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Kaylin Krahn 1I
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Postby Kaylin Krahn 1I » Fri Nov 03, 2017 11:01 pm

Why wouldn't ClO2 have double bonds instead of single bonds? (The solution manual shows the answer with single bonds)

Dylan Davisson 2B
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Re: 3.67(b)

Postby Dylan Davisson 2B » Fri Nov 03, 2017 11:55 pm

The molecule sticks to single bonds because if its Lewis structure was oriented any differently, the total formal charge would deviate from 0. And since the molecule has no charge, this would not be possible. If one double bond is used, the total formal charge would be +2 (this has resonance, but that does not change the level of total formal charge). And two double bonds cannot be used since that would give Chlorine an octet, but because there are an odd number of electrons and Oxygen obeys the octet rule strictly, there would be no place to put the electron with unpaired spin. Thus, the only possible Lewis structure configuration would be one of single bonds.

peytonruiz 1H
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Re: 3.67(b)

Postby peytonruiz 1H » Sat Nov 04, 2017 12:01 am

Hi Kaylin.
If you calculate formal charge, ClO2 does not have double bonds.
With single bonds (the model shown in the solutions manual), the formal charges of each oxygen is -1 (so together, their charge is -2). The central chlorine atom has a formal charge of +2.
This gives a total charge of 0, which cannot be achieved through double bonds.

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