View Full Version : Phosphate binding using Sr(NO3)2
dartier
09/09/2014, 04:28 PM
I have a question related to the possibility to bind phosphate into SrP (Orthophosphate).
I happened on a thread (http://reefcentral.com/forums/showthread.php?t=2381987) about the a Natureef Denitrator DIY build. In the discussion it was mentioned that the addition of Sr(NO3)2 into the anaerobic reaction chamber would cause the bacteria media to scavenge the Nitrogen and the Strontium would then bind the Phosphate leaving SrP as a result.
This got me to thinking, I am also running a bacteria driven reactor, for the purpose of lowering phosphate. I dose nitrate salts (Mg,Ca,K) currently into the system (along with Carbon) to spur the growth of the bacteria for the purpose of export.
Do you think that I could subsitute Strontium Nitrate - Sr(NO3)2 instead of my existing nitrate sources, and would the bacteria be able to access the nitrate. The most important question is, would the same binding occur (Sr(NO3)2 -> SrP) in an aerobic environment?
My thinking is that since the bacteria in the anaerobic process use up the oxygen first, making the Nitrogen available in the process, that this behaviour may be missing when running aerobically.
Anyway I am not a chemist, so I thought I would pose the question to the people who are.
Thanks,
Dennis
bertoni
09/09/2014, 04:38 PM
Do you mean Sr<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>? I suppose some might form. I don't know how much, though. I wouldn't want to spike the strontium level in a tank very much, so I'd be very careful.
shermanator
09/09/2014, 04:52 PM
Strontium phosphate is quite insoluble in water so it would work. As bertoni said, you'll be spiking your Sr high. I didn't re look at that thread but my recollection is that the natureef is done offline? You could carefully dose strontium nitrate based on phosphate amount (use sub-stoichiometric amount, bertoni has the right formula) and it might be safe.
It might be a reasonable alternative to lanthanum chloride (?) but while I've done that to recycle rock, I'm not sure I would do that in a full blown reef tank either.
disc1
09/09/2014, 05:12 PM
I wouldn't want to spike the Sr level in my tank. It can get incorporated into coral skeleton where it disrupts normal formation. That happens naturally at the normal level of Sr. I wouldn't want to try to cause more of that.
You're talking about this post?
The phosphagone in the natureef product uses strontium nitrate to remove phosphates. I still havent figured out the concentration of their solution.
Strontium Nitrate is widely used and can be purchased in the pure powder form. It burns crimson red and is used in road flares, sparklers and pyrotechnic shows.
The reaction:
Without delving into the chemistry equations.....You add the strontium Nitrate to the reaction chamber. When the chamber goes anoxic and the bacteria start using Nitrogen, the strontium gets removed from the Nitrate and then binds to phosphate. The Strontium Phosphate precipitates out of solution.
I'm not sure if it just falls to the bottom of the chamber or if it also get skimmed out after you flush. I have a salifert Sr test kit and the concentration of Sr in my tank is normal. I did have my water tested (aquariumwatertesting.com) back in Nov 2013 and my strontium levels were in the normal range but on the higher side of the range. Again, I've been using the phosphagone for over 2 years.
In my experience GFO works better to remove phosphates than their phosphagone product. However, I recently ordered some SrNO3 powder and will be experimenting with different concentrations to remove phosphates when added to the reaction chamber.
dartier
09/09/2014, 05:48 PM
You're talking about this post?
Yes, that is the post that gave me the idea of perhaps being able to get to my end goal, lowering PO4, without having to rely on dosing nitrate, as much as I currently do.
I was worried about strontium being elevated as result, but if I am understanding shermanator, as long as I do not exceed the amount of PO4 that is available to bind to the strontium, then the risks would be less, though not without risk.
If I am going to do this, I will need to get a strontium test ...
Thanks for the input guys.
Dennis
dartier
09/09/2014, 07:33 PM
Thinking this through (forgive me but I have never studied chemistry), in layman's terms, is this what occurs:
(Sr(NO3)2)3 = (Sr)3 + (NO3)6 + (PO4)2 = Sr3(PO4)2 + (NO3)6
... with the 6 nitrate ions being used by the bacteria, and the 2 phosphate ions binding with 3 strontium ions.
What has my confused is that in randy's article on Strontium he mentions that most of the strontium in seawater exists as Sr2+, so why do those strontium free ions not bind with the PO4 in our tanks on their own?
Obviously they do not as phosphate would not be such a problem for so many of us, but I am wondering why this is not the case?
Dennis
disc1
09/10/2014, 08:22 AM
The reason you still have Sr in your seawater despite the low solubility of Sr3(PO4)2 is because that is an equilibrium reaction. That means the reaction you are talking about goes both ways. The Sr phosphate also dissociates to free Sr and phosphate. The solubility product defines how much will stay free and how much will stay bound.
disc1
09/10/2014, 08:26 AM
I can't find the ksp for strontium phosphate very easy and don't really have time to search. Maybe Randy has it.
dartier
09/10/2014, 09:15 AM
The reason you still have Sr in your seawater despite the low solubility of Sr3(PO4)2 is because that is an equilibrium reaction. That means the reaction you are talking about goes both ways. The Sr phosphate also dissociates to free Sr and phosphate. The solubility product defines how much will stay free and how much will stay bound.
Thanks David. That sounds ominous, that the PO4 can dissociate from the Sr, which would be about the worst case scenario I can think of. That would both provide the PO4 to problematic consumers (algae), and spike the srontium in the process. Sounds like a dead end. Would this describe it, PO4 behaves with Sr, the same as it does when bound to calcium carbonate? Meaning it will leach until it reaches equilibrium, so it is not really made unavaialble as long as it is in the system ... just unavailable to be the tested for.
Is this what you were looking for in your later message:
Sr3(PO4)2 Ksp = [Sr--]^3 [PO4+++]^2
Ksp = (3x)^3(2x)^2 = 108x^5
1.00×10^-31 = 108x^5; x = 2.48x10^-7 M.
MW = 357.8 g/mole
(2.48x10^-7) (357.8) / (0.1) = 8.87x10^-4 g/100 mL.
(I did not work that out, I simply copied it from another site) :hmm6:
Dennis
shermanator
09/10/2014, 09:22 AM
I can't find the ksp for strontium phosphate very easy and don't really have time to search. Maybe Randy has it.
According to an old CRC I have, it is 1 x 10^-31 for Sr3(PO4)2. For comparison, LaPO4's Ksp is 1 x 10^-25.
It's been a long time since I have done these types of calculations, but 10 ppm Sr2+ (from Randy's article, that is an average found in aquaria) is ~113 µmol/L or 1.13 x 10^-4 M.
For phosphate, let's say you have 0.09 ppm PO4, that is 9.5 x 10^-5 M.
So... if you have 1 x 10-4 M Sr2+ and you have 9.5 x 10^-5 M, the solubility quotient (Q) is 1.1 x 10^-8. Q is > Ksp for Sr3(PO4)2 and it should form a precipitate.
I'm not sure what exactly this means or how the free Sr2+ concentrations are being measured. If Sr2+ is in equilibrium with something else (incorporation into coral skeleton, perhaps via interaction with calcium carbonate), then the free concentration of Sr2+ might be << 10 ppm.
shermanator
09/10/2014, 09:24 AM
Thanks David. That sounds ominous, that the PO4 can dissociate from the Sr, which would be about the worst case scenario I can think of.
Everything is reversible. On what timescale is the question... For example, calcium carbonate rock dissolves in water (just very, very, very slowly).
Randy Holmes-Farley
09/10/2014, 10:48 AM
The Ksp of strontium phosphate is less than calcium phosphate, so it is less soluble, but not by as much as the concentration of strontium is lower in seawater.
So strontium phosphate is more undersaturated (or less supersaturated) than is calcium phosphate under normal reef conditions, but not by a lot.
Remember, Ksp for these is [Sr++]^3 x [PO4]^2 so the concentration difference between Sr and Ca is raised to the third power!
It seems like an odd method to me.
shermanator
09/10/2014, 11:48 AM
Remember, Ksp for these is [Sr++]^3 x [PO4]^2 so the concentration difference between Sr and Ca is raised to the third power!
Even then, you have a Q = 1.3 e20 for Sr3(PO4)2 using 10 ppm Sr2+ and 0.09 ppm PO4. With a Ksp of 10^-31, it should essentially all be precipitate. So it should work... in theory.
I agree it's an odd method. I wouldn't do it in a reef, but I suppose for offline use like in a Natureef, it would work. But even in the offline Natureef case, why not use lanthanum? Lanthanum would be cheaper and probably has fewer side effects. Perhaps one benefit of strontium (vs lanthanum) is the ability to test for it prior to adding the processed water back into your reef.
dartier
09/10/2014, 11:56 AM
It seems like an odd method to me.
Yes, it does seem like a stretch in my particular application that it would be a benefit. This is however what is used with the Natureef denitrators with successful results. The user in the the thread I posted earlier (DrThompson) has been able to keep his PO4 in line using this method of dosing (albeit with the Natureef denitrator, not how I am suggesting it be used).
DrThompson is discharging his reactor directly into his skimmer to export the bacteria at the earliest possible point, and because the effulent from the reactor is very low in oxygen. In his photos, the reactor contents appears almost thick, so I can see why this would be a good idea (to discharge directly to the skimmer).
Do you think that the Sr(PO4)2 that should be present in the discharge is being skimmed? Kind of like the suggestion that dosing Kalwasser seems to lower PO4 for some people. The expectation that this has something to do with the bound PO4 being skimmed out after binding to Calcium.
Dennis
Randy Holmes-Farley
09/10/2014, 12:04 PM
Even then, you have a Q = 1.3 e20 for Sr3(PO4)2 using 10 ppm Sr2+ and 0.09 ppm PO4. With a Ksp of 10^-31, it should essentially all be precipitate. So it should work... in theory.
The values in seawater are a lot lower than can be calculated that way for several reasons:
1. It is the concentration of PO4--- that are used, not total phosphate. Maybe 20% of the total, depending on pH. Less at low pH in a denitrator.
2. PO4--- will be heavily ion paired in seawater, reducing its activity way further. Only about 0.2% is free phosphate. 73% is ion paired to calcium and 27% to magnesium.
3. The activity coefficient of Sr++ in seawater is only about 0.25 (same for calcium) so one must adjust for these effects by effectively dropping the concentration by a factor of 4.
All these things tend to lower the saturation state.
That said, calcium phosphate will be more likely to precipitate than strontium phosphate in sea water. So why does adding strontium help and calcium didn't already precipitate? I'm not sure it does. Perhaps there's a kinetic issue, but I doubt it. :)
Randy Holmes-Farley
09/10/2014, 12:06 PM
This is however what is used with the Natureef denitrators with successful results.
And not adding Sr into the same system has been shown to result in substantially higher phosphate?
dartier
09/10/2014, 12:38 PM
And not adding Sr into the same system has been shown to result in substantially higher phosphate?
Hmm, that would make for an interesting way of testing this hypothesis (that the addition of Sr(PO4)2 is having an effect).
Dennis
disc1
09/10/2014, 12:48 PM
Hmm, that would make for an interesting way of testing this hypothesis (that the addition of Sr(PO4)2 is having an effect).
Dennis
It's called the scientific method. It seems to have been lost in recent years.
shermanator
09/10/2014, 01:12 PM
This is however what is used with the Natureef denitrators with successful results.
And not adding Sr into the same system has been shown to result in substantially higher phosphate?
I also wondered this. It could be entirely a nitrate effect (dosing nitrate increases bacterial growth (and that growth requires phosphate)).
It's been a while since I read that thread, but I don't recall anyone tried dosing a different nitrate salt.
dartier
09/10/2014, 01:26 PM
I also wondered this. It could be entirely a nitrate effect (dosing nitrate increases bacterial growth (and that growth requires phosphate)).
It's been a while since I read that thread, but I don't recall anyone tried dosing a different nitrate salt.
Interesting idea. Because the (Natureef) unit uses an offline method to process a batch, there would only be an absolute amount of available nitrate in the batch. So adding the extra nitrate source could be essentially altering the ratio between NO3/PO4 and there by using up the PO4?
I would never thought of that ... even though that is what I am essentially doing myself in my online method (dosing nitrate to try and grown bacteria and lower PO4 in the process).
That would be ironic if that what was occuring, since that is my current method.
Dennis
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