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Would this work as a chiller?

Started by Guest_EricaWieser_* , Apr 08 2011 12:06 AM

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#1 Guest_EricaWieser_*

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Posted 08 April 2011 - 12:06 AM

I was reading this webpage about a DIY water change system and it occurred to me that it might also work as a chiller for the water because it taps into the cold water main. If you had a decent flow rate in the tank, I bet the water would stay mighty cold.
http://www.aquaticpl...nge-system.html

Now that the drainage part is constructed the only other thing to do is tap into your house water main! Now, if you are like me, you will probably be nervously shaking your head right now scrambling to click the back button to get out of here. But! Have no fear! A company called Watts makes an ice-maker installation kit that is designed to clamp onto your main copper water pipe and make a small hole in it, allowing adjustable water flow. All you need to do is buy the kit, screw it onto your a *cold* water pipe somewhere in the house (this can be a sink, bath, boiler copper water pipe of ~1 inch diameter) and turn the knob for fresh clean water!

Picture of kit in action: http://farm3.static....7e5a6ac45_o.jpg
That kit seems to cost about $15 to $20.

Thoughts?

Edited by EricaWieser, 08 April 2011 - 12:10 AM.

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#2 Guest_Skipjack_*

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Posted 08 April 2011 - 06:58 AM

And as a chlorinator. At least if you try to run enough water to cool the tank. The idea is a slow drip, which is not enough to cool the tank very much.

Edited by Skipjack, 08 April 2011 - 07:02 AM.

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#3 Guest_EricaWieser_*

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Posted 08 April 2011 - 07:56 AM

And as a chlorinator. At least if you try to run enough water to cool the tank. The idea is a slow drip, which is not enough to cool the tank very much.

I was thinking about the chlorine issue and decided to look some stuff up. Is there enough chlorine in municipal water to harm fish, I wondered? Or is it the chloramine formed from contact of chlorine gas with nitrogenous solutes in the aquarium water that hurts fish? I did some research.

Data: Water coming out of the tap is typically 0.5 to 2.0 ppm depending on how far you are from the water treatment facility (source: http://www.edstrom.c...clib/mi4174.pdf page 3).

Data: The chlorine toxicity in fish seems to be pretty high. This source http://books.google....epage&q&f=false claims 0.2 ppm causes low level mortality in fish, with 0.4 ppm being sometimes 100% lethal. This source http://www.isws.illi...I/ISWSRI-85.pdf claims a TL50s of .18 to .33 mg/L for bluegill and 0.09 for channel catfish, with an acceptable level of chlorine being 0.04 mg/L.

So I guess yeah, chlorine is very toxic. *sighs* It's too bad that chlorine gas itself is so toxic, because my water would not have formed any chloramine. Because I have a Walstad setup, my ammonia, nitrite, and nitrate are all 0 ppm, and 0 chloramine would have formed. Le sigh.

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#4 Michael Wolfe

Board of Directors

North Georgia, Oconee River Drainage

Posted 08 April 2011 - 08:26 AM

I cant remember where I read this, so it is going to sound crazy...

When I set up my first Elassoma tank I read (somewhere) that you should stuff it with plant and snails (of course)... and that you should do you water changes with straight chlorinated water... do not treat it with de-chlor... sounded crazy to me, but I did it and never had a problem... I believe the theory was that in the tannin stained low pH waters that some of these fish come from, there is very little microbial like in the water... rambling now but the rangers in the Okefenokee told us that the tea colored water there was the safest to drink anywhere because the low pH meant that there were virtually no pathogens in the water... anyway, all of that to say that I have done 1 and 2 gallon water changes in a 10 gallon tank of Elassoma with straight tap water and had no ill effects and a breeding population of E.okefenokeei

Either write something worth reading or do something worth writing. - Benjamin Franklin

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#5 Guest_Newt_*

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Posted 08 April 2011 - 08:44 AM

I'm sure you could use the cold water line to chill your tank water as long as you kept the two water supplies separate- sort of an exchange system. It would be somewhat elaborate and possibly no less expensive than a store-bought chiller, but fun to make if you're a DIY type of person. What you'd need is a chamber of some kind on the cold water line, a long coil of stainless steel tubing, and a pump. You would pump aquarium water through the tubing, the greatest length of which would be inside the chamber that the cold water is flowing through. Warm aquarium water would lose heat to the cold water outside the tubing, and cooler aquarium water would return.

The trouble would be keeping a steady supply of cold water. Most of the time it's just sitting in the pipes, which in this case would mean the cold water would approach the temperature of the aquarium water and so become less effective between water useages. The incoming line for the water heater seems like a good candidate, especially if you have a multiple-person household and use a lot of hot water on a daily basis, or someplace well up the main line so all the fixtures in the house draw off it.

Just thinking out loud, really. It's probably not a worthwhile project.

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#6 Guest_mywan_*

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Posted 08 April 2011 - 10:21 AM

The principle is the same as what I have in mind for a paludarium using a large external reservoir where the water is merely recycled between inside and outside. If the pumps goes down the circulation simply stops. As long as the temperature of the external reservoir is cool, and/or cascaded over a waterfall for evaporative cooloing, it keeps the tank cool as well as adding a far greater total oxygen capacity among other things.

A similar more limited sort reservoir might be used in your situation to lower the chlorine levels before going into the tank. Even if it does not hurt the Elassoma I would still be concerned about the blackworms, nitrosomonas, nitrobacters, etc. Chlorine tends to get reduced pretty quickly naturally if allowed access to air. Perhaps a reasonably large watertight trashcan as a reservoir which overflows into the tank would allow the chlorine to get reduced sufficiently before contact with the tankwater. Exactly the same overflow control could be used to feed the reservoir water to the tank. Reduction, as well as evaporated cooling and oxygenation, could also be greatly increased by spraying the tap water onto a drain board covered with filter media like what is used for home HVAC return air intakes. About $2 dollars at Wal Mart, just take the paper edging off. This would massively increase surface area of air on the chlorinated water. At the very least the chlorine content should be well below tap water levels.

It would be nice to do water changes simply by turning the tap on.

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#7 Guest_gzeiger_*

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Posted 08 April 2011 - 10:42 AM

Chlorine concentrations as you said may be fairly variable, but I will note that my water (Charleston SC) killed Gambusia affinis in a matter of about an hour.

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#8 Guest_EricaWieser_*

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Posted 08 April 2011 - 12:08 PM

I did some calculations to find the maximum flow rate into my 55 gallon (208 L) tank, given that the concentration of Chlorine (Cl2) shouldn't rise above 0.04 mg/L.

Known: Cl2 concentration varies from 0.5 mg/L to 2.0 mg/L. Assume an average value of (0.5+2.0)/2 = 1.25 mg Cl2 / L water
Known: Concentration of chlorine in the tank should be 0.04 mg/L or less.

Accumulation = In - Out + Generation - Consumption
Assuming that the system is in steady state, accumulation = 0.
Let the consumption term equal the evaporation rate out of the tank.
Let generation term equal zero because no chlorine is being generated in the tank. Due to the zero mg/L ammonia, nitrite, and nitrate concentration, the chloramine generation (and its equilibrium exchange with chlorine) is zero, and this assumption is probably valid.

Accumulation = In - Out + Generation - Consumption
0= (1.25 mg/L)*(X L/min) - (0.04 mg/L)*(X L/min) + 0 - Evaporation rate
Evaporation rate = (1.25 mg/L)*(X L/min) - (0.04 mg/L)*(X L/min)
That is on a per time (rate) setup. If we assume that this is happening over the course of one minute, then time cancels out across both sides of the equation.
n in mg = (1.25 mg/L)*(X L) - (0.04 mg/L)*(X L)
n in mg = X L *(1.25 mg/L - 0.04 mg/L)

The value of n can be found (according to Dr. Dai of the Chemical Engineering department of CWRU, I had to ask for help on this part) from the partial pressure. According to http://pubs.acs.org/...021/ie50378a014 , the partial pressure of Cl2 in water at 25 degrees C equals 0.973 atm at 0.6 g Cl2 / 100 g water. If we convert that to mg/l, that is
(600 mg Cl2 / 100 g water)*(3787.5 gram / gallon)*(1 gallon / 3.7854 L ) = 6003 mg/L
The data comes from a solution of 6003 mg/L and we are talking about 0.5 mg/L, so this is a bad fitting region to get the volatility data. If anyone has any evaporation data for lower concentration chlorine gas in water, that would be helpful.

Recalling the ideal gas law, n = P V / R T
n = (0.972 atm * 208 L)/(298 K * 0.08206 (L atm / mol K))
n = 8.267 moles
n in grams = 8.267 moles * 71 grams Cl2 / mol
n in grams = 587 grams

587,003 mg = X L *(1.25 mg/L - 0.04 mg/L)
X = 485126 L/min
X = 128157 gal/min

That's a ridiculously high flow rate, which makes me think there has to be an error. I bet it's because the volatility was really, really off. But I think that getting the volatility from a region with higher concentration chlorine gas in water would inflate the evaporation term to an artificially higher value? I'm not sure on that way. I need better data, particularly the partial pressure of chlorine at something less than 10 mg/L. I'm going to try to find that and see if I can redo the math.

Edited by EricaWieser, 08 April 2011 - 12:27 PM.

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#9 Guest_EricaWieser_*

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Posted 08 April 2011 - 12:29 PM

Aha! I had to actually read the source where I got my original chlorine partial pressure data.
At 25 degrees Celsius, the partial pressure of chlorine is 0.0621 atm and 0.104 g Cl2 / 100 g H2O.
That translates into 1041 mg Cl2 / L.

It's closer to our range of .04 to 2 mg/L, but still not that close. Still, maybe I can get a closer estimation.

Recalling the ideal gas law, n = P V / R T
n = (0.0621 atm * 208 L)/(298 K * 0.08206 (L atm / mol K))
n = 0.5282 moles
n in grams = 0.5282 moles * 71 grams Cl2 / mol
n in grams = 37.5 grams

37,503 mg = X L *(1.25 mg/L - 0.04 mg/L)
X = 44121 L/min
X = 11,656 gal/min

It went from 130,000 gal/min using the higher concentration evaporation data to 12,000 gal/min using the lower concentration evaporation data. I'm going to do a fit line to this data on this pdf here and see if I can get an estimate at around 2 mg/L.

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#10 Guest_EricaWieser_*

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Posted 08 April 2011 - 12:51 PM

Righto, I got a fit line.

Chlorine partial pressure fit line.bmp 1.78MB 3 downloads
http://gallery.nanfa...t line.bmp.html

y = 0.0002x - 0.148, R² = 0.9962

Partial pressure (atm) = 0.0002 * solubility (mg Cl2/ L) - 0.148
At 0.04 mg/L, the partial pressure would be -0.147992 atm. Hmm. Well drat, that's a negative number. My guess is that 0.0002 wasn't enough significant figures. I'm going to plot them with flipped axes and see if I get more sig figs.

flipped axes.bmp 1.76MB 1 downloads
http://gallery.nanfa...d axes.bmp.html

The new fit line is
y = 5653.7x + 848.15, R² = 0.9962
solubility = 5653.7 * partial pressure + 848.15
0.04 mg/L = 5653.7 * X + 848.15
X = -.15 atm

Drat. That's just not working.
New idea: Calculate partial pressure (using first fit line) for 100 mg/L:
Partial pressure (atm) = 0.0002 * solubility (mg Cl2/ L) - 0.148
Partial pressure (atm) = 0.0002 * 100 (mg Cl2/ L) - 0.148
Partial pressure (atm) = -.128 atm
That's less negative but still not above zero. Oh great. This data is not accurate enough to extend a fit line all the way down to the low concentrations I need. Well phooey.

I'm off to search for partial pressure data of Cl2 at the .04 mg/L region. Ta ta.

Edited by EricaWieser, 08 April 2011 - 12:53 PM.

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#11 Guest_EricaWieser_*

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Posted 08 April 2011 - 01:04 PM

Okay, I can't find any better data. I'm just going to assume it's around 0.01 atm and work with that.

n = (0.01 atm * 208 L)/(298 K * 0.08206 (L atm / mol K))
n = 0.0851 moles
n in grams = 0.0851 moles * 71 grams Cl2 / mol
n in grams = 6 grams

6000 mg = X L *(1.25 mg/L - 0.04 mg/L)
X = 5000 L/min
X = 320 gal/min

There's an estimation for ya. Unless the flow rate is 300 gallons per minute, the chlorine concentration is going to be above the 0.04 mg/L that is considered completely safe.

Now, let's say our tank volume was about 400 gallons (the tank I plan on building in the near future). How much flow rate would be needed for 0.04 mg/L chlorine?

n = (0.01 atm * 1514 L)/(298 K * 0.08206 (L atm / mol K))
n = 0.619 moles
n in grams = 0.619 moles * 71 grams Cl2 / mol
n in grams = 44 grams

I'm surprised to see the grams increase. I guess I was thinking about it the wrong way before. Continuing with the calculation,

44000 mg = X L *(1.25 mg/L - 0.04 mg/L)
X = 36329 L/min
X = 9597 gal/min

Wait, what? A larger tank needs a larger flow rate?! I'm so confused! Something in my original equation must have been wrong. I wonder if I'm even going about this in the right way. *head hurts*

I am going to set this down for a little while and look at it later and see if it makes any more sense.

Edited by EricaWieser, 08 April 2011 - 01:06 PM.

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#12 Guest_gzeiger_*

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Posted 08 April 2011 - 02:49 PM

Why not run the cold water through a submerged pipe as a heat exchanger rather than running it into the tank? Then you could keep the chlorine to zero.

By way of comparison, Gambusia from the same site as that mentioned above generally survive being bitten by a pickerel, and so far I've seen zero mortality from feeder collection even in the heat of a SC summer. I wouldn't mess with chlorine.

If you insist, though, an aeration source such as a sponge filter or just an airstone would help establish a lower equilibrium concentration, and it might help to have the inlet be a spray or drip rather than a flow from a submerged pipe.

Edited by gzeiger, 08 April 2011 - 02:52 PM.

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#13 Guest_EricaWieser_*

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Posted 08 April 2011 - 03:29 PM

Ah! I understand now. Those are the flow rates required to get the chlorine up to 0.04 mg/L. Anything below those flow rates means a chlorine concentration below .04. That's why the flow rates increased with increased volume.
I'm going to do a nice pretty write up when I get home to thoroughly explain everything.

Edited by EricaWieser, 08 April 2011 - 03:33 PM.

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#14 Guest_EricaWieser_*

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Posted 08 April 2011 - 04:55 PM

So, there's good news and there's bad news. The good news: I made a nice writeup. The bad news: Evaporation was calculated, Dr. Q of the Chem E department points out, without a surface area term in it.

chlorine schematic.png 9.82KB 0 downloads
http://gallery.nanfa...ematic.png.html

Assume steady state
dX1/dt = 0
Limit: X1 < 0.04 mg/L

Mass balance:
0 = In - Out + Generation - Consumption
0 = XF - X1F - Evaporation
Evaporation = F (X-X1)

The only issue is, Dr. Q pointed out that the previous way I'd been calculating evaporation (the way suggested by Dr. D) did not have a surface area term in it, which can't be right because evaporation is highly dependent upon surface area. But I wrote it up nicely so I'll transpose my calculations here:

Using the ideal gas law P V = n R T and the partial pressure of chlorine, Par Press, standing in for P.
n evap = (Partial Pressure of Cl * Vol) / (R T (X - X1))
F = Par Press * (V / (R T (X-X1)))
where V / (R T (X-X1)) is a constant.
The flow rate, X1, and partial pressure can be varied (flow rate and X1 as design parameters, Partial Pressure being modified as more accurate estimates come in).

The equation F = Par Press * (V / (R T (X-X1))) can be solved given
V = 208 L
R = 0.08206 L atm / mol K
T = 298 K (or less, if you want to set it lower and make your tank colder. This is a design parameter)
X = 0.5 to 2 mg/L, 1.25 mg/L on average.
X1 < 0.04 mg/L
Par Press < 0.05 atm, not sure the exact value because I think the real fit line tapers out at low concentration, like a parabola's bowl.

Right, so, the question is now, how do I incorporate surface area into the evaporation term? Something tells me that should be important. It's too bad it's Friday, because all the professors have gone home now and aren't there for advice. I can probably ask a mass transport teacher next week.

*sighs* Fish are hard.

Edited by EricaWieser, 08 April 2011 - 04:58 PM.

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#15 Guest_EricaWieser_*

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Posted 08 April 2011 - 05:01 PM

Why not run the cold water through a submerged pipe as a heat exchanger rather than running it into the tank? Then you could keep the chlorine to zero.

I want to know how much chlorine would accumulate as a steady state concentration in the tank if I have an automatic water changer. It has been my plan for nearly a year now to build a tank with an automatic water changer in it, and a second wish of mine is to inexpensively have the capability to make aquarium water colder. It occurred to me that the cold water main is pretty cold, which would make inducing a winter period really easy if you just cranked up the flow rate from the cold water pipe to decrease the tank's temperature.

So since this technique has two different useful applications (water changer and cooler), I want to know whether or not it's feasible. What is the chlorine concentration in the tank, I wonder? And how high can the flow rate be before the chlorine concentration rises above 0.04 mg/L (the safe concentration given on http://www.isws.illi...I/ISWSRI-85.pdf )

With the way evaporation is being calculated now, I'd have to have 320 gal/min to get the chlorine up to 0.04 mg/L. But that's probably wrong because evaporation is being calculated without considering surface area.

Edited by EricaWieser, 08 April 2011 - 05:04 PM.

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#16 Guest_EricaWieser_*

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Posted 08 April 2011 - 05:13 PM

If you insist, though, an aeration source such as a sponge filter or just an airstone would help establish a lower equilibrium concentration, and it might help to have the inlet be a spray or drip rather than a flow from a submerged pipe.

Yes, that's true, and I'll keep it in mind if the chlorine ends up being too high in the tank. Can't tell now though, without a good estimate on evaporation.

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#17 Guest_mywan_*

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Posted 08 April 2011 - 05:31 PM

The analysis and graph looks ok for what is is but I am still trying to work through how you are conceptualizing its application. Without knowing this I can only guess where to begin. I am not very adept at chemistry, but the ideal gas law is fine. The primary issue is almost always a problem of application rather than validity.

In general the partial pressure or Cl2 concentration does not define any particular evaporation rate, but merely its potential to do so under various circumstances. Such as a sealed container in which evaporation does not occur irrespective of partial pressures or concentration of the components in it. But first I think we should rethink the assumption that the accumulation rate is zero, even though it technically is in some sense. Yet in some sense accumulation of anything is an illusion if you take the entire system. internal an external, into account. Accumulation is merely an arbitrary designation of a region or form.

So we have a tank of water which we can presume effectively has no Cl2. When we begin adding tap water Cl2 will begin to 'accumulate' at a rate equal to the total Cl2 volume and flow rate of the tap water. The overflow itself insures that the accumulation has a limiting factor equal the Cl2 input rate from the tap water, while the same limiting factor exist for Cl2 concentrations as a result of the source tap water concentrations. It is the fish tank we are concerned about Cl2 accumulation in, not the total Cl2 from both source (tap) and tank water. Given that if the only loss of Cl2 that accumulates in the tank is by way of the overflow it is easily demonstrated that the tank accumulation of Cl2 soon reaches a concentration limit equal to the concentration of Cl2 in the tap. What we are then concerned with is if the Cl2 loss rate, through reduction and/or evaporation, is great enough to prevent the accumulation (within the fish tank) from going over some small set percentage of the limit factor defined by the Cl2 concentrations in the tap.

Now back to partial pressures. In chemistry when all the reaction components are contained in the medium partial pressures are more directly applicable. The reduction and evaporation of Cl2 are much more strongly dependent on contact with an outside medium (atmosphere). If the system is 'enclosed', beyond the tap input and overflow, the tank concentrations of Cl2 will soon rise to equal the tap water irrespective of any partial pressure. Much like the Cl2 concentrations in the water lines delivering it. Hence only the partial pressures on the boundary between the liquid medium and an the external atmosphere are important. This includes the absorption of reducing agents from the atmosphere as well as evaporation to the atmosphere. Thus the maximum Cl2 accumulation in the tank depends on the ratio of the accumulation rate (provided from tap) and the total area of atmospheric contact. Which is why I suggested a drain board with an HVAC filter medium to greatly increase this surface area. The reactive area on which these partial pressures work must be limited to this area of atmospheric contact and the turbulence, temperature or mechanical, which speed the turnover of water in contact with the air.

A precise answer requires partial differential equations. But two graphs that shows the Cl2 increase rate in the tank with only drain loss of Cl2 at a given flow rate in one graph and another graph showing the Cl2 reduction/evaporation rate under a given surface area and turnover rate (turbulence) then where these two lines cross should be a reasonably close estimate of the maximum Cl2 concentration buildup in the aquarium. The first is pretty easy, the second is beyond me in any reasonable period of time. Instead of the math I generally go for overkill on the surface area/turbulence because any shortfall in Cl2 reduction/evaporation rates can be fixed by increasing it.

I hope this made sense?

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#18 Guest_mywan_*

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Posted 08 April 2011 - 05:37 PM

I see you came across the surface area problem while I was writing the last post. The pair of graphs I mentioned I think is the key, unless you want to get into partial differential equations. If you have some data on empirical Cl2 reduction rates under predefined conditions it would work as a starting point for the second graph.

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#19 Guest_mywan_*

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Posted 08 April 2011 - 05:58 PM

Here is another possible solution that can be done on a smaller scale while insuring sufficient surface air contact before going into the fish tank. Rig a float in a 5 gallon bucket, similar to a toilet float, that keeps the bucket full. Set this bucket at a somewhat higher level than the tank. Use a bubble column to lift water over the lip of or drain of the bucket to flow into the tank. From there it overflows in the manner described in the original link.

The bubble column insures that the water that overflows into the fish tank is limited to the water that was most intensely in contact with the reducing agents in the air. Since this is now going directly into the fish tank no other aerators are really needed. The water change rate may be fairly slow but in the long run far greater than the most extreme water change regimes.