Chapter 10: Radical Reactions

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Did you ever wonder how some chemicals, you know, the ones in fancy fire extinguishers, can put out fires way better than just water?

I mean, fire itself, that combustion reaction, it's actually a chemical process, a chain reaction, really.

And it involves these super reactive little things called free radical intermediates.

Fleeting, but powerful.

Okay, so let's unpack that a bit.

Today, we're doing a deep dive into this whole world of radicals in organic chemistry.

And our main source, our guide, is a really solid chapter from David Klein's Organic Chemistry, the third edition.

Exactly.

And our goal here for you is to really get a feel for these radicals, not just the definition, but their structure, how they react, why they're so reactive sometimes.

We'll look at where they pop up industry, food, even in our own bodies.

And yeah, we'll definitely circle back to that fire question.

We'll see how certain chemicals can actually hijack that radical chain reaction into fire.

So this deep dive, it's about clarifying the key stuff, the mechanisms, the real world side of it, so you can quickly grasp what's, well, what's really important.

All right, let's jump in, starting right at the beginning.

What is a radical?

We usually think of bonds breaking unevenly, right?

Heterolytic cleavage makes ions.

One atom gets both electrons, but radicals come from the other way, homolytic cleavage.

The bond splits like perfectly evenly.

Each atom walks away with one electron.

Precisely.

And that single unpaired electron is the defining feature.

Now visually, this is crucial.

For ionic reactions, you're used to those double barbed arrows showing two electrons moving.

Forget those for radicals.

Here we use fishhook arrows, single barb, just like a tiny fishhook.

It signals the movement of one single electron.

It seems small, but it's fundamental.

It tells you immediately you're in radical territory.

Okay, fishhook arrows, single electron, got it.

So what do these things, these carbon radicals actually look like, shape -wise?

Well think back to carbocations.

They're flat, right?

Yeah.

SP2, trigonal planar.

Yeah.

And carbanions are more pyramidal.

It's P3.

Carbon radicals, they're kind of in between.

Experimentally, they look pretty much trigonal planar, or maybe a very, very shallow pyramid that's flipping back and forth incredibly fast.

So fast, it basically averages out to being planar.

So for practical purposes, we treat them as flat.

Yeah, for most situations in organic chemistry, thinking of them as trigonal planar works perfectly.

And that flatness, it has big implications for stereochemistry later on, especially if you're forming a new chiral center.

Okay, interesting.

Now,

just like we saw with carbocations, not all radicals are equally stable, right?

Some are more chill, others are desperate to react.

Absolutely.

There's a definite pecking order.

And it follows the same trend as carbocations, which is helpful.

Tertiary radicals are the most stable, then secondary, then primary, and finally methyl radicals are the least stable.

And is the reason the same, that hyperconjugation thing?

Exactly right, hyperconjugation.

Those neighboring alkyl groups help to spread out or delocalize that single unpaired electron just a little bit.

Stabilizes it.

We can actually see this in the numbers using bond dissociation energies, or BDEs.

Right, BDEs tell you how strong a bond is, how much energy to break it homilitically.

Yep, lower BDE means an easier break, which means the radicals formed are more stable.

So when you see that a tertiary CH bond has the lowest BDE, around 381 kilojoule, that confirms it forms the most stable radical.

Okay, tertiary is good, but is there anything even better, a way to make a radical super stable?

Oh yeah, resonance.

That's the real game changer.

If you have that unpaired electron right next door to a pi bond, think of lollik positions or bedensilik positions, you get resonance stabilization.

Ah, so the electron can be delocalized over multiple atoms.

Precisely.

And you draw this using three Fischach arrows, moving the single electron and rearranging the pi bond.

This sharing of the electron load makes these radicals incredibly stable.

Look at the BDEs again.

A benzilik CH bond is down around 356 kilojoule.

That's significantly more stable than even a tertiary radical.

Wow.

Now, here's a super common mistake, something you really need to watch out for.

Don't confuse allylic radicals, which are resonance stabilized, with vinylik radicals.

Phenylik is where the radical is on one of the double bond carbons.

Exactly.

And those are not resonance stabilized.

In fact, they're really unstable.

Less stable than even primary radicals.

Their BDE is way up at 464 kilojoule.

Big difference.

Easy mistake to make.

Good tip.

So practically speaking, knowing this stability order helps you pinpoint the weakest CH bond in a molecule, the one most likely to break and form a radical.

That's the key takeaway.

Find the spot that forms the most stable radical tertiary.

Or even better, allylic or benzilik.

And that's usually where the initial hydrogen abstraction will happen in a radical reaction.

Makes predicting reactions much easier.

Alright, we know what they are, how stable they are, how do they actually react, what are the moves they make?

You mentioned they have their own patterns, different from ionic ones, and they don't usually rearrange like carplications.

That's right.

Rearrangements are rare for radicals.

They have six main moves or patterns.

Let's quickly list them.

One,

homolytic cleavage.

That's how you make radicals, breaking a bond.

Usually needs energy, like heat or UV light.

Uses two fishhook arrows.

Okay, creation.

Two, addition to a bond, a radical attacks a double or triple bond, forms a new radical.

Three fishhooks.

Three,

hydrogen abstraction, a radical grabs a hydrogen atom, key point.

The whole atom, electron included from another molecule, makes a new radical.

Three fishhooks again.

Not a proton, an H atom.

Got it.

Four,

halogen abstraction, similar idea.

Radical grabs the halogen atom, like Cl or Br.

New radical formed.

Three fishhooks.

Five, elimination.

This is kind of the reverse of addition.

A radical causes a bond to break nearby, often forming a double bond and fragmenting the molecule.

Three fishhooks.

The beta position cleaves.

Yeah, the bond beta to the radical cleaves forms a pi bond between alpha and beta.

And finally, six, coupling.

Two radicals just meet up and form a bond.

Radicals are destroyed.

Two fishhooks.

Creation and destruction.

It's interesting how some are opposites, like cleavage and coupling, or addition and elimination.

It really highlights the dynamic nature of these reactions.

And these six steps fit into the three main stages of a radical chain reaction.

You've got initiation, making the first radicals, usually homolytic cleavage.

You've got termination destroying radicals, usually coupling.

And then the most important part, propagation.

This is where the chemistry happens.

One radical reacts to form a product, and another radical.

That new radical keeps the chain going.

And crucially, if you add up all the propagation steps, you get the overall net reaction.

That's the chain part.

One radical triggers potentially thousands of reactions.

That's exactly it.

It's incredibly efficient, which is why radical reactions can be so powerful and sometimes hard to control.

Let's make this real.

The classic example.

Chlorinating methane.

CH4 plus Cl2 makes methyl chloride, CH3Cl.

How does that work with these steps?

Okay, so first, initiation.

You need UV light or heat to break the ClCl bond.

Homolytic cleavage gives you two chlorine radicals, CLA.

Step one, make radicals.

Then propagation.

This has two steps that repeat.

Step one.

A chlorine radical bumps into methane and does a hydrogen abstraction.

It grabs an H atom, forming HCl, and leaving behind a methyl radical, CH3.

Step two.

That methyl radical bumps into a different Cl2 molecule and does a halogen abstraction.

It grabs a chlorine atom, forming our product, methyl chloride, CH3Cl.

And here's the key, regenerating a chlorine radical, CLA.

And that new Cl radical goes back to step one of propagation.

Bingo.

That's the chain.

It cycles round and round.

Step one.

Step two.

Step one.

Step two.

Making product each time.

Until?

Until termination.

Eventually, two radicals find each other instead of finding a reactant molecule.

Maybe two Cl radicals couple to make Cl2, or two CH3 radicals make ethane, or CH3 in the CLA and make methyl chloride.

Any coupling event removes radicals and slows or stops the chain.

Now you mentioned control issues.

With methane chlorination, you can get too much chlorination, right?

Dichloromethane, chloroform.

Yeah, polychlorination.

If you have lots of Cl2 around, the methyl chloride product can react further with another Cl radical via hydrogen abstraction, and the whole process repeats.

So how do you stop that?

How do you favor just adding one chlorine?

The trick is stoichiometry.

Use a large excess of methane relative to chlorine.

That way, a chlorine radical is statistically much more likely to bump into a fresh methane molecule than an already chlorinated product molecule.

Simple, but effective.

Makes sense.

Now, you need energy to start these.

You said UV or heat.

Are there chemical ways to kick them off?

Initiators?

Definitely.

Radical initiators are compounds designed with deliberately weak bonds that break easily with just a bit of heat to form radicals.

Alkyl peroxides with that ROR structure have a weak OO bond.

Aesoperoxides are even better.

Why are they better?

Their OO bond is incredibly weak, only about 121 kilojmol.

Plus, when it breaks, the radicals formed are resonance stabilized, making the initiation even easier.

They're very common initiators.

And the flip side.

Things that stop radical reactions?

Inhibitors.

Right, radical inhibitors or scavengers.

They intercept and destroy radicals to break the chain.

Molecular oxygen O2 is actually a diradical itself, which is pretty unusual, and it can with other radicals.

A more classic lab example is hydroquinone.

It has easily abstractable hydrogens.

It reacts with a radical, forms its own very stable resonance stabilized radical.

This new radical is much less reactive, or it can even couple with another radical, effectively killing the chain reaction.

Okay, so we can start them, stop them.

What about the choice of halogen itself?

You said chlorination and bromination work, but iodination doesn't, and fluorination is too crazy.

Why?

It boils down to thermodynamics,

specifically the enthalpy change for the overall reaction.

We look at the energy you need to break bonds versus the energy released when new bonds form.

Fluorination releases a huge amount of energy, negative 431 kilojoule is just too exothermic, too violent.

Like an explosion waiting to happen.

Pretty much.

Chlorination, negative 104 kilojoule, and bromination, negative 33 kilojoule, are nicely exothermic, making them favorable and controllable.

But iodination, it's actually endothermic, plus 55 kilojoule, it requires energy input overall so it just doesn't happen spontaneously under these conditions.

The equilibrium lies way over on the reactant side.

And you also mentioned bromination is slower than chlorination.

Yes.

Even though both are exothermic overall, the first propagation step for bromination, hydrogen abstraction by PROM, is actually slightly endothermic.

For chlorination, that first step is exothermic.

An endothermic step means a higher activation energy barrier to overcome.

So bromination just proceeds more slowly.

Think climbing a small hill versus going downhill.

Okay, that brings us to a really important practical point.

Selectivity.

You gave the propane example, chlorination gives a mix, but bromination heavily favors the secondary position.

Why such a big difference?

This is where the Hammond postulate comes in handy.

It connects the transition state structure to the energy of the step.

For chlorination, the rate determining hydrogen abstraction is exothermic.

Fast.

The Hammond postulate says the transition state will look more like the reactants.

At that point, the CH bond is only slightly broken.

The carbon has very little radical character.

So the energy difference between forming a primary, secondary, or tertiary radical transition state isn't huge.

Chlorine is reactive.

Not very picky.

It just grabs whatever hydrogen is available.

More or less, yeah.

There's some preference for the weaker bond, but it's not dramatic.

Now for bromination, that rate determining hydrogen abstraction is endothermic.

Slower.

The transition state looks more like the products, meaning the radical intermediate.

So in the transition state, the CH bond is almost fully broken and the carbon has significant radical character.

This makes the transition state energy very sensitive to the stability of the radical being formed.

Ah, so because the tertiary radical is much more stable than primary, the transition state leading to it is much lower in energy for bromination.

Exactly.

The energy gap between the transition states for forming primary versus secondary versus tertiary is much larger for bromination.

This leads to high selectivity.

Bromine is far more selective, waiting to abstract the hydrogen that leads to the most stable radical intermediate.

The numbers were pretty stark, like 1 ,600 times more selective for tertiary over primary for bromine.

Something like that, yeah, compared to only about five times for chlorine.

It's a massive difference.

So the practical advice is,

if you have a molecule with different kinds of hydrogens and you want to put a halogen on a specific spot, usually the most substituted, use bromination, not chlorination.

Precisely.

Bromination gives you control over the regiochemistry, avoids messy product mixtures.

Okay, what about stereochemistry?

We said the radical is basically planar.

What happens if we form a chiral center or react at one?

Right, that planar or rapidly inverting nature is key.

If the halogenation reaction creates a new chiral center where there wasn't one before.

Like halogenating butane at C2.

Exactly.

The incoming halogen atom can attack either face of that planar radical intermediate with equal probability.

So you get a 50 .50 mixture of the RNS and antiomers, a racemic mixture.

And if you start with a chiral molecule and the reaction happens at the existing chiral center.

Same outcome.

The reaction goes through that planar radical intermediate, which erases the original stereochemistry Again,

you end up with a racemic mixture.

The original configuration is lost.

What if there's another chiral center nearby that isn't involved in the reaction?

Now it gets more complex.

If you form a new chiral center, or react at one, but there's another pre -existing untouched chiral center elsewhere in the molecule, you'll form diastereomers.

And diastereomers have different physical properties.

Are they formed in equal amounts?

Generally no.

That nearby chiral center makes the two faces of the planar radical intermediate non -equivalent.

The approach of the halogen might be slightly hindered on one side compared to the other.

So you typically get an unequal mixture of diastereomers.

Okay, that's a lot on alkenes.

What if we try to halogenate an alkene?

Where does the radical attack?

Good question.

In an alkene, the weakest CH bonds are usually the ones at the allylic position, the carbon, right next to the double bond.

Because removing that H gives a resonance stabilized allylic radical.

You got it.

That allylic radical is nice and stable, so you'd expect hydrogen abstraction to happen there preferentially, allylic halogenation.

But there's a catch.

If you just throw in Br2 and light with an alkene,

you get a competing reaction.

Br2 can add across the double bond via an ionic mechanism.

That's often the major pathway, giving you the dobromol -halkene, not the allylic bromide you wanted.

Messy.

So how do chemists get around that problem?

Clever reagent design.

They use N -bromosacinamide or NBS.

NBS.

What's special about it?

NBS provides a source of bromine.

But it does so in a way that keeps the concentration of actual Br2 molecule extremely low throughout the reaction.

Basically, a tiny amount of HBr is formed during the radical process.

And that HBr reacts with NBS to generate just enough Br2 to keep the radical chain going, but not enough for the ionic addition across the double bond to compete effectively.

So NBS favors the radical pathway by starving the ionic pathway of Br2.

Exactly.

It's a neat trick.

Though, be aware, if the intermediate allylic radical has multiple resonance forms, NBS can still sometimes lead to a mixture of allylic bromides, depending on where the bromine ends up attaching.

Wow.

Okay, let's zoom out a bit.

These radicals aren't just lab tools.

You mentioned they have big real -world impacts, like in the atmosphere.

Huge impacts.

Think about the ozone layer in the stratosphere.

Ozone O3 is constantly being formed and broken down up there, and it absorbs most of the sun's harmful UVB radiation, the natural sunscreen for the planet.

And this involves radicals.

Oh yes.

The absorption of UV light by ozone causes it to break apart, often involving radical species.

It's a dynamic equilibrium.

But then, humans introduced chlorofluorocarbons, CFCs, freons, used in old refrigerators and spray cans.

These things are super stable down here.

But not in the stratosphere.

Right.

Up there, intense UV light can break their strong CCL bonds, homolytic cleavage, again generating chlorine radicals.

Cielo.

And these chlorine radicals are the villains in the ozone story.

They are devastatingly effective villains.

A single chlorine radical can catalyze the destruction of thousands of ozone molecules through a chain reaction.

Kelo reacts with O3 to make keo at an O2, then the kelo reacts with a free oxygen atom, which are also present up there, to regenerate CO and make another O2.

So the chlorine radical gets recycled, ready to kill more ozone.

Precisely.

It's a catalytic cycle.

That's what lead to the ozone hole, especially over Antarctica.

Which prompted the Montreal Protocol, the international treaty, to phase out CFCs.

Yes, a major environmental success story, really.

We switched to alternatives like hydrofluoroalkanes, HFAs.

Because they don't have the chlorine.

They have CF bonds instead.

Exactly.

CF bonds are much stronger and don't break easily under UV to form radicals.

So HFAs don't deplete the ozone layer.

But, there's always a but, isn't there?

Well, yeah.

HFAs, while ozone -safe, turned out to be potent greenhouse gases.

So now the search is on for third -generation replacements that are safe for ozone and don't contribute significantly to climate change.

It's an ongoing challenge.

Let's connect back to the very beginning fire extinguishers.

You said combustion is a radical chain reaction, and some chemicals are better than water.

How do they work?

Right.

Fire needs fuel, oxygen, and heat, the ignition source.

Standard extinguishers attack one of those.

CO2 smothers the fire, removing oxygen.

Water cools it, removing heat.

But the most effective ones, historically, were Halen's compounds containing bromine and sometimes chlorine and fluorine, like certain CFCs or BFCs.

Why were they so good?

They attack the fire chemically, on top of physically.

First, they're heavy gases, so they can displace oxygen.

Second, they absorb heat, as their CBR or CCL bonds break homiletically.

But the killer move is that the bromine, or chlorine, radicals produced then interfere with the radical chain reactions of the combustion itself.

They act as radical scavengers, coupling with the radicals that propagate the fire, effectively terminating the flame chemistry.

So they interrupt the fire's own radical cycle.

Exactly.

It's a very efficient way to extinguish a fire, especially in places where water or powder would cause damage, like around electronics.

But again, the ozone problem.

Halen's containing chlorine or bromine were phased out under the Montreal Protocol, too, right?

Correct.

Their production is banned, though existing stockpiles are sometimes maintained for critical uses like aircraft engine fire suppression, where alternatives aren't yet as effective.

Finding replacements with the same chemical firefighting power, but without the environmental downsides, is tough.

Okay, shifting gears slightly.

What about everyday things, like food going bad?

Is that radicals, too?

That rancid smell?

Yep, that's often auto -oxidation.

It's the slow reaction of organic compounds, especially those with CH bonds, next to double bonds, allelic, or benzene rings, benzilic, with atmospheric oxygen.

Oxygen itself.

Yeah, remember O2 is a diradical.

It can initiate radical chains, leading to the formation of hydroperoxides, ROH.

These can break down further into aldehydes and ketones, which often have unpleasant smells and tastes.

That's rancidity.

And this happens in our bodies, too.

Oxidative stress.

It does.

Similar processes can damage lipids, proteins, DNA.

That's why antioxidants are important, both in food preservation and in our biology.

How do antioxidants work, like BHT and BHA added to food?

They are radical inhibitors.

They typically have a reactive hydrogen atom, often on an OH or NH group, that they donate readily to a reactive radical in the chain reaction.

This stops the original damaging radical, and the antioxidant itself forms a new radical.

But the key is, the antioxidant radical is usually much more stable, often resonance stabilized,

and less reactive, so it doesn't effectively continue the chain.

It breaks the cycle.

Like vitamin E and vitamin C in our bodies.

Exactly.

Vitamin E is lipid -soluble, protecting cell membranes.

Vitamin C is water -soluble, working in the bloodstream and cytoplasm.

They sacrifice themselves to neutralize more dangerous radicals.

There's a striking medical example related to this, isn't there?

Tylenol overdose.

Yes, acetaminophen, Tylenol overdose.

Normally the liver metabolizes small amounts safely.

But in an overdose, a toxic metabolite is formed.

Our liver uses a natural antioxidant called glutathione to neutralize this metabolite.

Glutathione has a phial group, TGSH, with an easily abstracted hydrogen, acting as a radical scavenger.

But in an overdose, you run out of glutathione.

Precisely.

The glutathione stores get depleted.

Then the toxic metabolite, which likely acts via radical mechanisms, runs rampant and causes severe liver damage, potentially fatal.

And the antidote.

Is N -acetylcysteine or NFC?

NFC replenishes the body's supply of cysteine, which is a building block a liver needs to synthesize more glutathione.

By restoring glutathione levels, NAC allows the liver to safely neutralize the toxin.

It's a direct intervention based on understanding radical scavenging.

Fascinating.

So radicals cause problems, but chemists also use them deliberately as tools, right?

Oh, definitely.

We've talked about halogenation.

Another key one is the radical addition of HBr to alkenes.

You mentioned this gives the anti -Markovnikov product bromine on the less substituted carbon.

Yes.

If you add HBr to an alkene in the presence of peroxides, which act as radical initiators, the mechanism changes from ionic to radical.

In the radical mechanism, the bromine radical adds first to the double bond in a way that forms the more stable carbon radical intermediate, tertiary or secondary rather than primary.

Then that carbon radical abstracts hydrogen from HBr.

The net result is Brnaur on the less substituted carbon.

The opposite of normal HBr addition.

Does this work with HCl or HI2?

No.

Interestingly, it only works well for HBr.

The thermodynamics just don't line up favorably for the propagation steps with HCl or HI under radical conditions.

HBr is the sweet spot.

What about making plastics?

Polymers?

Radical polymerization is huge, making polymers like polyethylene from ethylene, PVC from vinyl chloride, teflon from tetrafluoroethylene, polystyrene.

Many common plastics are made this way.

An initiator creates radicals, which add to the double bond of a monomer molecule.

This creates a new radical, which adds to another monomer, and so on, building up a long chain.

And you mentioned branching.

Yeah, sometimes side reactions can cause the polymer chain to branch out.

Controlling the amount of branching affects the material's properties, like the difference between flexible LDPE plastic film and rigid HDPE plastic bottles.

And in the giant petrochemical industry.

Radicals are everywhere there too.

Processes like breaking down large, less useful alkane molecules from crude oil into smaller or valuable alkanes, and alkenes for gasoline relies heavily on high temperature radical reactions.

Reforming takes straight chain alkanes and uses catalysts, often involving radical intermediates, to make branched alkanes, which have better combustion properties and engines.

Hydrocracking combines cracking with hydrogenation, also using radical pathways.

It's fundamental to fuel production.

OK, so wrapping it all up for synthesis.

If you're in the lab wanting to use radical halogenation, what's the bottom line?

It's a useful tool, mainly for introducing a functional group onto an otherwise unreactive alkane.

That first halogenation can be the entry point to a whole range of other reactions.

The key decision is chlorination versus bromination.

Use chlorination if your starting material is symmetrical, meaning all the hydrogens are equivalent, or if you don't mind getting a mixture, it's faster, but unselective.

Use bromination if you have different types of hydrogens, primary, secondary, tertiary, and you want to selectively functionalize the position that forms the most stable, radical, usually tertiary, secondary, primary.

Its selectivity is its major advantage for controlled synthesis.

So quite a journey we've taken here, from thinking about fire, to tiny fissure arrows, ozone layers, food spoilage, and back to practical chemistry in the lab.

It really shows how understanding these fundamentals,

radical structure, stability, the patterns they follow, unlocks so much.

It's not just theory.

It explains industrial processes, biological functions, environmental issues.

It lets you predict what might happen, maybe control it, maybe even use it constructively.

These fleeting intermediates are powerful players.

Makes you wonder, doesn't it?

How many other everyday things that seem simple on the surface are actually driven by these hidden complex chemical dances, often involving this silent, super -fast world of radicals?

Something to think about.