Alright, this is lecture 30.2 and we're continuing to talk about atomic properties and the first atomic property we're going to take into consideration and talk about in any detail is atomic size. And atomic radii are the same thing as atomic size. This follows two different trends. And one of them is... pretty easy to understand and the other takes a little bit more information to really understand what's going on.
So, first and foremost, atomic radii increases going down a group. So that's going from top to bottom in the periodic table, i.e. the arrow pointing down, and the radii decrease going across a period. And that's the one that's going to take a little bit more information for you to really understand what's going on with that.
So atomic radius is measured as half the distance between adjacent nuclei in a molecule. So we're basically going from the nucleus to the nucleus. So the radii is the distance between them as shown in the image here. All right, and so we're not going to be measuring any atomic radii per se. We're going to just assume that...
We will be given measurements if we're asked to determine those radii, but we want to talk about why atomic radii change as we are looking at different elements in the periodic table. So to start with, let's go with the easy one first. The trend going down a group is that it increases. And the reason it's increasing is, well, we're adding shells, essentially. If you think of the Bohr model, and you're going from n equals 1 to n equals 2, n equals 3, and so forth, you know that these shells are higher and higher energy and getting further and further away from the nucleus.
And so, basically, as a principal quantum number changes, or the principal energy level increases, and electrons move further and further from the nucleus, obviously, this is going to increase. increase the radius of the element. And so basically more and more electrons, bigger and bigger atoms as we go down the periodic table. So atomic radii increases.
So that's easy enough to understand. And here's just an example. And as you can see as we are going down, the n number is increasing. The shells, you know, n2, 3, 4 and so forth.
And as you see, the radii increase as we go down the list and that makes sense. We've got more electrons, we've got a larger atom and so that's not anything, you know, real difficult to understand. Now, the thing that's a little bit trickier to understand is why atomic radii decrease, right?
We're talking about decreasing as we go across the period. So these are the rows, so we're going across a row in a periodic table, and as we go from the alkali metals all the way over to the halogens, the atomic radii decrease. And that's a little bit harder to understand, but it makes sense. Let's just say that. So what we're going to be looking at is both looking at the atomic number.
And of course, in a neutral atom, atomic number tells you not only the proton count, but also the electron count. And where are we actually adding the additional electrons? And that's where the crux of this comes in. So as we move across a periodic table, we notice that the atomic numbers increase, right?
That's how it's organized. And each element is one. proton larger than the one before it. And this means that the positive nuclear charge is increasing. So you're getting more and more positive charge as you go across the period.
Now we know that the periodic table is organized such that, you know, n equals one is the top row, n equals two is the second, n equals three, four, and so forth. So this means that each row in the periodic table is filling a shell. That means that the electrons that we're placing into these shells, whereas when we're going down the periodic table, the electrons are being placed in successively larger n values.
going across the periodic table, they're all going into the same n value. So they're all going sort of into the same general area around the nucleus. So the electrons are not going into higher levels. They're going into the lower or the same levels.
Now let me explain this nomenclature, so to speak, here. As you know, n equals 1. only has two electrons in it. N equals two has two and then six, so eight possible electrons. And so what these numbers are here, two and eight, well this would be the N equals one and N equals two shells.
So this is telling us that these are full and the electrons now this is the n equals 3, the electrons, there's one going in there, and then two going in there, and then three going in there. And so as we go across the periodic table, the core electrons are remaining exactly the same, and we're only placing electrons in the outermost shell. And so what's happening as we go across the periodic table, the number of protons is going up, right?
It's increasing. But when we're putting in electrons, so we're putting electrons in these shells, we're not going out to put them in, we're putting them all in at the same level. So what's happening is the pull, in other words the nuclear pull is increasing, whereas all of the electrons that are around here, are going into the same level. Now the electrons that are in here, these are what we call shielding electrons.
And so this is the amount of shielding right here in the inner circle. These are the shielding electrons. And what they do is they shield the outermost electrons from the nuclear charge. Well, as we go across a period in the periodic table, the number of shielding electrons, as you see, is not changing. It's basically staying exactly the same.
So they're not getting any more shielding, but the number of protons is going up. So the charge, the effective nuclear charge is what we call it, is increasing steadily going to the right. So think of it like a magnet.
We're getting a stronger and stronger magnet. The shielding between what the magnet's pulling on is staying exactly the same. So as we go across the periodic table, the shielding becomes weaker and weaker because the effective nuclear charge is becoming stronger and stronger. And that means things at the outermost that count on that shielding are being pulled closer and closer in as the A magnetic charge, so to speak, increases.
And so as the atomic number increases across the row, the additional electrons are all added just to that same or even lower. As you see here where we're showing this, these are where we're actually filling in electrons in even lower fields. This is, remember, they don't all fill sort of sequentially. They jump around when we get to the higher levels. And so you're basically going to have an effective nuclear charge that increases, and that means that these electrons are being pulled more strongly towards the nucleus.
And so effective nuclear charge, this is one of the few equations that you have to know. It's really more of a relationship, but the effective nuclear charge, which is symbolized as Z-effective, is just the nuclear charge, this is the number of protons, okay, and from that you're going to subtract, this is the core electrons, okay. So these are the shielding electrons and these would be the electrons from the previous n-value shells. So in this case this is of course n equals 1 and n equals 2 shells. completely full right that's what these shielding electrons are so if we were calculating this we would have 11 So, you know, we'd have 11 as our number of protons minus 10, so our effective nuclear charge would be plus 1. And as you go across, of course, the effective nuclear charge increases a lot.
So here's the sodium we were just looking at. This is the 2, 8 with the 1 electron in the n equals 3 shell. That's its electron configuration. And so as it goes across, right, you know, neon.
This would be neon 3s1 would be the electron configuration of sodium. And so when we look at this, you notice that, you know, there's just the one electron after the noble gas. The noble gas configuration is the core electrons. So again, the effective nuclear charge is plus one. As we go across that period, when we get to...
the next element here, this is your sulfur, you'll notice that we've added more electrons to this shell and this electron configuration, but the effective nuclear charge is only the core electrons. And so the core 10 electrons have not changed. and that means that the effective nuclear charge is plus six.
So no doubt a plus six charge is going to pull these outermost electrons, these outer six electrons, closer in because the core shielding electrons have not changed. And so this is what causes the atomic radii as you go across periods in the periodic table to decrease. And here's basically sort of a schematic of these trends, and as you can see as you go across, These things do shrink, so to speak. They increase going down, and that makes sense because you're getting more and more shells and more and more electrons. But going across, because of this effective nuclear charge, you see this reduction in the size of the element's radii.
Okay, so these are the two trends, and you need to understand the trends and what causes those trends. Okay? And you should probably be able to calculate an effective nuclear charge for an element or elements. All right. Now, we've done the neutral atoms.
So what happens when we turn these atoms into ions? And so we want to go through how the ionic radii will vary. So positive ions. We know that positive ions... occur when they lose their outermost electrons they always pull from the outermost shell and we talked about that and you make your electron configuration you have to pull from the outermost shell aluminum's electron configuration is essentially represented here again this is in you know so this would be 1s2 and then 2s2 2p6 and then we have 3s2 3P1 basically, okay?
And so this is the electron configuration for aluminum. If I want to make an ion out of aluminum, it likes to lose three electrons, right? Because then it can be like the nearest noble gas.
So this would be neon's electron configuration with just the two and the eight filled. And so These are how positive ions form, and this is just a review, so you should be able to do this and know what electrons in an electron configuration need to be lost to make a stable ion out of an element. So we're going to get element aluminum 3+, and so when we look at this, we have positive ions and they're going to essentially remove the outermost shell electrons.
So you're basically losing a shell. And if you lose a shell, obviously it's going to shrink. So positive ions get smaller, right? So they get smaller, all right? And their trend, you know, has this...
more positive charge and so you have a stronger effective nuclear charge because think about it, you basically now have 13 protons for aluminum, but you only have 10 electrons. And so this is going to pull even those core electrons closer to the nucleus. So positive ions are very much smaller. They're smaller than their element.
And so here's a table, and even the ionic... electron configuration is the same as the ionic radius gets progressively smaller as you go across. So sodium is plus one. Notice that it has an atomic radius of 190 because they're fairly large on the positive side of things but the ionic radius is is much smaller and your ionic configuration is the same.
But again why is this happening? So if our ionic Okay, so this is the electron configuration, right? For each of these elements, they're attempting to match their configuration to the nearest noble gas, which is neon.
And so they have the same exact same number of electrons. But they have different numbers of protons. And so what's happening is the proton is the pull towards the nucleus.
If you have the same number of electrons and you have a stronger pull as you go, It's going to get smaller and smaller and that's exactly what it has, you can tell what's happening. And so the effect of nuclear charge also means positive ions get smaller and smaller as you go to the right in the periodic table. Okay.
And so here are some comparisons and you look at this, look at the change that takes place in terms of the ionic nucleus. You know, this is... really a very very dramatic change in the in the radii because of that charge and that effective nuclear charge now you'll notice in these you'll see a trend starting over here in these negative ions and you can tell when you start adding electrons what happens of course you get a larger ionic radius and so if you look at the nonmetals here we got chlorine We're going to go from 287 to 288. We've added electrons. And, of course, that means the effective nuclear charge is decreasing.
Okay. So we've gone to a, you know, a full shell, so to speak. And so we are basically just making a bigger atom.
So if you add more electrons, you make a bigger atom. All right. And so, again, you know, if you go down the principal energy level, everything gets bigger, right?
We're going down the, you know, in the periodic table, if you're going down, size increases. And that's because you're adding shells. But going across periods, you're looking at... two different trends with respect to ion charge and size and think about it this way so when you're talking about positive ions right where do we switch from positive to negative ions well we do that when we go from metals to nonmetals right so as long as you're in the metals you're going to see the size decrease steadily as we go across the periodic table.
So if we go from left to right, you're going to see that the size decreases steadily, and that's because the charges are increasing, right? They're becoming more positively charged, and we just showed why that would happen. You know, the effect of nuclear charge is getting greater and greater as we go from left to right.
But then as soon as we get to the negative ions, then you're going to be seeing a difference, right? We're going to be seeing that the size goes up when we start going into the negative ions. And so here's sort of a, I guess, a summary you can see in this sort of table here. So let's go across.
Let's try this row. So we see that potassium plus is 138, and then calcium 2 plus is 99. Allium is 62. And again, we're still in positive ions, this is 69 which is a drop from here, but it's not as much of a drop. You see it starts, and this is in the kind of the metalloid section around here, you'll see it starting not to drop as much.
And then as we go up, what you're going to see is that it's going to jump up, but then it's going to start going down again. And the reason for that is we add three electrons here. You know, we add two here and add one here.
So you're going to go kind of to a low, and then you're going to have a, so it's going to go down, and then it's going to jump up, and then it's going to start going down again, okay? And the reason it's going down is because you only added one electron here, you added two here, and you added three there. So these are the trends that you see in the periodic table.
But the important part is you know why those trends exist. You know, if you need to go back over this, you need to listen to somebody else describe it, you definitely want to read the book and get your mind around why these trends exist. Knowing the trends, you know, it goes up, it goes down, it's increasing, decreasing, that's great. Knowing why the trends exist is the most important thing.
All right, so that's the end of this one.