phosphate leaching
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cutegecko3
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Im sure its leached out by now.
Seems to me that "leaching" is based on how much has been absorbed over the yrs which is often determined on the water quality of that the rock has sat in. Phosphates tend to bind to calcium carbonate but its hard to predict what circumstances they would leach back into the water column .. think its fair to say that live rock which has a higher phosphate level than the water column would tend to "leach".
wayne in norway
New member
Can somebody explain how the chemistry for this absorbtion/leaching from tankwater works without gross shifts in pH? Ditto nitrate absorbtion/leaching. Live rock contains phosphate in the first place, I suspect a lot of so called leaching is just P being released as the carbonate material slowly dissolves.
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Can somebody explain how the chemistry for this absorbtion/leaching from tankwater works without gross shifts in pH?
There are many issues here.
Calcium carbonate will bind phosphate to its surface, just like phosphate binders do (GFO, aluminum oxide, etc).
The amount bound is a function of the phosphate concentration. The more in solution, the more binds.
As the phosphate concentration drops, phosphate will be released.
As the phosphate concentration rises, more will bind.
This is well documented for natural waters.
Ditto nitrate absorbtion/leaching.
Nitrate does not bind to calcium carbonate ins seawater. Organic materials stuck to the rock, like dead organism tissues, detritus,e tc, will break down and slowly release nitrogen and phosphorus compounds that can end up as nitrate and phosphate.
Live rock contains phosphate in the first place, I suspect a lot of so called leaching is just P being released as the carbonate material slowly dissolves.
There may be some of that, but it is not the only process.
There are many issues here.
Calcium carbonate will bind phosphate to its surface, just like phosphate binders do (GFO, aluminum oxide, etc).
The amount bound is a function of the phosphate concentration. The more in solution, the more binds.
As the phosphate concentration drops, phosphate will be released.
As the phosphate concentration rises, more will bind.
This is well documented for natural waters.
Ditto nitrate absorbtion/leaching.
Nitrate does not bind to calcium carbonate ins seawater. Organic materials stuck to the rock, like dead organism tissues, detritus,e tc, will break down and slowly release nitrogen and phosphorus compounds that can end up as nitrate and phosphate.
Live rock contains phosphate in the first place, I suspect a lot of so called leaching is just P being released as the carbonate material slowly dissolves.
There may be some of that, but it is not the only process.
kzooreefer
New member
I also have some questions as to if and when phosphate actually leaches out in aquariums because from what I understand desorption does not always occur. If the bond formed is through chemisorption it can't be broken by simply decreasing the concentration. For instance the State of Michigan Dept. of Environmental Quality uses alum (aluminum sulfate) to bind phosphate in lakes. The bound phosphates settle out at the bottom of the lake and do not leach back out. The same is supposed to be true of GFO and lanthanum chloride (PhosBuster), that the bond is so strong that phosphates also do not leached back out. GFO used for arsenic even says you can put it in a regular landfill after use as the arsenic won't leach out. It don't believe Severn Trent would make such a claim without having proof as they are one of the top company's performing environmental remediation of wastesites for the EPA under SARA (superfund). If on the other hand the bond is through physisorption which are weak Van der Waal forces then the bond can be broken easily, say by a drop in concentration. Now I'm not sure how phosphate is bound to live rock (calcium carbonate). It may be through physisorption which could then lead to leaching.
If the bond formed is through chemisorption it can't be broken by simply decreasing the concentration.
Why couldn't it? The binding to all such materials is reversible. You are in a sense saying that calcium phosphate is totally insoluble? That isn't the case. In some cases it is slow, but it is thermodynamically possible and certainly happens in the time periods that we are concerned with. I've shown it myself with GFO (bind phosphate, filter it off, then put it into fresh seawater and watch the phosphate desorb; IMO, I'v e totally debunked the claims made that GFO binding of phosphate is irreversible, and anyone can demonstrate it for themselves easily enough).
I also do nto think that what forms on the surface of calcium carbonate is pure calcium phosphate. It is a mixed crystal of calcium, magnesium, carbonate, phosphate, and likely organics and other ions. Phosphate can slip into and out of a hydrated layer on the surface a lot easier, IMO, than forming actual calcium phosphate crystals.
Here's Millero's article showing it is reversible on claicum carbonate:
Frank Millero1, Fen Huang1, Xiaorong Zhu1, Xuewu Liu1 and Jia-Zhong Zhang2
(1) University of Miami/RSMAS, 4600 Rickenbacker Cswy, Miami, FL 33149, U.S.A
(2) AOML, National Oceanic and Atmospheric Administration, 4301 Rickenbacker Cswy Miami, FL, 33149, U.S.A
Abstract The adsorption and desorption of phosphate on calcite and aragonite were investigated as a function of temperature (5â€"œ45 °C)and salinity (0â€"œ40) in seawater pre-equilibrated with CaCO3. An increase in temperature increased the equilibrium adsorption; whereas an increase in salinity decreased the adsorption. Adsorption measurements made in NaCl were lower than the results in seawater. The higher values in seawater were due to the presence of Mg2+ and Ca2+ ions. The increase was 5 times greater for Ca2+ than Mg2+. The effects ofCa2+ and Mg2+ are diminished with the addition of SO4 2- apparently due to the formation of MgSO4 and CaSO4 complexes in solution and/or SO4 2- adsorption on the surface of CaCO3. The adsorbed Ca2+ and Mg2+ on CaCO3 (at carbonate sites) may act as bridges to PO4 3- ions. The bridging effect of Ca2+is greater than Mg2+ apparently due to the stronger interactions of Ca2+ with PO4 3-.
The apparent effect of salinity on the adsorption of PO4 was largely due to changes in the concentration of HCO3 - in the solutions. An increase in the concentration of HCO3 - caused the adsorption of phosphate to decrease, especially at low salinities. The adsorption at the same level of HCO3 - (2 mM) was nearly independent of salinity. All of the adsorption measurements were modeled empirically using a Langmuir-type adsorption isotherm
[ [PO4]ad = KmCm[PO4]T/(1 +Km [PO4]T) , ]
where [PO4]ad and [PO4]T are the adsorbed and total dissolved phosphate concentrations, respectively. The values of Cm (the maximum monolayer adsorption capacity, (mol/g) and Km (the adsorption equilibrium constant, g/(mol) over the entire temperature (t, °C) and salinity (S) range were fitted to
[ Cm = 17.067 + 0.1707t - 0.4693S + 0.0082S2 ( = 0.7) ]
[ ln Km = - 2.412 + 0.0165t - 0.0004St - 0.0008S2 ( = 0.1) ]
These empirical equations reproduce all of our measurements of[PO4]ad up to 14 mol/g and within ±0.7 mol/g.
The kinetic data showed that the phosphate uptake on carbonate minerals appears to be a multi-step process. Both the adsorption and desorption were quite fast in the first stage (less than 30 min) followed by a much slower process (lasting more than 1 week). Our results indicate that within 24 hours aragonite has a higher sorption capacity than calcite. The differences between calcite and aragonite become smaller with time. Consequently, the mineral composition of the sediments may affect the short-term phosphate adsorption and desorption on calcium carbonate. Up to 80 % of the adsorbed phosphate is released from calcium carbonate over one day. The amount of PO4 left on the CaCO3 is close to the equilibrium adsorption. The release of PO4 from calcite is faster than from aragonite. Measurements with Florida Bay sediments produced results between those for calcite and aragonite. Our results indicate that the calcium carbonate can be both a sink and source of phosphate in natural waters.
Why couldn't it? The binding to all such materials is reversible. You are in a sense saying that calcium phosphate is totally insoluble? That isn't the case. In some cases it is slow, but it is thermodynamically possible and certainly happens in the time periods that we are concerned with. I've shown it myself with GFO (bind phosphate, filter it off, then put it into fresh seawater and watch the phosphate desorb; IMO, I'v e totally debunked the claims made that GFO binding of phosphate is irreversible, and anyone can demonstrate it for themselves easily enough).
I also do nto think that what forms on the surface of calcium carbonate is pure calcium phosphate. It is a mixed crystal of calcium, magnesium, carbonate, phosphate, and likely organics and other ions. Phosphate can slip into and out of a hydrated layer on the surface a lot easier, IMO, than forming actual calcium phosphate crystals.
Here's Millero's article showing it is reversible on claicum carbonate:
Frank Millero1, Fen Huang1, Xiaorong Zhu1, Xuewu Liu1 and Jia-Zhong Zhang2
(1) University of Miami/RSMAS, 4600 Rickenbacker Cswy, Miami, FL 33149, U.S.A
(2) AOML, National Oceanic and Atmospheric Administration, 4301 Rickenbacker Cswy Miami, FL, 33149, U.S.A
Abstract The adsorption and desorption of phosphate on calcite and aragonite were investigated as a function of temperature (5â€"œ45 °C)and salinity (0â€"œ40) in seawater pre-equilibrated with CaCO3. An increase in temperature increased the equilibrium adsorption; whereas an increase in salinity decreased the adsorption. Adsorption measurements made in NaCl were lower than the results in seawater. The higher values in seawater were due to the presence of Mg2+ and Ca2+ ions. The increase was 5 times greater for Ca2+ than Mg2+. The effects ofCa2+ and Mg2+ are diminished with the addition of SO4 2- apparently due to the formation of MgSO4 and CaSO4 complexes in solution and/or SO4 2- adsorption on the surface of CaCO3. The adsorbed Ca2+ and Mg2+ on CaCO3 (at carbonate sites) may act as bridges to PO4 3- ions. The bridging effect of Ca2+is greater than Mg2+ apparently due to the stronger interactions of Ca2+ with PO4 3-.
The apparent effect of salinity on the adsorption of PO4 was largely due to changes in the concentration of HCO3 - in the solutions. An increase in the concentration of HCO3 - caused the adsorption of phosphate to decrease, especially at low salinities. The adsorption at the same level of HCO3 - (2 mM) was nearly independent of salinity. All of the adsorption measurements were modeled empirically using a Langmuir-type adsorption isotherm
[ [PO4]ad = KmCm[PO4]T/(1 +Km [PO4]T) , ]
where [PO4]ad and [PO4]T are the adsorbed and total dissolved phosphate concentrations, respectively. The values of Cm (the maximum monolayer adsorption capacity, (mol/g) and Km (the adsorption equilibrium constant, g/(mol) over the entire temperature (t, °C) and salinity (S) range were fitted to
[ Cm = 17.067 + 0.1707t - 0.4693S + 0.0082S2 ( = 0.7) ]
[ ln Km = - 2.412 + 0.0165t - 0.0004St - 0.0008S2 ( = 0.1) ]
These empirical equations reproduce all of our measurements of[PO4]ad up to 14 mol/g and within ±0.7 mol/g.
The kinetic data showed that the phosphate uptake on carbonate minerals appears to be a multi-step process. Both the adsorption and desorption were quite fast in the first stage (less than 30 min) followed by a much slower process (lasting more than 1 week). Our results indicate that within 24 hours aragonite has a higher sorption capacity than calcite. The differences between calcite and aragonite become smaller with time. Consequently, the mineral composition of the sediments may affect the short-term phosphate adsorption and desorption on calcium carbonate. Up to 80 % of the adsorbed phosphate is released from calcium carbonate over one day. The amount of PO4 left on the CaCO3 is close to the equilibrium adsorption. The release of PO4 from calcite is faster than from aragonite. Measurements with Florida Bay sediments produced results between those for calcite and aragonite. Our results indicate that the calcium carbonate can be both a sink and source of phosphate in natural waters.
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T Man
~PPPPPPP~
Thanks everyone for your informative replies! From my interpretation, please tell me if this is correct.
- The LR will hold the phosphate level of the water column it's in.
- If the WC's levels drop/rise , the levels in the rock follow suit.
- This happens over a very short period of time, not months or years.
Now what about a mature DSB...... does it behave in the same manner? T
- The LR will hold the phosphate level of the water column it's in.
- If the WC's levels drop/rise , the levels in the rock follow suit.
- This happens over a very short period of time, not months or years.
Now what about a mature DSB...... does it behave in the same manner? T
kzooreefer
New member
Cool, good stuff in there Randy. It helps me out with a problem at work with an iron-calcium media we are looking at for arsenic. I've been using phosphate as a kind of serrogate as my graphite furnace bit the dust and I was getting desorption of phosphate fairly rapidly. Which makes me worry about arsenic desorption. I have been comparing our medium side by side to GFO but didn't run the tests long enough so far to see desoprtion from the GFO (1 pound of GFO at 0.25 gpm for 500 gallons at 1.5 ppm PO4). I'm going to run another PO4 test (all my tests following NSF Standard 53) and maybe this time I will get to run the GFO to exhaustion and then see if PO4 desorbs. I'll do an extraction in the middle of testing to see if PO4 desorbs during an 8 hour stagnant period when the filters are resting thats part of NSF protocols anyways.
wayne in norway
New member
Excellent answers (again) Randy, thank you.
Thanks. 
This happens over a very short period of time, not months or years.
Now what about a mature DSB...... does it behave in the same manner?
The top of the sand bed will be have the same. Deeper portions may be different. If the phosphate is in less accessible areas, then it will be slower. Like deep in sand beds or in pores in live rock. Other processes will also happen over longer periods, such as actual dissolution of CaCO3 in sand beds.
The degradation of organics in sand beds is another complicating factor, and that can happen over various time periods.
This happens over a very short period of time, not months or years.
Now what about a mature DSB...... does it behave in the same manner?
The top of the sand bed will be have the same. Deeper portions may be different. If the phosphate is in less accessible areas, then it will be slower. Like deep in sand beds or in pores in live rock. Other processes will also happen over longer periods, such as actual dissolution of CaCO3 in sand beds.
The degradation of organics in sand beds is another complicating factor, and that can happen over various time periods.
kzooreefer
New member
I did a little more reading up on adsorption. The type of adsoprtion that occurs at ambient temperatures is normally physical adsorption and can be reveresed by placing the adsorbate in "pure water". I'm not sure how much desorption would occur if you used any thing other than DI water but I would imagine it to be substaintally less. This is why GFO used for arsenic removal can be disposed of in a landfill. Chemical adsorption on the other hand occurs at high temperatures and, I'll use this wording again, is irreversible. Chemical adsorption is not precipitation so it is not ruled by the solubility of the product but it doesn't normally occur in the situtations we are discussing so that is kind of a mute point. These same prinicpals hold true for GAC, you can get the organics to desorb by placing it in DI water. Though to truelly reactivate it requires high temperature steam. It would seem to me that from the article Randy submitted that PO4 leaching from live rock could be influenced more by the solubility of calcium carbonate monohydrate in which it is bonded.
Cooney, D.O., 1998. Adsorption Design for Wastewater Treatment. Lewis Pub., NY. Pgs 28-29.
Cooney, D.O., 1998. Adsorption Design for Wastewater Treatment. Lewis Pub., NY. Pgs 28-29.
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Well, the experiments I did with GFO were adsorption from artificial seawater, and then desorption in artificial seawater. I allowed 24 h to bind, and then tracked desorption.
No. It binds to minerals and specially made polymeric resins, but not appreciably to plastics or glass or adhesives/sealants.
You're welcome.
Happy reefing.
Happy reefing.
kzooreefer
New member
Now I'm wondering if the GFO in my canister filter is leaching phosphates? What I would hope is that a bed of GFO undisturbed in a canister wouldn't desorb phosphate to any appreciable degree. It is supposed to eventually reach an equillibrium where as much is adsorbed as desorbed so no net gain in seen either way. So if it in fact desorbs more this would really surprise me but right now I don't know for a fact it isn't. In the past I used to clean GFO by removing it from the canister and swirlling it in a bucket of seawater and I do get phosphates off when I do this. Which is expected as this is akin to backwashing a fixed bed media. The swirlling action and grain to grain scrubbing is enough to break the bond holding the phosphate to the GFO. Thats nothing new. But if it does in fact leach phosphates while it is still in the canister running that makes me worry. Then I'd have to ask myself what is even the point of using it? This is part of the reason I am running it in cartridges at work, to see what happens when GFO gets fully loaded. Does it leach and how much?
I'm also toying around with going with macro-algae as a way to control phosphates. I don't use a sump or refugium but have been adding turtle grass to my display and want to add some shaving brush plants. I may take the GFO out of the canister now and see what happens.
I'm also toying around with going with macro-algae as a way to control phosphates. I don't use a sump or refugium but have been adding turtle grass to my display and want to add some shaving brush plants. I may take the GFO out of the canister now and see what happens.
The three ways that I can see GFO release phosphate in a real aquarium use setting are:
1. When the phosphate concentration has dropped below the level required in solution to bind that much. That would only likely happen if something else was also lowering it, like algae or bacteria taking it up. In that case, leaving GFO in the tank with a big load of phosphate already on it might be counter productive. For example, if you started with a tank at 0.25 ppm phosphate, used GFO to take it to 0.05 ppm, and then drove bacteria to take it down to 0.01 ppm. That last step might cause some release of phosphate.
Note that if you use GFO to take phosphate to undetectable levels, and it stays there, then there is not much concern, because you have not "overloaded" it with phosphate by binding an excessively large amount due to excessive phosphate concentrations.
2. When something else in the water is slowly displacing it. For example, certain organic species that might be fairly uncommon by are strongly bound to GFO might displace the phosphate as they arrived at and took up binding sites.
3. If the ph dropped enough to tip the balance toward less absorbed phosphate for the same concentration in solution. That is, lower than when it initially bound. Diurnal ph cycles would probably not do much, but a long term trend or spike downward might help release it.
So if it in fact desorbs more this would really surprise me but right now I don't know for a fact it isn't. In the past I used to clean GFO by removing it from the canister and swirlling it in a bucket of seawater and I do get phosphates off when I do this. Which is expected as this is akin to backwashing a fixed bed media. The swirlling action and grain to grain scrubbing is enough to break the bond holding the phosphate to the GFO.
In the experiments I did the GFO was in a packed bed and I don't think it was moving around appreciably.
To be honest, I'd be a bit surprised if ordinary grain to grain rubbing broke many such surface chemical bonds, as opposed to breaking off whole bits of GFO through mechanical cracking sorts of actions.
1. When the phosphate concentration has dropped below the level required in solution to bind that much. That would only likely happen if something else was also lowering it, like algae or bacteria taking it up. In that case, leaving GFO in the tank with a big load of phosphate already on it might be counter productive. For example, if you started with a tank at 0.25 ppm phosphate, used GFO to take it to 0.05 ppm, and then drove bacteria to take it down to 0.01 ppm. That last step might cause some release of phosphate.
Note that if you use GFO to take phosphate to undetectable levels, and it stays there, then there is not much concern, because you have not "overloaded" it with phosphate by binding an excessively large amount due to excessive phosphate concentrations.
2. When something else in the water is slowly displacing it. For example, certain organic species that might be fairly uncommon by are strongly bound to GFO might displace the phosphate as they arrived at and took up binding sites.
3. If the ph dropped enough to tip the balance toward less absorbed phosphate for the same concentration in solution. That is, lower than when it initially bound. Diurnal ph cycles would probably not do much, but a long term trend or spike downward might help release it.
So if it in fact desorbs more this would really surprise me but right now I don't know for a fact it isn't. In the past I used to clean GFO by removing it from the canister and swirlling it in a bucket of seawater and I do get phosphates off when I do this. Which is expected as this is akin to backwashing a fixed bed media. The swirlling action and grain to grain scrubbing is enough to break the bond holding the phosphate to the GFO.
In the experiments I did the GFO was in a packed bed and I don't think it was moving around appreciably.
To be honest, I'd be a bit surprised if ordinary grain to grain rubbing broke many such surface chemical bonds, as opposed to breaking off whole bits of GFO through mechanical cracking sorts of actions.
kzooreefer
New member
From what information I've been able to find the bond between GFO and phosphate is more than likely a phsysical bond, weak Van der Walls forces, that are easily broken. Severn Trent recommends a backwash rate for GFO of 9 to 11 gpm per square foot but says nothing about what is realesed during backwashing. They just recommend it to prevent compaction of the bed. This is an extrememly slow rate which if it did release phosphate would be a problem in the upflow phosphate reactors [ 9 gpm/sq ft is about 0.4 gpm (22 gph) ]. Don't know if anyone has ever checked their reactor to see if over time it starts to release phosphates. You would think that there would be a thread on it by now if it was happening.
I'm not all that impressed with GFO anyways and PO4 has been the one thorn in my side. This has been a really good discussion thanks Randy for all the information.
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