Leah here, from Leah4sci.com.
And in this video we're going to look at simple resonance structure practice.
These questions are from the resonance practice quiz, which you can find on my website along
with the entire tutorial series Leah4Sci.com/Resonance.
The first question asks for the resonance structures of Nitrate.
If you follow my Lewis tutorial video linked below, you'll see that we start out with a
nitrogen surrounded by 3 oxygen atoms.
Where nitrogen should be double bound to one oxygen and single bound to two others.
The double bound oxygen has two lone pairs.
The single bound oxygens each have three lone pairs.
Giving us a formal negative charge on a single oxygen and a positive charge on nitrogen.
When you're thinking how to start resonance, you'll always want to look for your least
stable most negative electrons and you wanna push them towards the least stable positive
atom.
In this case we'll take the electrons from the oxygen on the right or left, it doesn't
matter since the molecule is symmetrical, they're exactly the same.
We show our arrow starting at the electrons and moving towards the positive atom.
This will cause a double bond to form between oxygen and nitrogen giving nitrogen too many
electrons in its octet.
To compensate, we'll take the pi electrons binding nitrogen to oxygen and kick them off
where?
Onto the oxygen atom below.
Don't forget the double headed arrows to show that the resonance can happen back and forth
and redraw the structure.
Oxygen on the upper right now has a double bond to nitrogen.
And oxygen on the bottom now has a third lone pair moving the negative charge from the upper
right oxygen to the lower oxygen.
Nitrogen still has a positive charge because it still has 4 bonds.
If we show the oxygen that just became negative attacking the nitrogen kicking out those electrons,
we'll simply go back to the structure we started with.
Instead, we now want to involve the other negative oxygen and show its electrons coming
out to attack the nitrogen causing the bond between the double bound oxygen and nitrogen
to collapse onto oxygen.
Once again a double headed arrow to show that it's moving back and forth.
We now have a double bond between the upper left oxygen and nitrogen and a third lone
pair on the upper right oxygen atom.
Once again, the single bound oxygens are negatively charged and nitrogen is still positively charged.
These are the three major contributing structures but there's one more that you can show for
your minor structure.
If we start with this structure at random, instead of showing oxygen attacking nitrogen,
we can simply show the electrons from a pi-bond collapsing onto oxygen without reforming a
double bond anywhere.
This will give us the structure that has nitrogen single bound to three oxygen atoms where each
oxygen has three lone pairs and a negative charge.
You can follow the colors to see where each pair came from.
The problem with the structure and the reason why this is a minor structure is the nitrogen
in the middle.
If you look at the nitrogen, it only has three bonds for a total of six instead of eight
electrons in its octet.
It gets a formal charge of positive two and has an incomplete octet.
These are two very very unhappy situations making this molecule very very unstable.
Each of the above structures has a positive Nitrogen plus one rather than a plus two and
every atom has a complete octet so they're stable in that respect.
That gives us a total of three major and one minor resonance structures for the nitrate
ion.
Next example we'll look at is the carbonate ion, which is CO3 2-minus.
We start by drawing the skeleton very similar to nitrate except that here, given two negative
charges, something is going to be different.
That difference comes from the central carbon.
Nitrogen with four bonds will have a positive charge, carbon with four bonds is happy and
has no charge.
This means when we have two negative oxygen atoms we have nothing to cancel on the positive
side and so the net charge is negative two.
The resonance for carbonate will look exactly like the resonance for nitrate.
We'll start with the lone pair on a negative oxygen atom and bring their bond towards carbon,
collapsing the double bond between carbon and oxygen onto the oxygen atom.
Double headed arrow to show that its resonance.
And let's see what changed.
We now have double bond between carbon and the upper oxygen and an extra lone pair on
the lower oxygen.
Add your formal charges and we’re good.
Now move off the third oxygen and bring its electrons down forming a bond between carbon
and oxygen and collapsing that newly formed bond back off onto the oxygen atom.
This gives us a double bond between carbon and the upper left oxygen and a extra lone
pair on the upper right oxygen.
Once again with two negative charges.
These are your three major contributing structure but like nitrate, we can also find a minor
structure.
If we take this structure and simply break the pi-bond, collapsing the electrons unto
oxygen, we get a carbon atom single bound to three oxygen atoms which each have a negative
charge and the central carbon has only three bonds, six in its octet with a formal charge
of plus one.
You'll see carbocations in organic chemistry but their relatively unstable, especially
when bound to three negative oxygen atoms.
The negative oxygen would want to come back down and attack the carbon making this structure
very unstable and only a minor contributing structure.
Our final example is the chlorate ion.
Are you starting to see a pattern?
We have a central atom bound to three oxygens.
Given a net negative charge, the question is how many pi-bonds do we make?
And with chlorine, we have the exception to the octet rule and that means we can have
more than four bonds or eight electrons on that central atom.
This gives us a chlorine double bound to two oxygen atoms and single bound to one.
Oxygen with a double bond and two lone pairs has no formal charge, oxygen with a single
bond and three lone pairs has a charge of negative one, but chlorine is where it gets
interesting.
If you count up the total number of electrons you'll see where we’re actually missing
two and that's because in addition to five bonds on chlorine, we have a lone electron
pair.
This is an exception to the octet rule because with five bonds and two lone pairs, chlorine
has a total of twelve electrons in it's octet.
Even with the extra electrons, it still has a formal charge of zero as explained to the
tutorial linked below.
And that’s the trickiest aspect of this question: once you figure out how to draw
your starting structure, the rest of it will look exactly like the nitrate resonance except
that we have two double bonds to account for.
We'll show the electrons from the negative oxygen reaching towards chlorine to form a
pi-bond and as a result, chlorine can only have twelve.
Adding two more electrons is too much so we have to break one of the pi-bonds onto a nearby
oxygen atom.
The new structure looks exactly the same except that the upper oxygen is negative and the
lower oxygen is neutral.
The negative oxygen will now do the same thing, reach down and form a bond to chlorine but
this time, we'll show the left oxygen breaking it's pi-bond with chlorine and the electrons
collapsing unto oxygen.
What do we have?
A double bond between the upper oxygen and chlorine, a single bond between the lower
left oxygen and chlorine, and a extra lone pair with a negative charge.
In each of these structures we have a chlorine atom that's neutral and one negative oxygen.
These are each major contributing structures and just like before, we can show a minor,
less stable structure let's justify that.
If I take this structure and simply break one of the pi-bonds between chlorine and oxygen,
without reforming another pi-bond, I'll get a structure that looks like this: an extra
lone pair on the lower right oxygen with a negative charge and chlorine now with two
single bonds, one double bond, and one lone pair.
It looks like it should be happy but pay attention to the formal charge, chlorine should have
seven, I have one, two, three, four, five, six!
Chlorine has a charge of plus one.
Think about what you know of halogens, do electronegative halogens want to have a positive
charge?
Absolutely not!
That makes this molecule unstable and that makes this a minor contributing structure.
If you want to take this to step further and break that second pi-bond unto oxygen, this
structure’s even worse: we have three negative oxygens and a central chlorine with a formal
charge of plus two.
In resonant structures, you want to avoid separation of charge wherever possible and
you want to avoid giving an atom more than one charge.
A positive two on an electronegative chlorine makes this an absolutely terrible, terrible,
structure.
We'll call it a super minor resonance contributing structure.
This was just questioned one on the resonance quiz.
If you found this helpful, make sure you try the entire quiz which you can find on my website
along with the resonance tutorial series, Leah4Sci.com/Resonance