Ok Aaron, the buffering ability of the bicarbonate ion is well known. The
solubility of CaCO3 decreases with increased temps above normal. This does
not mean that the solubility increases with decreasing temps below ambient
(look up buffers). In cooler temps, CaCO3 precipitates since the solubility
coefficient is so low. The major factor here is the acidity which negates
LeChatelier and the above discussion since CO2 is no longer one of the major
players. Acetates, humic acids, chlorides, and nitrates of calcium are
soluble and plentiful in peat moss mixes which overwhelms the buffering
capacity of the bicarbonate ion, which was the point. I am a high school
chemistry teacher of 34 years and put the 50's in my pocket, thank you. But
as any other high school teacher knows, they leak back into the classroom
anyway.
Gary
"Aaron Hicks" wrote in message
...
"Gareth Wills" spaketh thusly:

Sorry to disagree, in part, Steve,but Fowlie is wrong. The only chemicals
(barring a few weirdos) that dissolve better in cooler temps than warm
ones are gases such as oxygen. Calcium compounds are notoriously
insoluble except in acids.

Then, unfortunately for your system, you are going to have to lump
calcium carbonate in with the "few weirdos" as its solubility decreases as
temperature increases. I base my claims as regards to this paradox on
statements made in two separate copies of the CRC Handbook of Physics and
Chemistry, my attending graduate school in stable isotope geochemistry
(even if I wasn't all there at the time), and perhaps most importantly,
the laws of physics and chemistry.

Poor, neglected calcium carbonate, every caver's best friend. The
solubility of calcium carbonate is tied to the concentration of the
bicarbonate anion, which is present in solution in low concentrations
thanks to atmospheric carbon dioxide. An increase in the temperature
causes carbon dioxide to become less soluble in solution and, along with
it, the solubility of calcium carbonate also decreases. You may thank
LeChatelier's Principle for this (although, if you wish to thank me
instead, 20's and 50's are welcome- send them to your local high school
chemistry teacher). An excellent primer, if you're bored to death, is
found he

http://members.aol.com/profchm/comonion.html

Back to carbonate solubility!

Behold! The caver's formula:

CO2 + H20 -- H2CO3

Carbon dioxide plus water gives you carbonic acid.

Carbonic acid loses its first proton- a moment we should all pause
to reflect upon, but not he privately. When you're done, come back;
we'll still be here. Really.

Pure water has pH 7.0; let it sit around and pretend it's Perrier,
next thing you know, those water molecules are all over the carbon dioxide
like a dog on a Dreamsicle, and the pH plunges to 5.6 with atmospheric
concentrations of carbon dioxide (about 380 ppm or, under new Bush
administration projections, 2,280 ppm by 2040). Higher concentrations of
atmospheric carbon dioxide drops the pH even more, and heaven only knows
what this means for the planet's coral reefs; the ocean has the bajeezus
buffered out of it, but when you're talking about something that massive,
anything goes. Maybe upping the temperature will throw the equilibrium in
the opposite direction. :-P At least the "coral calcium" crap at the
health food store will be cheaper. Try not to let too much too much
strontium build up in your system. "Tums" are made with synthetic calcium
carbonate; corals accumulate Sr, as well as lead and other heavy nasties
at levels that you do not want to be ingesting.

On the bright side, more CO2 = more caves, the very thought of
which makes us troglodytes revel in ecstasy. But the really big caves are
formed with sulfuric acid, which is probably biogeochemical in origin.
That's another lecture. Bring a "blue book," there'll be a short quiz
afterwards.

Cheers,

-AJHicks
Chandler, AZ