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re: pH hell [2nd reply]

First a little background philosophy to support my stance. Forgive me if you
find this repetitive or basic.

I mentioned that an acid and its salt are required to form a buffer. It's a
two-part requirement in order to attain the stability imparted by the
buffering agent. Both sides of an acid reaction revolve around hydrogen ion
exchanges, leaving both proton donors and acceptors available to "absorb"
changes of state. Stability of the buffer can be increased or softened
through varying concentrations while pH or pOH controlled by means of the
mix's proportional ratios.

(For those with a penchant for numbers, stability is optimized at
concentrations in the range of 0.1 - 1.0 Molar. A working pH range of
plus/minus 1.0 around the equilibrium can be achieved by varying the
acid-to-salt ratio from 1:10 through 10:1.)

Using two or more competing buffers to increase the working range of pH
values also increases the amount of change between the endpoints, as you
have discovered. This makes true stability a fragile thing, balanced
essentially on knife-edge. And since we're still working within log scales,
it's not at all difficult to go wrong. Imagine the havoc this would create
in natural systems, then ask yourself, "Well then, just why do *we* try to
work that way?"

Natural buffering and pH control depends chiefly on one simple ingredient:
carbon dioxide. Not solely in its gaseous state, though, but in vast stores
of carbonates and bicarbonates and the carbonic acid providing the system's
driving force. Think about the implications of that for a moment.

We're all aware of the carbonic acid cycle, but I bring it up for
illustrative purposes. Introducing gaseous carbon dioxide to an aqueous
environment produces carbonic acid, bicarbonates and carbonates along the
lines of

CO2 + H2O <=> H2CO3 <=> H+ + HCO3- <=> 2H+ + CO3--

This is a *really* handy equation that can be entered into and altered at
any stage to give you almost complete control over a wide range of values.
The form of carbon dioxide that you add or remove in these distinct stages
couldn't make adjustments much easier. And best of all, it's totally
self-reliant in a subtle way - while the surrounding atmosphere will
continue to contain all of the CO2 you can pump into it, water can dissolve
only limited amounts. The other end is equally limited in the amount of
carbonates that can remain dissolved without precipitating. Excesses can be
removed from the reaction at either end and stored, automatically.

We see the practical application of this in the majority of freshwater
systems (I haven't looked into the marine side in enough detail to make the
same statement there). We see specific use made of it in natural processes
like biogenic decalcification. We even plot out its predictable patterns in
things like the pH/KH charts that aquatic gardeners believe holy. But for
some reason, the idea just must seem too simplistic for everyday uses such
as water changes. We've just "gotta" be able to walk that fine edge,
balancing on the knife again by creating chemical soups rather than
normalized, naturalized environments.

Phosphates aren't necessary, even were they desirable. Dual-buffer
tug-of-wars can, as you've stated, lead to anal-retentive behavior in the
pursuit of chemical "nirvana". "Bull'sEye", ProperpH and similar products
exist for the same reason diet pills do - we want an answer in a bottle.

You already have all of the control you need, as well as the only investment
necessary, in the higher-quality DI unit you're using. Especially since
you're trying for a degree of hardness less than your source water. The
"pure" water coming from the DI means that you can pinpoint a specific pH or
KH for the maintenance tanks with just some simple math. Hardness values are
reduced in direct proportion to the amount of pure water added to the mix.
The 2:1 ratio you were using would cut your starting values to 1/3, meaning
your GH of 190 ppm (10.6 DH) and KH of 130 ppm (7.3) would be cut to 63
ppm (3.5) and 43 ppm (2.4) respectively. These values are certainly very
close to natural - perhaps a touch extra buffering. It would go further in
explaining the vibrancy of the fishes' coloring than would a slightly lower
pH value.

I have to throw my hat into the "don't mess with it further" ring, if for
nothing else than the nature of the resultant water.

Should you find it necessary to alter the chemistry a little more toward the
"black water" end of things, say for the breeding period, then Dave Gomberg
is correct - use the acid of choice when Nature *does* decide to play
chemist. Black water systems that use humic acids for buffering don't
compete with carbonate buffers. They virtually replace them as the
predominate buffering agent through their dissociative free hydrogen ions,
in effect exhausting the "alkalinity" to gain control. The easiest, most
efficient method for us to introduce these humics is, of course, peat
filtration. The control here is gained through the amount of peat exposed to
the water, controlling the amount of hydrogen introduced. Since the peat
also provides cationic exchange services for the divalent hardness cations,
it performs "time-and-a-half" for us in other areas, too. Again going with
his suggestion, start with small quantities and work your way up until you
reach your intended goal.

I will add this to what has got to be a much simpler way of "experimenting".
When working with peat, especially when trying to determine the proper
amounts, its effectiveness tapers off drastically after 48 hours' exposure.
In other words, the difference between 2 days and 2 weeks is not of any
significance, so you can safely draw your conclusions on the second day.


David A. Youngker

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