States of matter
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States of Matter
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States of Matter Follow-Up
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Specific Heat, Heat of Fusion and Vaporization
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Chilling Water Problem
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Phase Diagrams
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Van Der Waals Forces
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Covalent Networks, Metallic, and Ionic Crystals
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Vapor Pressure
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Suspensions, Colloids and Solutions
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Solubility
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Boiling Point Elevation and Freezing Point Suppression
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Change of State Example
Van Der Waals Forces Van Der Waals Forces: London Dispersion Forces, Dipole Attractions, and Hydrogen Bonds.
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- Throughout our journey through chemistry so far, we've
- touched on the interactions between molecules, metal
- molecules, how they attract each other because of the sea
- of electrons and water molecules.
- But I think it's good to have a general discussion about all
- of the different types of molecular interactions and
- what it means for the boiling points or the melting points
- of a substance.
- So I'll start with the weakest. Let's say I had a
- bunch of helium.
- Helium, you know, I'll just draw it as helium atoms. We'll
- look in the Periodic Table, and what I'm going to do now
- with helium I could do with any of the noble gases.
- Because the point is that noble gases are happy.
- Their outer orbital is filled.
- Let's say, neon or helium-- let me do neon, actually,
- because neon has a full eight in its orbital so we could
- write neon like neon and it's completely happy.
- It's completely satisfied with itself.
- And so in a world where it's completely satisfied, there's
- no obvious reason just yet-- I'm going to touch on a reason
- why it should be-- if these electrons are evenly
- distributed around these atoms, then these are
- completely neutral atoms. They don't want to bond with each
- other or do anything else, so they should just float around
- and there's no reason for them to be attracted to each other
- or not attracted to each other.
- But it turns out that neon does have a liquid state, if
- you get cold enough, and so the fact that it has a liquid
- state means that there must be some force that's making the
- neon atoms attracted to each other, some force out there.
- Because it's in a very cold state, because for the most
- part, there is not a lot of force that attracts them so
- it'll be a gas at most temperatures.
- But if you get really cold, you can get a very weak force
- that starts to connect or makes the neon molecules want
- to get towards each other.
- And that force comes out of the reality that we talked
- about early on that electrons are not in a fixed, uniform
- orbit around things.
- They're probablistic.
- And if we imagine, let me say neon now, instead of drawing
- these nice and neat valence dot electrons like that,
- instead, I can kind of draw its electrons as-- it's a
- probability cloud and it's what neon's atomic
- configuration is.
- 1s2 and it's outer orbital is 2s2 2p6, right?
- So it's highest energy electron, so, you know, it'll
- look-- I don't know.
- It has the 2s shell.
- The 1s shell is inside of that and it has the p-orbitals.
- The p-orbitals look like that in different dimensions.
- That's not the point.
- And then you have another neon atom and these are-- and I'm
- just drawing the probability distribution.
- I'm not trying to draw a rabbit.
- But I think you get the point.
- Watch the electron configuration videos if you
- want more on this, but the idea behind these probability
- distributions is that the electrons could be anywhere.
- There could be a moment in time when all the electrons
- are out over here.
- There could be a moment in time where all the electrons
- are over here.
- Same thing for this neon atom.
- If you think about it, out of all of the possible
- configurations, let's say we have these two neon atoms,
- there's actually a very low likelihood that they're going
- to be completely evenly distributed.
- There's many more scenarios where the electron
- distribution is a little uneven in one
- neon atom or another.
- So if in this neon atom, temporarily its eight valence
- electrons just happen to be like, you know, one, two,
- three, four, five, six, seven, eight, then what does this
- neon atom look like?
- It temporarily has a slight charge in
- this direction, right?
- It'll feel like this side is more negative than this side
- or this side is more positive than that side.
- Similarly, if at that very same moment I had another neon
- that had one, two, three, four, five, six, seven, eight,
- that had a similar-- actually, let me do that differently.
- Let's say that this neon atom is like this: one, two, three,
- four, five, six, seven, eight.
- So here, and I'll do it in a dark color because it's a very
- faint force.
- So this would be a little negative.
- Temporarly, just for that single moment in time, this
- will be kind of negative.
- That'll be positive.
- This side will be negative.
- This side will be positive.
- So you're going to have a little bit of an attraction
- for that very small moment of time between this neon and
- this neon, and then it'll disappear, because the
- electrons will reconfigure.
- But the important thing to realize is that almost at no
- point is neon's electrons going to be completely
- distributed.
- So as long as there's always going to be this haphazard
- distribution, there's always going to be a little bit of
- a-- I don't want to say polar behavior, because that's
- almost too strong of a word.
- But there will always be a little bit of an extra charge
- on one side or the other side of an atom, which will allow
- it to attract it to the opposite side charges of other
- similarly imbalanced molecules.
- And this is a very, very, very weak force.
- It's called the London dispersion force.
- I think the guy who came up with this, Fritz London, who
- was neither-- well, he was not British.
- I think he was German-American.
- London dispersion force, and it's the weakest of the van
- der Waals forces.
- I'm sure I'm not pronouncing it correctly.
- And the van der Waals forces are the class of all of the
- intermolecular, and in this case, neon-- the
- molecule, is an atom .
- It's just a one-atom molecule, I guess you could say.
- The van der Waals forces are the class of all of the
- intermolecular forces that are not covalent bonds and that
- aren't ionic bonds like we have in salts, and we'll touch
- on those in a second.
- And the weakest of them are the London dispersion forces.
- So neon, these noble gases, actually, all of these noble
- gases right here, the only thing that they experience are
- London dispersion forces, which are the weakest of all
- of the intermolecular forces.
- And because of that, it takes very little energy to get them
- into a gaseous state.
- So at a very, very low temperature, the noble gases
- will turn into the gaseous state.
- That's why they're called noble gases, first of all.
- And they're the most likely to behave like ideal gases
- because they have very, very small
- attraction to each other.
- Fair enough.
- Now, what happens when we go to situations when we go to
- molecules that have better attractions or that are a
- little bit more polar?
- Let's say I had hydrogen chloride, right?
- Hydrogen, it's a little bit ambivalent about whether or
- not it keeps its electrons.
- Chloride wants to keep the electrons.
- Chloride's quite electronegative.
- It's less electronegative than these guys right here.
- These are kind of the super-duper electron hogs,
- nitrogen, oxygen, and fluorine, but chlorine is
- pretty electronegative.
- So if I have hydrogen chloride, so I have the
- chlorine atom right here, it has seven electrons and then
- it shares an electron with the hydrogen.
- It shares an electron with the hydrogen, and I'll
- just do it like that.
- Because this is a good bit more electronegative than
- hydrogen, the electrons spend a lot of time out here.
- So what you end up having is a partial negative charge on the
- side, where the electron hog is, and a
- partial positive side.
- And this is actually very analogous to
- the hydrogen bonds.
- Hydrogen bonds are actually a class of this type of bond,
- which is called a dipole bond, or dipole-dipole interaction.
- So if I have one chlorine atom like that and if I have
- another chlorine atom, the other chlorine
- atoms looks like this.
- If I have the other chlorine atom-- let me copy and paste
- it-- right there, then you'll have this
- attraction between them.
- You'll have this attraction between these two chlorine
- atoms-- oh, sorry, between these two
- hydrogen chloride molecules.
- And the positive side, the positive pole of this dipole
- is the hydrogen side, because the electrons have kind of
- left it, will be attracted to the chlorine side
- of the other molecules.
- And because this van der Waals force, this dipole-dipole
- interaction is stronger than a London dispersion force.
- And just to be clear, London dispersion forces occur in all
- molecular interactions.
- It's just that it's very weak when you compare it to pretty
- much anything else.
- It only becomes relevant when you talk about things with
- noble gases.
- Even here, they're also London dispersion forces when the
- electron distribution just happens to go one way or the
- other for a single instant of time.
- But this dipole-dipole interaction is much stronger.
- And because it's much stronger, hydrogen chloride is
- going to take more energy to, one, get into the liquid
- state, or even more, get into the gaseous state than, say,
- just a sample of helium gas.
- Now, when you get even more electronegative, when this
- guy's even more electronegative when you're
- dealing with nitrogen, oxygen or fluorine, you get into a
- special case of dipole-dipole interactions, and that's the
- hydrogen bond.
- So it's really the same thing if you have hydrogen fluoride,
- a bunch of hydrogen fluorides around the place.
- Maybe I could write fluoride, and I'll write hydrogen
- fluoride here.
- Fluoride its ultra-electronegative.
- It's one of the three most electronegative atoms on the
- Periodic Table, and so it pretty much
- hogs all of the electrons.
- So this is a super-strong case of the dipole-dipole
- interaction, where here, all of the electrons are going to
- be hogged around the fluorine side.
- So you're going to have a partial positive charge,
- partial negative side, partial positive, partial negative,
- partial positive, partial negative and so on.
- So you're going to have this, which is really a dipole
- interaction.
- But it's a very strong dipole interaction, so people call it
- a hydrogen bond because it's dealing with hydrogen and a
- very electronegative atom, where the electronegative atom
- is pretty much hogging all of hydrogen's one electron.
- So hydrogen is sitting out here with just a proton, so
- it's going to be pretty positive, and it's really
- attracted to the negative side of these molecules.
- But hydrogen, all of these are van der Waals.
- So van der Waals, the weakest is London dispersion.
- Then if you have a molecule with a more electronegative
- atom, then you start having a dipole, where you have one
- side where molecule becomes polar and you have the
- interaction between the positive and the negative side
- of the pole.
- It gets a dipole-dipole interaction.
- And then an even stronger type of bond is a hydrogen bond
- because the super-electronegative atom is
- essentially stripping off the electron of the hydrogen, or
- almost stripping it off.
- It's still shared, but it's all on
- that side of the molecule.
- Since this is even a stronger bond between molecules, it
- will have even a higher boiling point.
- So London dispersion, and you have dipole or polar bonds,
- and then you have hydrogen bonds.
- All of these are van der Waals, but because the
- strength of the intermolecular bond gets stronger, boiling
- point goes up because it takes more and more energy to
- separate these from each other.
- In the next video-- I realize I'm out of time.
- So this is a good survey, I think, of just the different
- types of intermolecular interactions that aren't
- necessarily covalent or ionic.
- In the next video, I'll talk about some of the covalent and
- ionic types of structures that can be formed and how that
- might affect the different boiling points.
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At 5:31, how is the moon large enough to block the sun? Isn't the sun way larger?
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