I have a homework question that says:
"Three resonance structures of the following anion are possible. One is given below, but it is incomplete. Complete the given structure by adding non-bonding electrons and formal charges. Draw the two remaining resonance structures, including non-bonding electrons and formal charges."
The molecule they give us looks like this (but in a skeletal structure):
O(-)-c--c-c--c-c
|
H
The (-) is the formal charge on O, the - means a single bond, and the -- is a double bond.
I know there must be lone pairs on the O to give it its negative formal charge, but since the structure is in bond-line form and it says it's missing pieces, do I assume the there are Hydrogens attached to each one of those carbons or are there supposed to be lone pairs?
If there is supposed to be lone pairs, that would make each carbon have a formal charge which I don't think would be right. But if there are Hydrogens attached, when I try to form resonance structures, I don't know where to move electrons without exceeding the atoms' octet.
Could anyone help me with how to draw a resonance structure for this molecule??
Thanks!
That H is supposed to be attached to the first Carbon, it just got shifted when I posted the question!
To draw resonance structures for the given molecule, it is important to understand the concept of resonance. Resonance occurs when there are multiple ways to arrange the electrons in a molecule, resulting in the delocalization of electrons over different atoms or bonds.
In the given molecule, the double bond can actually be moved between the two carbon atoms on the left. This means that the structure can be represented in multiple ways, known as resonance structures.
To start, let's complete the given structure by adding non-bonding electrons and formal charges. Since the oxygen atom has a negative formal charge, it must have an additional lone pair of electrons. The complete structure is as follows:
H
|
O(-)=C-C-C=C-C
Now, let's draw the remaining two resonance structures. We can move the double bond to adjacent carbon atoms one at a time while rearranging the lone pair of electrons on oxygen or the position of hydrogen. Also, remember to keep the octet rule in mind; the atoms should not exceed their usual number of valence electrons.
Resonance structure 1:
H
|
O=C-C-C(-)=C
In this structure, the double bond has shifted to the rightmost carbon atom, and the oxygen atom now has a single bond. The carbon atom with the negative formal charge now has three lone pairs of electrons.
Resonance structure 2:
H
|
O=C-C(-)-C=C
In this structure, the double bond has shifted to the carbon atom on the left, and the oxygen atom has a single bond. Again, the carbon atom with the negative formal charge now has three lone pairs of electrons.
It is important to note that resonance structures represent the different ways electrons can be delocalized in a molecule, and the actual structure of the molecule is a hybrid of these individual resonance structures. This means that the molecule does not exist in any one particular resonance structure but rather represents an average of all possible resonance structures.
I hope this explanation helps you draw the resonance structures for the given molecule.