pH and Buffers

pH and Buffers

October 11, 2019 94 By Ewald Bahringer


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.