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Re: Aquatic Plants Digest V2 #833
Aquatic Plants Digest wrote:
>
> Aquatic Plants Digest Monday, July 14 1997 Volume 02 : Number 833
>
> In this issue:
>
> Re: Brown silt
> Re: Cheap Substrate Heating (G. Booth)
> Re: CO2-Kh-PH table: Reference Wanted
> Re: Aquatic Plants Digest V2 #832
> Cheap Substrate Heating
>
> See the end of the digest for information on subscribing to the
> Aquatic Plants mailing list and on how to retrieve back issues.
>
> ----------------------------------------------------------------------
>
> Date: Mon, 14 Jul 1997 08:09:59 -0600
> From: George Booth <booth at hpmtlgb1_lvld.hp.com>
> Subject: Re: Brown silt
>
> >Date: Sun, 13 Jul 1997 10:14:44 -0700
> >From: AL ROBINSON <amrsb at erols_com>
> >
> > Having several tanks, I recently started a plant tank. 29L, two floura
> > grow lamps, power heads etc. etc. ... I have encountered a brown silt
> > like substance covering the entire tank. I have tried gently wiping it
> > from the plants leaves only to have it rapidly return. Could it be the
> > peat breaking down? Help.
>
> I don't know what "floura grow" lamps are but I bet they aren't very
> bright. What you are seeing is probably "brown algae" or diatoms.
> These typically appear when the light levels are too low. "Silt"
> (fine dirt) would not stick to things very well and would become
> suspended when you stirred the water. Since you had to "wipe" it, I'm
> guessing diatoms. Does this stuff look like the crud that forms in
> water tubing and filter boxes that aren't lighted? Same thing.
>
> George
>
> ------------------------------
>
> Date: Mon, 14 Jul 1997 10:35:23 -0400
> From: Karsten Klein <kklein at uceng_uc.edu>
> Subject: Re: Cheap Substrate Heating (G. Booth)
>
> George wrote:
>
> >You would probably want this to be a sealed system. Pulling aqaurium water
> >in would bring detritus and clog the tube. And dumping the water back into
> >the tank might heat the tank too much IF all the heat was not transferred
> >to the gravel. I suspect the silicon tubing is a pretty good insulator.
>
> I do not quite agree with George saying silicon tubing
> would be a pretty good insulator, it is rather the opposite.
> Assuming Silicon tubing contains mostly Silicon, it actually
> is a very good thermal conductor. Here are some thermal
I thought that the tubing was made of silicone, or are they the same
thing?
> conductivities to compare (in units of W/cm*C):
> Silicon: 1.57
> Steel: 0.97
> Stainless Steel: 0.329
> Iron: 0.803
> If one wants to get a higher thermal conductivity than
> Silicon, there are only a few options, e.g.:
> Aluminum 2.36
> Diamond 20.0 (WOW!)
> Meaning one could use Carbon in order to get a very very high
> conductivity.
> In order to achieve a relatively even heat distribution in
> the gravel, one could split the tubing coming from the ballast
> into 4 or 5 tubes going to different places in the substrate
> and then merging them back together, guiding the cooled
> water back to the "ballast heater".
>
> Regards,
> Karsten.
>
> ------------------------------
>
> Date: Mon, 14 Jul 1997 09:31:01 -0600
> From: George Booth <booth at hpmtlgb1_lvld.hp.com>
> Subject: Re: CO2-Kh-PH table: Reference Wanted
>
> >Date: Sun, 13 Jul 1997 18:12:23 -0400
> >From: Sanjay Joshi <sjoshi at psu_edu>
> >
> >I have seen the CO2-Kh-pH tables in several places. I am looking for
> >a published reference that explains the chemistry behind these
> >tables, and also provides the formulas and equations used to derive
> >the tables.
>
> I can *almost* explain the derivation of the tables. Karla helped me
> with the derivation and calculations once upon a time. I still have
> all the data but I am missing a key conversion that allows me to
> perform the calculations. I need to convince Karla to help once
> again. Or perhaps one of the chemistry gurus can fill in the missing
> pieces.
>
> First some background. This was written quite a while ago and I'm a
> bit unsure of some of the details. Please feel free to correct me.
>
> ===== Begin: from an old posting ======================================
>
> Karla worked up some water chemistry last night to share with the net.
> This is detailed information about the carbonate buffering system in an
> aquarium and how CO2 and KH and pH are related.
>
> Buffering
> - ---------
>
> First of all, when strong acids like sulfuric or hydrochloric acid are
> added to water, they completely ionize into hydrogen ions and the
> corresponding salt (HCL -> H+ + Cl-). However, when weak acids like
> carbonic or phosphoric acid are added to water, they form "conjugate
> acid-base pairs" (HA and A-) that are in equilibrium. There is an
> "equilibrium dissociation constant" (K) for a weak acid defined as:
>
> [H+] [A-] where [H+] = concentration of hydrogen ions
> K = ----------- [A-] = concentration of the salt or conjugate base
> [HA] [HA] = concentration of the weak acid
>
> This relationship is for the reaction:
>
> HA <==> H+ + A-
>
> Note that K will vary slightly with temperature (and other things?).
>
> The acid-base equilibrium is described by the Henderson-Hasselbach equation
> (derived from the "equilibrium dissociation constant" relationship):
>
> [A-]
> pH = pK + log ---- or pK = pH when [A-] = [HA] {since log(1)=0}
> [HA]
>
> The point where pH = pK is the point at which the system has the most
> buffering capacity to handle additions of acids or bases and can occur at
> multiple pH values for different buffering systems such as carbonate and
> phosphate.
>
> For example, with the carbonate buffering system, this occurs at pH 6.37
> and pH 10.25. At pH 6.37, H2CO3 (carbonic acid) and HCO3- (bicarbonate)
> are present in equal concentrations. At pH 10.25, HCO3- and CO3--
> (carbonate) are present in equal concentrations.
>
> H2CO3 <==> H+ + HCO3- <==> H+ + CO3--
>
> Another common example is the phosphate buffering system (I believe
> products like pH-UP and pH-DOWN are based on this system). This has
> equilibrium points at pH 2.13, 7.21 and 12.32.
>
> H3PO4 <==> H+ + H2PO4- <==> H+ + HPO4-- <==> H+ + PO4---
>
> Conjugate Base to Weak Acid Ratio
> - ---------------------------------
>
> At any pH point, the Henderson-Hasselbach equation describes the ratio of
> [A-] (conjugate base) to [HA] (weak acid). For example, consider the
> carbonate system at pH = 7.0:
>
> [A-] [HCO3-]
> pH = pK + log ---- -> 7.0 = 6.37 + log ------- ->
> [HA] [H2CO3]
>
> [HCO3-] 4.27
> 0.63 = log (4.27) -> ------- = ---- -> 77% / 23%
> [H2CO3] 1.00
>
> [I'm afraid I can't figure out where 77%/23% came from! George]
>
> Therefore at pH 7.0, the carbonate system is 23% H2CO3 and 77% HC03-. If
> you make a graph of pH versus the relative base/acid concentrations, you
> will get a "S" curve due to the logarithmic nature of the equation. An
> important observation is that at the pK point the slope of the curve is
> nearly vertical, i.e., a large change in relative concentration produces
> only a small change in pH.
>
> 100% | ___ ___
> | / /
> | / /
> salt | | |
> concen. | | |
> 50% | H2CO3 + HCO3- + CO3--
> | | |
> | | |
> | / /
> | ___ / ___ /
> 0% |___________________________
> 6.4 10.3
> pH
>
> (A graph of this appears on page 32 of the _Aquarium_Atlas_ (Baensch, 1987).
>
> How this system relates to buffering can be seen from an example. Consider
> the process of protein ammonification (from "Water Chemistry in Closed
> System Aquariums", A.J.Gianoscol, 1987). One of the byproducts of this
> process is phosphoric acid (H3PO4). At a pH of around 7, phosphoric acid
> will rapidly dissociate completely to H+ and dihydrogen orthophosphate
> (H2PO4-) and dihydrogen orthophosphate will dissociate partially to H+ and
> monohydrogen orthophosphate (HPO4--) (note that monohydrogen orthophosphate
> won't dissociate to H+ and phosphate (PO4---) until the pH gets up around
> the third equilibrium point at pH 12.32). The free hydrogen ions (H+) can
> then combine with bicarbonate (HCO3-) to form carbonic acid (H2CO3),
> shifting the acid/salt balance slightly downward. Since the reaction takes
> place near the pK point, the slight shift in concentration does not
> measurably affect pH.
>
> To put it simply, the system is "buffered" since any free H+ ions can
> combine with bicarbonate without altering the pH much. Naturally, as
> more and more "buffering capacity" is used up, the pH will be able to
> shift more and more. Also, salts from the weak acids build up in the
> water. Both these consequences point to the need for occasional water
> changes to remove the salts and replenish the buffer.
>
> Since the phosphate system has an equilibrium point at pH 7.2, you
> would think it would be preferred to the carbonate system for
> freshwater aquariums. This is not true for three reasons. First,
> aquariums naturally produce organic acid compounds (metabolism
> by-products) so you are most concerned about buffering acids. If you
> keep your pH at around 7.0, you are already below the phosphate pK
> point and will keep getting further away as acids are buffered,
> reducing the amount of buffering potential. Using the carbonate
> system, pH 7.0 is above the pK so that any acid buffering will move
> the pH even closer to the best buffering point. Second, plants can
> use the carbon compunds which are part of the carbonate system.
> Third, phosphates tend to grow algae, which is not desired.
>
> Using the Henderson-Hasselbach equation and some chemical
> calculations, you can create a chart showing the relationship of pH to
> KH to CO2 (wink, wink, nudge, nudge).
>
> ===== End: from an old posting ======================================
>
> Sounds easy, right? OK, here are some scribbles from my old work
> sheet -- the results of which created a fine pH/KH/CO2 table.
>
> First, the "equation" that Karla gave me:
>
> [A-]
> pH = pK + log ----
> [HA]
> - -or-
> HCO3- CO2
> pH = 6.37 + log [ ------ = ------- ]
> H2CO3 1.64CaCO3
> - -or-
> HCO3- CO2
> pH - 6.37 = log [ ------ = ------- ]
> H2CO3 1.64CaCO3
>
> I don't remember how to interpret this. Is HCO3-/H2CO3 "the same as"
> CO2/1.64CaCO3? Or is the relation part of the equation? What's the
> deal with "1.64"? HCO3- is the KH which is 17.8 mg/l CaCO3. Where
> does this fit in? Are the concentrations in millequivalents (I
> suspect it is but I am meq impaired).
>
> Next, the results. I calculated two points for each of many pH values.
> The points are the CO2 concentration for that pH at KH=1 degree and
> KH=8 degrees. This makes a nice linear chart with CO2 on the vertical
> axis, KH on the horizontal axis and showing lines of constant pH.
>
> Here's some sample results (CO2 in mg/l):
>
> CO2@ CO2@
> pH log[A-/HA] [A-/HA] KH=1 KH=8
> - --- ---------- ------- ----- -----
> 6.4 0.03 1.07 12.01 95.98
>
> 6.7 0.33 2.14 6.00 47.99
>
> 7.0 0.63 4.27 3.01 24.05
>
> 7.7 1.33 21.38 0.60 4.80
>
> 8.0 1.63 42.66 0.30 2.41
>
> Notice the nice "power of 10" relationship between pH and CO2 values.
> (pH=6.7, CO2=6.0) and (pH=7.7, CO2=0.6).
>
> I made a notation at the top of the table showing a factor of "0.292"
> for KH=1 and "2.334" for KH=8 (2.334 = 8 * 0.292). I don't remember
> what this factor meant.
>
> Also note that [A-/HA] * CO2 is a constant value (with rounding):
>
> 1.07 * 12.01 = 12.8 1.07 * 95.98 = 102.7
> 4.27 * 3.01 = 12.8 4.27 * 24.05 = 102.7
> 21.38 * 0.60 = 12.8 21.38 * 4.80 = 102.6
>
> So, if someone could figure out how to calculate the [A-/HA] values
> for CO2 and KH, the equation is simple to rearrange to solve for
> any one of the variables given the other two.
>
> I feel sooooo stupid.
>
> I'll forward this to Karla. Perhaps she will be inspired to help out
> once again.
>
> George
>
> ------------------------------
>
> Date: Mon, 14 Jul 1997 06:58:24 -0500 (EST)
> From: "Roger S. Miller" <rgrmill at rt66_com>
> Subject: Re: Aquatic Plants Digest V2 #832
>
> > >Date: Sun, 13 Jul 1997 18:12:23 -0400
> > >To: aquatic-plants at actwin_com
> > >From: Sanjay Joshi <sjoshi at psu_edu>
> > >Subject: CO2-Kh-PH table: Reference Wanted
> > >
> > >I have seen the CO2-Kh-pH tables in several places. I am looking for
> > >a published reference that explains the chemistry behind these
> > >tables, and also provides the formulas and equations used to derive
> > >the tables.
> > >
> > >Thanks,
> > >
> > >sanjay joshi
>
> There are probably any number of references for the tables, as they are
> based on rather basic aqueous chemistry. My fav is
>
> Skougstad, M. J. et al, 1979. Methods for determination of inorganic
> substances in water and fluvial sediments. Techniques of
> Water-Resources Investigations of the United States Geological
> Survey, Chapter A1. U.S Government Printing Office, Washington,
> DC. 626pp.
>
> I believe there is a newer version of the same reference.
>
> The formula (on page 280) is rather simple:
>
> mg/L CO2 = 1.60 X 10^(6.0-pH) X mg/L HCO3-
>
> Note that HCO3- in mg/L is alkalinity in German degrees kH * 21.8.
>
> The formula is good for fresh water with less than 800 mg/l of solutes.
> Sea water will depart significantly from this relationship because of
> variations in activity coefficients and the importance of several
> complexes (CaHCO+, and so on) that form in concentrated solutions. It may
> be possible to use programs like USGS' PHREEQC code to produce a
> similar, semi-empirical relationship for sea water.
>
> Roger Miller
>
> ------------------------------
>
> Date: Mon, 14 Jul 1997 10:17:56 -0400 (EDT)
> From: Paul Chapman <chapman at SEDSystems_ca>
> Subject: Cheap Substrate Heating
>
> George Booth wrote:
>
> >Think about what would be going on here. Let's say that the water in the
> >tubing did pick up some heat from the ballasts (it should since the
> >ballasts typically run at 150 F or so). So the heated water in the tube
> >begins to heat up the gravel. To heat the gravel, the water in the
> >tubing has to give up heat. In a very short distance, the water in the
> >tubing has given up all it's heat to the gravel and becomes the same temp
> >as gravel. Over a very long time, the gravel around the first part of
> >the tubing gets very hot and gravel downstream gets cooler and cooler.
> >This does not produce a very even heat density in the gravel. Too hot in
> >one section, not hot enough in some sections, not hot at all in the last
> >sections.
>
> Seems to me this would be a problem with a very slow flow system
> only i.e. if the water is allowed to sit in the tubing long enough to give
> up all of its heat. Why not increase the heat available to the system by
> cranking up the rate of flow? Granted, the heat in the first sections
> would be higher, but you can make the difference between the
> first and final sections negligibly small with a high flow rate.
>
> >You would probably want this to be a sealed system. Pulling aqaurium water
> >in would bring detritus and clog the tube. And dumping the water back
> >into the tank might heat the tank too much IF all the heat was not
> >transferred to the gravel. I suspect the silicon tubing is a pretty good
> >insulator.
>
> Yes, it would have to be a sealed system. As far as the silicon tubing
> being a good insulator, maybe, but I can feel the heat off such a tube
> when I run hot water through it....Any idea on how much heat transfer is
> really necessary?
>
> >And, of course, you would have a problem with brown algae build up unless
> >you used sterlized water and somehow kept it sterilized.
>
> Could be a problem. Solutions might involve adding something to the water
> in the sealed system to keep the algae down.
>
> >How would you control the heating? Hopefully, the on-off cycle of the
> >lights would just perfectly balance the heating of the gravel so the tank
> >water would not get too hot.
>
> Put the pump on a timer, or under thermostatic control, the same way you
> do your Dupla cables.
>
> >And, boy, won't that be exciting when the tubing springs a leak, wets down
> >the ballasts and causes a fire! Cool.
>
> Seems like the risk of fire is negligible if a ground fault circuit
> interrupter is used. Other techniques could be used as well, such as
> bolting the ballast on top of a metal plate, and fixing the coils on the
> bottom.
>
> Paul Chapman
> Saskatoon, SK.
>
> ------------------------------
>
> End of Aquatic Plants Digest V2 #833
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