All language subtitles for Covalent bonds 6

PROFESSOR: Section Three.

Let's start with covalent bonds.

As you know, we have our atoms.

They have a little nucleus here.

They have a shell with some electrons.

Some more electrons.

This guy here is a carbon.

We have, let's say, a hydrogen.

And I'm not going to go into the details of chemistry.

But some of you know-- and if you don't know you'll learn elsewhere--

that atoms can share pairs of electrons.

This guy would like to have eight electrons in its outer shell here, but

it's only got four.

And so carbons are up for sharing four other electrons to make four electron

pairs shared.

We can represent that by carbon has four electrons on offer.

Hydrogen has one electron on offer.

And therefore, four hydrogens can get together with this carbon and share

that pair of electrons.

Truly share that pair of electrons.

Share that pair of electrons.

Share that pair of electrons.

And we'll write that often now like that.

This is covalent bonds.

They are shared electrons.

Pairs.

Shared pairs of electrons.

And they are strong.

How strong are these covalent bonds?

Well, there are three kinds of covalent bonds.

There's a covalent bond that involves sharing a single pair of electrons.

And very sensibly, those are called single bonds.

There are double bonds.

There are triple bonds.

Two pairs getting shared.

Three pairs getting shared.

We talk about how strong the bond is by how much energy it takes

to break the bond.

We measure it-- the energy- often in a unit of kilocalories per mole.

That's just a unit of energy over a count of atoms.

Kilocalories per mole.

A certain amount of energy per a certain number of atoms.

And most of you probably know what a kilocalorie is and what a mole is, but

if you don't, it doesn't matter.

It's a unit of energy over a count of atoms.

So in fact, to break a single bond requires 80

kilocalories, kcals, per mole.

Now, how you relate that to anything?

I've told you, don't worry about the units right now.

I'm not going to ask you to convert kilocalories into something else.

I gotta compare it to something.

What you might compare it to is thermal energy.

The random fluctuations at room temperature.

The random fluctuations at room temperature for a

molecule have an energy.

Random fluctuations, thermal fluctuations, random thermal

fluctuations at room temperature are on the order of 0.6

kilocalories per mole.

So, do you think with a 0.6 vibration level thermal energy here some

covalent bond that requires 80 is going to break it is

going to come apart?

No way.

Now, sometimes you're a 0.6.

Sometimes 1.21.

1.8.

Maybe you'll get a 4 or 5 occasionally.

You ain't gonna get an 80.

You're not to get enough energy.

Covalent bonds are extremely stable usually, unless something is attacking

them and breaking them.

So that's the first thing to know.

Second thing that I want to know about chemistry is--

a very limited list of things I want you to know about chemistry-- but I

want you know for our six favorite atoms, how many bonds

they can engage in.

So I've already told you that for two of our atoms the

maximum number of bonds.

Hydrogen, one.

Carbon, four.

Let me tell you about the others.

Nitrogen, usually three.

But you can actually get four.

Phosphorus, five.

Sulfur, six.

That's it.

You should know that.

I'm not gonna ask you to memorize a lot, but you should know those things.

STUDENT: No oxygen?

PROFESSOR: Sorry?

STUDENT: No oxygen?

PROFESSOR: Oh.

What did I do?

Oxygen.

How about oxygen?

What can we get from oxygen?

STUDENT: [INAUDIBLE].

PROFESSOR: Two.

Thank you.

I left out oxygen.

Sorry.

Good point.

Six.

It's good you're getting your money's worth in the course there.

I promised six atoms, you're getting six atoms.

Very good.

Now, I talk about sharing.

Sharing electrons.

You may have noticed the human phenomenon that when you share with

somebody else the sharing is not always equal.

We talk sharing, but some people are better sharers than others.

It turns out that is also true at the atomic level.

Some nuclei of some atoms pull harder, tug harder, on the electrons in that

outer shell.

So if I have two atoms that are sharing these electrons here in the

outer shell, it may be that this one has more positive charge, more protons

in its nucleus.

It'll depend a little bit on the protons and also the distance.

So how many protons there are there and the distance away it is will

determine how greedy that is for electrons.

How tightly it pulls them.

How tightly it holds them.

And if it's stronger, the electrons will spend more time over here than

over there.

So certain bonds, the electrons are not equally distributed.

They're probabilistically moving around, but not

necessarily equally shared.

If I give you a bond of carbon and carbon, will the two carbons be

sharing the electrons equally?

Of course they will, because they're the identical atom.

They don't have different properties.

They have the same property.

So the electron cloud will be distributed equally across the two.

But it turns out, in terms of general greediness--

greediness is referred to at the atomic level as electronegativity.

That's the technical term for greediness.

Electronegativity.

At least when applied to atoms.

How tightly they pull electrons.

Oxygen.

Nitrogen.

Much greedier than carbon and hydrogen.

Those are the ones we're gonna really care about here.

And so if I give you a bond here that involves, say, O-H,

this guy is much greedier.

The oxygen is much greedier than the hydrogen.

And marginally more negative charge.

I'm writing a little delta.

Little delta means a little quantity.

And I'm putting a negative.

And I'm putting a little delta here with a positive.

The negative is a little more over toward the oxygen, the positive a

little more--

since the electron's over there more, there's more positive charge there,

more negative charge there.

It's a little magnet.

It's a little bar magnet there, that little bond there.

Similarly, nitrogen greedy compared to hydrogen.

Minus.

We refer to these bonds up here, which have no polarity--

they're not a magnet.

They have no polarity.

We're going to refer to those as--

very sensibly, you'll like the terminology--

non-polar.

Because they have no polarity.

These guys that have a polarity will be called polar bonds.

That makes perfect sense.

Polar bonds.

Non-polar bonds.

All right.

That's pretty much it for atomic chemistry.

Now, hang on to those principles.

Covalent bonds.

They're really strong.

They don't break at random.

We have unequal sharing.

Some bonds are polar.

Some bonds are not.

All right.

Now let's go on and ask, what are the consequences of this all?

This business of unequal sharing?

It turns out that these covalent bonds are the backbones of molecules.

Big molecules are collections of atoms held together by covalent bonds.

Water, H2O.

Covalent bond from the oxygen to the two hydrogens.

A big molecule of protein.

Lots of atoms with covalent bonds.

And we're gonna get into their structures.

But it turns out that understanding biological properties requires

understanding things beyond these covalent bonds that draw the atoms as

you normally would draw them.

It actually involves understanding non-covalent bonds.

And other funny forces.

So let's try to cover the non-covalent bonds that are important to

understanding biology.

Do you feel like an expert on bonds yet?

To test your understanding, we've got a question for you about polar

covalent bonds.