Chapter 16: Acid–Base Equilibria

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Welcome curious minds to the deep dive.

Ever wonder why lemons are so tangy?

Or why soap feels slippery?

Or even just how your own body stays in balance.

Exactly.

It all comes down to acids and bases, which are, well,

fundamental in chemistry.

Absolutely central.

Today, we're doing a deep dive into acid -base equilibria.

We've basically mined a top chemistry text, chemistry, the central science, to pull out the really key stuff for you.

Yeah, we want to give you the core insights, kind of a shortcut to getting a solid handle on this topic.

We'll look at where the ideas came from, what acids and bases are, water's weirdly important role.

The pH scale, strong versus weak ones, and even how their shape, their structure affects what they do.

It really boils down to what makes something acidic or basic and why, well, why you should care.

It pops up in more places than you might think.

Ready to jump in?

Let's do it.

So thinking about acids and bases, it wasn't always about molecules and ions, right?

People noticed them way back.

Oh, for sure.

It started with the obvious stuff.

Acids tasted sour thing vinegar, citrus fruits, and they changed the color of certain plant And bases.

Bases were often bitter and they had that characteristic slippery soapy feel.

Right.

And the name base, that comes from somewhere specific.

Yeah, it's kind of interesting.

It comes from an old sense of to bring low.

The idea was that bases lowered the properties of acids when you mix them.

Ah, like they canceled each other out, neutralization.

Exactly.

There was a huge clue early on and mix them just right.

And the sourness, the slipperiness, it all seems to disappear.

Okay, so that's the observation part.

How do we get to like a chemical definition?

Well, the first big leap was in the 1880s.

Svante Arrhenius came along.

He proposed that in water, acids are things that make H plus ions, hydrogen ions.

H plus C.

Got it.

And bases make OH ions, hydroxide ions.

Simple enough, like HCl, hydrochloric acid, makes H plus in water.

Right, that's an Arrhenius acid.

Yeah.

And something like nekio, sodium hydroxide, dissolves to give necklace and OH.

So it's an Arrhenius base.

Pretty revolutionary for the time.

But I guess the in water part was maybe a bit limiting.

That was its main drawback, yeah.

It only really worked for aqueous solutions.

So about 40 years later, 1923, Brinstead and Lowry came up with a broader idea.

Brinstead -Lowry?

That sounds familiar.

Independently, they both focused on the proton, the H plus ion.

They said acid -based reactions are all about proton transfer.

Transfer.

Okay, so Brinstead -Lowry acid is a proton donor.

Gives away an H plus air.

And a Brinstead -Lowry base is proton acceptor.

Takes the H plus ion.

Okay, that seems more flexible.

But wait, that H plus ion in water, is it really just a bare proton zipping around?

Seems unlikely.

You're right, it's not.

A bare proton, H plus ion, is tiny and super reactive.

It doesn't exist on its own in water for any meaningful time.

So what happens?

It immediately latches onto a water molecule, bang, forms H3O plus ion, which we call the hydronium ion.

Hydronium, H3O plus ion.

Yeah.

Now for convenience, chemists often just write H plus AQ, meaning hydroproton.

But H3O plus AQ is, well, chemically more accurate.

It's the proton -riding piggyback on a water molecule.

Okay, that makes sense.

So back to HCl in water, HCl gives its proton to a water molecule.

Precisely.

Forming H3O plus and the chorded ion, Cl.

So HCl is the acid, the proton donor.

And water acts as the base, the proton acceptor, huh?

Exactly.

And this definition works beyond water, too.

Think about ammonia gas, NH3, reacting with hydrogen chloride gas, HCl.

Okay.

HCl donates a proton to NH3, HCl is the acid, NH3 is the base.

They form solid ammonium chloride, NH4Cl.

That hazy stuff you sometimes see on lab glassware.

Often that's it.

Ah, interesting.

So gas phase acid -base reactions fit this model, too.

What about things that can, like, swing both ways, be an acid or a base?

Good question.

Those are called amphiprotic substances, amphium meaning both.

Like amphidions, living in water and on land.

Kind of, yeah.

Water itself is the classic example.

We just saw it acting as a base with HCl.

But if you put it with ammonia and H3, water actually donates a proton to ammonia.

So water becomes the acid in that case.

It does.

Another common one is the hydrogen carbonate ion, HCO3, which you find in baking soda.

It can donate its proton or accept another one.

Depends what it's reacting with.

Okay, amphiprotic.

Got it.

Now, I keep hearing about conjugate acid -base pairs.

What's the deal there?

Right.

It's simple, really.

When a Brinsted -Lowry acid donates its proton, what's left behind?

The rest of the molecule?

Yeah.

And that rest is now capable of accepting a proton back.

That's its conjugate base.

Ah, perfect.

And when a base accepts a proton, it forms its conjugate acid.

So the pair just differs by one single proton, H plus Plut -Ros,

like acetic acid, CH3COH, and the acetate ion, CH3COO.

That's a conjugate pair.

Acetic acid and acetate.

Makes sense.

And their strengths are related somehow.

Very much so.

And it's an inverse relationship.

Super important.

The stronger the acid, the weaker its conjugate base is.

Weaker base, meaning it's not very good at accepting a proton back.

Exactly.

Take a really strong acid, like HCl.

It gives up its proton completely.

Its conjugate base, the chloride ion Cl, has practically zero tendency to grab a proton in water.

Okay.

It's negligibly basic.

But a weak acid, like that acetic acid, only gives up some of its protons.

Its conjugate base, the acetate ion, is actually a reasonably decent weak base.

It could accept protons back.

The strong acid, garbage conjugate base.

Weak acid, okay -ish, weak conjugate base.

You got it.

And this also explains something called the leveling effect in water.

Basically, any acid stronger than H3O plus just reacts with water to make H3O plus high.

And any base stronger than OH reacts with water to make OH.

So water kind of limits how strong an acid or base can effectively be in water?

Pretty much.

H3O plus and OH are the strongest acid and base that can exist at equilibrium in water.

That's actually really neat.

So does this help predict which way a reaction will go?

Will the protons actually transfer?

Yes, absolutely.

Equilibrium always favors the reaction direction that goes from the stronger acid and stronger base towards the weaker acid and weaker base.

Stronger things react to form weaker things.

Essentially, yes.

It's like the proton prefers to be on the weaker base.

If you mix acetic acid with water, we said acetate is a stronger base than water.

So the reaction prefers to go backward, keeping the proton on the acetic acid.

Most of it stays undissociated.

Okay, that logic follows.

Let's shift gears a bit and talk more about water.

It's not just a solvent.

It does something weird on its own.

It does.

It's called auto ionization, self ionization.

Even in perfectly pure water, a tiny fraction of water molecules react with each other.

One water molecule acts as an acid and donates a proton to another water molecule, which acts as a base.

So H2O plus H2O gives H3O plus an OH.

Exactly.

H3O plus an OH.

Now, the extent is incredibly small.

At room temperature, maybe only two molecules in a billion are ionized at any instant.

Wow, tiny.

Does it even matter then?

Oh, it matters hugely.

This tiny equilibrium is always there, and it's the basis for the pH scale and how we think about acidity in water.

Okay, how do we quantify that?

We use the ion product constant for water, KW.

It's the concentration of H3O plus times the concentration of OH.

KWH H3O plus OH.

Right.

And at 25 degrees Celsius, that value is always 1 .0 by 1014.

Always.

Always 1014.

Even if I add acid or base to the water?

Yes.

That product is constant in any aqueous solution at that temperature.

If you add acid, H3O plus goes up, but OH must go down proportionally to keep the product KW.

It's a fundamental balance.

That's powerful.

So if I know one, I know the other.

But those numbers, 107, 1014, they're kind of awkward to work with.

Extremely awkward, which is why chemists developed the pH scale.

It's a logarithmic scale to make these tiny numbers manageable.

Logarithmic.

Okay.

How does it work?

pH is defined as the negative logarithm, base 10, of the hydronium ion concentration.

pH equals log H3O plus.

Logative log.

So if H3O plus is 1 .0 by 10 to 7m, like in pure water.

The log of 10, 7 is negative 7.

The negative of that is 7.

And so the pH of neutral water at 25 degrees C is exactly 7 .0.

And if it's acidic, H3O plus is higher than 10, 7, like 10, 5.

Right.

Then the pH would be NAS log 10 to 5, which is 5.

Lower pH means higher acidity.

Okay.

Lower pH equals more acidic.

And higher pH above 7 means basic.

Correct.

A basic solution has H3O plus lower than 10, 7, maybe 10, 9, giving you pH of 9.

Higher pH means more basic, less acidic.

This scale is everywhere.

I know, for instance, our blood pH has to stay in a super tight range, right?

Like 7 .35 to 7 .45.

Incredibly tight.

And that's critical.

Almost all the biochemical reactions in our body, especially those involving enzymes, are extremely sensitive to pH.

What happens if it goes outside that range?

Bad things.

Proteins can denature, metabolic processes get disrupted, even small deviations can lead to serious illness or worse.

Our bodies have really sophisticated buffer systems.

Just to keep that blood pH stable, it's vital.

Wow.

Okay.

So pH is crucial.

Is there a pOH, too?

Yep.

Same idea.

pOH, it is log OH.

And because KwH3O plus OH is 10, 14, there's a simple relationship at 25 degrees Cp plus pOH always equals 14 .00.

Very handy sometimes.

Nice.

How do we actually measure pH, not just calculate it?

Well, the most accurate way is using a pH meter.

It has electrodes you put in the solution, and it measures a voltage difference that's directly related to the H3O plus concentration.

Okay.

And the less precise way.

Acid -based indicators.

These are special dyes that change color depending on the pH.

Lipmus paper is a common one, red in acid, blue in base.

Phenolphthalein is another colorless in acid, pink in base.

They change color over a specific pH range.

Right, like those pool testing kits.

Okay, so we know what acids bases are, how water autocationizes, and how pH measures acidity.

Let's dig into strength.

You mentioned strong acids earlier.

Right.

Strong acids are the ones that we assume ionize completely in water, 100 % dissociation.

They're strong electrolytes.

Like HCl again?

HCl, HBr, HI, nitric acid H103, perchloric acid HClO4, chloric acid HClO3, and the first proton from sulfuric acid H2SO4.

Those are the main ones to know.

Only seven common ones.

Pretty much, yeah.

And because they ionize completely, calculating the pH is usually straightforward.

If you have a 0 .1m solution of HCl, the H3O plus is also 0 .1m.

Just take the negative log.

Unless it's super dilute, I guess.

Ah, good point.

If it's extremely dilute, like 10 o a .m., then water's own autocontinization starts to contribute noticeably, but usually we ignore that.

Sulfuric acid, you mentioned it's produced in huge amounts.

Massive amounts.

Globally, it's one of the most produced industrial chemicals.

Fertilizers, batteries, refining petroleum, wastewater treatment.

Its strength makes it incredibly versatile.

Okay.

What about strong bases?

Similar idea.

Complete dissociation in water, producing hydroxide ions.

The main ones are the hydroxides of the alkaline metals group 1,

like NaOH,

sodium hydroxide,

and KOH, potassium hydroxide.

Alive, basically.

Yeah.

And also the hydroxides of the heavier alkaline earth metals group 2, like KOH2, SerOH2.

These are all ionic compounds that just dissolve and release OH.

Anything else act as a strong base?

Well, some metal oxides, like Na2O or CaO, quicklime, react very vigorously with water to form the hydroxide ions, so they effectively create strongly basic solutions too.

Okay.

So those are the strong ones.

Complete dissociation.

But you said most acids are weak.

By far, yes.

Most acidic substances we encounter are weak acids.

They only partially ionize in water.

They set up an equilibrium.

Equilibrium.

So not a one -way street like strong acids.

Exactly.

We use the acid dissociation constant, Ca, to describe that equilibrium.

Ca equals H3O plus AHA, where HA is the weak acid and A is its conjugate base.

And a bigger Ca means?

A stronger weak acid.

More dissociation, higher H3O plus for a given concentration.

Think vinegar, cetic acid, citric acid and lemons, scorbic acid, vitamin C.

These are all weak acids.

Crucial in foods and also in metabolism, like the Krebs cycle.

You mentioned percent ionization before.

How does that relate to weak acids?

Percent ionization is just the percentage of the original acid molecules that have actually donated a proton at equilibrium.

It's H3O plus at equilibrium, initial HA, 100%.

And here's a slightly counterintuitive thing.

For a weak acid, the percent ionization actually decreases as you make the solution more concentrated.

Really?

Why?

Le Chatier's principle, basically.

As concentration increases, the equilibrium shifts slightly back toward the undissociated acid to counteract the status.

The absolute H3O plus still goes up, but the percentage that dissociates goes down.

Huh.

Interesting.

Calculating pH for weak acids must be more complex, then.

A bit, yeah.

You typically set up an ICE table, Initial Change Equilibrium, use the CHI expression, and often make an assumption that the amount dissociating is small compared to the initial concentration, which simplifies the math.

Okay.

What about acids that have more than one proton to give?

Polyproduct acids.

Right, like sulfuric acid, H2SO4, phosphoric acid, HCPO4, or citric acid.

They lose their protons one step at a time.

Stepwise dissociation.

Yeah.

And each step has its own CHI value.

CHI1 for the first proton, K2 for the second, K3 for the third, if there is one.

Are the CHI values similar?

Not at all.

They get progressively much smaller.

CHI1 is usually way bigger than K2, which is way bigger than K3.

It gets harder and harder to pull off each subsequent proton because you're pulling it away from an increasingly negative ion.

So for pH calculations, does that mean one step usually dominates?

Generally yes.

For most polyprotic acids, the first ionization, K1, produces the vast majority of the H3O plus Alma.

So you can often approximate the pH just based on that first step.

The later steps contribute much less.

Okay.

That simplifies things.

Let's flip to weak bases now.

How do they work?

Weak bases are substances that accept protons from water, but only partially.

They react with water in an equilibrium to produce their conjugate acid and hydroxide ions.

So they make the solution basic by generating OH.

Correct.

And we quantify their strength using the base association constant, Kb.

Similar to the Ka, a larger Kb means a stronger weak base.

What kinds of things are weak bases?

Two main types.

First, neutral molecules that have a lone pair of electrons, usually on a nitrogen atom.

Ammonia in H3 is the classic example.

Amines, which are like organic derivatives of ammonia, are also common weak bases.

Amines, like in smelly things.

Sometimes yeah.

The smell of decaying fish or tissue is often due to amines like putrescine and cadaverine.

But many drugs are amines to codeine, amphetamine, lots of them.

They're often sold as their acid salts, like amphetamine hydrochloride, because that form is more stable and water soluble.

And the second type of weak base.

Amines that are the conjugate bases of weak acids.

We talked about this earlier, like the acetate ion, CH3COO, the fluoride ion, F, or the carbonate ion, CO32.

If the acid is weak, its conjugate base will be a weak base.

Okay, that connects back nicely.

Is there a mathematical link between Kali of an acid and Kb of its conjugate base?

There is, and it's super important.

For any conjugate acid -base pair,

K times Kb always equals Kw.

K by b, which is 1 .0 by 1014 at 25 degrees.

Exactly.

This equation powerfully shows that inverse relationship.

If K goes up, stronger acid, Kb must go down, weaker conjugate base, to keep the product constant.

It lets you calculate one if you know the other.

We can also talk about pKa and pKb, right?

Negative logs again?

Yep.

pKa equals log ka, and pKb equals log ka.

And the relationship becomes even simpler.

pKa plus pKb equals pKb equals 14 .0 at 25 degrees.

Smaller pKa means stronger acid, smaller pKb means stronger base.

Very useful.

Okay, this next bit surprised me when I first learned it.

Salt solutions aren't always neutral pH 7.

There was a huge point.

People hear salt and think NaCO neutral.

But most salts are strong electrolytes, meaning they break up completely into ions and water.

Like potassium acetate KCH3COO breaks into K plus and CH3COO?

Right now you have to ask,

can either of those ions react with water?

This reaction is called hydrolysis.

Hydrolysis.

Okay, so which ions react?

Think about where they came from.

K plus comes from KOH, the strong base.

Ions from strong bases like alkali metals, heavier alkaline earth, are just spectator ions.

They don't react with water, don't affect pH.

So K plus is neutral.

What about the acetate ions, CH3COO?

Acetate is the conjugate base of acetic acid, CH3COH, which is a weak acid.

Ah, so if the acid was weak, the conjugate base is?

A weak base.

So acetate ions will react with water, CH3COO plus H2O gives CH3COH plus OH.

It produces hydroxide.

So a potassium acetate solution would be basic.

Yes, slightly basic.

Because the anion hydrolyzes to produce OH.

What if the cation came from a weak base?

Like ammonium chloride, NH4Cl.

Good example.

Cl comes from HCl, strong acid, so it's spectator.

But NH4 plus E, ammonium, is the conjugate acid of ammonia, NH3, a weak base.

So the conjugate acid of a weak base is a weak acid.

You got it.

NH4 plus will donate a proton to water, NH4 plus plus H2O gives NH3 plus H3O plus blacca.

It produces hydronium ions.

Making the solution acidic.

So NH4Cl solution is acidic.

Correct.

What if both ions can react?

Like ammonium acetate.

Ah, then it's a tug of war.

You have the K of the cation NH4 plus trying to make it acidic, K, and the anion CH3COO trying to make it basic, KV.

The overall pH depends on whether the K of the cation is bigger or the KV of the anion is bigger.

Whichever is stronger wins.

Essentially, yes.

If K of the cation were acidic, if they happen to be roughly equal, it might be close to neutral.

And there's one more type of workation that makes things acidic, small metal ones.

Right.

Small, highly charged metal cations like aluminum 3, Al3 plus 3, or iron 3, F3 plus sobit.

They aren't proton donors themselves, but they are highly attractive to the oxygen end of water molecules.

They pull water close.

Very close.

They coordinate water molecules around them.

This strong attraction pulls electron density away from the OH bonds within the water molecules, making those hydrogens more acidic.

So the hydrated metal ion complex can donate a proton from one of its water ligands.

Exactly.

AlH2O63 plus plus H2O gives AlH2O5OH2 plus plus H3O plus.

It generates H3O plus fovint, making the solution acidic.

And the higher the charge and smaller the size of the metal ion, the more acidic it tends to be.

This explains the hydrangea thing, right?

Blue flowers and acidic soil.

Precisely.

Acidic soil often has soluble aluminum.

The aluminum ions get taken up by the plant, leading to the blue pigment formation.

In neutral or basic soil, aluminum isn't as soluble, and the flowers are pink.

Fascinating connection.

Okay, let's zoom in even further.

How does the actual structure of a molecule dictate its acid -base properties?

It's all about how easily that proton, H +, can be removed and how stable the resulting conjugate base is.

Three main factors matter for an acid HA.

Okay, factor one.

Bond polarity.

The HA bond needs to be polar, with H having a partial positive charge.

More polarity generally helps, but it's not the whole story.

Sector two.

Bond strength.

A weaker HA bond is easier to break, meaning the proton comes off more easily.

This is actually more dominant than polarity sometimes.

Think about the hydrogen halides, HF, HTL, HBr, HI.

HF is the most polar, but it's a weak acid.

Exactly, because the HF bond is incredibly strong.

As you go down the group, the HX barn gets weaker.

Larger atom size.

So HEL, HBR, and HI are all strong acids, with HI being the strongest.

Bond strength wins there.

And factor three.

Stability of the conjugate base.

If the anion A that's left behind is very stable, the acid HA will be more willing to lose its proton.

Stability often comes from spreading out that negative charge, maybe through resonance or onto a more electronegative atom.

So for simple binary acids like HA, going across a period like CH4, NH3, H2O, HF.

Acidity increases dramatically.

Electronegativity increases, making the bond more polar and stabilizing the negative charge of the resulting anion.

F is more stable than OH, et cetera.

And going down a group like H2O, H2S, H2SE.

Acidity increases because the HA bond strength decreases as the central atom gets bigger.

H2S is a stronger acid than H2O.

Okay, what about oxyacids, like HOY?

Here we look at the central atom Y.

If Y is very electronegative, it pulls electron density away from the OH bond, making it more polar and easier to break.

So HClO is stronger than HPO, which is stronger than HIO.

Electronegativity of Y matters.

What else?

The number of extra oxygen atoms attached directly to Y.

Compare hypochlorous acid HClO, chlorous acid HClO2, chloric acid HClO3, and perchloric acid HClO4.

More oxygens, stronger acid.

Much stronger.

Each extra oxygen is highly electronegative and pulls electron density away.

Weakening the OH bond and crucially, stabilizing the conjugate base anion by spreading out the negative charge through resonance.

Each HClO4 is one of the strongest acids now.

That makes sense.

And carboxylic acids, RCOOH, what makes them acidic?

Two things.

First, the oxygen in the CO double bond pulls electron density from the OH bond.

Second, and maybe more importantly,

the carboxylate anion, RCOO, formed when the proton leaves is highly stabilized by resonance.

Resonance stabilization.

The negative charge is shared between the two oxygen atoms.

Exactly.

That stability makes the proton easier to lose.

And this structure is key in biology.

Amino acids have both this acidic NECOH group and a basic NH2 group.

Right.

They're amphiprotic because of that.

They are.

And in neutral solution, they actually exist mainly as sweterians.

The protons in the NECO group transfers to the anti -NH2 group within the same molecule.

So you have COO and NH3 plus on the same molecule.

A molecule with both positive and negative charges.

Yep.

Like a tiny internal salt.

This explains why amino acids are typically solids with high melting points and are quite soluble in water.

Amazing how structure dictates function.

Okay, we've covered Arrhenius, Brinstead, Lowry.

But there's one more, even broader definition.

Lewis acids and bases.

That's right.

G and Lewis focus on electron pairs, not protons.

Electron pairs.

Okay.

A Lewis acid is an electron pair acceptor.

A Lewis base is an electron pair donor.

Donor and acceptor of electron pairs.

How does this relate to Brinstead -Lowry?

Well, any Brinstead -Lowry base proton acceptor must have a lone pair of electrons to form a bond with the proton.

So all Brinstead -Lowry bases are also Lewis bases, electron pair donors.

Okay, bases fit.

What about acids?

The Lewis definition of an acid is much broader.

It includes the proton, H plus, of course, because H plus needs an electron pair to form a bond.

But it also includes things that don't have protons at all.

Like what?

Molecules with an incomplete octet, like BF3.

Boron only has six electrons, so it readily accepts an electron pair from a Lewis base like ammonia.

BF3 is a Lewis acid, NH3 is a Lewis base.

No proton transfer needed.

Exactly.

Also, many simple metal locations, especially those highly charged ones like AL3 plus or BF3 plus voom, have empty orbitals and can accept electron pairs.

That's how they coordinate water molecules.

The water acts as a Lewis base, donating electrons to the metal Lewis acid.

Ah, so that ties back to why those hydrated metal ions make solutions acidic.

The metal ion acting as a Lewis acid weakens the OH bonds in the water.

Precisely.

The Lewis concept ties together a huge range of reactions,

including complex formation and reactions in nonaqueous solvents that don't fit the Brinstead -Lowry model.

It's a very powerful, unifying idea in chemistry.

Wow.

Okay, that really broadens the perspective.

We've gone from tasting lemons to electron pair donation.

It's quite a journey through acid -base equilibria.

It really covers a lot, from historical ideas to the nitty -gritty of molecular structure and bonding.

And it underlines how this balance, this equilibrium between H plus and OH, or between proton donors and acceptors, or even electron pair donors and acceptors, is just fundamental.

It's happening in your food, in the soil, in industrial processes, and critically, inside your own body.

These often tiny concentrations dictate so much of the world around us.

So maybe the next time you experience something sour or slippery, or just think about how life works, you can appreciate the hidden chemistry, that delicate acid -base dance happening at the molecular level.

What other chemical balances, maybe just as subtle, are shaping your everyday experience without you even realizing it?

It's definitely something to think about.

Chemistry is everywhere.

Keep asking those questions.

Keep observing.

Thanks so much for joining us on this Deep Dive.

Thank you.