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PROFESSOR CIMA: OK, next one is, with these weak acids, we're going to use,
again, an equation that gets a name, Henderson Hasselbach, a couple of
physical chemists. They took this equation for a weak acid and
rewrote it in terms of log K_a.
So I've simplified my hydronium. I've just put it in as H+.
Or you can multiply everything by minus 1. And you get that the pH is equal to the pKa
plus the log of A- over HA, the acid.
STUDENT: What is that, is pKa? PROFESSOR CIMA: pKa, so it's going to be just
like, minus log K_a. It's just the same thing we've been doing
with pOH. And that's the Henderson Hasselbach equation.
So if you know the pKa, you can calculate the pH by finding out what
those two concentrations are. And where this is most useful is in the context
of what are called buffers.
A buffer is a weak acid plus the salt of a weak acid.
So what do I mean by a salt of a weak acid? So if we were talking, a weak acid is acetic
acid, then the salt of a weak acid would be something like sodium acetate.
And this is freely soluble in water. Now why this is significant is that, if I
look at this equation, pH equals pKa plus log A- over HA, like I said, I can
approximate this bottom one as how much weak acid I add.
Let's say that it's 0.1 molar. The top one has two sources, one from the
dissociation of the acid, and the other one from the free dissociation of the
salt. And that's huge, right?
I'm going to add roughly 0.1 molar to that. That's going to totally dominate the top one.
And so this is, basically, the salt concentration. So in other words, this ratio fixes the pH.
What do I mean by fix? Well, let's say I add a strong acid to my
buffer, which has both of these. Well, obviously, as soon as I add a strong
acid, that hydrogen ion I'm going to add reacts with the acetate ion to
form acetic acid again-- oops, H2O--
but did it change the pH? Well, if I take my equation here, I'm going
to subtract some acetate ion, depending on how much acid I added.
And that's going to form more acetic acid. And I think you can see what's going to happen.
This is not going to matter, if this is small compared with that, and if
this is small compared with that. So in other words, if these are both 0.1 molar,
the pH doesn't change until I've added at least 0.1 molar, what's called
the buffer capacity. So that impact of that is pretty--
so here's the normal behavior of water, the curve
that we just had there. And this is moles of HCL I added.
So for every mole of HCL I add, I can calculate what the concentration
hydrogen ion is, because it fully disassociates. But here's my buffer.
You see it just stays flat, until I reached it's buffer capacity.
And then boom, the pH starts dropping. No, so we are adding HCL, let's say.
And for every HCL, I produce one of these, because it's a strong acid.
But as soon as I do that, I shove this equilibrium to the right.
Remember, I'm using the general notation. But this, of course, in my problem for acetate
ion, is that one. So I will decrease this.
That should be AcO and AcOH. I'll decrease this through this reaction.
And I will increase this one through this reaction.
But if this and this are larger than the amount of acid I've added, the pH
stays fixed. And that's what the importance of a buffer
is. OK.
Yep? STUDENT: Can you use a weak acid plus salt
of a weak base? PROFESSOR CIMA: No.
STUDENT: No? PROFESSOR CIMA: So the salt of the weak acid
is-- well, it's the salt of that same acid, basically.
Any other questions?