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- Let's think about what type of reaction we might be able to get going
- if we had some methane and
- some molecular chlorine.
- So if we just let this be and we didn't heat it up
- or put in any UV light into this reaction,
- pretty much nothing will happen.
- Both of these molecules are reasonably happy
- being the way they are.
- But if we were to add heat into it, if we were to start
- making all the atoms and molecules vibrate more and
- bump into each other more, or we were to add energy
- in the form of UV light, what we could start doing is
- breaking some of these chlorine-chlorine bonds.
- Out of all of the bonds here, those are the weakest.
- That would be the most susceptible to breakage.
- So let's say we were to add some heat, what would happen?
- So let's see.
- Let me draw the valence electrons of
- each of these chlorines.
- This chlorine has 1,2,3,4,5,6,7 valence electrons,
- and this chlorine over here has
- 1,2,3,4,5,6,7 valence electrons.
- Now, when you add heat to this reaction, enough for these
- guys to vibrant away from each other, for this bond to break,
- what's going to happen, and we haven't drawn an arrow like
- this just yet, but what's going to happen is that each
- of these chlorines, this bond is going to break.
- Each of these chlorines are just going to take their part
- of the bond.
- So this guy on the left, he's just
- going to take his electron.
- And notice, I draw it with this half arrow.
- It looks like a fish hook.
- It's just half an arrowhead.
- This means that this electron is just going to go back to
- this chlorine, and this other magenta electron is going to
- go back to the right chlorine, so we can draw it like that.
- If it was up to me, I would have drawn it more like this.
- I would have drawn it more like this to show that that
- electron just goes back to the chlorine, but the convention
- shows that you can show that half of the bond is going back
- to the entire atom.
- Now, after this happens, what will everything look like?
- Well, we're still going to have our methane here.
- It hasn't really reacted.
- So we still have our methane.
- Let me draw it a little bit.
- So we still have our methane here.
- And all that's happened is, because we've put energy into
- the system, we've been able to break this bond.
- The molecular chlorine has broken up into two chlorine
- atoms. So we have the one on the left over here, and then
- we have the one on the right.
- And let me draw the left's valence electrons.
- It has one, two, three, four, five, six, seven.
- I just flipped it over so that the lone electron is on the
- left-hand side right here.
- And then you have the guy on the right.
- He has one, two, three, four, five, six,
- seven valence electrons.
- Now that each of these guys have an unpaired electron,
- they're actually very, very, very reactive.
- And we actually call any molecule that has an unpaired
- electron and is very reactive a free radical.
- So both of these guys now are free radicals.
- And actually, the whole topic of this video is
- free radical reactions.
- Both of these guys are free radicals.
- And you've probably heard the word free radical before.
- In the context of nutrition, that you don't want free
- radicals running around.
- And it's the exact same idea.
- It's not necessarily chlorine that they're talking about,
- but they're talking about molecules that
- have unpaired electrons.
- They'll react with some of your cell's machinery, maybe
- even with your DNA, maybe cause mutations that might
- lead to things like cancer.
- So that's why people think you shouldn't have free radicals
- in your body.
- But as soon as we form these free radicals, in this step
- right here, where we put energy in the system to break
- this bond, we call this the initiation step.
- Let me put this.
- We used energy here.
- This was endothermic.
- We use energy.
- This right here is the initiation step.
- And what we're going to see in general with free radical
- reactions is you need some energy to get it started.
- But once it gets started, it kind of
- starts this chain reaction.
- And as one free radical reacts with something else, it
- creates another free radical, and that keeps propagating
- until really everything has reacted.
- And that's why these can be so dangerous or so bad for
- biological systems.
- So I've told you that they react a lot.
- So how will they react now?
- Well, this guy wants to form a pair with someone else.
- And maybe if he swipes by this methane in just the right way,
- with just enough energy, what will happen is he could take
- the hydrogen off of the carbon, and not just the
- proton, the entire hydrogen.
- He will form a bond with the hydrogen using the hydrogen's
- electrons, so they'll get together and
- they'll form a bond.
- The hydrogen will contribute one electron.
- Notice, I'm drawing the half-arrow again, so the
- hydrogen isn't giving away the electron to someone else.
- That would be a full arrow.
- The hydrogen is just contributing its electron to
- half of a bond.
- And then the carbon, the carbon would do the same.
- I'll do that in blue.
- So the carbon, this valence electron right here, could be
- contributed to half of a bond, and then they will bond, and
- this bond over here will break.
- And so the carbon over here on the left, this carbon over
- here will take back its electron.
- So what does it look like?
- What does everything look like after that's done?
- So our methane now, it's no longer methane.
- It is now, if you think about it-- so
- we have three hydrogens.
- It took its electron back.
- It is now a free radical.
- It now has an unpaired reactive electron.
- The hydrogen and this chlorine have bonded.
- So let me draw the chlorine.
- It has this electron right over here.
- It has the other six valence electrons: one, two, three,
- four, five, six.
- And we have the hydrogen with its pink electron that it's
- contributing to the bond.
- And so we have them bonded now.
- This chlorine is no longer a free radical, although this
- one out here is still a free radical.
- Let me copy and paste it.
- So it's hanging around.
- Copy and paste.
- And now, notice we had one free radical react, but it
- formed another free radical.
- That's why we call this a propagation step.
- So this right here is a propagation step.
- When one free radical reacts, it created
- another free radical.
- Now, what's that free radical likely to do?
- You might be tempted to say, hey, it's going to just react
- with that other chlorine, but think about it.
- These molecules, there's a gazillion of them in this
- solution, so the odds that this guy's going to react
- exactly with that other free radical is actually very low,
- especially early on in the reaction where most of the
- molecules are still either methane or molecular chlorine.
- So this guy is much more likely to bump into another
- molecular chlorine than he is to bump into one of these
- original free radicals that formed.
- So if he bumps into another molecular chlorine in just the
- right way-- so let me draw another molecular chlorine.
- So that's another molecular chlorine.
- And each of these one, two, three, four, five six, seven;
- one, two, three, four, five, six, seven.
- There is a bond here.
- If they bump in just the right way, this chlorine electron
- might get contributed, and this free unpaired electron
- will be contributed and then this CH3, I guess we could
- call it, this free radical, this carbon free radical, or
- this methyl free radical, will then form a
- bond with this chlorine.
- What's everything going to look like after that?
- Well, after that happens this is now bonded to a chlorine.
- It's now chloromethane.
- Let me draw it.
- So it's carbon, hydrogen, hydrogen, hydrogen.
- Now, it's bonded to a chlorine.
- Let me draw the electrons so we can keep track of
- everything.
- We have that magenta electron right over there.
- And then we have the chlorine with its one, two, three,
- four, five, six, seven valence electrons.
- They are now bonded.
- This is chloromethane.
- And now you have another free radical because this guy-- and
- I should have drawn it there.
- This guy, that bond was broken, so he
- gets back his electrons.
- So he's sitting over here.
- He is now a free radical.
- So this is another propagation step.
- And we still have that original free radical guy
- sitting out over here.
- So we keep forming more and more free
- radicals as this happens.
- Now, eventually we're going to start running out of methanes
- and we're going to start running out of
- the molecular chlorines.
- So they're going to be less likely to react and you're
- actually going to have more free radicals around.
- So once the concentration of free radicals gets high
- enough, then you might start to see them
- reacting with each other.
- So when the concentration of free radicals get high enough,
- you might see, instead of this step happening-- this will
- happen a long time until most of the free radicals or most
- of the non-free radicals disappear.
- But once we have a soup of mainly free radicals, you'll
- see things like this.
- You'll see the methyl free radical.
- So let me draw it like this.
- You'll see him maybe reacting with another methyl free
- radical, where they both contribute an
- electron to form a bond.
- And then, once the bond forms, you have ethane.
- I could just write as CH3, H3C.
- So you might have something like this.
- And so this type of a step where two free radicals kind
- of cancel each other out, this is a termination step because
- it's starting to lower the concentration of free radicals
- in the solution, but this is only once the concentration of
- free radicals becomes really high.
- You might also see some of the chlorines cancel out with each
- other again, so a chlorine free radical and another
- chlorine free radical.
- I'll only draw the unpaired electron.
- They can bond with each other and form
- molecular chlorine again.
- That again is a termination step.
- Or you could see something like the methyl free radical.
- Just for shorthand, I'll write it like this: H3C.
- The methyl free radical and a chlorine free radical might
- also just straight-up react and form chloromethane, And
- form H3C-Cl.
- So this will all happen once the concentration of free
- radicals gets really high.
- Now, another thing that might happen once this reaction
- proceeds, and we have a lot of the propagation steps, is that
- you might have a situation where you already have a
- chloromethane, so it looks like this.
- You already have a chloromethane.
- And once you have enough of these, it then becomes more
- likely that some free radical chlorine might be able to
- react with this thing, so it might actually add another
- chlorine to this molecule.
- And the way it would do it, this chlorine over here-- I'm
- just drawing the free electron pairs.
- It would form a bond with this hydrogen right over there.
- They would both contribute their electrons.
- And then the carbon would take back its electron.
- Notice, all of the half-arrows.
- You'd be left with-- the hydrogen and the chlorine
- would have bonded.
- And now, this guy's going to be a free radical, but he's
- going to be a chlorinated free radical.
- So it's going to look like this.
- He has a free electron over there: hydrogen, hydrogen.
- And then he might be able to react with
- another chlorine molecule.
- He contributes an electron.
- Maybe this guy contributes an electron.
- This guy-- I don't want to draw a full arrow-- he
- contributes an electron to a bond, and then this guy takes
- his electron back and becomes a free radical.
- And then we're left with what?
- We're left with a doubly chlorinated methane.
- So then we have Cl, Cl, and then a
- hydrogen and a hydrogen.
- And this could actually keep happening.
- As the concentration of these get higher, then it becomes
- more likely that this can react with another chlorine.
- Of course, this chlorine over here becomes
- another free radical.
- But the general idea here that I wanted to show you is that
- once a free radical reaction starts-- the first step
- requires some energy to break this chlorine-chlorine bond,
- but once it happens, these guys are highly reactive, will
- start reacting with other things, and as they react with
- other things, it causes more and more free radicals, so it
- starts this chain reaction.
- And actually, all in all, this required energy to occur.
- This step right here, this propagation step, it requires
- a little bit of energy, but it's almost neutral.
- It requires energy to break this bond, but it creates
- energy when this bond is formed.
- It still requires a little net energy.
- And then things like this start to become exothermic.
- And especially once you start getting to the termination
- steps, you start releasing a lot of energy.
- So actually, all in all, this reaction is actually going to
- release energy, but it needed some energy to get started.