 So let's talk a little bit about a word you might have heard. And that is ion. Let's talk about what it is. And then we'll talk about trends in the periodic table on, I guess, how hard it is to make something an ion. In particular, how hard it is to make something a positive ion. So an ion is just an atom or molecule that has charge. And it'll have charge if the protons are not equal to the electrons. Neutrons are obviously also a constituent of atoms, but neutrons are neutral. What you're going to get your charge from are your protons or electrons. So you're going to have a net charge if your number of protons, and this is for an atom or molecule, molecules just a bunch of atoms bonded together, if the number of protons does not equal the number of electrons. And you can have positive ions if the protons are more than the number of electrons. Protons are positive, electrons are negative. And you could have negative ions if the number of electrons are greater than the number of protons. For example, if you just had hydrogen in its neutral state has one proton and one electron. But if you were to take one of those electrons away, then hydrogen would have a positive charge. And essentially, it would just be in its most common isotope, it would just be a proton by itself. And so when we talk about a positive ion like this, where our protons are more than our electrons, the number of protons are more than the number of electrons, we call these cations. Cation, once again, just another word for positive ion. Likewise, we can have negative ions. So say, for example, fluorine. So fluorine gains an electron. It's going to have a net negative charge. It's going to have a charge of negative 1. And a negative ion, we call an anion. And the way that I remember this is, a kind of means the opposite or the negation or something. So this is a negative ion where we're negating. You could somehow think we are negating the ion. So with that out of the way, let's think about how hard it will be to ionize different elements in the periodic table, in particular, how hard it is to turn them into cations. And to think about that, we'll introduce an idea called ionization energy. Ionization energy. And this is defined as the energy required to remove an electron. So it could have even been called cationization energy, because it really is the energy required to remove an electron and make the overall atom more positive. So let's think about the trends. And we already have a little bit of background on the different groups of the periodic table. So for example, if we were to focus on, especially, we could look at group one. And we've already talked about hydrogens, a bit of a special case in group one. But if we look at everything below hydrogen, if we look at the alkali metals here, we've already talked about the fact that these are very willing to lose an electron. Why? Because if they lose an electron, they get to the electron configuration of the noble gas before it. So if lithium loses an electron, then it has an outer shell electron configuration of helium. It has two outer electrons. And that's kind of, we typically talk about the octet rule. But if we're talking about characters like lithium or helium, they're happy with two, because you can only put two electrons in that first shell. But all the rest of them, sodium, potassium, et cetera, et cetera, if you take an electron away from them, then their outermost shell, well, all of them, they're outermost shell, they're going to have the electron configuration of the noble gas before it. And for sodium on down, that outer shell is going to have that perfect eight. Lithium, if you remove an electron, it would get to helium. It will have two electrons in its outer shell. So you can imagine that the ionization energy right over here, the energy required to remove electrons from your alkali metals is very low. So let me just write it down. So when I say low, I'm talking about low ionization energy. Low. Now what happens as we move to the right of the periodic table? In fact, let's go all the way to the right of the periodic table. Well, if we go here to the noble gases, the noble gases we've already talked about, they're very, very, very stable. They don't want their electron configurations messed with. So it would be very hard. Neon on down has their eight electrons. That octet rule, helium has two, which is full for the first shell. And so it's very hard to remove an electron from here. And so it has a very high ionization energy. Low energy, easy to remove electrons, or especially the first electron. And then here you have a high ionization energy. You have trouble seeing that H. So this is high, high ionization energy. And that's the general trend across the periodic table. As you go from left to right, you go from low ionization energy to high ionization energy. Now what about trends up and down the periodic table? Well, within any group, even if we look at the alkali, if we look at the alkali metals right over here, if we're down at the bottom, if we're looking at, say, cesium right over here, that electron in the one, two, three, four, five, six, in the sixth shell, that's going to be further from that one electron that lithium has in its second shell. So it's going to be further away, it's not going to be as closely bound to the nucleus, I guess you could say. So this is going to be even, that one electron is going to be even easier to remove than the one electron in the outermost shell of lithium. So this one has even lower, even lower. Even lower. And that's even going to be true of the noble gases out here. That xenon, it's electrons in its outermost shell, even though it has eight valence electrons, they're further away from the nucleus. And so the energy required to remove them is still going to be high, but it's going to be lower than the energy from, say, neon or helium. So this is low. So once again, ionization energy, low to high as we go from left to right, and low to high as we go from bottom to top. Or we could say a general trend that if we go from the bottom left to the top right, we go from low ionization energy, very easy to remove an electron from these characters right over here, to high ionization energy. Very hard to remove an electron from these characters over here.