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Electronegativity | Chemistry of life | Biology | Khan Academy

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Voiceover: What I want to talk about in this video are the notions of Electronegativity, electro, negati, negativity, and a closely, and a closely related idea of Electron Affinity, electron affinity. And they're so closely related that in general, if something has a high electronegativity, they have a high electron affinity, but what does this mean? Well, electron affinity is how much does that atom attract electrons, how much does it like electrons? Does it want, does it maybe want more electrons? Electronegativity is a little bit more specific. It's when that atom is part of a covalent bond, when it is sharing electrons with another atom, how likely is it or how badly does it want to hog the electrons in that covalent bond? Now what do I mean by hogging electrons? So let me make, let me write this down. So how badly wants to hog, and this is an informal definition clearly, hog electrons, keep the electrons, to spend more of their time closer to them then to the other party in the covalent bond. And this is how, how much they like electrons, or how much affinity they have towards electrons. So how much they want electrons. And you can see that these are very, these are very related notions. This is within the context of a covalent bond, how much electron affinity is there? Well this, you can think of it as a slightly broader notion, but these two trends go absolutely in line with each other. And to think about, to just think about electronegativity makes it a little bit more tangible. Let's think about one of the most famous sets of covalent bonds, and that's what you see in a water molecule. Water, as you probably know, is H two O, you have an oxygen atom, and you have two hydrogens. Each of the hydrogen's have one valence electron, and the oxygen has, we see here, at it's outermost shell, it has one, two, three, four, five, six valence electrons. One, two, three, four, five, six valence electrons. And so you can imagine, hydrogen would be happy if it was able to somehow pretend like it had another electron then it would have an electron configuration a stable, first shell that only requires two electrons, the rest of them require eight, hydrogen would feel, hey I'm stable like helium if it could get another electron. And oxygen would feel, hey I'm stable like neon if I could get two more electrons. And so what happens is they share each other's electrons. This, this electron can be shared in conjunction with this electron for this hydrogen. So that hydrogen can kind of feel like it's using both and it gets more stable, it stabilizes the outer shell, or it stabilizes the hydrogen. And likewise, that electron could be, can be shared with the hydrogen, and the hydrogen can kind of feel more like helium. And then this oxygen can feel like it's a quid pro quo, it's getting something in exchange for something else. It's getting the electron, an electron, it's sharing an electron from each of these hydrogens, and so it can feel like it's, that it stabilizes it, similar to a, similar to a neon. But when you have these covalent bonds, only in the case where they are equally electronegative would you have a case where maybe they're sharing, and even there what happens in the rest of the molecule might matter, but when you have something like this, where you have oxygen and hydrogen, they don't have the same electronegativity. Oxygen likes to hog electrons more than hydrogen does. And so these electrons are not gonna spend an even amount of time. Here I did it kind of just drawing these, you know, these valence electrons as these dots. But as we know, the electrons are in this kind of blur around, around the, around the actual nuclei, around the atoms that make up the atoms. And so, in this type of a covalent bond, the electrons, the two electrons that this bond represents, are going to spend more time around the oxygen then they are going to spend around the hydrogen. And these, these two electrons are gonna spend more time around the oxygen, then are going to spend around the hydrogen. And we know that because oxygen is more electronegative, and we'll talk about the trends in a second. This is a really important idea in chemistry, and especially later on as you study organic chemistry. Because, because we know that oxygen is more electronegative, and the electrons spend more time around oxygen then around hydrogen, it creates a partial negative charge on this side, and partial positive charges on this side right over here, which is why water has many of the properties that it does, and we go into much more in depth in that in other videos. And also when you study organic chemistry, a lot of the likely reactions that are going to happen can be predicted, or a lot of the likely molecules that form can be predicted based on elecronegativity. And especially when you start going into oxidation numbers and things like that, electronegativity will tell you a lot. So now that we know what electronegativity is, let's think a little bit about what is, as we go through, as we start, as we go through, as we go through a period, as say as we start in group one, and we go to group, and as we go all the way all the way to, let's say the halogens, all the way up to the yellow column right over here, what do you think is going to be the trend for electronegativity? And once again, one way to think about it is to think about the extremes. Think about sodium, and think about chlorine, and I encourage you to pause the video and think about that. Assuming you've had a go at it, and it's in some ways the same idea, or it's a similar idea as ionization energy. Something like sodium has only one electron in it's outer most shell. It'd be hard for it to complete that shell, and so to get to a stable state it's much easier for it to give away that one electron that it has, so it can get to a stable configuration like neon. So this one really wants to give away an electron. And we saw in the video on ionization energy, that's why this has a low ionization energy, it doesn't take much energy, in a gaseous state, to remove an electron from sodium. But chlorine is the opposite. It's only one away from completing it's shell. The last thing it wants to do is give away electron, it wants an electron really, really, really, really badly so it can get to a configuration of argon, so it can complete it's third shell. So the logic here is that sodium wouldn't mind giving away an electron, while chlorine really would love an electron. So chlorine is more likely to hog electrons, while sodium is very unlikely to hog electrons. So this trend right here, when you go from the left to the right, your electronegativity, let me write this, your getting more electronegative. More electro, electronegative, as you, as you go to the right. Now what do you think the trend is going to be as you go down, as you go down in a group? What do you think the trend is going to be as you go down? Well I'll give you a hint. Think about, think about atomic radii, and given that, pause the video and think about what do you think the trend is? Are we gonna get more or less electronegative as we move down? So once again I'm assuming you've given a go at it, so as we know, from the video on atomic radii, our atom is getting larger, and larger, and larger, as we add more and more and more shells. And so cesium has one electron in it's outer most shell, in the sixth shell, while, say, lithium has one electron. Everything here, all the group one elements, have one electron in it's outer most shell, but that fifty fifth electron, that one electron in the outer most shell in cesium, is a lot further away then the outer most electron in lithium or in hydrogen. And so because of that, it's, well one, there's more interference between that electron and the nucleus from all the other electrons in between them, and also it's just further away, so it's easier to kind of grab it off. So cesium is very likely to give up, it's very likely to give up electrons. It's much more likely to give up electrons than hydrogen. So, as you go down a given group, you're becoming less, less electronegative, electronegative. So what, what are, based on this, what are going to be the most electronegative of all the atoms? Well they're going to be the ones that are in the top and the right of the periodic table, they're going to be these right over here. These are going to be the most electronegative, Sometimes we don't think as much about the noble gases because they aren't, they aren't really that reactive, they don't even form covalent bond, because they're just happy. While these characters up here, they sometimes will form covalent bonds, and when they do, they really like to hog those electrons. Now what are the least electronegative, sometimes called very electropositive? Well these things down here in the bottom left. These, over here, they have only, you know in the case of cesium, they have one electron to give away that would take them to a stable state like, like xenon, or in the case of these group two elements they might have to give away two, but it's much easier to give away two then to gain a whole bunch of them. And they're big, they're big atoms. So those outer most electrons are getting less attracted to the positive nucleus. So the trend in the periodic table as you go from the bottom left, to the top right, you're getting more, more electro, electronegative.

Video Details

Duration: 9 minutes and 54 seconds
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Language: English
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Views: 10
Posted by: jake.goldwasser on Nov 19, 2015

Defines electronegativity and compares to electron affinity. Explains group and period trends in electronegativity using atomic radii.
Watch the next lesson: https://www.khanacademy.org/science/biology/chemistry--of-life/chemical-bonds-and-reactions/v/electronegativity-and-chemical-bonds?utm_source=YT&utm_medium=Desc&utm_campaign=biology
Missed the previous lesson? https://www.khanacademy.org/science/biology/chemistry--of-life/chemical-bonds-and-reactions/v/ionic-covalent-and-metallic-bonds?utm_source=YT&utm_medium=Desc&utm_campaign=biology
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