Electronegativity and bonding

Electronegativity is probably the most important concept to understand inorganic chemistry. We're going to use a definition that Linus Pauling gives in his book, The Nature of the Chemical Bond. So Linus Pauling says that electronegativity refers to the power of an atom in a molecule to attract electrons to itself. So if I look at a molecule, I'm going to compare two atoms in that molecule. I'm going to compare carbon to oxygen in terms of the electronegativity. And to do that, I need to look over here in the right at the organic periodic table, which shows the elements most commonly used in organic chemistry. And then in blue, it gives us the Pauling scale for electronegativity. So Linus Pauling actually calculated electronegativity values for the elements and put them into the table. And that allows us to compare different elements in terms of their electronegativities. For example, we are concerned with carbon, which has an electronegativity value of 2.5. And we're going to compare that to oxygen, which has an electronegativity value of 3.5. So oxygen is more electronegative than carbon. And the definition tells us that if oxygen is more electronegative, oxygen has a greater power to attract electrons to itself than carbon does. And so if you think about the electrons and the covalent bonds between carbon and oxygen that are shared, they're shared unequally. Because oxygen is more electronegative, oxygen is going to pull those electrons in red closer to itself. And since electrons are negatively charged, the oxygen is going to get a little bit more negative charge. And so it's going to have what we call a partial negative charge on it. So partial negative. Its partial sign is a lowercase Greek letter, delta. And so the oxygen is partially negative. It's pulling the electrons in red closer to itself. Another way to show the movement of those electrons in red closer to the oxygen would be this funny arrow here. So the arrow points in the direction of the movement of the electrons in red. So carbon is losing some of those electrons in red. Carbon is losing a little bit of electron density. Carbon is losing a little bit of negative charge. So carbon used to be neutral, but since it's losing a little bit of negative charge, this carbon will end up being partially positive, like that. So the carbon is partially positive. And the oxygen is partially negative. That's a polarized situation. You have a little bit of negative charge on one side, a little bit of positive charge on the other side. So let's say it's still a covalent bond, but it's a polarized covalent bond due to the differences in electronegativities between those two atoms. Let's do a few more examples here where we show the differences in electronegativity. So if I were thinking about a molecule that has two carbons in it, and I'm thinking about what happens to the electrons in red. Well, for this example, each carbon has the same value for electronegativity. So the carbon on the left has a value of 2.5. The carbon on the right has a value of 2.5. That's a difference in electronegativity of zero. Which means that the electrons in red aren't going to move towards one carbon or towards the other carbon. They're going to stay in the middle. They're going to be shared between those two atoms. So this is a covalent bond, and there's no polarity situation created here since there's no difference in electronegativity. So we call this a non-polar covalent bond. This is a non-polar covalent bond, like that. Let's do another example. Let's compare carbon to hydrogen. So if I had a molecule and I have a bond between carbon and hydrogen, and I want to know what happens to the electrons in red between the carbon and hydrogen. We've seen that. Carbon has an electronegativity value of 2.5. And we go up here to hydrogen, which has a value of 2.1. So that's a difference of 0.4. So there is the difference in electronegativity between those two atoms, but it's a very small difference. And so most textbooks would consider the bond between carbon and hydrogen to still be a non-polar covalent bond. All right. Let's go ahead and put in the example we did above, where we compared the electronegativities of carbon and oxygen, like that. When we looked up the values, we saw that carbon had an electronegativity value of 2.5 and oxygen had a value of 3.5, for difference of 1. And that's enough to have a polar covalent bond. Right? This is a polar covalent bond between the carbon and the oxygen. So when we think about the electrons in red, electrons in red are pulled closer to the oxygen, giving the oxygen a partial negative charge. And since electron density is moving away from the carbon, the carbon gets a partial positive charge. And so we can see that if your difference in electronegativity is 1, it's considered to be a polar covalent bond. And if your difference in electronegativity is 0.4, that's considered to be a non-polar covalent bond. So somewhere in between there must be the difference between non-polar covalent bond and a polar covalent bond. And most textbooks will tell you approximately somewhere in the 0.5 range. So if the difference in electronegativity is greater than 0.5, you can go ahead and consider it to be mostly a polar covalent bond. If the difference in electronegativity is less than 0.5, we would consider that to be a non-polar covalent bond. Now, I should point out that we're using the Pauling scale for electronegativity here. And there are several different scales for electronegativity. So these numbers are not absolute. These are more relative differences. And it's the relative difference in electronegativity that we care the most about. Let's do another example. Let's compare oxygen to hydrogen. So let's think about what happens to the electrons between oxygen and hydrogen. So the electrons in red here. All right. So we've already seen the electronegativity values for both of these atoms. Oxygen had a value of 3.5, and hydrogen had a value of 2.1. So that's an electronegativity difference of 1.4. So this is a polar covalent bond. Since oxygen is more electronegative than hydrogen, the electrons in red are going to move closer to the oxygen. So the oxygen is going to get a partial negative charge. And the hydrogen is going to get a partial positive charge, like that. All right. Let's do carbon and lithium now. So if I go ahead and draw a bond between carbon and lithium, and once again, we are concerned with the two electrons between carbon and lithium. The electronegativity value for carbon we've seen is 2.5. We need to go back up to our periodic table to find the electronegativity value for lithium. So I go up here, and I find lithium in group one of my periodic table has an electronegativity value of 1. So I go back down here, and I go ahead and put in a 1. And so that's a difference in electronegativity of 1.5. So we could consider this to be a polar covalent bond. This time, carbon is more electronegative than lithium. So the electrons in red are going to move closer to the carbon atom. And so the carbon is going to have a little bit more electron density than usual. So it's going to be partially negative. And the lithium is losing electron density, so we're going to say that lithium is partially positive. Now here, I'm treating this bond as a polar covalent bond. But you'll see in a few minutes that we could also consider this to be an ionic bond. And that just depends on what electronegativity values you're dealing with, what type of chemical reaction that you're working with. So we could consider this to be an ionic bond. Let's go ahead and do an example of a compound that we know for sure is ionic. Sodium chloride, of course, would be the famous example. So to start with, I'm going to pretend like there's a covalent bond between the sodium and the chlorine. All right. So I'm going to say there's a covalent bond to start with. And we'll put in our electrons. And we know that this bond consists of two electrons, like that. Let's look at the differences in electronegativity between sodium and chlorine. All right. So I'm going to go back up here. I'm going to find sodium, which has a value of 0.9, and chlorine which has a value of 3. So 0.9 for sodium and 3 for chlorine. So sodium's value is 0.9. Chlorine's is 3. That's a large difference in electronegativity. That's a difference of 2.1. And so chlorine is much more electronegative than sodium. And it turns out, it's so much more electronegative that it's no longer going to share electrons with sodium. It's going to steal those electrons. So when I redraw it here, I'm going to show chlorine being surrounded by eight electrons. So these two electrons in red-- let me go ahead and show them-- these two electrons in red here between the sodium and the chlorine, since chlorine is so much more electronegative, it's going to attract those two electrons in red so strongly that it completely steals them. So those two electrons in red are going to be stolen by the chlorine, like that. And so the sodium is left over here. And so chlorine has an extra electron, which gives it a negative 1 formal charge. So we're no longer talking about partial charges here. Chlorine gets a full negative 1 formal charge. Sodium lost an electron, so it ends up with a positive formal charge, like that. And so we know this is an ionic bond between these two ions. So this represents an ionic bond. So the difference in electronegativity is somewhere between 1.5 and 2.1, between a polar covalent bond and an ionic bond. So most textbooks we'll see approximately somewhere around 1.7. So if you're higher than 1.7, it's generally considered to be mostly an ionic bond. Lower than 1.7, in the polar covalent range. But that doesn't always have to be the case. Right? So we'll come back now to the example between carbon and lithium. So if we go back up here to carbon and lithium, here we treat it like a polar covalent bond. But sometimes you might want to treat the bond in red as being an ionic bond. So let's go ahead and draw a picture of carbon and lithium where we're treating it as an ionic bond. So if carbon is more electronegative than lithium, carbon's going to steal the two electrons in red. So I'll go ahead and show the electrons in red have now moved on to the carbon atom. So it's no longer sharing it with the lithium. Carbon has stolen those electrons. And lithium is over here. So lithium lost one of its electrons, giving it a plus 1 formal charge. Carbon gained an electron, giving it a negative 1 formal charge. And so here, we're treating it like an ionic bond. Full formal charges here. And this is useful for some organic chemistry reactions. And so what I'm trying to point out here is these divisions, 1.7, it's not absolute. It's a relative thing. You could draw the dot structure above, and this would be considered be correct. Right? You could draw it like this. Or you could treat it like an ionic bond down here. This is relatively close to the cutoff. So this is an overview of electronegativity. And even though we've been dealing with numbers in this video, in future videos, we don't care so much about the numbers. We care about the relative differences in electronegativity. So it's important to understand something as simple as oxygen is more electronegative than carbon. And that's going to help you when you're doing organic chemistry mechanisms.