nitrate lewis structure

Moderators: Chem_Mod, Chem_Admin

705340227
Posts: 86
Joined: Wed Sep 30, 2020 10:08 pm

nitrate lewis structure

Postby 705340227 » Sun Nov 08, 2020 9:37 pm

Why couldn't nitrate have two double bonds with oxygen molecules and one single bond with the last oxygen molecule that would have a -1 charge? I believe that all the electron totals still work and all of the atoms have an octet.

Lucy Wang 2J
Posts: 86
Joined: Wed Sep 30, 2020 10:09 pm
Been upvoted: 2 times

Re: nitrate lewis structure

Postby Lucy Wang 2J » Sun Nov 08, 2020 9:48 pm

I think it has to do with formal charge. If two of the oxygen had a double bond, then the nitrogen would have 5 bonds which is not very stable.

Ivan Chen 2H
Posts: 88
Joined: Wed Sep 30, 2020 9:48 pm

Re: nitrate lewis structure

Postby Ivan Chen 2H » Sun Nov 08, 2020 9:49 pm

The nitrogen would have 5 bonds, which violates the octet rule. Second period elements always have to follow it without exceptions.

Simrah_Ahmed1J
Posts: 88
Joined: Wed Sep 30, 2020 9:49 pm

Re: nitrate lewis structure

Postby Simrah_Ahmed1J » Sun Nov 08, 2020 9:51 pm

I believe it is because that is to many bonds for nitrogen, 10 instead of 8. That energy level would technically be favorable but not possible based on nitrogen bonding abilities

SashaAnand2J
Posts: 85
Joined: Wed Sep 30, 2020 9:37 pm

Re: nitrate lewis structure

Postby SashaAnand2J » Sun Nov 08, 2020 9:54 pm

Two oxygen double bonds doesn't work for nitrate because we want to ensure that the most stable version of the structure is being formed. Having a 5 bond nitrogen would break the octet rule, because nitrogen doesn't have a d-orbital to accommodate more valence electrons. This is why instead, a formal charge of +1 is placed on the nitrogen atom, and 2 charges of -1 on the single bonded oxygen atoms.

Natalie 3k
Posts: 86
Joined: Wed Sep 30, 2020 10:11 pm

Re: nitrate lewis structure

Postby Natalie 3k » Sun Nov 08, 2020 9:58 pm

Hi! I believe this is because it would cause nitrogen to have 5 bonds, the two double bonds and a single bond, which wouldn't be very stable.

Adrienne Chan 1G
Posts: 87
Joined: Wed Sep 30, 2020 9:40 pm

Re: nitrate lewis structure

Postby Adrienne Chan 1G » Sun Nov 08, 2020 9:58 pm

There's a problem with formal charge in this case! Although it might technically follow the rules, the individual atoms don't have the optimal formal charge.

Ryan Hoang 1D
Posts: 95
Joined: Wed Sep 30, 2020 9:49 pm

Re: nitrate lewis structure

Postby Ryan Hoang 1D » Sun Nov 08, 2020 10:04 pm

Everyone here is correct: but the reason why it breaks the rule and something like sulfur doesn't is because nitrogen, being n=2 only has l=1, and ml=-1,0,1. Because of this, the maximum amount of electrons that can be filled into nitrogen are the p orbitals, the px, py, and pz orbitals (2 electrons in each orbital, 6 in total+ 2 from s orbital=8 electrons max). Because there are only 8 electrons max, we can only have a maximum of 4 covalent bonds in total. However, with something like sulfur, it can break this rule because it's in the n=3 period, with l=2, ml=-2,-1,0,1,2. Because l=2, this means sulfur can have d orbitals. While sulfur, in completing its octet doesn't actually normally use these d orbitals, when it needs to have 5 bonds (10 electrons) for example and expand its octet, it'll access these d orbitals and use d orbitals to bond with other elements.


Return to “Lewis Structures”

Who is online

Users browsing this forum: No registered users and 1 guest