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	FORUM ON BEER, HOMEBREWING, AND RELATED ISSUES
		Digest Janitor: pbabcock at hbd.org


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Contents:
  Excess CaCO3 ("A.J deLange")


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Date: Fri, 1 May 2009 13:55:32 -0400
From: "A.J deLange" <ajdel at cox.net>
Subject: Excess CaCO3

Matt posted yesterday about storing brett over an excess of CaCO3 and  
wondering what the pH of the solution might be. This is a problem that  
can be handled by the Nearly Universal Brewing Water Spreadsheet (you  
knew I was going to say that, didn't you?). What may be surprising to  
most is that the independent variable here and the thing that sets the  
pH is the partial pressure of CO2 over the container. I'm going to  
describe briefly how the answers to be given were obtained from NUWBS  
(obtainable at www.wetnewf.org). If you want to try this the remarks  
should suffice to get you started. Otherwise just look at the answers.

Start by setting up the NUBWS for deionized source water (carbo mode  
"C", value 0, pH 7.0, temperature 20 C) and zero all inputs to the  
Synthesis section. Now ask the Solver to set proton deficit to 0 while  
varying added CaCO3, CO2 and pH subject to the constraints that the  
pressure of CO2 is 0.0003 atmospheres and that the saturation pH  
equals the actual pH. Solver will find pH 8.33.  This is the pH DI  
water over chalk will come to if it is allowed to stand in air at  a  
partial pressure of 0.0003  but it may take days for this to be  
reached. IOW we can calculate the equilibrium pH from thermodynamics  
but can't say how long it may take to reach it. NUBWS will also  
calculate that 55 mg/L of CaCO3 will be dissolved.

Should Al Gore be right and the partial pressure of CO2 double to  
0.0006 atm the equilibrium pH would drop to 8.13 and believers worry  
about, for example, what the effects on marine animals might be  
(generate more shell material and send it to the bottom or die - shell  
material is more soluble - and really louse up the system).

Now when we add the brett and they start to chug away and produce acid  
the situation isn't quite so simple. Evolved CO2 will accumulate in  
the container so that the partial pressure of that gas over the water  
will be higher.  Lets assume that you do this in an open beaker with a  
fan blowing over it so that evolved CO2 is swept away and the partial  
pressure of CO2 is again 0.0003 atm. Reset the CO2 pressure constraint  
to 0.0003 atmospheres in NUBWS, choose citric as the 'other' acid and  
enter 10 mg/L of it to simulate the acid production of the yeast. Now  
have the Solver do its thing again. It should calculate a pH of 8.31  
and show 61 mg/L chalk dissolved. For 20 mg/L citric produced the pH  
will drop to 8.29  and 67 mg/L chalk will dissolve. Thus it appears  
that the extra acid produced by the yeast dissolves a bit more chalk  
than the CO2 from the air does by itself.

Now suppose that you had the yeast and chalk in a flask with an  
airlock. Eventually the CO2 would fill the flask to the point where  
the partial pressure of CO2 over the liquid was 1 atm. Put that in for  
the CO2 pressure constraint ad the Solver will find a pH of 5.83 for  
20 mg/L citrate with 572 mg/L chalk dissolved. With 20 mg/L citrate the
20 mg citrate the pH changes by less than 0.01 unit (i.e. still 5.83  
to 2 decimal places) and the dissolved chalk is up to 576 mg/L.

Thus you can conclude that the acid produced by the yeast has  
relatively little effect on the pH relative to the exposure of the  
broth to CO2 with whether or not you allow the evolved CO2 to  
accumulate or not having a major effect.

Remember again that these calculations are based on thermodynamics and  
may not reflect what you observe unless you take steps to see to it  
that equilibrium conditions are being approached. Shaking the flask to  
liberate dissolved CO2 would be an example of one action that will  
move the system towards equilibrium. Sweeping away evolved CO2 with a  
fan as mentioned above would be another.

A.J.


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