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RE: Carbon Dioxide Concentrations in Water

Date: Sun, 16 Dec 2001 16:54:46 -0000
From: George Booth
Subject: Re: ph-KH-CO2 (Vol4 #1474)

> If I remember right, the theoretical equilibrium
> value of CO2 in water at a specific temperature,
> sea level and 350 ppm in the air is 0.498 mg/l.
> This varies with the three conditions, so your
> results WILL vary.

Date: Mon, 17 Dec 2001 10:31:44 -0500 (EST)
From: Paul Sears
Subject: Re: CO2 concentrations in water (Vol4 #1475)

> The equilibrium concentration of CO2 in water in
> contact with outdoor air at about 23C is, as various
> people have pointed out, about 0.5 ppm.  The outdoor
> air CO2 concentration is about 350 ppm.  _Indoor_
> air CO2 concentrations are rather higher.  I think
> that below 1000 ppm CO2 is recommended, but it
> isn't at all uncommon to see it higher.  That could
> easily explain 2 ppm of CO2 in water.


If we make the "standard" assumption that carbonic acid and hydrated carbon
dioxide are essentially the same beast...

and use a solubility constant of Kco2 = 3.38x10-2...

and move the atmospheric CO2 concentration up to the current 369 ppm (which
would give us a CO2 percentage of about 0.000369)...

and consider that, since we are below 5 atmospheres' pressure, Henry's Law
holds true...

then the naturally- occurring concentration of CO2 in the water should be

[CO2(aq)] = Kco2*Pco2
          = 3.38x10-2 M/l Atm * 0.000369 Atm
          = 1.25x10-5 M/l

and at 44.01 g/M, then by weight we have

44.01 g/M x 0.0000125 M/l = 0.000550 g/l

or 0.55 ppm.

Since I based my prior statement on math I'd performed a couple of years
ago, I thought in light of Paul's posting (particularly) I'd better go back
and try the numbers again. I now see I've been a whole decimal place off
(don't ask me - I've "slept a few times" since then - or perhaps was
dreaming of a job at JPL at the time...)

Up until a few minutes ago, I was perfectly willing to take up Steve's cause
and argue the point, but I don't really have a leg to stand on now ;-)
Sorry, Steve - but I _was_ game there for a moment.

Date: Mon, 17 Dec 2001 09:47:29 -0500
From: James Folsom
Subject: Re: Aquatic Plants Digest V4 #1474

> ...you must understand that distilled water has
> no buffer capacity whatsoever.  A ph drop of 0.8
> in distilled water is caused by much less CO2
> than what it would take to lower the pH of tapwater.
> Either that or CO2 is more soluble in distilled water.

The same amount of CO2 added to the water should create the same pH drop no
matter the starting point, but only if you stay within the carbonate family.
Adding another acid would kick in the buffering effect of the bicarbonates
and cause a different shift in equilibria points through removal by
consumption of one of the defining components.

> Simple experiment:  Put your distilled water,
> and your tapwater under aeration and measure the
> pH.  The deionized water in my lab usually rings
> in at a pH of 4.5 wheras the tapwater is about 7.1.

Ah, but wouldn't this come a lot closer to explaining the difference in
bicarbonate concentations than absorption? And if both had the same
concentrations, whether zero or a particular amount, wouldn't their pH
values be the same within the same room? (BTW, your tap sounds like it's at
about the same as mine here in Tennessee - see
http://www.tawc.com/ourwaterqual/data.html for most of the pertinent

Date: Mon, 17 Dec 2001 14:23:25 -0800 (PST)
From: Scott Hieber
Subject: Re: CO2 and it's sources (Vol4 #1477)

>> Hmm...would that be what's known as "equilibrium"?...

> Yes, once the states of the water and air are stable.

I apologize for that, Scott - I was being facetious at the time. Although
the site's seriously in need of an update (and a couple minor corrections),
if you'll check http://www.mindspring.com/~nestor10/bcrb-fE4.htm (among
others), you'll find I'm already well- versed in the necessary chemistry...


David A. Youngker
nestor10 at mindspring_com