Sal: In the last video, we learned that if you just leave
water to itself, it autoionizes.
So let me see.
I have molecules of water.
There's some probability, and it's an equilibrium, that one
of the hydrogens on one of the waters bumps off and just
joins the other one.
So you could end up with a hydronium ion.
It's an ion because it has an extra hydrogen.
So I'll put a plus charge there.
And obviously, everything you hear is
in an aqueous solution.
I could write aqueous everywhere.
aq.
I could write aq.
It's kind of obvious.
Water is obviously an aqueous solution.
It's obviously dissolved in water.
And then plus a hydroxide ion.
This is the water molecule that lost its hydrogen atom,
or its proton.
Because we know that a hydrogen atom-- or at least a
hydrogen ion-- if you get rid of its electron, all you have
left is a proton.
And that's also in an aqueous solution.
And we learned that people really care--
We learned about the equilibrium
constant of this reaction.
And we also learned that people really care about the
concentration of this right here.
Of the hydronium.
And let me write this again.
And I went over this in the last video.
But sometimes the same exact autoionization reaction will
be written like this.
H2O, disassociating in itself, really, into a hydrogen ion in
an aqueous solution plus a hydroxide anion in an aqueous
solution or a negative ion.
These are the same thing, and I write it this way because
this is actually the state that happens.
That you don't actually have these protons just
independently, they actually do join onto another water
molecule and form a hydronium ion.
But these are the same exact concept.
But we learned in the last video that people really care
about the concentration of-- you could pick
one of these two.
They normally write this one, right here.
So people really care about the concentration of your
hydrogen ions.
And we learned in the last video that just pure water at
25 degrees Celsius, which is room temperature.
So let me write that.
At 25 degrees Celsius.
The concentration of your hydrogen protons-- because
that's really what they are, or hydrogen ions-- is 10 to
the minus 7 molars.
And then we even said, hey, chemists, for whatever reason,
don't like dealing with negative exponents.
So they defined the pH as equal to the minus log base 10
of your hydrogen concentration.
And of course, that's equal to minus log base 10 of 10 to the
minus 7 for pure water, at 25 degrees, which is equal to 7.
Fair enough.
And you can probably imagine, if people took the trouble of
constructing this pH thing and saying it's the negative log,
they must really care about what the hydrogen
concentration of water is.
And it does matter to-- in a whole bunch of-- especially
actually biological systems-- all the time, people worry
about the pH of this or that.
And obviously I talked in the last video
about pH-balanced cosmetics.
All sorts of bad things might happen to your skin if the pH
of the cosmetic is different than your skin.
I don't know if that's a fact, but I guess that's the that's
the theory behind it.
So there's obviously things in the world that change the
hydrogen concentration.
That change this concentration.
So, when we learned about Le Chatelier's Principle, we said
hey, maybe if we add more hydrogen on this side, the
equilibrium will change and this concentration will go up.
And so we have more hydrogen concentration.
I guess on the other side, you could add more hydroxide and
then you would have more OH concentration, and maybe a
little less hydrogen concentration.
So if you if you care about things that are going to do
that to this reaction, you probably want to
have a name for them.
And the name for them is acids and bases.
And I'm sure you've heard both of these words before.
So, let me write them both down.
Acid and base.
Now.
Just to over-complicate your life a little bit more, there
are actually multiple definitions of an acid and
multiple definitions of a base.
And they kind of become more and more broad
in what they imply.
The definition that you're probably going to use the
most-- the definition that I always use in my head whenever
someone says an acid-- is it increases hydrogen
concentration when you place it in water.
And a base is something that increases your hydroxide
concentration in water.
This is the definition that you're probably going to see
in 90%-- especially in a first year chemistry class.
And this is called the Arrhenius acid or base.
I hope I get his spelling right.
You should never take any of the my spelling
words at face value.
You should look them up because I'm clearly not a
master speller.
Arrhenius.
Arrhenius, with an H.
These are Arrhenius acid, Arrhenius base.
Let me give you an example.
If I were to take some hydrochloric acid-- that's
actually what it's called, but you could say hydrogen
chloride-- in an aqueous solution.
It actually disassociates completely.
This is not an equilibrium reaction.
So let me put it in this nice-- let me
use a bolder color.
This nice one-way arrow, which says, hey, this isn't one of
these wishy-washy reactions that go in both directions.
This is a strong acid.
It completely disassociates in water into-- let me put the
hydrogen in a different color.
It completely disassociates in water into a hydrogen ion.
An aqueous solution plus chloride.
So it's just a chloride negative ion
in an aqueous solution.
And they just float in there.
So you could imagine if you place hydrogen chloride in
water, you're going to increase its overall hydrogen
concentration.
And the same exact reaction could have been
written like this.
Just to hit the point home.
Hydrogen chloride in an aqueous solution.
And then you can say plus h-- let me do it
in a different color.
Plus H2O.
It's a one-way reaction.
And then you end up with, and I think you understand.
Obviously not all 3 of them are-- H3O, that's too dark.
H3O plus-- it donated its hydrogen proton
to this water molecule.
And then you have plus the chloride anion
in an aqueous solution.
And these acids that completely disassociate-- in
future videos, we're going to look at ones that are a little
bit more wishy-washy.
That don't completely diaassociate.
But these acids that completely disassociate are
called strong acids.
And strong isn't just a nice word
that they use in chemistry.
It literally means-- when someone says a strong acid,
something that completely disassociates into water.
It's a one-way reaction.
And all of a strong acids-- you could probably even guess
based on their chemical make-up.
But hydrochloric acid is one of them.
You also have hydrogen bromide.
Another strong acid.
Hydrogen iodide.
You have nitric acid.
I think I showed you that in a previous--
All of these, when you when you put them into water, this
little h is going to pop off.
You join another water molecule, form hydronium.
And then this molecule right here, in these cases-- well,
in the case of chloride-- these are halogens.
But they're going to go and form negative ions.
Whatever's left over.
Nitric acid.
Then you have sulfuric acid.
You've probably heard of that.
Very strong acid.
Sulfuric acid and then perchloric acid.
These are the strong acids.
O4.
And these are good to know, because if you see them on a
chemistry exam, you know that these are going to completely
disassociate.
And remember, we keep using this word acid.
What does acid mean?
It means that these are strong acids.
These completely disassociate.
If you like Arrhenius's definition of what an acid is,
they increase your hydrogen concentration.
Because obviously when you throw this in some water, all
of these new hydronium ions are going to be formed, and
they're going to overall increase your
concentration of them.
And that's why they're called an acid.
And now if you went on the strong base side, by
Arrhenius's definition, the base is something that creates
hydroxide ions, or anions in water.
If you look at the periodic table, any of your group 1
elements, and group 1 elements are your alkali earth metals.
Your alkali earth metals are bonded with hydroxide.
If you put them in water, the hydroxide pops off.
So let me just-- has to do with lithium or sodium.
So for example, these are strong bases.
So if I take some lithium the hydroxide
in an aqueous solution.
This completely disassociates.
No equilibrium here.
This is a strong base.
So what's it going to disassociate to?
Well the little hydroxide ion is going to pop off.
That's going to be a minus.
It's in water still.
Plus a lithium ion.
Right there.
So this will always disassociate in water.
The same thing is if you have sodium.
Sodium's going to do the same thing.
Sodium hydroxide in an aqueous solution.
One-way reaction.
You produce hydroxide and some sodium cations.
And so you imagine, when you put these in water, this
concentration is going to go up.
It's just like what we did in Le Chetalier's Principal.
Where we said, oh, what if you add some of this or some of
that to an equilibrium reaction.
Well, how do you add it?
Well, in this case, you can add a strong
acid or a strong base.
Now, everything I've done so far-- this is the Arrhenius
definition-- where an acid increases your hydrogen
concentration, a base increases your hydroxide
concentration.
Now that is what you're going to see 90% of the time.
But there is a slightly broader definition out there.
And I don't know the correct-- I always say
just Bronsted, but.
If you don't have fancy fonts, it's spelled like this.
Bronsted-Lowry acid or base.
If you if you have good fonts, there's usually this little
cross at the o.
So I don't know.
Well, Bronsted-Lowry acid.
And all of these would also be Bronsted-Lowry acid.
But the broadening of the definition is an acid is a
proton donor.
And a base is a proton acceptor.
So let's look at this definition in the context of
everything we've just done so far.
So if you live buy the Bronsted-Lowry definition,
what is a proton donor here?
Well, this hydrochloric acid is-- let me look at this
reaction right here.
This hydrochloric acid is donating a proton to this
water molecule, right?
A hydrogen atom is just a proton.
That's an important thing to remember, because it had-- a
hydrogen ion.
Because if you get rid of its electron, it has no neutron.
It's just a proton sitting in space.
So this hydrochloric acid is donating a proton to this
water molecule to make a hydronium molecule, and it
gave it away.
So this is a Bronsted-Lowry acid as well.
Just as it was an Arrhenius acid.
Let's see the Bronsted-Lowry version of a base.
OK.
A base is someone who-- proton acceptor.
So this guy, he had this nice little relationship with this
hydroxide ion.
The hydroxide ion got disassociated, and it becomes
a little bit fungible right here.
You can kind of say hey, this guy accepted
this positive charge.
And, you know, it's a little bit wishy-washy here.
Because it's not like someone gave it a proton.
You can kind of think of it as he
actually gave away electrons.
But if you just look at the final product.
OK, he's got a positive charge when everything was done.
So you say OK, maybe that's a Bronsted-Lowry base.
Now you're saying, OK well, why did people even make the
trouble of defining a Bronsted-Lowry base when all
of the Arrhenius acids and bases could also be
Bronsted-Lowry.
Well that's because Arrhenius is always--
You're always dealing with water.
Everything's in an aqueous solution.
But I drew an example here of a Bronsted-Lowry base.
It doesn't have to be an aqueous solution.
So if you have acetic acid right here--
that's what's in vinegar.
Plus ammonia-- it doesn't have to be in an aqueous solution--
what happens?
This hydrogen gets donated to the
ammonia ion to make ammonium.
So this becomes positive, this because negative.
It donated this proton.
And Arrhenius has nothing to say for this.
Because everything he deals with is hydronium and water.
But the Bronsted-Lowry definition
works in the situation.
Now the broadest definition-- although they're all these
definitions, but you're going to be 90% of the time good if
you just know the Arrhenius definition and just know the
Bronsted-Lowry and what I'm about to say
right now also exists.
And that's Lewis acids and bases.
Now Lewis cares about electrons.
Bronsted-Lowry cared about protons.
So Lewis, instead of saying an acid is a proton donor, Lewis
acid says it's an electron acceptor.
And a base is an electron donor.
Now, let's us look at in the context of everything we've
done so far.
If this is really an acid, it should
be an electron acceptor.
And the way you can kind of think of it is you had
hydrochloric acid before, when this guy gave away this atom
or this proton right there.
He kind of kept his elecron.
So he was kind of an electron acceptor.
It's a little bit of a grey area there.
It's not like he took an electron from somebody else.
But this would be considered still a Lewis acid.
And if you think about a Lewis base, a Lewis base is what?
It's an electron donor, right?
So if you have it right here, Lewis base
is an electron donor.
This lithium hydroxide, when it goes into water-- There's a
couple of ways to think of it.
You can say hey, it's donating this OH, which has these 2
extra-- this negative charge, so it's an electron donor.
I don't lik it.
It's kind of squishy that way, so I'm not a big fan of that
definition.
In my mind, Arrhenius makes the most sense.
It's kind of the purist thing.
Are you creating hydronium or not?
Or are you creating hydroxide or not?
But just to show that the Lewis acid base definition is
the broadest definition.
Here's a case of a Lewis acid base reaction.
That would not be considered either a Bronsted-Lowry or an
Arrhenius base.
Because this doesn't have to happen in water.
And what you have here is boron trifluoride, with a
fluoride anion.
It has a negative charge.
It has this little extra electron here.
In this situation, this fluoride right here, or this
fluorine anion can donate these 2 electrons.
So let's say that this one right here is
this one right here.
And what it does is, it donates these 2 electrons to
this boron complex right here.
And so if it's the electron donor, what is it?
The fluoride is the base.
So this is Lewis base.
And then the electron acceptor is the Lewis acid.
So this right here is the Lewis Acid.
It's good to know that these definitions exist, and
especially so you don't get confused in the future.
But for a first-year chemistry class, if you know the
Arrhenius definition really well, you'll
do very, very well.
And frankly in my mind, that's the easiest one to
contemplate.
And that's the one that's going to matter in most of the
reactions where you deal with the changes in pH to water.
Because that's what all the acids and bases do, especially
in the Arrhenius case.
Anyway, see you in the next video and watch as we do some
math, figuring out the changes in pH due to some acid.