Hi. It's Mr. Andersen and this is chemistry essentials video 69. It's on pH and buffers.
The proteins in our blood have a problem. They have to have a specific pH. And if it
changes radically out of this range between 7.35 and 7.45, they start to denature and
they can't do the job that they're intended to do which is to carry oxygen and carbon
dioxide. Thankfully we can use a buffering system. So what happens is the carbonic acid
that is created when we add carbon dioxide to the water is a weak acid. And it has a
conjugate base. And so that creates what's called a buffer solution. What does that mean?
If we add more protons to it it will simply push toward the left. And if we add more hydroxide
to it it will push it more towards the right. And so it keeps our pH fairly stable. And
that's how buffer solutions work. So pH remember is based on the proton availability. It's
the concentration of that proton in solution. And so we want to keep that as stable as we
can. And so we use a buffer solution to do it which is essentially a weak acid and its
conjugate base. And so what's going to affect the pH of that reversible reaction? Well the
first thing is the pKa, which is going to be the equilibrium constant. And so if we
can keep that equal to our pH or around our pH that's going to keep our pH stable. And
also we could look at the concentration of the acid, that weak acid, to its conjugate
base. If we can keep those values equal as well, that's going to increase the buffer
capacity. And we'll look at that algebraically in just a second. But big picture, what are
we doing here? Well we've got a weak acid and a reversible reaction that forms this
hydronium ion and then its conjugate base. And so in a good buffer solution we want the
weak acid and the conjugate base to be equal in values. And so what happens? Let's say
we add hydronium to that. Let's say we add a strong acid to that. Well LeChatelier's
principle tells us, if we add more of it on this side it's simply going to push it in
the other direction. So it's going to push it more towards that weak acid side. But since
those values are equal, it's not going to change it that much. And our pH value is not
going to change very much as well. If we look at adding a base now, if we add a hydroxide
to it, what is it going to do? It's simply going to push it more toward the right. And
so what happens is we can add strong acids and we can add strong bases and it's going
to keep that pH around a stable set point. And so let's look at this as an equilibrium
equation. And so if we look at our equilibrium constant, if we were to write it out, how
do we do that again? It's simply going to be the concentration of our two products over
the concentration of our reactant. And so if we do a little big of manipulation algebraically
what we can do is isolate the concentration of those hydronium ions on the left side.
What is that? Remember we take the negative log of that. That's going to be our pH. And
we want to keep that as stable as we can. So if it's a good buffer solution how do we
keep it as stable as we can? Well if we keep our Ka value equal to our concentration of
our hydronium ion or if we keep our pH equal to pKa value, that's going to create a good
buffer. If one of those is much larger than the other one, changes in one will change
the other. Also we want to look at equal concentrations of that weak acid and its conjugate base.
And so if we can keep those equal to 1 we can have large changes in that. Ten-fold changes
in that will only change the pH value a total of 1. And so we want to keep those values
of the weak acid and conjugate base equal to each other. And also we can use our pKa
values, which is remember looking at the concentration of reactants and products. And we can figure
out what's going on in the reaction. So if our pH value is less than our pKa that means
we have more of this weak acid. And if it's greater than our pKA that means that we have
more of the base. And so if you think of it like this, if pH goes down we've got more
of the acid. And if pH goes up then we've got more of the base over on this side. So
what are some good applications of that? Well an acid-base indicator is a great example
of that. So if we're looking at bromothymol blue, so what color is that going to be? If
we're in a neutral solution it's going to be right at 7. So our pH and our pKa values
are essentially equal to each other. And so what happens if our pH value goes down? Well
that's going to shift it more towards the left. And so we're going to have more of this
form of bromothymol blue which is going to give us that yellow color. What happens if
we go to the right, that's going to give us more of this blue color in relation to that
neutral. And so we can see changes in the color of that indicator and what that's telling
us is changes in the pH. This is also important biologically. Remember proteins are made up
of an amino acid. And each of those amino acids are going to have a different side chain
which is going to be if we look at everyone of these amino acids, the top part is identical.
But each of the side chains that drop off the bottom are going to be different. And
so each of these have a different pKa value. And so if we change the pH of the overall
protein, so this is the myoglobin, for example, it's going to change the behavior of each
of those side chains and the amino acid inside it. And so did you learn that when we're creating
a good buffer solution we want to keep our pH and pKa values equal to each other? And
did you learn that changes in the pH related to the pKa tells us if we're moving more towards
the left, more of the acid or more of the base? And then could you design a good buffer
solution? Remember what we've simply got is a weak acid on the left side and its conjugate
base on the right side. I hope so. And I hope that was helpful.