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


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Contents:
  Re: Mike Maag pump problem (David Towson)
  hardness ("Darrell G. Leavitt")
  March pump (was heat sticks) ("Mike Maag")
  Re: water (M Lewandowski)
  Water ("A. J. deLange")


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Date: Tue, 16 Feb 2010 09:32:27 -0500
From: David Towson <davidtowson at verizon.net>
Subject: Re: Mike Maag pump problem

I haven't been following this thread, but based on the description 
given in HBD 5658, I have to question whether the pump rotor is 
actually turning.  It sounds as though the rotor is seized-up, and 
all you're getting is gravity flow when the outlet hose is low 
enough. If this is the case, the motor can still run, but the 
magnetic drive will slip. I suggest taking the thing apart and 
checking for a stuck or broken rotor. Also check that the rotor 
magnet is still attached to the rotor.

Dave



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Date: Tue, 16 Feb 2010 09:52:20 -0500 (EST)
From: "Darrell G. Leavitt" <leavitdg at plattsburgh.edu>
Subject: hardness

Ok, let's do talk.

I am not an expert here, but think that hardness has to do with the amt of
minerals in one's water (Calcium, and Magnesium)  My water analysis says
that Total Hardness ( as CaC03/ Calcium Carbonate) is 178.  And, I believe
that this is close to Munich water, and that this means that I can more
easily brew red ales, and darker ales, but have to use some distilled
water, or pre-boil some of my water, so as to get rid of some of the
temporary hardness.

This is a start, no?

Darrell



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Date: Tue, 16 Feb 2010 12:14:32 -0500
From: "Mike Maag" <mikemaag at comcast.net>
Subject: March pump (was heat sticks)

I raised the mounting point (still below the source) and swiveled the pump 
housing to vertical, outlet up.  Works like a champ.  All the air gets out 
of the pump and lines just fine.  Now I should be able to get my hot sparge 
water up into the "cooler" before it cools too much.  Thanks to all for the 
input, and thanks to the archives for additional pump mounting tips.

Cheers!

Mike Maag
Staunton, VA  (Shenandoah Valley) 



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Date: Tue, 16 Feb 2010 15:26:42 -0500
From: M Lewandowski <m-lew at comcast.net>
Subject: Re: water

 From a theoretical water chemistry standpoint, hardness is the 
concentration of polyvalent cations in the water.  In English, this means 
all of the positive ions with a charge greater than one.  From a practical 
standpoint, the most common contributors to hardness (by-far) are calcium 
and magnesium.

Alkalinity measures how resistant a water is to a pH change.  It takes much 
more acid to lower the pH of a high alkalinity water.  Basically, high 
alkalinity buffers the system against rapid pH changes.  The most common 
source of alkalinity is the carbonate system.  At normal drinking water pH, 
the bicarbonate ion (HCO3-) is the most common source of alkalinity.

Here's why you may be confused about the two parameters.  Hardness is due 
to dissolved positive ions.  Alkalinity is due to dissolved negative 
ions.  In any natural water, the positive ions should balance the negative 
ions.  However, there's the measurement of these ions is olny so good, so 
the numbers don't always match "exactly" but they should be pretty close.

There's more than one way to measure hardness.  In a lab, you can use an 
atomic absorption spectrophotometer to measure each positive ion.  Add the 
totals and you have a very exact measurement of hardness.  Many companies 
also sell a kit, where you add some reagent and measure how much of another 
chemical you have to add to make the color change.

When I was in school, we measured alkalinity by titration.  You measured 
how much acid it took to change the pH to a certain level.

Here's some trivia for you.  Some soaps have a very hard time forming a 
lather in water with high concentrations of calcium and magnesium.  That's 
where the term "hardness" came from; it's hard to wash with water like that.

I hope this helps.  If you want more information, feel free to ask 
follow-up questions.


Mike L.



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Date: Tue, 16 Feb 2010 15:28:06 -0500
From: "A. J. deLange" <ajdel at cox.net>
Subject: Water

Haven't had a question like this here in years!

Hardness is defined as the concentration of Calcium and Magnesium (only 
 - this is important because some tests will also respond to strontium, 
iron and so on which are specifically excluded from the definition by 
Standard Methods for the Examination of Water and Wastewater which is 
incorporated by reference into the ASBC's MOA's). Separate hardness for 
Calcium and Magnesium is defined as is a total hardness which is the sum 
of the two. If they are to be added they must be in the same units and 
that is either milliequivalents per liter (sometimes called vals) or 50 
times milliequivalents per liter which is called "parts per million as 
calcium carbonate" for reasons which will become apparent in a moment. 
Hardness is measured by adding an indicator dye (such as Eriochrome 
Black-T) which is bright red in the presence of calcium or magnesium 
ions and bright blue in their absence to a specific volume of sample 
(commonly 100 mL). To measure total hardness the sample is buffered to 
an appropriately high (but not too high) pH, the dye added and a 
titration with a chelating agent that grabs both calcium and magnesium 
such as EDTA (but which also chelates strontium and other metals and 
thus these can fool the test) carried out until the solution turn blue. 
The chelating agent is calibrated in milliquivalents it can chelate per 
unit volume and thus the total hardness depends on the volume of the 
sample, the strength of the titrant and the volume of the titrant 
required to effect the color change.

If it is desired to know the calcium and magnesium harndnesses 
separately several approaches exist. The most common one is to buffer 
the solution to a pH higher than that used for the total test. This 
causes the magnesium to precipitate as the hydroxide so that a 
subsequent titration with EDTA responds only to the calcium. The value 
so determined is the calcium hardness which can be subtracted from the 
total hardness to give the magnesium hardness. Or a chelating agent 
which takes out only calcium (EGTA) can be added in excess. A subsequent 
titration with EDTA is then responding only to the magnesium hardness 
which can be subtracted from the total hardness to give the calcium 
hardness.

If separate values for magnesium and calcium hardness are available 
these can be converted to concentrations of the ions in mg/L. If the 
hardness numbers are in ppm as CaCO3 divide by 50 to get mEq/L and then 
multiply the calcium hardness mEq/L by 20 to get mg/L calcium and the 
magnesium hardness by 12.15 to get mg/L magnesium.

Calcium and magnesium hardness can be determined by atomic absorbtion 
spectroscopy. The sample is sprayed into a flame through which a beam 
from a calcium vapor lamp is shone. The amount of light absorbed by the 
flame is proportional to the amount of calcium in the sample. Note that 
strontium, magnesium etc have no effect because the light is of the 
frequency which excites only calcium. The same is done for magnesium 
using a magnesium vapor lamp. mg/L so determined can be divided by the 
equivalent weights (20 for calcium, 12.15 for mangnesium) to give the 
hardnesses in mEq/L which can be multiplied by 50 tto give ppm as CaCO3 
and so be added for total hardness in either units.

Alkalinity is defined as the amount of acid (expressed in mEq) which 
must be added to a liter of sample to lower its pH to an arbitrary pH 
which is an important part of the definition but opinions vary as to 
what that pH should be. There are guideline values which can be found in 
various sources which depend on the amount of bicarbonate in the water 
(but you need to know the alkalinity first to figure that out). Another 
approach is to bring the sample to the pH where the concentration of 
hydrogen ions equals that of bicarbonate ions (the "equivalence end 
point") and others (including me) just use 4.3. Standard Methods says 
that any pH can be used as long as the report says what it is. Few 
laboratories put the end point pH on their reports (I've never seen it).

Practical details: 100 mL of sample is placed in a flask with an 
indicator which turns color at around pH 4.3 (methyl orange turns at 4.3 
and was the standard for a long time to the point where the alkalinity 
so determined is called the M alkalinity). Or, if you are color blind as 
is the writer, a pH electrode is inserted into the sample. The sample is 
then titrated with 0.1N acid (hydrochloric or sulfuric usually) until 
the end point pH is reached. The number mL of titrant is the number of 
mEq/L required to move the sample to end point and so the mL reading 
from the buret is the alkalinity in mEq/L. Sample size and/or strength 
of titrant can be varied for high or low alkalinity samples. Alkalinty 
in mEq/L can be multiplied by 50 to give alkalinity in ppm as CaCO3 and 
thus be in the same units as hardness and the obvious reason for wanting 
to do this is so that hardness and alkalinity can be directly compared. 
If 100 mg CaCO3 is placed in a beaker through which CO2 is bubbled until 
it is dissolved (this is how limestone is dissolved underground) and the 
pH reaches 8.3 the hardness of that treated sample will be 100 ppm as 
CaCO3 and its alkalinity will also be close to 100 ppm as CaCO3. 100 mg 
of CaCO3 is 1 mMol  (molecular weight of CaCO3 is 100) which means  1 
mMol of Ca++ which is 2 mEq/L. Multiply that  by 50 and you have 100. 1 
mMol CaCO3 contains 1 mMol of carbonate which, reacted with 1 mMol of 
carbonic acid would give 2 mMol of bicarbonate which is 2 mEq because 
bicarbonate carries a single charge. Multiply by 50 again and again you 
get 100 for the alkalinity. Thats where the times 50 thing comes from.


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