Welcome to the second video of chapter 4.2 covalent bonding. In this video our learning objective is to define electronegativity and assess the polarity of covalent bonds. As we jump in let's go ahead and first define what's the difference between a pure covalent versus a polar covalent bond.
So a pure covalent bond is a bond with perfectly equal sharing of electrons between the atoms that make up that bond. which is to say the atoms both have equal electron density or an equal distribution of electron density around them. A polar covalent bond is a bond with unequal distribution of electrons around the two atoms and this is characterized by a positive end and a negative end which is gives us the sort of terminology of poles right so polar covalent there's one that's positively charged and the other end is negatively charged. So we can see this depicted here with an HCl atom or molecule, sorry an HCl molecule, hydrogen chloride or hydrochloric acid if you've dissolved it in water.
It's really important to note however that we're not dealing ever with chlamylobondes, we're not dealing with full charges, we're dealing with partial charges and so we need to use some notation that tells us that these are partial charges as opposed to full charges like we see with ionic compounds. and that notation is this lowercase Greek delta. So this basically just means that this chlorine atom has a partial negative charge and this hydrogen atom has a partial positive charge. This is shown here with this difference in electron density around each of these atoms.
So the chlorine side has more electron density around it. The electrons are spending more of their time around this chlorine atom. And again because electrons are negatively charged that means that the chlorine has a partial negative charge. And when the electrons are spending most of their time around the chlorine, they're not spending very much time around the hydrogen which leaves that with a partial positive charge.
And so we have these two poles. This arrow notation is a nice way for us to express this quickly. The positive end of the molecule gets the tail of the arrow which looks like a little positive. and then we draw the arrow towards the more negative atom, the one that has the partial negative.
So we need a way to sort of describe this in if we can come up with some numbers that's really helpful. And so a very famous chemist called Linus Pauling came up with electronegativity. which is a measure of the tendency of an atom to attract electrons or rather electron density towards itself when it's in a bond. So essentially this is a measure of exactly what we're hoping to describe here, this ability of an atom to pull the density towards itself and away from its bonding partner in a bond. This is a chart of electronegativities, so there's some important things to note here.
First off, there's no noble gases, right? We don't see helium, neon, and the reason for that is because they don't tend to form covalent bonds. You can force noble gases to form covalent bonds if you put them in extremely, well, extreme circumstances, but in general they don't tend to form covalent bonds, and so they don't fit neatly into the model that we have for electronegativity, and so we just don't include them when we're discussing electronegativity.
The trend here is that electronegativity increases as you go towards the right across the periodic table and it decreases as you go down towards the bottom of the periodic table. I like to describe this as a diagonal trend, which means that if you can remember that fluorine is the most electronegative atom on the periodic table, then you know both directions, right? You know that it increases towards the right and you know that it decreases as you go down. So we need to keep in mind the difference between electronegativity and electron affinity. They are related but they're very different terms and they mean very different things.
They sound very similar but they are very different. So electron affinity, which we talked about earlier, is the energy released or absorbed during an electron capture. So it is a physical measured thing. It's it's a measured physical quantity and it represents an actual measurable physical quantity.
Electronegativity is not a measurable physical quantity. It's calculated. It's unitless.
and it's on an arbitrary scale between 0 and 4, and it essentially just describes how tightly an atom pulls electrons towards itself as it is in a bond. Alright, so we want to talk about types of bonds. So what makes a bond pure covalent versus polar covalent versus all the way to ionic when we have complete electron transfer? And essentially it's the difference in electronegativity.
And we're going to think about the absolute value, but it's the difference in electronegativity. And so we're going to represent that a lot by delta En. And so delta is change or difference, and then En is just a representation of electronegativity.
So the larger the difference in electronegativity, that is to say the larger the difference in the ability of these atoms to hold electrons tightly towards themselves in a bond. So the larger that difference, the more likely it is that one is going to essentially win the tug of war more often. And when that happens, if you have a really large difference in electronegativity, you're going to have either a polar covalent bond or a full ionic bond where the electron is just transferred.
So one example of these polar covalent bonds is things between hydrogen and oxygen. So... Hydrogen has an electronegativity of 2.1. Oxygen has an electronegativity, if we look on this chart, of 3.5.
So the difference in these guys, the difference between those guys is 1.4. And what this tells us is that one of these atoms is significantly more electronegative. It holds the electrons in the bond much more tightly than the other one.
And it will basically claim the electron density much more often than the hydrogen is able to. when they are sharing electrons in a bond. So what this looks like is, for example, if you have water and we will talk about geometry a little bit later, but this is a depiction of a water molecule, an oxygen with two hydrogens bonded to it.
We can see that oxygen has a greater electronegativity and so we're going to draw this sort of arrow thingy up towards the oxygen and actually each bond has this sort of, we can represent this polarity on each bond. So the oxygen is winning winning the tug of war for electrons versus the hydrogen, the density of electrons is going to be focused around the oxygen and the density will be much lower around the two hydrogens in that water molecule. If you have atoms that have a more similar electronegativity, like for example two identical atoms, then you have a pure covalent bond where there's no difference in electronegativity.
We don't draw any polar representations. There's no partial positives or partial negatives and that essentially happens when the difference in electronegativity is very small or the atoms are identical. So this table represents some of the values that we have sort of set there.
It's guidelines, there's exceptions to this. For example HF, so hydrogen has an electronegativity of 2.1, fluorine has an electronegativity of 4.0. The difference in these guys is 1.9. If we just go by the table, the difference is so big that this would actually predict that the electron would just be transferred and this bond would be completely ionic.
In fact, that's not the case. This is a polar covalent. It's just very polar.
And the same thing is true. There are ionic bonds that you would not predict would be ionic. You would predict they'd be polar covalent because the difference in electronegativity is too small.
So this is just guidelines, but in general you can think about how big the difference is between the bonding atoms. So if there's no difference, they're pure covalent bonds. If there's a medium difference, then they're probably experiencing a polar covalent bonds.
And if there's a huge difference in electronegativities between those atoms, like for example between the metals and the non-metals that we see up here, the metals have very low electronegativity values, then you are likely to see ionic bonds in those compounds. Alright, the last thing we're going to talk about in this video is polyatomics. So it's important for us to note that you can actually have ionic compounds that contain covalent compounds.
And what that looks like is you can have these polyatomic ions, like for example, nitrate, where the bonds between these atoms are covalent bonds. But you can also have an ionic. or an ionic compound form if you have an attraction between a sodium plus ion and your nitrate anion. You can form sodium nitrate.
which is a ionic compound but it contains this covalent ion.