Use formal charges to predict which of the resonance forms is most plausible for (a) SO3 and (b)

step 1 All right, problem number 63. This goes right along with problem number 61 in the book that we d

Use formal charges to predict which of the resonance forms...

Key Concepts

Formal Charge Calculation

Formal charge is determined by assigning electrons in a molecule’s bonds to the atoms based on an equal split, then comparing that number to the atom’s valence electrons. It’s calculated using the formula: formal charge = (valence electrons) – (nonbonding electrons) – (bonding electrons/2). This concept is essential because minimizing formal charges generally leads to more stable and realistic resonance structures.

Resonance Structures

Resonance structures are different Lewis structures for a molecule that cannot be represented by a single conventional Lewis formula. They illustrate the delocalization of electrons across a molecule, indicating that the actual electronic structure is a hybrid of these forms. Choosing the most plausible resonance form often involves comparing formal charges between the different possible structures.

Electronegativity Considerations

Electronegativity is a measure of how strongly an atom attracts electrons in a bond. In the context of resonance structures, atoms that are more electronegative are better able to accommodate negative formal charges, while less electronegative atoms are more appropriate for carrying positive formal charges. These considerations help in evaluating the stability of different resonance forms.

Resonance Stabilization

Resonance stabilization refers to the lowering of the overall energy of a molecule due to the delocalization of electrons over several atoms, which is depicted by resonance forms. More stable resonance structures typically have minimized formal charges, full octets where possible, and appropriate placement of charges in relation to electronegativity. This stabilization is key in predicting the most plausible structure among the available resonance forms.

Transcript

00:01 All right, problem number 63.

00:03 This goes right along with problem number 61 in the book that we did just a moment ago.

00:09 So if you need more in -depth background, check out problem number 61 from the same chapter.

00:15 It'll go a bit more in depth.

00:17 Which resonant structure is the most plausible? and first we're looking at so3.

00:21 All right.

00:23 So first off, we're going to draw our lewis structure.

00:26 Lewis structure is based on how many valence electrons are in our compound.

00:30 We can use the periodic table for a quick guide on how many valence electrons something has.

00:37 For sulfur, we've got six valence plus oxygen, which also has six, and there's three of those.

00:48 So we're working with a total of 24.

00:51 We're going to put sulfur in the middle, fill it with, surrounded by oxygens.

00:56 It doesn't matter if it's kind of slanted like that, or if you, you know, straight corners.

01:01 Doesn't matter.

01:03 Three oxygens attached to that sulfur.

01:07 Two bonds or two electrons are in each bond.

01:10 So so far we have used up six electrons, which means we've got 18 more to go.

01:18 Each sulfur, or sorry, each oxygen currently has two electrons.

01:23 It needs to have at least eight.

01:24 So we're going to put six more on each of those.

01:30 Six times three is 18.

01:33 We are out of electrons, but that sulfur is not satisfied.

01:39 That sulfur only has six, so we need to use a double bond.

01:43 Choose one of the oxygens and go ahead and move it.

01:47 One of its lone pairs to create that double bond.

01:54 All right.

01:55 And i'll draw the electrons that survive.

01:57 Okay.

01:59 Now, without looking at formal charge, just based off of the octet rule, this would be our final structure.

02:07 All right.

02:08 Sulfur has eight, oxygens have eight, everybody's happy.

02:11 But in fact, if we look at the formal charge of this compound, this is actually not a great structure.

02:17 So really quick, formal charge of an atom equals how many valence electrons it starts with, minus half of its bonding electrons, which are the electrons in bonds that are attached to it, minus the non -bonding.

02:34 I find that to be a mouthful and hard to remember, especially when teachers tend not to give this equation on exams.

02:41 And so i remember it as the formal charge is the valence electrons minus how many sticks are attached to an atom, meaning like bonds, how many sticks or how many lines, minus how many dots are on that atom, which are your individual electrons.

03:00 All righty, so if we look at that sulfur to start with, sulfur, as we know, starts with six valence electrons.

03:07 So its formal charge would be six minus.

03:10 There are a total of four, sticks, attached sticks, lines, whatever you want to say.

03:17 Attached to that.

03:18 One, two, three, four.

03:21 And there's no dots on it.

03:22 So it has a formal charge of plus two.

03:25 We really want that to be a zero.

03:28 Plus two is not super stable.

03:31 Okay.

03:31 If something isn't very stable, it means that it's going to find a different configuration that is more stable and it's going to spend more time there.

03:39 Right? these oxygens, the ones that have a single bond, would have a formal charge of six for oxygen, minus one stick, and six dots for a formal charge of minus one.

03:54 So two of those are both minus one.

03:58 And then the oxygen that had the double bond, that would be six valence minus two sticks, four dots for zero.

04:10 So the only really great formal charge and that whole thing is the oxygen that was double bonded.

04:16 So we're going to take advantage of the fact that sulfur can have more than eight electrons.

04:22 This is true for any of the atoms that are in at least the third row of the periodic table or lower.

04:29 They still want to get to eight electrons.

04:31 They still have that octet rule.

04:33 But if it means forming a more stable molecule, they will have up to 12 electrons.

04:40 Okay.

04:41 So helium and hydrogen want to have two.

04:45 Lithium, beryllium, boron, carbon.

04:48 Those are all aiming at eight.

04:50 And then once you get to the third row, it's eight all the way up to possibly, it's really like eight to 12.

04:57 They can have up to 12.

04:59 I just don't have room to write right there.

05:00 I didn't leave myself enough room.

05:02 So eight, ah, eight plus.

05:07 All right.

05:08 So what we're going to do is try adding another double bond to this and see how it does to the formal charges if it's better.

05:14 So instead of just one oxygen, having a double bond.

05:18 Let's try putting it on two.

05:20 That would mean we still have that first one.

05:24 Now we do another one, and then one of them is left over like this.

05:31 Okay.

05:32 Now, based on our previous understanding then, the oxygens with double bonds had a formal charge of zero, and then the oxygens that had a single had a formal charge of minus one.

05:44 So what's left is to look at that sulfur.

05:47 Sulphers have six valence electrons to start.

05:51 Minus one, two, three, four, five sticks and no dots.

05:58 We're down to a plus one, which is better than a plus two.

06:03 But you may be able to guess at this point, if we do one more double bond, everybody's going to have a formal charge of zero, which is the most ideal situation.

06:14 So let's try that instead.

06:16 Remember, each time we're creating a double bond, the lone pair that's creating that is coming from one of the oxygens.

06:23 So that's why these oxygens are going from having six valence electrons, like six dots around them.

06:28 That's why they're dropping to four, because the two extra electrons went into the double.

06:36 All right, our oxygens are six minus two sticks and four dots.

06:42 So all the oxygens are zero.

06:46 And that sulfur is six minus one, two, three, four, five, six sticks.

06:53 And no dots for a formal charge of zero.

06:58 That is your ideal situation, which means that is your most stable form.

07:02 The molecule is capable of jumping between all the ones that we have drawn.

07:06 It can form all of these.

07:09 But if you are able to take a picture of a molecule over time, the vast majority of the time it's going to be in this last one here because it's more stable.

07:21 Okay, so that's looking at that one.

07:23 Let's look at the next one.

07:25 I already had a periodic table.

07:27 We'll bring it down a little bit more.

07:29 So we can jump into this next one.

07:31 Okay.

07:32 Now we're looking at hno3.

07:40 All right.

07:41 Now the way that these are generally formed is the nitrogen is going to be in the middle...