00:01
We've got two different compounds that we're going to be doing the lewis structures of.
00:05
I figured we'd start with the easier one, which is sulfate, so4, with a minus 2 charge.
00:11
I've already got the periodic table up labeled with how many valence electrons each column has, and then also the equation for formal charge down at the bottom.
00:21
All right, so to begin, we need to add up how many electrons, how many valence electrons are in our compound.
00:27
For starters, we have sulfur.
00:29
Sulfers in the sixth column.
00:31
So it has six valence electrons.
00:34
Oxygen is right on top of it, so it also has six, and there are four of those.
00:40
So six times four.
00:41
Lastly, the minus two charge means that two additional electrons have been added to the system.
00:48
So we'll add two more to that for a total of 32.
00:53
All right, we need that because we need to track how many electrons we have as we go.
00:59
Next, we'll go ahead and set up sulfur in the middle.
01:03
The atom written first generally goes in the middle unless it's hydrogen.
01:07
The rule behind that is that the least electronegative atom goes in the center.
01:14
And then it'll have four oxygens around it.
01:19
Now, if we count up electrons so far, if there's two in each bond, then we have used up eight electrons.
01:26
So that leaves us with 24 to go.
01:30
We want to worry about the outer atoms before we focus on the inner.
01:33
So we're caring about the oxygens.
01:36
Each oxygen currently thinks it has two electrons from sharing two in each bond, so they each need six more.
01:45
Five six, six, and one, two, three, four, five, six.
01:50
Six times four, that is 24 electrons.
01:55
So we are out of electrons.
01:57
Now, it's not uncommon at this point to think that you're done, and based on the normal rules, everybody has an octet, everybody is fine.
02:04
So it's not a bad assumption.
02:06
But if you start to check formal charges on this thing, it actually isn't very stable as is.
02:15
So let's see the formal charges, and then i'll show you we're going to have to add some double bonds to get this thing stabilized a bit.
02:22
So down in the bottom right, i've written that the formal charge of a single atom equals how many valence electrons that atom has.
02:29
That's just from the periodic table.
02:30
Minus half of its bonding electrons.
02:34
Bounding electrons are those in bonds, minus the non -bonding electrons, which are the individual dots.
02:42
An easier way to remember this, if this is helpful to you, formal charge equals the valence electrons, minus the sticks on that atom, minus the dots on that atom.
02:56
It's just a bit easier to remember.
02:59
So, starting with our outer oxygens here, that oxygen would start its life with six valence electrons.
03:07
It's got one bond on it, so that's one stick, minus six dots equals a minus one.
03:15
And each of those oxygens are all in the same situation at the moment.
03:19
So all four of them would be minus one.
03:24
And then we can look at the sulfur.
03:27
Sulfur starts with six valence minus four sticks and no dots.
03:34
So it's a plus two.
03:37
Now, an ideal situation has zero for the formal charges, and if it has to deviate from zero plus and minus one, this is about as far as you want to go.
03:48
You only want to do plus two or minus two if you absolutely have to.
03:53
And in this case, we can make this better.
03:57
And we're going to play off the idea that if you are at least in row three or lower, if you are in row three or lower.
04:10
You can have more than eight electrons.
04:15
You can actually go up to 12.
04:21
Yes, you could technically go up as high as 12.
04:27
And the reason for that is because these atoms have access to de -orbital at this point.
04:33
This is a little bit of a deviation.
04:35
But if you want to know the background, that sulfur there, his electron configuration would be neon.
04:42
And then 3s2 and 3p4.
04:48
But at this point, he does have access to the 3d orbitals.
04:53
There's nothing in them to start with.
04:55
But he does have access to them.
04:56
So if he's offered more electrons than will fit in his p orbitals, then he can put them up into the d block.
05:02
All right.
05:03
So because of that, we can break the octet.
05:05
We can have more than 8.
05:07
That's what we're going to do.
05:09
So since that sulfur is a plus 1, okay.
05:12
What we want to do is if we, or that sulfur was a plus two, i'm sorry.
05:17
If we take one of the oxygen's lone pairs and move them down there, okay, the resulting molecule, let's draw it down here underneath.
05:36
Looks like this.
05:42
Now, the oxygens with one bond have not changed.
05:46
They're not zero, they're minus one.
05:51
So these guys are still the same.
05:55
This oxygen now that has the double bond is now six minus two.
05:58
Bonds minus four dots is zero.
06:02
So that's actually better for the oxygen.
06:04
It's more stable when it has a double bond rather than just a single.
06:08
And now the sulfur in the middle is six minus one, two, three, four, five minus zero.
06:16
So it's now a plus one.
06:19
So we can actually do this one more time and get the sulfur all the way down to zero...