Chapter 16: Acids and Bases

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Welcome back to the Deep Dive.

It's really great to have you with us today.

I actually want to start with something a little, well, a little mundane.

I was standing in the shower this morning and I was reading the back of my shampoo bottle because, you know, that's apparently what I do now to avoid waking up.

And I saw that phrase we have all seen a thousand times, pH balanced.

Oh yeah, it is everywhere.

Deodorants, skin creams, bottled water.

Exactly.

And it kind of struck me.

We use that term constantly.

We talk about acid reflux, acid rain, maybe even, I don't know, acidic coffee.

We know that lemons are sour because of citric acid.

But if you actually stop someone on the street and ask them, okay, but what physically makes an acid,

like what is actually happening on the molecular level?

You would probably get a lot of people talking about burning sensations or maybe remembering the litmus paper turning red from their high school science class.

Right, right.

We know the effects, but we don't really know the engine driving it.

And that is exactly what we are doing today.

We are taking chapter 16 of our general chemistry text, which is titled Simply Acids and Bases.

And we are just, we're tearing it down to the studs.

And it's a fascinating chapter because it is really the story of a tug of war.

It is not just about liquids and beakers changing colors.

It's about a literal battle for protons.

A battle for protons.

I like that image.

So our mission today for this deep dive is simple, but pretty deep.

We want to move past the memorized definitions.

We're going to understand the molecular architecture that makes a lemon sour, and more importantly, the math that tells us exactly how strong that sourness is.

It is arguably one of the most vital concepts in the entire textbook.

If you don't grasp how protons move and how equilibrium works in these systems, you really cannot do biology.

And you certainly cannot do organic chemistry later on.

So quick ground rules for our deep dive today.

We are sticking strictly to the text of chapter 16.

We aren't going to wander off into outside industrial applications or alternate methods that aren't in the book.

We want to build this narrative logically for you, exactly as the authors intended, starting with the definitions and moving all the way through to molecular structure.

Which means we should start right at the beginning,

sections 16 to 1,

acids, bases, and conjugate acid base pairs.

Let's set the stage a bit.

Throughout history, chemists have they've tried to pin down what an acid actually is.

I feel like the definition has shifted a few times over the century.

It really has.

It's actually a great example of how scientific theory evolves to become more and more useful.

If you go back to the 1700s, you have Antoine Lavoisier.

He's a father of modern chemistry, right?

But he actually got this one wrong.

He thought oxygen was the key element.

Against the name oxygen, right?

Exactly.

The word oxygen comes from Greek roots, meaning acid former.

He saw that things like sulfur and phosphorus burned in air to make acidic compounds.

So he just he blamed the oxygen.

But he was wrong.

Later on, Humphry Davy showed that hydrogen was actually the common element in all these acids.

And then fast forward to the late 1800s, and we get Svante Arrhenius.

Right.

Arrhenius.

He gave us the first really workable definition, but it was, well, it was limited.

He focused entirely on water.

He essentially said an acid is something that increases the concentration of hydrogen ions, so H plus, in water.

And a base is something that increases the concentration of hydroxide ions, OH minus, in water.

Which is true, isn't it?

I mean, hydrochloric acid in water does exactly that.

It is true, yeah, but it restricts you.

It treats water as the only solvent in the universe.

It's kind of like saying the word football only refers to the NFL.

The real breakthrough and the definition that dominates this entire chapter came in 1923 from two chemists working completely independently.

Johannes Brunstedt in Denmark and Thomas Lowry in Great Britain.

Ah, the Brunstedt -Lowry theory.

This is the big one.

This is the absolute workhorse of general chemistry.

They realized that you don't need water to have an acid -base reaction.

You just need a transfer.

They shifted the focus entirely to the proton, the hydrogen ion, H plus.

Okay, let's unpack this for you listening.

In the Brunstedt -Lowry world, what exactly is an acid and what is a base?

It is very simple, but you have to visualize the transaction.

An acid is a proton donor.

It's the philanthropist.

It holds a proton and it actively wants to give it away.

And the base?

The base is a pro -con acceptor.

It is the thief.

It wants to take that proton.

So every reaction is essentially a handoff.

One gives, one takes.

Precisely.

Acid -base chemistry is fundamentally about the transfer of proton from one species to another.

Let's look at the specific example the text gives right away.

Acetic acid, which is CH3COOH reacting with water.

Okay, visualizing this.

We have an acetic acid molecule.

This is the stuff in vinegar, right?

Right.

Picture the molecule.

It has a carboxyl group.

That's a carbon double bonded to an oxygen and single bonded to an OH group.

That hydrogen on the OH is the ionizable hydrogen.

It's loosely attached, kind of like a loose tooth.

So that specific hydrogen is the proton waiting to be donated.

Exactly.

Now in floats a water molecule.

You know, water is H2O.

The oxygen atom in water has what we call lone pairs.

These are two sets of electrons just hanging out, not involved in bonding.

They're highly negatively charged regions.

And the proton on the acid is positive.

Attraction.

If you look at figure 16 to 1 in the text, it uses those curved arrows we often see in mechanism diagrams.

This is really important for visualizing the how of the reaction.

The arrow starts at the lone pair on the water molecule.

Because that's the base.

And it points directly at the hydrogen atom on the acetic acid.

The water is actively reaching out, grabbing that proton and ripping it off the acid molecule.

That is a pretty violent image, but it really helps.

So the acetic acid donates the proton, making it the acid.

The water accepts the proton, making it the base.

Correct.

And this leads to a really important point about water that the text emphasizes.

In the old Arrhenius definition, water was just the solvent, just the background noise.

In Brinkstead -Lowry, water is a player.

It is an active participant in the chemical reaction.

So what happens after the handoff?

The acetic acid lost a proton and water gained one?

What are we left with?

The acetic acid becomes the acetate ion, CH3COO minus.

It now has a negative charge because it kept the electron that used to bind the hydrogen and the water.

It picked up a positive charge along with that proton.

It becomes H3O plus.

H3O plus.

The hydronium ion.

Yes, the hydronium ion.

And the text makes a massive point here that we need to stop and appreciate.

You will very often see chemists write just H plus in chemical equations because it's faster.

But physically, in a solution, a bare proton, a stripped hydrogen nucleus, does not exist.

Really?

Why not?

Think about what it is.

It's a nuclear particle with a focused positive charge.

But it has zero electrons to buffer it.

It is incredibly tiny and has a massive, massive charge density, like a magnet on steroids.

It will instantly latch onto the nearest source of negative charge.

Which, in an aqueous solution, is the lone pair of a water molecule.

Exactly.

So H plus immediately becomes H3O plus.

The text even references figure 16 -3 to show that it actually goes even further.

The hydronium ion itself gets surrounded by other water molecules through hydrogen bonding, forming these larger hydrated structures.

But for our calculation purposes, just remember, whenever we say proton transfer, we mean forming hydronium.

Okay, so we have the transfer.

Acetic acid gave to water.

But this reaction is reversible, isn't it?

It is.

And this brings us to the concept of conjugate acid -base pairs.

This is something students often trip up on, but it is really just looking at the before and after snapshots of the molecules.

Walk us through it.

How do we spot a pair?

A conjugate pair consists of two species that differ by exactly one proton.

Take our acetic acid, CH3COOH.

When it acts as an acid and loses that one proton, it turns into acetate, CH3COO minus.

So acetic acid and acetate are a pair?

Yes.

Acetic acid is the acid.

Acetate is its conjugate base.

Why do we call it a conjugate base?

Because if you ran the movie backward, if you reverse the reaction,

the acetate would be the one accepting a proton to turn back into acetic acid.

So in the reverse direction, it is acting as a base.

That makes total sense.

And the water?

Water H2O accepted a proton to become hydronium, H3O plus.

So H3O plus is the conjugate acid because it now has the proton to give, and water is its base.

So in every single Brensted -Lowry acid -base reaction, you will always have two pairs?

Always.

Acid 1 becomes base 1.

Base 2 becomes acid 2.

Let's try another example from the text to really cement this, because I think seeing it from the other side helps.

Let's look at ammonia, NH3.

Good choice.

Ammonia is a base.

When you put it in water, it reacts completely differently than the acetic acid did.

Ammonia has a lone pair on its nitrogen atom.

So it is actively looking to score a proton.

Right.

It finds a water molecule.

But in this case, water acts as the donor.

Water gives up one of its protons to the ammonia.

A second ago, you said water was the base with the acetic acid.

Now it's the acid.

That is the beauty of water.

It is amphiprotic.

That's the technical term the text uses.

It means a substance can act as either an acid or a base, depending on who it's dancing with.

It's really versatile.

It is.

If water meets a bully, like a strong acid, it acts like a base and takes a proton.

If it meets a proton -hungry base, like ammonia, it acts like an acid and gives one up.

Back to the ammonia reaction, ammonia, NH3, picks up a proton from water.

It becomes ammonium, NH4+.

NH3 and NH4 plus are a conjugate pair.

NH3 is the base.

NH4 plus is the conjugate acid.

And the water.

Water, H2O, lost a proton.

What is left behind?

OH-, the hydroxide ion.

Water is the acid and hydroxide is its conjugate base.

The rule of thumb seems to be, just count the hydrogens.

The species with the extra H is always the acid side of the pair.

That's it.

It's a simple bookkeeping method that tells you exactly how the protons are moving in the system.

So we've established that water is this chemical shapeshifter.

It's amphiprotic.

It can give or take.

This leads us right into section 16 -2, the self -ionization of water.

Because if water can be both, what actually happens when water is alone in a room with nothing but other water?

It reacts with itself.

It's called auto -ionization.

Imagine a dark room filled with nothing but water molecules bouncing around.

Occasionally two of them will bump into each other with just the right geometry and enough energy.

And one acts as the acid, one acts as the base.

Exactly.

One water molecule literally rips a proton off another.

So you end up with one hydronium ion, H3O plus, and one hydroxide ion, OH -.

Does this happen a lot?

Is my glass of drinking water just a seething soup of ions?

No, thankfully not.

It is a very, very limited reaction.

Water is incredibly stable.

In pure water at room temperature, only about two out of every billion water molecules are ionized at any given moment.

That is a tiny amount.

It is.

But that tiny amount is constant.

And because it is a dynamic equilibrium process, we can describe it mathematically with a number.

An equilibrium constant.

This is KW.

The ion product of water, KW.

It is defined simply as the concentration of hydronium ions multiplied by the concentration of hydroxide ions.

And at 25 degrees Celsius, the standard temperature, what is that number?

1 .0 times 10 to the negative 14.

That is incredibly small.

It is.

Since the auto -wilkwoods -horonization reaction produces exactly one hydronium for every one hydroxide, their concentrations have to be equal in pure water.

The square root of 10 to the negative 14th is 10 to the negative 7th.

So in pure water, the concentration of H3O plus is 1 .0 times 10 to the negative 7th molar, and the OH minus is also 1 .0 times 10 to the negative 7th molar.

But the real magic of KW is that it's a constant, right?

It doesn't matter if the water is pure or if it's full of lemonade mix.

That is the key takeaway for section 16 -2.

In any aqueous solution at 25 degrees, whether it's pure water, gastric juice, or drain cleaner, the product of those two concentrations must always equal 1 .0 times 10 to the negative 14.

So it is essentially a seesaw.

If I add acid and the hydronium concentration goes up, the hydroxide concentration must go down proportionally so that when you multiply them together, you still get 10 to the negative 14th.

You can never have high acid and high base at the same time in water.

They destroy each other until that equilibrium constant is satisfied.

Now, dealing with numbers like 1 .0 times 10 to the negative 7th is pretty annoying.

It is hard to print that on a shampoo label.

It's very cumbersome.

So in 1909, a Danish chemist named Søren Sørensen introduced the pH scale to make it manageable.

The definition is elegant in its simplicity.

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

Let's actually do the math on this so we can see how it works.

The formula is pH equals negative log of H3O+.

So if the concentration is 10 to the negative 7.

The log of 10 to the negative 7th is simply negative 7.

The negative of that is 7.

So pure water has a pH of exactly 7.

That is where our concept of neutral comes from.

And if you have a lot of acid, say a concentration of 10 to the negative 1 molar.

The log of 10 to the negative 1 is negative 1.

So the pH is 1.

This is why the scale feels backward to a lot of people, right?

A high concentration of protons actually gives you a low pH number.

Exactly.

And because it's a logarithmic scale, every single step is a factor of 10.

A pH of 3 isn't just a little bit more acidic than a pH of 4.

It is exactly 10 times more acidic.

A pH of 2 is 100 times more acidic than pH 4.

There is a really specific math tip in the text about significant figures here that I think is worth mentioning because it always trips people up on chemistry exams.

Oh, the log rule, yes.

It's a quirk of how logarithms work with significant figures.

The rule from the text is this.

The number of decimal places in your final pH value corresponds to the number of significant figures in the original concentration.

Can you give us a concrete example of that?

Sure.

Say your hydronium concentration is 1 .0 times 10 to the negative 3 molar.

The 1 .0 has two significant figures.

Okay, I fall.

So when you calculate the pH, which is 3, you must write it as 3 .00.

Two decimal places to match the two sig figs.

If you just wrote 3, you'd technically be wrong according to the rules of precision.

Good catch.

Now, just like we have pH, we also have pOH.

Right.

Just as pH handles the acid part, pOH is the negative log of the hydroxide concentration.

And because of that Kiabo relationship we just talked about, where the product is always 10 to the negative 14th, there's a beautiful simple equation that ties pH and pOH together.

pH plus pOH equals 14 .00.

At 25 degrees Celsius, yes.

This is incredibly useful for problem solving.

If I tell you the pOH of a solution is 4, you don't even need a calculator to find the pH.

You just ask yourself what plus 4 equals 14.

It's 10.

So the pH is 10, which means it is basic.

Exactly.

The text references figure 16 to 5 to ground these numbers in reality.

It shows common substances on the scale.

Gastric juice, which is stomach acid, is around pH 1 to 2.

That is incredibly acidic.

Carbonated drinks around pH 3, right?

Yes.

And then on the other side of the scale, household ammonia is up around 11 or 12.

And concentrated lye, sodium hydroxide, is practically pH 14.

It really helps you realize that this 0 to 14 logarithmic scale covers a massive range of chemical reality.

All right.

Moving on to section 16 to 3, ionization of acids and bases.

We know what acids are, and we know how to measure their concentration with pH.

But here's the thing.

Not all acids are created equal.

We classify them fundamentally as strong or weak.

This is a critical distinction in the chapter.

It's not about how dangerous they are to touch or how concentrated they are in the bottle on the shelf.

It is entirely about ionization.

It's about how willing the molecule is to actually give up that proton to water.

Explain the difference for us.

Think of it like a breakup.

A strong acid, like hydrochloric acid, HCl, undergoes a complete breakup.

If you put 100 molecules of HCl in water, essentially all 100 will split apart completely into H plus and Cl minus.

They don't stay together at all.

No.

The reaction goes entirely to completion.

In the chemical equation, we use a single arrow pointing strictly to the right.

It is one -way street.

And a weak acid.

A weak acid, like our friend acetic acid, is hesitant.

It's a complicated relationship.

If you put 100 molecules of acetic acid in water, maybe only one or two will actually break apart into ions.

The other 98 stay together as the whole intact molecule.

So the vast majority of the acid is actually just floating around intact doing nothing.

Yes.

And because the ions that do form can just rejoin to form the acid again, it reaches a dynamic equilibrium.

We use the double arrow pointing both ways.

And because it's an equilibrium, we can describe it with an equilibrium constant.

Exactly.

Kappa,

the acid ionization constant.

It is calculated just like any equilibrium constant you've seen before.

Products over reactants.

So chi equals the concentration of hydronium times the concentration of the conjugate base, all divided by the concentration of the intact weak acid.

So logically, looking at that fraction, if chi is a big number.

That means the numerator is big.

Lots of products.

Lots of ions.

That means it is a stronger acid relative to others.

And if chi is a tiny number?

It means the denominator is big.

Most of the stuff is still sitting there as reactants.

Very few ions are produced.

It's a weak acid.

But just like with pH, chemists really don't like writing out scientific notation like 1 .8 times 10 to the negative fifth all the time.

So we use the p scale again.

pKa.

Just like pH is the negative log of H, pKa is the negative log of Ka.

And because of that negative sign in the logarithm, the trend flips again.

Right.

And this is where students almost always get confused.

A high Ka means a strong acid.

But a low pKa means a strong acid.

Think of it like golf scores.

You want a low pKa to be a strong player.

That definitely helps keep it straight.

So section 16 to 4 focuses specifically on the strong acids and bases.

The text actually just gives a list.

And honestly, the best strategy for students here is straight memorization because the list is very short.

There are six common strong acids listed in table 16 .3.

Let's run through them so everyone has them.

First, the binary acids.

Hydrochloric HCl.

Hydrobromic HBr.

And hydriotic HI.

Notice that HF, hydrofluoric acid, is missing from that list.

We will get to exactly why that is later.

Then we have the oxoacids.

Nitric acid, HNO3.

Prochloric acid, HClO4.

And sulfuric acid, H2SO4.

But with sulfuric acid, there is a catch, right?

Because it has two protons.

Right.

Only the first proton is strong.

Removing the second one is actually a weak acid process.

But for the others, they are one and done strong acids.

And for strong bases, what's the list there?

They are typically the group 1 and group 2 metal hydroxides.

Lithium hydroxide, sodium hydroxide, which is lye.

Potassium hydroxide, calcium hydroxide.

These are ionic solids that just dissolve and dissociate completely into metal taintions and hydroxide anions.

Calculations of these strong guys are pretty straightforward, aren't they?

Because they ionize completely.

It is the easiest math in the entire chapter.

If you have a 0 .1 molar solution of HCl because it splits 100%, you instantly know you have a 0 .1 molar concentration of HCl.

You just plug 0 .1 directly into your pH calculator.

Done.

No equilibrium tables.

No IC tables.

Just direct plug -in.

Is there ever a time where it's not that simple with a strong acid?

Only if the solution is ridiculously dilute.

The text mentions that if the strong acid concentration is near 10 to the negative 7th molar, so barely any acid in there at all, you can't ignore the fact that the auto -optonization of water itself is also providing some protons.

But for typical laboratory concentrations, you just assume the strong acid completely runs the show.

Okay, that was a good warm -up.

Now we are heading for the main event of the chapter, section 16 -5, calculating pH of weak acid solutions.

This is where the math gets real.

This is where we need the ICE table.

The dreaded ICE.

ICE, initial, change, equilibrium.

I feel like we really should walk through the logic of example 16 -8, step -by -step, because this is the absolute bread and butter of this chapter.

If you don't get this, you're stuck.

Let's do it.

So the problem gives us 0 .100 molar acetic acid.

We want to find the final pH.

And we are given that the Ka for acetic acid is 1 .8 times 10 to the negative 5th.

First step.

Always write the balanced chemical equation.

CH3COH plus H2O is in equilibrium with H3O plus and CH3COO minus.

You have to see the reaction first.

Now we set up the spreadsheet, the ICE table.

Let's do the IRO, initial.

We start with 0 .100 molar of the intact acid.

What about the ions on the product side?

At time zero, right before any reaction happens, we assume the product ions are essentially zero.

Okay.

Next row, change.

Now the reaction runs.

Some of the acid breaks apart.

We don't know exactly how much, so we call it X.

The acid is being consumed as it breaks apart, so its change is negative X.

The ions are being produced.

Looking at the stoichiometry, for every one acid molecule that breaks, you get exactly one hydronium and one acetate.

They both change by positive X.

Final row, equilibrium.

This is just the initial row plus the change row.

For the acid, it's 0 .100 minus X.

For hydronium, it's zero plus X, so just X.

For acetate, it's also X.

Exactly.

We take these equilibrium expressions and we plug them into the Ca formula.

Ca equals products over reactants.

1 .8 times 10 to the negative fifth equals X times X divided by 0 .100 minus X.

Which simplifies on top to X squared over 0 .100 minus X.

Right, and here is where students usually panic during an exam, because if you cross multiply that out, you get an X squared term, an X term, and a constant.

That is a full quadratic equation.

And nobody wants to use the quadratic formula during a timed chemistry exam if they don't have to.

It's messy and prone to calculator errors.

So chemists use a little cheat code here.

We call it the simplifying assumption.

How does that work?

This is so important.

We look at the physical nature of a weak acid.

By definition, it barely ionizes.

The Ca is 10 to the negative fifth.

That is one hundred thousandth.

So the amount that actually ionizes, which is our X, is going to be a tiny, tiny number.

So small that it doesn't even matter.

Think about like a bank account.

If you have $100 ,000 in the bank and you lose five cents, you basically still have $100 ,000.

So mathematically, in the denominator term, 0 .1 over 0 minus X, we assume X is negligible compared to 0 .1 over 0.

We just drop the minus X entirely.

Wow, that makes the math so much cleaner.

The equation just becomes 1 .8 times 10 to the negative fifth equals X squared divided by 0 .1 over 0.

Now, it is incredibly easy algebra.

You multiply both sides by 0 .100 to get 1 .8 times 10 to the negative sixth equals X squared.

Then you just take the square root of both sides.

X equals 1 .3 times 10 to the negative third.

And since our ICE table says X is the concentration of hydronium H3O plus, we just take the negative log of that number to get the pH.

Which is 2 .89.

Done.

But surely we cannot just assume X is small willy -nilly on every problem.

There must be a rule.

When are we actually allowed to do this?

There is a validity test.

The text suggests checking two things to be safe.

First, before you even start the math, look at the ratio of the initial acid concentration to the Ka value.

If the initial concentration divided by the Ka is greater than 100, the assumption usually holds up fine.

Or method two, calculate your X using the assumption and then check the percent ionization.

Which is defined as just X divided by the initial concentration times 100 percent.

Right.

If your calculated X is less than 5 percent of the original concentration, the assumption is considered mathematically valid.

In our acetic acid example, it was 1 .3 percent.

So we are totally good.

And if it is more than 5 percent, what if the acid is a bit stronger or maybe very dilute?

Then, unfortunately, the shortcut fails.

You have to go back, keep the minus X term in the denominator, and you have to use the quadratic formula.

The text shows exactly this in example 16 to 9 with chloroacetic acid, which has a larger Ka.

You just have to grind through the math.

Speaking of percent ionization, there is a totally counterintuitive nugget here that really surprised me when reading.

As you dilute a weak acid solution, meaning you just add more water to it, the percent ionization actually increases.

It does seem weird, right?

You dilute the whole thing, but a higher percentage of it ionizes.

Yeah, why does that happen?

It's pure Le Chatelier's principle from the equilibrium chapter.

Think about the balanced equation.

Acid plus water is in equilibrium with hydronium plus anion.

You have one sloop particle on the left turning into two sloop particles on the right.

When you add water, you dilute everything.

The system wants to counteract that sudden drop in concentration by filling the space with more particles.

So it shifts to the side with more particles, the product side.

Exactly.

It shifts right to produce more ions.

So while the total concentration of everything decreases because of the added volume, the percentage of the original acid molecules that decide to split apart goes up.

That is a great conceptual check.

Okay, let's make things a little more complicated now.

Section 16 -6, polyproduct acids.

Acids with more than one proton to give.

Like phosphoric acid, H3PO4.

It has three protons.

Do they all fall off at once like some kind of cluster bomb?

No, it is strictly a sequential process.

It happens in discrete steps.

Step one, the neutral H3PO4 loses its first proton to become H2PO4 minus.

Step two, that new ion loses a second proton to become HPO4 2 minus.

And step three, that ion loses the final proton to become phosphate, PO4 3 minus.

And each of these steps has its own separate equilibrium constant, K1, K2, and K3.

And it's a massive trend in those values.

K1 is much, much larger than K2, which in turn is much larger than K3.

We are talking factors of 100 ,000 difference between each step.

Why does it get so much harder to remove the subsequent protons?

Think about the electrostatics of it.

In step one, you're pulling a positive proton away from a neutral molecule.

That takes some work, but it's very doable.

But in step two, you are trying to pull a positive proton away from a negatively charged ion.

Ah, opposites attract.

That negative ion wants to keep the positive proton.

Exactly.

It holds on tight.

And then step three, you're pulling a positive proton from a doubly negative ion.

It is holding on for dear life.

So ionization gets progressively harder and harder.

How does this affect the math we have to do?

Do we have to solve three simultaneous ICE tables to find the pH?

Thankfully, no.

Because Ka1 is so incredibly much bigger than the others, almost all of the hydronium ions in the solution come from that very first step.

The acid contribution from the second and third steps is mathematically negligible.

So for calculation purposes, you usually just treat it like a monoprotic acid.

You use Ka1 to find the pH, and you basically ignore the rest.

That definitely simplifies things.

But there is an exception to this ignore the second step rule in the book, isn't there?

Sulfuric acid?

Ah, yes.

H2SO4, the problem child of the chapter.

The first proton in sulfuric acid is strong.

It comes off completely, 100%.

But the resulting ion, HSO4 minus the hydrogen sulfate ion, is a weak acid.

It has a Ka.

So you have a strange mixture of strong and weak behavior in the same molecule.

Right.

You can't just plug numbers in simply.

You have to assume the first step gives you a baseline of H plus equal to the initial acid concentration.

Then you set up an ICE table for the second step.

But your initial H plus for that table isn't zero.

It is that baseline amount from step one.

The text walks through this very carefully in example 1611.

It is the one common case where you really have to pay attention to both steps.

Section 16 -7 briefly talks about a systematic approach.

This seems to be the sort of nuclear option for when our simple assumptions fail.

It is.

If you have a very dilute weak acid solution or a really complex mixture where you just can't simplify things, you need a rigorous mathematical method.

What's in the toolkit for that?

You basically set up a system of equations and equations for n unknowns.

You have your standard equilibrium expressions, your Ka, KBKW.

Then you add two entirely new types of equations.

First, the material balance equation or MBE.

This is just the conservation of mass.

If I put 0 .1 moles of phosphoric acid into the beaker, all the various phosphate species floating around must add up to exactly 0 .1.

The atoms don't just disappear.

And the second new equation.

The charge balance equation, the CBE.

Solutions must be electrically neutral overall.

So the total concentration of positive charge must perfectly equal the total concentration of negative charge.

So you end up with a giant algebra problem.

Yes, you do.

You solve it with a computer program usually.

It is the perfect exact way to solve it.

But it's usually overkill for general chemistry exams unless the professor specifically asks for it.

Let's move to section 16 to 8.

Ions as acids and bases.

This is where we learn that salt water isn't always neutral pH 7.

This introduces the concept of hydrolysis.

When you dissolve the salt in water, the solid breaks apart and the ions float away.

But some of those individual ions can interact with the water molecules to actually change the pH.

We need a way to predict when this happens.

The text derives a crucial equation linking the acid and base strength of a conjugate pair.

Yes, K times KB equals Ktelio.

This strictly applies to conjugate acid -base pairs.

So if I look up the K of an acid in a table, I automatically know the KBB of its conjugate base.

Exactly.

You just take caradoly, which is 10 to the negative 14th, and divide it by the cat.

This leads to a really interesting inverse relationship.

If the acid is strong...

The conjugate base is incredibly weak.

In fact, for a strong acid like HCl, the conjugate base chloride, Cl-, is so weak, it is essentially useless as a base.

We call it a spectator ion.

It does absolutely nothing to the pH in water.

So NaCl, standard table salt, is perfectly neutral because the sodium ion is a spectator and the chloride ion is a spectator.

Correct.

Neither reacts with water.

But take a weak acid like hydrofluoric acid, HF.

Its conjugate base is the fluoride ion, F -.

Since HF is a weak acid, F - is actually a decent active base.

It's not a strong base, but it's strong enough to do something.

If you dissolve sodium fluoride, NaF, in water, the fluoride ions will act as a base.

They will literally grab protons from the water molecules, generating OH-,.

So a solution of sodium fluoride is actually basic.

Yes.

Anions derived from weak acids make solutions basic.

Conversely, caracations derived from weak bases, like the ammonium ion and Nase 4 +, make solutions acidic.

There is one more category of acidications mentioned in this section that I thought was strange.

Metals.

Small, highly charged metal -cations, like aluminum, Al3 +, or iron, Fe3+.

They don't have any protons themselves to donate, so how are they acidic, right?

Well, when they dissolve in water, they get heavily hydrated.

Six water molecules tightly surround the aluminum ion in a complex.

And the aluminum is so highly positive.

It forcefully pulls electrons away from those attached water molecules.

It weakens the OH bonds in the water so much that a proton just falls off into the solution.

So a simple solution of aluminum nitrate is actually quite acidic, similar in pH to vinegar.

Section 16 -9 deals with qualitative aspects.

Basically, how to predict the direction of an acid -base reaction without doing all the math.

I call this the survival of the weakest rule.

I like that.

Explain how that works.

The principle is that an acid -base equilibrium always favors the formation of the weaker acid and the weaker base.

Why is that the case?

Think about chemical stability.

A strong acid is highly reactive.

It's stressed out.

It desperately wants to get rid of its proton.

A weak acid is stable.

It's perfectly happy holding on to its proton.

Nature always prefers stability.

So the reaction naturally moves away from the strong reactive side and settles toward the weak stable side.

So if I mix a strong acid and a strong base together.

They're both highly unstable and reactive.

They drive the reaction violently toward the products, which are just water and a spectator salt, which are very stable.

The reaction essentially goes 100 % to completion.

Now, section 1610.

This is the part I find really, really cool.

Molecular structure and acid strength.

Up until now, we've just been talking about K values as random numbers in a table.

But why is one acid stronger than another?

Why is HF weak, but HCl strong?

This is where the chemistry becomes physics.

It all comes down to bond strength and bond of polarity.

Let's look at binary acids first.

The simple ones like HF, HCl, HBr, HI.

The text mentions two competing factors here, bond strength and polarity.

Right.

Let's look at the group trend first, going straight down the halogen column on the periodic table from fluorine down to iodine.

As you go down, the atoms get physically bigger.

Fluorine is tiny, iodine is huge.

So the bond between hydrogen and iodine is very long and mismatched.

It is a very weak bond.

And a weaker bond is easier to break.

Correct.

So it is much easier to donate the proton.

That means acidity actually increases as you go down the group.

This is highly counterintuitive for many students.

They think, well, fluorine is the most electronegative.

It's scary, so HF must be the strongest acid.

No.

The HF bond is short and very strong.

It holds the proton too tight.

So HF is a weak acid.

HI is a strong acid.

OK.

What about going across a period on the table, like from methane, CH4 over to HF?

There, the polarity factor wins.

As you go across, electronegativity increases.

Fluorine is the most electronegative element.

It creates a very polar bond, pulling all the electrons away from the hydrogen, leaving the proton partially exposed and positive.

That makes it much easier for the H to fall off as an ion.

So acidity increases as you go across the period.

OK.

What about oxoacids?

Acids containing oxygen, HOZ, like comparing a hypochlorous acid to hypouretate acid.

Here, everything depends on that central Z atom.

If Z is very electronegative, like chlorine, it pulls electrons through the entire molecule.

It pulls from the oxygen, which in turn pulls from the hydrogen.

We call this the inductive effect.

So it weakens the OH bond from a distance.

Exactly.

It makes the proton fall off easier.

So HOCl is a stronger acid than HOI, because chlorine is more electronegative than iodine and pulls harder.

And what about the number of oxygens attached, like prochloric acid, HClO4 versus hypochlorous HClO?

The rule is more terminal oxygens equals a stronger acid.

Prochloric acid has three extra oxygen atoms attached to the central chlorine.

Oxygen is highly electronegative, so they all pull electron density away from the central atom, which pulls it away from the OH bond.

But there is also a stability argument for the conjugate base here, right?

That is the absolute golden rule, resonance.

When prochloric acid, HClO4, loses its proton, the resulting negative charge isn't just stuck on one single oxygen atom.

It is spread out, delocalized, over all four oxygen atoms equally.

Spreading out that charge makes the ion incredibly stable, and because the final product is so stable, the reaction really wants to happen.

So prochloric acid is an absolute beast of an acid, because its conjugate base is just so relaxed and stable.

Exactly.

The text also mentions organic acids in this section, carboxylic acids, with the no -key UOH group.

It is the exact same logic.

When they lose a proton, the resulting carboxylate anion, COO-, is stabilized by resonance.

The negative charge is shared nicely between the two oxygen atoms.

That stability makes it worth it for the acid to lose the proton.

What if you add electronegative atoms nearby on the carbon chain, like comparing regular acetic acid to chloroacetic acid?

The chlorine atom exerts that inductive effect again.

It pulls electrons through the carbon chain, stabilizing the anion even more.

So chloroacetic acid is significantly stronger than regular acetic acid.

You can see this visually in Figure 1611.

Finally, we reach Section 1611, and here we completely break the paradigm we've been using.

Lewis acids and bases.

The Brinsted -Lowry definition is fantastic for protons, but G .N.

Lewis, the same guy who gave us Lewis dot structures, looked at this and said, what about reactions that act exactly like acid -base chemistry, but don't involve any hydrogen at all?

So how did he redefine it to fix that?

He looked at the electrons instead of the protons.

A Lewis acid is defined as an electron -pair acceptor.

A Lewis base is an electron -pair donor.

Does this still fit with our old definition?

It does.

Look at our old friend H+.

It has no electrons.

It desperately wants a pair.

So it is an electron -pair acceptor.

It is a Lewis acid.

Look at ammonia, NH3.

It has a lone pair.

It donates it.

It is a Lewis base.

But this definition is so much broader.

It covers things with zero hydrogen.

Right.

Take boron trifluoride, BF3.

Boron only has six valence electrons in that molecule.

Its octet is incomplete.

It is hungry for two more electrons.

It can react directly with ammonia.

Ammonia donates its lone pair right to the boron.

They form a new bond, a cordon -covalent bond.

So BF3 is the acid in that reaction.

Yes.

It accepted the electron pair.

The product they formed together was called an adduct.

This entire concept of electron donors and acceptors is absolutely massive.

Almost all organic chemistry reactions can be viewed fundamentally as a Lewis base, a nucleophile, attacking a Lewis acid, an electrophile.

Which brings us to the end of our journey through Chapter 16.

We have gone from the illiteral taste of lemons all the way to the electron -orbital interactions of boron.

It really is a journey from the macroscopic right down to the microscopic.

We start with simple proton swapping.

Then we did the heavy accounting math of equilibrium, looked at how simple salts can mess with pH, and finally explained exactly why these things happen using molecular structure and electrons.

And as the text says in the outro, acid -base concepts are arguably among the most important of all concepts in chemistry.

It is entirely true.

If you can learn to spot the acid, the electron acceptor, and the base, the electron donor, you can predict the outcome of almost any chemical reaction.

You basically have the keys to the kingdom.

A huge thank you to everyone for sticking with us on this deep dive into Chapter 16.

Hopefully the next time you check the pH balance of your shampoo, you won't just see a marketing number.

You will actually picture those hydronium ions waging war against the hydroxide ions.

And remember, whether it is exchanging protons or sharing electron pairs, chemistry is just the universe trying to find a comfortable, stable arrangement.

From the Last Minute Lecture Team, thanks for listening.

Goodbye.