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PROFESSOR: So a couple of announcements.
Weekly quiz Tuesday.
And also on Tuesday, I want to draw your attention to the
event in the slide.
We have a poster here for a lecture that will occur right
in this room, at 4:00 on Tuesday.
And it's entitled the Wolf Lecture.
The Wolf Lecture was established here about 30
years ago in honor of Professor John Wolf.
John Wolf was an antecedent of mine.
He was the person that invented 3091.
And in 1961, he dared, as a metallurgist, to launch a
variant of freshman chemistry.
And you're the beneficiaries of John Wolf's initiative.
And he also was a spectacular teacher, quite a showman.
And in his honor, they instituted in the Department
of Material Science and Engineering, the Wolf Lecture,
which, as the poster says, is the entire community is
invited to attend, but it's geared towards freshman.
So this is a lecture that the rules are.
Talk about whatever you want in terms of material science
and engineering, but make it entertaining, and make it
engaging and accessible to freshmen.
You should go to lectures.
If you can't go to this one, then by all
means, go to other seminars.
If all you do when you come to MIT is go to class, you're
missing out on some of the richness here.
We have all sorts of people coming through here every day.
And you can go to these seminars, you can learn a lot.
You might say, well, I'm just a freshman, I'm not going to
understand everything.
I go to these things.
I don't understand everything.
But I learn something, because the first few minutes of the
lecture, if the speaker is any good, he or she is going to
set up the topic for you.
If you follow the first five minutes, you'll have something
that, you know, gets you oriented to the topic.
And then the other thing, the other reason to go-- you see
this on the poster--
room 10-250, reception immediately following.
That means there will be refreshments!
Food, for that, you know,
pick-me-up in the late afternoon.
So you go to the seminar, you know, sometimes they have
refreshments before the speaker.
I'm not going to tell you have to go into the seminar, having
eaten the food and partaken of the refreshments.
You didn't hear me say that.
But anyways, there's always food around here.
Go to this one.
I think you'll enjoy it.
And the speaker is Professor Michael Rubner.
He's a faculty member in course 3.
An excellent teacher.
As you see, MacVicar Faculty Fellow, which means he's been
acknowledged as a good teacher.
I think you'll learn a lot.
And he's talking about nature-inspired, so
biomimetics, how we see things in nature and then mimic that
design in advanced material science.
OK.
Enough said.
Let's get on with the lesson.
Last day we started looking at solutions, and we recognized
that bonding is the key to understanding, and it's
encapsulated in that catch-all phrase, like dissolves like.
And towards the end, we talked about Ksp as a metric that
helps us understand common ion effect.
Common ion effect, again, if I tell you you have, say, 100
units of sodium chloride that can go into solution, you
would say, that's the solubility limit.
But then if I tell you, I've got a solution that already
contains 25 units of sodium chloride, that's a no-brainer.
You've only got another 75 units to go.
Now I say I've got potassium chloride in there.
So now you stop and think, wait a minute.
It's not sodium chloride, but it is a chloride.
How do I think about this?
Common ion effect through Ksp helps you reason through it.
It's when you have more than one solute, and one of the
constituents of the solute is already present.
So today I want to talk about a subset of solutions, and in
particular, I want to talk about acids and bases.
And this is important, not only in materials processing,
but we're going to have to understand this if we're going
to go forward later and talk about biochemistry.
So, you know, they're around us everywhere.
You know, you probably got up this morning, you washed your
hair with a pH-balanced shampoo, maybe had orange
juice or grapefruit juice with citric and ascorbic acid.
Your radio is powered by zinc alkaline batteries.
I started my car, it's got the lead acid battery in it, and
electricity for all my appliances came from
coal-fired power plants, spewing out SO2,
turning out acid rain.
So we're off to a good start!
It's Friday.
So now we want to go back and understand
this from the beginning.
So we're going to start with a history lesson.
And the history lesson starts in ancient times.
Acids were known all the way back in early times for
processing of food and materials.
So who among us hasn't eaten something
that has been pickled?
Pickle means to have been processed in acid solution.
In fact, the word acid comes from the word acidus in Latin,
and acidus means either sour or tart.
But the modern chemistry of acids and bases starts with
Lavoisier in France in 1779.
And I'm going to be naming many scientists from different
countries today, so I'm going to use the
international symbol.
You know, these ovals that you put on the back of the car, so
now if you see a car with this on it, you know
it's a French registry.
So he's probably got a Peugeot, and he's driving
around with his Peugeot.
And what Lavoisier said, in trying to understand acids and
bases, is--
it's a really interesting story.
He said that oxygen is present in all acids.
And why?
Because he spent most of his career studying combustion, so
he was very concerned about oxygen.
And the interesting thing is that oxygen, the name of the
gas, oxygen, actually comes from the greek oxy, which
means sharp, and, you know, the particle gen, as in to
generate or to be born.
OK?
So it's interesting that this gas is named incorrectly.
It's named for an attribute that is ascribed to acid, and
it's wrong.
And it's in other languages, too.
The word oxygen translates into other languages as
meaning something sharp, and it has nothing to do with it.
But while we're on the topic of Lavoisier, Lavoisier did
study oxygen.
And there was an intense rivalry between Lavoisier in
France, Joseph Priestley in Britain,
and Scheele in Sweden.
And all three of them were working simultaneously to
study combustion and understand oxygen.
And here I've got a very nice image.
This is a portrait of Lavoisier with his wife.
And it's hard to see.
If the lights were a little bit dimmer, you'd see--
notice here, you see all of this chemical apparatus.
There is bell jars, and various glass apparatus.
His wife was his partner in the laboratory.
She was unique among French women in
that she read English.
And in his rivalry with Priestley, he relied on his
wife in order to read Priestley's writings.
So they worked together.
She even helped him in the laboratory.
They married at the time he was 27 and she was 13, and
they were--
what are you so shocked about?
It's France, and it's 1779!
Get over it!
Anyways, she helped him a lot, and we'll say a little bit
more about that later.
OK.
So now let's go on to the next part of the history lesson.
So so far, we've got one explanation, and it's wrong.
So let's go to the next one, and we'll go to Britain.
Sir Humphry Davy in London in 1810.
So we'll put GB here, and he's driving around in his little
Jaguar, I suspect.
And he said something that was correct.
He said that it's hydrogen present in all acids.
And that was pretty much it.
He was a great scientist, did some very good work, but
really didn't do anything quantitative here.
So for the quantitative stuff, we have to wait for almost the
end of the century.
And we go up to Sweden again, to Arrhenius.
And Arrhenius, in 1887--
so I'll give him a Volvo.
And he said that the acid is a substance that disassociates
to produce protons.
So the acid is defined as a solution that dissociates to
give protons.
I'm going to say H plus, meaning the hydrogen ion, or P
plus, and we're going to keep using the term proton.
I'm not going to say hydrogen ion, I'm going to say proton.
So proton in solution.
So this is what he defines.
And he further defines the base.
He defines the base as the complement to the acid.
And he says that the base is something that dissociates to
give us hydroxyl.
The base dissociates to give OH minus the hydroxyl.
And so now we've got this whole concept of electrolytic
disassociation, which is what wins him the Nobel prize in
1903 for this thing.
And then we saw last day how the addition of ions to water
gives charge carriers.
And so the presence of protons and hydroxyl is, in fact, the
way we have any charge carried through water.
So up until now, when I said electrical conductivity, we
were pretty much referring to electronic conductivity.
But there's a second kind of conductivity, and we can have
ionic conductivity.
Ionic conductivity, as the name implies, is not by
electrons, but by ions.
And this is typically ions in solution.
Ionic conductivity in solution, and a solution that
is an ionic conductor is called an electrolyte.
So we have electrolyte in our bodies,
saline, about 5% chloride.
We have electrolyte in batteries that are ionically
conductive.
And the term simply means, we have conduction by ions.
And so it was the theory of electrolytic dissociation that
won the Nobel prize.
So how does this work?
So we can start with something like HCl gas.
I'm going to dissolve it in water, and this will give me
proton, H plus.
And I'm going to write aq, meaning that it's dissolved in
water, and the chloride ion, also dissolved in water.
And so this gives me an acid.
This is a proton donor.
And now let's look at
something like sodium hydroxide.
Sodium hydroxide at room temperature is an ionic
compound, so it is a solid, but it is soluble in a polar,
hydrogen-bonded liquid, so H20, to give hydroxyl aqueous
plus sodium ion aqueous.
So again, we see the dissociation.
And then we can have, from here, a
neutralization reaction.
And the neutralization reaction is simply
reconstitution of the solvent.
So neutralization, another way to think about it, because
we're not going to be confined to water by the end of the
lecture, neutralization is simply a reaction that results
in reconstitution of the solvents.
So the solvent, in this case, is water.
So how would we reconstitute water?
We combine proton with hydroxyl.
So let's do that and see the result of that reaction.
So if I take, and run it in the vertical direction, proton
plus hydroxyl, we'll give water again--
let's give H20 liquid--
and now you can see Na plus Cl will give me
NaCl, aqueous dissolved.
So this is what you probably learned in your high school
chemistry, that acid plus base ca give you salt plus water.
So you see all of this resulting from just the simple
Arrhenius definition.
So this is good.
We've got off to a decent start here with Arrhenius.
But then the theory has its limitations, as they all do.
And how do we discover the limitations?
With some new data.
So let's look at some data.
So it had been known for a long time that ammonia, when
it's dissolved in water, can act as a solution capable of
neutralizing an acid.
So if ammonia dissolved in water neutralizes an acid,
then this must be a base.
But look, there's no hydroxyl here.
So the Arrhenius definition of base is inadequate
to account for this.
So let's just get that down.
Ammonia, which you know is a gas at room temperature, will
neutralize acids.
But no hydroxyl present.
So something's going on here that we can't account for by
the simple Arrhenius definition.
So to get us out of this conundrum, we had to wait
until 1923.
And two scientists simultaneously enunciated the
same ideas.
So I'm going to put them both down.
1923.
Bronsted, who was in Denmark, and Lowry, who was
working in the UK.
Also 1923.
And I'll give him a Jaguar as well.
And so the two of them proposed a broader definition
to account for what's going on here.
And what did they say?
They said that acid--
they'll keep the same definition as Arrhenius.
So an acid is going to be a proton donor.
So that's good.
Let's even put here, same as Arrhenius.
But now here's the difference.
The base is no longer confined to hydroxyl chemistry.
They call it a proton acceptor.
And that's different.
So hydroxyl-free.
Now, let's be careful here.
If I propose this new definition, I
can't throw out hydroxyl.
So you've got to watch to make sure that by broadening the
definition, we don't exclude hydroxyl.
So the theory has to encompass what we already know, and then
continue to encompass things that tend to contradict what
we already know.
So let's take a look at what this gives us.
So I'm going to write a broader equation here.
I'm going to write this as the acid.
So what do I have?
I have every acid has to have a proton in
it, plus some residual.
So I can rewrite the acid in this form.
I'm going to react it with a base.
So according to this definition, this has to be a
proton donor.
And this, you're going to watch me on this one.
This is going to be a proton acceptor.
And we'll put some identities here.
Eventually we're going to write this with ammonia in it,
but let's just do the broad definition.
So what happens is, if this is a proton acceptor, on this
side of the equation, it takes the proton from here.
This is the proton donor.
It gives it away, and we end up with a BH plus.
So this proton acceptor has accepted the proton, leaving
behind the deprotonated A minus.
That's a nice little reaction.
But now, if you'd nodded off for the last 90 seconds, and
then opened your eyes and said, well, he said acid is a
proton donor, this thing can give up a proton.
So OK, I'm going to call this thing a proton donor.
And this is most certainly a proton acceptor.
Look, it has accepted the proton.
So this is going to be a proton acceptor.
So now I've got in this equation two proton donors and
two proton acceptors.
And they're linked.
This proton donor, BH plus, is linked to B.
So I'm going to put a little yolk over this one.
And this HA, the original acid, is linked to
this base, A minus.
So I'm going to put a yoke over that.
And the Latin word for yoke, to yoke, is jugere, from which
we get the word conjugate.
So I'm going to call these things
conjugate acid-base pairs.
And every equation has two.
Because you've got a proton donor proton acceptor.
So now this is, follow the bouncing ball.
If you want to understand acid-base chemistry, just
follow the proton.
Here's proton here, goes over here.
Here's something deprotonated, now it's protonated.
That's the rhythm here.
So the whole thing is, acid-base reactions can be
defined as a proton transfer reactions.
Remember we talked about ionicity and electron transfer
in order to achieve octet stability?
So we saw the whole life and times of
electron transfer reactions.
Now life and times of proton transfer reactions.
OK.
And last thing I want to do, is to just make sure that we
see the definitions.
You know, all queens are female, but not all females
are queens.
So this one here, what's this?
This is definitely, it's a proton donor.
So it's an acid.
So it's definitely an Arrhenius acid, because it's
got a proton, and it's also a Bronsted-Lowry acid.
Now, what about this one?
This is a base.
It may not have hydroxyl in it.
May or may not.
So this is definitely a Bronsted-Lowry base, but I
can't say for sure that it's an Arrhenius base, because
Arrhenius has to be O minus.
This one here, what about this?
Unless it's OH minus, this one here is, what?
It's Bronsted-Lowry base.
And this one here, this can be both Bronsted-Lowry acid and
it can be an Arrhenius acid.
So this is how we can differentiate them.
OK, good.
Now let's go to the ammonia, see if we've got ourselves out
of the conundrum so we can write now NH3.
And I'm going to say that this is ammonia
that's already dissolved.
So ammonia is already in aqueous solution, and it
reacts with water, H20 liquid, to do what?
This is supposed to be a--
if this is a base, remember, this is the thing that we're
trying to demonstrate.
If it's a base, according to this, it's a proton acceptor.
So I'm going to stick a proton on this, and I'm going to get
NH4 plus the ammonium ion.
And how did I get the ammonium ion?
I took a proton from water, leaving behind OH.
So can you see that how this thing works?
By gobbling up protons, by becoming a proton vacuum
cleaner, it takes protons out of water, leaving
an excess of hydroxyl.
And in effect, now we've got something that's tantamount to
an Arrhenius base.
But it all started off with this.
So now we can link these two.
I'm going to yoke ammonium and ammonia as conjugate
acid-base, and there is water and hydroxyl
as acid-base pairs.
So again, let's get the colors going.
Here we've got Bronsted-Lowry--
the bases are always going to be in blue.
Blue and base both begin with the letter B.
So this is definitely Brosted-Lowry base, but it's
not an Arrhenius base.
This is a Bronsted-Lowry base, and it's also
an Arrhenius base.
This one here, we can call this one--
in this case, it's a proton donor, so it's both an
Arrhenius acid and a Bronsted-Lowry acid.
But look at this.
You'd say, OK, I know what he's doing.
Arrheius, Bronsted-Lowry, Arrhenius, Bronsted-Lowry.
This is Bronsted-Lowry only, this must be
Bronsted-Lowry only.
But look!
This is a proton donor.
It dissociates to give protons.
So this is both Bronsted-Lowry and Arrhenius.
Most people will miss that.
We'll get to see, maybe, on a subsequent celebration, how
many of you miss that.
OK.
So.
So now, let's go into the chemistry here.
Because I said up here at the beginning of the lecture--
what is it-- bonding is the key to understanding.
So what is it about bonding here that defines the
Bronsted-Lowry base?
So I'm going to write this reaction in this way, now.
I'm going to put-- here's the proton, which
is coming from water.
And I'm going to write that with NH3.
I'm going to use the Lewis structure.
So, you know, Lewis structure looks like this, Nitrate has
got five valence electrons, three of them are bound, and
now I've got this long pair.
So what do we know about the electronic
structure of proton?
What does it look like?
What's the Lewis structure of proton?
How many electrons on proton?
None.
This is nothing.
This is very needy.
It's like the neediest friend you have. You know, the one
that's always taking stuff from you, taking your
emotional energy, giving back nothing?
That's proton.
That's the human equivalent of proton.
So if proton wants to make a bond to form NH4, proton is
going to come over here and exploit both electrons.
So if this is going to be a proton acceptor, the only way
it can be a proton acceptor is to have two
unused electrons available.
This is going to be a dative bond, agreed?
Dative, because both electrons from the bond
come from the nitrogen.
The proton doesn't contribute anything.
So that's the hallmark.
The proton acceptor axiomatically must have an
available non-bonding pair.
So wherever you see non-bonding pairs, wherever
you see a compound that has this capability, it could
serve as a Bronsted-Lowry base.
We know that water, because it has hydroxyl in it, has to
satisfy this.
So let's take a look.
H20 we can do similarly.
Oxygen.
It's got 6 electrons, 2 are in the, bonds and we have 2
non-bonding pairs.
So then when we attach proton here, we make
something called hydronium.
We make hydronium ion.
So it's a little fancier than just saying, I've got this
naked proton with no electrons swimming around in water.
In fact, it's more coordinated like this.
H30 plus, and while we're in the neighborhood, we
might as well show.
This is sp3 hybridized.
We've got one, two, three, four orbitals.
Hydrogen's on three.
It's got a net charge of plus 1, and lone pair
on the fourth side.
So there's the structure of hydronium.
And we can now write the equation.
We can say H20 liquid plus H20 liquid gives me H3O
plus, plus OH minus.
So this is now the self-ionization or
self-dissociation reaction.
So now I can pair these as well.
I can yoke these.
I have acid-base pair.
So this is acid here, and the conjugate base is water, H20.
Hydroxyl, we know, has to be a base, and its
conjugate acid is water.
So you see, in the same equation, same place, same
time, water is acting as both acid and base relative to
these two species.
And so we call such a such a compound that acts as both
acid and base, we call it amphipathic.
Has two moods.
You know, amphi, like in amphitheater.
Have you ever seen a Roman theater?
It looks like this from the top down.
It's got all this, you know, like this, and all the people
are here, and there's the stage up here.
But if I make a theater in the round, that is amphitheater.
Or as it's commonly mispronounced, ampitheater.
3091ers do not say ampitheater, they say
amphitheater.
So this reaction here is self-dissociation.
Right?
This reaction is self-dissociation of water.
Or because we're forming ions, it's also called
self-ionization.
Self-dissociation or self-ionization of water.
Now, turns out that the chemical's power of H3O plus
and OH minus is very high.
So this reaction doesn't go very far to the right.
The amount of the dissociation is tiny, and to be specific,
at 25 degrees C, the amount of, if you take deionized,
deaerated water, absolutely pure and pristine.
The amount of native H30 plus and OH minus is on the order
of one part in 10 million.
So H3O plus concentration would be 10
to the minus 7 molar.
And obviously, from the stoichiometry of the equation,
you get the identical amount of OH minus.
Well, that's in the self-dissociation.
And so you have very, very few charge carriers, which is why
high-purity water is a very poor conductor of electricity,
even as an electrolyte.
And that's, I've told you in the past, that could be one
indicator of certain toxins.
If you measure the conductivity of water and you
discover it's abnormally high, and it's putatively supposed
to be pure water, that's an indication that
it's not pure water.
There's some other charge carriers present.
So we can write a Ksp we analogy.
And it's called the water ionization constant, which is
the product H3O plus and OH minus.
And you can do the math.
That's 10 to the minus 14 is the product.
And you can use the common ion effect here.
If I introduce some other acid, look at how this
equation helps you the same way that Ksp did.
If I have acid, in other words, protons donated from
some other source, that means this number is going to be
higher than 10 to the minus 7.
If this number is higher than 10 to the minus 7, and the
product must be 10 to the minus 14, this must be lower.
So when I have, H3O plus goes up, then OH minus must go
down, and under these circumstances, we have
something that is called proton-rich.
And we call this acid, simply because the proton
concentration exceeds the hydroxyl.
And likewise for the other.
And so we can make a plot of this, and that's shown here.
And all we've got is OH versus H3O plus.
And you can see that it's a right hyperbola.
But it goes back to that comment that I made in the
past about how you graph data.
I don't know what to do with that, because it's curved.
I can't tell whether the data are good or bad.
So instead of looking at something like this, y versus
x, I want to transform so that can get a straight line.
Because then I can make a judgment about goodness
of fit and so on.
So I need some f of x versus g of y to, quote unquote,
straighten this out.
And we have that, thanks to another Dane,
by the name of Sorensen.
Sorensen was the one that gave us another way
to think about it.
Sorensen, who in 1901, and he's got the DK on his--
he's probably got a Volvo, too.
Or maybe he's driving a Saab.
So he was a biochemist at the Carlsberg brewery.
Yes, they had biochemists, because they wanted understand
the chemistry of beer production.
What a concept.
Having people that understand the
technology of the business.
So he was working at the Carlsberg brewery, and he
decided to take this acid-base business and turn it into
something that's much more readily recognizable.
And so he defined a concept called the chemical potential,
which, if you like, is the reactivity, the chemical
potential of hydronium.
And he gave it the symbol.
Lowercase p for potential, and uppercase H as, obviously, for
the hydrogen ion.
And he said, I'm going to make this a logarithm.
So it's logarithm, and in those days, everybody was
using slide rules, so it's log base 10 of the concentration
of hydronium.
And being a good engineer, he recognized since this starts
off at 10 to the minus 7, the log of a number less than one
is going to be negative.
And who wants to deal with negative numbers?
So you put a minus sign here.
That way, pH normally is a positive number.
And now you can see the results of Sorensen's
straightening things out for us.
So now you go over a wider range, and you can see, you
have nice straight line relationship, and then you can
get data on there, and so on.
And this is taken from your book, and it shows the pH
values of some common substances.
If you start here at neutral pH 7, you have milk, you have
human blood, normal things that you would expect, fluids
and so on, are hovering around 7.
If you take some coffee, you're going to go into the
acidic region.
You see tomatoes down here at about 4.
Wine, you'll see in the reviews of wines, they'll say
it has balanced acidity and so on.
Yeah, it's down around pH 3.
And carbonated soft drinks, some of them get down to 2.
Vinegar is wine that has spoiled.
Vin aigre, which means eager.
The Middle French eager meant
impetuous, or tart, or something.
So this is spoiled wine.
And the pH changes.
So by measuring the pH of wine, you can tell if it's
changing or not.
There's lemon juice.
Gastric juices in the stomach can get down to 1, pH of 1.
But you want that happening only at one point in the
digestive cycle.
If you run around for long periods of time at pH 1, you
will ulcerate, in other words, you'll puncture the walls.
And if that's about to happen, you need to neutralize.
So how do you neutralize?
You go up with something that has a high pH.
So you can start with something like Alka-Seltzer,
sodium bicarbonate at about 8.4, or milk
of magnesia, 10.5.
Why is it milk of magnesia?
From last day, because it's not a solution, it is a
suspension.
The magnesia is in suspension.
The magnesium hydroxide.
And that's one of those mandatory shake well before
using, because all the magnesia is on the bottom, and
you've got this almost clear, colorless liquid on the top.
So you're just drinking water.
Not going to help you if you've got stomach pain, but
gastric juice is down here.
If you're really desperate, don't reach for ammonia.
You have to be patient.
They have to work with milk of magnesia.
OK.
So that's the range.
So what else do we have here?
Let's take a look.
So far we've been assuming that when we add acid to the
system, we get 1 to 1 dissociation.
So the next message that I want to give you here, is that
not all acids are of equal strength.
See, I could look up here and say, well, maybe the reason
that the lemon juice is down it at 2 and the gastric juices
are at 1 is simply a concentration effect.
How much acid was introduced.
No, there's another explanation for it.
And that is that certain acids don't fully disassociate.
So for example, if we look at 1 molar HCl, hydrochloric
acid, you look at 1 molar hydrochloric acid, that goes
100% into solution and gives us 1 molar H30 plus.
So 1 to 1 correspondence between how much hydrochloric
acid we introduce, and how much
hydroxyl that we generate.
So this is total dissociation of the HCl.
And as a result, we call this a strong acid.
In contrast, let me show you a weak acid.
So a weak acid is going to dissolve, but it doesn't
dissociate.
So there's really two steps here.
First you've got to get the stuff in solution, and then
once you get it in solution, it has to break apart.
It's possible to go into solution and not break apart.
In which case, you don't have the protons, but you've got
the solution.
That's not an acid.
So a weak acid would be, let's, in contrast,
look at a weak acid.
A weak acid would be something like acetic
acid, which is CH3COOH.
And this is for historical reasons.
Normally, the proton that's going to dissociate is at the
front of the formula, but this was written long ago, before
people understood this theory.
And that's the way you'll see acetic acid written.
So really, this is the proton, and the CH3COO is really the
acetate anion, Ac, acetate anion.
So we're going to put that into solution, like this.
I'm going to put it into water, H20 liquid.
So this is the salt out of solution, and all we're going
to do is make it into the aqueous solution of same COOH,
and I'm just going to write aq.
Now, if this were strong acid with impunity, I would write
whatever the concentration of acetic acid is, I break it
apart, and I get the full amount of the proton.
What happens is that this, now plus H20, gives me H3O plus
plus the remaining acetate, which is CH3COO minus.
And now, here's where the weakness is manifested.
The degree to which the acetic acid grabs protons from here,
and then ends up being protonated,
is very, very small.
It turns out that in the case of 1 molar--
I'm going to write acetic acid this way, now.
HAc.
So this thing here is the CH3COO minus, all right?
It turns out that in 1 molar acetic acid, we
get only 0.4% reaction.
0.4% reacts and dissociates.
So to give protons, it hangs onto most of its protons.
So we can then represent this in the form of an acid
dissociation constant.
And that it looks like this.
K sub a for acid.
So on the right side, what we're going to do is make a
mass balance here.
We're going to take the product of the proton, the
acetate divided by the undisassociated acetic acid.
So we'll put H3O plus concentration times the
acetate concentration divided by the
undissociated HAc aqueous.
So I'm trying to represent what the ratio is here.
And instead of being one to one, it's 10 to the minus 5.
Very, very weak.
You get a little bit, but not much.
Now, you could say, well, 10 to the minus 5, this is, just
for reference, it's 100 times 10 to the minus 7,
which is the Kw.
So the point here is, if you put in acetic acid, you'll end
up with 100 times the proton population that you would have
had by just self-dissociation.
So it isn't acid.
It is donating protons.
But it's not doing it a lot.
So this is called a weak acid, because it's
a poor proton donor.
And so we can then look at a comparison to,
say, over here, HCl.
Let's go back here, and we can write the acid dissociation
constant for HCl, for the strong acid.
And in that case, Ka over here, which is going to be the
H3O plus Cl minus, H3O plus.
In the case of HCl, we're going to get proton and
chloride over undissociated HCl dissolved in water.
And that's 10 to the plus 6, which for all intents and
purposes is infinity.
You put in 1 molar HCl, you get 1 molar H plus.
And look at the ratio.
10 to the sixth, 10 to the minus 5.
There's a 10 to the eleventh relationship between the weak
acid and the strong acid.
And here's the cartoon that shows this strong acid, put in
this amount, 100% dissociation, weak acid gives
you small amounts of proton, and a very weak acid is
imperceptible amount of disassociation.
And here's a table that quantifies it.
So the strong acids up here with the Ka.
This value of Ka has to be greater than 1, because it's
saying that we're getting a lot of dissociation, the
reactions moving to the right.
And you have these numbers here.
10 to the ninth, 10 to the eighth, 10 to the sixth.
Those are differences without distinction.
if I say the ratio is a million to one, or a billion
to one, for all intents and purposes, you have
dissociation.
And then here are the week acids.
Look.
10 to the minus 3.
There's phosphoric, hydrofluoric, and so on, all
the way down.
There's carbonic acid, 10 to the minus 7.
So there's the list.
And just another cute demonstration of the
relationship between bonding and acidity, the acid strength
increases from left to right, and the bond strength
increases from right to left.
You know, HF is very strong.
It's a hydrogen-bonded.
Very tight HF bond.
But what determines if it's going to be a
strong acid or not?
It's how willing HF is to give up its H.
But its H is tightly bound.
So oddly enough, in HF, you don't give up much of the H.
HF is over here.
It's less than one.
It's a weak acid.
HCl is a pretty strong acid.
If you want to go heavy-duty, HBr and HI are even stronger
acids, because they have a weaker hold on the hydrogen
within them, so they are much, much more
generous proton donors.
So you can see that from there.
OK.
So what's the takeaway message here?
The takeaway message is that equal acid concentration does
not mean equal acid strength.
You have to be mediated by the Ka.
Equal acid concentration, which means, how
much did you dissolve?
Does not equal acid strength.
And what's the reason?
Because equal acid concentration, this is a
function of solubility, whereas this is a function of
dissociation.
So many dissolve, few dissociate.
How about that?
There's a nice tagline you can use.
Use that at a party this weekend.
Many dissolve, but few dissociate.
All right.
Now I want to go very, very high.
Big concept, all right?
Last definition of acid-base comes from the United States.
G.
N.
Lewis.
Same G.
N.
Lewis who gave us the Lewis cross and dot structures,
covalent bonding, didn't stop inventing.
So we're going to put U.S.A., with a Chevrolet out in
California.
And what did G.
N.
Lewis tell us?
He said, I want to extend the Bronsted-Lowry concept.
But you know, let me give you an analogy.
Have you ever stood on a bridge and
watched the traffic go?
You see it starts from a stop sign or a traffic light.
Light turns green, the first car pulls away, the second car
pulls away, the third car pulls away.
The other way you can look at it is, when the first car
pulls away, there's a car vacancy.
And instead of watching the cars go from right to left,
watch the car vacancy move from left to right.
And the two are linked.
And point of fact, if you're the fourth car from the
traffic light, light turns green and you
hit the horn, why?
You can't drive.
You have to drive only into a car vacancy, and it takes time
for the car vacancy to get to you.
When the rate of vacancy flux is not matched by the rate of
car flux, then we have a collision, and then we
exchange papers.
So it's a mass transport problem.
So now what G.
N.
Lewis said was, instead of looking at this reaction as a
proton transfer reaction, and looking at proton attachment,
or proton acquisition, he said, why don't we look at
this from the other side of the relationship?
And so he said, let's look at something like, for example,
we've got NH3 here.
So let's put the N with the 2 electrons.
And we've been talking about this from the perspective of
the bond being formed here through the
acquisition of the proton.
Lewis said, I can view this from the other perspective,
and say, I'm going to view it from the
perspective of the nitrogen.
From the perspective of the nitrogen, this reaction
represents the donation of the electron pair.
A very high concept.
So this base is not about proton attachment.
It's about donating an electronic pair.
So let's get that down, because that's really good.
So I'm going to say something that's a base looks like this.
This is what a base is.
It's a lone pair.
This is electron pair, very high level base.
And what's the proton give us?
What we're doing, is we're talking about
electrons and donation.
So what's this other thing in that same category?
An electron pair can match up with an empty orbital.
That's the only place electron pairs can go.
They go to the empty orbitals.
So electron pair is a base, and this is just a circle
around nothing.
OK?
This is nothing.
But it's a special kind of nothing.
This is called a vacant orbital.
And I'm going to call a vacant orbital an acid, in the most
general sense.
So now I'm going to take an electron pair plus vacant
orbital, and what did they react to give me?
Covalent bond.
You'd expect that from G.
N.
Lewis, because G.
N.
Lewis enunciated covalency.
That's why you'd expect it from G.
N.
Lewis.
So this is, vacant orbital plus electron pair gives
covalent bond.
Very high level.
So what is a base?
A base is now not just a proton acceptor.
It's an electron pair donor, and the acid is an electron
pair acceptor.
So now there's no chemistry here.
No chemical identities.
Which means, I can even go from--
I don't even have to be in solution.
I don't have to be in a liquid.
This could be a gas and a solid.
Could be a gas plus a gas.
So this is, like, Darwinian.
You know, we've come out of the primordial ooze, and now
we can fly, because we can talk about any chemical
reaction we want.
Any state of matter.
It's fantastic.
Absolutely fantastic.
So here, let me do one last color drawing.
We want to do this one here.
This is Lewis and the biggest concept.
So we're going to go like this.
Base like this.
Base with its lone pair hanging out
plus acid gives us--
three colors.
This is great.
We've got all this colored chalk, it's Friday,
what could be better.
All right.
So now we're going to end up with A, B, there it is.
This is the Lewis concept.
Lewis acid-base concept.
In the broadest thing.
Now I'm going to show you an example of
how you can use this.
I'm not going to stop here.
OK, let's go.
All right.
So we're going to talk about acid rain
from burning of coal.
You know, about half of the electric power in this country
comes from burning coal, and coal contains about 1% sulfur.
You can purify it of sulfur, but it takes money to do so.
A ton of coal will give you about 25 million British
thermal units, and that's the number in SI units and joules.
So 3 tons of coal will give you 1 megawatt per day.
By the way, it takes about one gram of uranium
to do the same thing.
So someday, when you're sitting there as a
policymaker, you've got a choice of 3 tons of coal or
one gram of uranium, keep this fact in mind.
A 10 megawatt plant burns 30 tons of coal per day,
containing a third of a ton of sulphur, which makes 2/3 of a
ton of SO2.
And SO2 is a precursor to acid rain.
Now we can reduce that SO2 emissions by reacting SO2 with
a lime, CaO, according to this reaction, to
make calcium sulfite.
And now I'm going to bring in this.
Lewis.
Oh, here's from your textbook, there's the scrubbers with the
calcium oxide, and water missed, and so on, and
eventually they trap the SO2.
They do nothing to the CO2.
So all the CO2 is going up.
So here's calcium oxide, which is a solid.
It's lime.
You know, stuff you sprinkle on football fields.
And this is SO2, and it's got this structure with the
resident bond double single.
The oxygen here acts the way it did in the
glasses, as a modifier.
It goes in and breaks this double bond, and now we have
three single bonds.
So we can look at this from Lewis acid-base.
You know, calcium oxides, electron pair donor.
If it's an electron pair donor, it's a base.
So it's a base.
And you know SO2 better be an acid, or your theory is nuts,
because SO2 is a precursor to acid rain.
So it has to be an acid.
It's an electronic pair acceptor,
which is what we said.
It's an electron pair acceptor.
So this is a Lewis acid-base reaction that
is a gas-solid reaction.
So we started with aqueous solutions, and we end up
generalizing all of these types of reactions.
So that's pretty good.
Here's a plot of acid concentration
up in the sky, here.
You're Here.
Yeah.
Last thing I'll do, I talked a lot about Lavoisier and
Scheele and Priestley.
This is a play, if you've got the little bit
of time this weekend.
You might try reading it.
This is Carl Djerassi and Roald Hoffman, both Nobel
Prize winners in Chemistry.
Not the same year.
They won independent Nobel Prizes, and actually speak to
each other, and they collaborated on this play.
Here's the setup, the premise.
The premise of the play is that the Nobel committee
decides to give Nobel Prizes retrograde.
Before 1901.
You see, you can't get a Nobel Prize posthumously.
You've got to do something great, and you've got to keep
on living until you get the prize.
So they decided they're going to go backwards, you know,
maybe give prizes to people like, I don't know, Maxwell or
Faraday or whatever.
And they decide, it should be simple in the old days.
Because we didn't have all this dog-eat-dog competition
in science, and big grants, and corporate interests.
Should be simple.
Until they hit the Nobel Prize for the discovery of oxygen.
And then they hit the wall with these three different
competitors.
And so the story goes that the three competitors and their
wives end up in Stockholm, and the
interactions the take place.
And it's really fascinating, because you have Lavoisier,
who was the political conservative but the chemical
radical, Priestley, who was the chemical radical but the
political conservative--
by the way, Lavoisier.
You know how he died?
He was guillotined during the French revolution.
And people think it's because he was a tax collector.
Not true.
He was a tax collector.
The real reason is Marat, who was one of the chief justices
of the Jacobin terror, was also a scientist. And around
1740, Marat had submitted an article for publication, and
it was rejected.
And Lavoisier was on the editorial board.
And so sometimes when I get an article to review, and I have
to give it a negative review, I think about this story.
And sometimes, I will just say, you know, I'm really busy
teaching 3091 and I won't get to this on a timely basis.
OK, people.
Have a nice weekend.