4.17

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Marcela Udave 1F
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Joined: Fri Sep 28, 2018 12:16 am

4.17

Postby Marcela Udave 1F » Tue Nov 27, 2018 10:38 pm

Predict the bond angles at the central atom of the following molecules and ions: (a) ozone, O3; (b) azide ion, N3 ; (c) cyanate ion, CNO ; (d) hydronium ion, H3O .

For part a.) For the lewis structure of O3, why isn't there double bonds connecting both the oxygen to the oxygen with the lone pair? All the examples I see have a double bond and a single bond.

Neil Hsu 2A
Posts: 61
Joined: Fri Sep 28, 2018 12:16 am

Re: 4.17

Postby Neil Hsu 2A » Tue Nov 27, 2018 10:57 pm

There isn't double bonds on both terminal oxygens because if you count the number of electrons, 6x3 = 18 electrons. This means that the structure that you proposed is impossible; if there were double bonds on both terminal oxygens, there would only be a total of 16 electrons. Thus, there must be a double and a single bond with a lone pair on the central oxygen to have 18 electrons (having the correct structure).
Last edited by Neil Hsu 2A on Tue Nov 27, 2018 10:58 pm, edited 1 time in total.

KylieY_3B
Posts: 31
Joined: Fri Sep 28, 2018 12:24 am

Re: 4.17

Postby KylieY_3B » Tue Nov 27, 2018 10:57 pm

The formal charge is closer to 0 for the central oxygen atom when it has a lone pair and one double bond and one single bond.

Christopher Tran 1J
Posts: 77
Joined: Fri Sep 28, 2018 12:15 am

Re: 4.17

Postby Christopher Tran 1J » Tue Nov 27, 2018 11:04 pm

Couldn't there be double bonds connecting the two oxygen atoms with two lone pairs to the central atom, along with one lone pair on the central oxygen? That would lead to 18 total electrons and a formal charge of zero for all three oxygen atoms. Or does the central oxygen also have to obtain a full octet?


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