Let's do the ClO2- Lewis structure.
Chlorine has 7 valence electrons.
Oxygen has 6, we have two Oxygens; plus, we
have a valence electron up here we need to
add in for a total of 20 valence electrons.
Chlorine's the least electronegative, it goes
in the center.
We'll put an Oxygen on either side.
Next, two valence electrons between atoms
to form chemical bonds; we've used 4, 6, 8,
10, 12, 14, 16, now back to the center to
complete the octet, 18 and 20.
So this looks pretty good.
Each of the atoms has an octet.
We've used all 20 valence electrons.
But Chlorine can be found in the third period
of the periodic table.
That means it can have more than eight valence
electrons, so we need to check our formal
charges.
When we calculate the formal charges, we can
see the Chlorine has a plus one formal charge,
and the Oxygen has a minus one.
And they're both symmetrical, so they both
have minus one.
That makes sense: minus one plus one, minus
one, that gives us this negative charge here.
But we want our formal charges to be as close
to zero as possible and still have this negative.
So let's see what we can do.
If we move a pair of electrons from the outside
here to form a double bond, that should get
rid of this +1 charge here for Chlorine.
So let's recalculate our formal charges and
see how that worked.
So when we recalculate, after moving these
valence electrons here into the center to
form a double bond, the Chlorine now has a
formal charge of zero.
The blue Oxygen right here didn't change,
it's still -1.
And then the Oxygen in black right here now
has a formal charge of zero.
This is a better structure because the formal
charges are closer to zero while still retaining
that negative one right there.
One last thing: we do need to put brackets
around this to show that it is a negative
ion.
And that's it.
That's the best structure for ClO2-.
This is Dr. B., and thanks for watching.
