Write Lewis structures that obey the octet rule for each of the following molecules and ions. (I

Write Lewis structures that obey the octet rule for each of...

Key Concepts

Lewis Structures

Lewis structures are diagrams that represent the arrangement of electrons among atoms in a molecule or ion. They are used to show how atoms are bonded together by covalent bonds and to depict any lone pairs of electrons. These drawings provide a simplified visualization of the electron distribution, which is foundational for predicting molecular geometry and reactivity.

Octet Rule

The octet rule is a guideline in chemistry which states that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell, achieving a noble gas configuration. This rule is particularly useful for constructing Lewis structures of many main group elements, ensuring that each atom reaches a stable state through shared or lone pairs.

Valence Electrons

Valence electrons are the electrons present in the outermost shell of an atom and are crucial for chemical bonding. In creating Lewis structures, the total number of valence electrons available—accounting for any extra electrons due to negative charges or the loss of electrons from positive charges—determines how electrons can be distributed to satisfy the octet rule.

Formal Charge Calculation

Calculating formal charges helps in determining the most appropriate Lewis structure among possible alternatives. It involves comparing the number of electrons assigned to each atom in the Lewis structure with the number of electrons the atom would have in its neutral, isolated state. Structures with formal charges closest to zero are generally more stable and are preferred when multiple valid Lewis structures exist.

Structural Variability with the Same Electron Count

Even when species have the same number of atoms and valence electrons, they can exhibit different Lewis structures due to variations in the arrangement of bonds and lone pairs. This concept highlights that identical overall electron counts can lead to isomers or alternative bonding scenarios, each with distinct formal charge distributions, which in turn can influence the physical and chemical properties of the molecules or ions.

Transcript

00:02 All right, so today we're going to be looking at lewis structures again, more specifically looking at lewis structures, and then looking at molecules that have some form of formal charge to them as well.

00:16 And so to get started, we'll just do a little quick run on a structure, on a molecule, sorry, just to practice.

00:26 Just to practice these lewis structures, drawing them.

00:29 This one's going to be pretty simple.

00:31 With these that we're going to draw always the first molecule that, or the first atom that's found on the molecule will be the central molecule.

00:39 So we have our phosphorus there at the center.

00:42 The first step, just a reminder when we're doing this is we want to count how many valence electrons are found in this molecule.

00:49 So we have phosphorus that's going to contain five valence electrons looking at the periodic table.

00:56 Oxygen will have six and then chlorine will have seven.

01:01 And we're going to have three of those.

01:04 So we have 21 found right there, plus 6, going to make that 27, plus 5 is going to make that a32, valence electrons that we need to put here somewhere.

01:15 Okay, so i'm just going to start drawing these out as such.

01:23 So we've got our chlorines and our oxygen right there with our phosphorus found in the center.

01:29 And again, remember, we want to have these follow the octet rule, meaning that each one of these will have eight valence electrons in its outer shell when it's all finished, said, and done.

01:44 So i've got right here eight electrons that i need to subtract, leaving us with 24.

01:50 And so let's just draw out these 24 electrons on here right now.

01:59 Those lines are two.

02:03 Okay, so there we go.

02:05 We've got our 24 valence electrons.

02:08 We have six right there, so 6, 12, 18, 24.

02:11 And let's just see make sure it follows the octet rule so we have six seven eight and all these will then follow that so right there we've got our eight valence electrons and then our phosphorus has one two three four five six seven eight and so right there is the structure for p ocel three properly drawn loose structure okay all right now let's look at uh now that we've kind of reviewed that let's go and look at another one we'll just do it down here i guess, looking at one that now has a formal charge on it.

02:49 Let's look at so4 with a 2 minus charge.

02:54 And so when we're doing this, we need to incorporate that negative 2 formal charge on this overall structure.

03:02 And to do that, it kind of makes it a little fun.

03:04 So we have, again, let's just go back.

03:06 Let's follow the same procedure.

03:08 How many valence electrons are there? so we have our sulfur that contains six valence electrons.

03:12 And our oxygens that do as well.

03:18 Now what's fun is we have this negative two formal charge.

03:21 We actually add two electrons onto this total thing.

03:25 And so we have six times, if we add that altogether, it should be 32 electrons, valence electrons, that we're going to have to incorporate into our structure.

03:40 So i'm going to start out.

03:42 Got a sulfur in the center with our four oxygens outside.

03:45 Outside.

03:46 And that makes eight.

03:47 So we're going to subtract the eight from that, making that, well, not 30.

03:52 That's going to be 24.

03:54 And so we're seeing kind of the similar trend here.

03:56 We're just going to draw our valence electrons on the outside, maybe, like so.

04:05 And so we've got our valence electrons there.

04:08 Now, let's see and make sure it follows the octet rule.

04:12 Okay.

04:12 So we go in, one, two, three, four, five, six, seven, eight.

04:15 So we're good there.

04:16 We're good there.

04:17 And we're good there.

04:18 And so we're good there.

04:18 And so we're sulfur also follows the octet rule.

04:21 So you guys are thinking, oh, we're good to go, right? no, we're not, because we've got this formal charge.

04:29 And how you kind of calculate the formal charge is as such.

04:33 Okay, so we want to look at the formal charge of all of these atoms.

04:38 So for example, let's look at this oxygen right here.

04:41 Okay.

04:41 Now this oxygen right there, it has, we know, six valence electrons.

04:51 Normally, you know, because we added that up right there.

04:55 Same thing right here.

04:56 Six valence electrons.

04:58 We know that, and we have that drawn right here.

05:01 We have one, two, three, four, five, six.

05:03 So we have our six valence electrons.

05:05 And we're going to subtract that from the non -bonded valence electrons that it also has.

05:12 Okay.

05:12 And so we know from right here, for example, that it has, let's see, the non -bonded, sorry, so the non -bonded valence electrons are found right here.

05:25 So we've got our six.

05:26 So we're to subtract six from that, leaving that with a zero formal charge if we looked at that one specifically.

05:36 However, if you notice, we've actually got these bonded electrons as well.

05:41 And so we're not quite done here.

05:43 And so i'm just going to take this away.

05:46 I kind of jump the gun.

05:47 I'm going to actually redraw this because it's kind of hard to see what i'm, doing okay so let me draw this a little bit better so you guys can follow so we're going to be looking let's specifically look at maybe this oxygen right there and so we're going to look and see how many valence electrons does it have so we know it has six valence electrons we're going to subtract that from non -bonded we know right here it's got six so we're going to there's six right there and then the next thing we want to subtract is bonded electrons but we only take half of that and so for example so we take one half bonded.

06:41 So right here, it's got two bonded electrons, but we're going to subtract just one.

06:47 So we start with six minus six, puts it at zero minus the one, puts us out a negative one.

06:53 So realistically, looking at these, all of these oxygens right now have an overall formal charge of negative one.

07:02 Okay? let's look at sulfur.

07:05 Let's do it in blue.

07:07 So sulfur in blue.

07:10 So sulfur also has six valence electrons to start out with.

07:15 We showed that right there.

07:17 How many non -bonded electrons does it have? hopefully you can see that there are none there.

07:24 But we do have some bonded electrons, but we can only take half.

07:27 So we're going to take one, two, three, four.

07:30 So six minus four is two.

07:34 And so this actually has a positive two formal charge right there in that.

07:40 Now, there's a couple things that we can do here to fix this.

07:44 So right now, the 2, the plus 2, would go in and maybe get rid of some of those, leaving us with a formal charge of negative 2, right? and so we're good to go.

07:54 So this is one structure that we could take.

07:58 And sometimes what you'll see actually happen, though, is we can see these bonded electrons maybe come in and form double bonds there to make this a 0.

08:09 And so let me just redraw what this kind of would look like.

08:15 There we go.

08:17 And let's see what that would cause.

08:20 We have our sulfur bonded to an oxygen.

08:23 Bonded to an oxygen.

08:26 There we go.

08:28 And the reason why i would do this is because we don't, all these have a formal charge on them.

08:34 And it doesn't really like that.

08:36 So it's going to try and minimize and make some as many as possible have a zero formal charge to make them kind of just as happy as possible.

08:47 They want to be in that happy state of that when we're looking at these charges.

08:53 And so let me draw our valence electrons on here.

09:05 And let's count the formal charge that we have on these different molecules now.

09:11 So these ones we didn't change any.

09:13 So these will still have a negative one formal charge.

09:16 But let's look at these now that are double bonded.

09:19 So our oxygens that are double.

09:21 Double bonded.

09:22 They still have six valence electrons.

09:24 Non -bonded, they only have four now.

09:27 And then we have two, one, two.

09:31 And so when we calculate that for six minus four, minus two, is going to give us a zero formal charge on those.

09:39 Now sulfur right here, let's see what it's got.

09:43 So we start with six.

09:44 We have one, two, three, four, five, six.

09:48 Oops, not right there, though.

09:50 So zero and then six, making that have a zero formal charge as well.

09:56 So when we look at this structure, this is the most simple structure, and it's probably the one it would form, still has a negative two formal charge, but now we have three molecules that don't have a charge on them.

10:08 They're pretty happy...

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