[ RadSafe ] LSC Method Assistance Needed - CO2

Fri Jan 8 16:16:01 CST 2010

Dear Enoch:

There are two questions here: first is the efficiency of reacting the CO2
gas bubbling through an aqueous solution, and then once reacted, the ability
of the solution to retain the carbon. Gas bubbling is fairly efficient, but
if the gas is recycled through a high pH solution several times, virtually
all of the CO2 would be reacted to form CO3-2.

There is also a third implicit question: the concentration of CO2 in the
person's breath.  If a patient is exercising hard, I would imagine that the
concentration of CO2 in their breath would be elevated compared to that of a
resting person.  Also, assuming that the patient is not suffering from
metabolic or respiratory acidosis, and the blood pH was normal, then the
exhaled concentration of CO2 in the breath would be relatively constant.
But that's a little out of my territory as a geologist...

CO2 behaves like other gases. When a constant pressure of CO2 is exposed to
water, a portion of the CO2 dissolves in the water according to Henry's law:

CO2(g) <=> CO2(a)       kH=29.76 atm/(mol/L) at 25°C

A similar reaction occurs with nitrogen and oxygen, however, these two gases
do not chemically react with the water, CO2 does react to form a weak
diprotic acid; 

H2CO3 <=> H2O + CO2(a)     

The first ionization of that acid is a follows:

H2CO3 <=> HCO3- + H+    pK1 = 6.37

The second ionization has the following reaction:

HCO3- <=> CO3-2 + H+    pK2 = 10.25

The "other" important reaction is:

H2O <=> H+ + OH-

With distilled water, virtually all of the CO2 that is reacted causes the
water pH to drop.  

At atmospheric pressure, the concentration of CO2 is 10^-3.5 atmospheres. 
At equilibrium, the endpoint pH of the solution is 5.6 or so.  Thus all
rainwater is slightly acidic. For a person, the exhaled partial pressure of
CO2 is higher, so the solution would be even more acid if it were in
equilibrium. For distilled water in these pH ranges, virtually none of the
bicarbonate (HCO3-) dissociates to CO3-2, so the dominating equilibrium
reactions that can occur are:

H2CO3 <=> HCO3- + H+ ;
H2CO3 <=> H2O + CO2(a) ; and 
CO2(a) <=> CO2(g) 

Which can and will lose  CO2(g) on exposure to the ambient air concentration
of CO2 until equilibrium is reached.

So if a distilled water solution were to be reacted with a volume of air
from a person's breath, and then exposed to a lower concentration of CO2
from the atmosphere, the net result would be loss of CO2 from the solution
until it reached equilibrium conditions.  

At a pH above about 8.4, most of the carbon would be in the phases: 
HCO3- <=> CO3-2 + H+ 

And very little would exist in the equilibrium reaction:
H2CO3 <=> HCO3- + H+  as H2CO3 phase. 

Thus the sample would lose very little carbon as CO2.

Now, if the aqueous solution is kept very alkaline with the addition of
sufficient NaOH to always maintain the pH above say 11, by Le Chatelier's
principle, more protons would be released and neutralized, and virtually all
of the CO2 that reacted with the solution would be forced to form carbonate
(CO3-2) with virtually no HCO3- present.  Looking at it another way, all of
the protons generated by the reaction would be neutralized by the OH- from
the NaOH. So the dominant reactions would be:

HCO3- <=> CO3-2 + H+  and H2O <=> H+ + OH-

with the bulk of the original carbon in the CO3-2 phase.  Thus, no loss of
CO2.  Some, perhaps most of the CO3 will probably react to form an aqueous
neutral species, Na2CO3.  There will also be contributions from NaCO3- etc.
I don't think that an end-product - Na2CO3(a)- is as important as the fact
that once the CO2 is reacted to form CO3-2, regardless of the carbonate
aqueous speciation, Na2CO3(a), NaCO3-, CO3-2, etc., the total carbon is
conserved and the carbonate stays in solution.

I would imagine that whatever "cocktail" is prepared, the pH would have to
be adjusted to be 9 or greater. One issue with the high pH cocktail would be
sorption of ambient atmospheric CO2 and ambient C-14.  If this concentration
is far below that of the sample from the patient (I assume) then the
addition could be could be dismissed.

Sorry to be so long-winded!  And I still didn't pullout my computer codes to
model the reaction... Yippee!

Dan ii

--
Dan W McCarn, Geologist
2867 A Fuego Sagrado
Santa Fe, NM 87505
+1-505-310-3922 (Mobile – New Mexico)
HotGreenChile at gmail.com (Private email)

-----Original Message-----
From: Tung, Enoch [mailto:Enoch.Tung at PERKINELMER.COM] 
Sent: Friday, January 08, 2010 08:55
To: Dan W McCarn; Dale Boyce; radsafe at radlab.nl
Subject: RE: [ RadSafe ] LSC Method Assistance Needed - CO2

Thanks Dale and Dan.  I think you two were able to accurately gauge what
my ambiguous question was asking.
My question via Barb/Radsafe probably wasn't as clear as it should have
been.

The elixir that we use comprise of scintillation fluid and other
constituents, which adds up to the chemicals I/Barb listed.

I should have asked how other people capture CO2, and then what they
used to count the samples.

Considerations for it would include efficiency of the CO2 capture, and
type of scintillation fluid/chemicals used to minimize any collateral
effects (quench, chemiluminesce, etc) that the trapped CO2 solution
would have on the counting.

How efficient is 2NaOH + CO2 = > Na2CO3 reaction?  A chemist told me it
would be pretty efficient (assuming adequate flow rate/surface to area
ratio). Does the reaction occur immediately on contact between the
molecules?

Thanks,
Enoch Tung

-----Original Message-----
From: radsafe-bounces at radlab.nl [mailto:radsafe-bounces at radlab.nl] On
Behalf Of Dan W McCarn
Sent: Friday, January 08, 2010 12:11 AM
To: 'Dale Boyce'; radsafe at radlab.nl; blreider at aol.com
Subject: RE: [ RadSafe ] LSC Method Assistance Needed - CO2

Dear Dale:

I was about to step into this discussions with exactly the points that
you
raised regarding the mineralization of CO2 to carbonate in a high pH
environment.  I'm glad that you explained the reason for maintaining a
high
pH environment in order to prevent loss of CO2.  This saves me from
running
my mineral-equilibria codes to demonstrate the reason!

Dan ii

--
Dan W McCarn, Geologist
2867 A Fuego Sagrado
Santa Fe, NM 87505
+1-505-310-3922 (Mobile - New Mexico)
HotGreenChile at gmail.com (Private email)
-----Original Message-----
From: radsafe-bounces at radlab.nl [mailto:radsafe-bounces at radlab.nl] On
Behalf
Of Dale Boyce
Sent: Thursday, January 07, 2010 21:16
To: radsafe at radlab.nl; blreider at aol.com
Subject: Re: [ RadSafe ] LSC Method Assistance Needed - CO2

One comment that I haven't seen addressed. The NaOH in the original
cocktail

will capture CO2 as carbonate. Solutions with pH on the basic side tend
to 
cause many, if not all cocktails to chemoluminesce.

If I were starting from scratch, I would look for cocktails with a high 
aqueous capacity, as you may need to wait awhile before counting (keep
the 
samples in the dark as well). Setting a lower level discriminator above
the 
chemoluminescence will help assuming you have control over this.

In developing the procedure, I would recommend counting individual
samples 
as a function of time. Prepare any standards from a C-14 carbonate
solution.

You don't want to calibrate with an organic C-14 standard, as the
organic 
will follow the organic phase during organic aqueous phase separation,
while

your C-14 of interest in your sample will follow the aqueous phase.

Do not acidify either your samples or standards at any point. You will
lose 
CO2, and not be able to rely on your results.

Make quench standards specific to your sample and cocktail formulation. 
Maybe a little overkill, but you are developing a new procedure for a 
technique that is on the edge of "normal" lsc usage. Also, look in the 
literature for C-14 carbon dating using LSC. It has been used, and there

should be some papers out there. Typically these were also CO2 samples.

Have fun, and possibly get a paper out of the effort!

Dale 

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