FORUM ON BEER, HOMEBREWING, AND RELATED ISSUES
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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
------------------------------
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
------------------------------
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)
------------------------------
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.
------------------------------
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.
------------------------------
End of HOMEBREW Digest #5659, 02/16/10
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