Electronegativity

That last post was pretty long, huh? I'll bet it was too long for you. Well, you're in luck. This post is going to be a short one because I'm tired and electronegativity seems pretty easy to explain. It should be easy to explain, even to someone like you. That's why I'm skipping some stuff in this book that looks like it's about molecular orbital theory. Do you think you can handle molecular orbital theory? Yeah, maybe you can. I don't know. But for now, we're not going to worry about it.

Electronegativity is a property atoms have. Atoms that are highly electronegative have a high affinity for electrons. Atoms that are not very electronegative have a low affinity for electrons. Is that simple enough for you? Here. I'll personify the atoms for you, just like every high school science teacher on the planet has ever done. Atoms that are highly electronegative want electrons. They will take them. Atoms that are not very electronegative are more likely to lose their electrons to those mean, electronegative bullies. Now I feel all dirty.

Anyway, electronegativity is inversely proportional to atomic radius. So smaller atoms tend to be more electronegative than bigger ones. There are various explanations for this, but I took general chemistry a long time ago. I usually look at it in terms of "shielding" although technically, there's more to it. For an example, let's compare oxygen to sulfur. They're in the same group and have similar chemical properties, after all. Now, oxygen has eight protons, but sulfur has sixteen. And if you were good and read the second post I made here, you know that protons reside in the nucleus of the atom. So all those protons in sulfur are presumably twice as attractive to any electrons as the only half as many protons in oxygen. If that were the whole story, we'd expect sulfur to be more attractive to electrons. But sulfur also has twice as many electrons sitting in its orbitals. And they take up more space. Hence the mention of atomic radius. The sulfur atoms doesn't just have more subatomic particles. It is spatially bigger. And that large, negatively charged space has a repelling effect on other electrons. It's not strong enough to override the attractive effect the nucleus has on electrons. It doesn't shield the nucleus that much. But the effect is enough to make sulfur's electronegativity lower than oxygen's.

Of course, not all elements are in the same group as oxygen. And that matters because the number of valence electrons is also very important when it comes to electronegativity. One very simple way to look at electronegativity is that it increases as one moves to the right on the periodic table (ignoring the very last column for reasons that are obvious if you take general chemistry) and decreases as you move down the periodic table. The most electronegative element of all is fluorine. Second is oxygen, followed closely by chlorine. Fourth is nitrogen followed closely by bromine, with iodine being sixth. Whenever one of these elements forms a covalent bond with some other element I didn't just name, the electronegative one will have more of the electron density from the covalent bond.

By the way, carbon is more electronegative than hydrogen, but only by a little bit. It's not usually enough to concern us. But when oxygen or nitrogen (or any of the halogens) is in a molecule, electronegativity becomes important.