Let's say I have a bucket of aqueous solution, and it
contains some weak acid in it.
And so weak acid.
We've done this multiple times.
It's in equilibrium with its dissociated state.
So it would be in equilibrium.
Let me write the aqueous.
So it's a weak acid, where the A could be any other kind of
molecular group.
That could be the conjugate base of the weak acid.
So that's in equilibrium with some hydrogen ions, so it'll
produce some of them, but some of this is going to
be around as well.
Plus its conjugate base.
A minus.
Aqueous.
And just like we did with the strong acid titration, let's
measure it's pH.
So I'll do that right here.
Actually, I want to do it a little higher.
So let's do it right there.
Let me make a little scale here.
So this is 0.
Let's say that right there is 7.
That this right here is 14.
So our first measurement of the pH of our solution-- we
have a maybe a liter of it.
Let's say we have 1 liter of our solution.
We don't know the concentration of this thing.
But let's say our first measurement, we go and we say
OK, it's got a pH of-- I don't know.
Something here.
Let's say it's got a pH of 3.
So it's acidic, not ultra acidic, but it's acidic.
So it's got a pH of 3 right there.
This whole scale on the left-hand side.
This is all pH, just like in the last video.
This is pH.
The pH is going up as you go up the scale.
That's our first measurement.
Now we're going to titrate this weak acid, just like we
titrated the strong base.
And normally when you do this-- I mean, you can, in
theory, titrate with anything.
But you normally titrate-- if you're titrating an acid,
whether it's weak or strong-- you always titrate with a
strong base.
And if you are titrating a base, whether it's weaker
strong, the reagent, or the thing that you're going to
add, is going to be a strong acid.
So in this case, even though we're going to drop some drops
into a weak acid, we're going to drop
drops of a strong base.
So let's use sodium hydroxide again.
NaOH.
And just to make things interesting, to make the math
a little interesting, let's say it's a 0.5 molar solution.
And we're going to add it in increments of little
100-milliliter increments.
A tenth of a liter.
But the interesting thing is really the shape of the curve
that forms, or the titration curve, or the pH curve.
Let me write what the reaction of NaOH.
Just so you have that in your head.
NaOH, sodium hydroxide, and I did this in
the previous video.
It disassociates completely in water.
And so it turns into Na plus-- all this is aqueous, of
course-- plus OH minus.
So when you throw even a little bit.
Let's say you throw in-- this is 0.5 molar solution-- 100
milliliters.
That's equal to one tenth of a liter.
This 0.5 molar, that's its molarity.
So it says hey, I'm going to have half a
mole of this stuff.
Or really, of this stuff.
I'm going to have half a mole of hydroxide
ions for every liter.
So in one tenth of a liter, that means that I'm
going to have what?
I'm going to have half of that.
So I'm going have 1/20, or 0.05 moles.
So let's say this point right here's is 100 milliliters.
I've added 100 milliliters of our solution.
That's going to have 0.5 moles of OH.
Because remember, we could say it's 0.5 moles of this, but
whenever this is in solution, it disassociates
completely into this.
So if you have 0.5 moles of this, you really have 0.5
moles of this.
This will be in this form completely, because it's a
strong base.
So if I have 0.5 moles of this and I throw it into the
solution, what's it going to do?
It's going to start sopping up some of our hydrogen ions.
So it's going to increase the pH, and it's going to keep
increasing.
Right?
And this is a buffer.
And this is a key thing to remember.
This is going to sop some of this thing up.
But this is a buffer.
This is a weak acid.
So it's in equilibrium.
So it has this reserve over here of the acid form.
So when you start lowering this by adding a strong base
to it, to sop this up, what's going to happen?
The reaction is going to move in this direction to kind of
backfill, to fill or to almost replace.
It won't get quite to where it was before, but it's going to
dampen the stress on the equilibrium.
So what's going to happen is the concentration of your weak
acid is going to go down.
It's going to try to backfill this stuff.
But every time the reaction goes in this direction, this
is going to go up.
Right?
Remember, nothing is sopping up the conjugate
base up here, right?
We're just sopping up the hydrogen ions.
As we sop up more and more hydrogen ions, this reaction
goes more and more in the rightward direction.
But as it goes more and more in the rightward direction,
we're just adding more and more to this.
And the base just takes more and more of this out.
So essentially when the base is taking this out, it's
really taking this out-- or it's taking half of this out.
It's taking the H part of this out.
Because that converts to that, and it gets sopped up by the
by the reagent.
The strong base.
And all of this ends up on this side, as
the conjugate base.
So at some point, when we have added the exact same amount of
OH as there is of this stuff, as there is of this stuff plus
the hydrogen.
When we've sopped up all of the hydrogen,
what's going to happen?
Well, right when you get above that point, when you start
adding even more, you're going to start getting very basic.
It's just going to skyrocket up the pH.
Because then you're adding OHs, and they're not being
canceled out by any hydrogen over there.
And there's nothing to kind of backfill the hydrogen.
So you're going to get to a very high reaction.
When we add enough of the sodium hydroxide that this
completely neutralizes this and this, and all we
have left is that.
Well, you might say, oh, we have a pH of 7 because all of
the acid is neutralized.
But we don't have a pH of 7 like we did with the strong
base, because here, as we sopped up stuff, we kept
adding conjugate base over here.
Remember, and it's a real base.
It actually has basic properties.
If this is a strong acid, if this was hydrogen chloride,
I'd still have the chloride over here.
But chloride, even though it's the conjugate base, it has no
basic properties.
It does not increase the pH of a solution.
But this stuff right here does increase the pH of a solution.
And we've done that in previous videos.
So once all of this is sopped up, your equivalence point, or
the point at which the number of moles of this is equal to
the number of moles of this, is going to have a high pH.
So it's going to be like right over there.
And then as you add more and more of your reagent, or your
strong base, your pH is going to get higher and higher and
approach 14 or even go above it.
So your titration curve is going look
something like this.
And once again, you look at the steepest
point on the curve.
It's right there.
And you say, OK, that's the equivalence point.
That's the point at which I had an equal amount of
hydroxide having sopped up all of the hydrogen, and I've run
out of all of this stuff to backfill the hydrogen.
So you look at your equivalence point.
And you say, OK, if that equivalence point was when I
added, I don't know.
Maybe that's when you added 2 liters of your reagent.
So you added 2 liters of your reagent.
Your basic reagent, right here.
Or your titrant.
Or titrator.
I think that's another word for it.
How many moles have you added?
Well it's 2 liters, it's 0.5 molar, so you've added 1 mole.
So here you've added 1 mole of OH.
And when you added 1 mole of OH, you sopped up all of the
whatever mole you had of this.
So you clearly had 1 mole of that.
So over here, your original concentration of your-- I
mean, you really could say your concentration of your HA
plus your initial equilibrium hydrogen is equal to 1 molar.
Now in most chemistry classes, this number is way bigger than
this number for most weak acids.
And you could look at that.
Because if you look at some of the pKa's for some weak acids,
you'll see that this is some-- well, I won't go into the
math, but this number's a lot lower.
So this is an indication, essentially, of your initial
concentration of your weak acid.
So you essentially just keep titrating it, figure out the
inflection point, you say, the inflection point happened when
I added 2 liters of the titrator to it, which
corresponded to 1 mole.
So therefore I must have had 1 mole of my original acid in
the equation.
Another interesting thing.
So if you said OK, I had 1 mole of my original acid in
the equation.
So you say, this was 1 molar originally, before I started
titrating it all.
And if I said that this is a 1 liter solution-- this is 1
mole, not 1 molar.
But if I know it's a 1-liter solution, now I know that
originally I had 1 mole in 1 liter, so I had
a molarity of 1.
Now, let me go over something else that's interesting about
a basic reaction.
Actually, let me do that in the next video, because I
think this one's getting a little long.