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Current time:0:00Total duration:11:30

The previous video,
we saw some steps for drawing dot structures. In this video we're going to
use the same steps to draw a few more structures. But we're also going to talk
about how formal charge relates to dot structure. So we'll get back to this
definition in a minute. Right now, let's draw
a quick dot structure for the ammonium cation. So NH4 plus. The first thing you do
is find the total number of valence electrons. And so to do that you
look at a periodic table and find nitrogen,
which is in group five. Therefore, nitrogen has
five valence electrons. Right, each hydrogen has one. And we have four of them. So we have 5 plus 4,
giving us 9 electrons. However, there's
a plus 1 charge. Meaning this is
a cation, meaning we're going to lose
an electron here. So instead of representing
nine in our dot structure, going to represent
eight electrons. And so let's go ahead and put
the nitrogen at the center. Remember you put the
least electronegative atom at the center,
except for hydrogen. So nitrogen is going to
go in the center here. And we know it's going to
have bonds to four hydrogens, so we go ahead and put in
those hydrogens right here. And let's see how
many valence electrons we've used up in
our dot structure. Two, four, six, and eight. So that takes care of all
eight valence electrons that we were supposed
to represent. So this is the structure. And I can go ahead and put
some brackets around it here. And also a plus 1 charge to
indicate that this is an ion. And so that's the dot structure
for the ammonium cation here. Let's see if we can
assign formal charges to the nitrogen
and the hydrogen. So I'm going to go ahead and
redraw our dot structure here. And I'm also going to draw
in the electrons, right? We know that each of
those covalent bonds consists of two electrons. I'm going to go ahead and put in
those two electrons right here. And if I want to find a
formal charge for, let's say, the central nitrogen. What I would do is
think about the number of valence electrons
in the free atoms. So if you had a nitrogen
all by itself, right? You look at the periodic
table, it's in group five. And so therefore we're talking
about five valence electrons in the free atom for nitrogen. From that number we're
going to subtract the number of valence
electrons in the bonded atom. And the way to approach that is
to look at your dot structure here, and think about
those two electrons in those covalent bonds. One of them we're going
to assign to the hydrogen, and one of them are going
to assign to the nitrogen. And so we go around,
we do that for each one of our covalent bonds like that. And so now, we can
see that nitrogen is surrounded by four valence
electrons in the bonded atom. So let me go ahead
and write that. So it's 5 minus 4. And so 5 minus 4 is
of course plus 1. So we have a plus 1 of formal
charge on the nitrogen. So this nitrogen as a
plus 1 formal charge. Now let's do it
for hydrogen here. So hydrogen's in group
one on the periodic table. So let me just point this out. This is for . Nitrogen and then for hydrogen. It's in group one. So one valence electron
in the free atom. And from that we're
going to subtract a number of valence
electrons in the bonded atom. So if we look here, we assigned
one valence electron here to each hydrogen. So therefore, it's
just 1 minus 1, or 0. So there's 0 formal charge
for all the hydrogens in it the ammonium cation. And so that's how to
assign formal charges. Let's see how that applies
to actually affecting our final dot structure. And an example of that
would be something like sulfuric acid here, so. The first step, of course, is
to calculate the total number of valence electrons
we need to worry about in our dot structure. So, once again, each hydrogen
is one, I have two of them. Sulfur is in group six
on the periodic table, so therefore I have
six valence electrons. Oxygen is also in group six. And so we have six and we
have four of them right here. So 6 times 4 is 24. So we have 24 plus 6
is 30, plus 2 is 32. So we need to worry about 32
valence electrons in our dot structure for sulfuric acid. All right, next thing we do
is choose the central atom. And once again you
ignore hydrogen so it's between
sulfur and oxygen. And if you look at
the periodic table, you'll see that oxygen is higher
in group six than sulfur is. Therefore oxygen is
more electronegative. And so therefore, we're going
to put sulfur at the center. So we're going to put
sulfur right here. Once again, look at the
rules from the previous video if that didn't make
quite sense to you here. So we have sulfur
attached to four oxygens. And I'll go ahead and put
my four oxygens in there like that. And then I have two hydrogens. And by experience, you are
talking about an acid here. You're going to put your
hydrogens on oxygens. And so we're going to go ahead
and put our hydrogens here. And let's see how
many valence electrons we've used up a drawing
this skeleton here. So we have two, four, six,
eight, ten, and twelve. So we've used up 12 valence
electrons so 32 minus 12 gives us 20 valence electrons
left to worry about. And so, remember
the next step is to assign some of those
left over electrons to some of the terminal atoms. But again, we're not going
to assign those electrons to hydrogen because
hydrogen's already surrounded by two electrons. And so we're going
to try to assign some electrons to oxygen. And oxygen's going to
follow the octet rule. So let's examine, let's
say the top oxygen here. And we could see the top oxygen
is surrounded by two electrons already right there in green. And so if we're going to give
it an octet it needs six more. So we have one, two,
three, four, five, six. Same thing for this oxygen
down here, it needs an octet. So we go ahead and give it
six more electrons like that. Right so, we also have these
other oxygens over here to worry about. So let me go ahead
and use green again. So let's say this
oxygen over here on the left, the one bonded
to this hydrogen here. Here's two electrons and
here's another two for four. So for that option to have
an octet, it needs four more. So that means we're going to
put two lone pairs of electrons on this oxygen. And then we're going to do
the same thing for this oxygen as well. So let's see, how many electrons
did we just represent there? Well we had six
on the top oxygen, six in the bottom oxygen. That's 12. And then we had four on the
left and four more on the right. So that's eight. So 12 plus 8 is 20. So we've now represented
all of the valence electrons that we needed to show. And let's think about this as
possibly being the final dot structure. So we have an octet around
sulfur, an octet around oxygen, and hydrogen's fine. So you might think
that we are done here. However, let's go ahead and
assign some formal charges. And let's see what that does. So let me go ahead and draw
in some electrons here. So we know that each bond
consists of two electrons. I'm going to go ahead and
make them red here like that. And let's assign a formal
charge to the top oxygen here. All right, so this top oxygen. I'm going to, and the bond
between oxygen and sulfur. I'm going to give
one of the electrons to oxygen and one of the
electrons to the sulfur. And so I can see
that oxygen is being surrounded by 7 electrons. In the free atom, right? We would expect oxygen to
have six valence electrons, 6 minus 7 in this case gives
us a formal charge of -1. And so this top oxygen
has a formal charge of -1. It's the same situation
for this bottom oxygen here, so that one has a
formal charge of -1 as well. Let's look at the sulfur. So if we examine
the sulfur here, and we know-- right
sulfur is in group six on the periodic table. So normally six electrons
for the three atom. And in this bonding situation,
right let's go ahead, we know this one oxygen-- this
one electron I should say, goes to sulfur. And then in this bond
between oxygen and sulfur, sulfur is going to
get one electron. And so on, all the
way around here. So sulfur is surrounded
by four electrons in the bonded atom here. And so it's 6 minus 4, which
is a formal charge of plus 2. So this dot structure
might look like we're done, but we have a lot
of formal charges. We have -1, plus 2, and -1. And usually molecules
like to have-- like to minimize
the formal charge. And so if there's any way
to get this formal charge as close to 0 as possible, that
would be the preferred dot structure. And so let's go ahead and
redraw this really quickly. And let's see if we can
move some electrons around to minimize our formal charges. Now we can't add
any more electrons because we've
already represented all 32 valence electrons
that we were supposed to. So the only thing
that we can do is to share some more electrons. And so, if I took two electrons
from this top oxygen here. So if I took these
two electrons, and I move them into here. And if I took these two
electrons right here, and I move them into here
to form double bonds. Let's go ahead and look and see
what our dot structure would look like and assign
some formal charges. So now, I would have
sulfur double bonded to this top oxygen, and double
bonded to this bottom oxygen. Top oxygen has only two lone
pairs electrons are on it. Same thing for
this bottom oxygen. And then these oxygens are the
same, with the OH on the left. And then we have the OH
on the right, right here. OK so now let's look
at our formal charges. So we'll put in our electrons. So let's go ahead and do that. And once again, we're
going to do the same thing that we did before. So assigning formal charges will
start with the top oxygen here. So the top oxygen, right? Six valence electrons
in the free atom. And then if I go like that,
you can see there are six here. So 6 minus 6 gives us
a formal charge of 0. So the top oxygen is 0 now. Same thing with this bottom
oxygen, this one's 0. Let's look at the sulfur now. So normally, we're talking
about 6 for the sulfur. And let's see how
many valence electrons are on the bonded atom here. So we do the same
assigning of electrons that we've done before. And so now we can see that that
sulfur is surrounded by six. So we would go 6 minus 6
gives us a formal charge 0. And if you assign
a formal charge to one of these other oxygens. So let's go ahead and
do that really quickly. So one of these other
oxygens right here. So I'm doing the
oxygen on the right. OK, so let's see how many
valence electrons are surrounding this atom. So we have a total
of six is well. So 6 minus 6 is a
formal charge of 0. So this dot structure actually
gives us formal charges of zero for everything. And so this would be
the preferred one. And let's finally
talk about octets. Right so we know that hydrogen--
we know the hydrogen's happy surrounded by two electrons. We know that oxygen is
happy with an octet. And so you can see that all of
these oxygens have an octet. Sulfur in this case, is
not surrounded by an octet. It has an expanded outer shell. Right, it's OK for sulfur have
an expanded valence shell, because it's in the third
period on the periodic table. And so we talked
about why that's OK in the previous video. And so this is
just one more thing to think of when you're
drawing dot structures. Sometimes formal charge will
affect the final structure of your molecule or ion.