Re: Stuck on a Biochemistry question

On 2006-04-30, raconte@xxxxxxxxxxx <raconte@xxxxxxxxxxx> wrote:

Jimchip wrote:
On 2006-04-29, ///Owen\\\ <mo-jo@xxxxxxx> wrote:
Ash wrote:
"Jimchip" <jimchip@xxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxxx> wrote in
message news:1256hendogsal09@xxxxxxxxxxxxxxxxxxxxx
On 2006-04-29, Ash <mail@xxxxxxxx> wrote:
If anyone out there has done this before, and is bored enough to
have a look
at it again, any help would be greatly appreciated (its the last
question and i cant do it :P):

"A newly discovered insect squirts out formic acid (Ka = 1.8 x
10-4) as a defensive mechanism. An analysis of this liquid shows
the concentration of
formate ion to be 0.015M and the total concentration of formate plus
formic
acid to be 1.45M. What is the pH of this liquid? Hint: pH = pKa +
log([A]/[HA])"

Formic acid= HA
Formate = A-

pH = pKa + log([formate])/[formic acid])

--
Henderson-Hasselbach equation

How come such a simple equation (it's only a rearrangemt of HA === H+ + A-
and logged) earned two people immortality?

Well, Henderson's work is almost 100 years old (1908). Think about what it
was like back then in terms of the quantitative description of matter,
chemical bonding theory, and especially acid-base theory. Arrhenius was only
~20 years before and Lowry and Bronsted were 15 years later. Hasselbalch's
log formulation was ~9 years before Lowry and Bronsted.

Add to that the fact that it is an elegant mathematical description of a
rather complex system involving coupled equilibria, that it is a model for
the modern analysis of pharmacological "intrinsic activity", and that it is
so simple that anyone can use it, or a rearranged version, in very important
applications.

--
Elegant and profound.
'Simple' is a compliment when the subject is deep.

That, and it does come in handy. If you want to mix up a buffer, and
you have the acid and base forms handy, you can just mix the correct
ratio, and the pH will be very close. That was probably more important
years ago, when pH meters were cranky devices that never seemed to be
stable -- or even non-existant.

I could have added above that it is still being used 100 years later because
no one has improved on it.

Sometimes I don't trust the electronic calibration of our pH meters and a
little bit of a quickly made buffer in a constant temp. water bath can be a
very nice check. Sometimes I don't trust "dilut-it" buffer.

There's also the nice little diddly

[A-]
(pH-pKa) ------
10 = [HA]

for estimating relative protonation using a pH meter. Very handy with many
aqueous reactions.

--
Maybe I'm just not a trusting person
.

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