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In the last video we talked about ionization energy, or

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the energy required to remove an electron.

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And we saw the general trend in the periodic table, that

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when you're in the bottom left-hand side close to

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cesium, cesium really wants to give up electrons.

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It's a big atom.

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It only has one extra electron in its sixth shell.

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It can just give it up, and then it'll have

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five complete shells.

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So it really wants to give it away, so it requires very

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little energy to ionize.

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On the complete other side of the spectrum, helium requires

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a lot of energy to ionize.

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It's very happy.

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I has a full shell at the first shell.

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It's a very small atom.

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The electrons are very close to the protons.

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So the coulomb force is super-duper-duper strong.

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So it takes a lot of energy to remove that incremental

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electron, and we learned that.

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And the one thing I want to cover before moving on to

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other types of trends or properties amongst the

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different atoms is the idea of a second ionization energy.

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And I want to do this because sometimes it's covered on some

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chemistry exams or some chemistry standardized tests.

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And it's just the idea that ionization energy is the

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energy required to remove the first electron, to go from a

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neutral state to popping one electron off of it.

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The second ionization energy is, then, the energy required

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to remove the very next electron.

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And the reason why this is interesting, is sometimes

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they'll say, OK, what elements have a very high second

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ionization energy?

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And your temptation would be, OK, high ionization energy,

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that also probably means high second ionization energy.

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And that might be true.

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For example, neon has a very high ionization energy, It

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really wants to keep that 10th electron, because it fills out

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the second shell.

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And then, of course, even if you were able to remove that

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electron, to remove the ninth electron, when now its

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configuration looks a lot like fluorine,

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that's still very difficult.

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So you'd say its second ionization energy

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is still very high.

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But if you think about it, the elements with the highest

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second ionization energies are going to be some of the

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elements with the lowest ionization energy.

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So, think about it.

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And that might be kind of confusing.

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Lithium, for example.

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Very low ionization energy.

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It's got that extra electron.

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It just wants to give it away.

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But once it gives it away, it's in a very stable

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situation, Then its electron

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configuration looks like helium.

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So to remove that second electron is

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super-duper-duper difficult.

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So lithium has a very high second ionization energy.

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And so you might run into a question where they're like,

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which of these elements has the biggest difference between

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their ionization energy and their second ionization

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energy, where their second ionization energy is higher

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than their ionization energy.

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And lithium, or anything in group one, that would be true,

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because as soon as you remove one electron, its electron

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configuration becomes super stable, so removing that

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second one is super-duper difficult.

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And you also see this in this chart.

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This is, of course, the first ionization energies.

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But let's say, the case with lithium, you

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removed that electron.

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It was very easy.

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You only needed five electron volts to do it.

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But then your configuration looks a lot like helium.

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So that second ionization energy is going to look a lot

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like helium's first ionization energy.

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Anyway, I don't want to confuse you too much.

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But that's an interesting point that might pop up every

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now and then.

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Now another property, which is in a lot of ways, in my mind,

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related is the idea of electronegativity.

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The concept came up by Linus Pauling.

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I always remember him.

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He was a famous chemist. What I always remember is that he

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was famous for being convinced that Vitamin C was kind of the

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key to living forever.

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And he would take huge doses of Vitamin C.

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I should probably read up on that again.

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I don't want to spread lies about Linus Pauling.

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But I remember reading that when I was in high school.

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But anyway, he came up with the idea of electronegativity.

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And the idea is that when two atoms form covalent bonds--

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and I haven't taught you what a covalent bond is, and I was

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planning on doing that in a couple of videos from now--

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but the idea of a covalent bond is really just atoms

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sharing electrons.

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Let me draw that out.

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So if I have oxygen, oxygen looks something like this.

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I could draw it like that.

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I could also draw oxygen like this, just because I'm going

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to use these extra electrons to bond.

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And if you take oxygen like that and you add it to two

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hydrogens-- hydrogen has one electron--

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what's going to happen?

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You might not know yet, if you haven't seen a covalent bond.

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But the atoms will actually share electrons.

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So this oxygen, you put it in the center.

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You have these, over here.

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Let me draw it like that.

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The electrons from oxygen I'll do in green.

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And then hydrogen, I'll just do it in this orange color.

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So we have two of these hydrogens.

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So one hydrogen will be there.

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And then the other hydrogen will be there.

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Now what just happened?

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Well if this hydrogen can pretend that both of these

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electrons, it has to kind of share this green

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one with the oxygen.

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And the deal is, hey, I share the green one and you let me

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borrow the green one, and I'll let you borrow the orange one,

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we both can kind of feel like we have a stable electron

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configuration.

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Hydrogen feels good because the one s-shell

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is completely filled.

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Oxygen fills great because it's valence shell is

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completely filled with eight electrons, two

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of which are borrowed.

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So it feels great.

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This is a covalent bond, where the

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atoms are sharing electrons.

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And so this sometimes will be drawn like this.

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Oxygen.

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Those are the extra electron pairs of oxygen.

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And they'll just draw a line like that.

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And that line implicitly is saying, look, there's two

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atoms on either end.

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There's the oxygen electron there.

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And then you have the hydrogen electron there.

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And they're kind of shared.

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These two things mean the same thing.

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But that line just means a covalent bond.

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Now my whole point behind talking about covalent bonds a

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little prematurely is so that I can touch on

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electronegativity.

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And the idea that Linus Pauling came up with is that

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in these covalent bonds, the sharing is not equal.

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That some of the atoms will hog the

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electrons a little more.

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So in this case, oxygen.

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We learned about oxygen.

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Oxygen is way over here.

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It loves to grab electrons.

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It has a very high ionization energy.

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It's only two away from having an electron configuration

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similar to neon and being super-duper happy.

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So oxygen loves electrons.

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Hydrogen is a little bit here or there.

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It could gain an electron and then it'll

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have a stable 1s orbital.

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Or it could lose an electron and it will essentially just

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turn into a positive ion.

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It can go either way.

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So it's a little bit more ambivalent about what happens

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relative to the electrons.

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But oxygen really wants the electrons so

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that it can get completed.

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So in this relationship between oxygen and hydrogen,

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oxygen is more electronegative.

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It's more electronegative, which means it kind of hogs

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the electrons a little bit more.

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So if you were to draw this relationship here, it might

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look something-- If you were to draw this bond.

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This is all abstract.

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Maybe you would draw it a little bit

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heavier on that side.

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And this is not really at all a convention, but I

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just made that up.

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Or if you just drew the hydrogen and the oxygen part

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of it, maybe the electrons spend most of their time

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around-- this is a probability distribution-- and less of

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their time around hydrogen.

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And that would be true for the other hydrogen.

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They spend less of their time around the hydrogen and a lot

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more of their time around the oxygen.

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The idea of electronegativity is just that one atom is going

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to hog the electrons more when you form a covalent bond.

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Now if we wanted to figure out the trend of electronegativity

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on the periodic table, what do you think's going to happen?

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Which elements are likely to hog electrons?

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Well, the ones that love electrons.

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The ones that it's very hard to take

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electrons away from them.

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The ones that are super-close to completing a full eight

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valence electrons in their outermost shell.

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So the most electronegative atoms are

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going to be right here.

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They're going to be the halogens, especially the

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fluorine, because the small ones want the electrons even

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more because they're a small atom.

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The electrons are going to get closer to the nucleus.

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And the reason why I'm not talking about the noble gases

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here is because these don't form covalent bonds.

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They're always happy.

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They're all these inert gases.

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Inert just means that they don't do anything.

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A similar word is inertia.

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Inertia means the tendency to want to stay at rest, not do

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anything, or stay in motion, but I won't go

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into that too much.

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But these are inert.

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They don't do anything.

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So these guys react.

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They form covalent bonds up here.

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And when they form covalent bonds, they hog the atoms.

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Likewise, when these guys down here form covalent bonds,

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they're like, you know what, you can have the atoms. I

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don't need them.

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I'm actually happier without them altogether.

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In fact, sometimes these guys actually just

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give away the atom.

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They don't even form a covalent bond.

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It's called an ionic bond.

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We'll talk about that in the next video.

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But as you can see, the trend is the same as it is for

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ionization energy.

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These guys, a lot of energy required

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to remove an electron.

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That's because they love electrons.

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So these guys are also very electronegative.

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They're going to hog the electrons in a covalent bond.

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These guys, very low ionization energy.

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Very easy to take an electron away from them.

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And that's why they have very low electronegativity.

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They're very unlikely to hog an electron in a bond.

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Now the other trend that some people sometimes talk about is

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the metallic nature of an element.

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And so, there's a lot of things that, in my mind, I

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imagine when someone talks about metallic nature, I

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imagine it should conduct electricity, it should be

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shiny, it should be malleable.

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I can bend it without it cracking.

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That's how I imagine metallic nature.

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But when people talk about it in chemistry, they're really

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just talking about a willingness

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to give away electrons.

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That's metallic nature.

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And that is important.

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If you talk about something that's going to conduct

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electricity or be malleable or have this sea of electrons

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available that the atoms can sit in.

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But the same trend.

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Which atoms are very likely to give away electrons?

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Well, the bottom left, right?

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As you go down, the atom gets bigger, so the electrons are

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further away from the nucleus.

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So the coulomb force is weaker, so those electrons are

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more weakly bound.

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And also, if you just have one extra electron here or two

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extra electrons there in your outermost shell, you're just

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like, hey, let me get rid of them and then I'll have a

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complete outer shell.

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So these guys want to give away electrons.

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So they have a very high metallic nature.

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These guys want to keep electrons.

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And they want to take more.

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So they have a very low metallic nature.

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In fact these are completely non-metallic in any way.

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And if you were to say, within a group, the trend-- I mean, I

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did the diagonal, but that's in general true-- is that the

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further you go down a group, the size of the atom is

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increasing and the outer electrons are

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further from the nucleus.

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So the electron force is going to be weaker-- or the coulomb

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force is going to be weaker.

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So you're more likely to give away electrons.

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So your metallic nature will increase as you go down.

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And your metallic nature will increase as you go to the

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left, because when you only have a couple of electrons in

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your outermost shell, you want to give them away.

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So metallic nature, it goes in the opposite direction.

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It goes like that.

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But for the same reason.

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These guys love to hog electrons.

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These guys love to give them away.

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Right?

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So ionization energy increased to the top right.

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Electronegativity increased to the top right.

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Metallic nature increased to the bottom left.

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The last trend we could talk about is just atomic radius.

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And there's a lot of different ways to actually measure this.

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And there's no one best way, because obviously, we already

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talked about it, an atom doesn't have a fixed radius.

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The electron could show up pretty much anywhere.

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So you could just kind of do a hard boundary.

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OK, 90% chance of finding the electron.

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That's your sphere of the atom.

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Or you could say, OK, if this atom bonds with another atom,

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what is half the distance between the two nucleuses.

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Right?

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If you make a bond like that.

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This is the distance between the two nucleuses and then you

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can say the atomic radius is that.

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So there's a lot of ways.

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But I think you get the general idea.

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It's just the size of the atom.

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And you could already imagine that as you go down any one

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group, the size of the atom increases.

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You're adding on more and more energy levels,

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more and more shells.

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The atom is just getting larger and larger.

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In fact, we've used that as an argument as to why, as you go

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down, ionization energy goes down, or

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electronegativity goes down.

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So the atoms become larger as you go down.

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Now, the one thing that might be a little un-intuitive is

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what happens as you go to the right?

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You're adding electrons as you go to the right, but you're

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adding them all in the same shell, right?

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So if this is the nucleus, right there, and you're in

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some shell, some orbital shell.

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And obviously, they're not all spheres.

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But let's say you're in some orbital shell.

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As you go to the right in a period, you just keep adding

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electrons to that shell.

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Right?

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This is a super-gross oversimplification.

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And as you go to the right, you have more

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protons in the nucleus.

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So this is only getting more and more positively charged.

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So what happens is that these electrons get pulled inwards.

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They get pulled inwards.

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So as you move to the right on the

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periodic table, size decreases.

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And then you say, OK, but what about when you

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go to the next period?

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You're getting more protons there.

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Won't that decrease?

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You are.

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But at the same time, you're now adding the electrons in a

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new shell that's further from them.

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So it gets larger when you go to the new period.

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So electron size, as you go down, large.

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And as you go to the left, you get larger.

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So electron size goes from the bottom right to the top left.

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Although in general, the things that are in a lower

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period will have a larger size than most things in a higher

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period, regardless of what group it's in.

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But the general trend within a group, the higher the number,

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the larger the atom.

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Within a period, the more protons you have,

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the smaller the atom.

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Anyway, I hope you found those interesting.

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In the next few videos we'll start with bonding.

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