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Understanding the Trend of Atomic vs Ionic Radii
Let's talk about understanding the trend of atomic radii versus ionic radii. This trend comes in handy across the board, when talking about many different aspects of the atom. When talking about the chemical behavior of the atom, even some in physical behavior of the atom. So you really have to understand why the atomic radius behaves the way that it does, and it will help you with other trends in the periodic table as well.
So when you have the periodic table, this box on top represents the periodic table. We know that the atomic radius decreases as we go across, and increases as we go down.
I do want to note, however, I did not include the transition metals in this. This is only the representative atoms, because Transition metals typically are all relatively the same size within a period. Yes they increase as we go down too, but they do not decrease as we go across. So they aren't really within this trend at all. They have their own little entity. So this is just representative atoms. That's a good test question by the way. I've seen that on many tests before.
So let's look at the atomic radii, and why it does behave the way that it does. So I've put up here Sodium, Phosphorous and Argon. I drew these, but not drawn to size or drawn to perspective. So they're all just drawn. So we have 11 protons in the nucleus of Sodium. 15 in Phosphorous, and 18 in Argon. Around Sodium we have one electron, then have 5 here, and then 8 here. Then these little lines represent the electron. Atoms don't look like this, but it helps us understand what's going on.
So electrons are kept close to the nucleus by the pull of the protons. So the more pull we have from the protons, the closer this outer electron is so and then smaller the atom.
So here looking at Sodium, we having 11 protons pulling on that one outer shell electron. I'm not really caring about the inner shell electrons right now. With that one outer shell electron, so you have 11 protons pulling on it. You have 15 protons pulling on this guys. This guys we have 15. This guy we have 11. This guys is probably going to be smaller. This guy is going to be small. Then they have 18 protons pulling on this one.
They have the same amount of energy levels between them. They're in the third period, so they have three energy levels. All of these too. So this guy is 18, this guy is the smallest. This is because there is 18 protons pulling on the outer shell electron. So this is due to effective nuclear charge. You might know that.
So the higher number of protons in the nucleus, it's where the nuclear comes from, higher number of protons is able to pull that outer shell electron towards it making it very, very small. So even though you think it weighs more and has more electrons, it has a higher number of protons, that's able to shrink it down. That's why it decreases as you go across.
What about as you go down? So we have Sodium. Right underneath Sodium, in the periodic table, is Potassium. Potassium also has one electron in the outer shell. Going by what we just discovered, or what we just talked about, Potassium has 19 protons, Sodium has 11, you'd think it be even smaller, because 19 protons is more of a pull on the outer shell electron. But it has an extra energy level, extra shell so it's not able to pull in as much. These inner shells, I know you've heard this before, create a shield between the protons and that outer shell electron, allowing that outer shell electron to expand or get bigger as well. This is a shield, and this is what we call the shielding effect. It's when the electrons are able to get further, and further away from the nucleus. Nucleus doesn't have a great pull on outer shell electrons due to inner shell electrons blocking its effective nuclear charge. So this is big. It's bigger with each shell we put on it.
The bigger the shield, the bigger the atom gets, so the more shielding effect. It's very important. This is an answer on a lot of questions. Shielding effect is the answer, it can explain ionization energy. It can explain electronic activity. It can explain a lot of other phenomenon within the periodic table, because it talks about how the outer electrons work, and why they behave the way they do. Shielding effect is a massive influence. So we left atomic radius, the atom.
Now we're going to learn about ionic radius. The ion. What is the difference? It's the size of ions, which we know ions are charged particles. They're charged, meaning we took away or gained electrons.
So looking at metals, metals always they lose electrons. They become cations. So things in group 1 lose one electron. Things in group 2 lose two electrons. Things in group 3 lose three electrons. If they lose electrons, they essentially have a greater effective nuclear charge. So we have more protons than we do electrons. So we're able to squeeze that in. So these in general are smaller than the atom. If they lose 1, it's small. If they lose two, it's even smaller. If they lose three, it's the smallest.
Looking at non-metals, they gain electrons. They become anions. Things in group 7 gain 1. Things in group 6 gain 2, things in group 5 gain 3. These overall they're gaining electrons. Think about weight. When you gain weight, you become bigger. Overall bigger than the atom. If they lose 1, it's big. They lose 2 it's even bigger. If they lose 3 it's the biggest.
So it doesn't have a definite trend. Notice that it's the same kind of idea, but they don't have a definite trend as we go across. It changes a little bit. It kind of has this kind of effect here. As we go down, the trend is the same. We're still adding on layers of energy levels, and extra electron shells. So it does get bigger, because of the same reason they got bigger here due to shielding effect. But as we go across it's different.
Understand the difference between ionic radius, and atomic radius and why it works the way that it does. Because if you do, it could you understand lots of different things, as I mentioned before. So hopefully this helps you understand it just a little bit better.