Stable lewis structures

(Polar molecules, Non-polar molecules, etc.)

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Husnia Safi - 1K
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Joined: Fri Sep 29, 2017 7:04 am

Stable lewis structures

Postby Husnia Safi - 1K » Sun May 20, 2018 10:58 pm

So my TA mentioned something about when there is a formal charge of a -1 and a +1, a bond can be added to "even it out". Can someone explain to me why this is? Also Im not sure if I remembered this correctly so sorry if what I said is inaccurate.

Chris Fults 1C
Posts: 27
Joined: Wed Nov 15, 2017 3:01 am

Re: Stable lewis structures

Postby Chris Fults 1C » Sun May 20, 2018 11:06 pm

Do you mean the way of double or triple bonds? For example in single bonded N-N there would be 6 lone pairs for each carbon which means the formal charge would be -2. If you were to triple bond nitrogen there would be 2 lone pairs and 3 shared electrons decreasing the formal charge to zero, therefore, making it more stable.

Jennifer Ma 1G
Posts: 31
Joined: Fri Apr 06, 2018 11:03 am

Re: Stable lewis structures

Postby Jennifer Ma 1G » Tue May 22, 2018 12:19 pm

By adding another bond, you are able to make the formal charge of the molecule closer to 0, making it more stable. To help you figure out whether two atoms need another bond, just calculate the formal charge on each atom, and keep adding bonds wherever it is necessary in order to make the formal charge of the molecule the closest it can to 0, or even at 0.

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