Lecture 07. Periodic Trends Continued
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Transcript: English(auto-generated)
00:07
So, last time we had started talking about our different periodic trends and we had started with effective nuclear charge because that's one of the main reasons behind all of our other trends. So, today we're going to continue with that and we're going to start up where
00:20
we left off doing examples of ionic ones. So we had talked about a few of them, so it's time to do a few more so you can see exactly the sorts of things that you're expected to know. So here's a little periodic table for you to be able to use as we go through. As always, when you're listening to these lectures, you should have the periodic table in front of you so that you can follow along even when I don't have it up on the screen.
00:42
So let's do these two first, or these three first. So we have lithium plus, beryllium 2 plus, and fluorine minus. So you can find all of these on the periodic table, you can say, well there's lithium 1 plus, so that has the same sort of electron configuration as helium. Beryllium 2 plus, which would also have the same electron configuration as helium.
01:02
And then F minus. So when we go through and we put these all in order, we would want to put our 2 plus charge first because that's going to be our smallest, and then go through and put these in our 2 plus first because it's the smallest, then our lithium 1 plus because now we have
01:21
a little bit less of an effective nuclear charge. And then our F minus, because with F minus now we have that whole extra row of, that the fact that it's a minus charge, so the effective nuclear charge is lower. And so that gives the electrons a little bit more room to spread out, the combination of which makes fluorine definitely the largest.
01:43
So these two are isoelectronic, so we're just looking at the charges. Fluorine, now that has a whole other row of electrons, and it has a lower effective nuclear charge with the minus charge. Okay, next ones. So we have copper 1 plus, copper 2 plus, and potassium.
02:04
So we go to our periodic table and we find them, and so there's copper. So copper 1 plus and copper 2 plus would just have their S electrons removed. Remember, always removing from the S electrons first. So then we look at where potassium is, same column, or same row, but all the way over
02:21
to the left. So that's going to be isoelectronic with argon. So when we look at these and we put them in order, the first thing to look at is our charges, which means that, or which one's going to be our smallest. Well the Cu2 plus, that's definitely going to be our smallest. So then we go through and we say, well what about copper and potassium?
02:44
Those have the same charge, and they're in the same row of the periodic table. So which one's going to be smaller there? At that point we backtrack to what we did last class, and we look at where they are in the periodic table, and we do it just like we would do atomic radius. So what happens as we go across the table this way to our radius?
03:05
Well to know that, we have to think about what happens to our effective nuclear charge. So as we go across the periodic table that way, our effective nuclear charge becomes greater. So our protons in our nucleus are pulling in those electrons harder, and it's making it smaller.
03:20
So something on this side of the periodic table is going to be smaller than something on this side of the periodic table. So between Cu1 plus and potassium, the Cu1 plus is going to be smaller, because it's further to the right. And so we have this ordering. This is the smallest because it's both furthest to the right, and the highest charge.
03:42
And then this one, because it's further to the right than potassium. So even though they're the same charge, at that point you backtrack to doing it how we did atomic nuclei. Next one. So now these are all ions, they're all our halogen ions. They're all a minus one. So there's nothing here really to look at as far as charges go.
04:01
So the only thing we need to look at is our trend as we go down the periodic table. So as we go down the periodic table, we're adding shells of electrons, and each time we add a shell of electrons, it's going to get larger. And so because of that, our smallest is going to be up here, and our largest is going to be down here. So that makes fluorine our smallest, and then bromine our largest.
04:24
And so our ordering is like that. So this is probably the trickiest one of the group, but keep that in mind. It's really no different than this. Just like here we can say, well, we're going down the periodic table. They're all the same charge, and so all we're looking at is the trend.
04:41
We're doing the same thing here. We just looked at the trend going across the periodic table. Okay. Let's do one more where we kind of do some matching. So I'm going to give you a picture with these four ions, and I want you to match up their pictures. So we have sodium, magnesium, chloride, and oxygen in their ionic forms.
05:03
So let's think about how we could do this. So again, you're looking at your periodic table, and you're thinking, what are the size difference here? So I have them as relative size differences to each other. We know in this case that I have a one-to-one ratio. So if we look and we say which one's going to be our smallest of these ions, and which
05:23
one's going to be our largest of these ions, we can say, well, magnesium's definitely going to be our tiniest, right, because it's only a plus two, or it's a plus two, where oxygen's going to be our largest because it's a minus two. So there's not nearly as many protons per electron. So that would make this our magnesium oxide, and this our sodium chloride, since the
05:43
sodium and the chloride are sort of in between. And so we can go through and we can label and say, well, this is our biggest of the grouping. And not only that, but we know that magnesium oxide would form a one-to-one ratio, and that sodium chloride would form a one-to-one ratio. So that's one where we can kind of go through and match up.
06:02
Okay, so that's sort of it for our atomic radii and our ionic radii. Now we're going to move on to first ionization energy. So what ionization energy is, well, let's think about the wording here, ionization, so the amount of energy it takes to ionize something, to pull off an electron, to take
06:22
an electron from an atom and pull it away. So we have some trends here that we already sort of generally showed you earlier, but as you go across the periodic table this way, the ionization energy is going to increase. And as you go up the periodic table, it increases.
06:41
Now this one has a ton of exceptions that we're going to talk about in a minute, but first of all, let's explain why this trend happens, because it's never enough to just memorize the trends. I want you to know why they happen, and I will ask you that. So as we go across the periodic table, there's something else that increases. So think about what we talked about when we talked about atomic radii and why it increased.
07:02
And then the trend before that that we talked about being effective nuclear charge. So that effective nuclear charge has to, is the reasoning behind a lot if not most of the other trends. And this one's no exception to that. As you go across the periodic table, your effective nuclear charge increases, right?
07:21
It was like adding more strength to that inside nuclear magnet and the outside electrons, but not adding another shell. So we weren't moving the electrons further away, they were the same distance away. So that effective nuclear charge increased. So if effective nuclear charge is increasing from this side to this side, well, now you
07:42
have to look at it and you have to say, okay, well, now you're saying my electrons are being held on tighter to my nuclei. So if I want to remove one of those electrons, I have to pull it away with more strength. If you think about putting two magnets together, which is harder to yank the magnets apart with? Something that's a very weak magnet or something that's a very strong magnet?
08:01
Of course, something with a very strong magnet. So this is the same sort of idea. You're trying to pull an electron away from something with a very high effective nuclear charge if you're over here. Over here you have a lower effective nuclear charge, so in comparison it's going to be relatively easy, so your energy will be relatively low. Now with that being said, there's a whole bunch of exceptions.
08:24
But they follow the exceptions, follow their own little trends. So we can look at these in a minute and see why they are. And they're all going to be from half and fully filled shells. So this comes from an older book from Chang, but it's also highly edited. So I've added in some things, so I want you to make sure you copy this down.
08:45
Now with this, we have each of the elements as we go across the periodic table. And I want you to explain why there's these exceptions. Because as you start on your second row and you work your way across, you'd see that we have helium and then lithium.
09:02
Now you would be expecting it to just go beryllium, boron, carbon, nitrogen, oxygen, fluorine, neon, all the way up straight. Yet, you have an exception here and an exception here. So to do this, oh, and then of course the same thing over here. So to do this, we have to go back a little bit and talk about electron configurations again.
09:20
So now that you're good at electron configurations, write out the electron configurations for these. So for beryllium, you'll go to helium and you'll say, okay, well, now the two s's are filled. And you'll have two s's too. Now if you write down boron, you'll see you're one further on the periodic table.
09:41
So it's going to be two s two, two p one. So what happens here and here is that you reach this sort of level of stability here that you don't have here. And so boron has this one little electron sitting in its p orbital. So that's a lot easier to remove
10:01
than trying to remove this. So this one's going to be lower than beryllium because with beryllium you're removing from a fully filled shell. And there's a relative stability here. With boron, you just have two p one. There's no sort of added stability that comes from that one little electron being there. And so it doesn't take much energy to remove it.
10:24
So now, let's see if you can do nitrogen and oxygen. So you're looking at your periodic table, you're finding nitrogen on it, and then you write down the electron configuration. So it's going to start off the same as these. We're going to start with helium, and then two s two,
10:41
and then two p three. Now, this is something we haven't completely talked about yet. But when you have a half filled subshell, there's a level of stability there too. It's not quite as good as having a fully filled shell, but there's still extra stability. And so what happens there is that then when you move to the oxygen,
11:01
and you add this fourth one, you get a little bit of a repulsion from that. When you add in that one extra electron, there's sort of this combination, extra stable half filled shell, and then added a little bit of repulsion from adding in that next electron and having it be the opposite spin. So that combination of the two issues means that oxygen is going to be a little more willing
11:23
to lose that one extra electron and become the same sort of electron configuration as nitrogen. So this one, it might be good to also draw out an electron configuration diagram on your own and see how this works. That this, you would have your p orbitals and you would draw your one electron per orbital.
11:41
You'd see that it's half filled. So adding this one extra electron destabilizes it a little bit and you end up with this little exception. Now, on your own at home, what you're gonna do is I want you to write out the electron configurations for magnesium and aluminum and phosphorus and sulfur and prove to yourself that it's the exact same issue.
12:02
So when you do that, what you'll see is that basically just the numbers, the n value will change. And so go home and do that and prove to yourself why this is like it is so that you could explain it to me if it showed up on an exam or something. So let's do a bunch of examples now with this. And remember, you're always keeping this in mind
12:22
that when you get to this half filled shell, you're gonna have, you might have some exceptions that you need to look at. And when you have here, you're gonna have some exceptions you have to look at. Now, as we're out here, you can look at the exceptions, pay a little bit of attention to them, see where they are, they're interesting, but I'm not gonna test you on all of them.
12:41
That goes into a little bit more inorganic chemistry that we're not gonna get into in this class. But these, we can definitely explain just using electron configurations. So that's something that you should be able to explain. Okay, so same idea as before. Let's rank these.
13:01
So let's first do some ones that don't have exceptions. Now, of course, when you're given this and you sort of examination, you wouldn't be told whether they're normal or not. But in this case, we're gonna start that way. So you have helium, neon, and argon. So you're gonna go to your periodic table and say, okay, well, these are definitely noble gases. So they're gonna be over here.
13:21
Helium, neon, and argon. And rank them in order. So which one's gonna be the easiest to pull off? And for that matter, why? We didn't necessarily talk about the trend going down. So going down, we know that it's easier to pull off from things lower on the periodic table. Why is that? Well, that's because they're further away, right?
13:41
So trying to, if you have two magnets like this and you try to yank them apart, it's gonna be a lot easier to pull them apart than if you have two magnets that are nearly touching and pulling them apart. Same idea here. The electrons are further away and they have all of those inside electrons. So remember, there was a phrasing for that too. You have all the inside electrons that are blocking the outside electron
14:02
from feeling the nuclei. So the buzzword that went along with that one was shielding, right? The inside electrons are shielding the outside electrons. And so they're able to be pulled off easier. They're further away and they're shielded. So with that in mind, helium is gonna be our smallest, and then neon, and then argon.
14:21
So, sorry, I said that backwards. So the furthest down one, I did decreasing. So we want lowest to highest. So argon is gonna be our lowest, and then neon, and then helium. So we start with the lowest one down here because the electron and the nuclei are furthest apart, and it has the most shielding.
14:40
So it's gonna be able to be pulled off very easily. It's not gonna take very much energy to pull off that electron. Something like helium though, now there's no shielding, right? There's no electrons between the helium's electrons and the nuclei. So that's gonna be very hard to pull off. There's no shielding and they're very close to each other.
15:01
Okay, next one, boron, lithium, and beryllium. So now, if we look at this, we have lithium, boron going this direction, and then neon. So they're all in the same row. And they're all going just straight across.
15:20
So our trend going that way has to do with effective nuclear charge. And which one has the smallest effective nuclear charge? Well, lithium does, right? Because it's over on this side. And so if it has a very small effective nuclear charge, then it's not gonna take as much energy to pull off the electrons. And so the ordering of it will have to do with that.
15:42
It doesn't take much energy to pull off the electrons. The ionization energy is very low. So because lithium has a low effective nuclear charge, it also has a low ionization energy. And then so on across the periodic table. Okay, so those were our normal ones.
16:00
So those are ones that followed the rules, right? There is no weird things going on with their electron configurations or anything like that. That's not always gonna be true, especially when you start looking at boron. When you see a boron, beryllium, or nitrogen, oxygen, and anything below it, that's when you wanna start looking at their charge, or their electron configurations.
16:21
So now let's do lithium, beryllium, and boron. Lithium, beryllium, and boron. So let's say there wasn't any exception going on, and it just went straight across the periodic table. We would know to rank this as our smallest, and then beryllium, and then boron. Now we have to think back to the last slide, though, and we have to say, okay,
16:40
how are the electron configurations factoring in here, though? Well, beryllium here is a fully-filled shell. It has this relative stability. Boron has this one extra electron in this p orbital that isn't in any sort of half-filled or fully-filled state. So with that electron being out on its own, it doesn't necessarily mind losing it.
17:01
It doesn't take much energy to lose it. Where beryllium is saying, well, I'm kind of at this level of stability. I don't wanna lose that electron. So instead of going lithium, beryllium, and boron, the beryllium and the boron flip-flop. So you get lithium, boron, beryllium.
17:21
And then our next one, C, N, and O. So right here, across the periodic table. And again with this one, if we were looking at this and we said, okay, this is gonna follow the trend, well, we would do it in exactly this order. We would go C, N, and O, because C has the lowest effective nuclear charge, oxygen has the highest,
17:42
so oxygen would be the hardest to remove. But it's not only based on the trends, it's also based on the electron configurations. And if we look at the electron configuration for nitrogen, it has a half-filled shell. Oxygen, you add in one extra electron. So you're not making it fully filled. You're just adding, or you're just taking away the one.
18:03
So in this case, oxygen doesn't really want that extra electron necessarily. It adds a little bit of a repulsion. It doesn't add any stability. So oxygen's okay with losing it. It's not gonna take a lot of energy to remove that one. Where nitrogen has a half-filled shell,
18:21
so it says I don't really wanna lose another one. I don't wanna lose an electron and be stuck with two in my p-orbital. And so instead of going with the trend, the oxygen and the nitrogen flip-flop. It doesn't take as much energy to remove the electron from the oxygen as it does the nitrogen. Okay, so those are our examples,
18:41
some following the trend, some not. And you should be able to recognize any of these regardless of whether I tell you if they're exceptions or not. That was just for an intro. Okay, next thing then that we're gonna talk about. Second and third ionization energies. Now this is something that I could expect you
19:00
to kind of rank within an atom, or have you look at the electron configurations to figure it out. So what a second and third ionization energy is is after you pull off the first electron, well how much energy does it take to pull off the second one? How much energy does it take to pull off the third one? That's what you're going through as you do second, third, and so on
19:21
with ionization energies. So your first one is always gonna be your smallest, which sort of makes sense, right? If you take one off, now you have more protons than you do electrons, right? You have a cation, you have a positively charged ion. And so of course it's gonna take more energy to pull off the second one
19:40
because now you have more positive charge per negative charge. So each one that you remove, the energy's gonna go up a little bit. Now, you'll have some homework problems on trying to rank these, and you end up having to have a lot of data for them, so that's why we'll leave that for the homework. But what happens is, is that as you go through,
20:03
you'll be able to see big jumps. So let's go back here so I have a periodic table to explain with. So for instance with beryllium, you'll see a certain ionization energy. And then that means it would be down to a 2s1. So you're pulling off that lithium,
20:21
it's gonna, that second 1s electron, now it's gonna be a little bit bigger, but it's still gonna be relatively small. At which point you'd be isoelectronic or you'd have the same electron configuration as helium. And when you go to pull off that one, it's gonna be a little bit larger than the first and the second one, but it's gonna be a huge jump. So you're gonna see a little bit of a difference,
20:42
a little bit of a difference, a huge jump. And that's how you know how these work, is you look at the differences between them. First there's always the smallest and then they get bigger and bigger and bigger, but you can tell what sort of electron configuration you're dealing with based on the differences between them. And you'll have some examples of that in your homework.
21:00
Okay, so that sort of wraps up ionization energy. Now electron affinity, and then we'll also talk about electronegativity even though it doesn't technically fall in this chapter. The two are very similar and so I wanna talk about them together. So electron affinity is sort of the opposite of ionization energy. With ionization energy, we had an atom
21:21
and we took away an electron. With electron affinity, we're gonna be doing the opposite. We wanna know how likely is that atom to take on an electron. So to one atom by itself to bring on an electron. Now, so for instance, something like chloride taking on an electron to become chloride minus.
21:40
So with ionization energy, we were making cations. We were taking away electrons. With electron affinities, we're adding electrons, so we're making anions, we're making negative electrons or negative ions. Now, with these, the trend sort of follows
22:00
a lot of the other ones we've been looking at where it increases as you go up the periodic table and increases as you go to the right on the periodic table. Now, I wanted to just show you this and I'll show you this in a few different ways for electron affinity, but this website is really great for looking at periodic trends, especially if you're a visual person.
22:21
You can go in and you can graph any of the trends that we've been talking about in the periodic table format. You can do it in a whole bunch of different ways. If you like numbers, it'll fill in the numbers for you. If you like seeing it in the sort of like cityscape idea, it'll do that. You can do it squares, circles, so that you can see it in a lot of different ways and you can kind of see where the trends fall.
22:41
And it's, I think, helpful for a lot of people to see it that way. So, you know, take a look. I picked two different ways of showing you electron affinity in it, but there's tons of different ways to display the data. Okay, so now let's talk about why these trends are. So at this point, you want to still be thinking about all your other trends and how we explain those because they're pretty similar explanations.
23:01
So as we go up on the periodic table, we're increasing, or I guess we're getting smaller, right? So if you think about going down the periodic table, we'd be getting bigger. And so our electrons are already further away from our nucleus. Our outside electrons aren't able to feel
23:21
the nuclei very well because of one distance, they're just further away, and two, shielding, right? The fact that those electrons on the inside are blocking the electrons on the outside from feeling the nuclei. So as we go down the periodic table, the electron affinity gets smaller because it's going to be further away
23:41
and there's more shielding. As we go across the periodic table this way, now what is increasing? Well, our effective nuclear charge is increasing, right? So as we go across the periodic table, our effective nuclear charge increases. So our electrons are able to feel our nuclei a little bit better.
24:00
And so a new electron will be able to feel the nuclei a little bit better too. So as we go across the periodic table, it gets larger. Now you'll notice there's a lot of exceptions to this, just like there was with ionization energy. And if we had graphed ionization energy like this, it would actually look kind of similar.
24:20
So go play around with the website and graph it and see. So you'll notice where your exceptions are here. Well, there's this exception here with this 2s2, or excuse me, s2, depending on what n you're in. And the s2p1, again, depending on what n level you're in. And then there's this other exception right here, right?
24:42
Which is your nitrogen row, where this is that same idea as when we talked about ionization energy. So there's that set of exceptions. And we'll do these in more detail later, but this is just so you have a picture going into it. And now, same idea as every other trend.
25:02
This trend works well until you get to this halogen. So if you count over, you'll see that this is the seventh column, right? This is your halogens. So what's this hidden block here that you can't see? Well, that's all of your noble gases. And they're really low. They're really low for electron affinity, ionization energy, all of that. And that's just because they're already really stable.
25:22
They don't want another electron. They don't want to lose an electron. And so they're going to be very low and they're going to break all of those rules. So the rules work really well, except for when you have these exceptions because of electron configurations, right here, right here, and then of course, your noble gases. Okay.
25:42
Now, I want to show you some numbers too, in case you like numbers, so that you see that they are quite a bit different. So this shows you the exact same thing that I showed you here, except in number form so that you can see it. You'll see that these were actually so low, they really were pretty much zero, maybe even negative, depending. And you'll see the same thing with nitrogen.
26:02
So we have those same sorts of exceptions we talked about with ionization energy. And it's for the exact same reason. It's because of the electron configurations. Here, you're already a little bit stable, not going to want to add another electron to have this electron configuration. With nitrogen, you're already at a half-filled shell. You don't really want to add another electron
26:22
and become like an oxygen. And so because of that, you get these exceptions. And then you'll notice your noble gases all the way down. Your noble gases aren't going to take on another electron. They're not going to become an ion. Okay.
26:41
So, let's go through and actually walk through some of these electron configurations so that we can see how this works. Okay. Why do group 1A and 7A form stable atomic anions? So your 1A is your sodium row and your 7A is your halogen.
27:02
So first of all, let's start with your halogens, mostly because those are the ones that we're used to looking at, being, having an anion. To be honest, we're not super used to looking at this as ions, or as anions, I should say. Okay. So we write down our electron configuration for fluorine. Now quickly write down your electron configuration
27:21
for a fluorine minus. So of course, you'll add an electron here and you'll end up with this. So now, let's think about why this would be. Why is our 7A going to form this anion? Well, if you give it a negative one charge,
27:43
you add an electron, fluorine has a high electron affinity, right? It has a very high electron affinity. It wants to take on that electron. And up until now, we've sort of said, well, it wants to be like a noble gas. So that's true. This definitely gives it a noble gas configuration. And so that's one of the reasons behind
28:01
why it has a high electron affinity. It's on that right-hand side of the periodic table. It has a high effective nuclear charge, thereby having a high electron affinity, and then you fill its shell so it becomes like a noble gas. So all of those reasons are good reasons for it to want to be a negative one ion.
28:20
Now let's look at our next one. Sodium. So if we write down our sodium, we have this, and we write down our sodium minus, we get this. Now, we may not be used to looking at sodium as a minus,
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but it's possible. It's not quite normal. Normally, what do we do? Yeah, we just take away this electron and we make it a plus charge. But we could in theory do that. It's not the worst thing in the world for it. And it gives it a full subshell. So it's not gonna fill its shell like with fluorine,
29:02
but it does give it a little bit of added stability. So if you put it under the exact right circumstances, it could do this, and it would be a little bit more stable than a sodium on its own. Now, of course, it's also more stable to just remove that electron and become a plus ion, but either one is more stable than just sodium.
29:21
Okay, now our noble gases. Why don't they form stable atomic anions? Why is their electron affinity so low? So let's pick one. I picked neon. You could pick any of them though. And write out the electron configuration for it. So we write out the electron configuration, and everything's filled, everything's stable.
29:40
So it's not going to wanna go ahead and take on another electron. It's not gonna wanna be a negative charge because it doesn't add any stability to it. And so its electron affinity is basically zero. Now, let's look at one more, and let's look at nitrogen.
30:01
Why would this one not wanna form an ion? So why would we not have a nitrogen negative one? Well, if we did that, we would end up putting a 2p4 here. And that extra electron causes a little bit of repulsion. So there's this extra repulsion
30:20
that comes from having to pair one electron. Now, if you pair all three of them and you get yourself to a noble gas configuration, well, okay, that has a level of stability to it. But just adding one, that lowers the stability of it. And so that's not going to happen. So that's why nitrogen, if we go back to this page,
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has basically zero. And that's also why, then if we wanna look at the one before it, we can see that this row breaks that trend. It breaks that trend and it goes lower because of that half-filled shell. Okay, now we're gonna move on
31:02
to what technically is chapter two, but I wanna cover in chapter one and you'll be tested on it in chapter one. And that's electronegativity. So this is very similar to electron affinity and people get it confused a lot. So be careful that you realize the difference. This is the ability of an atom to attract electron density in the chemical bond.
31:24
So it's not attracting a whole electron and bringing it to itself. That's electron affinity. This is, it has to be in a bond. Now, the nice thing about it being in a bond is it takes away most of the exceptions. So make sure that you can identify the differences between these. And this is why this is technically in chapter two.
31:42
Technically, we haven't learned about bonding yet, but you wouldn't be in this class if you didn't know what bonding was and we talked about it in the fundamentals. So this is gonna follow the exact same trend as electron affinity and ionization energy and it's gonna do so for the exact same reason. So you don't really have to memorize the reason
32:01
behind these for every single one of these. They all have the same reasoning. As you go across the periodic table this way, what's increasing? Well, your effective nuclear charge is increasing. And if your effective nuclear charge is increasing, it's holding onto those electrons a little bit more.
32:20
And so since it's holding onto the electrons a little bit stronger, it's going to be a little bit less likely to lose an electron. It's gonna be a little bit more likely to an attract an electron. Or in this case, it's gonna be a little bit more likely to pull the electron density toward it if it's in a chemical bond. Now, as we go down the periodic table then,
32:41
it gets smaller. And that's for the exact same reasoning too. As you go down the periodic table, the electron shells that you're adding electron shells, you're making those outside electrons further away from the nuclei, which means that the outside electrons have less of an effective nuclear charge because of distance and because of shielding.
33:01
And so as you go down the periodic table, your ionization energy decreases because it's easier to pull off an electron. Your electron affinity decreases because it's less likely to attract an electron. And your electronegativity decreases because it's already having enough, it's already less likely to hang onto its own electrons.
33:22
It's a lot less likely to attract electrons in a bond. So all of the same reasonings behind the other ones are true for this as well. There are some exceptions to this and those exceptions are in the D block and fall into the realm of more intense inorganic chemistry that you can learn if you take inorganic chemistry.
33:41
So we're not really gonna mess around with exceptions here. So let's look at this so that you can get sort of a visual for it. And you can see that yeah, this looks a lot different than our electron affinity and our ionization energy pictures, right? As you go across, it's almost a pure gradient. And sure, you can see some exceptions here.
34:03
And maybe after our discussion about electron configurations, you may actually be able to figure out some of them. We're really not gonna worry about it too much. Just know that as you go this way, it increases. Now let's talk about why there's not exceptions. Cause there's, you may not need to know the exceptions for this,
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but you should know why there isn't exceptions in electronegativity, but there is an electron affinity and ionization energy. So what's the main difference between the two? Well, let's first just talk about the differences between the two and maybe at that point, you'll know why.
34:41
So they have the same trends, right? We just looked at the trends for both. They're the same trends. They're both talking about the ability of an atom to attract electrons. Now your main difference comes in, in the fact that electronegativity is the ability to do it in a bond. So it's attracting electron density. It's not attracting a electron.
35:02
It's still sharing the electrons, but it's just attracting more of it. It's being unequal sharing. Now electron affinity is the ability of it to do it in an isolated atom. So an atom on its own, not bonded to another atom, what is the likelihood of it attracting a whole electron
35:20
on its own and becoming an anion? As opposed to electronegativity, where you're talking about the ability of it to attract it in a shared bond. So this example would be something like CO2. Something where oxygen, if you find that on your periodic table, is much more electronegative than carbon. So if you were to look at where all the electron density is sitting,
35:41
most of it's being stolen by the oxygen. The oxygen's taking most of the electron density and not sharing it with the carbon. With something like electron affinity, you're gonna be talking more like this, where you have a chlorine atom on its own, not bonded to another atom, taking an electron and becoming a Cl minus.
36:00
Okay, so now with all of that in mind, why would you have exceptions for one and not the other? Well, if you're talking about electron affinity, and for that matter ionization energy too, why is it that we had the exceptions? What is it that made those exceptions happen?
36:20
How did we explain them? We did that with electron configurations, right? We said, well, the electron configurations say that we're at a fully filled subshell, or we're at a half filled subshell. So why wouldn't that matter for something in a bond? They already have an octet, right? Everything in a bond, you formed that bond for a reason.
36:41
With CO2, we formed the CO2 bond so that this would have an octet. It's sharing with the oxygen to form an octet. And the oxygens are sharing with the carbon to form an octet. So there's no exceptions here because the electron configuration is already stable. We've given everything its octet.
37:00
We've made everything stable. So the only thing that plays any role in how much electron density an atom is gonna pull toward itself and steal from the other is its effective nuclear charge. And of course, as you go down the periodic table, how much shielding there is and all of that, which affects the outside shell's effective nuclear charge.
37:20
So that's why electronegativity doesn't have as many exceptions as electron affinity. Okay. So now I wanna go through and explain something that we talked about in the fundamental section. When we talked about it in the fundamental section, I said, here's a little trick to help you memorize and we'll explain it later.
37:42
Now it's later, so let's explain it. The inert pair effect. So we're kind of, I should make mention, we're moving on from periodic trends. This was the last of the periodic trends that we're gonna talk about. So we're sort of done with that. And now we're gonna move on to a couple of different things that we need to talk about
38:01
that come up with the periodic table. So it's still sort of in the realm of it, but just not quite as general. So the inert pair effect. So this was, I pulled this up during our fundamentals talk to say, hey, this will help you memorize how to do these ions because you can look and you can say, well, look, these are off by two. These are off by two, off by two.
38:23
And so it gave you some hints for memorization. Now let's go through and explain why these are how they are. So we have this little miniature section of the periodic table there. So let's, to do this, write out the electron configurations for everything.
38:41
So first we're gonna do two. We're gonna do antimony and lead, so that we'll do two examples. And then I'll leave these other four for you to do at home for practice. So we'll start with SB. So we start here. So you should have your big periodic table in front of you so you can see where this little section is.
39:02
And let's write out the electron configuration for just the neutral compound. So if we do that, we're left with the noble gas, five S two, four D 10, five P three. So if we go to remove our electrons, where are we gonna remove it from? Well, our first step is to remove our P electrons, right?
39:25
And this sometimes gets confused in the midst of the different rules for where you remove from when. When you're in the S block, you remove from the S block first. When you're in the P block, you remove from the P block first. That switch comes when you're in the D block.
39:41
Then, and only then, you switch around and then you remove from the S block first. So we are well within the P block of the periodic table. So we're gonna remove the P electrons first. And it becomes the noble gas core, five S two, four D 10.
40:01
So what charges does that, that gives us a three plus charge, right? We just took away the P electrons, leaving, taking away three electrons and electrons are negative. So if we take three of them away, we're left with the cation. So we take the cation and we make it SB three plus. So now we're basically in the D block, right? We're basically sitting here
40:20
at the cadmium electron configuration. So now where do we remove from? Now is where we remove from the S block. So when we remove from the S block, we remove both of these electrons and we get the noble gas core plus four D 10, okay? So these are valence electron configurations.
40:42
So you'll see, so the reason that this is too different is because at this point you removed from your S block. So you removed two of them. Now you wouldn't remove just one of them, right? Cause now you're at that sort of in between, you're not stable, so you wouldn't wanna just remove one so you're either gonna remove all of your Ps
41:01
and none of your Ss or all of your Ps and all of your Ss. Okay, so that's SB. Now let's move on to lead. So if we go through and we write out our electron configuration for just neutral lead, we get this. When we need to remove, we're sitting in our P block.
41:24
So we remove from our P block first and that gives us what charge? Two electrons, we take those two electrons away, we have a two plus charge. So at that point, now we're sitting in our D block, we're sitting at like mercury.
41:41
So how do we, can we remove more? Well sure, we can remove these two S electrons without too much difficulty, which leaves us with four plus and it gives us this electron configuration. So all of these are gonna end up with a difference of two because one set of the electron configurations, the lower one, is gonna be
42:00
if you just remove the P electrons. The second set comes into play when you remove the S electrons, which there's two of them, so it's always gonna be off by two. And so that gives you these gaps. So up until now, we were just using this as a help to memorize. Now I could ask you to explain it and you could say well, here, here's the electron configuration for one of them,
42:21
here's the electron configuration for another one of them. So you can see that we're always gonna be removing these P electrons first and then the S electrons. So there's always gonna be a difference of two. So then why the difference charge between this column, this group and this group? Well, they have a different number of P electrons, right? This only has one P electron, so you're gonna get a plus one
42:41
and then when you remove the S electrons, you're gonna get a plus three. How many P electrons do these have? Two. So you remove the two P electrons and then you remove the two S electrons. These have three of them, so you're gonna take away the three P electrons first and then the two S electrons.
43:01
So go through for the other four and write out all the electron configurations. It's good practice writing out electron configurations and you can use, you can explain this. Okay, so now we have one more of these sort of periodic groupings to talk about. This is called the diagonal relationship. So this is something that sort of comes out
43:23
of all of the other trends and what happens here is let's, so let's look at this part of the periodic table real quick. And you have this sort of diagonal trend going this way. Now, you may think back to questions that people said, well, can you ask me this?
43:41
And maybe the question was, which ones, you know, more has a higher electron affinity, phosphorus or selenium? Because in one case, the trend goes this way and in one case, the trend goes up. So it was these diagonals against the trends that I said, I can't ask you. I can't ask you which one has it, which one's bigger, phosphorus or selenium?
44:02
Because going down the periodic table, they get bigger. So that would say that selenium is bigger, but going across the periodic table this way, well, they get smaller, so selenium would be smaller. So according to one trend, it's bigger and according to one trend, it's smaller. So which one is it? And I said, I don't know, I'd have to look it up.
44:22
And I also had said that it'd be close, right? Because those two trends, they're kind of battling with each other. One's making it bigger than the other one, one's making it smaller than the other one. And so they're actually very similar. And that works for the other trends too, that works for ionization energy and electron affinity, all of the trends that have that same direction.
44:41
So if we're talking as we go down the periodic table this way, that each time we're saying, okay, well, this one increases as this one decreases, what happens there is that they happen to be very similar. They have similar ionization energy, similar electron acidities, similar sizes. And so they end up with similar properties.
45:04
So here's that in sort of picture figures from your book, that's what the figures from your book were trying to get at. If we look here at our sizes, as we go this way, there's an obvious trend, right? And that's those trends that we've been talking about. Fluorine is our smallest, fluorine is our most electronegative.
45:21
And then as we come down here, cesium would be our largest, it would be our least electronegative. So when you go this direction though, the diagonal this way, now it's harder to see. You look at this and this and this, and they're pretty close to each other. This and this and this, and they're pretty close to each other.
45:41
Same thing as you go this way on the periodic table. And that makes it so that you end up with really similar properties. So that's called the diagonal relationship. Now that's obviously very different from the fact that we have similar properties as we go straight down a group, or straight down this group, right? That's based on electron configurations. That's based solely on, okay,
46:01
you have this many electrons in each of your shells. This just comes out of the fact that, hey, these have similar numbers for all of these things that we've been talking about in this last quarter of the chapter. And so this is what the diagonal relationship is.
46:21
Okay. So that pretty much wraps up chapter one. That kind of ends the first exam material. And all of the sort of relationships that go along with it. So at this point we'll end for the day. And then next time we'll start up with chapter two,
46:41
and starting to talk about how to draw Lewis structures and how bonding works and things of that sort.