00:01
We're looking at of4 or oxygen tetrafluoride.
00:05
We're going to draw the lewis structure for this of a structural formula.
00:09
So first we need to count up how many valence electrons are going to be in this image.
00:14
Quick reminder, if you don't remember, looking at the periodic table, the number of valence electrons an atom has is based on which column it's in, or at least that makes it easy to figure out.
00:26
It's actually based on its electron configuration.
00:28
But that's a shortcut.
00:31
The only exception to that is helium, which has.
00:35
All right.
00:36
So for oxygen, we're looking at six, valence electrons.
00:42
So we're going to put six.
00:44
And then fluorine has seven.
00:47
So seven times four, 28 plus six gives you 34 electrons to play with.
00:55
So that's how many electrons we can put in our image.
00:58
We're going to put oxygen in the center.
00:59
You generally put the atom listed first in the middle.
01:04
The scientific or the chemistry reason behind that is the least elective atom goes in the center, which in this case is oxygen.
01:14
We're going to surround that by four fluorines.
01:20
There we go.
01:23
To make sure we don't put too many electrons or that we get them all into the image, we're going to subtract off the electrons as we go.
01:29
So each bond has two electrons in it.
01:32
And there's four bonds there.
01:34
So that would be eight thus far.
01:40
34 minus 8 gives us 26 left to play with.
01:45
All right.
01:46
Next, you want to make the terminal atoms have full octets before you worry about the central.
01:52
So the ones on the outside, that's the fluorines.
01:54
Everybody, for the most part, wants to have eight electrons at least.
02:00
A couple of exceptions, hydrogen and helium.
02:03
We'll just want two.
02:06
Boron can have six or eight.
02:09
Either is fine.
02:11
And then anything in column three or lower can really do eight or more.
02:16
But really the goal here is eight.
02:17
Fluorine's a normal, typical atom.
02:20
It wants eight electrons.
02:21
So the bond touching each fluorine counts as two electrons.
02:25
So each fluorine currently thinks it owns two electrons, which means it wants six more.
02:32
So we're going to surround each flooring with six.
02:34
More electrons.
02:39
Six times four, that's 24, which drops us down to, oh, that doesn't equal to.
02:47
26, minus 24 is 2.
02:50
We've got two electrons left.
02:54
Those two electrons will go in the only place it has available, and that's on the central atom.
03:01
So those two electrons are going to go to the oxygen right there, which knocks out the rest of our electrons.
03:09
So there you have your structural formula or your lewis structure.
03:12
If we draw this to actually look like the actual shape of it, we have got five electron domains around this thing.
03:20
We've got four bonds and a lone pair.
03:24
That means that the real way this would look would be like a trigonal bipramidal.
03:39
And i suppose we should work our way into that.
03:41
So i'm not going to redraw it quite yet.
03:43
Let's do this in a decent order here.
03:46
Okay, so first off, the vesper class of the molecule.
03:55
I believe what it's asking for is something like this, where a stands for the central atom, x is how many bonds there are on that central atom, and e is how many electron pairs.
04:10
So we've got four bonds and an electron pair...