Periodic table, trends, and bonding
Periodic Table Trends: Ionization Energy What an ion is. Using the periodic table to understand how difficult it is to ionize an atom.
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- I think we're now ready to start talking about
- some general trends in the periodic table.
- And the first one that normally comes up
- in a lot of first-year chemistry classes
- and it even shows up
- on some chemistry standardized tests
- is the notion of ionization energy.
- Ionization Energy.
- And what this is, is essentially the energy required
- to remove an electron
- from the neutral version of that atom.
- So let's see.
- Energy to remove an electron.
- And just in case you haven't been exposed to
- the idea of what an ion is
- I guess this is a good time to explain it.
- And then this will make sense.
- So an ion is essentially any atom or molecule
- and we'll talk about molecules.
- Molecules are just combinations or groups of atoms
- that are all bonded together in some way.
- And we'll talk about bonding
- in a few videos from now.
- But it's an atom or a molecule
- in which the protons don't equal the electrons.
- And if the protons don't equal the electrons
- then you have some charge.
- If you have more protons than electrons
- you have a positive charge, right?
- So let's say if you have hydrogen.
- Let's say you have hydrogen.
- Normally it's neutral, right?
- It has no charge.
- But let's say you were to remove an electron from it.
- So now you have hydrogen plus an electron.
- Hydrogen only has one electron outside.
- So now it just has a proton on the inside
- no electrons on the outside.
- So now it has a plus one charge.
- We could write a minus one there.
- And now, what we've done in this,
- I guess we could say,
- this procedure or this reaction is, we've removed
- we've ionized hydrogen.
- We've removed an electron.
- And in this case, this type of ion that hydrogen is
- where your protons are greater than your electrons
- and this is just a nice word to know.
- This is called a cation.
- A cation is an ion with a positive charge.
- Now you just as easily could have a situation
- where you have a, let's say
- I don't know, let me take an element
- let's say we start with chlorine.
- Cl, right?
- In a stable form it has seven valence electrons.
- Why don't we add an electron to it.
- And now this will actually be pretty stable.
- It will have eight valence electrons
- but it'll have a negative charge.
- So this, right here, this is a negative ion.
- And that's known as an anion.
- Anion.
- The way I remember it is, in a lot of words,
- "a" means the opposite or a negation.
- Right?
- The "a" prefix.
- So "a", anion, means a negative ion.
- And then, obviously, cation is the other.
- But ionization energy
- it really should maybe be called cationization energy
- because it's the energy required
- to remove an electron
- not remove or add an electron.
- So it's really the energy required
- to turn something into a cation.
- So we've already discussed the periodic table
- and we'll get to this chart in a second.
- But just based on what we already know
- which elements will it be harder
- to remove an electron from
- and which ones will it be easier?
- We already talked about -- let's start with group one,
- just because it is group one.
- Group one is right here.
- And we'll start especially with the alkali metals.
- We can put hydrogen aside for now.
- But all of these, guys, we've talked about a lot.
- In order for them to get to the magic number eight
- in their outermost shell
- the easiest way for them to do it
- is just to get rid of that one valence electron
- they have in their outermost shell.
- Let's say potassium, right there.
- Potassium has one valence electron in its fourth shell.
- If it just got rid of it
- then it has eight in its third shell.
- Then it looks a lot like argon
- from an electron configuration point of view,
- which would be nice.
- So these, guys, really want to give away electrons.
- So it requires very little energy to ionize them
- or to cationize them, to take away their electron.
- So this is low ionization energy.
- And I think you see where this is going.
- What about these guys?
- What about neon?
- How hard is it to remove an electron from neon.
- Well neon is completely satisfied.
- It doesn't want to even deal with any of this
- reaction business and bonding business.
- It's like, you know what?
- I've achieved happiness in life.
- Don't mess with my electron.
- So it really doesn't want to give away its electrons.
- Neon, or krypton, or argon
- or any of these noble gases.
- So to remove an electron from one of these guys
- requires a lot of energy.
- So this is a high ionization energy.
- So in general
- as you go from left to right across the periodic table
- it goes from low to high.
- And some people memorize this,
- but you really don't have to memorize it.
- You just say, look,
- these guys have one extra electron
- that they're always trying to get rid of.
- These guys have two.
- Ideally they'd want to get rid of, maybe,
- both of their electrons.
- But the first electron doesn't want to jump off
- as much as the first electron here
- because the first electron here
- you get rid of it
- you immediately get to the super stable state.
- And that trend just becomes more and more true.
- This guy, definitely, under no circumstances
- wants to give away an electron.
- This guy is so close to being like neon
- that he definitely doesn't want to go a step backwards
- and look more like oxygen.
- So he doesn't want to get rid of an electron.
- So the trend is pretty clear.
- Anytime you're confused about the trend
- just look at the extreme cases.
- This guy wants to give away electrons.
- This guy wants to get electrons.
- So if you say, what's the energy required
- to take away an electron?
- Well this guy's almost going to give it to you.
- While this guy's going to be very hard,
- especially neon,
- to take away an electron.
- Now what happens as you go down?
- As you go down in this direction,
- let's say you go along a group, right?
- We already established that
- these alkali metals like to give away electrons.
- But as you go down,
- the electron cloud gets bigger and bigger.
- And you could say this 55th electron
- there's 55 protons,
- there's also in a neutral cesium atom,
- there's also 55 electrons
- that the 55th electron is
- a lot further away from the nucleus of this atom
- than the third electron is in the case of lithium.
- So this, the 55th electron,
- not only does it want to be given away,
- but it has even a weaker attraction to the nucleus
- than the third electron does in lithium.
- So because you're getting larger and larger
- as you go down a group,
- and the electrons are getting
- further and further away from the nucleus,
- this guy wants to give away his electrons
- even more than lithium does.
- Right?
- So ionization energy decreases as you go down.
- Even though xenon really wants to keep his electrons
- he's a little bit more willing to give them away
- than neon is.
- Right?
- So in general, the ionization energy,
- or the energy required to ionize an atom,
- will increase as you go up.
- Right?
- And if you ever forget it, don't memorize these things.
- Because that might be useful for just one test
- but then later in life
- when you're 42 and someone asks you,
- hey, what has a higher ionization energy,
- cesium or fluorine?
- You might have forgotten it.
- But then if you look at a periodic table, you'd say,
- you know, cesium has this one electron
- that it's just dying to give away.
- It's super far away from the nucleus.
- That 55th electron just really wants to leave.
- While fluorine, that ninth electron
- just needs one more electron.
- Well let's say neon.
- Neon is super happy.
- All the electrons are in the stable configuration.
- They're close to the nucleus.
- There's a lot of attraction
- with the protons in the nucleus.
- They definitely don't want to
- give away their electrons.
- So if you talk about the
- energy required to remove them
- very low energy at cesium
- very high energy at helium
- So that's the trend.
- And this is just, this trend we'll see often
- this is your willingness to give electrons.
- And this is how much you want to
- hog electrons or keep them to yourself.
- And this right here, this is actually
- I got this off Wikipedia
- this is the actual, measured ionization energies.
- And people looked at patterns like this.
- I'm not sure if they actually
- looked at the ionization energy,
- but they looked at patterns like this,
- and this is actually
- where they came up with the periodic table.
- Because they said, look,
- as we increment up
- the number of protons that we have in an atom,
- and likewise, the number of electrons,
- we see these repeated patterns, or these periods,
- in the elements.
- So this is hydrogen.
- So hydrogen's ionization energy
- is around 13 1/2 electron volts.
- Right?
- Which is a unit of energy.
- It can be converted into joules, if you like.
- But then all of a sudden, helium is a lot more stable.
- It takes almost double the energy
- to remove that second electron from helium
- because it's so stable.
- But as soon as you do that,
- as soon as you go from helium,
- this point right here,
- that point right here is lithium.
- Right?
- Lithium is atomic number of three.
- Let me put that in.
- This is lithium.
- All of a sudden, to remove an electron from lithium
- it only requires five electron volts.
- So less than half of what it required for hydrogen.
- Then as you go to the right of the periodic table,
- the ionization energy keeps increasing.
- These little divots are interesting.
- We could talk about that, maybe, in a future video.
- But the ionization energy increases
- all the way to neon.
- And then you get to neon
- and then you add one more
- see, neon is a noble gas, it's completely happy
- then you add one more electron
- and you get to sodium.
- And you say, oh, now sodium,
- it's really easy to take away that electron.
- And the ionization energy drops.
- Not only did it drop,
- but here it dropped slightly below lithium.
- So you see this general trend.
- So these are the noble gases right here.
- Let me make sure
- you can see them all in the video screen.
- Helium.
- Very hard to remove that eighth electron.
- Neon.
- It's very hard, but a little easier than helium.
- And that's because neon is a little bit bigger.
- The electrons are a little bit
- further away from the nucleus.
- Argon.
- Same pattern, but it's actually interesting.
- Argon is actually not that different than hydrogen.
- Krypton.
- Same thing.
- It's hard to remove that electron.
- But it's actually no harder
- to remove the electron on krypton
- than on hydrogen.
- And xenon.
- And you go all the way out to radon.
- And you can see, you know,
- why is the distance between these increasing?
- Well if you remember the periodic table,
- all of a sudden when you go after argon,
- then you have all of the d-block elements that go in.
- So after you get to argon, notice.
- This is where you only had s- and p-block elements.
- Now all of a sudden the d-block elements show up.
- So you have more subshells to fill up.
- And now all of a sudden, before you get to radon,
- you're starting to fill out the f-block, too.
- And that's why that distance is increasing,
- because you have to also fill out the f elements.
- But the general trends that we just discussed
- they apply here.
- That as you go to the right along the periodic table,
- it becomes harder and harder to remove an electron
- because everyone wants to
- get to the magic number eight.
- And it's very hard to remove them from a noble gas.
- But then as you go one more above that,
- this guy would love to give away his electron
- and get back to a configuration like neon.
- And then the other trend is, as you go down a group
- in this case, these are the noble gases
- the ionization energy decreases.
- And that's because
- the overall radius of the atom increases,
- so the valence electrons in radon are
- a little bit less strongly attracted
- than the valence electrons in helium.
- Anyway, I think that's it for ionization energy.
- I'll continue this in another video.
- We'll talk about
- metallic character and electronegativity.
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