Determining CO2 Concentrations in Natural Waters

["David A. Youngker" <nestor10@mindspring.com>]

> From: David Sanchez
> Sent: Monday, January 28, 2002 3:56 PM

> I think you just don!'t understand the relationship
> between pH and carbonic acid...

David -

I'm afraid that it is you that has confused a few issues. Let's see if we
can walk through this in an academic manner and arrive at some sort of
"consensus", shall we?

> KH and pH both determine the level of CO2 in a
> body of water. This comes from a basic
> understanding of the pH equation.

There is no "pH equation", per se - only the _definition_ of pH, which is an
inverse log of the actual *count* of hydrogen ions in a solution. It's a log
scale because we're dealing with _huge_ numbers here, and the numbers we
_do_ use are in effect the *exponent* of the value.

The level of free hydrogen in the solution is totally dependent upon what
the various solutes _release_ into the solvent. Since the most commonly
found buffer in natural systems, from stream  beds to bloodstreams, is the
bicarbonate ion, we usually have to look no further than the dissociation
constant of that ion to see how many hydrogen ions it will provide.

The formula you're probably thinking of is the bicarbonate equilibrium
described by the commonly-referenced

H2CO3 <=> H+ + HCO3-

which shows the products and reactants in relation to the equilibrium point.
The actual point itself, determined through about 150 years' of empirical
data collected by analytical chemists, is described as a ratio of products
to reactants in a comparison of their creation rates. When there are as many
hydrogen ions and bicarbonate ions being produced as there are being
consumed, we have achieved "equilibrium" around the value

([H+][HCO3-])/ [H2CO3] = 4.3 x 10-7

Therefore, to determine the pH of the solution involved, we solve this for
the hydrogen concentration as

[H+] = ([H2CO3][4.3 x 10-7]) / [HCO3-]

which ties the hydrogen content to the ratio of carbonic acid- to-
bicarbonate ions.

This gives us a starting point for the rest of the conversation.

> This whole buffer solution system is so misunderstood
> by many hobbyists in my opinion. Why will pH fluctuate,
> just because we have a KH of 0 do we assume the pH will
> simply bounce all over the place?

If you're using bicarbonates as the predominant buffer, then that puts KH,
carbonate hardness, or alkalinity (or whatever the popular hobby term is
presently) right square in the denominator of the above ratio. What happens
when you try to divide by zero? (By the way, since the charts on Erik's site
is based on the Hendersen-Hasselbach equilibrium you'll find that same
"divide by zero" problem as the reason the charts break down at the
extremes. A zero in the denominator also represents reaction *completion*.)

> The relationship of KH to pH, is not if we have 0 KH
> then we no longer have a !'buffer!( this is a incorrect
> way of looking at this...

It is _not_ the incorrect way of looking at this - it _is_ the *only* way to
look at this. The "KH" is THE primary natural buffer within the equation. So
if you have no buffering agent, you can have no buffering. There _are_ other
buffers at work in our Apisto tanks down at the ranges we work on, and their
related to the humics most of us introduce through peat/tannin filtering

> A high KH will help maintain a Higher pH and vice
> versa a low KH (for example a 0 KH)will help maintain
> a lower pH...

A buffering agent is one that absorbs or lessens a stressor applied to the
system. In this case, the stressor is the concentration of hydrogen, which
in our tanks is usually on the rise thanks to the processes of
nitrification.

Buffers are established with a combination of an acid and its salt or a base
and its salt. In this case, the carbonic is the acid and the bicarbonates
are its base salt. The difference between the two is the added hydrogen,
which turns HCO3- into H2CO3. It is the "change of state" in capturing or
releasing the hydrogen that provides the buffering. If there is an influx of
hydrogen through, say, acids, then the product side of the equation relieves
the added pressure by combining the hydrogen with a bicarbonate to produce
carbonic. If we destroy the hydrogen through, say, the addition of a
hydroxide base, then the pressure is on the reactant side to produce more
hydrogen and re-establish a balance, and it does this by splitting more
carbonic molecules.

> ...I am not advocating adding CO2 to a system with
> sufficient CO2 present as can be found in a system
> with a low KH and pH...

If you consider the mechanics of infinite sources in contact with finite
bodies, then you'll find that our tanks tend to equalize their carbon
dioxide content to match that of the surrounding atmosphere. That leaves
(unaccounting for biological activity) the CO2 concentration as essentially
fixed. In the solution to pH presented above, since the CO2 content is
multiplied by the equilibrium constant, the entire denominator of the
equilibrium equation becomes fixed.

Did you catch that "fixed" part? That means it doesn't change. Therefore,
the only thing really influencing the final pH is the bicarbonate
concentration.

Speaking of concentration, mine's beginning to fade right about now - been
_very_ sick the last few days (sorry I was off-line, Teresa). And as this is
rather detailed, I'll pause here momentarily to entertain any questions up
to this point. We'll continue once this is established.

But be forewarned - my reading includes people like Arrhenius, Bronsted,
Lewis, Dalton, Henry, LeChatlier, Boyle, Charles and their ilk. And my
practical experience is approaching the fourth decade...

-Y-

David A. Youngker
nestor10@mindspring.com



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