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HOMEBREW Digest #5493
HOMEBREW Digest #5493 Mon 02 February 2009
FORUM ON BEER, HOMEBREWING, AND RELATED ISSUES
Digest Janitor: pbabcock at hbd.org
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Contents:
Re: Aluminum kettle (John & Joy Vaughn)
Osiris HBC Results posted (Dion Hollenbeck)
Diatomacious Earth for Bug control (slaycock)
How Much Alkalinity ("A. J. deLange")
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Date: Sun, 01 Feb 2009 22:29:22 -0900
From: John & Joy Vaughn <hogbrew at mtaonline.net>
Subject: Re: Aluminum kettle
Fred,
As long as he has removed all of the oil from the surface of the kettle,
it will work great as a boil kettle.
John Vaughn
Wasilla. Alaska
> Date: Sat, 31 Jan 2009 18:55:12 +0000 (UTC)
> From: Fred Scheer <fredscheer07 at comcast.net>
> Subject: Aluminum kettle
>
> HI:
>
> A friend of mine got some time ago one of
> this Turkey Fryers, 30 qt, Aluminum.
> Now, he made fish fry in there three times.
> He cleaned it very well, I can't see any
> spots in the kettle.
> I have no experience in homebrewing in Aluminum
> kettles. Any advise is appreciated?
>
> Cheers,
>
> Fred
>
>
>
>
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Date: Mon, 02 Feb 2009 09:30:16 -0700
From: Dion Hollenbeck <hollen at woodsprite.com>
Subject: Osiris HBC Results posted
For the complete results to the Osiris Homebrew Competition sponsored
by SKA Brewing in Durango, CO., please go to:
http://hstrial-cmorrow8.homestead.com/
and click on the "Awards" link.
Congratulations to Chris Vest who won BOS and will have his beer
brewed and entered by SKA in the GABF Pro/AM Competition this year.
dion
- --
Dion Hollenbeck
Email: hollen at woodsprite.com Home Page: http://www.woodsprite.com
Brewing Page: http://hbd.org/hollen Toys: 98 4Runner, 86 4x4 PU
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Date: Mon, 2 Feb 2009 11:36:31 -0600 (CST)
From: slaycock at discoverynet.com
Subject: Diatomacious Earth for Bug control
Greetings,
I have been researching the uses of Diatomacious Earth lately for personal
health benefits.
Already knowing that it is amazing as a natural "food grade" bug killer.
(I hesitate calling it an insecticide, because it isn't a poison but the
microscopic diatoms sharp edges simply cuts the insects/parasites as they
land/walk on the DE or ingest it and they dehydrate or "bleed out" if you
will, and die. It's much like what would happen if we walked & rolled
around on sharp knife blades...not a pleasant thought!)
All that said, can I safely add this to my barley grain and even Bulk DME
stock for bug control WITHOUT harming the health of my yeast during
fermentation?
Keep in mind that DE used for Pool filtering IS NOT FOR HUMAN CONSUMPTION!!
If you use DE, it MUST BE MARKED "Food Grade or Codex Food Grade"
Thanks for any feedback!
Steve in KC
Highwater Brewhaus
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Date: Mon, 02 Feb 2009 18:59:59 +0000
From: "A. J. deLange" <ajdel at mac.com>
Subject: How Much Alkalinity
Kai and I have been discussing this off line. In summary for those who
may be interested:
It would be well to start with the definition of alkalinity: the
amount of acid required to move the pH of a water sample from whatever
it is to pH 4.3 (which is a somewhat arbitrary, in the sense that 4.2
or 4.4 would serve just as well, number but the one usually accepted
in brewing).
If one puts 100 mg (1 mMol) of chalk in a beaker and adds water to
make up a to a little less than a liter and then begins to add a
strong acid (i.e. one whose highest pK is at least 2 units less than
4.3) until pH 4.3 is reached he will have to add 2.05 mEq of acid.
The first mEq goes to convert the carbonate to bicarbonate: CO3-- + H
+ --> HCO3- and the second to convert all the bicarbonate to
carbonic (actually at pH 4.3 not quite all but certainly most will
convert): HCO3- + H+ --> H2CO3. The .05 is needed just to establish
the H+ ion concentration of 10^(-4.3) mol/L required for a pH of 4.3
so that even completely deionized water has an alkalinity of .05 mEq/
L or 2.5 ppm as CaCO3 when 4.3 is the defining pH. As 2.05 mEq of acid
total have been used in reaching pH 4.3 the alkalinity of this sample
is 2.05 mEq/L or 102.5 ppm as CaCO3 (multtiply by 50).
If one adds the acid incrementally the pH will decrease from a value
around 10 to 4.3 as the acid is added. Let r1 = 10^(pH-6.38) be the
ratio of bicarbonate to carbonic and r2 = 10^(pH - 10.38) be the ratio
of carbonate to bicarbonate ion concentration. Then f1 = 1/(1 + r1 +
r1*r2) is the fraction of total carb species which are carbonic, f2 =
r1*f1 is the fraction which is bicarbonate and f3 = r2*f2 = r1*r2*f1
is the fraction which is carbonate. The amount of acid consumed in
going to a particular pH from solid calcium carbonate is the increase
in carbonic plus the decrease in carbonate. If starting with chalk
(f1 = 0, f3 =1.0 and 100 mg = 1 mMol) so the amount of acid required
is (1)*(f1 + 1 - f3). Thus at pH 4.3 where f1 is approximately 1 and
f3 approximately 0 the acid required is approximately 2 mEq as above.
In this simple calculation we are ignoring the 0.05mEq for the pH
value itself. At pH 8.3, by contrast, f1 and f3 are each about 1% and
f2 about 98% so the acid required is approximately (1)(.01 + .99) or
about 1 mEq/L. To take this sample the rest of the way to pH 4.3 would
thus require another mEq (plus the .05) and the alkalinity of this
sample, prepared by putting 100 mg of CaCO3 into a bit less than 1L of
water, adding acid to pH 8.3 and adding water to make up to exactly 1L
would be 1.05 mEq or 52.5 ppm as CaCO3. If I added enough acid (1.19
mEq) to get the sample to pH 7 where 19% of CO3-- has converted to
carbonic and 81 to bicarbonate and you sent the sample off to a lab
they would come back with a report of about 2.05 - 1.19 = 0.86 mEq/L
or 43 ppm as CaCO3. To reach pH 6 would require 1.71 mEq of acid and
the alkalinity would then be about 0.34 mEq/L or 17 ppm as CaCO3. Thus
the alkalinity is about half the number of mg/L CaCO3 added at pH 8.3
and drops rapidly as the pH of the prepared sample is made lower.
Now things are different if a weak (at least one pK close to the
working values of pH of interest) acid, in particular carbonic, is
used to supply the protons. The reaction is
CO3-- + H2CO3 --> 2HCO3-
Here a carbonic molecule has given up a proton becoming a bicarbonate
ion. The proton is immediately taken up by a carbonate ion which
becomes a bicarbonate ion. At pH 8.3 again 98% of the added carbonate
has been converted to bicarbonate but now for each bicarbonate from
the added chalk there is a second from the acid which supplied the
proton to convert it. It should be clear that the alkalinity will, in
this case, be double what it was when strong acid was used and it is
at about 101.9. But at lower pH as more protons are required more
carbon dioxide must dissolve to supply them and this increases the
alkalinity above what it would be were strong acid (hydrochloric,
sulfuric) used. At pH 6 1.71 mEq of protons are still needed but as
only 29% of carbonic converts to bicarbonate one requires 5.82 mMol of
carbon dioxide tot be dissolved. As one bicarbonate is produced for
each proton there will thus be 1.71 mMol of bicarbonate from carbonic.
Of the 1 mMol of added carbonate 29% is converted to bicarbonate as
well and there is thus a total of 2.00 mMol of bicarb which will
require 2 mEq of acid to convert totally to carbonic (near pH 4.3) and
the alkalinity of this water is thus 2 mEq/L or about 100 ppm as
CaCO3. Thus if the CaCO3 is dissolved with carbonic acid, which is
nature's way (surface water, limestone caves etc.) The alkalinity will
be about 1 part per million "as CaCO3" for each mg/L CaCO3 dissolved
and it is relatively insensitive to pH as long as you stay above pH 6
(at pH 5 the alkalinity drops to 0.82ppm/mg/L and it takes 1.4
atmospheres of CO2 to get enough to dissolve). That's why we use "as
CaCO3 units (though I really prefer mVal/mEq because they don't lead
to the kind of confusion we often see over this very point).
The bottom line here is that if you are doing calculations concerning
calcium carbonate additions to brewing water you MUST consider the pH
of the water you are formulating and the means that will be used to
dissolve the calcium carbonate. Any spreadsheet, brewing software,
etc. that does not ask about both of these things has the potential to
lead you astray.
All the calculations alluded to above were done with the spreadsheet I
developed for a class last fall which is available at
http://www.wetnewf.org/Brewing_articles/BURP_OCT08
and which, as a consequence of my correspondence with Kai I have
updated to fix a couple of minor things, added fields which show
saturation WRT carbon dioxide and calcium carbonate (this can be an
issue - water in equilibrium with air and calcite can only dissolve 55
mg/L) and appreciably augmented the instructions. It is probably
pretty intimidating but it does cover most aspects of brewing water
chemistry.
A.J.
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End of HOMEBREW Digest #5493, 02/02/09
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