These have two or more resonance structures. Write all the resonance structures for each molecul

step 1 We're drawing the Lewis structure for SO3. To begin, we want to know how many valence electrons

These have two or more resonance structures. Write all the...

Key Concepts

Resonance Structures

Resonance structures are different valid Lewis formulas for a molecule or ion that depict the delocalization of electrons. They do not represent multiple coexisting forms but rather a hybrid that averages over the electron distributions shown in the individual drawings.

Electron Delocalization

Electron delocalization refers to the spreading out of electrons over two or more adjacent atoms, which increases the stability of the molecule or ion. It is a key factor that allows resonance structures to exist, as electrons can move between positions without altering the overall connectivity of the atoms.

Formal Charge

Formal charge is a bookkeeping tool used to assign charges to atoms in a molecule based on the difference between the number of valence electrons in the free atom and the number assigned in a particular Lewis structure. In resonance structures, the distribution of formal charge helps determine which forms contribute most to the resonance hybrid.

Lewis Structures

Lewis structures are diagrams that represent the arrangement of atoms, bonds, and lone pairs of electrons in a molecule or ion. They form the basis for constructing resonance structures by showing how the electrons can be redistributed without changing the atom connectivity.

Octet Rule

The octet rule is a guideline that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell. When drawing resonance structures, ensuring that atoms (especially central atoms) satisfy the octet rule is important for obtaining a stable electron configuration.

Transcript

00:01 We're drawing the lewis structure for so3.

00:04 To begin, we want to know how many valence electrons we've got that's going into our image.

00:09 So i've already written them out on a periodic table to the right based on columns.

00:14 So anything in the first column has one valence electron, second column has two, etc.

00:19 So first we've got a sulfur, which has six valence.

00:24 And then we have oxygen, which also has six, and there's three of those.

00:29 So we're looking at a total of 24 electrons to go in our picture.

00:34 We'll put sulfur in the middle and oxygens around it.

00:39 It doesn't matter if you do it like on, you know, it can be on slanted sides or you could have it perfect like this.

00:45 It doesn't make any difference as long as they're all connected.

00:51 And that will have taken since there's three bonds and two in each bond, we've used up six electrons so far.

00:58 So we've got 18 to go.

01:00 Next, we'll fill up the outer oxygens before we worry about the sulfur.

01:05 Everybody wants to have at least eight electrons.

01:08 Each oxygen currently thinks it has two.

01:12 So they need six more.

01:19 Six times three, that's 18.

01:21 We are now out of electrons.

01:24 We've used them all.

01:25 All right, but sulfur is not satisfied yet.

01:29 It needs at least eight, which means since we can't add more electrons, we've got to do a double bond.

01:34 So we're going to take one of these loan pairs and move them to there creating a double bond.

01:52 All right.

01:54 So now, talking about resonance structures, we could have put the double bond in any of those three spots.

02:03 Right? we could have put it in the bottom right and the top or the bottom left, which means that this gives us our first three resonance structures.

02:13 We could have put it here instead, which would have made it look like this, or we could have put it down here.

02:31 All three spots are equally likely.

02:33 There's no difference between them in terms of stability.

02:39 In reality, they probably all jump back and forth between the three locations, meaning the two electrons that made up the double part of the bond, right? they all have single bonds.

02:49 But the two electrons that we use to create the double.

02:53 Those electrons, instead of being associated with just one location, they're probably jumping between all three locations.

03:03 But this isn't necessarily our endpoint.

03:06 Technically, sulfurs can break their octets.

03:10 And with this molecule, it's actually likely.

03:13 And we're going to bring in formal charge for this.

03:16 Formal charge, as a reminder from earlier problems, we went over extensively in number 57, for example.

03:24 Formal charge equals the number of valence electrons on the atom minus half of its bonding electrons and all of its non -bonding.

03:41 To put this easier, formal charge equals the number of valence electrons an atom has minus how many sticks are attached to it and how many dots.

03:53 So if we look at the sulfur then, and our first three, they'll have the same situation for all three sulfur.

04:02 Sulfur starts with six valence electrons based on the periodic table.

04:07 It has four sticks attached to it.

04:10 If you want to think about that in terms of half of the bonding electrons, the bonding electrons are the ones tied up in bonds.

04:18 So if there's two electrons in each of those bonds, four in the double, then half of that would be eight divided by two.

04:26 But for simplicity, just saying how many sticks are on the sulfur works too.

04:31 So that would be 6 minus 4 minus 0 gives you a formal charge of plus 2.

04:37 Plus 2 is not great.

04:40 Zeros are much better.

04:43 On top of that, some of the oxygens also are charged.

04:49 The oxygens that have single bonds have a formal charge of 6 minus 1 stick and 6.

04:59 So not only is the sulfur plus two, at least two of the oxygens are minus one.

05:09 That last oxygen is actually a zero.

05:11 He's six valence minus two sticks and four dots.

05:15 Okay.

05:16 So these aren't super stable structures.

05:20 In terms of having all the valences complete and things, they're fine.

05:23 But if you want the best format or if we're trying to find the rest of the resonant structures, this really isn't great.

05:30 So let's try making one where we have two double bonds.

05:34 Now remember, this is possible because of how low in the periodic table sulfur is.

05:38 If you're at least in row three or lower, you have access to the d orbitals, which means while you'll be happy to have eight electrons, if somebody offers you more, those atoms can just put them up into the d block.

05:59 They have empty orbitals where those electrons can reside, that it's not hard to reach.

06:06 So sulfur can break an octet.

06:09 It can go higher.

06:12 Okay.

06:13 So we're going to try making this now where we have two double bonds.

06:16 So that means that we would take one of the lone pairs off of one of the oxygens and now create another one.

06:26 So we could have, keep that original one.

06:30 And now let's say we make this one a lone pair as well.

06:40 Now our sulfur, let's double check his.

06:45 That would be six for sulfur minus one, two, three, four, five sticks and no dots.

06:51 He's now a plus one, which is better than plus two.

06:55 Two of those oxygens are zero and the top one's a minus one.