We are now with this learning objective ready to bring all the pieces together. We've looked at trends of ionization energy, size, electron affinity across a period, up and down a family or group. We have seen those trends, now we're going to try to understand how the electron configuration, the ionization, the electron affinity, all play out in the reactivity of the elements. And we're going to focus on the representative elements. because that's what we've mostly been examining, when we've looked at these trends on the periodic table. What we see is that elements in the same group resemble one another. Now, remember, when we talked about the design of the periodic table, they knew nothing about the electron configurations. They just started putting things together according to putting them from smallest to largest, and putting them in to where they overlapped if they had similar chemical behaviors. Now we're going to see that that information that we obtained gives them similar, why? And now we're going to see why all of that plays together to give us that similar behavior. Now the fact of the matter is, when we are going down a group, the elements in a group resemble each other, but the elements in the second period are quite different from the rest of theirs. So there is a bigger difference between th, the element in the second period, and the elements below it. What about the first period? Well, the only one that has ones in the first period would be hydrogen. And hydrogen is unique all of its own. There is something that is called the diagonal relationship, okay. And when you're talking about that element in the second period it is more similar to the element to its diagonal right. So this diagonal relationship results in, if you have a, your periodic table, look at lithium and magnesium, they're diagonal to each other. Lithium is here and magnesium is diagonal to the right. Okay? Beryllium is here and aluminum is diagonal to its right, boron is, find it on the periodic table, silicon is diagonal to its right. So in many ways the element that's in that second period is more similar to what's diagonal to the right, and we can really see that play out here. Look at boron and silicon on your periodic table. These are both metalloids. Okay. That is not true for any other elements that are in the group with boron. Okay? All the other elements are metals down there. Boron is the only metalloid. Silicon is a metalloid, so boron is in many ways more like silicon than it is the rest of its family. So that is called the diagonal relationship. Now we're going to go through various representative elements. Remember, our representative elements are the elements on the tall towers on either side of the periodic table, but not the noble gases. They're not considered representative elements. Representative elements have incompletely filled s and p subshells, so it had to be incompletely filled. And our noble gases don't fit that criteria. Now I mentioned that hydrogen is the most unique of the representative elements. There's really no suitable place to put hydrogen on the periodic table. And I've seen some periodic tables where they'll just stick it up off the periodic table by itself. I've seen periodic tables put it above the alkali metals, though it's not a metal. And put it above the halogens, it's not a halogen. So they've, I've seen them locate it in two locations, take it out of the periodic table altogether, but most commonly it sits above the 1A family because it has one valence electron. Hydrogen is unique in that in, if it bonds with metals, it tends to gain an electron, and we call this the hydride ion. But if it were to bond with non-metals, it would become an H plus, okay. So we have acids, for example HCl. If we were to dissolve that into water, we will be seeing eventually that this is what they do. We have those ions associated it. It is actually a non-metal, okay? Hydrogen is a non-metal though it sits with the 1A family. So it's quite unique. But let's look at the real 1A family, the alkali metals, okay? What do you, what do they all have in common? What they all have in common is their valance shell is going to be some number and then s1. So for example, lithium is 2s1, so it's helium 2s1 and so forth. They all, because they have only one valance electron, they form plus 1 cations. And we think about the trends of ionization energy. It gets larger as we move across the periodic table. Ionization energy increases as you go across the periodic table. So these guys over here on the left-hand side have the lowest of the ionization energy. Not only do they form plus one cations but it's really easy, it still requires energy to do that. 'Kay, so don't think that it doesn't require energy just because it wants to get rid of its electron. It still requires a little energy to remove it, but it is on sale very cheap, it doesn't take much energy. Because it's a low ionization energy, these are highly reactive. Now what happens to ionization energy as you go down the family? So as we're going down the family on the periodic table, that valance electron is far, farther and farther down. So, this is going to have the lowest of the low. The lowest of the low down here. So, these are highly reactive. What's the reactivity going to do as you go down the family? Well as you go down the family, the reactivity is going to increase. Okay. So, the question is why, and the rean, the answer is, the reactivity is going to increase because it's easier and easier to remove that electron. Now, these metals are so reactive that they will easily react with water. You can take the metal, and you can place it in water, and it will react with water. I want you to watch this demonstration, you're not only going to see the fact that they're reactive, but you're also going to see that they're more reactive as you move down the group. Within our unit we have been talking about ionization energy, and we've been talking about how the alkali metals have a very low ionization energy. They're highly reactive. Let's talk about trends of that ionization energy as you move down the periodic table. So as you go down, it's easier and easier to pull an electron off the atom and turn it into a cation. I'm going to demonstrate for you how that reactivity increases as you go down that group, and I want to demonstrate it with a piece of sodium in water, and then a piece of potassium in water. If you look at the equation here behind me, it's not a balanced equation, but we see that sodium is reacting with water, and it's going to produce sodium hydroxide. And this is the base, we haven't talked about bases yet, but we'll see the evidence of the base. It's also going to produce hydrogen and that's a gas, and we'll obviously see that gas being produced during the course of this reaction. Then we will carry it out with potassium instead, cause potassium is further down the periodic table, it is a little more reactive, and therefore we'll see that play out in our demonstration. So the first metal I'm going to put in here is a piece of sodium metal. And as a 1A metal we know that it wants to become Na plus. And we will place that in the water. And watch what it does. It is dancing around on a bed of hydrogen gas that's being generated. And I don't know if you can hear it sizzling, but it's sizzling. You also notice that it's turning a little bit of a pink color there, and that is a result of the sodium hydroxide that is being produced. Now let's do potassium. Potassium is further down the periodic table. Its valence electrons are farther away from the nucleus. And it is easier to take an electron, it takes less energy to remove an electron. Therefore it's more reactive. And we're going to see that play out in this reaction. Notice that the hydrogen gas that's being produced is catching on fire. We have a flame occurring, and it reacts more rapidly. [SOUND] It's done. [LAUGH] So we see the two reactions. They're both reacting with the water. We see them happening easily, but one is more reactive. The potassium is more reactive than the sodium. Okay, so in the demonstration we saw that between sodium and potassium, potassium burst into flames when it reacted with the water. Sodium reacted fine, but potassium burst into flames. It's more reactive with it. And the reason for all of that is that the ionization energy is going to decrease down the family as we turn potassium into potassium plus. It's really easy to do that. It's generating a lot of hydrogen gas that, and a lot of exothermic process overall, and it bursts into flames. The other thing that we can know, that we ought to take note of with these 1A metals, is that they not only tend to form oxides, they form peroxides and what are called superoxides. So what's a superoxide? A superoxide is got two oxygens held together with a minus one overall charge, okay? And so if you saw something like this, KO2, you would know you're dealing with a superoxide. Now how would you know that? Well, potassium has to be plus one. That's its only charge available to potassium. So if potassium is plus one, then all of this has to be minus one, okay? So we have a potassium ion and an O2 minus ion. And that's called a superoxide. If you saw something written like this, you might say, well, why would they write it like that? Why not the lowest whole number ratio between them and write it just KO? Well, the reason they do that is because this is a group of two atoms that together have a minus 2 charge. That's a peroxide. So then you'd have to have 2 pluses to go with it. Okay, let's go to the next group. We've got our elements that are in the alkaline earth metals. So they're all metals, and if we look at their electron configuration, we know it's ns2. Okay so we know that beryllium was helium 2s2, for example. As we go down it'd be 3s2 and 4s2 and so forth. Now these are not going to be quite as reactive, their ionization energy is not as low, and you not only have to remove one electron, but you have to remove both electrons. So you're not going to have them quite as reactive. Because, not only do you remove one, but you have to remove two. And, we know that the second ionization energy is higher than the first ionization energy. You have to remove the two, but we will see the same trend. It's going to become more reactive as you move down a family. Now these, some of them will react with water. Beryllium won't. Okay, it won't react with water. But magnesium would react as long as it's hot enough. So it would react with steam. So we would see that beryllium is not very similar to the other guys in its group, is it? Because these guys will react. Calcium will react, and so will any of the elements below it will react. And it would do it at room temperature water. So you would end up being, forming calcium hydroxide when you dissolved it in water. So that same trend is harder. It's not as reactive, so it's not as, it would be no fun to watch a demonstration of these guys reacting with water. But we can see that they are reactive, not as reactive as the 1A elements. We see the same trend of reactivity decreasing as you go down, and this is all driven by ionization energy. [SOUND] If you want a metal to react, it's going to want to form a cation. To form a cation, you have to remove an electrons. It's going to be more reactive if it's easier to remove those electrons. Okay lets move to the group 3A elements, just look in the periodic table, and we could do aluminum for example. All of these will have this basic electron configuration, so we have helium 2s2, I said aluminum, I should've said boron, 2s2, 2p1. And so these tend to form tend to lose electrons. If we do aluminum, It's neon 3s2 3p1. And aluminum. The only ion it will have is a three plus charge. This is not going to be true for all the other metals, but that is the only ion it can have. It is going to remove all of them in order to get a plus three charge. Now as we move down that periodic table, that is not going to be true for the other metals. For the other metals, they're going to able to form plus one ions as well as plus three ions. Let's think about why that would be. To do this, it's best to think about, whenever there's breaks in what you expect, think about the electron configuration. So let's have you review those skills. What's the electron configuration of gallium? Well, if you said it was number 3, you are correct. Look at that electron configuration. We've got 4s2, 3d10, 4p1. So these are my valence electrons. And they're being separated by a bunch of d's, okay? So if you think about gallium, why would you expect it to form plus ones and plus threes? Well, let's write down that configuration we came up with, 4s2, 3d10, 4p1, okay. If it only lost that p, that'd be okay, it still has a completely filled edge, yes, and all those d's in there, and so we'd like to do that, but then you could also remove these as well, and that would be three electrons. Remember, to always remove the highest end when you're removing electrons to form cations, so we won't touch those d's, they're an inner shell. But when you have, those d's separating the s's from the p's, then you can expect two possible charges for those metals. Let's move onto the 4A elements, okay? So they're going to have that electron configuration. We're going to have four valence electrons and we have see, we see in this family, this is the first family where we see non-metals up there at carbon. We come down to and we see some metalloids with silicon. We go further down and we have metals. We have metals such as tin and lead in there. With tin and lead, we've got these electron configurations, and you might want to try writing their electron configurations out before you proceed. But what, once, once you write those, what positive charges would you expect to see for tin and for lead? Well let's see what you come up with. Write the electron configuration and choose your answer. Well if you choose number three, you would be correct. Because when you're doing the electron configuration you're going to see that both for tin and for lead, you're going to have two electrons in an s. That's going to be followed by some d's. And there'll be ten of them. And then we're going to have a couple p's. Oh, I put a 2 there, that should be an n. Okay? So this number's going to be different for tin versus lead, but everything else is going to be the same. So if it lost these two only, that'd give you a plus two charge. If it lost these and these, that would give you the plus four charge. So those are common ions for the metals. So we start up at carbon at the top. Carbon doesn't like to become a cation or an anion. It has four electrons, it's very small, doesn't like to gain or lose in an ion sense, it tends to form covalent bonds with other non-metals. We'll talk about that much later. But these metals down below like to become cations and you will see plus two and plus four charges. Moving on over to the next group, we're at the 5A elements, okay? These are going to have a half filled s, I mean p subshell. So we've got our nitrogen up there that we had looked at an electron configuration not too long ago. Again, this group we see non-metals at the top. We have an area of metalloids in the middle and then, I call them metalloids, and then we have some metals down towards the bottom of the periodic table. The one unique element of this group, that's worth noting, we'll mention nitrogen. One thing about nitrogen is it takes on a whole variety of oxides. It can have one oxygen, a two-to-one ratio, a one-to-two ratio. This is very unique for this. I would not expect you to memorize that. But you're going to run across many different oxides of nitrogen. And later on in this course you're going to learn about something called oxidation numbers, and you're going to be able to determine those oxidation numbers and see how they change for these various elements. Moving on over to the 6A elements, okay. All of these will have six valance electrons, and so they'd be two in the s and four in the p's. These, on the anion end of things, if they are a non-metal, they tend to form minus two ions. And you just have a very large variety of compounds formed from these elements. They play out in a whole gamut. Biological systems use it al, use oxygen and sulfur a lot. And, so we won't get into specific details because there's too many to mention. Our last group that we want to look at here are the halogens. These halogens are in the 7A elements. We know that they all have this number, two in the s and five in the p. We learned that these are, well, if you look over there at your periodic table, they're all non-metals. They all would tend to want to gain one electron if they were to become an ion because they could get that plus I mean that minus one ele, with a minus one charge they would become like the noble gases. These are your diatomic elements. They all exist as diatomic. So, F2 Cl2 Br2 I2 in their elemental form, they are all diatomic. And this plays out because their electron configuration, but we will reserve that for later when you talk about covalent compounds in a later unit. Now all of these really, really would benefit from receiving one more electron and getting that noble gas core, that noble gas electron configuration. So these elements are highly reactive. Fluorine, you'll never find in nature as F2. Chlorine, any of these, they, they will react easily with things around them, they want to gain an electron to give it that noble gas electron configuration. So the high reactivity plays off the fact that they have a large positive electron affinity. They would like that electron so much that it will provide energy, an ex, an export of energy that can be used for something else. They have a large, positive electron affinity, making them very reactive. Now, that reactivity increases as you go up the family. And so your fluorides, or your fluorine elements are the highest reactivity of all the elements. Now, this discussion we really ought to at least mention, it's not in terms of chemical reactivity, we ought to at least mention the noble gases. We know that they have finally got that completely filled s-n-p subshell. We know they're not really considered representative elements. We can call them main group elements. They all exist very happily as independent atoms, monoatomic atoms. And this is because it's got a very high ionization energy. It would be very, very hard to remove an electron. They're very, very stable and all of this is completely because of their electron configuration. There are no compounds formed between helium, neon, argon and any other element. You cannot, we have not at this point, and do not believe we will make compounds with those elements. However you can easily make compounds with xenon, not easily, I shouldn't say easily. You can make compounds with xenon and krypton with the most highly reactive non-metals. Fluorine, oxygen can form bonds with xenon and krypton. So for reasons we won't get into at this point, but maybe later in this unit, see why and how fluorine and oxygen can react with these noble gases. So this ends our last learning objective of this unit. We have seen how electron configuration, ionization energy, electron affinity plays out in the reactions of our different groups within the representative elements. This ends the entire unit for us and you'll now have an opportunity to do some practice problems and take your assessment. Good luck with that.
