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Hi! I’m Deboki Chakravarti and welcome to Crash Course Organic Chemistry!
Antioxidants have gotten a lot of hype as a superfood,
possibly helping us fight illnesses like heart disease and cancer.
On a chemical level, antioxidants are fairly straightforward:
they react with radicals, which are molecules with a single unpaired electron.
We need radicals to stay alive.
They combine with oxygen as part of our normal metabolism.
But sometimes, these radical reactions produce too many reactive oxygen species, or ROS,
like the superoxide radical.
It's not great to have too many chaotic ROS around, so our bodies have ways to control them,
like an enzyme that turns superoxide back into oxygen.
But sometimes, along the way, a more dangerous ROS is produced.
The hydroxyl radical can damage cell membranes, proteins, and DNA.
When our enzymes are overwhelmed, antioxidants can come to the rescue,
giving up single electrons to neutralize cell-damaging radicals.
For example, Vitamin C can donate a hydrogen with a single electron to neutralize ROS.
But nutrition isn't as simple as chugging gallons of green tea or blueberry smoothies to live a long life.
ROS have a purpose:
signaling to our cells that something is wrong, so the cell can correct it… or die.
Very high doses of antioxidants can overwhelm those natural signals that too much ROS are being produced in the first place.
So, basically, we need balance!
Now that we know a little bit about neutralizing radicals,
let’s learn how to make them and explore some of the reactions they perform.
[Theme Music]
So far in this series, we've talked a lot about pairs of electrons.
There are lone pairs, shared pairs of electrons in covalent bonds,
and pushing pairs of electrons around in reaction mechanisms.
We've seen many examples of heterolytic cleavage,
where a bond breaks and a pair of electrons migrate to one of the two atoms.
Like when a base is reacted with a hydronium ion, the two electrons from the OH bond both end up on the water molecule.
But electrons don't always have a buddy.
A radical is an atom or group of atoms with a single unpaired electron.
Some radicals are stable, like the chemical nitric oxide, which ends up with a radical on nitrogen.
Radicals can also form during chemical reactions and make some interesting stuff happen,
like the reactive oxygen species we were just talking about.
This process is called homolytic cleavage,
where electrons in a broken bond go in equal but different directions, forming two radicals.
For example, with heat or light, the bond in diatomic chlorine is weak enough to be split equally between the two chlorine atoms,
with each one getting a single electron.
Instead of using a full arrowhead to push the pair of electrons,
we use half-barbed arrows to split the bond, pushing single electrons onto each chlorine atom.
We're finally using those fishhook arrows we talked about in episode 13!
Radical reactions take place in three stages: initiation, propagation and termination.
The initiation stage is where a reactive radical forms.
We need this reaction to get the party started, like the first person who breaks the ice and starts dancing their heart out.
Propagation is where a few radicals bounce around, reacting with other molecules.
It's the step that keeps people on the dance floor,
but it's sort of important to limit the number of high-energy guests so things don't get too rowdy.
Importantly, propagation also regenerates the reactive radical we made in the initiation step.
Finally, the termination stage is when the radical reaction stops.
We’ve had enough high-energy dancing.
It's time to end the party so everyone can go to sleep.
Talking metaphorically about radical reactions and parties might be fun,
but actually looking at examples will help us get more comfortable with them.
To start, let's look at the radical halogenation of alkanes.
To keep things simple, we'll use methane as our alkane
(the thing that's going to be radically halogenated, like the name of the reaction).
Our initiation step makes chlorine radicals, our high-energy party people.
The propagation stage kicks off when a chlorine radical bumps into a methane molecule.
We can think of the chlorine radical as being so charismatic that a hydrogen electron wants to party too.
Or, in reaction terms, make a bond.
The hydrogen only needs to donate one of its two shared electrons to make a bond and form hydrochloric acid.
So, as a result, a methyl radical forms with a single electron on the carbon, sitting happily in a p-orbital.
Then, this methyl radical can react with diatomic chlorine to make chloromethane plus another chlorine radical,
which can continue the propagation stage by reacting with another molecule of methane, generating a new chlorine radical.
The steps of the chain reaction repeat, and the dance party rages on!
Eventually, most of our reagents are used up and it's time to get all the high-energy radicals out of here, so we enter the termination stage.
Basically, two radicals make a bond using their single electrons,
all radicals are used up, and the party is officially over.
Here, there are three possible termination steps: we can end with two chlorine radicals,
or instead one methyl radical and one chlorine radical, or two methyl radicals can combine.
Radical reactions happen quickly, and can create side products,
like how two methyl radicals can make ethane when we’re trying to make chloromethane.
And there are some patterns in radical reactivity that we might recognize from other organic reactions.
First of all, more substitution makes a radical more stable.
We've learned that tertiary carbocations are the most stable carbocations,
and tertiary radicals are the most stable radicals too.
So, for example, if we do radical bromination of propane, we can create a primary or secondary radical.
Because the secondary radical is more stable, we get way more of the product that comes from it: 2-bromopropane.
But here's a weird thing:
if we do radical chlorination of propane, it's much closer to a 50/50 split of the primary and secondary chloropropane products,
even though we still see more 2-chloropropane.
To explain why bromine radicals and chlorine radicals act so differently,
even though they follow the same general pattern, we’ll need to revisit thermodynamics.
The big difference between these two reactions is that the first propagation step in the radical chlorination of propane is exothermic,
which means it loses heat to the surroundings.
Like a chemical hot pack.
On the other hand, the first propagation step in the radical bromination of propane is endothermic,
which means the reaction takes in energy from its surroundings.
Like a chemical cold pack!
We need to look closely at the transition state of these reactions:
the peak of the hill on an energy diagram, where bonds are partially formed.
Hammond’s Postulate is the idea that the transition state of a reaction resembles the species it's closest to in energy.
In other words, the transition state for an exothermic reaction will look more like the reactants.
While the transition state for an endothermic reaction will look more like the products.
So in the first propagation step of radical chlorination, which is exothermic,
the transition state looks a lot like the reactants and happens earlier along the reaction path.
It doesn’t matter too much that the secondary radical is more stable than the primary one,
because the transition state doesn’t resemble the propyl radical that much.
Now, our chlorine radical doesn't just pull off the first hydrogen it collides with on propane.
The formation of the secondary propyl radical is slightly favored because the activation energy is just a bit lower.
But plenty of primary and secondary radicals go on to make the final products.
But in the first propagation step of radical bromination, which is endothermic,
the transition state happens later in the reaction, and resembles a propyl radical much more.
In this late transition state, the secondary carbon radical has a significantly lower peak to overcome its activation energy.
So much more of the secondary propyl radical forms and goes on to make the major product.
We've talked about the differences,
but one key thing these reactions have in common is taking our couch potato alkanes and adding reactive groups onto them.
This is so important to involve them in more exciting organic chemistry.
So with alkanes covered, let’s move on.
Like we saw with superoxide and antioxidants, radical reactions can be in compounds with double bonds too.
For example, the allylic bromination of alkenes.
This is a more carefully-planned party.
Under the right conditions, alkenes can be brominated at the carbon next to the double bond,
which is called the allylic position.
As with all radical reactions, the first stage is initiation.
We need a bromine radical!
To make one, we'll take a bromine-containing source called N-bromosuccinimide,
or NBS for short, and irradiate it with light.
This splits the molecule into radicals, and some of these radicals form a small amount of molecular bromine.
An energetic bromine radical is irresistible to an allylic hydrogen electron on cyclohexene,
so the stage is set for propagation.
HBr is formed, and an allyl radical is left on the alkene.
Basically, allylic hydrogens are so ready to party because this allyl radical is in a p-orbital and is stabilized by resonance with the double bond.
Resonance stabilization in radicals is super important.
In fact, the extensive resonance stabilization in the radical anion formed from Vitamin C is partially why it's such a good antioxidant!
Anyway, to go back to allylic bromination, the propagation stage continues as the allyl radical reacts with the small amounts of molecular bromine we formed.
Even though we have an alkene and molecular bromine,
which are ingredients in the addition reactions we've learned in the past few episodes, the radical propagation steps are really fast so they take control.
The radical reaction keeps propagating until the reagents are all used up and all the radicals pair up, terminating the party.
Now we've seen alkanes and alkenes, but we can't forget about alkynes!
Specifically, the dissolving metal reduction reaction produces E-alkenes from alkynes.
The "dissolving metal" part of the name comes from two reagents that you can see above the reaction arrow:
sodium metal dissolves in liquid ammonia to produce solvated electrons
that are floating around and stabilized by the ammonia solvent.
These solvated electrons form a beautiful dark blue solution.
We'll actually talk about colors in organic chemistry much later in this series,
so this is just a sneak peek.
For now, we'll focus on how solvated electrons do radical reactions.
It's a little tricky, so we'll use a reaction mechanism diagram with orbitals to really see what's happening.
To kick off the reaction, a solvated electron can jump into one of the plain old p orbitals of a pi bond in the alkyne.
As the bond homolytically cleaves, this solvated electron pairs up with one electron from the triple bond.
So now we have a new molecule.
The new electron pair causes one carbon to have a negative charge,
since it has five valence electrons.
Meanwhile, the other electron from the triple bond forms a radical on the other carbon.
When this very basic anion finds a source of protons, like the ammonia floating around,
the negative charge will actually remove a proton from ammonia in an acid-base reaction!
To finish this reaction, another solvated electron comes along, adds to the orbital holding that radical,
and makes another negative charge, putting five valence electrons on the formerly-radical carbon.
At this point, the R groups reorganize themselves to different sides of the double bond to make a more stable anion.
Finally, another molecule of ammonia is deprotonated, and we get an E-alkene!
And the party can end.
Even though we explored antioxidants to solvated electrons,
we didn’t even get to how radicals help us make different kinds of plastics.
So don't worry, this won't be the last time we hear of them!
In this episode, we learned that:
Radicals are highly reactive single electron species, which are important in antioxidant chemistry
There are three steps in radical reactions: initiation, propagation, and termination
Hammond’s postulate helps us predict transition states and explain product distributions
And radical reactions can be used to halogenate alkanes and alkenes, and reduce alkynes
In the next episode of Crash Course Organic Chemistry, we’ll use these halogenated alkanes in a new type of reaction: substitution.
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