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  • - [Instructor] Let's talk a little bit about ionic solids,

  • which you can imagine are solids formed by ions.

  • So, let's think a little bit about these ions.

  • So, for example, we could look at group one elements here,

  • especially things like lithium, or sodium, or potassium.

  • And in many other videos we have talked

  • about these elements wanting maybe not so much

  • to keep their outermost electron

  • because they only have one electron

  • in their outermost shell.

  • And it'd be pretty easy for them to loose that electron

  • to get to a noble gas configuration

  • to have a full outer shell.

  • So, these characters like to lose one electron

  • the group two elements like to lose two electrons.

  • While if you go on the other side of the periodic table,

  • if you look at the halogens right over here,

  • they're one electron away

  • from having a noble gas electron configuration,

  • from having a full outer shell.

  • So, they really like to grab electrons.

  • And, if you look at elements like oxygen and sulfur,

  • they really like to grab two electrons, if they can.

  • So, what do you think happens if you have some metals

  • on the left end here mixed with some nonmetals

  • on the right end here?

  • Well, you might imagine there would be a reaction.

  • So, for example, if you mixed sodium with chlorine,

  • the sodiums might lose an electron to the chlorines,

  • in which case you're going to have sodium cations,

  • positively charged ions.

  • And if the chlorines are now taking those electrons,

  • they then become chloride anions.

  • And now if you have a bunch of positive ions

  • hanging around a bunch of negative ions,

  • what do you think is going to happen?

  • They're going to get attracted to each other.

  • And they're going to get attracted to each other

  • and form a lattice structure, like this.

  • I like to use sodium chloride as an example

  • because this is probably the one

  • that we see most in our life, this is table salt.

  • If you were to lick it, it'd taste salty.

  • But, there's many other ionic salt solids,

  • many of them would actually be categorized

  • as salt, generally.

  • You could have a potassium chloride.

  • You could have a sodium chloride.

  • You could have, for example, a magnesium oxide.

  • What's going on there?

  • Well, in that situation, the magnesium, each magnesium

  • might lose two electrons,

  • so they become a ion with a positive two charge,

  • and each of the oxygens would gain two electrons.

  • So then they are anions with a negative two charge.

  • And these characters once again are going to be attracted

  • to each other and form an ionic solid in a regular

  • lattice structure like this.

  • So let's think a little bit about their properties.

  • So first of all, let's think about the melting points.

  • So, these solids, the electrostatic attraction between

  • these ions is strong.

  • And so they tend to have high melting points.

  • Now what if we were to compare melting points

  • between ionic solids?

  • So for example, if you wanted to compare the melting point

  • of sodium chloride to the melting point

  • of magnesium oxide which one do you think

  • has a higher melting point?

  • Pause this video and think about it.

  • Well, as you can imagine the electrostatic attraction,

  • it's going to be dependent on two things.

  • The magnitude of the charge and the radius

  • of the atoms that ae forming this lattice structure.

  • And the magnitude of the charge here is clear.

  • Here you have a plus two charge being attracted

  • to a negative two charge so this has a stronger

  • electrostatic attraction and so you're going to have

  • a higher melting point right over here.

  • The melting point of magnesium oxide?

  • 2,825 degrees Celsius, while the melting point of table salt

  • or sodium chloride is 801 degrees Celsius.

  • You could also try to compare sodium chloride

  • to something like sodium fluoride.

  • Which one do you think is going to

  • have a higher melting point?

  • Sodium chloride or sodium fluoride?

  • Well fluorines are smaller than chlorines and each of them

  • gain an electron, then the fluoride anion

  • is still going to be a reasonable bit smaller

  • than the chloride anion.

  • Or when you have smaller constituent ions,

  • the electrostatic attraction is actually stronger.

  • Remember, we've seen in Coulomb's law,

  • that the closer two charges are to each other,

  • the stronger the attractive or the repulsive force,

  • and if they're opposite charges,

  • it's going to be an attractive force.

  • So, sodium fluoride is actually gonna have a higher

  • melting point than sodium chloride, by a little bit.

  • It actually turns out that the melting point

  • of sodium fluoride is 996 degrees Celsius.

  • But if you're comparing these three,

  • the highest melting point is magnesium oxide,

  • followed by sodium fluoride, followed by sodium chloride.

  • So charge is what's really dominating over here.

  • Now the next question you might be wondering

  • is all right I can imagine these solids are really hard,

  • but what would happen if I were to try to break it?

  • Would it bend like a lot of the metals we know

  • and we'll study that in other videos,

  • or would something else happen?

  • And to understand that, let me draw a two dimensional

  • representation of this.

  • So let me draw the chlorine,

  • or I should say the chloride anions.

  • And this is just a two dimensional version of that lattice.

  • Obviously not drawing it to scale.

  • And then let me draw the sodiums.

  • Sodium cations.

  • As you can see, the positives are attracted to the negative,

  • that's why they're next to each other,

  • the negatives aren't next each other

  • because they repel each other.

  • The positives aren't next to each other,

  • but what would happen if I were to try to,

  • or I were to press down really hard on this side

  • and if I were to press really hard up on this side?

  • So what would happen if I press had enough

  • that this side begins to budge?

  • So it begins to budge.

  • Would it just bend, or what do you think's gonna happen

  • when I get right about there?

  • Well, when I get right about there,

  • all of a sudden I've, not only have I broken the lattice,

  • but the negatives are next to the negatives

  • and the positives are next to the positives

  • and so it's not just going to bend,

  • and be malleable like a lot of the metals we've seen,

  • it's just going to break.

  • So this is going to be, even though it's going to be hard,

  • it is going to be brittle.

  • Now the last question we'll address in this video

  • is how good do you think ionic solids conduct electricity?

  • Pause this video and think about that.

  • Well, in order to conduct electricity,

  • either electrons or charge generally has to be able

  • to move about.

  • And when it's just in its solid form like this,

  • the, even though you do have these ions

  • they're not going to move about.

  • So ionic solids in their solid form,

  • they aren't good at conducting electricity.

  • They can be good at conducting electricity

  • if you were to dissolve it in a solution.

  • For example, if you were to dissolve this salt

  • in water, now the ions can move around

  • and then they're good at conducting electricity.

  • Or, if you were to heat this sodium chloride

  • up beyond 801 degrees Celsius and it turns into a liquid,

  • then once again the ions can move around

  • and you can actually conduct electricity.

  • Take everything I say with a grain of salt.

  • Sorry, I know, I couldn't help it.

  • But hopefully you know a little bit more

  • about ionic solids now.

- [Instructor] Let's talk a little bit about ionic solids,

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B2 中高級

離子型固體 (Ionic solids)

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    林宜悉 發佈於 2021 年 01 月 14 日
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