Chapter 8: Acidity, Basicity, and pKa

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Welcome to the Deep Dive where we take your most interesting sources and really extract the essential nuggets of knowledge, giving you that shortcut to being truly well -informed.

Today, we're diving head first into, well, a cornerstone of organic chemistry for any serious student.

We're focusing on Chapter 8, Acidity, Bacicity, and Picae from Organic Chemistry, the second edition by Clayton, Greaves, and Warren.

Our mission for you, our listener, is to untangle these fundamental principles of acid -based chemistry in organic compounds.

We'll really emphasize the mechanistic reasoning, the why behind how molecules behave, you know, how this impacts reactions, functional group transformations,

make it all predictable.

We'll try to cut through the complexity with clarity using language, hopefully accessible to upper level undergrads so you grasp not just what happens but why it's so incredibly important.

That's absolutely right.

Understanding picae isn't just like memorizing a table.

It's foundational, truly.

It's the key to predicting how molecules will behave in pretty much any reaction, how they'll dissolve,

and critically, how they can be intelligently designed for things like drug discovery.

This deep dive will hopefully eliminate the why behind these crucial properties, make them feel more intuitive.

Okay, let's unpack this then.

Let's start with a very practical sort of real -world challenge.

How do we get organic compounds, which are often, let's face it, stubbornly insoluble, to actually dissolve in water?

Or maybe even more practically, how do we efficiently separate them out from mixtures?

Ah, yes.

The elegant solution really lies in temporarily converting them into ions.

Water, you see, being a polar product solvent, is excellent at surrounding and stabilizing charged species, both positive and negative.

A perfect example is aspirin.

The acidic form.

Barely soluble.

But if you add a weak base, something simple like sodium hydrogen carbonate, it readily forms its sodium salt.

And that salt is highly soluble in water.

What's really fascinating, though, is how reversible this is.

In the very acidic environment of your stomach, that soluble aspirin just flips right back to its insoluble acid form.

Huh.

So it's like we're just giving the molecule a temporary charge to make it play nice with water.

Precisely.

Exactly that.

And we see the same principle with organic bases.

Think about codeine, for instance.

As a neutral compound, it's sparingly soluble.

But just by lowering the pH of the solution, its basic nitrogen atom happily picks up a proton, gets protonated, and forms a salt that's much, much more soluble.

This principle is absolutely vital in how many drugs are formulated and how they get absorbed in the body.

Okay.

That's brilliant for solubility.

I get that.

But what other practical uses jump out?

Thinking about manipulating these charges,

can we use this to actually separate compounds?

Oh, absolutely.

And it's one of the most elegant applications of this chemistry, I think.

Adjusting pH is just a fundamental strategy in lab separations.

Imagine you've got a mixture, say, benzoic acid and toluene.

They're both dissolved in dichloromethane, maybe.

Now, if you add an aqueous solution of sodium hydroxide and AOH and you shake it up, here's what happens.

The toluene, it's neutral, right?

It stays happy in the organic layer.

But the benzoic acid, being an acid, reacts with an AOH.

It converts into its water -soluble sodium salt and, poof, it migrates into the aqueous layer.

Ah, clever.

Then you just separate the layers, take that aqueous layer, add some acid back to it, acidify it, and boom, your insoluble benzoic acid crashes right back at a solution.

It's a classic purification technique.

And it's a constant reminder in the lab to always ask, you know, what charge does my compound have right now?

Which layer will it prefer?

It can save you from accidentally throwing away your product.

Okay, so we've seen how incredibly practical controlling these properties is in the lab, in drug design.

But to really control them, we need the language, the measurements.

That brings us to pH and pKa.

Indeed.

At its simplest, an acid is something that tends to lose a proton and a base tends to accept one.

We often forget, I think, that an isolated proton is incredibly reactive, like a tiny ball of positive charge.

When we say HCl, hydrochloric acid, is strong, it's not because free H plus ions are floating around, no.

It's because HCl almost completely hands over its proton to water, forming stable hydronium ions, H3O plus Az and chloride ions.

Water acts as the perfect sort of chaperone for that proton.

Right.

So pH is basically just measuring that H3O plus concentration.

It tells us how acidic or basic the solution is on that log scale.

Neutral seven below is acid above base.

Simple enough for the solution itself.

Exactly.

But pH measures the solution's acidity.

It doesn't tell us about the inherent personality of a compound, its own tendency to give up a proton.

That's where pKa comes in.

The pKa is the specific pH value where an acid and its conjugate base are present in exactly equal amounts, 50 -50.

So for acetic acid, its pKa is about 4 .8.

That means if you adjust the solution to pH 4 .8, you'll find exactly half acetic acid molecules and half acidate ions.

And technically, formally, pKa is the negative log of CHI, the acid dissociation constant.

So lower pKa means stronger acid.

Makes sense why HCl pKa around minus seven is way stronger than acetic acid.

Precisely.

And if we look at the bigger picture, even water has a pKa around 15 .7.

This value more or less defines the practical pH range we can work with in water, roughly from say, negative 1 .7 up to 15 .7.

So every assay has a conjugate base, every basic conjugate acid.

What does that tell us about their relative strengths?

Is there a connection?

Oh, absolutely.

It's an inverse relationship.

The stronger the acid, the weaker its conjugate base.

It has to be.

Think about hydrogen iodide, HI, super strong acid pKa about minus 10.

Its conjugate base, the iodide ion, I is incredibly weak, practically non -basic in water.

Conversely, look at methane, CH4.

Its pKa is sky high, maybe 48.

It's unbelievably non -acetic, which means its hypothetical conjugate base, the methyl anion, CH3, would be an extraordinarily strong base, almost too strong to really exist in water.

That makes sense.

But, hmm, that brings up a question.

If something is too strong an acid or base for water, how do we even measure its pKa if it just reacts with the water immediately?

Yeah, that's a great point.

That's where we run into solvent limitations.

In water, like we said, we're pretty much stuck measuring pKa's between about one acoma up to 1 .7 and 15 .7.

For acids stronger than H3O plus R, or bases stronger than hydroxide, you can't just dissolve them in water and measure their true strength.

They'll just protonate or deprotonate the water itself.

It levels them out.

So for those extreme cases, the values are determined in other solvents, non -aqueous ones, where the solvent doesn't get in the way.

Then they're carefully extrapolated back to what they would be in water, like deprotonating acetylene, pKa25.

You can't use hydroxide in water.

You need something much stronger, like sodium amide and liquid ammonia.

The solvent matters.

Okay, so now we know what pKa is.

Let's get into the why.

Why does a compound have its specific pKa?

What makes one acid stronger than another?

What are the key factors?

Right.

The absolute most crucial factor, the sort of golden rule, is the stability of the conjugate base.

The more stable that anion is after the proton leaves, the more willing the acid is to let that proton go.

Makes sense, right?

More stable product means easier reaction.

And this comes from a few key things.

First, there's electronegativity.

As you go across a row in the periodic table, say from carbon to nitrogen to oxygen to fluorine, electronegativity increases.

Exactly.

It increases.

Atoms get better at pulling electrons towards themselves and, crucially, better at handling a negative charge.

So the pKa's of methane, ammonia, water, and HF just plummet from like 48 for methane, way down to about three for HF.

Why?

Because the fluoride ion, F, is way happier holding that negative charge than a metal anion CH3 is.

Second factor,

bond strength.

Now, this is interesting.

As you go down a group, like the halogens F, Cl, Br, I...

The bond to hydrogen gets weaker.

It gets weaker, yes.

And the anion gets bigger.

So even though fluorine is the most electronegative, the combination of the weaker H -I bond and the larger size of the iodide ion, which spreads out the charge, makes H -I a much stronger acid than HF.

H -I has a pKa around minus 10.

Remember, HF is about three.

So size and bond strength win out over electronegativity on going down a group.

Electronegativity across bond strength size down.

Got it.

And here's the one that seems really powerful in organic chemistry,

delocalization.

This feels like a game changer.

Oh, it absolutely is.

Huge effect.

Consider the perchloric acids, HClO, HClO2, HClO3, HClO4.

As you add more oxygen atoms, the negative charge on the conjugate base, the anion, can be spread out, delocalized, over more and more oxygen atoms through resonance.

Spreading out charge is stabilizing, always.

So the anion gets much more stable and the acid gets much, much stronger.

The pKa drops like a stone from about 7 .5 for HClO all the way down to minus 10 for HClO4.

And the fact that all the ClO bonds in the perchlorate ion are identical, that's beautiful physical evidence of that charge being perfectly shared.

So how does this play out in common organic acids, like comparing an alcohol to a carboxylic acid?

The difference is just massive.

An alcohol -like ethanol, pKa16,

it's conjugate base, ethoxide,

that negative charge is stuck, localized, on that one oxygen atom.

But acetic acid, pKa, around 4 .8.

It's conjugate base, the acetate ion.

That negative charge is delocalized over two oxygen atoms through resonance.

That stabilization makes acetic acid about a billion times stronger as an ethanol.

It's incredible.

Even just delocalizing into a benzene ring, like in phenol, pKa, about 10, makes it way more acidic than, say, cyclohexanol, pKa16, where there's no resins possible for the anion.

Wow.

Okay, so when looking at acidity, the checklist is electronegativity, size bond strength, and especially can that negative charge be spread out through delocalization?

That's a great way to think about it.

And to simplify, maybe we can keep some rough benchmarks in mind, like protonated acids,

pKa or maybe zero goes, carboxylic acids around 5, phenols around 10, alcohols themselves around 15, 16.

Those are excellent benchmarks to have, very useful reference points.

All right, let's pivot a bit.

Let's talk nitrogen.

Amines, amines.

Nitrogen's less electronegative than oxygen.

How does that fundamental difference shape their acid base properties?

Right.

It means, generally speaking, amines are less acidic and more basic than their oxygen cousins, the alcohols.

Protonated amines, like ammonium salts, typically have pKa values around 10, actually quite similar to phenols, interestingly.

This means amines are often protonated at physiological pH, around 7 .4, which is really important for how they function biologically.

Trying to pull a proton off a neutral amine, though, that's tough.

You need a really strong base, like an organobotallic region, to form the amide anion, which is pretty unstable.

Okay, but I've definitely heard chemists say things like, oh, the pKa of triethylamine is about 11, but triethylamine doesn't have any acidic protons to lose, right?

So what do they mean by that?

That always confused me a bit.

Yes.

You've hit on a really common, slightly confusing bit of jargon.

What they mean is the pKa of the conjugate acid of triethylamine.

Sometimes you'll see it written as pKaH, to be clear.

Triethylamine itself is a base.

Its conjugate acid, the tricholomonium ion, ET3NH +, is the species that has a pKa of 11 .0.

So that number, 11, is really telling you about the basicity of the triethylamine, how readily it picks up a proton.

Higher pKaH means stronger base.

Got it.

pKaH for basicity.

Makes sense.

And aniline, the aromatic aniline, shows this sort of interesting duality.

It's way less basic than simple amines like ammonia.

Why?

Because its nitrogen lone pair is busy being delocalized into the benzene ring, less available to grab a proton so that the pKaH is low, maybe 4 .6.

But for exactly the same reason that delocalization aniline is more acidic than ammonia, if you do manage to deprotonate it, the resulting anion is stabilized by that same resonance into the ring.

So aniline's pKaH for losing the NH proton is around 28, whereas ammonia's is much higher, like 33.

Delocalization cuts both ways.

Fascinating.

And hybridization matters too, right?

Like comparing pyridine and piperidine.

Absolutely.

Hybridization has a big effect on basicity.

Pyridine has its lone pair in an SB2 orbital.

Piperidine, the saturated version, has its lone pair in an SB3 orbital.

Electrons in its B2 orbitals are held tighter, close to the nucleus, because there's more karatis character, makes them less available for donation.

So pyridine is a much weaker base, pKaH around 5 .5 compared to piperidine, pKaH around 13.

Okay.

What about amides then?

They seem different again.

Amides are really quite different beasts, yes.

That nitrogen lone pair is heavily involved in resonance with the carbonyl group next door.

This makes them, perhaps surprisingly, significantly more acidic than amines pKaH around 15, so you can actually deprotonate them with strong bases, like hydroxide sometimes.

But it also makes them much weaker bases.

The lone pair isn't really available on nitrogen.

Protonation, if it happens, usually occurs on the carbonyl oxygen, and the pKaH is way down near zero, so not basic at all, really.

And then you have the powerhouses, amidines and guanidines?

Indeed.

They are incredibly strong bases for neutral organic molecules.

Amidines are stronger than amines because you get resonance stabilization of the protonated form involving both nitrogens.

And guanidines, with three nitrogens involved in stabilizing the positive charge after protonation, they are among the strongest neutral organic bases known.

pKaH values up around 14 .5.

This makes arginine, the amino acid with the guanidine group, the most basic amino acid.

Super important biology.

Wow, okay.

So beyond the main functional group, how do other bits of the molecule, you know, substituents, tweak these pKa values?

It feels like small changes can sometimes have pretty big effects.

You're absolutely right.

Substituents can have dramatic effects, especially if they are either involved in conjugation or if they're strongly electron withdrawing or donating through inductive effects.

For instance, take phenol, pK10, stick a nitro group on the other side of the ring in the pair position, you get p -nitrophenol.

That nitro group is strongly electron withdrawing through both resonance and induction.

It helps pull electron density away and stabilize the negative charge when the phenol loses its proton.

So the pKa drops significantly down to about 7 .1, much stronger acid.

And inductive effects alone can be powerful too.

Oh yes, inductive effects pulling electron density through the sigma bonds can be very strong, especially with highly electronegative atoms like fluorine.

Compare acetic acid, pKa 4 .8, with trifluoroacetic acid, TFA.

You just replace the methylhydrogens with fluorines.

Fluorine is intensely electronegative.

It just yanks electron density through the bonds, away from the carboxylate group, stabilizing the anion enormously.

The pKa of TFA plummets to around negative one.

It becomes a very strong acid purely due to that inductive pull.

Right.

Okay, now for something maybe a bit surprising.

Carbon acids.

We usually think of CH bonds as, well, not acidic at all.

Methane, pKa 48, right?

Basically impossible to deprotonate.

But some CH bonds can be acidic.

Yes, it's a fascinating area and incredibly important for synthesis.

The prime example we run into is terminal alkynes.

Acetylene, the simplest alkyne, has a pKa of about 25.

Now, it's still very weak compared to water, but it's vastly more acidic than methane.

You can actually deprotonate acetylene using a strong enough base, like sodium amide and NH2, in liquid ammonia, what's the difference?

Hybridization again.

The CH bond in acetylene involves in spi -hybridized carbon.

When it loses the proton, the resulting negative charge, the lone pair, resides in the TAT orbital.

And spot orbitals have more sin character.

Exactly.

50 % character.

That means electrons are held closer to the nucleus, more stabilized, compared to the negative charge in a methylanian, which is in an sp3 orbital, only 25 % character.

So hybridization makes a huge difference for carbanion stability.

A massive difference.

And then, just like with oxygen and nitrogen acids, conjugation makes carbon acids even stronger.

Put a CH bond next to a carbonyl group, like an acetaldehyde.

The pKa drops down to around 13 .5.

When it loses that proton, it forms an enolate anion.

And that negative charge isn't really on the carbon.

It's delocalized, mostly onto the more electronegative oxygen atom via resonance.

Huge stabilization.

Ah, enolates.

Key intermediates.

Absolutely fundamental.

Even just putting a CH next to a nitro group, like a nitro -methane, makes it acidic enough, pKa around 10, that it will actually dissolve in aqueous sodium hydroxide.

So yes, while most CH bonds aren't acidic, those adjacent to pi systems or electron withdrawing groups can be.

And forming those carbanions is how we make most new carbon bonds in organic synthesis.

Okay, let's connect all this back to the real world again.

You mentioned drug design earlier.

One of the most compelling stories, I think, where pKa was just central is the development of the anti -ulcer drug samitidine.

Take a minute, right?

Oh, absolutely.

It's a fantastic case study in rational drug design.

A real triumph of medicinal chemistry.

And pKa was at the heart of it.

So, peptic ulcers, often caused by too much stomach acid.

And acid secretions, often triggered by histamine binding to specific receptors, H2 receptors, in the stomach lining.

The goal for the chemists, said Smith Kline, our French, led by Sir James Black, was to find an antagonist, something that would bind to that H2 receptor, but not activate it.

Block histamine from doing its job.

And histamine itself has that imidazole ring, so early attempts tried to mimic histamine.

They did.

And histamine at physiological pH is mostly protonated on that imidazole ring.

So one of the early leads was a molecule containing guanidine group, which as we know is very basic, pKaH around 14 .5.

But here's the problem.

That highly basic guanidine analog actually acted as an agonist.

It stimulated acid secretion, the opposite of what they wanted.

Oh dear.

Not ideal.

Not ideal at all.

But it led to a crucial insight.

They hypothesized that maybe it was the positive charge itself, the fact that it was protonated at physiological pH, that was causing this unwanted agonist activity.

Okay, so the new goal became, make a molecule that fits the receptor, but make it less basic so it's not protonated.

Exactly.

A brilliant strategic shift.

They needed to lower the basicity, so they started systematically modifying the structure.

They replaced the CNH part of the guanidine with ACS, making a theria.

Theres are much, much less basic than guanidines.

This led to compounds like bermamide and then mediamide.

Mediamide was actually quite effective as an antagonist, but unfortunately it had some side effects thought to be related to that theria group.

So close, but not quite there.

Close.

The final leap was amazing.

They needed to get rid of the theria but keep the lower basicity.

And here's the clever part.

They actually went back to a guanidine structure, but they engineered it.

They attached a strongly electron withdrawing group, a cyano group, next CN, directly onto the guanidine part.

Ah, using that substituent effect we talked about.

Precisely.

That electron withdrawing cyano group pulled electron density away, dramatically lowering the basicity of the guanidine group, ensuring it wouldn't be protonated at physiological pH.

And the result was cementidine.

The result was cementidine.

Incredibly effective, much safer, and it became the world's first blockbuster drug.

The first drug to reach a billion dollars in annual sales.

Sir James Black won the Nobel Prize for this work.

It's just such a powerful example.

None of that lifesaving innovation would have happened without a deep practical understanding of PK and how to manipulate it.

Incredible story.

Okay, before we wrap up, just one quick but important distinction.

We've spent all this time talking about acids and bases in terms of proton transfer Brinstead -Lowry theory, but there's another definition that's really important in organic chemistry too, right?

Yes, absolutely.

Lewis acids and bases.

It's a broader definition.

Lewis acids aren't proton donors.

They are electron pair acceptors.

And Lewis bases aren't proton acceptors.

They are pair donors.

Often Lewis acids are things like boron trifluoride BF3 or aluminum chloride LCl3.

They have an incomplete octet, an empty orbital that's hungry for an electron pair.

So like BF3 reacting with an ether, the ether's oxygen uses its lone pair to donate into BF3's empty orbital, forming a bond.

No protons involved at all.

Exactly that.

The ether is the Lewis base electron donor.

BF3 is the Lewis acid electron acceptor.

These Lewis acid -base interactions are fundamental to so many organic reactions.

They often act as catalysts, activating molecules by accepting electrons.

Think about Friedel -Crafts reactions, for example.

Absolutely vital concept.

Got it.

Okay, what a deep dive that was.

From something as seemingly simple as making aspirin dissolve all the way to Nobel Prize -winning drug design and really getting under the hood of molecular behavior.

These concepts, acidity, specificity, PKA, they really do underpin so much of organic chemistry and how we use it.

It really is like having a molecular crystal ball sometimes.

It really is.

And its foundational understanding lets us predict, control, and ultimately innovate.

So maybe a final thought for you, our listener, to mull over.

How might this fundamental grasp of acid -base properties continue to drive innovation?

Maybe in areas we haven't even touched on much.

Think about designing new materials with specific properties, or creating better catalysts for green chemistry, or even decoding really complex biological pathways.

The principles we've discussed today, they're just the starting point for so much more.

Absolutely.

Well, we really hope this deep dive has given you a clearer, maybe more connected understanding of these crucial concepts and perhaps a new appreciation for the power packed into that little term PKN.

Thank you so much for joining us on this exploration, and thank you as always for being part of the Last Minute Lecture family.