At the end of this video, you’ll be a pro
at drawing the basic Lewis structures. There really are only five main steps that you need
to know. Let’s get started. Step 1: find the total number of valence electrons for the
molecule you’re drawing the Lewis structure for. Let’s try one. We’ll use the periodic
table to figure out how many valence electrons each element has. Hydrogen is in Group 1 on
the periodic table, so it has 1 valence electron, but the subscript tells us we have two Hydrogens,
so we’ll need to multiply that by 2. Oxygen is in group 6 or 16; it has 6 valence electrons.
When we add that up, 2 + 6, we get a total of 8 valence electrons for H2O. Remember:
valence electrons are the outer shell electrons. They’re the ones that form chemical bonds
and interact with the world around them. Let’s try one more: cyanide. For cyanide, we can
look on the periodic table again. Carbon has 4 valence electrons and Nitrogen has 5. The
trick for this one, is, though, it’s a negative ion. This negative sign right here tells you
there’s going to be one extra valence electron. So 4 + 5 + 1: 10 valence electrons for the
cyanide ion. Step 2: put the least electronegative atom in the center. So how do you know it’s
the least electronegative atom? Just remember that Fluorine is the most electronegative,
followed by Oxygen. So as you move away from Fluorine and Oxygen, atoms become less electronegative.
That means those atoms will go at the center. Let’s see how it works: When I draw the Lewis
structure for something like NO2, I can see that Nitrogen is further away from Fluorine
than Oxygen. That means when I draw my Lewis structure I’ll put Nitrogen at the center,
and Oxygens on the outside. For something like PCl3, I can see that Phosphorus is further
away from Fluorine than the Chlorine atom. So I’ll put Phosphorus at the center and the
Chlorine atoms on the outside, like this. Some Lewis structures, like HCl, only have
two atoms. In that case, you don’t have to worry about it, because there’s no center.
An important note: Hydrogen always goes on the outside of Lewis structures. Always. Seriously.
Let’s look at HCl again. We know it has 8 valence electrons, and we’ll put the first
2 in our structure between the H and Cl to form a chemical bond. That’s important: by
putting those two electrons there, we’re showing that the Hydrogen and the Chlorine are chemically
bonded. Step 4: Complete the octets on the outside atoms. Octet means 8. For HCl, we
have a total of 8 valence electrons–one for the Hydrogen, 7 for the Chlorine. So we’ve
used 2 already to form the chemical bond. For the Chlorine, we’ll put 4, 6, and then
8. We’ve used all the valence electrons and we’ve completed the octet for Chlorine–it
has 8 valence electrons. For basic Lewis structures, the big exception is going to be Hydrogen.
It only needs 2 valence electrons for a full outer shell. So Hydrogen has a full outer
shell with two valence electrons. And Chlorine, its octet is full with 8, and we’ve used only
the 8 valence electrons that we have for the HCl molecule. We’re done with this structure.
There’s one more rule that we need to look at. It doesn’t apply to HCl. Actually, it
doesn’t apply to a lot of molecules. But what happens when you’ve used all your valence
electrons and you still haven’t completed the octets for each atom in the molecule?
That’s step 5. So sometimes you run out of valence electrons and still haven’t filled
the octets. Hey, it happens, right? So, in Step 5 what we’re going to do is, if we run
out of valence electrons, we’ll move some of the valence electrons from the outside
to the inside and share them, and form double and triple bonds. Here’s how it works: So
for O2, Oxygen, if I follow all the rules, I count the valence electrons. I have 12.
I put a bond between the atoms, and then fill the outer shell, I end up with a situation
where I have an octet for this Oxygen–it has 8–but this Oxygen only has 6 valence
electrons. It doesn’t have an octet, and I don’t have any more valence electrons. To
solve the problem, I can take these valence electrons right here and move them to the
center to share them and form a double bond between the two Oxygen atoms. I’m still only
using 12 valence electrons, but now this Oxygen has 2, 4, 6, 8 valence electrons. And this
Oxygen has 2, 4, 6, 8 valence electrons. So by sharing those valence electrons, I can
get the octet on each of the Oxygens, and still use the 12 valence electrons that I
had for the O2 molecule. If I wanted, I could replace these right here with two lines to
show the double bond, where each line represents a pair of electrons. For Lewis structures
like N2, I would even use a triple bond to achieve octets. So how do you get good at
drawing Lewis structures? The first thing is, know the steps. They’re your guide. The
second thing is practice. You need lots of practice to get good at this, for it to be
second nature. Use the links in this video, or go to my website and try more videos. First,
try it yourself and then watch the video. Pause when you’re stuck, see what’s going
on there, and try to correct the problem. This is Dr. B., and thanks for watching.