Re: Determining CO2 Concentrations in Natural Waters

["Mike Jacobs" <mjacobs2@tampabay.rr.com>]

.......................Youngker's in my canoe.................that was the
damnest explanation I've ever seen.  I've read it 4 times and will have to
do it 2-3 more times and talk to our chemistry teacher at school.  I'm not
picking sides to cause trouble..............really I'm not.............!!!!!
And all I wanted to do was grow some foxtail.............just add another
fluorescent tube Mike..........heh, heh, heh.........wow!  Somewhere in
there I should be able to use differentiation or integration to grow that
foxtail.................remember in there about trying to divide by
zero...................I followed that.....;-) ;-) ;-).....you know the
definition of a "black hole"??????..................................it's
where God tried to divide by zero......................

...............I think we all ought to go to Java moss................;-)
;-) ;-)

....hey Big John........I didn't see anything about "blue goop" in there.
did you!!

Mike

Mike Jacobs
Center for Advanced Technologies
High School Math Instructor
St. Pete, Fl.
----- Original Message -----
From: "David A. Youngker" <nestor10@mindspring.com>
To: "Apistogramma Mailing List" <apisto@listbox.com>
Sent: Monday, January 28, 2002 6:42 PM
Subject: Determining CO2 Concentrations in Natural Waters


> > 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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