Write Lewis structures for the following. Show all resonance structures where applicable. a. NO2

Write Lewis structures for the following. Show all resonance...

Key Concepts

Lewis Structures

Lewis structures are diagrammatic representations that show the arrangement of electrons in molecules, including bonding pairs and lone pairs. They serve as a useful tool to determine how atoms are connected and to predict molecular geometry and reactivity by illustrating which atoms share electrons and how many electrons participate in each bond.

Resonance Structures

Resonance structures represent different ways to draw a molecule or ion where the distribution of electrons can vary while the placement of atoms remains constant. They are used to depict delocalized electrons and provide a more comprehensive picture of electron distribution, which is essential for understanding the stability and reactivity of molecules that cannot be adequately represented by a single Lewis structure.

Formal Charge

Formal charge is a bookkeeping tool used to estimate the distribution of electrons in a molecule. It is calculated by comparing the number of valence electrons of an atom in its free state with the number of electrons assigned to it in the Lewis structure, considering both bonding and lone pairs. Assigning formal charges helps in selecting the most significant resonance structures, typically the ones with the smallest formal charges, and in ensuring the overall neutrality or correct ionic charge of the molecule.

Octet Rule

The octet rule is a guideline stating that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell, achieving a stable noble gas electron configuration. This concept is fundamental to drawing Lewis structures since it informs the placement of electrons around atoms. However, there are exceptions, especially for molecules with electron deficiencies or expanded valence shells, which must be taken into account during structure determination.

Transcript

00:01 Problem 93 says to draw the lewis structures for each of the following molecules showing all resonance forms where they exist.

00:10 So for part a, we have a few different nitrogen and oxygen containing either ions or molecules.

00:20 The first one is no2 with a negative one charge.

00:25 And so if we think first about figuring out how many valence electrons we need to be able to show on our leo structure, we first consider that nitrogen has five valence electrons and each oxygen has six.

00:40 So if we add everything together and include the additional electron that makes this ion have a negative one charge, we have 18 total electrons.

00:55 So i'll first draw the resonant structures down here.

01:02 So n sits in the middle, nitrogen sits in the middle of this ion with two oxygen surrounding it, and we know that it has a negative one charge.

01:13 So what we can first do is we can draw bonds between the nitrogen and the oxygen, accounting for four electrons.

01:21 And because we know that nitrogen prefers to have to have.

01:26 Five valence electrons.

01:28 One of the easy ways to accomplish that is to double bond between one of the oxygens, giving nitrogen three valence electrons, and then give nitrogen a lone pair.

01:42 So we have a total of one, two, three, four, five.

01:46 Now oxygen, like i said, has six valence electrons typically.

01:52 And so giving this oxygen over here, two lone pairs, gives this one's six.

01:58 And so both of these have formal charges of zero since their valence electrons in this molecule match the number of valence electrons in the lone atom.

02:13 Now giving the negative one charge for this ion overall is the fact that this oxygen has three lone pairs of electrons, giving it seven total, seven total valence where the atom alone has six.

02:33 So the negative one formal charge over here is what gives this ion its charge.

02:40 So as you can imagine, then there are two different structures that you can draw.

02:46 The other one's simply altering the position of the double bond.

03:02 So for no3, we have three oxygens, each having six valence electrons, so that's 18, 19, and then nitrogen, as i said before, it has five.

03:13 So that gives us 24 total.

03:16 Maybe i should do that in a different color.

03:18 I'll do this in green.

03:23 So i'm thinking about no3 with a negative charge.

03:27 It's the same sort of concept where we have nitrogen bound to...

03:35 Three oxygens.

03:40 So in this case, it's a little bit difficult, a little bit more difficult to think about where each of the electrons are on the nitrogen is going to sit because there are three bonds, or three, three oxygens around the central nitrogen.

04:03 And that sort of lends to the possibility of either putting a lone pair of electrons here, but that would require that there are no double bonds or nitrogen, meaning that each of these individual oxygens would have a negative one formal charge, making this an overall negative 3 ion, which is not the case.

04:28 So we know that the nitrogen can't have a lone pair of electrons, which means that it's more likely that one of the oxygens has a double bond.

04:41 What this means for nitrogen is that while its octet is filled, it has a plus one formal charge.

04:47 So what would make sense then is there's a single double bond with one of the oxygens, and then each of the remaining oxygens has a negative one formal charge, giving the ion its negative one overall charge.

05:03 So the three resonance forms of this ion then include the three different possible positions of the double bond.

05:29 And finally we have n204.

05:32 So because there are two nitrogens, that means there are 10 valence electrons from the nitrogen and 24 total from the oxygens, making a total of 34 electrons.

05:44 Have to account for.

05:47 So because we know that according to the problem, this molecule is essentially 2 -02 is bound to each other.

05:59 That gives us a hint about how to draw this molecule.

06:07 So we know that there's a nitrogen -nitrogen bond and each of the nitrogen is independently bound to two oxygens...

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