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Ionic, covalent, and metallic bonds | Chemistry of life | Biology | Khan Academy

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Everything we've been dealing with so far has just been with the individual atoms, but atoms bond. Or another way of saying it is, they stick together. Because if atoms didn't stick together, then we'd all be essentially just a collection of atoms and this video wouldn't be being produced. So atoms stick together and they form molecules. You take a bunch of atoms together and they'll stick together. And they'll form molecules. And then obviously molecules start building up and you get other structures. And if we started talking about organic chemistry, you'd have a bunch of atoms, a lot of carbons and hydrogens and other things, fitting together and they'd be forming proteins. And then proteins would fit together to form organic structures. And you fit enough of those together, and you'll eventually get someone recording a YouTube video. So this is where it all starts. Atoms bond. Or they stick together. And the purpose of this video is to think about the different types of ways that an atom can stick to another atom. So the first, and kind of the most powerful way-- or I think of it as the most powerful way-- is if you take an atom that really wants to give an electron, and then you have another atom that really wants to take an electron. Right? And we've talked about this before. An atom that wants to give an electron wants to give it because it's trying to get into a stable configuration in its outer shell. Everyone wants to look like a noble gas. They're all envious of the noble gases, because the noble gases have eight electrons in their outer shell. So who wants to give? Well if you look at the period table, the people who want to give really badly-- and we've talked about this a good bit-- are the alkali metals. These guys just really want to offload an electron. And there are other people who want to give them. But let's take the extreme example. So these guys really want to offload an electron. And who wants to take an electron? Well, the halogens. We've talked about it. These guys love taking electrons. They're not the only ones. But they have a very high electronegativity. They really want to take electrons. So if you put these around each other, what happens? Let's say, sodium and chlorine. And let's say we wanted to flavor some of our food. So you have some sodium and you have some chlorine. So sodium-- let me draw its valence shell-- sodium's valence shell looks like this. It's got one electron sitting there that it would really just like to get rid of. And then chlorine looks like this. It has seven valence electrons. One, two, three, four, five, six, seven. So what happens is this guy wants to escape. This little blue electron right here really wants to escape the sodium and essentially move into the chlorine. And obviously, it's not like a one-for-one. You'd have billions and trillions of these atoms rolling around, and these electrons jump off, then they go to one, then they jump to another. But for the sake of our purposes, let's say we just have these two atoms. And what you have is that that electron jumps off. And then if that electron jumps off, what happens to sodium? Well then the sodium has no electrons in its valence shell. Although it does. Now its valence shell is one lower, but we can say it's lost that one electron that was out there. And now its atomic configuration will look a lot like neon. Right? Sodium, you lose an electron, now it looks a lot like neon, at least its electron configuration. But now it has one fewer electrons than protons. So now, it has a positive charge. It was neutral back here. Right? Now it's positive. And now, what does a chlorine look like? And I'm kind of mixing up notations, but that's really just to give you the idea. So chlorine before had seven electrons. One, two, three, four, five, six, seven. That electron had jumped onto the chlorine. So now it's happy. It looks a lot like argon now. It has a completely filled valence shell. And what's the charge now? Well in has one more. Now it will have 18 electrons instead of 17. Right? So what is its charge now? It has 17 protons, 18 electrons. It has a negative one charge. So I'll just put a negative up there. It has a negative charge now because it got that electron from sodium. So now these guys are both happy from an electron configuration point of view. They both have these stable valence shells. But they're attracted to each other, right? Coulomb forces. Positive is attracted to negative, negative is attracted to positive. And it can be very strong, this electrostatic force, so they stick to each other. And so this force of attraction, this is an ionic bond. So they essentially will form NaCl. They're not sharing electrons. This guy wanted the electrons so badly, and this guy wanted to give them away so badly, he just handed the electron over. But then he says, oh, by the way, now that I handed you the electron, you're negative, I'm positive. I want to stick to you. And then we formed table salt and we're ready to season our food. Now that's the situation where one guy really wants to offload an electron, one guy really wants to take it. What happens in the situation where they're both not as extreme in their views in whether or not they want to give or take electrons? So let's think of a few other examples. The best example is elemental oxygen. Right? Let's see, elemental oxygen. So this right here is an ionic bond. Not to jump back and forth, but I'm not sure if I just mentioned that. Why is it called an ionic bond? Because we formed ions. When we donated the electron from sodium to chlorine, we formed an ion. The sodium, this became a cation, because it's positive. And this became an anion, because it's negative. And then they stuck to each other, so this is an ionic bond. Fair enough. Now what happens, like I was just starting to say, if we have two elements that aren't that different in how much they want electrons. Their electronegativity is very similar. And the best example of that is we had two of the same element. So let's say I had oxygen. Let's have one oxygen there. Let's look at the periodic table to make sure that we're not-- oxygen has six valence electrons, right? One, two, three, four, five, six valence electrons. Right? It's 2s2, 2p4. So on the second shell it has six electrons. So oxygen has one, two, three, four, five, six. And then let's say we have another oxygen. It has one, two, three, four, five, six electrons. Now both of these oxygen atoms would love to have eight electrons. They'd be stable. They could start pretending like they're a noble gas. But clearly, they don't have eight electrons. And let's say in this, all they have around each other is other oxygen atoms. So what they can do is say, this oxygen goes to that oxygen, and says, hey, why don't we share some electrons and then we can both pretend that we have eight electrons. And this guy says, oh, sure enough. So we can just bring him over here. And I'll just write him in blue. Oxygen doesn't necessarily have to change colors. I'm joking. So I'm just going to draw this guy over on this side just so you recognize that this is different than this guy. And then they share these electrons. So they share these electrons. And we could do it by drawing a line here. So they're sharing two pairs of electrons. So this guy right here, he had six electrons, but he can kind of pretend that he has this electron and that electron. So he has eight in his valence shell. And this guy, he's going to do the same thing. He has one, two, three, four, five, six, but he also can kind of pretend that these guys are also in his valence shell. So he's happy. And this notion, where you're actually sharing electrons, where these electrons are going to go and both electron probability distribution clouds of both atoms. This is called a covalent bond. And this is typical when you're dealing with two elements that aren't very different in terms of their electronegativity or their desire to attract electrons. Now, when we talked about ionization energy, I think, we talked about when oxygen and water bond, right? And oxygen-- oxygen, we've drawn that-- is six. Not oxygen. Water. Oxygen and hydrogen to form water, and hydrogen looks something like this. Right? You have a hydrogen atom there. You have a hydrogen atom there. They said, hey, why don't we get together. Let's share some atoms. And the hydrogen atoms say, oh, OK, let's share some atoms. Let me rewrite this oxygen like this, so it becomes clear that we're sharing. So if I rewrite this oxygen like this. I essentially split up one of these pairs. And these hydrogens come along and they share one hydrogen there, one hydrogen there. This guy can pretend like he has his first shell filled, because you can only put two there. That's where the eight rule breaks down in the first shell. This guy can pretend, too. And now oxygen can pretend like he's got eight electrons in his valence shell. And everyone's happy. So this is also a covalent bond. Another way we could have written this, and I think I did this in the last video, I could have written it like this. Where the implication of this line, each of these lines involve two electrons. These are equivalent statements. But in this situation, oxygen is more electronegative than hydrogen. It wants to get the electrons more than hydrogen. So in this situation, the electrons are going to spend more time around oxygen than they will around hydrogen. So hydrogen will experience, I guess you could call it, a partial positive charge on this side of the molecule, while the oxygen side will experience a partial negative. I'm going to draw it real small, because it's a partial negative. This is called a polar covalent bond. Because it's still covalent. We're sharing electrons. But it's polar, because the electrons are getting pulled to spend most of their time at one side of the atom. And since that is the case, the molecule as a whole, the collection of atoms, is going to have polarity. One side of the molecule is going to be more negative than the other side, which will be more positive because the electrons are spending more time on that side. Now the last bond we can talk about, and I've touched on this a little bit, is the metallic bond. I was in a metallic bond in high school, but anyway, that's a subject for another video. But with metals, you can't really draw the electron structure there. But what happens with, let's say we have iron, right? And you have just a bunch of neutral iron atoms sitting around. And we established the one commonality of metals, what makes something metallic or have metallic characteristics is that they have a bunch of electrons in their outer orbital that they're very giving. They're very happy to share. So if you put a bunch of these guys together, what happens is they share their electrons. So they all become positive. They're very communal this way. The metallic atoms. And then their electrons kind of just form this sea out here. But they all share. E minus. E minus. E minus. E minus. And because their electrons are all on the sea and they've kind of gotten this positive charge, they're attracted to the sea that they've created. They're attracted to their shared electron pool that all of the atoms have donated to. And this is essentially what allows, well definitely, metals to be conductive, because you have this pool of electrons that are very easy to move around. And also it's what makes them malleable. Because even if you have visually-- it's a little intuitive. There's nothing exact here. But you can kind of move these. You can imagine that this is kind of a big pudding of electrons or big glue of electrons. And you can move, you can bend the rod or flatten the rod without having it break or get brittle. While if you're talking about salts that have a very strong but rigid bond, if you were to try to bend a bar of salt, the bond will just be broken. There's no, kind of, squishy electron mush that you can kind of bend around and play with. Anyway, so those are the three bonds. And hopefully that gives you a little intuition. And this is super useful, because in the rest of chemistry, everything we do will essentially involve some combination of these bonds. And we'll start talking about what these bonds mean in terms of the temperature at which they boil, or the properties of the molecules themselves. Anyway, see you in the next video.

Video Details

Duration: 13 minutes and 22 seconds
Language: English
Genre: None
Views: 3
Posted by: jake.goldwasser on Nov 19, 2015

Introduction to ionic, covalent, polar covalent and metallic bonds.
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