxygen or nitrogen and carbon,
there's going to be a dipole in that direction, towards the more polar one.
Then finally, lone pairs are also polar.
Lone pairs are going to have their own dipole moment that pulls in the direction of wherever
the lone pair is facing.
Finally, we have our last thing which is the net dipole.
Net dipoles exist when atoms have asymmetrical dipole moments.
So what that basically means is remember when I talked about with charges, those formal
charges, and how if you have a bunch of formal charges that equals a net charge, it's the
same thing with a dipole.
A bunch of individual dipole moments will make an overall net dipole.
Now I do want to let you guys know that for the sake of this course, I'm not going to
make you guys calculate out what the net dipole is.
Why?
Because that would take a lot of math.
That would actually take vectors and decomposition of axes and stuff like that.
So instead what we're going to do is we're just going to try to visualize, okay, what
would that net dipole kind of look like.
In fact, many times I'm not going to ask you to draw the net dipole, I'm just going to
ask you does it have a net dipole.
So if all of the dipole moments perfectly cancel out, then I will not have a net dipole.
But if it's asymmetrical, meaning let's say I have a really big dipole in one direction
and a really small dipole in the other, that will have a net dipole.
Does that make sense?
Now we're going to do some practice problems because I know that you guys are eager to
apply this.
So let's go ahead and do that.