As we continue on in this unit called Periodic Trends, our next trend we're going to be looking at is ionization energy. We're going to define ionization energy, and we're going to look at the role of effective nuclear charge, in terms of the trends of ionization energy, as we go across the periodic table, or up and down a group in the periodic table. So let's begin by defining ionization energy. Ionization energy is the minimum energy required to remove an electron from a gaseous atom in its ground state. So I've got in I guess that's pink there, various words: remove, gaseous atom and ground state, because these are important components of this. So let's imagine a simple atom of lithium. Okay, so we had drawn lithium in our last learning objective in which we had the three protons. Okay? Then there's two electrons in the 1s orbital, and then out here in the 2s orbital, there is an electron. 'Kay? We're going to take one of those atoms in the gaseous state, okay, and that is just so that when we compare from one atom to another, we're making a similar comparison. Because removing an electron from an atom that's in the solid state is different than the energy to remove it from the gas state. So, we have it in the gas state. It's a ground state electron, so it's already in its lowest energy state. We haven't promoted this electron up to some higher energy level and it's sitting there. Okay? So that's not where it is. It's in it's ground state sitting in that 2s orbital. The minimum energy would be the energy to remove the furthest away ones. So we're going to take this electron, and we are going to completely remove it from the atom, and measure the energy it takes to do that. Now it always requires energy. So there's another word that's an important word. It requires it. Well, why would it require energy? Well, that electron is being attracted by an, by the effective nuclear charge, to that nucleus. It is feeling a pull by that nucleus, and we have to pull against that in order to remove that electron. So it since it's a require energy, it's an endothermic process, so it's always a positive value. Ionization energies are always positive. So what is this a measure of, if you're measuring ionization energy? It's got to be a measure of how tightly that electron is held to that atom. Well we know that size was also a measurement of that. The more tightly the atom was drawn towards the nucleus, the smaller it was going to be. The more tightly it is drawn toward the nucleus and the smaller it is, the harder it's going to be to remove an electron from that atom. So, if we look at it in an equation form, this would be the equation. We start with an atom in its gas state. And we remove the electron from it. What will that leave us with? It'll leave us with a cation and its electron. Okay. So, that's how you would represent that. So, it's the energy associated with that. Sometimes we would say that the, well what we would call this is the ionization energy. And we might call it I1, because we're removing the first electron. Because we could remove more electrons and we will do that as we proceed. Okay? Oh, that's right here on this slide. First ionization energy, which is I1, is the energy to remove the first electron. But we could also talk about the removal of a second electron, so if we start out with a cation already, we've already removed one electron, and we're going to go in there and pull the next available electron out. This would be the second ionization energy. So you think about that. In this case, we are removing an electron from a neutral atom. In this case we are removing an electron from something that's positively charged. Can you imagine which one would require more energy? Well, you think about that and answer this question. Well if you chose number two, you'd be correct. Okay. It always requires more energy, no matter which atom you're talking about, it's always going to require more energy to remove the second. And would could keep that going, and going, and going. We could remove more and more electrons as we go through and it would keep requiring more and more and more energy to do so. Okay, let's look at this table here. Now these are reported in and I've got a little bit of a word here that we have not yet defined, maybe you've seen it in some previous class you've had, but a mole is a certain quantity of these atoms, okay? So this is how much it would take to actually remove the electron from and the number is 6.02 times 10 to the 23rd of these atoms or ions. Okay? Just as a note, that's what that is referring to, how many you have. This is the amount of energy to remove an electron, and it's the first ionization energy, the second ionization energy, third ionization energy, and fourth ionization energy. And we can see as for boron, that it is getting higher and higher as we move across. You take any one of those, and the ionization energy is going to increase as you move across, because you're pulling off more and more electrons, okay? And what remains behind is being pulled in more tightly by the nucleus. You take off another electron, the remaining electrons are pulled in more tightly, and it keeps getting smaller and harder to remove an electron. Now we can kind of tell how many electrons are in a valence shell of an atom by looking at the trend. Let's choose, for example, this one here. Let's see what's happening to our, our ionization energy. So we start off with fro, four protons, okay, and we know we have those four protons in the nucleus. We know we have two electrons in the 1s orbital and then we have two electrons in the 2s orbital. So there is its electron configuration. I'll write it out. That was helium 2s2 or 1s2 2s2. Now when we move our first electron and take this one away, it's going to require this much energy to do that for a mole of these atoms. That is going to make it a little smaller, as we have less electrons vying for those four proton pull. So it gets a little bit smaller and it gets a little bit harder to remove the next electron, so I pull that electron away, it gets a little bit smaller. But now we've removed the valence electrons, and we're starting to go to the third, we're going to try to remove this one. Well, this one has very little shielding and it's really close to the nucleus, so we see a huge, while it increased some here, it is a huge increase as we go from the second to a third. That is an indication that the second was still in the outer shell, and we have now reached in and started taking electrons out of the inner shell. And so you can always look where the cutoff and the big jump is, and say, this is how many outer shell electrons we have and this is the, and know what group the element is in. If you didn't know it was beryllium, if you just saw these numbers you could say this guy's got two valence electrons. If we look at boron and didn't know where boron sat on the periodic table, we see an increase, an increase, an increase, and then a huge increase. Okay? So this increase, we could put a line right here and say that these must be valence electrons. And now we've reached in for this one and pulled out an electron in an inner shell. Now let's examine more closely just the first ionization energy, and let's look at the trends. Let's begin by looking at the trend that we see right here. This is going across one period. So if we were to draw a rough sketch of a periodic table here, okay? Here it is. We are looking, as we go across the period, starting at lithium, and working our way over to neon. We see, as a general trend, in the first ionization energy, it always increases, as you go across the periodic table. That's a general trend. Now, we see some zigs and zags in it. We'll talk about those as we go, but, as a whole, it increases as you go across. Well, let's think about why that would be. If we go back to our last learning objective, we talked about the effect of nuclear charge. And we talked about as you go across the periodic table, the effect of nuclear charge is going up. It is getting, is drawing those electrons in more tightly, and we saw that play out as a smaller and smaller atomic radius. As they get pulled in tighter and tighter to make the atoms smaller, it's going to get harder and harder to pull that electron off. So that fits what we understand. Now let's look at the trend, down a family or group, okay? So we start with neon, we're just looking at the, sorry start with helium here. As we go down the noble gases, we see that it is decreasing its first ionization, okay? So we see a increase as you go up the family, or decrease as you go down the family. Whichever way you want to think about it. Now a lot of times this is the way we think about it. Ionization energy increases as you go up, and it increases as you move to the right. Okay. First, ionization energy increases up, and it increases across. So why would it do that? Well, as you go up a family, again the atoms are getting smaller and smaller, because they're being drawn more closely to that nuclear, that nucleus. And it's going to be harder and harder to remove the electron as you move up a family. So it's more difficult to take off an electron if it's high on the periodic table that it is if it's down low, because its valence electrons are very far away from the nucleus, there are a lot of electrons shielding it between in its outer shell and that nucleus and it is much easier to pull off the electron to remove it from the atom. Okay let's focus a little bit more closely on this area here, and let's see the the drop that we see at boron and at oxygen. If we look really closely at this period, okay. Question I have is, why are there dips at boron and oxygen? We did not see those dips in terms of the effective nuclear charge. We didn't see those dips in terms of the atomic radius. It got smaller and smaller as we went across. And what these dips are saying is, yes, it requires enerdy, energy to remove my electron, but it's less than you would expect. Okay? They're kind of on sale. They're on sale a little cheaper than what you would expect. Now, why would that be? Well, always to understand the exceptions to these trends, it's helpful if we look at the electron configurations. Okay, so let's do that. Let's look at the electron configuration of these two atoms. For beryllium, the electron configuration, and that's what sits right here, okay, I don't have it written down, but there's beryllium. Beryllium, the electron configuration is helium, 2s2. Boron's electron configuration, which dips, it's the one that's on sale. It's a little cheaper than you would expect. It's 2s2, 2p1. All right, so could you figure out why boron has a lower ionization energy? Well, what are you going to remove? You're going to remove this electron. And we learned in a previous unit that there is stability to a half-filled or a completely filled subshell. So if we remove that electron we are left with only filled subshells, so that is going to be, still requires energy, but it's on sale, it's cheap. It says, I will let you have this electron and I won't make you pay as much for it in energy if you'll take it away from me. Now let's look at the next drop. The next drop goes between these two elements, okay? So the first one is nitrogen and the next one is oxygen. So we could look at the electron configuration of those two. So what is nitrogen? It's 2s2, 2p3. Whereas oxygen is 2s2, 2p4. So if you look at those two electron configurations, can you figure out why oxygen has a lower ionization energy than nitrogen? Why is it on sale? Okay, why does it say yes, it requires energy, it still requires energy, but I will not require as much energy to remove that electron. Well, if it takes that one electron away it drops it down to three and that's a half-filled. So with ionization energy, anytime you see a break in the trend, you can imagine that the reason behind it is going to be that you are going to be either cleaning out a subshell, or you are going to be half-filling a subshell. So at right under boron it's aluminum, right under oxygen is sulfur, so we see the same drops that are occurring at those regions. So this is the end of learning objective number four. We've looked at ionization energy, we know its definition, we see its trends up and down the group, and across the periodic table, and we have tied that to two things. We have tied it to effective nuclear charge, and the attraction that the electrons have for the nucleus. And we've also tied it to its electron configuration when we're looking at breaks in the trends that you would expect.