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• - [Instructor] In this video, we're gonna start talking

• about exceptions to the octet rule,

• which we've talked about in many other videos.

• The octet rule is this notion

• that atoms tend to react in ways

• that they're able to have a full outer shell.

• They're able to have eight valence electrons.

• things like hydrogen.

• Its outer shell is that first shell

• which gets full with two electrons.

• So it's trying to get to that duet rule.

• But as we'll see, there are other exceptions.

• Boron and aluminum, for example,

• they can form stable molecules where the boron

• or the aluminum only have six valence electrons, not eight.

• And there are exceptions in the other direction.

• As you get to the third period and beyond,

• we'll actually see atoms that can maintain more

• than eight valence electrons.

• And we're actually going to see an example

• of that with xenon.

• So let's just go into a few examples.

• Given what I've told you,

• see if you can come up with the Lewis diagram

• for aluminum hydride.

• So aluminum hydride has one aluminum

• and three hydrogens.

• See if you can draw the Lewis diagram for that.

• All right, now let's do this together.

• So the first thing you wanna do is account

• for all of the valence electrons.

• Aluminum's outer shell is the third shell,

• so the third period here,

• and it has one, two, three valence electrons.

• And then we have three hydrogens,

• and each hydrogen has one valence electron.

• And so you add all of this up together,

• three plus three is equal to six valence electrons

• in aluminum hydride.

• Now, the next step after that is

• to try to draw the structure with some covalent bonds.

• We don't wanna make hydrogen our central atom.

• That would be very atypical.

• And so let's put aluminum in the center.

• And then we're gonna have three hydrogens.

• So one, two, and three.

• And then let's put some covalent bonds in here.

• And so let's see, how many valence electrons

• have we now accounted for?

• This is two in this covalent bond.

• Another two gets us to four.

• Another two gets us to six.

• So we have just accounted for all six valence electrons.

• So we have no more valence electrons to play with.

• Let's think about how the various atoms are doing.

• So the hydrogens are all meeting their duet rule.

• These two electrons in this bond are hanging around hydrogen

• and around the aluminum.

• But from hydrogen's point of view, it has a full duet,

• that hydrogen as well and that hydrogen as well.

• But notice the aluminum over here,

• it has two, four, six electrons,

• valence electrons around it, and so it's not a full octet.

• But aluminum hydride is actually something

• that has been observed.

• Let's think about another example.

• Let's think about xenon pentafluoride.

• Xenon pentafluoride cations,

• a positively charged ion here.

• Pause this video and see

• if you can draw the Lewis diagram for this.

• All right, now let's do this together.

• If any of this seems unfamiliar,

• I encourage you to watch the video

• on introduction to drawing Lewis diagrams.

• But what we'd wanna do is first think

• So xenon right over here, it's actually a noble gas.

• It already has a full octet in its outer shell,

• so it has eight valence electrons.

• So xenon has eight valence electrons.

• And then fluorine, we've seen this multiple times,

• has one, two, three, four, five, six,

• seven valence electrons, but there's five of them.

• So five times seven.

• I'm gonna be drawing a lot of electrons in this.

• So this gives us a total of eight plus 35,

• which is 43 valence electrons.

• But we have to be careful.

• This is a cation.

• It is a positively charged molecule.

• It has a positive one charge.

• So we have to take one electron away because of that.

• So let's take away one valence electron to get that cation.

• And so we are left with 42.

• 42 valence electrons.

• So the next step is to try to draw its structure

• with some basic single covalent bonds.

• And xenon would be our preferred central atom

• because fluorine is more electronegative.

• It's actually the most electronegative element.

• So let's put xenon in the middle,

• and then let's put some fluorines around it,

• five of them to be specific.

• So one, two,

• three, four.

• I'm having trouble writing an F.

• Four and then five fluorines.

• And now let me make five covalent bonds.

• One, two, three, four, five.

• So just like that, I have accounted for 10 valence electrons

• because you have two valence electrons

• in each of these covalent bonds,

• two, four, six, eight, 10.

• So let me subtract 10 valence electrons.

• And then we are left with 32 valence electrons.

• Now, the next step is to try to allocate some more

• of these valence electrons to the terminal atom

• so that they get to a full octet.

• So let me do that to the fluorines.

• Each of these fluorines already are participating

• in a covalent bond, so they already have

• two valence electrons hanging out with them,

• so let's give 'em each six more.

• So let's give that fluorine six,

• and that fluorine gets six,

• and that fluorine gets six valence electrons,

• and that fluorine gets six valence electrons,

• and then last but not least,

• this fluorine gets six valence electrons.

• So I have just given away six valence electrons

• to each of five fluorine atoms.

• So that is 30 valence electrons that I have just allocated.

• And then what does that leave me with?

• That leaves me with two valence electrons

• that have gone unallocated.

• And the only place to now put them is on the xenon.

• And as I said, things that are lower down

• in that periodic table of elements,

• especially as we get below the third period,

• these can defy the octet rule.

• Xenon already has 10 valence electrons,

• and I'm about to allocate it two more to it just like that.

• So you allocate those two more.

• And then we have allocated all of our valence electrons.

• And I wanna make sure I remind myself and everyone

• that this is a cation.

• So I have to put that plus charge just like this,

• but this is something that has been observed

• where you can actually have a central atom like this

• that goes beyond an octet number of valence electrons.

• In this case, it has two, four, six,

• eight, 10, 12 valence electrons.

• Now, an interesting question is how do these atoms

• that are in the third period or beyond handle more

• than eight valence electrons?

• And it is a matter of debate,

• but some chemists believe that it's possible

• because they're able to place their electrons

• in their empty valence d-orbitals.

• But once again, this is controversial

• in the chemistry community.

- [Instructor] In this video, we're gonna start talking

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# Exceptions to the octet rule | AP Chemistry | Khan Academy

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林宜悉 posted on 2020/03/31
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