00:01
We're doing the lewis structure and formal charges of three different molecules.
00:05
The first one being ocs.
00:07
We've got a periodic table up so that we can add up valence electrons.
00:13
Remember valence electrons, you can just tell from which column it's in as a shortcut.
00:19
So for this compound ocs, let's see.
00:22
Oxygen would have six.
00:26
Carbon would have four.
00:29
And sulfur has another six.
00:32
So we're working with a total of 16 valence electrons.
00:36
That's going to matter because it tells us when we run out of electrons, and if we don't have octets, we have to start doing double bonds.
00:45
All right.
00:45
So we'll connect these in the way it was written, o -c -s.
00:55
Partly because that's the order in which it was written, but also because commonly the least electronegative atom goes in the middle.
01:04
Electron negativity increases, the closer you get to fluorine.
01:07
So that means that carbon is our least election negative atom.
01:11
So i'm putting him in the middle.
01:13
All right, two bonds.
01:14
That means we've already used up four electrons, two per bond.
01:19
We've got 12 electrons left.
01:22
We can fill out our outer atoms first.
01:26
Those are them more electronegative.
01:28
So you give those electrons first.
01:31
We use up the rest of our 12, and that leaves us with a carbon that doesn't have a full octet.
01:37
Now, you might be tempted to put loan pairs on that carbon, and if you didn't count your electrons, but commonly is what happens, you cannot do that because you've run out of electrons already.
01:50
You cannot create them.
01:51
Instead, you can move electrons that you've already placed, right, which means double bonds.
01:58
What you're going to do is take one of the loan pairs that came from the left -hand side, and you're going to move it down and create a double bond.
02:11
That puts carbon up to six.
02:13
Then you can do the same thing with the other side.
02:19
That creates a structure that looks like this.
02:27
Now you had some choices you could make there.
02:31
You could have done it like this, or you could have taken two loan pairs from one side or the other and done a triple bond that looks like this or this.
02:50
Now, choosing one or the other for the sake of symmetry because it looks even is not a valid reason to make that choice.
02:57
Instead, you've got to do your formal charges.
03:00
Formal charges represent how many electrons each atom is still in possession of compared to how many valence electrons it started with.
03:09
An ideal lewis structure has zero as all the formal charges.
03:14
So that's what we're going to try to get to.
03:16
Now, the way you calculate formal charge, the full definition, it's a little bit wordy, but the full definition is you take the number of valence electrons from the periodic table for that atom.
03:33
Each atom will have its own formal charge.
03:35
You subtract off all of the non -bonding electrons on that atom.
03:41
Those are the individual dots.
03:44
And then you subtract off half of the bonding electrons.
03:48
So half of the electrons that come from the bonds where each bond has two.
03:54
Now, most teachers don't put that on the equation sheet when it comes time to take the exam.
04:00
And it's really easy to not remember that.
04:03
So an easier shortcut if this helps you is formal charge equals the number of valence electrons minus how many dots minus how many sticks are attached to that atom.
04:18
So valence minus the dots minus the sticks.
04:21
I'll show you how this works.
04:24
We'll look at our first structure on the left with the two double bonds.
04:29
First, we'll look at the oxygen.
04:32
Okay, this oxygen would have started with six valence electrons from the periodic table.
04:41
Right? it now is in control of four dots and two sticks, which gives it a formal charge of zero.
04:54
That is a good thing.
04:56
We want zeros.
04:57
So i'll put a little zero next to it.
04:59
All right.
05:02
Next, we look at the carbon.
05:06
Carbon started with four valence.
05:10
There are no dots on that carbon, but there are four sticks.
05:17
So again, we have a formal charge of zero.
05:19
Also a good thing.
05:23
Lastly, we look at the sulfur.
05:26
Sulfur starts life with six valence.
05:32
Six minus four dots on that sulfur and two sticks for a formal charge of zero.
05:39
So this is actually your most ideal format.
05:41
We'll do the other.
05:42
Just to show you the difference, but all zeros is your best lewis structure.
05:47
What that means is it's most likely to be the form this molecule takes the majority of the time.
05:53
Okay, these are all different resonance structures.
05:56
Most likely, most of the time, if you take a snapshot of molecule, it's in this format.
06:01
Now let's quickly look at the others.
06:03
Okay.
06:04
For the one where the triple bond is between the c and the o, that oxygen would have six valence, minus two dots and three sticks.
06:14
That gives you a plus one.
06:17
Already, that's a red flag.
06:20
You should never have a plus charge on your most electronegative atom.
06:24
Remember, electronegativity increases as you get closer to flooring.
06:28
So oxygen is our most electronegative atom here.
06:31
It does not want a plus charge.
06:33
What that means is it means it's lost an electron to somebody.
06:36
He came in with six, but he's only in sole position of five.
06:41
Now, don't get me wrong.
06:42
That doesn't mean he doesn't have a full octet.
06:44
He does have eight if you include who he's sharing with.
06:48
But one of them that he came in with has now moved to a different atom.
06:54
And it's not the atom he's sharing with.
06:57
He's lost control.
06:58
Since he's the most electronegative atom, he has the power to hold on to all of them.
07:03
And usually, or sometimes, steal from others.
07:06
He's not going to sit idly by while he's lost an electron.
07:11
He will just snatch it back.
07:13
So this lewis structure, while it probably exists, probably is not the form most of the time.
07:21
Okay, we'll do the others.
07:23
That carbon there would have four valence minus zero dots and four sticks, so that's a zero.
07:29
So that's not who stole the electron from them.
07:33
It must be the sulfur.
07:35
Sulfur has six valence, minus six dots and one stick.
07:39
That's where it is.
07:40
He's got a negative one charge.
07:42
He stole an electron from oxygen.
07:44
That's why these charges are the way they are.
07:47
Oxygen, however, is stronger, so it will steal it right back.
07:51
So this is not a good lewis structure...