Understanding Electronegativity and Polarity

Hello, this is a lecture topic video discussing electronegativity and polarity. And what I'm going to give you is just a very brief overview of these topics because what we want to start to do is think about chemical reactions happening inside of an aqueous environment. In other words, happening in water. And as you might expect, there are many, many, many different reactions that work inside of an environment that is chocked full of water. And so what we want to do is build up this idea of why water is such a good solvent. Okay. So that's a term we're going to talk about in the next video. Okay. And in order for us to do that, we need to discuss electronegativity and how electronegativity starts to lead to this idea of polarity. All right. So we've talked about two different types of bonding so far. Okay. We've talked about ionic bonding and covalent bonding. So let's talk about ionic bonding first. So if we think about Na. and Cl, okay? So sodium and chloride. So we have a metal and a nonmetal. What we like to think about, all right, is that that sodium atom can essentially lose an electron and adopt a plus one charge, okay? So again, remember, our metals love to lose electrons, okay? And then what we can think about is that that chlorine, because that electron just doesn't dissipate, appear, all right, we can think about that chlorine wanting to adopt that electron. And when it does so, it adopts a negative charge. Okay. So we end up with, at the end of the day, we get a sodium cation and a chlorine anion. Okay. So we've essentially transferred an electron from the sodium to the chloride. And then what can happen? Well, our sodium and come together and interact with our chloride in an ionic bond, okay? An ionic bond is nothing more than a charge-charge interaction, okay? And so we can get an ionic compound, all right? So it's a charge-charge interaction. And remember, the strength of this interaction, this is a very, very strong interactions, okay? They are dictated by Coulomb's law, all right? So in this case, I've talked about monatomic cation interacting with a monatomic anion. But remember, we can also have our polyatomic ions come into play here too. We got a couple of polyatomic cations and many, many, many polyatomic anions. They can also take part in ionic bonding. All right. Then the other type of bonding is our covalent bonds where we're sharing electrons. And we're sharing electrons because what are we doing? We're getting some amount of orbital overlap. All right. So remember hydrogen. Hydrogen has a single proton. So I'm just going to circle that inside of its nucleus. And then we have a single electron that is kind of zipping around outside of the nucleus that in an orbital, okay, we call them orbitals, and they're defined by some amount of probability. Okay, so probability defines where we might be able to find this electron. All right. And what we can start to do is think about bringing two hydrogens together to make a hydrogen-hydrogen bond. So again, we like to think about this in terms of energy, where we have energy on our y-axis here and we have distance on our x-axis. And remember, the further away these hydrogens are, the less they're interacting. But as we start to bring these guys closer together, our energy starts to drop until we get to this energy minimum. So this is where our energy... Well, will be. Okay. So it's the lowest energy of our system. Okay. And what is that? What defines that point? All right. It's the optimal distance between these two hydrogen atoms. All right. And why is that the optimal distance? Well, remember, it's like a balancing act. We have repulsive forces. We have attractive forces. At that optimal distance, what we've done is we've maximized our attractive forces. And we've minimized our repulsive forces. So let's think about what those are. So again, we get orbital overlap. That's why we get a bond in the first place, okay? Because that bond, that orbital overlap starts to focus, all right? Starts to localize those electrons between the two nuclei, between those two hydrogen nuclei, okay? And then now what we have is we have some attractive forces, but we also have repulsive forces, all right? So remember our attractive forces, right? Clearly, those protons wanting to interact with those electrons and electrons with the protons. And that works for both of our atoms. Okay. So those are our attractive forces that we're trying to maximize. Okay. But we also have some repulsive forces as well. So remember, we still have some probability that those electrons are not in between those two nuclei. All right. So we still have these orbitals. Okay. These electron clouds, right? And those electron clouds lead to some repulsion. And we also get some charge-charge repulsion from... our two protons okay but again at this minimum energy what we've done is we have maximized okay we've maximized the attractive forces okay and minimized the repulsive forces Okay. And what that does is it leads us to this energy minimum. So a lot of energy has been given off that again, something else can harness and use to do some sort of work. Okay. And we've decreased the energy of our system. And again, that is a favorable thing. Okay. But what we've done now is we've made a bond. All right. And so what do we do? We kind of represent that with this stick. Okay. So we've made a hydrogen hydrogen bond. All right. You can also think about now moving those two nuclei closer to each other. All right. And then what happens with our energy then? Okay. Well, as we move them closer to each other, of course, we start to maximize those repulsive forces. Right. And so, of course, what happens here is our energy just kind of skyrockets. Okay. So the energy skyrockets. And once we get above that zero energy mark, what do we have to do to make that happen? we have to add energy into that system. In other words, we have to add energy into that system to move those two nuclei closer to each other. Because again, what do we have to do? The energy is being supplied to overcome those repulsive forces. All right, so what we can also start to think about is doing the same sort of thing with fluorine. So fluorine has nine protons and nine electrons. In this case, fluorine is one period below. So... it has two different shells that are holding those electrons, okay? There's two in this inner shell. I'm just going to kind of highlight this guy. And we're going to talk a lot about this. There's two on this inner shell. And then the other seven are in this outer shell, okay? This outer shell, those electrons are higher energy than the ones within that inner shell, okay? And this outer shell is called... our valence shell, okay, our valence shell. And when we do bonding, all right, the electrons that take part in that bonding are in the valence shell, okay? So what we can, again, start to do is we can start to move two fluorines, okay, from being really far apart from one another to actually coming into contact with each other, all right? And when we find that optimal distance, what do we get? We get orbital overlap. we start to focus electron density between those two nuclei. That focus density allows us to make a bond, which again, we can represent as a stick, okay? And at this point, all right, this is the point where we've maximized our attractive forces, okay? Minimize those repulsive forces. And this is also that point where we're at the lowest energy. So energy has been released, okay? Because this is a favorable process, all right? But what I want you to notice is that, again, the way this picture is set up, it's showing you that it's those valence electrons, those electrons in that outer shell that are taking part in the bonding. And that's where the orbital overlap is also happening. Okay. So what we can start to do is think about these ionic bonds, right, where we can essentially think about this as being an electron transfer event to give us two charged species that can interact with each other through charged-charge interactions. Or we can start to… bring nuclei together, get some orbital overlap where we can start to localize electron density between those two nuclei, essentially sharing those electrons, and we can get covalent bonds that result from that. So essentially what I've done now is I've described to you two extremes to bonding. There's two extremes to bonding where at one extreme we have ionic and again ionic We're talking about full charges here interacting with each other. So again, if we're thinking about our sodium chloride, you can think of this M as being our sodium. and this X being our chloride, okay? So our sodium cation and our chloride anion, those guys coming together, making charge-charge interactions, and making very strong ionic bonds with one another, okay? So getting, which would lead us to these ionic solids that we've talked a lot about, okay? But on the other end, all right, we would get perfect sharing. We'd get covalent bonds, okay? Where those electrons are shared perfectly between those two nuclei that are bonded together. So we get covalent bonds, and more specifically, we call these nonpolar covalent bonds. So we get these nonpolar covalent bonds. Something like hydrogen bonding with hydrogen, or fluorine with fluorine. There's an example down here, right there. Chlorine with chlorine, so on and so forth. Perfect sharing. So these are defining the spectrums, the two extremes to bonding. And what we can start to do is to represent these things with something we call electrostatic potential maps. And that's what these images are. These show us where the electrons are within. these molecules. Okay. So you see a lot of colors. If you think about our rainbow, we have Roy G. Biff. So our colors, red, orange, yellow, green, blue, indigo, violet. Okay. Well, when we're down here and the blues and the violets, those indigo colors, this is trying to tell us about the positive character of a molecule. Whereas these more aggressive colors, the red, oranges, and yellows, those are where we have negative character. Well, how are we getting these builds up of negative and positive character? Well, again, remember, electrons carry negative charges. So when you get a buildup of electron density, you start to get a buildup of negative character, negative charge. All right. Versus those blue colors. Well, what happens if you take away some negative character, you take away that electron density, there's less electron density to help neutralize that positive character coming from the protons. So you start to get a buildup of positive charge. Okay. So what you can see here is let's take a look at this electrostatic potential map for sodium chloride. That example we've been talking about. Well, what you can see here, there's one particular color that's focused around that sodium. So let's just get that sodium here. And then there's another color focused around those chlorides. Well, I just told you that those blue indigo violet types of colors, that's telling us that there's positive character. Whereas those red, orange, yellow colors, those more aggressive colors, that's where we have a buildup of negative charge. And so that makes perfect sense, right? We already said that we can think about this as being a transfer of an electron. OK, from the sodium to the chlorine. And when that happens, we get a sodium cation and we get a chlorine anion. OK, and so it makes perfect sense that around that sodium we get these blue indigo violet colors. But around the chlorine, we get this more aggressive red, orange and yellow colors. OK, and the fact that this transition between them, notice there's very little green here. Green is kind of like our neutral color. Okay. There's very little green there, which suggests that this is a very, very, very sharp transition, which makes perfect sense because there's literally no sharing of electrons. Okay. And so that's why we see a sharp transition between those colors and a very pronounced transition between those colors. Okay. And what that does, all right, is it also yields to an... asymmetrical distribution of charge. All right. That's why it, this, this picture, this image of sodium chloride, this electrostatic potential map does not appear to be symmetrical. Okay. And that's because again, what have we done? We've essentially taken an electron from sodium, given it to the chlorine, giving us a sodium cation and a chlorine anion. Now, on the other hand, if we're talking about these non-polar covalent molecules, right, we said that there's perfect sharing. There's perfect sharing. If there's perfect sharing, we expect there to be very little buildup of negative charge and very little buildup of positive charge. In fact, we expect there to be a symmetrical distribution of charge. And that's because the electrons are symmetrically distributed across this molecule. And when that happens, we get something that looks symmetrical. That's why... This image, all right, this electrostatic potential map of chlorine, chlorine of this chlorine gas in this example is very symmetrical. And in fact, it's more than that, right? It's all more or less these greens and yellows. Okay. And so that's telling us, all right, again, that there's really no buildup of charge anywhere on this molecule. Okay. Now, what happens with everything in between our two extremes? All right, because again, one extreme, we have ionic bonding. The other extreme, we have nonpolar covalent molecules, okay? Well, everything else in between, we call polar covalent molecules, all right? And polar, polar is telling us that we have unequal distribution of electron density, okay? I wouldn't compare that to ionic, all right? Because ionic is literally... A transfer of an electron, or at least for us, that's what I want you guys to think about. A transfer of an electron from something to something else that's going to accept that electron. Again, leading us to this sharp transition between our colors of the rainbow when we're talking about these electrostatic potential maps. How do we know what leads to these polar covalent molecules? Well, that's electronegativity. Okay, so what's electronegativity? This is the ability of an atom in a molecule to attract shared electrons to itself. Okay, so two things here. All right, this does not, the concept of electronegativity does not apply to atoms of any given element that exist by themselves. We have to be talking about molecules. Okay, we're talking about a molecule. All right. And what it is, is when we've bonded atoms of different elements together, the idea is which one of those atoms wants the electron density more within that bond, okay? And again, when we've talked about F2, fluorine is fluorine, okay? Fluorine has some propensity to want to withdraw electron density away from whatever it may be bonded to, okay? But in this case, it's... bonded to itself. All right. So neither of those fluorines want electron density more than the other. Again, so what leads, so what do we get from that? What's yielded from that? Okay. Well, we get a symmetrical electron distribution and that color is mostly green. You can imagine that one fluorine nuclei is here and the other one is here. And what you see is again, green color suggesting no buildup of negative or positive charge. And we get a symmetrical distribution of electron. density. Okay. All right. Well, what we can do, all right, is start to think about things that are not non-polar covalent, because again, that's what this is. This is non-polar covalent. Okay. And again, we have to think about what, we have to start to define some scale of electronegativity. Okay. And so that's what this little table down here is. Okay. So again, this is the periodic table. And what we've done, all right, is they have put values of electronegativity. All right. This is a relative scale. So in other words, we have to think about how much, okay. So again, it has to happen within a molecule. So we set up here and it's a relative scale, right? So how much does one thing want electron density relative to the other atom that it is bonded to, right? And so we have The scale, okay? And you can see all these numbers. And the greater the number an element has, the more that it wants that electron density, okay? So we go down here to like cesium, for example. This is our lowest electron, electronegativity, excuse me, all the way up here to fluorine, okay? Which is the highest electronegativity. So that should make some sense, right? Metals love to donate electrons. So it kind of makes sense that cesium have very little electronegativity versus our non-metals want to accept those electrons. might make some sense that fluorine is going to want electron density, okay? And in fact, there's a trend, okay? There's trends for electronegativity. Electronegativity increases from left to right on the periodic table, okay? And from bottom to top on the periodic table, okay? So again, what you can see here, is that as we transition from these teal colors, which again, most of those are metals, all right, to these blue colors, all right, that you see that right after those teal colors, those are essentially our metalloids. They have some sort of electronegativity that's intermediate, all right, from the metals, and then finally to those nonmetals. And those nonmetals have the highest electronegativities, all right, and in fact, fluorine has the greatest electronegativity, okay? So this kind of makes some sense relative to the things we've already talked about. Okay. And so what we can start to do is to define a, essentially a scale. Okay. A scale. And again, you see all these numbers here. I'm not expecting you to memorize all these values. Okay. If I wanted to do a specific calculation, all right, I would just give you the table and you could do a specific calculation. Right. But generally what we say is that if we. have a difference in our electronegativity. So again, we have to bond one atom of an element to another. All right. And if the difference in those electronegativity values that you see on this table, okay, if they're less than 0.5, we consider this to be nonpolar covalent. Okay, so in other words, we essentially say that this is perfect sharing of electrons within that covalent bond. On the other hand, if we have an electronegativity difference that's between 0.5 and 2.0, we call this a polar covalent bond. The polar meaning that, yes, we are sharing electrons because it is a covalent bond. All right. But it's an unequal sharing. OK. Of that electron density. OK. And then again, if we get above a 2.0 difference. All right. We call this essentially an ionic bond. OK. Where they're not really sharing electrons. We can again think about this as being an electron transfer event. from one species to the next, then they interact via these charge-to-charge interactions, okay? But again, if we're between a difference of 0.5 and 2.0, all right, this is a polar covalent bond. Again, that's defining this middle region that's between those two extremes of our ionic bonds and our polar, nonpolar covalent bonds, excuse me. So, okay, so let's take an example real quick. Well, let's look at fluorine bonded to hydrogen. Now, again, notice where hydrogen is in this particular case. Generally, people put it above those group 1A metals. It is not a metal, okay? It is a nonmetal, all right? The reason people put it there is because, again, it tends to form or it has the potential to form plus one cations. So again, what's that? That's our protons. So again, that's our sign for an acid, right? But what we can also do is think about placing hydrogen in between boron and carbon. Because again, it's got an electronegativity value that's somewhere intermediate to those two. All right. Okay. So what we can start to do is think about making a bond from fluorine to hydrogen. Okay. And we can think about the electronegativity. And this molecule is kind of like. written a little bit weird. I normally wouldn't write it like that, but it allows me to show the math a little bit easier. So we have an electronegativity of four. And again, we want to subtract from that the electronegativity of hydrogen, which is 2.1. All right. And what do we get? We get 1.9. Okay. And again, that falls between that 0.5 and that 2.0. Okay. So this is a polar covalent bond. All right. And it manifests itself in that electrostatic. potential map all right because again looky here okay you can imagine our hydrogen nuclei is here And our fluorine is here, all right? And we have a distribution of charge, an asymmetric distribution of charge, all right? Because that fluorine is withdrawing that electron density out of that bond, away from the hydrogen, it's left with a partially positive region of that molecule. And over here on the fluorine, we get a partial buildup of... negative charge. Okay. And again, this isn't a complete transfer. Notice that we have a gradual transition from our blues, our indigos and our purples through the green, through the red. Okay. So we have a gradual transition. All right. But it is still an asymmetrical distribution of electron density. This is called buildup of partial charge. Okay. So our symbol for partial is right here. Okay. So this Greek Delta, anytime you see that, that's telling us that we have a partial charge that's involved, not a complete charge that like what we see with those ionic species. Okay. So this is a partial buildup of charge. And what you'll see here for this example, right, is that we have a partial buildup of positive charge on that hydrogen, because again, this isn't an ionic. bond. We're not talking about a complete transfer of an electron. We're talking about unequal sharing and a partial buildup of negative charge. Well, when you sum the partial positive and the partial negative, they should still equal zero. But the point is this electrostatic potential map is showing unequal sharing. And the other important thing is that it shows you that there is an asymmetrical distribution of electronic character of electron density across that molecule. That's why it doesn't look symmetrical, like what we see with this chlorine gas example. Okay. So the last thing that's important to recognize here, all right, is that we can symbolize this in another way as well. Okay. So again, we have this asymmetrical buildup of electron density across this molecule, gives us an asymmetrical buildup of charge across that. molecule. All right. And so what we can do is we can start to represent this like this. All right. With an arrow, with this little hash mark towards the bottom end. So essentially making a plus, right? Okay. And so this plus end of this arrow always goes with the positive charge and the arrow head is always pointing towards the negative charge. Okay. So that's how you'll see this represented. Okay. And what this leads to, all right, because again, our charges are not canceling out with each other. In this case, it's an asymmetrical distribution that leads to something we call a dipole. Okay, a dipole or dipole moment. Okay, and whenever you have that, okay, a dipole moment, again, that is suggesting that there is an unequal distribution of. electron density across that molecule, which means that we have an asymmetrical distribution of charge across that molecule. In other words, we have one end of this molecule that's got a positive character and the other end of this molecule is more negative in character. And so with that, I hope you all have a great rest of your day and I'll be seeing you soon.

And as you might expect, there are many, many, many different reactions that work inside of an environment that is chocked full of water. And so what we want to do is build up this idea of why water is such a good solvent. Okay. So that's a term we're going to talk about in the next video. Okay.

And in order for us to do that, we need to discuss electronegativity and how electronegativity starts to lead to this idea of polarity. All right. So we've talked about two different types of bonding so far.

Okay. We've talked about ionic bonding and covalent bonding. So let's talk about ionic bonding first.

So if we think about Na. and Cl, okay? So sodium and chloride. So we have a metal and a nonmetal. What we like to think about, all right, is that that sodium atom can essentially lose an electron and adopt a plus one charge, okay?

So again, remember, our metals love to lose electrons, okay? And then what we can think about is that that chlorine, because that electron just doesn't dissipate, appear, all right, we can think about that chlorine wanting to adopt that electron. And when it does so, it adopts a negative charge.

Okay. So we end up with, at the end of the day, we get a sodium cation and a chlorine anion. Okay. So we've essentially transferred an electron from the sodium to the chloride.

And then what can happen? Well, our sodium and come together and interact with our chloride in an ionic bond, okay? An ionic bond is nothing more than a charge-charge interaction, okay? And so we can get an ionic compound, all right?

So it's a charge-charge interaction. And remember, the strength of this interaction, this is a very, very strong interactions, okay? They are dictated by Coulomb's law, all right?

So in this case, I've talked about monatomic cation interacting with a monatomic anion. But remember, we can also have our polyatomic ions come into play here too. We got a couple of polyatomic cations and many, many, many polyatomic anions.

They can also take part in ionic bonding. All right. Then the other type of bonding is our covalent bonds where we're sharing electrons.

And we're sharing electrons because what are we doing? We're getting some amount of orbital overlap. All right. So remember hydrogen. Hydrogen has a single proton.

So I'm just going to circle that inside of its nucleus. And then we have a single electron that is kind of zipping around outside of the nucleus that in an orbital, okay, we call them orbitals, and they're defined by some amount of probability. Okay, so probability defines where we might be able to find this electron.

All right. And what we can start to do is think about bringing two hydrogens together to make a hydrogen-hydrogen bond. So again, we like to think about this in terms of energy, where we have energy on our y-axis here and we have distance on our x-axis. And remember, the further away these hydrogens are, the less they're interacting. But as we start to bring these guys closer together, our energy starts to drop until we get to this energy minimum.

So this is where our energy... Well, will be. Okay.

So it's the lowest energy of our system. Okay. And what is that?

What defines that point? All right. It's the optimal distance between these two hydrogen atoms.

All right. And why is that the optimal distance? Well, remember, it's like a balancing act.

We have repulsive forces. We have attractive forces. At that optimal distance, what we've done is we've maximized our attractive forces. And we've minimized our repulsive forces.

So let's think about what those are. So again, we get orbital overlap. That's why we get a bond in the first place, okay?

Because that bond, that orbital overlap starts to focus, all right? Starts to localize those electrons between the two nuclei, between those two hydrogen nuclei, okay? And then now what we have is we have some attractive forces, but we also have repulsive forces, all right? So remember our attractive forces, right?

Clearly, those protons wanting to interact with those electrons and electrons with the protons. And that works for both of our atoms. Okay.

So those are our attractive forces that we're trying to maximize. Okay. But we also have some repulsive forces as well.

So remember, we still have some probability that those electrons are not in between those two nuclei. All right. So we still have these orbitals. Okay.

These electron clouds, right? And those electron clouds lead to some repulsion. And we also get some charge-charge repulsion from...

our two protons okay but again at this minimum energy what we've done is we have maximized okay we've maximized the attractive forces okay and minimized the repulsive forces Okay. And what that does is it leads us to this energy minimum. So a lot of energy has been given off that again, something else can harness and use to do some sort of work.

Okay. And we've decreased the energy of our system. And again, that is a favorable thing.

Okay. But what we've done now is we've made a bond. All right.

And so what do we do? We kind of represent that with this stick. Okay.

So we've made a hydrogen hydrogen bond. All right. You can also think about now moving those two nuclei closer to each other. All right. And then what happens with our energy then?

Okay. Well, as we move them closer to each other, of course, we start to maximize those repulsive forces. Right.

And so, of course, what happens here is our energy just kind of skyrockets. Okay. So the energy skyrockets. And once we get above that zero energy mark, what do we have to do to make that happen? we have to add energy into that system.

In other words, we have to add energy into that system to move those two nuclei closer to each other. Because again, what do we have to do? The energy is being supplied to overcome those repulsive forces. All right, so what we can also start to think about is doing the same sort of thing with fluorine. So fluorine has nine protons and nine electrons.

In this case, fluorine is one period below. So... it has two different shells that are holding those electrons, okay? There's two in this inner shell.

I'm just going to kind of highlight this guy. And we're going to talk a lot about this. There's two on this inner shell.

And then the other seven are in this outer shell, okay? This outer shell, those electrons are higher energy than the ones within that inner shell, okay? And this outer shell is called... our valence shell, okay, our valence shell.

And when we do bonding, all right, the electrons that take part in that bonding are in the valence shell, okay? So what we can, again, start to do is we can start to move two fluorines, okay, from being really far apart from one another to actually coming into contact with each other, all right? And when we find that optimal distance, what do we get? We get orbital overlap.

we start to focus electron density between those two nuclei. That focus density allows us to make a bond, which again, we can represent as a stick, okay? And at this point, all right, this is the point where we've maximized our attractive forces, okay? Minimize those repulsive forces. And this is also that point where we're at the lowest energy.

So energy has been released, okay? Because this is a favorable process, all right? But what I want you to notice is that, again, the way this picture is set up, it's showing you that it's those valence electrons, those electrons in that outer shell that are taking part in the bonding. And that's where the orbital overlap is also happening.

Okay. So what we can start to do is think about these ionic bonds, right, where we can essentially think about this as being an electron transfer event to give us two charged species that can interact with each other through charged-charge interactions. Or we can start to… bring nuclei together, get some orbital overlap where we can start to localize electron density between those two nuclei, essentially sharing those electrons, and we can get covalent bonds that result from that. So essentially what I've done now is I've described to you two extremes to bonding. There's two extremes to bonding where at one extreme we have ionic and again ionic We're talking about full charges here interacting with each other.

So again, if we're thinking about our sodium chloride, you can think of this M as being our sodium. and this X being our chloride, okay? So our sodium cation and our chloride anion, those guys coming together, making charge-charge interactions, and making very strong ionic bonds with one another, okay? So getting, which would lead us to these ionic solids that we've talked a lot about, okay? But on the other end, all right, we would get perfect sharing.

We'd get covalent bonds, okay? Where those electrons are shared perfectly between those two nuclei that are bonded together. So we get covalent bonds, and more specifically, we call these nonpolar covalent bonds. So we get these nonpolar covalent bonds. Something like hydrogen bonding with hydrogen, or fluorine with fluorine.

There's an example down here, right there. Chlorine with chlorine, so on and so forth. Perfect sharing. So these are defining the spectrums, the two extremes to bonding. And what we can start to do is to represent these things with something we call electrostatic potential maps.

And that's what these images are. These show us where the electrons are within. these molecules.

Okay. So you see a lot of colors. If you think about our rainbow, we have Roy G. Biff.

So our colors, red, orange, yellow, green, blue, indigo, violet. Okay. Well, when we're down here and the blues and the violets, those indigo colors, this is trying to tell us about the positive character of a molecule.

Whereas these more aggressive colors, the red, oranges, and yellows, those are where we have negative character. Well, how are we getting these builds up of negative and positive character? Well, again, remember, electrons carry negative charges. So when you get a buildup of electron density, you start to get a buildup of negative character, negative charge. All right.

Versus those blue colors. Well, what happens if you take away some negative character, you take away that electron density, there's less electron density to help neutralize that positive character coming from the protons. So you start to get a buildup of positive charge. Okay. So what you can see here is let's take a look at this electrostatic potential map for sodium chloride.

That example we've been talking about. Well, what you can see here, there's one particular color that's focused around that sodium. So let's just get that sodium here.

And then there's another color focused around those chlorides. Well, I just told you that those blue indigo violet types of colors, that's telling us that there's positive character. Whereas those red, orange, yellow colors, those more aggressive colors, that's where we have a buildup of negative charge. And so that makes perfect sense, right? We already said that we can think about this as being a transfer of an electron.

OK, from the sodium to the chlorine. And when that happens, we get a sodium cation and we get a chlorine anion. OK, and so it makes perfect sense that around that sodium we get these blue indigo violet colors. But around the chlorine, we get this more aggressive red, orange and yellow colors. OK, and the fact that this transition between them, notice there's very little green here.

Green is kind of like our neutral color. Okay. There's very little green there, which suggests that this is a very, very, very sharp transition, which makes perfect sense because there's literally no sharing of electrons.

Okay. And so that's why we see a sharp transition between those colors and a very pronounced transition between those colors. Okay.

And what that does, all right, is it also yields to an... asymmetrical distribution of charge. All right. That's why it, this, this picture, this image of sodium chloride, this electrostatic potential map does not appear to be symmetrical.

Okay. And that's because again, what have we done? We've essentially taken an electron from sodium, given it to the chlorine, giving us a sodium cation and a chlorine anion. Now, on the other hand, if we're talking about these non-polar covalent molecules, right, we said that there's perfect sharing.

There's perfect sharing. If there's perfect sharing, we expect there to be very little buildup of negative charge and very little buildup of positive charge. In fact, we expect there to be a symmetrical distribution of charge.

And that's because the electrons are symmetrically distributed across this molecule. And when that happens, we get something that looks symmetrical. That's why... This image, all right, this electrostatic potential map of chlorine, chlorine of this chlorine gas in this example is very symmetrical.

And in fact, it's more than that, right? It's all more or less these greens and yellows. Okay. And so that's telling us, all right, again, that there's really no buildup of charge anywhere on this molecule. Okay.

Now, what happens with everything in between our two extremes? All right, because again, one extreme, we have ionic bonding. The other extreme, we have nonpolar covalent molecules, okay? Well, everything else in between, we call polar covalent molecules, all right?

And polar, polar is telling us that we have unequal distribution of electron density, okay? I wouldn't compare that to ionic, all right? Because ionic is literally...

A transfer of an electron, or at least for us, that's what I want you guys to think about. A transfer of an electron from something to something else that's going to accept that electron. Again, leading us to this sharp transition between our colors of the rainbow when we're talking about these electrostatic potential maps.

How do we know what leads to these polar covalent molecules? Well, that's electronegativity. Okay, so what's electronegativity?

This is the ability of an atom in a molecule to attract shared electrons to itself. Okay, so two things here. All right, this does not, the concept of electronegativity does not apply to atoms of any given element that exist by themselves. We have to be talking about molecules. Okay, we're talking about a molecule.

All right. And what it is, is when we've bonded atoms of different elements together, the idea is which one of those atoms wants the electron density more within that bond, okay? And again, when we've talked about F2, fluorine is fluorine, okay? Fluorine has some propensity to want to withdraw electron density away from whatever it may be bonded to, okay? But in this case, it's...

bonded to itself. All right. So neither of those fluorines want electron density more than the other.

Again, so what leads, so what do we get from that? What's yielded from that? Okay.

Well, we get a symmetrical electron distribution and that color is mostly green. You can imagine that one fluorine nuclei is here and the other one is here. And what you see is again, green color suggesting no buildup of negative or positive charge.

And we get a symmetrical distribution of electron. density. Okay. All right.

Well, what we can do, all right, is start to think about things that are not non-polar covalent, because again, that's what this is. This is non-polar covalent. Okay. And again, we have to think about what, we have to start to define some scale of electronegativity. Okay.

And so that's what this little table down here is. Okay. So again, this is the periodic table.

And what we've done, all right, is they have put values of electronegativity. All right. This is a relative scale.

So in other words, we have to think about how much, okay. So again, it has to happen within a molecule. So we set up here and it's a relative scale, right? So how much does one thing want electron density relative to the other atom that it is bonded to, right?

And so we have The scale, okay? And you can see all these numbers. And the greater the number an element has, the more that it wants that electron density, okay? So we go down here to like cesium, for example.

This is our lowest electron, electronegativity, excuse me, all the way up here to fluorine, okay? Which is the highest electronegativity. So that should make some sense, right?

Metals love to donate electrons. So it kind of makes sense that cesium have very little electronegativity versus our non-metals want to accept those electrons. might make some sense that fluorine is going to want electron density, okay?

And in fact, there's a trend, okay? There's trends for electronegativity. Electronegativity increases from left to right on the periodic table, okay? And from bottom to top on the periodic table, okay?

So again, what you can see here, is that as we transition from these teal colors, which again, most of those are metals, all right, to these blue colors, all right, that you see that right after those teal colors, those are essentially our metalloids. They have some sort of electronegativity that's intermediate, all right, from the metals, and then finally to those nonmetals. And those nonmetals have the highest electronegativities, all right, and in fact, fluorine has the greatest electronegativity, okay? So this kind of makes some sense relative to the things we've already talked about. Okay.

And so what we can start to do is to define a, essentially a scale. Okay. A scale. And again, you see all these numbers here. I'm not expecting you to memorize all these values.

Okay. If I wanted to do a specific calculation, all right, I would just give you the table and you could do a specific calculation. Right.

But generally what we say is that if we. have a difference in our electronegativity. So again, we have to bond one atom of an element to another. All right.

And if the difference in those electronegativity values that you see on this table, okay, if they're less than 0.5, we consider this to be nonpolar covalent. Okay, so in other words, we essentially say that this is perfect sharing of electrons within that covalent bond. On the other hand, if we have an electronegativity difference that's between 0.5 and 2.0, we call this a polar covalent bond.

The polar meaning that, yes, we are sharing electrons because it is a covalent bond. All right. But it's an unequal sharing.

OK. Of that electron density. OK. And then again, if we get above a 2.0 difference. All right.

We call this essentially an ionic bond. OK. Where they're not really sharing electrons.

We can again think about this as being an electron transfer event. from one species to the next, then they interact via these charge-to-charge interactions, okay? But again, if we're between a difference of 0.5 and 2.0, all right, this is a polar covalent bond.

Again, that's defining this middle region that's between those two extremes of our ionic bonds and our polar, nonpolar covalent bonds, excuse me. So, okay, so let's take an example real quick. Well, let's look at fluorine bonded to hydrogen. Now, again, notice where hydrogen is in this particular case.

Generally, people put it above those group 1A metals. It is not a metal, okay? It is a nonmetal, all right?

The reason people put it there is because, again, it tends to form or it has the potential to form plus one cations. So again, what's that? That's our protons.

So again, that's our sign for an acid, right? But what we can also do is think about placing hydrogen in between boron and carbon. Because again, it's got an electronegativity value that's somewhere intermediate to those two.

All right. Okay. So what we can start to do is think about making a bond from fluorine to hydrogen. Okay. And we can think about the electronegativity.

And this molecule is kind of like. written a little bit weird. I normally wouldn't write it like that, but it allows me to show the math a little bit easier. So we have an electronegativity of four.

And again, we want to subtract from that the electronegativity of hydrogen, which is 2.1. All right. And what do we get? We get 1.9. Okay.

And again, that falls between that 0.5 and that 2.0. Okay. So this is a polar covalent bond.

All right. And it manifests itself in that electrostatic. potential map all right because again looky here okay you can imagine our hydrogen nuclei is here And our fluorine is here, all right? And we have a distribution of charge, an asymmetric distribution of charge, all right? Because that fluorine is withdrawing that electron density out of that bond, away from the hydrogen, it's left with a partially positive region of that molecule.

And over here on the fluorine, we get a partial buildup of... negative charge. Okay.

And again, this isn't a complete transfer. Notice that we have a gradual transition from our blues, our indigos and our purples through the green, through the red. Okay.

So we have a gradual transition. All right. But it is still an asymmetrical distribution of electron density. This is called buildup of partial charge.

Okay. So our symbol for partial is right here. Okay. So this Greek Delta, anytime you see that, that's telling us that we have a partial charge that's involved, not a complete charge that like what we see with those ionic species.

Okay. So this is a partial buildup of charge. And what you'll see here for this example, right, is that we have a partial buildup of positive charge on that hydrogen, because again, this isn't an ionic.

bond. We're not talking about a complete transfer of an electron. We're talking about unequal sharing and a partial buildup of negative charge. Well, when you sum the partial positive and the partial negative, they should still equal zero.

But the point is this electrostatic potential map is showing unequal sharing. And the other important thing is that it shows you that there is an asymmetrical distribution of electronic character of electron density across that molecule. That's why it doesn't look symmetrical, like what we see with this chlorine gas example. Okay. So the last thing that's important to recognize here, all right, is that we can symbolize this in another way as well.

Okay. So again, we have this asymmetrical buildup of electron density across this molecule, gives us an asymmetrical buildup of charge across that. molecule. All right.

And so what we can do is we can start to represent this like this. All right. With an arrow, with this little hash mark towards the bottom end. So essentially making a plus, right?

Okay. And so this plus end of this arrow always goes with the positive charge and the arrow head is always pointing towards the negative charge. Okay. So that's how you'll see this represented. Okay.

And what this leads to, all right, because again, our charges are not canceling out with each other. In this case, it's an asymmetrical distribution that leads to something we call a dipole. Okay, a dipole or dipole moment.

Okay, and whenever you have that, okay, a dipole moment, again, that is suggesting that there is an unequal distribution of. electron density across that molecule, which means that we have an asymmetrical distribution of charge across that molecule. In other words, we have one end of this molecule that's got a positive character and the other end of this molecule is more negative in character.

And so with that, I hope you all have a great rest of your day and I'll be seeing you soon.

Transcript for:Understanding Electronegativity and Polarity

Transcript for:
Understanding Electronegativity and Polarity