Published in 1969, The Andromeda Strain was the book that
placed Michael Crichton on the map. He
was still in medical school, but managed to blend a biological crisis with true
life undertones to create his first (of many!) best-selling thrillers. The story opens with the discovery of
decimated Piedmont, Arizona – only the town drunk and a crying baby are found
alive. As revealed at the end, the baby
lived through a biological attack that killed nearly all others because his incessant crying slightly changed the pH of his blood and rendered him immune.
The importance of pH to blood, viruses, infection, and life are
numerous. Human blood must remain
between 7.35 and 7.45 at all times. However,
we are constantly placing acids (for example, orange juice!) and bases (yum – milk!)
into our bodies. How do we maintain
pH? How does anything maintain pH? The answer is through the use of buffers.
In my
post Salty Water, I discussed how an acid and base neutralize each other to form
water and salt (aptly named post, huh?).
That post was quite specific about acids and bases so I’m going to
simplify matters greatly in this post to make it easier to understand.
Let’s
begin with looking at an acid. An acid
has two parts: the H+ (orange circle) and the rest of the molecule (gray rectangle). The H+ is able to leave the rest
of the molecule (Figure 75.1). It is also able to come back together.
pH is a
measure of how many H+ ions are floating around in solution. The more H+ around, the lower the
pH; the less H+ around, the higher the pH (Figure
75.2). The key to visually
judging the pH is to look at how many free
orange circles are floating around.
Let’s
pretend we dropped some of our acid into a cup of water. This particular acid we are going to call a
weak acid. This means that some of the
molecules we dump in there are going to separate into H+ (orange circles)
and the rest of molecule (gray rectangle), but some are going to stay
together. You will end up with a mixture
of both (Figure 75.3). The pH of this solution is determined by how
many free orange circles are floating around (orange circles still bound to the rest
of the molecule do not count).
Okay. We’re all set! The picture in Figure
75.3 of the weak acid in water is a buffer. I know, right? You didn’t think it would be so simple. A buffer is simply a mixture of acid (H+
bound to the rest of the molecule) and its conjugate base (the rest of the
molecule alone). I’ve labeled
these parts in Figure 75.4.
Our
blood maintains its pH by having a mixture of weak acids and its conjugate
bases floating around. But, what happens
when there’s an influx of H+ from another place? How does the buffer maintain pH?
The
conjugate bases floating around are capable of binding to H+,
yes? I showed you that in Figure 75.1.
So, if a large amount of H+ are thrown in, then those
conjugate bases are capable of picking up those extra H+. Remember, pH is a measure of how many free H+
are floating around. If the conjugate
bases pick them up, they won’t be contributing to the pH. I’ve tried to visually show this in Figure 75.5. In this way, buffers are helping to maintain a constant concentration of H+ in solution.
This
goes on all day long – addition of H+, removal of H+ -
and the buffer responds accordingly such that the concentration of H+
(and thus
the pH) remains constant.
Obviously, though, there are limits to how well a buffer will work. What if you add more H+ than there
are conjugate bases? Well, you’re going
to end up with more H+ in solution and thus a lower pH. Scientists will talk about buffer capacity
when this problem crops up. This is just
a way for us to express how much acid and conjugate base are in solution and
how much acid/base we can add to it before it stops resisting change in pH.
The
small letter p is short-hand for the mathematical function –log. That means that pH = -log [H+], where brackets means concentration. pH isn’t the only kind of p around. You can talk about pOH (-log of the [OH-])
and pKa (-log of the Ka). Small
ps are used a lot in chemistry.
Why
does crying change the pH of blood? A
common weak acid used as a buffer in our bodies is carbonic acid, which is
directly related to how much carbon dioxide we breathe out. Crying alters our breathing patterns and
changes the concentration of carbonic acid in our blood, which changes our
buffer capacities and results in slight changes in pH.
NOTE: This is a VERY simplistic explanation of buffers. The main concepts are there, but I’ve chosen
to present it in this way so that the concepts could be understood without the
mess of strong acids, strong bases, weak bases and considering the whole
situation from both viewpoints. It’s
overwhelming if you aren’t familiar with these species or the idea of
equilibrium constants. Yes, what truly determines to pH of a buffer
is the ratio of weak acid:conjugate base or weak base:conjugate acid, as
explained by the Henderson-Hasselbach equation.
Yes, only certain weak acids or weak bases are useful at different pHs,
which is determined by the pKa or pKb, but I decided that
was beyond the scope of this post.
REFERENCES
Zumdahl, Steven S. “Chemical Principles, 4th
Edition” (2002) Houghton Mifflin Company, Boston, MA.
Crichton, Michael. “The Andromeda Strain.” (1969) Harper
Collins Publishers, New York, New York.
