Chapter 4: Acids and Bases

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Welcome to the Deep Dive.

Today we're plunging into the core of chemistry,

acids and bases.

Ever wondered why some chemical reactions are just incredibly dramatic, fizzing and popping, while others are maybe more subtle, but still profoundly powerful?

Yeah, and it often comes down to these unsung heroes of the molecular world, acids and bases.

Absolutely.

So what's the plan for today?

Okay, so for this deep dive, our mission is basically to guide you through a key chapter from Shriver and Atkins' Inorganic Chemistry, the fifth edition.

We're going to try and demystify

these fundamental concepts.

Step by step, right.

So even without the book in front of you.

Exactly.

So you can really grasp the ideas.

By the end, hopefully you'll have a much clearer understanding of what acids and bases really are, how we measure their strength and, you know, why they're just so crucial everywhere in science.

Sounds like quite a journey.

Where are we starting?

Well, we'll kick off with how people first understood these terms, which was surprisingly risky, actually.

Yeah, then we'll quickly move to the modern definitions like Brunsted -Lowry and Lewis.

We'll look at how to quantify strength, how solvents totally change the game.

Change the game.

I like that.

And applications too.

Oh, yeah.

We'll touch on some pretty wild real -world stuff from, say, industrial catalysts right through to these mind -bending super acids.

Okay, cool.

So this deep dive is really for anyone, right?

If you're maybe prepping for an exam or just trying to connect some chemical dots.

We're just curious, yeah.

We want to help you grasp the core ideas quickly and maybe have a few of those aha moments along the way.

Excellent.

Let's get started then.

Okay.

So this chemical evolution, you mentioned the early understanding was a bit risky.

Primitive is maybe a better word.

Early chemists literally identified acids by a sour taste.

Seriously.

And bases by their soapy feel.

Not exactly good lab practice today.

Definitely not.

A more scientific approach really started with Svante Arrhenius around 1884.

He defined acids as things that produce hydrogen ions, HU, when you dissolve them in water.

Okay, that sounds more like chemistry.

A good start.

A good start, yeah, but limited.

Because it only really worked for water.

Chemistry, you know, it moved on pretty fast.

As it does.

So what was the next big leap?

That came in 1923.

Johannes Brunstedt and Thomas Lowry, working independently, proposed a much broader idea based on proton transfer.

Proton transfer.

Okay, so moving away from just making hay and water.

Exactly.

In their view, a Brunstedt acid is simply a proton donor, and a Brunstedt base is a proton acceptor.

And by proton, we mean heishi, a hydrogen ion.

Correct.

But here's the crucial bit.

This proton, this heishi, it never actually exists all by itself in solution, especially not in water.

Right, it's too reactive.

It latches onto water molecules immediately.

We often write it simply as the hydronium ion for convenience.

Okay.

But really, you should picture it more like, well, like the proton is rapidly hopping between water molecules, almost like a hot potato.

It gets passed around, forming bigger clusters like asho or even larger structures.

It's a dynamic network, not just a simple asho ion sitting there.

That's a much better picture.

Like it dissolves into the water structure.

Can you give us a concrete example of this proton transfer?

Sure.

Let's take hydrogen fluoride, HF, in water.

HF is the acid.

So it donates its proton to a water molecule.

Okay.

Water, in this case, acts as the base, accepting the proton.

What you end up with is asho, the hydronium ion, and A, the fluoride ion.

HF gives asho, got it.

What about a base example?

Ammonia and asho, in water.

Ammonia acts as the base.

It accepts a proton from a water molecule.

So this time, water is the acid.

Exactly.

Water donates a proton, forming the ammonium ion, NHO, and the hydroxide ion, OA.

Interesting.

So water can play both roles.

Precisely.

That makes water what we call an amphiprotic substance.

It can be an acid or a base, just depends on what else is in the beaker, you know, who it's dancing with.

Amphiprotic.

Good term.

And this whole proton transfer thing, that leads to conjugates, right?

Conjugate acids and bases.

Yes, absolutely.

When an acid, like RhF, donates its proton, what's left behind the Ever ion is called its conjugate base.

Okay.

Acid loses proton, becomes conjugate base.

And when a base, like water in that first example, accepts a proton, it becomes AJO, which is its conjugate acid.

Base gains proton, becomes conjugate acid.

It's like they swap identities.

It is.

There's always this pairing.

Acid sharos plus base sharos reacts to form acid post -Brisos.

Every acid has a partner base, and vice versa, linked just by that transfer of a single proton.

Okay, so we've defined them by proton swapping, but like we said, some do this much more readily than others.

How do we measure that?

Put numbers on strength.

Right, we need to quantify it.

That's where equilibrium constants come in.

For an acid reacting with water, we use the acidity constant, Ka.

Ka.

Ka basically tells you how much the acid wants to donate its proton in water.

So a large Ka means it's a strong acid, it dissociates a lot, gives up the proton easily.

And a small Ka.

A small Ka, say much less than one, means it's a weak acid.

It prefers to hold on to its proton.

The equilibrium lies more towards the undissociated acid.

Makes sense.

And there's a similar thing for bases.

Yep.

The basicity constant, Kb, measures how readily a base accepts a proton from water.

Large Kb, strong base.

Small Kb, weak base.

Right.

And water itself.

You mentioned it can be both acid and base.

Does it react with itself?

It does.

It's called autoproteolysis.

One water molecule can donate a proton to another water molecule.

So water acts as both acid and base in the same reaction.

Exactly.

Two HO forms Heisho and Oho.

This happens even in pure water, but only to a very tiny extent.

How tiny?

Well, this reaction has its own equilibrium constant, the autoproteolysis constant of water, Kw.

At 25 degrees C, Kw is 1 .00 times 10 to the minus 14.

Wow, that is tiny.

10 -0.

Very small.

It means in pure water, the concentrations of HO and OO are extremely low.

Just 10 -OO over each.

That K value feels really important.

Does it link up with Ktob somehow?

It links them fundamentally.

For any conjugate acid -base pair in water, K time Kb always equals Kw.

Always.

Ktob Kw.

Always at a given temperature.

So if you know the K of an acid, you immediately know the Kb of its conjugate base, just by dividing Kit -Obiobico.

That's neat.

What does that tell us about strengths?

It tells us something crucial.

The stronger an acid is, meaning larger Ka, the weaker its conjugate base must be.

Smaller Ka.

And vice versa.

They have an inverse relationship.

Oh, okay.

Strong acid means its leftover bit, the conjugate base,

isn't very good at grabbing a proton back.

Exactly.

And for convenience, chemists often use p -values, pK, pKb, tKw.

The p just means take the negative logarithm.

Right, like pH.

Precisely.

So because Ka -S -Kalbi equals Ka -W, it follows that pKk plus pKb equals pKw.

And since pKw is 14 .00 at 25 degrees C, we have pKi plus pKb equals 14.

So knowing pK immediately gives you pKb.

Yeah.

That makes comparing strengths much easier, especially across big ranges.

It really does.

It avoids dealing with lots of fiddly powers of 10.

So based on these pK values, how do we draw the line between strong and weak?

Generally, a strong acid is one with a pKi less than zero.

That means its Ka is greater than one.

Think HCl, sulfuric acid.

In water, they're essentially 100 % deprotonated, existing as asho and the enum.

Fully ionized and weak acids.

Weak acids have pKa greater than zero, so Ka is less than one.

A sootic acid, like in vinegar, is a classic example.

It only partially deprotonates in water.

There's an equilibrium between the acid form and its ions.

Okay.

And bases follow suit.

Pretty much.

Strong bases react almost completely with water to generate a while.

Oxide ions, for example.

Weak bases, like ammonia, only partially accept a proton from water, setting up an equilibrium.

Gotcha.

Now some acids, like sulfuric acid you mentioned, have more than one proton they can donate, right?

Correct.

Those are polyprotic acids.

Sulfuric acid, AESOO, is dipartic, two protons.

Phosphoric acid, HPO, is tripartic.

And is it just as easy to remove the second or third proton?

Almost never.

The second acidity constant, KaOI, is nearly always much smaller than the first, KaO, and KaO is smaller still.

Why is that?

Think about it electrostatically.

Once HHEEKO loses one proton, it becomes HSO, which is negatively charged.

Pulling off another positive proton from something already negative is much harder.

Ah, that makes intuitive sense.

More attraction holding it back.

Exactly.

You can actually visualize this with something called a distribution diagram.

If you plot the fraction of each species, like HPO, HPO, HPO versus pH, you see distinct regions where each form dominates.

HPO dominates at very low pH, PRO at very high pH, and the intermediate ones in between centered around their respective pico -A values.

Okay, that helps picture the success of deprotonations.

So we know how to define them, how to measure strength.

But what makes an acid strong or weak at the molecular level?

What's the fundamental reason?

Good question.

To really dig into that, it helps to first think about acidity in the gas phase without any solvent getting involved.

Okay, isolating the molecule itself.

Right.

Here we talk about proton affinity, often symbolized AP.

It's basically a measure of how strongly a base binds to a naked proton in the gas phase.

High proton affinity means it's a strong base.

It really wants that proton.

So low proton affinity for the conjugate base means the original acid was strong in the gas phase.

Exactly.

If the conjugate base doesn't hold the proton tightly, low AP, the acid must have been good at donating it.

We see trends across the periodic table here.

Going across a period, gas phase acidity generally increases.

Why is that?

Because the atoms get more electronegative, they stabilize the negative charge on the conjugate base better.

Think about comparing CH, NH, HO, HF.

HF is the strongest acid in that series.

Okay.

And down a group?

Down a group, acidity also generally increases.

But here, the dominant factor is usually the bond strength.

As you go down, the bond between hydrogen and the other atom gets longer and weaker, easier to break.

So HI is a much stronger acid than HF in the gas phase.

Easier HA bond breaking dominates.

Okay, that's gas phase.

But chemistry usually happens in solution.

How does the solvent, like water, change things?

The solvent changes everything and introduces solvation.

When an ion is formed, solvent molecules surround it, interact with it, stabilize it.

This is usually a very favorable process, energetically speaking.

Like the solvent gives the ion a comforting hug.

Huh.

Sort of, yeah.

A very energetic hug.

Especially for polar solvents, like water interacting with small, highly charged ions.

Water is particularly good at this because of its polarity and its ability to hydrogen bond.

So the stability of the ions in solution is key.

Absolutely critical.

The overall strength of an acid in solution depends on a balance.

How easy it is to break the HA bond and how well the resulting HU and A ions are stabilized by the

Water is exceptionally good at stabilizing HU as HO and larger clusters, and also very good at stabilizing small ions like Esch or OA through hydrogen bonding.

Sometimes the solvation energy completely outweighs the gas phase trends.

So water's stabilization power can reshuffle the order of acidity compared to the gas phase.

It definitely can.

Take the hydrogen halides.

In the gas phase, HI is stronger than HBR, which is stronger than HCL, which is stronger than HF.

But in water,

HI, HBR, and HCL are all equally strong acids.

They're completely leveled.

And HF is actually a weak acid in water.

Whoa, wait.

HF is weak in water, even though F is the most electronegative.

Yes.

Because the fluoride ion, A, is so small and forms such strong hydrogen bonds with water, it gets incredibly stabilized.

And the HF bond itself is very strong.

Those factors combined make HF reluctant to donate its proton in water, even though it's quite acidic in the gas phase.

That really highlights the solvents power, which brings us to this leveling effect you mentioned.

Right.

Water exerts a leveling effect on strong acids.

Any acid that's intrinsically stronger than HAO just ends up donating its proton completely to water, forming HO.

So in water, you can't tell the difference in strength between, say, HBR and HI.

They both just look like solutions of HO and the Halida ion.

Water sets the upper limit for acid strength you can observe.

Pretty much.

HO is the strongest acid that can effectively exist in water.

If you want to actually differentiate between acids like HBR and HI, you need to use a solvent that's a much weaker base than water.

Like what?

Like, say, pure acetic acid.

In acetic acid, HBR and HI don't fully dissociate, and you can measure their different equilibrium constants and see that HI really is stronger than HBR.

Okay, so choosing the right solvent lets you see the true relative strengths.

Does the same leveling happen for bases?

It does.

Any base stronger than the hydroxide ion, OA, will simply react completely with water to form OA.

So OA is the strongest base that can exist in water.

Correct.

Bases like the amide ion, NHRO, or hydride ion HO are incredibly strong, but in water, they just instantly rip a proton off water, generating OA.

Water levels them down to the strength of hydroxide.

This range of strengths a solvent can distinguish, you called it the discrimination window.

Yeah, the discrimination window.

For water, it runs roughly from the pKa of HO, around zero, to the pKa of HO acting as an acid, which relates to OA's formation around 14.

So a 14 pKa unit window.

And other solvents have different windows.

Exactly.

Liquid ammonia, for instance, is much more basic than water.

Its autoprotolysis constant, PCOM, is about 33.

So it has a much wider discrimination window, particularly on the basic end.

It can differentiate bases that would all be leveled to OA strength in water.

Fascinating.

Okay, but what about solvents that don't really do proton transfer, a product solvents?

How do we think about acids and bases there?

Good point.

For those, we often use the solvent system definition.

It's based on the idea that many of product solvents also auto -anionize, even if not via protons.

Like water -forming HOE and OO?

Analogous to that, yeah.

Take liquid bromine trifluoride, BRF.

It auto -anionizes slightly.

2 -BRF gives BRFU and BRFU.

Okay.

Anacation and anion from the solvent itself.

Right.

So in the BRFU solvent system, an acid is defined as anything that increases the concentration of the solvent's cation, BRFU, and a base is anything that increases the concentration of the solvent, anion BRFU.

So you're defining acidity -basicity relative to the solvent's own ions, even without protons.

Exactly.

It's a way to extend acid -based concepts into systems where Brensted -Lowry doesn't really apply directly.

Clever.

Okay, let's circle back to Brensted acids, the proton donors.

We can group them based on where the proton comes from, right?

We can.

It's useful to think about three main classes, especially those with protons on oxygen atoms.

First, aqua acids.

Aqua acids?

Proton on a water molecule.

Yes.

A water molecule that's coordinated, bonded, to a central metal ion, like in hexaquaron, 30 -arrow.

Those coordinated water molecules are more acidic than free water because the positive metal ion pulls electron density away.

Makes the protons easier to remove.

Okay, what's next?

Then hydroxo acids.

These have an acidic proton on a hydroxyl group, but there's no double -bonded oxygen atom right next to it on the same central atom.

TOH, telluric acid, is an example.

Okay, just AOH groups.

And the third type?

Oxo acids.

These are probably the most familiar.

They have the acidic proton on an OH group, and there's also at least one oxo group, a double -bonded oxygen attached to that same central atom.

Sulfuric acid, 8 -ARSO, which is O -SOH, is a classic oxo acid.

Right.

HSO, HNA, HPO, lots of common ones.

Exactly.

Generally found when the central element is in a high oxidation state.

Now, for those aqa acids, you mentioned the metal ion makes the water acidic.

Does the type of metal matter?

Oh, absolutely.

There are clear periodic trends.

Acidity increases as the positive charge, Z, on the metal ion increases, and as the ionic radius decreases.

So higher charge density means stronger pull on electrons, more acidic protons.

That's the main idea, yeah.

A small, highly charged ion like aloe makes its coordinated waters much more acidic than a larger, less charged ion like NaO, but it's not purely ionic.

For some metals, especially later in the D block or P block,

covalent bonding effects can make the aqa acid stronger than you'd predict just from charge and radius alone.

Interesting.

So covalency plays a role, too.

Yeah.

What about the oxo acids?

We said more oxo groups generally mean stronger acid.

Yes.

Those double bonded oxygens are strongly electron withdrawing.

They pull electron density away from the OH bond, weakening it, making the proton easier to lose.

They also help stabilize the negative charge on the conjugate base through resonance.

Makes sense.

And this is where Pauling's rules come in handy, right?

They do.

Pauling gave us two simple rules to estimate pico values for oxo acids.

Rule one is for the first pico.

If you write the formula as OPEOHQ, where P is the number of oxo groups.

The double bonded oxygens.

Right.

Then pico is approximately 8 minus 5 times P.

H5P.

Let's try sulfuric acid, OPEOH, so P2.

Rule says pico here E552.

Exactly.

Predicts it's a very strong acid, which it is.

For, say, phosphoric acid, OPEOHP equals 1, so pico 8551 yields 3.

Also pretty close to the experimental value of about 2 .1.

Wow.

That's surprisingly good for a simple rule.

What's rule two?

Rule two just says that for polyprotic oxo acids, the successive pico values increase by about 5 units each time.

So pico is roughly pico plus 5.

picos are roughly picos plus 5.

Also a useful guideline.

Are there exceptions where these rules don't quite work?

Oh, definitely.

Chemistry loves exceptions.

Harbonic acid and sulfurous acid are famous ones.

Their measured pico values are much higher, much weaker acids than Pauling's rules predict.

Why is that?

It turns out that when you dissolve kico or so in water, most of it just stays as dissolved gas molecules hydrated by water.

Only a tiny fraction actually forms the true HE or OO molecule.

The measured pK reflects the equilibrium with the dissolved gas, not the true acidity of the small amount of actual acid molecules present.

Ah.

So the concentration we think we have isn't the real concentration of the acid species.

Clever.

It highlights that we need to be careful about what's actually in the solution.

Okay.

Now, related to oxal acids are the anhydrous oxides.

Things like CaO or Aso.

How do they fit in?

Right.

These are oxides without any water involved initially.

The general trend is that metal oxides tend to be basic.

They react with water to form hydroxides, or they react directly with acids.

Think calcium oxide, CaO, reacting with water to give KOH.

Okay.

Metal oxides are basic.

And nonmetal oxides tend to be acidic.

They react with water to form oxoacids, or react with bases.

Sulfur trioxide, Sessio, reacts vigorously with water to form sulfuric acid, Ho.

Nonmetal oxides are acidic.

Is there anything in between?

Yes.

Amphoteric oxides.

These can react with both acids and bases.

Aluminum oxide, Allerose, is a classic example.

It dissolves in strong acid, and also in strong base.

You find amphoterism often near the diagonal line separating metals and nonmetals on the periodic table, and also for some D -block elements.

Like a chemical split personality.

You could say that.

And these acid -base properties lead to interesting structures.

Aqua ions can link up, losing water, to form polymers' polycations, especially as you raise the pH.

Like metal ions clumping together.

Yeah, often forming extended networks, or eventually precipitating as hydroxides.

Conversely, many simple oxoanions, like phosphate or chromate, can link up into polyoxoanions as you lower the pH.

So, anions link up in acidic conditions.

Right.

Think of chromate, C -R -O -S -G -O, turning into dichromate, orange, in acid, or phosphates linking into chains and rings.

This is hugely important biologically, the energy currency of life, ATP and ADP, involves polyphosphate chains breaking and reforming.

That energy transfer is fundamentally an acid -base driven condensation hydrolysis process.

Wow, connecting bench chemistry right to life itself.

Okay, we've spent a lot of time in water.

Why look at non -aqueous solvents?

Several key reasons.

Sometimes your molecule just reacts with water hydrolysis, so you need a different solvent.

Sometimes, as we discussed, water's leveling effect hides the true strength differences, so you need a less basic or less acidic solvent to see them.

And sometimes, it's just about solubility, your compound might dissolve better elsewhere.

What are some common non -aqueous solvents used for acid -base chemistry?

Oh, there are many.

Liquid ammonia is a classic, and hydrogen fluoride, HF, and hydrosulfuric acid, even things like dinitrogen tetroxide.

Let's take liquid ammonia.

What's it like?

Well, it's liquid below minus 33 degrees C, so you need to cool it.

It's a surprisingly good solvent for many salts and organic compounds, and, like water, it autoserinizes.

How does ammonia autoserinize?

Two NH molecules react to form the ammonium ion, NHO, the characteristic cation, the acid, and the amide ion, NHO, the characteristic anion, the base.

So acids in liquid ammonia increase NHO, bases increase NHO.

Exactly.

And because ammonia is inherently more basic than water, it's a great solvent for studying acids that are very weak in water.

Acetic acid, for example, which is weak in water, acts as a strong acid, almost fully ionized, in liquid ammonia.

It boosts their apparent acidity.

It does.

Ammonia is also famous for dissolving alkali metals, like sodium, to give these amazing deep blue solutions containing solvated electrons a really unusual form of reducing agent.

Cool.

What about hydrogen fluoride, HF, you said, and hydros?

Yes, pure liquid HF.

It boils around room temperature.

It's extremely corrosive, dangerous stuff, etches glass, penetrates tissue deeply.

It's a very acidic solvent.

How does it autoserinize?

It's not quite like water or ammonia, is it?

It's a bit more complex.

It's thought to be three HF molecules reacting to form HF, the solvated proton, the acid, and the hydrogen bonded HFO ion, the bifluoride ion, the base.

HF here is the base.

Interesting.

Because HF is so acidic itself, only extremely strong acids can actually donate a proton to HF.

Most things, even strong acids like nitric acid, actually act as bases in liquid HF.

They accept a proton from HF.

Wow.

Turns the tables.

And an hydrosulfuric acid.

Another highly acidic, very viscous solvent.

High dielectric constant complex autoserinization involving ions like HSO and HSOO.

It's so acidic that even things we normally consider acids like phosphoric acid act as bases they get protonated by the sulfuric acid.

Everything looks like a base compared to sulfuric acid.

Pretty much.

Its most famous use is probably in making the nitronium ion, NaOO, by reacting nitric acid with sulfuric acid.

That NaOO is the key ingredient for nitrating benzene rings in organic synthesis.

A classic example of using a non -Iqueous solvent to generate a reactive species you couldn't easily make in water.

Right, this really broadens the picture.

Which leads us perfectly into the Lewis definition of acids and bases.

This one doesn't even need protons, does it?

Not at all.

G .N.

Lewis proposed this around the same time as Brinstead and Lowry, 1923, but it's much more general.

A Lewis acid is an electron pair acceptor.

A Lewis base is an electron pair donor.

Electron pair acceptor, electron pair donor.

Simple, but broad.

Very broad.

The fundamental reaction is the formation of a coordinate covalent bond, where the base donates both electrons for the shared pair, forming an adduct or complex.

So how does this relate to Brinstead -Lowry?

Well, a proton has an empty orbital, so it's definitely an electron pair acceptor, it's a Lewis acid.

And a Brinstead base, by definition, has a lone pair available to accept that proton.

So it's an electron pair donor, a Lewis base.

So Brinstead -Lowry is just a special case of Lewis theory, focused on the proton as the acid.

Exactly.

But Lewis theory covers so much more.

Think about boron trifluoride, BF year.

Boron only has six electrons around it, an incomplete octet.

It desperately wants an electron pair.

So it's a Lewis acid.

Classic Lewis acid.

It reacts readily with ammonia Nase, which has a lone pair on the nitrogen.

Ammonia is the Lewis base.

They form an adduct, FUNH.

No protons involved at all.

Okay.

Incomplete octets are Lewis acids.

What else?

Metallocations are great Lewis acids.

Think of cobalt ions in water.

They accept lone pairs from six water molecules to form the complex ion co -uac.

Water is the Lewis base here.

Right.

Coordination complexes.

Also, molecules that might look like they have complete octets, but can rearrange electrons, like coase.

The carbon atom can accept a pair from a Lewis base, like OTO, to form the bicarbonate ion, HCO.

Ah, even coase acts as a Lewis acid.

And molecules where the central atom can expand its valence shell go beyond an octet.

Silicon tetrafluoride, CFA, can accept two more fluoride ions, Lewis bases, to form the hexafluorosilicate ion, CNO.

Silicon expands its coordination number.

So incomplete octets, metal ions, rearranging octets, expanding octets, lots of types of Lewis acids.

Are there trends here too?

Definitely.

In group 13, the boron trihalides are classic Lewis acids.

Interestingly, the acidity trend is BFALBCL -bibros.

Wait, fluorine is the most electronegative.

Shouldn't BFEAs be the strongest Lewis acid, pulling electrons the most?

You'd think so, but there's a competing effect.

Pi -bonding.

There's some double -bond character between boron and the halogens, strongest between boron and the small fluorine atom,

PP -pi overlap.

This internal pi -bonding stabilizes BFR and has to be partially broken when it forms an adduct.

The weaker pi -bonding in BCL and bibros means they are actually stronger Lewis acids more ready to accept that external lone pair.

Wow, that's counterintuitive, but makes sense of the pi -bonding.

Fascinating.

What about other groups?

Group 14, silicon compounds like Cifeli we mentioned.

Tin -2 -chloride SN -CLO is interesting.

It can act as both a Lewis acid, excepting quiu to form SN -olero, and a Lewis base, using its own lone pair to bond to other metals.

Amidextrous.

Group 15, phosphorus pentafluoride PFO is a strong Lewis acid.

Antimony pentafluoride SBFO is an extremely powerful Lewis acid, key for making super acids.

We'll come back to those.

Group 16, sulfur dioxide SO -uro can be both acid and base.

Sulfur trioxide zero is a very strong Lewis acid.

Even the halogens themselves, like boron as well, can act as mild Lewis acids forming complexes.

So Lewis theory really covers a huge range of interactions.

How do these acids and bases interact?

What kinds of reactions happen?

Three main types.

First, simple complex formation.

Acid plus base adduct,

like BF -euro plus NH.

Second, displacement reactions.

One Lewis base can displace another from an adduct, or one Lewis acid can displace another.

Ba plus B plus Ab.

Actually, Brinsted proton transfer is just a displacement reaction where the acid is hay.

Or acid kicks out acid.

Third is metathesis, or double displacement.

Ab plus Ab, Ab plus Ab partners get swapped, often driven by forming a particularly stable product or something insoluble.

Swapping partners, got it.

Now, within Lewis theory, is there a way to predict which acid prefers which base, beyond just saying they react?

Yes, and this is where the concept of hard and soft acids and bases, HSAB, comes in.

It's a really useful qualitative idea, developed mainly by Ralph Pearson.

Hard and soft, like textures.

Sort of an analogy.

Hard acids are typically small, highly charged, not easily polarized ions.

Think Allo, MgOu, X itself.

Soft acids are larger, have lower charge, or zero, and are easily polarized, their electron clouds are squishy.

Think HgU, PdOu, Allo.

Okay, small and stubborn versus large and squishy.

What about bases?

Same idea.

Hard bases are small, highly electronegative, not easily polarized.

They hold their electrons tightly.

Think Ao, Oa, Oaa.

Soft bases are larger, less electronegative, easily polarized.

Their electrons are more available, more soft.

Think Allo, Phosphenes.

Hard likes hard, soft likes soft.

Is that the idea?

That's the core principle.

Hard acids prefer to bind to hard bases, and soft acids prefer to bind to soft bases.

And what does that mean chemically?

Why does hard like hard?

Hard -hard interactions are thought to be primarily electrostatic or ionic in nature.

Positive attracting negative, charge driven.

Soft -soft interactions have a much larger covalent contribution, more about effective orbital overlap and electron sharing.

Ionic versus covalent preference, that makes sense.

Does this explain real world chemistry?

It explains a ton, like why certain metals are found as oxides, hard O -base in nature, while others are found as sulfides, soft SO -base.

Hard metal ions like aluminum and iron prefer the hard oxide.

Soft metals like mercury and lead prefer the soft sulfide.

It also helps predict reaction outcomes.

For instance, the thiocyanate ion, SCN, can bond either through the harder nitrogen or the softer sulfur.

Oh, an ambiditate ligand.

Exactly.

It bonds through N to a hard acid like silicon, but through S to a soft acid like platinum.

HSEB predicts that perfectly.

That's really elegant.

Is there a way to quantify this beyond just hard and soft?

There is.

One approach is the Drago -Wayland equation.

It's a semi -empirical equation that assigns two parameters to each acid, Ea and Ca, and two to each base, Eb and Cb.

E and C, electrostatic and covalent.

Roughly, yes.

Ea and Eb relate to the electrostatic contribution, while Ca and Cb relate to the covalent contribution.

The equation predicts the enthalpy change for forming the adduct.

Nefes He equals EaEb plus CaCb.

It works remarkably well for predicting bond strings for thousands of Lewis acid base pairs, especially in the gas phase or non -polar solvents.

A more quantitative handle on interaction strength.

Very useful.

Now, we keep mentioning solvents.

They can act as Lewis acids or bases too, right?

Absolutely.

Most common solvents are Lewis bases water, alcohols, ethers, DMSO.

They have lone pairs.

They're often hard bases and readily compete to coordinate to Lewis acids dissolved in them.

The solvent is always competing.

Always competing.

Some solvents can act as Lewis acids too, perhaps through hydrogen bonding like chloroform, CHCO, or other interactions.

Truly inert solvents like saturated fluorocarbons that have very weak Lewis acid base character are quite rare and specialized.

The solvent is never just a passive background.

Okay, let's wrap up with some cutting edge applications.

You mentioned super acids.

Right.

Super acids are systems that are much more acidic, much better proton donors than 100 % pure sulfuric acid.

More acidic than pure sulfuric acid.

How do you make that?

Usually by dissolving a very strong Lewis acid like antimony pentafluoride, SBFU, into a strong Brunsted acid like anhydrous HF or fluorosulfonic acid, HSOSF.

The Lewis acid pulls electron density away, making the proton on the Brunsted acid even more available.

So SBFU basically supercharges the HF or HSO here.

Exactly.

The mixture of SBFU and HSOS is famously called magic acid.

George Ola won the Nobel Prize in 1994, largely for his work using super acids like this to protonate things like alkanes, simple hydrocarbons, generating stable carbocations that could then be studied.

It opened up a whole new area of organic chemistry.

Protonating alkanes, wow.

And the opposite exists.

Super bases.

Yep.

These are proton acceptors, much stronger than the hydroxide ion.

Things like lithium nitride or metal hydrides like NOH or Coar, they react violently with water, ripping off protons to make OHAR.

They're used extensively in organic synthesis where you need an incredibly strong base to deprotonate very weak acids and also in areas like hydrogen storage research.

Extremes of the acid base world.

Finally, what about reactions happening not in solution but on surfaces?

Heterogeneous acid -base chemistry, hugely important, especially in catalysis.

Many solid materials like metal oxides, silica alumina, or mixed oxides like alumina silicates, think zeolites have acidic sites right on their surface.

Acid sites on a solid, like what?

They can have both Brunsted sites, like surface OH groups that can donate protons, and lewisites, like exposed metal ions that can accept electron pairs.

Zeolites are amazing examples.

They have these intricate channel structures with tunable, acidic sites inside.

And these are used as catalysts.

Massively.

Zeolites are workhorses in the petrochemical industry for cracking hydrocarbons and other reactions.

They are often much more environmentally friendly, greener than traditional solution -based catalysts because they are solid, easily separated, reusable, and can be highly selective, reducing waste.

Their acidity is key to their function.

So understanding surface acidity leads to better, greener catalysts.

Exactly.

It's a fantastic example of applying fundamental acid -base principles Brunsted, Lewis, even HSAB ideas about site preferences to design materials that solve real -world problems.

Wow, what a journey indeed.

We've gone from, well, literally tasting chemicals.

Not Rick Bended.

To the dance of protons in Brunsted -Lowry theory, then blew it wide open with Lewis's electron pairs.

It's amazing how the definitions evolved.

It really is.

We saw how to put numbers on strength with Ka and Kaib, how solvents aren't just backgrounds, but active players with leveling effects, and how structure dictates everything.

Then connecting it all with hard -soft theory and seeing the power in super acids in these crucial surface reactions.

Yeah, from the simplest proton transfer to complex catalysis, it all comes back to these fundamental acid -base interactions.

So maybe the final thought for our listeners.

The next time you see a chemical reaction, whether it's baking soda, fizzing with vinegar, or something complex in a lab, maybe pause and think.

What's the underlying acid -base chemistry?

Is it a proton hop?

An electron pair being shared?

Maybe even a subtle, hard -soft interaction driving things?

It's often the silent, fundamental driving force.

Always more to explore.

Hopefully this dive helped connect some dots.

It certainly did.

Thanks for guiding us through that.

And thank you all for joining us on this deep dive.

This has been a Last Minute Lecture Production.