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  • - [Instructor] So let's talk a little bit

  • about molecular solids.

  • So just as a little bit of review,

  • we've talked about ionic solids

  • where ions form these lattices.

  • So those might be the positive ions right over there,

  • and then you have your negative ions.

  • And the negative is attracted to the positive.

  • The positive is attracted to the negative.

  • And I'm just showing a two-dimensional version of it,

  • but it forms a three-dimensional lattice.

  • So that's an ionic solid.

  • We have also seen metallic solid

  • where you have metals that all contribute

  • some valence electrons to the sea of electrons.

  • So what you end up having

  • is essentially these positive cations

  • that are in this sea of electrons.

  • And we've talked about those properties,

  • very good at conducting electricity, malleable, et cetera.

  • Now, what we're gonna do is talk about what happens

  • when you have nonmetals.

  • So the nonmetals, you can see in yellow right over here,

  • also includes hydrogen.

  • Now, of course, noble gases are also nonmetals,

  • but they're not reactives.

  • So we're gonna talk about the reactive nonmetals.

  • They can form molecules with each other.

  • For example, one iodine can bond

  • to another iodine with covalent bonds.

  • So you could have a molecule like I2.

  • You have things like carbon dioxide.

  • Each carbon can bond to two oxygens.

  • These are each molecules formed

  • due to covalent bonds between nonmetals.

  • Now, when we talk about molecular solids,

  • we're talking about putting a bunch of these together.

  • So let's say putting a bunch of iodine molecules together,

  • and the intermolecular forces

  • at a sufficiently low temperature

  • are sufficient to hold together those molecules as a solid.

  • So what do I mean by that?

  • Let's look at a few examples.

  • This right over here is a picture of solid iodine,

  • and the way it's made up is you have these iodine molecules.

  • Now, each of these iodine molecules are formed

  • by a covalent bond between two iodine atoms.

  • Now, the reason why it's a solid

  • is there's enough dispersion forces.

  • We talked about these London dispersion forces

  • that are formed by temporary dipoles

  • inducing dipoles in neighboring molecules.

  • For example, just by random chance,

  • for a moment, you might have more electrons

  • on this end of this iodine molecule,

  • creating a partially negative charge.

  • And then that means some of the electrons

  • on this end of this neighboring iodine molecule

  • might be repulsed by that negative charge,

  • so it forms a partially positive charge.

  • And so you have a temporary dipole

  • inducing a dipole in the neighboring molecule,

  • and then they'll be attracted to each other,

  • and we've talked about that as London dispersion forces.

  • And at a sufficiently low temperature,

  • that can keep them altogether as a solid.

  • Now, it's important to point out,

  • I keep saying sufficiently low temperature

  • because these molecular solids,

  • because they are only held together

  • not by the covalent bonds,

  • the covalent bonds hold together each of the molecules,

  • but the molecules are held together

  • by these fairly weak dispersion forces.

  • They tend to have relatively low melting points.

  • For example, solid iodine right over here

  • has a melting point,

  • has a melting point

  • of 113.7

  • degrees Celsius.

  • And I know what you're saying.

  • That's not that low.

  • That's higher than the temperature at which water boils.

  • It would be quite uncomfortable for any of us

  • to be experiencing 113.7 degrees Celsius.

  • But this is relatively low when you talk about solids.

  • Think about the temperatures

  • it requires to melt, say, table salt.

  • We've talked about that.

  • Think about the temperatures it takes to melt iron.

  • There, you're talking about hundreds of degrees,

  • in certain solids, thousands of degrees Celsius.

  • This is much lower.

  • And so as a general principle,

  • molecular solids tend to have relatively low melting points.

  • Now, how good you think they're gonna be

  • as conductors of electricity?

  • Pause the video and think about that.

  • Well, in order to be conductors of electricity,

  • somehow charge needs to move through the solid.

  • And unlike metallic solids,

  • you don't have the sea of electrons

  • that can just move around,

  • so these tend to be bad conductors of electricity.

  • If you wanna see another example of a molecular solid,

  • this right over here is solid carbon dioxide,

  • often known as dry ice.

  • What you see here is each of these molecules, each carbon,

  • is bonded to two oxygens.

  • It has a double-bond with each of those oxygens.

  • These are covalent bonds

  • that form each of these molecules.

  • But what keeps all of the molecules attracted to each other

  • is, once again, those dispersion forces.

  • And these forces between the molecules are so weak

  • that solid carbon dioxide doesn't even really melt.

  • It doesn't even go to a liquid state.

  • If you heat it up enough

  • to overcome these intermolecular forces,

  • these dispersion forces, it will sublime,

  • which means it goes directly from a solid to a gas state,

  • and it does that at a very low temperature.

  • It sublimes

  • at negative 78.5

  • degrees Celsius.

  • And if you've ever handled a dry ice,

  • which I don't recommend you doing without gloves

  • because it will hurt your skin if you do touch it,

  • I actually did that recently at my son's birthday party,

  • we were playing around with dry ice,

  • you don't mess around with this thing

  • because it is so incredibly cold.

  • And at that temperature, it will go from a solid.

  • It won't even melt to a liquid state.

  • It will go straight to a gas state.

  • Now, the last thing I wanna do

  • is think about why different molecular solids

  • will have different melting points.

  • So let's compare, for example, molecular iodine

  • to molecular chlorine.

  • Each of these can form molecular solid.

  • We looked at iodine a few minutes ago.

  • Which of these would you think would form molecular solids

  • with higher melting points?

  • Pause the video and think about that.

  • Well, as we talked about it, each of these molecules,

  • they're formed by covalent bonds between two atoms,

  • and what keeps the solid together

  • are these dispersion forces.

  • In an earlier videos,

  • when we first talked about dispersion forces,

  • we talked about temporary dipoles and induced dipoles,

  • and they were likely to form between heavier atoms

  • and molecules because they have larger electron clouds

  • and are more polarizable.

  • So if you compare molecular iodine to molecular chlorine,

  • you can see that iodine is clearly made up of larger atoms

  • and is therefore a larger molecule,

  • which is more polarizable.

  • So it's larger,

  • which means it's more polarizable, generally speaking,

  • polarizable,

  • which means it has stronger,

  • generally speaking, dispersion forces,

  • stronger dispersion forces.

  • Now, just as a reminder, these dispersion forces

  • are between molecules.

  • Each molecule has a covalent bond between two iodines,

  • and then the dispersion forces are between the molecules.

  • But because it has stronger dispersion forces,

  • we would expect that a molecular solid formed by iodine

  • is gonna have a higher melting point

  • than a molecular solid formed by chlorine.

  • And I actually do have the numbers here.

  • So the melting point

  • of a molecular solid formed by iodine,

  • we've already talked about that,

  • that's 113.7 degrees Celsius,

  • while the melting point of a molecular solid

  • formed by molecular chlorine

  • has a melting point of negative 101.5 degrees Celsius,

  • which is very cold,

  • and so iodine has a higher melting point

  • because of the stronger dispersion forces.

  • Now, as I said, those dispersion forces

  • are still not that strong.

  • This is still not that high of a temperature

  • compared to melting points of other types of solids

  • we have looked at in the past.

- [Instructor] So let's talk a little bit

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B2 中高級

分子固體 (Molecular solids)

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    林宜悉 發佈於 2021 年 01 月 14 日
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