Chapter 6: Stereochemistry II: Chiral Molecules and Their Properties

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Welcome back to The Deep Dives, your shortcut to being truly well informed.

We sift through tons of research, pull off the key insights you need.

Today, we're plunging into something really fundamental.

Something that underpins how we build almost everything complex.

From drugs to materials, even life itself, we are talking about carbanions.

Now, I know that word might sound a bit chemistry textbook heavy, but don't worry.

Our mission today is to make it clear.

Show you the logic behind these really crucial molecular players.

At its core, a carbanion is just a carbon atom with a negative charge.

Simple, right?

Think of it like a super reactive Lego brick for chemists.

It lets them stitch together new carbon bonds, build the skeletons of organic molecules.

These are your key nucleophilic carbon species.

They're electron rich, always looking for somewhere positive to connect.

But here's the twist, and this is where it gets really interesting.

Carbanions are usually pretty unstable.

That negative charge makes them super reactive.

They're often hard to generate, even harder to control precisely.

And that challenge?

Well, it's driven decades of research into how to make them, stabilize them, manage them.

It's a puzzle, but solving it gives chemists incredible power.

Exactly.

And that's the journey for today's deep dive.

We'll start with the basics.

What are they?

How do we measure their acidity?

Which tells us a lot about their stability.

Then we'll shift gears a bit.

Look at their shapes, their structures, especially when they partner up with metals and organometallic compounds.

Very versatile stuff.

We'll explore how different bits of a molecule, different functional groups, can help stabilize these reactive carbons, make them more manageable.

We'll even touch on some related neutral players like enols and enemins that behave kind of similarly.

And finally, the payoff.

Their practical uses, how carbanions are just indispensable in synthesis, especially in things like SN2 reactions, letting us build complex molecules, sometimes even controlling their exact 3D shape.

So what does this all mean for you?

Whether you're deep in chemistry,

just curious, or want to grasp the science behind, well, everything.

Understanding carbanions is like getting a molecular construction superpower.

It's vital for pharmaceuticals, materials, science.

It reveals the logic behind how chemists actually make new stuff.

It's fundamental.

So yeah, let's dive in.

Okay, let's start right at the beginning.

Carbanions, carbon, negative charge.

But chemically speaking, what are they?

And why are they such a headache to work with, especially just to form them in the first place?

Well, fundamentally, they're conjugate bases.

Remember, acid base chemistry.

Acid loses a proton, you get the conjugate base.

Here, a proton, a positive hydrogen gets pulled off a carbon atom.

That leaves the carbon with a spare pair of electrons and crucially, that negative charge.

Now, because they have this concentrated negative charge in available electrons,

they are, well, electron rich by definition.

This makes them fantastic nucleophiles.

They're drawn to electron poor spots, ready to share electrons, form new bonds, especially carbon bonds.

Right, acids and bases.

That rings a bell for most people.

We think of things like vinegar, acetic acid.

You measure its acidity in water.

Pretty straightforward.

But here's where the carbanion puzzle really starts, isn't it?

Hydrocarbons, the basic building blocks, are incredibly weak acids, like orders of magnitude weaker than water or alcohol, which creates this huge practical problem.

If you take a strong base, say sodium hydroxide, and try to pull a proton off a hydrocarbon in water, forget it.

The base will just react with the water molecules way faster.

It's like trying to, I don't know, light a match under water.

Yeah, exactly.

That's the core difficulty.

So chemists, well, they had to get clever.

They developed these specialized solvents called a product solvents.

A product just means they don't have easily removable protons themselves, so they won't interfere.

Key examples are things like DMSO, dimethyl sulfoxide,

and cyclohexylamine.

Why do these work?

Two main reasons.

One, they're highly polar, which helps stabilize charged things.

And two, they're exceptionally good at solvating the positive counter ions, the cations.

They surround and stabilize them.

This strong solvation helps keep the newly formed carbanion separate from its positive partner, stops them just snapping back together immediately.

So our measurements actually reflect the intrinsic acidity of the carbanion itself, not muddied by ion pairing effects.

It's a neat chemical workaround.

Okay, so no water, special solvents.

But how do you actually measure the acidity of these ridiculously weak hydrocarbon acids?

You mentioned this basicity function H.

Sounds like a pH scale, but for like super bases in weird solvents.

That's a really good way to think about it.

Yeah.

He is exactly that.

It quantifies how strongly a basic non -acreous solution wants to grab a proton.

And the way they build this scale is quite ingenious.

It uses a series of what are called overlapping indicators.

Imagine you have a set of colored compounds.

Each one changes color or its light absorption spectrum changes when it loses a proton.

By carefully watching these color changes, as you make the solution more and more basic, you can link them together like stepping stones, building a continuous eight scale.

This lets us measure PK values over a huge range way beyond what's possible in water, like up to 35 TK units or even more.

Wow.

And how do you apply that directly?

Well, it's quite elegant, actually.

If you're a hydrocarbon and its carbanion absorb light differently, different colors or different UV vis spectra, you can measure how much of each is present in a solution where you know the H value.

So you see how much hydrocarbon gets converted to the carbanion.

And from that equilibrium, you can calculate its PK.

It's like using light to gauge acidity in these extreme conditions.

That is brilliant.

But it feels like we're pushing the limits here.

What if the acid is so weak, PK maybe over 35, that you can barely even see the carbanion at equilibrium?

Do chemists just give up then?

Oh no, chemists rarely give up.

They find another way.

In those cases, we switch tactics and look at kinetic acidity.

Instead of measuring the amount of carbanion at equilibrium, we measure the rate at which the proton gets removed.

How fast does it happen?

A common method for this is

Say you use a solvent where the hydrogens are deuterium or tritium heavy isotopes.

You can then track how quickly these heavy isotopes get incorporated into your hydrocarbon.

Faster exchange means the proton is removed more easily, the carbanion forms more readily, even if it's just for a fleeting moment.

Techniques like tritium NMR are super sensitive for this.

Okay, so measuring the rate.

But you mentioned a complication.

Internal return.

What's that about?

Sounds like it could mess up the measurement.

It absolutely can.

It's a really crucial, sometimes subtle point.

Imagine the instant the base pulls off the proton, you form this transient ion pair, the carbanion and the protonated base, still right next to each other, maybe with the positive metal ion hanging around too.

Now if this little cluster just immediately collapses back to the starting materials, the carbanion grabs its original proton back faster than it can swap it for a deuterium from the solvent,

then effectively that deprotonation event escapes detection.

You see no isotopic exchange, even though the carbanion did technically form for a split second.

Ah, I see.

So how do you test for that?

Well, one way is to look at chirality.

If you start with a chiral hydrocarbon, meaning it has a specific 3D handedness, and you see it lose that handedness racemise without picking up any deuterium, that's a smoking gun for internal return.

The carbanion formed, flattened out, or inverted, but grabbed its old proton back before exchange could happen.

So for kinetic acidity measurements to be truly reliable, you really need to establish a linear relationship between your measured exchange rates and known equilibrium acidities for similar compounds.

That gives you confidence you're actually measuring the rate of productive carbanion formation.

Okay, so these pK values,

they're not absolute numbers then?

You said solvent and the canation matter?

That's absolutely critical.

They are not universal constants like, say, the speed of light.

The pK value is specific to the solvent system and the positive counter ion used.

For example, hard cations, small, highly charged ones like lithium, tend to encourage tight ion pairing and aggregation, especially in less polar solvents like THF.

But in really good solvating polar or protic solvents like DMSO, ion pairing is much less of an issue.

The solvent molecules do a great job of separating the ions.

So while the numbers might shift a bit between systems, these pK values provide an incredibly useful relative scale.

And DMSO and cyclohexylamine have become key solvents for getting quantitative, comparable data across many different hydrocarbons.

They give us a solid benchmark.

So we know how they're measured, acknowledging the caveats.

So let's actually look at some trends.

Our sources have tables, like table 6 .2, showing pKs from maybe 16 up to over 40.

That's a massive range.

What causes such huge differences?

It is a huge range and it tells us a lot about structure and stability.

Let's dig into some examples.

Look at triphenylmethane versus diphenylmethane versus tellurine.

The acidity drops off dramatically.

PH3CH is way more acidic than

PH2CH2, which is more acidic than plain old PHCH3.

And the reason is clear.

Each extraphenol group provides more stabilization for the negative charge on that central carbon.

It's a powerful combination, really.

You get resonance where the charge is spread out over the pi systems of all those aromatic rings.

Think of it like diluting the charge.

And you also get polar effects or inductive effects where the electronegative phenol groups just pull electron density away through the sigma bonds.

More phenols mean more stabilization, easier proton removal.

OK, that makes sense, spreading the charge.

But then you see something like fluorine.

It looks related, but it's way more acidic than, say, diphenzocycloheptatring.

It seems odd at first glance.

What's fluorine's secret weapon?

That's a fantastic observation, and it leads to a real aha moment.

Fluorine is dramatically more acidic, and the reason is profound.

When it loses a proton, the resulting carbanion achieves aromaticity in that central five -membered ring.

It becomes a cyclopentadieny derivative.

And aromaticity, as you probably remember, confers this huge extra stability.

It's a massive energetic bonus.

You see the exact same thing with cyclopentadiene itself.

Entry nine in the table.

It's exceptionally acidic for a hydrocarbon PK around 16, almost like an alcohol.

Why?

Because its anion is aromatic.

It's a classic example of structure dictating reactivity in a big way.

So that five -membered ring becoming aromatic is key.

Absolutely.

But structure gets subtle, too.

Notice how fusing another benzene ring onto cyclopentadiene, like an indine or benzofluorine, actually decreases the acidity compared to cyclopentadiene itself.

It's counterintuitive, maybe, but it shows how tweaking the structure, even just adding more rings, can affect the electron distribution and stability in complex ways.

What about other types of CH bonds, like next to a double bond allelic?

Or right on a double bond, sp2 carbons?

Or even the hydrogens on plain alkanes, the really saturated ones?

Good questions.

Allelic positions, like in propene, they have PKs around 45.

The negative charge gets some stabilization by delocalizing over the three -carbon allyl system.

Hydrogens, directly on an sp2 carbon, like in benzene or ethane, are also very weakly acidic.

PKs estimated around 45 or 46.

Now, for saturated hydrocarbons like methane or ethane, an isotopic exchange is just way too slow to measure directly.

So there, chemists use more indirect methods.

They might combine electrochemical measurements like, how easily does a carbon radical accept an electron to become an anion, with known CH bond strengths to estimate the PK values?

These methods give values like, say, 48 for methane, maybe 38 for propene, that allelic stabilization helps.

39 for toluene, benzylic stabilization.

These numbers can vary a bit, depending on the method, but they give us a ballpark idea for these extremely weak acids.

Okay, but let's unpack one really standout case.

Terminolatines.

Triple bonds.

These are surprisingly acidic compared to everything else we just mentioned.

Phenylacetylene, PK, around 26 .5 in DMSO.

What makes that CH bond so much easier to break?

That's a fundamental concept, really important.

The key difference is S character in the alkyne, that carbon, is sp hybridized.

That means the hybrid orbitals it uses for bonding are 50 % sorbital and 50 % p orbital.

Compare that to an sp3 carbon in an alpane, 25 % character, or an sp2 carbon in an alkene, 33 % character.

Why does the ski's character matter so much?

Well, as sorbital are closer to the nucleus, they penetrate more effectively than p orbitals.

So electrons in an orbital with higher s character are held more tightly by the nucleus.

This makes the carbon atom effectively more electronegative from the perspective of hydrogen bonded to it.

It pulls electron density away from the hydrogen more strongly, making that proton easier to remove, easier to leave behind.

It's such a clear effect, you can even see correlations between pAK and percent's character, or even with NMR coupling constants like the J13CH value.

Higher s character, more acidic proton.

It also explains why CH bonds on strained rings like cyclopropane are also a bit more acidic.

The strain forces more St's character into those external bonds.

All right, we've got to handle on acidity, the numbers, the trends.

Now let's switch from numbers to shapes.

What do these carbonians actually look like in 3D?

Do they just flatten out when they lose a proton?

Or do they have a preferred geometry?

That's a great question.

And computational chemistry using computers to model molecules has given us pretty clear answers, especially for simple ones.

For methyl anions CH3 or ethyl anions CH3CH2, the calculations consistently show a pyramidal geometry at negatively charged carbon.

Imagine a squashed pyramid, not a flat triangle.

The HCH bond angles are calculated to be around 97 to 100 degrees, significantly less than the 120 degrees you'd expect for a flat, say, P2 hybridized center.

Okay, so pyramidal.

Why?

Why not flat?

It comes down to where that lone pair of electrons is most comfortable, energetically speaking.

The lone pair prefers to reside in an orbital that has more S' character.

In a flat planar carbanion, the lone pair would have to be in a pure P orbital.

But in a pyramidal geometry, it can occupy a hybrid orbital that mixes in some his character.

Orbitals with Morse character are lower in energy because they're closer to the nucleus.

So the pyramidal shape allows the lone pair to be in a more stable lower energy orbital.

It lines up nicely with basic Vius -Seppier theory, too.

Three bonding pairs, one lone pair usually gives a pyramidal shape.

Makes sense.

And does this hold up if you look at them without solvent effects, like in the gas phase?

Yes.

Gas phase measurements are really illuminating.

They strip away all the complications of solvent interactions in ion pairing.

If you look at gas phase acidities like the enthalpy change for proton removal, the absolute energy differences between different hydrocarbons are much larger than in solution.

That reflects the inherent instability of an isolated negative charge without solvent stabilization.

However, the relative trends generally mirror what we see in solution.

Hybridization effects, resonance stabilization, those fundamental factors still dominate.

And things like theoretically computed aqueous pK values, like those shown in table 6 .5, actually show pretty good agreement with experimental data where available, for instance, for ethane.

That gives us confidence in the models.

And can we break down why some hydrocarbons are more acidic than others even further?

Like what contributes to the overall energy change?

Yeah, researchers like Tupitzen and his group did some elegant work dissecting the overall deprotonation energy into two main components.

First, there's the energy needed just to break the C -H bond itself.

This factor tends to be dominant for things like strain ring compounds, where the bonds are already pre -stressed.

Second, there's the energy gained from structural reorganization of the carbenion.

This is the stabilization energy you get when the newly formed anion relaxes into its most stable geometry, often involving delocalization of the negative charge.

This reorganization energy is a major factor for compounds where the anion benefits a lot from resonance or hyperconjugation, like propane or toluene.

Interestingly, benzene has a very low relaxation energy, consistent with its anion having the charge mostly localized in an FP2 orbital.

So yeah, it boils down to two big things.

The inherent strength and character of the C -H bond itself, hybridization, strain, and how well the resulting anion can stabilize that negative charge, delocalization, aromaticity.

Okay, so we know their acidity, their shape,

but how do they behave in 3D space during a reaction?

Let's talk stereochemistry.

You mentioned the environment is critical ion pairing with the positive counterion and the solvent.

Sounds like a delicate dance.

It really is a delicate dance, and the conditions choreograph the whole thing.

Let's take a classic example from the literature.

The cleavage of compound one.

Compound one is chiral.

It has a specific 3D shape.

When its forms and then gets protonated again, we can look at the 3D shape, the stereochemistry of the product to phenylbutane to see what happened to the carbanion in the meantime.

Okay, so what happens in different environments?

Right, let's change the conditions.

In a non -polar solvent like benzene, using potassium TB toxide as the base, the reaction happens with 93 % net retention of configuration.

Whoa, what's that tell us?

It means the carbanion intermediate has a very short lifetime, trapped in a tight ion pair with the potassium ion.

It barely has time to move or reorient itself before it gets protonated again, likely by the protonated base, t -butanol, that's still right there in the solvent cage.

Since benzene itself isn't a good proton donor, the proton comes back from essentially the same direction it left, so the configuration is retained.

Okay, tight ion pair, retention.

What if you change the solvent?

Good question.

Switch to ethylene glycol.

This is a prognotic solvent, meaning it can donate protons and it's more polar.

Now we observe 48 % net inversion of configuration, so the outcome flips.

This suggests the solvent itself is now the main proton source.

The inversion tells us the protonation is likely happening on an unsymmetrically solvated carbanion.

It's perhaps shielded on one side by the counter ion, but exposed on the other side to the solvent, allowing protonation from the back side, flipping the stereocenter.

Fascinating.

And what about a solvent like DMSO?

Ah, DMSO.

That's a highly polar product solvent.

It's good at separating ions, but doesn't readily donate protons itself.

In DMSO, the product is completely racemic, a 50 -50 mix of both configurations.

This implies that in DMSO, the carbanion lives long enough to become symmetrically solvated.

It essentially forgets its original 3D orientation before it eventually finds a proton, perhaps from traces of water or the conjugate acid of the base.

Since protonation can happen equally easily from either face, you get race extremization.

So the solvent and the ion pairing completely dictate the 3D outcome, retention, inversion, or scrambling.

Exactly.

And we see similar patterns in other reactions, like hydrogen -deuterium exchange studies.

For example, exchange in 2 -phenylbutane proceeds with retention in T -butanol -potassium -T -butoxide, suggesting a tight solvent -coordinated ion pair mechanism, but with race extremization in DMSO suggesting a longer -lived, symmetrically -solvated carbanion.

So what does this all mean?

It means by carefully choosing your solvent, your base, your counter ion, you can actually control the stereochemical outcome of reactions involving carbanions.

That's incredibly powerful for making specific molecules, especially complex ones like pharmaceuticals, where the exact 3D shape is critical.

Precisely.

This fundamental of solvation, ion pairing, and carbanion lifetimes is the basis for designing highly stereoselective synthetic methods.

It's not just random, it's predictable and controllable.

We've talked a lot about how tricky it is to make carbanions just by pulling off a proton, but what if you could sort of start with a carbanion already made, just packaged differently?

That brings us to organometallic compounds, right?

Things with lithium or magnesium bonded to carbon.

Let's unpack this.

They act like carbanion salts, even if they're made differently, like carbanions in disguise.

That's a great way to put it, yeah.

They have significant carbanionic character, even though they're often prepared differently, usually by reacting the metal directly with an organic halide.

For instance, methyl iodide reacting with lithium metal gives you methyl lithium, CH3 -ali.

And these compounds, especially alkylithiums and Grignard reagents, organomagnesium halides, are incredibly strong bases.

Methyl lithium's conjugate acid, methane, has a pK estimated around 50.

This means they will rip a proton off almost anything, even remotely acidic alcohols, amines, thiols, even terminal alkynes.

They react instantly.

But here's the really curious part.

Despite being thermodynamically powerful enough to deprotonate many less acidic hydrocarbons, the reactions are often surprisingly slow.

There's a significant kinetic sluggishness.

Okay, that's weird.

Super strong base, but slow reacts sometimes.

Why?

What's slowing them down and how do chemists speed them up when they need to?

The main reason is aggregation.

These organometallic compounds rarely exist as simple single molecules in typical solvents like ethers or hydrocarbons.

Instead, they clump together into clusters, tetramers, four units, hexamers, six units, sometimes even larger aggregates.

N -beta lithium in THF, a workhorse region, is a mix of tetramers and dimers.

Crystal structures show these amazing, often complex arrangements like distorted cubes or stacks with the metal ions, like lithium, forming a core and the organic groups sticking outwards.

Wow, like little molecular fortresses.

Kind of.

And to make them more reactive, to break down these aggregates, chemists often add strongly coordinating solvating molecules.

A classic example is temita tetramethylamine.

Temita is great at chelating the lithium ions, basically wrapping around them using its two nitrogen atoms.

This pulls the clusters apart into smaller, more reactive units.

For example, phenolithium goes from being a tetramer to a dimer when temita is added, making it much more effective as a base or nucleophile.

Computational studies back this up, too.

They show individual methyl lithium is tetrahedral and confirm that these seli bonds are highly ionic, but the clustering helps them dissolve even in nonpolar solvents by hiding the ionic core.

Okay, so this compact character, the aggregation, explains the sluggishness.

It sounds like you need energy just to break the cluster apart before the tarbanion big can even reach out and grab a proton.

Exactly right.

The carbanionic center is buried within this electrostatic cluster,

and activation energy is needed to disrupt the aggregate before it can act as a base.

But what's fascinating here, and incredibly useful synthetically, is that this kinetic sluggishness as a base often doesn't stop them from being excellent nucleophiles.

They will readily attack electrophilic centers like the carbon atom of a carbonyl group in aldehydes, ketones, esters, much faster than they deprotonate the relatively acidic alpha protons of the same carbonyl compounds.

This kinetic preference nucleophilic attack faster than proton abstraction is why organolithiums and grignards are so fundamental for adding carbon groups to carbonals, a cornerstone of organic synthesis.

That's a perfect transition.

So besides packaging them as organometallics, what else makes carbanions easier to handle?

How can we stabilize that negative charge intrinsically within the molecule itself?

This is where electron withdrawing substituents or EWGs come in, right?

They sound like they anchor the charge.

They absolutely do.

EWGs are real game changers.

They dramatically increase the acidity of adjacent CH bonds, sometimes by many, many pK units, making carbonyl formation much easier.

Common powerful EWGs include carbonyl groups, CO found in ketones, aldehydes, esters, nitro groups, NO2, sulfonyl groups, SO2R, and cyano groups, CN.

And how do they work their magic?

What's the mechanism?

It's typically a combination of two effects working together.

First, polar effects.

These groups contain electronegative atoms, like oxygen or nitrogen, that pull electron density away from the carbon through the sigma bonds, just by electrostatic attraction.

This inductive withdrawal helps dissipate the negative charge.

Second, and often more powerfully, resonance effects.

If the EWG has a pi system, like CO, NO2, CN, the negative charge on the adjacent carbon can be delocalized into that pi system through resonance structures.

Spreading the charge over multiple atoms is hugely stabilizing.

You can see this clearly if you look at pK data for substituted methane, say in DMSO, like in table 6 .6.

The stabilization order is pretty clear.

Nitro is the strongest, then carbonyl.

Then groups like esters, sulfones, nitriles are roughly similar, followed by amides.

The resulting carbanions often get special names, too.

Carbanions next to carbonyls are called enolates, and those next to nitro groups are nitronates, highlighting their resonance stabilization.

And I guess having two such groups must be even better.

Oh, absolutely.

The effect is amplified.

Think about datetones, like acetylacetone, pentane 2, 4 -dione.

It has two carbonyl groups flanking a CH2 group.

Its pK in water is around 9.

That's acidic enough to be significantly deprotonated by weak bases, even water or hydroxide.

Compare that to a simple ketone needing a super strong base, like LDA, lithium disipropylamide, or maybe LeHMBS, lithium hexamethyldesilumide.

This ability to easily generate carbanions from compounds with one or especially two EWGs is just fundamental to synthetic organic chemistry.

It opens up countless ways to form new CC bonds.

How do people measure the rates of deprotonation for these stabilized systems?

Is it still isotopic exchange?

Isotopic exchange using DRT is definitely a common way.

But there's also a classic older technique specifically for carbonyl compounds, measuring the rate of halogenation.

See, neutral ketones or aldehydes don't react very fast with halogens like bromine or iodine, but their enolate forms react extremely rapidly.

So if you carry out the halogenation under conditions where enolate formation is the slow step, the overall rate of halogen consumption directly tells you the rate of deprotonation.

It's a clever kinetic trick.

Neat.

And what do those rates tell us about structure?

Well, looking at base catalyzed deuteration rates for alkyl ketones, people 6 .8, there's a clear trend.

CH3 protons are removed fastest, then RCH2, then R2CH slowest.

The dominant factor here is steric hindrance.

It's harder for the base to approach a more crowded proton.

For example, the CH2 group in 2 -butanone gets deprotonated about 100 times faster than the CH2 group in 4 -moronidimethyltupetanone.

That bulky neopetanol group just gets in the way of the base.

Sterics also hinder the solvation needed to stabilize the developing negative charge, further slowing the reaction for hindered sites.

Okay, this leads us directly into a really critical strategic concept in synthesis,

kinetic versus thermodynamic control.

Especially when forming ethylates from unsymmetrical ketones, this sounds like the chemist gets to choose which product they want.

How does that work?

It's like choosing path A or path B to get different outcomes.

It's absolutely a cornerstone of reaction control, a really powerful lever for chemists to pull.

Let's break it down.

Kinetic control means you set up conditions so the reaction goes through the lowest energy transition state, meaning the fastest reaction pathway dominates, regardless of final product's ability.

For enolate formation, this usually means using a very strong, bulky, non -nucleophilic base like LDA or KHNDS, potassium hexamethyldesylenolimide, at very low temperatures like medXD8°C, and a non -equilibrating solvent like THF.

Under these conditions, the bulky base removes the most accessible proton, the least sterically hindered one, most quickly.

This typically leads to the less substituted enolate as the major product, like shown in Scheme 6 .1.

It's all about speed and accessibility.

You can even influence EVSC enolate geometry sometimes under kinetic control.

Okay, so kinetic is fast, less hindered proton, less substituted enolate.

What about thermodynamic?

Thermodynamic control is about letting the system reach equilibrium, so the most stable product predominates, regardless of how fast it formed.

You achieve this by using conditions that allow the proton removal to be reversible, maybe using a weaker base or proprotic solvent, or using an excess of the starting ketones where protons can be exchanged back and forth, or sometimes adding specific additives like HMPA, hexamethylphosphoramide, that promote equilibration.

Under these equilibrium conditions, the more highly substituted enolate is usually the more stable isomer, and therefore the major product.

Why?

Because alkyl groups stabilize double bonds, hyperconjugation, etc.

There are exceptions, of course.

If the more substituted position is extremely crowded, like in 3 -methyl -2 -butanone, steric hindrance can destabilize that enolate enough that the less substituted one becomes thermodynamically preferred.

But the general rule is kinetic gives less substituted, thermodynamic gives more substituted.

You mentioned HMPA as an additive.

Can you give an example of how just adding something like that can dramatically change the outcome, maybe the geometry ZDAZE of the enolate?

Absolutely.

The 2 -methyl -2 -butanone again, deprotonate it with LDA in just THF at low temp.

You mostly get deprotonation at C1, the methyl group kinetic product.

For the C3 position, the methylene, you get some enolate, but the ratio of ZDE isomer is pretty low.

Now here's the kicker.

Add some HMPA to that LDHMF mixture.

You still favor C1 deprotonation overall, but for the C3 enolate, that does form the ZE ratio shoots way up.

It strongly favors the Z isomer.

Wow.

Why does HMPA do that?

It's all about the solvation of the lithium cation.

HMPA is incredibly good at coordinating to LiP+.

It wraps around it, pulling it away from the enolate oxygen during the deprotonation process.

This disrupts the normally preferred titulated transition state, where the LiP plus bridges the base and the ketone oxygen, which tends to favor the E enolate.

Instead, you get a looser, open transition state where the lithium is more solvated by HMPA, and this geometry favors formation of the Z enolate.

It's a beautiful example of how subtle changes in the reaction environment, specifically how the counter ion is solvated, can have profound effects on stereoselectivity.

We see similar effects like rate accelerations with other additives like triethylamine too.

It's all about manipulating that metal oxygen interaction.

We've talked about which proton goes, kinetically or thermodynamically, but what about the actual direction the base approaches from?

How does the 3D shape of the molecule itself, especially in rings, influence which proton is easier to grab?

This sounds like where stereoelectronic factors really come into play, the alignment of orbitals.

Precisely.

Stereoelectronics are all about how the spatial arrangement of orbitals affects reactivity.

A great example is norbornanone, that rigid, bicyclic ketone.

Calculations and experiments show the exoproton, the one pointing sort of outwards from the cage, is removed much faster than the endoproton, pointing inwards.

The energy barrier is about 3 .8 kilomole lower for exo removal.

Why?

Because the C -H bond of the exoproton is better aligned with the pi system of the carbonyl group.

This allows for more efficient overlap between the breaking C -H bond orbital and the developing p orbital on carbon as the analyte forms.

Better alignment means a lower energy transition state.

And you see similar things in simpler rings, like cyclohexanone.

Absolutely.

In a standard cyclohexanone share confirmation, removing an axial proton is significantly favored over removing an equatorial proton by about 2 .8 kilocommel, according to calculations.

Again, it's about better stereoelectronic alignment.

The axial C -H bond is oriented roughly parallel to the orbitals of the carbonyl pi system, allowing for good overlap during departination.

See, figure 6 .3 in the source.

Removing the equatorial proton requires more distortion and encounters more torsional strain in the transition state.

The molecule prefers the path of least electronic resistance.

Okay, this is fascinating.

Yeah.

But it raises a potential conflict, doesn't it?

Yeah.

We talked about nitroalkanes earlier.

You said adding alkyl groups makes them kinetically slower to deprotonate because of sterics.

Right.

But thermodynamically more stable because alkyl groups stabilize the nitrinate.

How can something be more stable but harder to form?

Doesn't that seem contradictory?

It seems contradictory, but it highlights the crucial difference between reaction rates, kinetics, and product stability, thermodynamics.

They aren't always correlated.

For nitroalkanes, the transition state for proton removal is sterically hindered by the alkyl groups.

So the activation energy barrier to get to the anion is higher, making the reaction slower, lower kinetic acidity.

However, once formed, the resulting nitrinate anion is more stable with more alkyl groups, likely due to hyperconjugation or inductive effects helping stabilize the negative charge, higher thermodynamic acidity.

So yes, you can absolutely have a situation where the more stable product forms more slowly.

It just depends on the relative energies of the starting materials, transition states, and products.

Got it.

Kinetics versus thermodynamics again.

What other groups are powerful stabilizers?

You mentioned cyano.

Yes.

The cyano group, CN, is extremely effective, as shown in table 6 .9.

It's strongly electron withdrawing, both inductively and through resonance.

Hydrocarbons with multiple cyano groups can become remarkably strong acids.

And we shouldn't forget stabilization by elements from the third row, particularly phosphorus and sulfur, table 6 .2.

A synthetically vital example is 1 -pharmate -3 -dithion.

The protons between the two sulfur atoms are surprisingly acidic, PK 36 .5 and THF, readily removed by N -butylytium.

The resulting anion is a fantastic nucleophile used in many CC bond formations.

The reasons for sulfur's stabilizing effect are debated, but likely include the polarity of the CS bond, sulfur's greater polarizability, its electron cloud is larger and more easily distorted to accommodate charge, and possibly some contribution from hyperconjugation involving CS sigma orbitals, or even D orbitals, although D orbital participation is now considered less important than previously thought.

Polarizability seems key.

A phenylpheo group, PHS, can boost acidity by maybe 15 PK units.

And this leads us to those unusual molecules with adjacent charges, FODs, phosphonium, sulfonium, whole -sides.

You mentioned their synthetic importance.

Exactly.

Yylides are neutral overall that have a formal negative charge on carbon right next to a formal positive charge on a heteroatom like phosphorus, sulfur, or nitrogen.

Their structure was debated for a long time.

Was it mainly the dipolar ylide form, C -P++, or an uncharged eileen form with a double bond, CP, which would imply D orbital involvement for P and S?

Modern evidence strongly favors the dipolar ylide structure as the main contributor.

The stability comes primarily from the strong electrostatic attraction between the adjacent opposite charges and the electron -withdrawing inductive effect of the positive heteroatom.

They are usually made by deprotonating the corresponding onium salt, like a phosphonium salt, R3P, plus a national CH2R.

Substituents that stabilize the negative charge on carbon make the halolide more stable and easier to form.

For sulfur ylides, sulfoxonium ylides with an extra oxygen on sulfur are generally more stable than sulfonium ylides because that extra oxygen is very electronegative and helps pull electron density away.

Okay, let's change perspective slightly.

We focused on negatively charged carbanions, but you mentioned related neutral molecules that can act as carbon nucleophiles too.

Talk about enols first.

What's really striking here is that carbonyl compounds can act as nucleophiles even under acidic conditions through their enol tautomer.

How does that work?

A neutral molecule acting nucleophilic in acid seems counterintuitive.

It does seem a bit odd at first, but it's down to tautomerism.

Most simple ketones and aldehydes exist in a rapid equilibrium with a small amount of their corresponding enol form.

This is the keto -enol equilibrium.

The enol has a CEC double bond and an iso -H group attached to one of the carbons.

This equilibrium can be catalyzed by either acid or base.

Now why is the enol nucleophilic?

Because of the pi electrons in its CEC double bond, just like an alkene.

But it's actually more reactive than a typical alkene.

The attached OH group acts as an electron donor through resonance, pushing electron density into the double bond, making the other carbon, the beta carbon, more electron -rich and nucleophilic.

Plus, during the reaction, forming the very strong CO bond in the product provides a thermodynamic driving force.

However, it's important to remember they are less nucleophilic than the corresponding negatively charged enolate anions, simply because they're neutral, not carrying that extra negative charge.

Still useful, but less powerful.

So how does this enolization happen in acid?

What's the crucial step?

Under acid catalysis, the generally accepted mechanism starts with protonation of the carbonyl oxygen.

This makes the carbonyl carbon more electron deficient, which in turn makes the alpha protons more acidic.

The rate -determining step is usually the removal of one of these alpha protons by a weak base, like water or the conjugate base of the acid catalyst, to form the enol CEC double bond.

And what really confirms this?

A significant primary kinetic isotope effect.

If you replace the alpha hydrogens with deuterium, the reaction slows down considerably.

KHKD is often around 5.

This tells us the CH or CD bond is breaking in the slow step.

OK, and does acid catalysis lead to the same enol as base catalysis if the ketone is unsymmetrical?

You mentioned base favors the less substituted.

Good question.

No, typically not.

Acid catalyzed enolization generally favors formation of the more substituted enol.

See table 6 .11.

For example, with 2 -butanone, acid favors enolization towards the internal methylene group C3 over the terminal methyl group C1 by about 4 .1 after correcting for statistics.

This is because alkyl groups stabilize double bonds.

The transition state leading to the more substituted enol resembles that more stable double bond, so it's lower in energy.

The overall reactivity differences between different sites are also smaller in acid compared to base catalysis.

So acid gives more substituted, base gives less substituted kinetically.

More tools for the chemist?

What about the actual amount of enol present at equilibrium?

Is it usually a lot or a little?

For simple aldehydes and ketones, it's usually very little.

The equilibrium constant for enolization, canal, is typically tiny, like 10 -4 to 10 -5 for aldehydes, even smaller for ketones, maybe 10 -8.

Table 6 .12.

The ketoform is strongly favored thermodynamically, and for esters and amides, it's practically negligible.

Canal can be as low as 10 -20.

Why?

Because the carbonyl group in esters and amides is already strongly stabilized by resonance donation from the adjacent oxygen or nitrogen lone pair.

That makes the ketoform super stable, and there's almost no driving force to carbonyl compounds earlier.

Exactly.

That's the big aha moment for enols.

For about diatones, like acetylacetone and ketosteres, like ethylacetoacetate, the enol form can be the major tautomer present at equilibrium, sometimes well over 50%.

There are two key reasons.

First, the enol form is stabilized by conjugation.

The COC double bond is conjugated with the remaining carbonyl group.

Second, and very importantly, they can form a strong intermolecular hydrogen bond between the enolic OH group and the oxygen of the nearby carbonyl group, forming a stable six -membered ring.

Structural data clearly shows this internal H bond.

It locks the molecule in the enol form.

And does the solvent affect this ketoenol balance, especially for those H bonding cases?

Yes, significantly.

Take ethylacetoacetate.

It has much higher enol content in nonpolar solvents like carbontate or benzene, maybe 15 -30%, compared to polar solvents like water, 1%, or acetone, 5%.

The intermolecular hydrogen bond in the enol makes the molecule less polar overall.

Nonpolar solvents favor this less polar form.

Polar solvents, especially those that can hydrogen bond themselves, like water, consolidate the ketoform effectively, stabilizing it and shifting the equilibrium away from the enol.

It's a competition between internal H bonding and external solvation.

Can you ever isolate these simple enols, like the enol of acetaldehyde?

It's tricky, but yes.

Chemists have managed to generate unstable enols, like vinyl alcohol, enol of acetaldehyde, and observe them at low temperatures, sometimes for hours.

They're tasteable.

At room temp, they isomerize back to the ketoform quickly, especially if any base catalyst is around.

Interestingly, polar product solvents like DMSO or DMF can actually slow down this isomalization by hindering the proton transfer steps, extending the enol's lifetime somewhat.

And acidity -wise, enols are way more acidic than their keto counterparts.

Acetophenone's enol has a pKa around 10 .5, while the ketoform is about 18 .4.

This difference directly reflects that tiny kennel value we talked about.

The ease of forming enols or enolates makes them central to many synthetic reactions.

That guy covers enols.

What about enamines?

Similar idea, but with nitrogen.

Exactly.

Adenomines are the nitrogen analogs of enols.

You have an amino group attached to a CQC double bond.

They're typically formed from ketones or aldehydes reacting with secondary amines.

The key difference.

Nitrogen is generally a better electron donor than oxygen.

This means the nitrogen lone pair pushes electron density into the double bond even more effectively than the oxygen in an enol.

As a result, the beta carbon of an enamine is even more nucleophilic than the beta carbon of an enol.

They are superb carbon nucleophiles for synthesis.

And do they have interesting structural quirks like enols?

They do, particularly regarding which isomer forms.

For enamines made from unsymmetrical ketones like two alkylcyclohexanones, the double bond preferentially forms towards the less substituted side.

This is especially true for enamines from cyclic amines like pyrolidine.

The reason is believed to be A13 allylic strain.

For maximum nucleophilicity, the nitrogen lone pair needs to be aligned, or coplanar, with the CQC double bonds pi system.

If there's a bulky substituent on the beta carbon, the C3 position in A13 strain terms, it clashes sterically with the groups on the nitrogen in this required planar conformation.

So the molecule avoids this clash by forming the double bond where there isn't a bulky group.

This steric effect also influences conformation.

Alpha alkyl groups often prefer an axial position to minimize interaction with the amino group.

It also affects specificity less strained.

Five and seven -membered ring enamines are often stronger bases than six -membered ones due to better conjugation.

Okay, we've built a solid picture of carbanions, inlets, enols, enemins, how they form, how stable they are.

Now let's focus squarely on their main job in synthesis, acting as nucleophiles to build new carbon -carbon bonds, especially via SN2 reactions.

Yes, this is where they really shine in constructing molecular frameworks.

Carbanions and their stabilized cousins, like enolets, are generally considered soft nucleophiles.

What does soft mean here?

It's part of the HSAB hard -soft acid -base principle.

Soft nucleophiles tend to have their reactive electron density in more diffuse polarizable orbitals.

They preferentially react with soft electrophilic centers, and carbon atoms, especially in typical alkyl halides used for SN2, are considered relatively soft.

This makes them excellent for SN2 reactions attacking the carbon and kicking out a leaving group, rather than attacking harder centers like protons or metal ions, though they can act as bases as we discussed.

This selectivity is key for C -C bond formation.

Can you give some examples of these SN2 reactions using organometallics?

Sure.

Classic examples involve reacting things like aryl, alkanol, or allyl lithium reagents with primary alkyl halides, or tessellates, schemes into point 2.

The lithium region acts as the carbon nucleophile.

Similarly, Grignard regions, like aryl magnesium halides, are often reacted with alkyl sulfates or sulfonates to form new C -C bonds.

These are bread -and -butter reactions for building up carbon chains or attaching groups to rings.

Do we have proof these are really SN2, like the stereochemistry we talked about earlier, backside attack inversion?

For some cases, absolutely yes.

When allyl or benzoyl lithium reagents react with a chiral substrate like 2 -bromobutane, you see clear, complete inversion of configuration.

That's the textbook SN2 signature.

However, it gets complicated.

Here's the twist.

If you use simple n -butyl lithium, you often get a largely racemic product, meaning the stereochemistry gets scrambled.

This suggests that with simple alkyl lithiums, isn't always a clean SN2.

Other pathways, possibly involving electron transfer leading to radical intermediates or complex interactions with the organometallic aggregates, might be competing and leading to loss of stereochemical information.

Remember those aggregates?

They complicate things.

Halide ions from the substrate can even get incorporated into the lithium clusters, changing their reactivity.

We also see things like allyl transposition, where, say, phenolithium reacts with allyl chloride, and the new bond forms at the end of the allyl system, C3, instead of where the chloride left, C1.

This might involve cyclic transition states with the organolithium dimer.

It's not always simple.

So while powerful, these direct alkylations of organometallics can be tricky.

Have chemists found better ways?

Yes.

For many applications, especially in complex molecule synthesis, direct alkylation of simple organolithiums or greenyards has been largely superseded or augmented by transition metal catalyzed cross -coupling reactions.

Using catalysts based on metals like copper, palladium, nickel allows for much milder conditions, broader substrate scope, better functional group tolerance, and often exquisite control over the reaction.

Think Suzuki, Heck, Nagishi couplings.

These are workhorses now.

But the fundamental principle remains the same.

Using a species with a carbionic character, often generated transiently in the catalytic cycle, to form a new carbon -carbon bond with an

That makes sense.

Let's zero in then on arguably the most widely used and perhaps most controllable version of carbanion chemistry in synthesis.

The alkylation of enolate anions derived from ketones, esters, and related compounds.

This seems like where the real precision engineering happens.

It absolutely is.

Enolate alkylation is incredibly versatile and allows for the construction of complex carbon skeletons with high predictability and control.

This applies to both the highly stabilized enolates from banocarbonyls and the less stable, more reactive ones from simple ketones or esters generated using strong bases like LDA.

To understand the control, we need to revisit their structure and solution.

Just like organolithiums, lithium enolates, for example, are often aggregated tetramers, dimers, even hexmers are common, depending on the enolate structure, solvent, and concentration.

And does this aggregation affect how they react?

You mentioned for organolithiums that the monomer might be the reactive species.

Exactly the same principle seems to apply here.

Studies on enolate reactivity strongly suggest that even if the aggregate, like a tetramer, is the dominant species in solution, the monomeric form of the enolate, even if present only in tiny amounts, like the 1 .3 % mentioned for isobufenone enolate, is often the actual nucleophile doing the reacting.

Why?

Possibly because the negative charge on the carbon is less stabilized by interactions with just a single lithium ion compared to being buried within a larger aggregate cluster.

The monomer is freer and more reactive.

So the aggregates are like reservoirs, slowly releasing the active monomer.

What do these aggregates look like from crystal structures?

The crystal structures are really beautiful and informative, like in figure 6 .4.

They often show these intricate cubic or ladder -like arrangements.

Lithium cations and the oxygen atoms of the enolates occupy alternating corners or positions.

A key feature is that the enolate oxygen atom is usually involved in multiple bridging contacts with lithium ions within the cluster.

This effectively ties up the oxygen.

Meanwhile, the nucleophilic carbon atom tends to be more on the periphery of the cluster, more exposed and available to react with an incoming electrophile.

This inherent structure biases the reaction towards C -alkylation over O -alkylation.

Computational models do a great job of predicting these structures and confirming their stability.

That structural bias is fascinating.

But chemists still need to worry about regioselectivity, right?

Getting C -alkylation versus O -alkylation.

What controls that choice?

Yes.

Controlling CVS -O -alkylation is crucial.

Generally, as we said, carbon is the softer nucleophilic site and typical alkylating agents, alkoholides, are soft electrophiles.

So C -alkylation is the norm.

Looking at it from a frontier molecular orbital FMO perspective, the highest occupied molecular orbital, HOMO, of the enolate, which is where the nucleophilic electrons reside, has large coefficients, electron density, on both the oxygen and the alpha carbon.

The reaction pathway is often dictated by which interaction leads to the lowest energy transition state, often involving that stereoelectronic preference for attack perpendicular to the enolate plane.

But the CO ratio is very sensitive to reaction conditions, which gives us levers to pull.

First, the degree of aggregation.

Adding things that break up aggregates strongly coordinating solvents like HMPA, DMF, DMSO, or additives like crown ethers that sequester the cation tends to increase the reaction rate and often favors O -alkylation.

Why?

Because freeing the oxygen makes it more nucleophilic, less tied down by the cation.

Second, the metal cation.

Reactivity generally increases Liena, correlating with weaker ion pairing and more dissociated ions.

Thirteen CNMR studies confirm higher electron density at carbon with more dissociated cations, like K -plus versus Li -plus, favoring c -alkylation.

Third, the leaving group on the electrophile.

Hard leaving groups, less polarizable, often oxygen -based like sulfonate sulfates, tend to favor reaction at the hard oxygen site, O -alkylation.

Softer leaving groups, more polarizable, like bromide, iodide, favor reaction at the soft carbon site, c -alkylation, HSOB principles in action again.

Okay, incredible control over CVSO.

Now for the ultimate challenge.

Serial selectivity of enolate alkylation.

Controlling the 3D outcome.

How do you ensure the new alka group adds to the correct face of the enolate?

This is where synthesis becomes truly three -dimensional chess.

The guiding principles combine sterics and stereoelectronics.

Principle one.

The electrophile usually approaches from the less sterically hindered face of the enolate, avoid the crowds.

Principle two.

There's that stereoelectronic requirement for the electrophile to approach roughly perpendicular to the plane of the enolate double bond in a way that minimizes torsional strain, twisting, as the new bond forms and the carbon re -hybridizes.

Orbitals need to overlap efficiently.

Can you give some examples of how this plays out?

Sure.

For simple cyclohexanone enolates, there's often a slight preference for the electrophile to come in from the equatorial direction, which is often less hindered.

For enolates within a ring system, endocyclic, like in two alkylcyclohexanones, attack often occurs to maintain the stable chair conformation if possible.

In fused ring systems like decolones, the preferences can be quite pronounced.

The one -menolate of one decolone gets alkylated to give a cis -ring fusion because attack from the top face avoids bumping into part of the other ring, but the two -enolate of trans -2 decolone prefers axial attack from the bottom and this preference gets even stronger if there's an alkyl group at C1, pushing the enolate conformation to favor that approach.

Theoretical work by Ken Huck and others really highlighted torsional effects.

Minimizing eclipsing interactions between bonds and the transition state can be decisive.

For example, in trans -2 ,3 -dimethyl cyclopentanone, attack syn to the methyl groups is slightly favored because it leads to a more staggered lower energy transition state.

Even in acyclic systems where things are floppier, alkylation usually happens anti -opposite side to the largest existing substituent on the chiral center.

Those selectivities might only be modest unless you design the system carefully.

A dramatic example is using a bulky, silly group.

Its sheer size can act as a very effective stereo directing group, forcing the electrophile to come in exclusively from the opposite face, leading to very high selectivity.

So the bottom line is, by carefully considering the 3D shape of the enolate, the steric environment around it, and these fundamental stereo -electronic preferences for attack trajectory, chemists can often predict and achieve remarkably high levels of stereo control in these crucial bond -forming reactions.

It allows for the precise construction of complex target molecules with defined 3D architectures.

So what does this all mean, this deep dive?

From basic acidity to complex aggregation and stereo control?

It really shows carbanions aren't just reactive intermediates, they're incredibly versatile tools.

Chemists can use this deep understanding to build new molecules with just amazing precision, controlling where bonds form and how they're arranged in 3D space.

It really is molecular engineering.

Absolutely, and connecting these fundamental principles, acidity, structure, ion pairing, sterile electronics, directly to their practical applications and synthesis is key.

This knowledge underpins so much of modern organic chemistry, pharmaceuticals, materials,

everything.

We hope this journey into the world of carbanions has sparked your curiosity about the power locked within these tiny charged species.

Think about how manipulating them allows us to create the world around us.

Keep exploring, keep asking why.

Thank you so much for joining us on this deep dive and for being part of our learning community here at Last Minute Lecture.

Stay curious, stay informed, and we'll catch you on the next one.