Question

Why is the negative charge on the oxygen atom with the single bond?

My question is that why is the negative charge on the oxygen atom with the single bond, why not on the oxygen atom with the double bond? I am confused.

Answer 1

This is a good question, and illustrates the formalisms we use when we assign `"formal charge...."`

Explanation:

Now ozone is a NEUTRAL molecule, and the Lewis structure indeed reflects this even tho there is charge separation. We follow the old convention that the TWO electrons of the covalent bond are SHARED EQUALLY between the atoms for the porpoises of charge assignment...

And so for `:stackrel(ddot)O=stackrel(ddot)O:`, EACH oxygen atom claims one electron from each covalent bond, and owns the lone pairs that it bears. And so for `O_2`, each oxygen atom has 6 valence electrons, and has 2 inner core electrons, that are not conceived to be involved in bonding (formally these are the `1s^2` set). And so each oxygen atom is neutral, because there are the eight electrons to balance the 8 positive charges in each oxygen nucleus....

And now for `"ozone"`....`""^(-)O-stackrel(ddot)O=stackrel(ddot)O:`, from left to right as we face the page, around each oxygen atom, there are NINE, SEVEN, and EIGHT electrons formally associated with that centre leading to FORMAL charges of `-1`, `+1`, and `0`.

The electronic structure of ozone is thus based on a trigonal plane, but the molecular structure is BENT, with `/_O-O-O<=120^@`. This site quotes `/_O-O-O<=116.8^@`, (i.e. reduced from the ideal trigonal planar angle of `120^@` with `O-O=1.278xx10^-10*m`, which is slightly longer than the bond length we observe in dioxygen, `1.210xx10^-10*m`.