I've been working on the textbook 2B problems and I want to know how we can determine where the double bond is placed in a Lewis Structure. For example, the double bond for molecule ONF is shown to be between N and O, but how come it cannot be between N and F. Similarly, when looking at ClNO2, the double bond resonance is shown between N and O, but can there not be a structure with N and Cl? Or is it that there can be a double bond in these places, yet formal charges are not favored?
Thank you!
Determining Double Bonds (2B textbook)
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Re: Determining Double Bonds (2B textbook)
I think the most stable form between ONF has a double bond between O and N because of formal charges, because a single bond between F and N has a formal charge of 0 for F but a double bond between F and N has a formal charge of +1.
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Re: Determining Double Bonds (2B textbook)
There are many different lewis structures you can draw for each molecule, but you need to keep in mind that you should draw the most stable structure. So you should calculate the formal charge for each atom (using the formula FC = # valence electrons - (lone pairs + shared electrons/2) ) and draw the double bonds so that the formal charges are closest to zero.
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Re: Determining Double Bonds (2B textbook)
Hello! It all depends on the formal charge for the atoms, and having the most stable amount of electrons in each atom's valence shell. F only needs 1 electron to be at its most stable point, so a double bond would give an extra, unneeded electron, whereas a double bond between O and N makes both formal charges equal to 0, which is the most stable point.
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Re: Determining Double Bonds (2B textbook)
It's best to keep in mind the formal charge of any particular molecule. Having a molecule with a collective charge of 0 or close to 0 indicates that that particular structure is more stable than any other possible structure. More stable structures are favorable, so having a F==N bond would produce a charge of 1+ while having the double between O==N results in a balanced 0 charge.
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