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PMDD
>
> Date: Fri, 04 Apr 1997 07:39:36 -0500
> From: Neil Frank <nfrank at mindspring_com>
> Subject: Re: chemicals and dosing
>
> Thanks Tim for the complete and interesting post. I have set it aside for
> study when I get the time. In the meanwhile, I thought I would share my
> draft writeup that is related to the general concept of dosing --
involving
> trace element mixes, commonly available chemicals and use of tools like
> measuring spoons. I thought people might find it useful (I hope the
columns
> are not messed up too badly and that people do not take offense to the
use
> of "American" units <g>)
>
> -
--------------------------------------------------------------------------
> Determining the Concentrations of Chemicals Added to the Aquarium
>
> Introduction
> Chemicals are often added to aquariums to provide food for plants or to
> establish other chemical parameters. Added nutrients may be macro
nutrients
> (e.g. potassium (K), magnesium (Mg), calcium (Ca)) or trace elements. In
> some cases, Nitrogen (N) may also be needed when the plant's requirements
> exceed the available supply or when plant growth is intentionally pushed.
> The chemicals may also be needed to correct nutritional deficiencies. In
> addition, bicarbonate levels in the aquarium water are often modified to
> achieve desired pH. With a pH controller, this will help establish
> particular CO2 concentrations in tanks with CO2 injection.
>
> Sometimes chemicals are provided through commercially prepared nutrient
> solutions, pH buffers or KH "builders." However, the aquarist may want
to
> create his or her special brews to replace or supplement the commercial
> products. Fine tuning may be useful in order to account for one's local
tap
> water and particular aquarium conditions. This discussion does not
> extensively focus on the specific circumstances indicating when and why
one
> adds these chemicals to the aquarium. This has been discussed elsewhere.
> Instead, the focus here is on determining the concentrations and
achieving
> target levels.
>
> Chemicals for the aquarium can be found as dry ingredients (crystals or
> powders), or in liquid form (solutions). Many useful chemicals can be
found
> in the grocery or local pharmacy. The dosing approach described here can
> involve a concentrated "stock" solution or dissolving the chemical in a
> small amount of water and then adding the potion directly to the
aquarium.
> Applications are included for adding macro nutrients (N, K, Mg, Ca),
> achieving target iron concentrations using trace element powders and
> increasing carbonate hardness (KH) or general hardness (GH). Avoidance of
> excess concentrations is discussed.
>
> Terminology and Computational Procedures
> Concentration units of chemicals in the aquarium are often expressed in
> parts per million (ppm), for example milligrams of nitrate per 1,000,000
mg
> of solution. (A liter (L) of water weighs 1000 grams (g) or 1,000,000 mg,
> and so one ppm is one milligram per liter). One ppm is also the same as 1
g
> (1,000 mg) per 1,000 L.
>
> For Americans and others who may not be comfortable with the metric
system
> or still think about their tanks in terms of U.S. gallons, concentrations
> can also be expressed in other units. Multiples of 10 gallons is a useful
> volume. For 1 gram of soluble material, the concentration in 10 gallons
of
> water can be determined by simple arithmetic. Because 10 gallons is 37.85
L,
> then 1 gram (g) / 10 gallons is 26.4 ppm. For 50 gallons, one gram
creates
> 5.3 ppm.
>
> Often we are not interested in achieving the concentration for a
particular
> chemical compound (like table salt, NaCl), but separately for its
chemical
> constituents or ions (i.e., the Na+ or Cl-). This is also true for the
> components of nutrient salts, like Sodium Nitrate (NaNO3) or mixtures of
> trace elements. In general, specific elements or ions represent a
fraction
> of the entire chemical. For example, Chloride represents 61 percent of
the
> weight of table salt.
>
> If the concentration of Fe in the trace element. mix is 7 percent, then
one
> gram of trace element powder added to 50 gallons of water is simply
(0.07) x
> (5.3) = 0.37 ppm.
>
> We can also work backwards to determine the amount of chemical needed to
> achieve a particular target concentration. With trace element powders,
iron
> is often used as the indicator variable and tells the aquarist if all
other
> trace elements are in proportion. For this purpose 0.1 ppm Fe is used.
>
> To determine the achieve 0.1 ppm Fe from the trace element powder, the
> number of needed grams, G, in our example volume of 50 gallons of water
is:
>
> (0.7) x (5.3) x G = 0.1, so, G = 0.27
>
> In other words, approximately 1/4 g of this trace element powder is
needed.
>
> The same principle can be used to determine the amount of any compound
(e.g.
> table salt) needed to achieve a desired concentration of any ion or
element.
>
> Example Chemical Concentrations
> The following table shows the percentage of various ions in different
> compounds and the resultant concentration for different constituents from
> dissolving 1 gram of the compound in 50 gallons of water.
> -
----------------------------------------------------------------------------
--
> Compound ion or percent (p) concentration (ppm)
> element of compound resulting from
> 1 g in 50 gal (189 L)
> -
----------------------------------------------------------------------------
--
> Sodium Nitrate nitrate 73 3.9
> (Nitrate of Soda) nitrate N 16 0.9
> NaNO3 sodium 27 1.4
>
> Ammonium Nitrate nitrate 78 4.1
> NH4NO3 nitrate N 17.5 0.9
> ammonium 22.5 1.2
>
> Ammonium Chloride ammonium 34 1.8
> NH4Cl
>
> Sodium Bicarbonate bicarbonate 73 3.9
> (Bicarbonate of Soda) sodium 27 1.4
> NaHCO3
>
> Potassium Chloride potassium 52 2.8
> (Muriate of Potash) chloride 48 2.5
> KCl
>
> Potassium Nitrate potassium 39 2.1
> KNO3 nitrate 61 3.2
>
> Calcium Carbonate calcium 40 2.1
> CaCO3 carbonate 60 3.2**
>
> Magnesium sulfate magnesium 10 0.5
> (Epsom salt) sulfate 39 2.1
> MgSO4*7H20 water 51 2.7
> -
--------------------------------------------------------------------------
> * concentration = 5.3p / 100
> **Note: the carbonate is converted into bicarbonate after the CaCO3
dissolves.
> -
--------------------------------------------------------------------------
>
>
> How to Measure a Particular Amount
> Not everyone has a gram scale, and sometimes it may be more convenient to
> prepare concentrated stock solutions and then add a portion to produce
the
> desired (diluted) concentration. On the other hand, precisely knowing the
> resultant concentrations is not critical and therefore some standard
> measuring devices (like fractional teaspoons) can be very useful to
> approximate these small weights. Sometimes, the chemical comes in nicely
> pre-packaged amounts, like Calcium tablets (Dietary supplement, pure
> calcium carbonate), but generally teaspoon measures are sufficient. I
have
> discovered that for most chemicals, 1/4 teaspoon (t) = 1 to 2 grams. Here
> are a few example concentrations resulting from 1/4 teaspoon of different
> compounds and from calcium tablets.
>
> -
----------------------------------------------------------------------------
--
> Approximate
> Compound Weight (g) Element Concentration (ppm)
> per 1/4 tsp. or ion of 1/4 t in 50 gallons
> - ----------------------------------------------------------------------
> sodium nitrate 1.8 NO3- 7.0
>
> sodium bicarbonate 1.3 HCO3- 5.1
>
> ammonium nitrate 1 NH4+ 1.2
> NO3- 4.1
>
> potassium chloride 1.5 K+ 4.2
>
> Potassium nitrate 1.4 K+ 2.9
> NO3- 4.5
>
> Magnesium sulfate 1.35 Mg++ 0.7
> (hydrate) SO4-- 2.8
> - -------------------------------------------------------------------
> Calcium carbonate Ca++ 3.2*
> * (600 mg Calcium tablet) CO3-- 4.8*
> Note: the carbonate is converted into bicarbonate after the CaCO3
dissolves.
> - -------------------------------------------------------------------
>
>
> Molarity of Solutions.
> Preparing a stock solution is another way to precisely provide the
> amounts of material to create concentrations in the aquarium. Stock
> solutions are often described in terms of molar concentration. This is
> determined from the atomic weight. For example, potassium has an atomic
> weight of 39.1 - one mole of potassium would weigh 39.1 grams and a 1
molar
> solution of KCl is 39.1 g K per liter.
>
> With solutions, milliliters (mL) are a standard unit of measurement.
These
> are found, for example, as the markings on the vials that come with some
> aquarium test kits. (By the way, one mL is the same as one cc - a cubic
> centimeter). Furthermore, the molarity of a solution is a standard way
to
> describe ionic concentration. A one molar solution of a molecule (or
ion)
> is a solution that contains the molecular (or ionic) weight, in grams,
> * of that molecule (or ion) per liter of solution. Therefore,
measurements
> in milliliters of a molar solution is another convenient way to produce a
> specific amount of material in milligrams. One milliliter of a one molar
> solution contains one thousandth of the amount of material in a mole. For
> example, one mL of a one molar solution of KCl contains 0.0391 grams or
39.1
> milligrams of K.
>
> Here is an example: Five mL of 1 molar KCl would be 0.005 liters,
which
> would contain 0.005 moles. Adding this KCl to 10 gallons of water (37.8
> liters), we now have 0.005 moles K in 37.8 liters, or 0.000132 moles per
> liter. Potassium has an atomic weight of 39.1, and so one mole of
potassium
> would weigh 39.1 grams. Multiplying the 0.000132 moles per liter times
39.1
> grams per mole gives 0.00517 grams per liter or 5.17 milligrams K per
liter
> or 5.17 ppm. This procedure is an alternative to directly adding 0.195 g
of
> K (0.005 moles) or 0.375 g KCl (also 0.005 moles). In 50 gallons, 5
> millimoles KCl produces approximately 1 ppm K.
>
>
> Concentrations Resulting from Trace Element Powders.
> A trace element mixture contains many different elements. The labels
usually
> indicate their composition in terms of their percentage by weight. For
> example, PP Ltd trace element powder has 6 different nutrient trace
elements
> as presented below. Knowing their weight by percent allows a direct
> calculation of the concentration resulting from adding a specific weight
per
> unit volume of water.
>
> As an example, the following concentrations would result from 1 gram (and
> 0.25 grams) in 50 gallons of water:
>
> - ---------------------------------------------------------
> Percent Concentration (ppm)
> Element by weight 1 g per 50 gal. 0.25 g per 50 gal.
> - ---------------------------------------------------------
> Fe 7 .37 0.1
> Mg 2 .11 0.03
> Zn 0.4 .02 0.005
> Cu 0.1 .005 0.001
> Bo 1.3 .07 0.02
> Mb 0.06 .003 0.001
> - ---------------------------------------------------------
>
> Similar tables can be produced for different volumes, including stock
> solutions.
>
> If one wants to use the powder directly and not have to worry about
storage
> problems associated with the prepared solution, I discovered that the
> measure that came with the Dupla test kits (e.g. iron kit) corresponds to
~
> 0.1 g of trace element powder. Therefore, 2 ½ measures in 50 gallons of
> water will yield the desired Fe and other concentrations. A more precise
> approach would utilize a stock solution and then a specific number of
> milliliters of solute can be decanted for each application. Let's say
that
> one wants a solution of 5 milliliters to contain 0.25 g of trace element
> powder. This means that one mL would contain 0.05 g and 1000 mL would
need
> 50 g.
>
> German, English and American Units for Hardness
> There is a lot of confusion caused by many different units to measure
> hardness. Sometimes hardness is expressed in terms of degrees. In the
UK, 1
> degree of hardness is equivalent to one grain of CaO per Imperial gallon
of
> water. In Germany, 1 deg (dH) is equivalent to 10 mg CaO per liter, while
in
> the U.S. water hardness is expressed in ppm of CaCO3. This is further
> complicated by the distinction between general hardness (GH) and
carbonate
> hardness (KH). General hardness considers both permanent hardness caused
by
> all calcium and magnesium compounds including their sulfates and
chlorides
> and temporary or carbonate hardness based on the carbonates. Strictly
> speaking, GH is always greater than or equal to KH. However, since KH is
> measured as a bicarbonate, one can appear to have KH without any Calcium
or
> Magnesium in solution.
>
> Here are some conversions:
>
> 1 degree of carbonate hardness (KH) = 17.9 ppm of carbonate
> (measured as CaCO3)
>
> 1 degree of general hardness (GH) = 7.14 ppm Calcium
> Or 17.9 ppm of CaCO3
>
>
>
> Chemicals Used to Increase Carbonate Hardness (KH) and General Hardness
(GH)
> Carbonate and general hardness can be increased by using several
chemicals
> which are available in grocery stores, pharmacies and the aquarium shop.
> These include Calcium Carbonate tablets (sold as a dietary supplement)
and
> sodium bicarbonate (baking soda).
>
> One g of CaCO3 yields 5.3 ppm in 50 gallons,
> so 3.4 g becomes 17.9 ppm or 1 KH.
>
> CaCO3 tablets sold in the pharmacy as a dietary supplement are 1.5 grams;
> thus 2 1/4 tablets are needed to raise one degree of KH. The tablets are
> pure calcium carbonate and dissolve very easily. I put them in a one
liter
> bottle of water to create a CaCO3 suspension. When this chalky liquid is
> poured into the water, it will take several hours before it will clear.
This
> is because it must react with CO2 to form the very soluble bicarbonate.
>
> A solution of 17.9 mg/L of CaCO3 in water gives KH of 1. This
> solution contains 0.179 mM Ca++, but does not contain any significant
CO3--,
> because this is converted to HCO3-, of which two are formed from each
CO3--,
> giving 0.358 mM HCO3-. We thus want 0.358 mM NaHCO3 in the solution.
The
> molecular weight of this is 84, so we want 30 mg/L sodium bicarbonate. In
50
> U.S. gallons, this amounts to 5.68 g or a little more than 1 teaspoon
(1.1 t).
>
> - -------------------------------------------
> Amount to achieve
> 1 degree KH in 50
> Compound Gallons of water
> - -------------------------------------------
> Sodium bicarbonate 1.1 teaspoon
> Calcium carbonate 2 1/4 tablets
> - -------------------------------------------
>
> How to Increase Calcium Concentration and GH
> Corresponding to the discussion of CO3, one g of CaCO3 yields 2.1 ppm of
Ca
> in 50 gallons of water. So, adding 2 Calcium tablets (3 g) will increase
the
> Ca concentration by 6.4 ppm. This is a little less than 1 degree of
general
> hardness. As with carbonate hardness 2 1/4 tablets are needed to get 1
> degree GH. You will note that an increase of 1 degree of carbonate
hardness
> will also cause an increase of one degree of general hardness.
>
> Concluding Remarks
> Adding chemicals should be done with caution. Unless nutrient
deficiencies
> are known or specific target concentrations are desired, these actions
are
> not needed and could even be harmful. Monitoring water chemistry is
> useful. This can be done by observing the behavior and appearance of the
> plants and fish (i.e. looking for symptoms of deficience or toxicity) or
by
> performing chemical testing. There are many commonly available general
(pH,
> KH, GH) and chemically specific tests (e.g. Ca, CO3, NO3, NH4, N, Fe)
which
> will ensure a stable system. Some elements like Potassium (K), however,
do
> not have common home test kits, so increases beyond the uptake by plants
and
> fish does warrants some attention. There is also concern about relative
> imbalance in concentrations because plants have the ability to consume
more
> chemical than they need and high concentrations of one element (e.g. Mg)
can
> block the uptake of other elements. Inhibition of nutrient uptake does
not
> appear to be a problem with other macro nutrients (N, P, K). Nitrogen can
be
> added in terms of Ammonium or Nitrate compounds. Although plants prefer
the
> former, the use of the latter is probably wiser, because of potential
> ammonia toxicity to fishes at relatively low concentrations.
>
> The directions for commercial nutrient preparations make assumptions
about
> fish load, amounts of fish food, plant density, growth rate and tap water
> chemistry. Nevertheless, they do provide an indication of safe
> concentrations, both in quantity and in their relative amounts.
Aquarists
> are advised to research their own water chemistry together with
available
> sources of information before they haphazardly start to dump stuff in the
> tanks. One excellent way to reduce the chance of overdosing from
routinely
> adding chemicals, however, is to add them at the time of a water change
and
> at a rate less than the desired concentration. Although many chemicals
are
> partially or completely consumed by actively growing plants, it is still
> theoretically possible that none may be used up between water changes.
> Therefore, lacking precise information on chemical uptake, adding
chemicals
> should (1) generally accompany an X percent water change, (2) be done
when
> the replacement water is lacking that substance, and (3) at a rate equal
to
> X percent of the desired target increment. The latter is needed to ensure
> that concentrations do not increase. For example, if the starting
> concentration is 25 ppm and 5 ppm are added with each 20 percent water
> change, then there will not be any increase in the final concentration.
> Knowing the target concentration is not always easy to ascertain. The
> range of safe concentrations is not always readily available. More
> information is desired on the consumption rates in the aquarium and
> desirable target concentration levels - both for individual chemicals and
> for their combinations. Differences between hard and soft water
situations
> are also important.
>
> Incremental concentrations of 1 ppm potassium and 0.5 ppm magnesium are
> utilized in some commercial preparations for weekly dosing together with
> biweekly water change. Potassium can probably be safely increased a few
> fold. Nitrogen and phosphates are often omitted from commercial
> preparations for aquarium plants, but aquarists have empirically
determined
> that 1-5 ppm NO3 are safe added amounts for nitrogen deficient tanks. As
a
> rough rule of thumb, these amounts would correspond to small quantities
of
> dry chemical: 1/4 tsp Potassium Chloride (Muriate of Potash) in 200
gallons
> of water, 1/4 tsp Magnesium Sulfate hydrate (Epsom Salt) in 70 gallons of
> water and 1/4 tsp of Sodium nitrate (Nitrate of Soda) in 70 gallons of
> water. Alternatively, the K and NO3 can also be roughly achieved by 1/4
tsp
> in 150 gallons of water. Assuming the replacement water does not have
any
> of these macro nutrients and there are no other suppliers of these
> chemicals, then dosing with water changes ensures that long term
> concentrations stay relatively low (0.5 to 2.5 ppm for Mg, 1-5 ppm for K
and
> 5 to 25 ppm for NO3). With actively growing plants, the steady state
> concentrations will invariably be much lower.
>
> Acknowledgments
> Thanks to Paul Sears for a general review and advice on carbonate
> chemistry; and to Paul Krombholz for providing the example calculation of
> ppm and molarity of solutions.
>
>
>
>
> Neil Frank Aquatic Gardeners Association Raleigh, NC
> The Aquatic Gardener - journal of the AGA - now in its seventh
year!!
>
>
> Date: Fri, 04 Apr 1997 07:24:55 -0500
> From: Neil Frank <nfrank at mindspring_com>
> Subject: Re: SAE's
>
> See <http://www.cco.caltech.edu/~aquaria/AGA/cyprinid.html>
>
>
>
> Neil Frank Aquatic Gardeners Association Raleigh, NC
> The Aquatic Gardener - journal of the AGA - now in its seventh
year!!
>
>
> Date: Fri, 4 Apr 1997 13:32:15 -0500 (EST)
> From: ChazzHess at aol_com
> Subject: Kitty Litter Revisited
>
> As an aside to the great Kitty Litter Controversy for the past two years
we
> have been working with a potting media/soil amendment called Profile.
This
> material, like many "kitty litters" is a fired clay. However, Profile
goes
> under a more rigerous firing process. The attributes of this material
are
> that it has some cation exchange capacity, contains some plant nutrients
> (e.g., calcium) and has been shown to retain its positive physical
attributes
> for several years in terrestrial conditions (I'm not sure if the tests
were
> performed in soil or as a potting media). It's also more attractive than
> kitty litter (i.e., colored other than grey). For those of you who are
> considering a kitty litter substrate, this medium may be a better choice.
>