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Chemistry - Atoms, Elements and the Periodic Table
So far, we've learned a little bit about determining electron configurations. Let's see if we can use that information to group elements on the periodic table and then guess as to what they might do when they react with other elements. So let's just figure out the electron configurations of a couple of elements just for a little bit of practice. So, lithium, right there. What does it look like? Lithium's electron configuration. You get the first shell, is 1s2. Two electrons there. And then you have 2s1. And sometimes, just to be quick, to get the notation, is you can imagine lithium's electron configuration is the exact same thing as helium's electron configuration-- this is helium's electron configuration-- plus the 2s1. This could have also been written as-- do that in light blue-- could also have been written as helium, 2s1. Which essentially means that lithium's electron configuration is exactly what you would have written for helium's electron configuration, and then you'd have written 2s1. You could do that a bunch of times. Let's say if we wanted to figure out the electron configuration of iron. Instead of going through the whole thing, you know, it's 1s2, and then it's 2s2, and 2p6. Instead of doing that whole thing, you could just say, OK, iron has the same electron configuration. So you could say iron's electron configuration is the same thing as argon's electron configuration. So I'll just put argon in brackets. And then you get 4s2. And then you have one, two, three, four, five, six. So d6. And we learned that when you're in the d subshell, or when you're in the d-block of the periodic table, you are actually backfilling the previous shell. So when we're in the fourth period, in the d-block, we're backfilling the third shell. So 3d6. And someone had asked-- and this is an interesting question-- why does it do that? Why does it not just continue? Why doesn't it fill the fourth d shell? And the way I think about it-- and this is all intuition, and things at the atomic level really start to become, on some levels, non-intuitive-- but the way I think about is as the atom grows larger and larger, there are more spaces between the previous orbitals. For example, this is just how I visualize it. If my first shell looks like this. Let's say the s looks like this. And then, if I just cut it out, let's say the p's look like something like this. This is maybe the second shell. The p's look like this. And then the next place an electron might want to be might be in the third shell, right? So the third shell would be like this. And then you fill out the third p shell. This is just an intuition. This isn't exactly what an electron would look like. Maybe the third p shell would look something like that. Look something like that. And then look something like that. And then you're in the fourth shell. So you're doing the fourth shell. The s subshell might look something like that. And then instead of immediately starting the next p shell, you're in the d-block now. So this is-- let me just write some labels-- 4s. This is 3s. This is 3p. This is 2p. This is 2s. And then 1s is inside of 2s. So you don't have to worry about that too much. But my intuition behind why the d orbital gets backfilled is because now, as the atom gets larger and larger, you have these spaces in between the previous orbital. So now, after filling the 4s subshell, or the 4s orbital-- so this is 4s here-- out here, we go back and we fill in the 3d orbital. So we're going back and we're filling these spaces right here. So this is a lower energy state than this. It takes more energy to cram an electron back into the 3d shell, back there. But then once you do that, now you're ready to then go to the 4p shell, which might look something like this. So an electron would rather go to another shell, which is the fourth shell, rather than backfill the 3d shells. But once it fills out the fourth shell, it fills in those spaces in between. And as the electron gets bigger and bigger, there's more and more spaces in between. So eventually, when the electron gets big enough, there's going to be spaces between the d shells, and that's where the d orbitals and that's where the f orbitals will go. That's my intuition behind its working. And obviously, when we're dealing at the atomic scale, as far as I'm concerned, that's the best that I can do. But fair enough. That's not what I want to do here, but that was a good question, as to why do you go and backfill the third shell when we're in the fourth period? Fair enough. This is an easy way to write iron's electron configuration. The reason why I'm doing all of this is to figure out how many electrons you have in the outermost shell. In the case of lithium, you have one electron in your outermost shell, right? This is your outermost shell right here. You have one electron. And you could have done the same thing right there. In the case of iron, how many electrons in the outermost shell? Remember, the outermost shell is the period you're in. And this is the outermost shell. So even though these are higher energy electrons-- it took more energy to backfill those into the lower energy shell-- it's these that are on the outside energy shell, the fourth shell, that are going to be the ones that are reacting. And how many are there? There are two. And this is an important thing. So there's two here. There's two on the outside shell here. And actually, there's going to be two for any of these in pink right here. Any of the ones in the d-block, what happens? You fill whatever period you're in. Let's say that you're in period five here. Right? You're going to have 5s1. 5s2. And then you're going to go back and you're going to fill the 4d shell. Right? But in terms of how many electrons you have on the outside shell, in this case the fifth shell, you are going to have two electrons. So all of these are going to have two electrons in their outermost shell. In the case of these, the outermost electrons are going to be 4s2, right? Because then you go back and fill the 3d, but the outer ones are 4s2. So this one also has two electrons in its outermost shell. How many does this group have? And I've just used a word that I don't know if I've defined before, but the group are the columns in the periodic table. And as you can see, they all have patterns to them. Everything in this first group has one electron in its outermost shell. If you don't believe me, look at hydrogen. Hydrogen's electron configuration is 1s1. Its outermost shell is 1s. It has one electron there. Right? And that's true for all of these. All of these guys have two electrons in their outermost shell. These guys have those same two electrons. We can view it that way, in their outermost shell, but then they go and backfill the d shell. But in terms of their outermost shell, only two electrons. Than once you fill the d-block, or you go backfill, in the case of the fourth period, you go and backfill the third d sub-orbital. Then you go back to filling the fourth shell again. Now the p block, right? So this one's going to have three electrons in its outside orbital. Or you could say three valence electrons. This is four, five, six, seven, and eight. Let me do one more, just in case you don't believe me. What's the electron configuration for Sn. This is, what, selenium? [tin] I'm not even sure. But let's say Sn. What's the electron configuration? It's going to have the same electron configuration as krypton. Yes, that element is krypton. There is such an element. So it will have the same electron configuration as krypton. So I could have figured out krypton's electron configuration just by going through the whole periodic table, but this is just a faster way of doing it. Same thing as krypton, and then it has 5s2. Then it goes back and backfills the d-block. So then there's 10 there. So 4d10. And then it starts filling up the p-block in the fifth shell again. So 5p2. So how many valence electrons does it have? Valence electrons, or electrons in the outermost shell? Well, what's the outermost shell? It's the fifth shell. So these and these. These electrons have a higher energy state than that. It took a little bit more energy to cram them back into that previous shell than it took to put these on the s orbital. But if you talk about the electrons that will react, and that's why I'm emphasizing these, these are the electrons that are going to react with other atoms. Or sometimes with just other electrons, even. This one has four outside electrons. And you see that right there. Four outside electrons. And since the outside electrons, for the most part, are the ones that you're going to care about, there's a-- I guess you could say, a notation where you only draw the outermost electrons. So, let's say, for hydrogen, you could write it like this. Where you're only drawing the outermost, valence electrons. Valence electrons are just the outermost electrons. You could write it like that. You could write it like that. But this says, hey, I just have one outside electron for hydrogen. If I wanted to draw it for iron? Iron, right here? How would I do that? I have two electrons in my outermost shell, so iron I could just do like this. And electrons, they tend to be paired. So if I have, let's say I wanted to take the example of, if this is Sn, this is selenium. Let me do carbon. Carbon, I have four electrons in my outermost shell. So carbon I could write like this. Or if I didn't want to pair them, in theory I could write them like that as well. And now they're ready to react with other things. Now what does this tell me about, you know, this one has one electron in its outermost shell. These blue, these noble gases-- and we'll talk a little bit about them in a second-- have eight electrons in the outermost shell. How does that help me when I'm actually trying to figure out how things react? Well, it turns out that all atoms want to have eight electrons in their outermost shell. And that number is important. Eight. They want to have eight electrons in their outermost shell. This is the most stable configuration for atoms. Or I guess you could say, to some degree, a better energy state for the atom. And why is it the number eight? Well, that's something to think about. This is another fundamental number that just pops out of nature. And I've thought a little bit about it. It must be something about the atoms in the outermost shell, when you have eight, they resonate well with each other. And they somehow don't get in the way of each other. Or don't want to push away from each other. I don't know the answer to that. And frankly, if someone could really answer the question of why eight, exactly why eight, they would make a good career for themselves in physics or chemistry. But through experimentation, it has been well established that atoms want to have eight electrons in their outermost shell. So the question is, if you're dealing with something like, let's say you're dealing with potassium. Right? Potassium has one electron in its outermost shell. Let's say you have stuff like chlorine, that has seven electrons in its outermost shell. What do you think's going to happen if you put some potassium near some chlorine? What's going to happen? Well, what's the easiest way for the chlorine to get eight electrons? Well it has seven in its outermost shell. What's the easiest way? Well, it'll want to gain an electron really, really badly. And what's the easiest way for potassium to have eight electrons in its outermost shell? Well, if it lost that one electron, then it will have eight electrons in its outermost shell, right? Its outermost shell won't be the fourth shell anymore. It'll be the third shell. But it'll have eight electrons in the third shell. Its configuration will then look like argon if it loses that one electron. So it'll be a more stable state. So if you put sodium in the presence of chlorine, what's going to happen? This electron wants to jump off of sodium real bad so that sodium can have eight electrons in its outermost shell, or have an electron configuration like argon. And that electron is going to jump to chlorine, and then chlorine will have eight electrons in its outermost shell, and also have an electron configuration like argon. And so, as you can imagine, this group right here, which are called the alkali metals. And we'll talk probably in the next video why they're called metals. This group here, alkali metals. And they tend to exclude hydrogen, and we'll talk about that. These really want to give away electrons. And because of that, they're highly, highly reactive, especially if you put them in the presence of these elements, these yellow elements right here, which are called the halogens. These really, really want to take electrons from other things, because they just need one to get to eight. They usually want to give away electrons, because they just have to give away one to get to eight. And the reason why hydrogen, actually, isn't included is because hydrogen doesn't want to give away its electron as bad as these guys. This rule that your outermost shell wants to get to eight, that's true for everything except for hydrogen and helium. Hydrogen and helium, just because they have one shell, they're happy with just two electrons. And so with hydrogen, sure, it could lose an electron, but could just as easily gain an electron and be happy, because it'll have a full first shell. But all of these other ones, these alkali metals, they want to give away electrons really bad. When people in chemistry talk about metallic nature, they're really talking about how badly something wants to give away electrons. Anyway, I'm all out of time now. In the next video, we'll continue discussing the groups in the periodic tables and any trends we can ascertain from them.