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> 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 --
> trace element mixes, commonly available chemicals and use of tools like
> measuring spoons.  I thought people might find it useful (I hope the
> are not messed up too badly and that people do not take offense to the
> 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
> (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
> 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
> water and particular aquarium conditions. This discussion does not
> extensively focus on the specific circumstances indicating when and why
> adds these chemicals to the aquarium. This has been discussed elsewhere.
> Instead, the focus here is on determining the concentrations and
> 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
> 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
> 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
> 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
> (1,000 mg) per 1,000 L.
> For Americans and others who may not be comfortable with the metric
> 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
> water can be determined by simple arithmetic. Because 10 gallons is 37.85
> then 1 gram (g) / 10 gallons is 26.4 ppm.  For 50 gallons, one gram
> 5.3 ppm.
> Often we are not interested in achieving the concentration for a
> chemical compound (like table salt, NaCl), but separately for its
> 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
> of the entire chemical. For example, Chloride represents 61 percent of
> weight of table salt. 
> If the concentration of Fe in the trace element. mix is 7 percent, then
> 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,
> is often used as the indicator variable and tells the aquarist if all
> 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
>                   (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
> The same principle can be used to determine the amount of any compound
> table salt) needed to achieve a desired concentration of any ion or
> 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
> -
> 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
> 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
> 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
> - -------------------------------------------------------------------
> 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
> solution of KCl is 39.1 g K per liter.
>   With solutions, milliliters (mL) are a standard unit of measurement.
> 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
> describe ionic concentration.  A one molar solution of a molecule (or
> is a solution that contains the molecular (or ionic) weight, in grams,
> * of that molecule (or ion) per liter of solution.  Therefore,
> 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
> milligrams of K.
>    Here is an example: Five mL of 1 molar KCl would be 0.005 liters,
> 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
> would weigh 39.1 grams. Multiplying the 0.000132 moles per liter times
> grams per mole gives 0.00517 grams per liter or 5.17 milligrams K per
> or 5.17 ppm. This procedure is an alternative to directly adding 0.195 g
> 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
> indicate their composition in terms of their percentage by weight. For
> example, PP Ltd trace element powder has 6 different nutrient trace
> as presented below. Knowing their weight by percent allows a direct
> calculation of the concentration resulting from adding a specific weight
> 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
> 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  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
> 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
> 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
> water. In Germany, 1 deg (dH) is equivalent to 10 mg CaO per liter, while
> the U.S. water hardness is expressed in ppm of CaCO3.  This is further
> complicated by the distinction between general hardness (GH) and
> hardness (KH). General hardness considers both permanent hardness caused
> all calcium and magnesium compounds including their sulfates and
> 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
> 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
> Carbonate and general hardness can be increased by using several
> which are available in grocery stores, pharmacies and the aquarium shop.
> These include Calcium Carbonate tablets (sold as a dietary supplement)
> 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
> 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.
> 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
> because this is converted to HCO3-, of which two are formed from each
> giving 0.358 mM HCO3-.  We thus want 0.358 mM NaHCO3 in the solution. 
> molecular weight of this is 84, so we want 30 mg/L sodium bicarbonate. In
> 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
> in 50 gallons of water. So, adding 2 Calcium tablets (3 g) will increase
> Ca concentration by 6.4 ppm. This is a little less than 1 degree of
> 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
> will also cause an increase of one degree of general hardness.
> Concluding Remarks
> Adding chemicals should be done with caution. Unless nutrient
> are known or specific target concentrations are desired, these actions
> 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
> performing chemical testing.  There are many commonly available general
> KH, GH) and chemically specific tests (e.g. Ca, CO3, NO3, NH4, N, Fe)
> will ensure a stable system.  Some elements like Potassium (K), however,
> not have common home test kits, so increases beyond the uptake by plants
> fish does warrants some attention. There is also concern about relative
> imbalance in concentrations because plants have the ability to consume
> chemical than they need and high concentrations of one element (e.g. Mg)
> block the uptake of other elements. Inhibition of nutrient uptake does
> appear to be a problem with other macro nutrients (N, P, K). Nitrogen can
> added in terms of Ammonium or Nitrate compounds. Although plants prefer
> 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
> 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. 
> are advised to research their own water chemistry together with 
> sources of information before they haphazardly start to dump stuff in the
> tanks.  One excellent way to reduce the chance of overdosing from
> adding chemicals, however, is to add them at the time of a water change
> at a rate less than the desired concentration. Although many chemicals
> 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
> should (1) generally accompany an X percent water change, (2) be done
> the replacement water is lacking that substance, and (3) at a rate equal
> 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
> 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
> that 1-5 ppm NO3 are safe added amounts for nitrogen deficient tanks.  As
> rough rule of thumb, these amounts would correspond to small quantities
> dry chemical: 1/4 tsp Potassium Chloride (Muriate of Potash) in 200
> 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
> in 150 gallons of water.  Assuming the replacement water does not have
> 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
> 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
> 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
> 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
> have been working with a potting media/soil amendment called Profile. 
> material, like many "kitty litters" is a fired clay.  However, Profile
> under a more rigerous firing process.  The attributes of this material
> that it has some cation exchange capacity, contains some plant nutrients
> (e.g., calcium) and has been shown to retain its positive physical
> for several years in terrestrial conditions (I'm not sure if the tests
> 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.