will I get ferric phosphate if I mix micro/macro? - The Planted Tank Forum
 
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post #1 of 12 (permalink) Old 06-12-2015, 08:48 PM Thread Starter
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will I get ferric phosphate if I mix micro/macro?

At what ppm of phosphate and iron will I get ferric phosphate? I'm asking because I mix a week's worth of PPS Pro into my auto top off reservoir. I'm wondering if this causes the phosphate and iron to precipitate.

What ppm of Iron and Phosphate will cause precipitation?

I add enough to my 5 gallon reservoir to raise my reservoir to:

* 3.8 PPM of phosphate from KH2PO4
* 13.8 PPM of Iron from CSM+B (EDTA Chelated)


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Last edited by RisingSun; 06-13-2015 at 10:25 PM. Reason: more details
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post #2 of 12 (permalink) Old 06-13-2015, 05:04 PM
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Why so much iron? Iron is a trace element, even though it is an important trace element. It should be in a much lower concentration than phosphate, which is one of the 3 basic nutrients (NPK). With the concentration of iron being that high, you would lock up the phosphate, leaving a significant deficiency for the plants.

I know this doesn't answer your question. I don't know at what concentrations you get the chemical reaction that combines iron and phosphate.

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post #3 of 12 (permalink) Old 06-13-2015, 05:15 PM
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I'm no chemist so take what I say with a grain of salt. There are variables that can alter the stability of the nutrients. A week or so? Maybe. Six months? Probably not.

Since Plantex uses EDTA the PH would be a major factor when predicting the de-chelation of the elements, primarily iron. EDTA is more effective at lower PH ranges i.e. 5-6. Lowering the PH of the container would help the iron stay chelated longer. When iron is chelated it won't react with phosphorus so maintaining chelation should be the goal. Once the iron loses it's "protection" it will react with PO4 as well as several other things and fall out of solution. There are MANY other variables that can affect the reactions. However, this is most likely the largest IMO. Even a top notch chemist couldn't predict the stability over time without much more information.

Why is the ratio between PO4 and Iron so large? I would expect you to need more PO4 than iron.

If iron is of more importance I would use DTPA iron instead. It's much more stable and can withstand higher PH ranges we typically see in planted tanks.
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post #4 of 12 (permalink) Old 06-13-2015, 06:09 PM
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If you are using that much CSM+B then you are probably over dosing the other micros to the point of being toxic.

I would dose way less CSM+B and add Fe-DTPA to increase the iron only, but not to the high levels you are targeting. Usually about 1ppm is a reasonable target for Iron. Perhaps some plants would need a little more.

Dose them separately: N-P-K on one day, Micros (CSM+B and Fe) the other.
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post #5 of 12 (permalink) Old 06-13-2015, 06:38 PM
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That concentration of metals will pretty much kill all of your plants, fish and shrimp.
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post #6 of 12 (permalink) Old 06-13-2015, 07:01 PM
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Aside from the chemistry, I don't think you'll get regular and accurate dosing using your ato to do the work. Pretty much you're only dosing when your float switch engages. There's no measurement of how much of what you're really putting in there. Also the time it takes for the tank to evaporate until the point the float switch kicks on could be widely variable.
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post #7 of 12 (permalink) Old 06-13-2015, 07:04 PM
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Quote:
Originally Posted by Solcielo lawrencia View Post
That concentration of metals will pretty much kill all of your plants, fish and shrimp.
That's not the concentration in his tank. It's an automatic top-off reservoir. How many ppm that equates to in the tank? I have no idea. We need tank size and amount of solution being added to know that. My only guess why the PO4 and trace ratio is so far off is that his fish/food are providing PO4. Just a guess.
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post #8 of 12 (permalink) Old 06-13-2015, 08:12 PM
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Yes, you will precipitate out the iron phosphate. The Ksp for FePO4 is something like 10^-22, so your phosphate will be consumed (to a concentration on the order of a few ppb) and your iron will be at roughly 10ppm. Precipitation reactions are generally fairly fast, and here you'll wind up with some brownish-gray sediment in whatever container you put this in. It will be a mix of ferric phosphate and phosphates of whatever other metals you put in there. Most phosphates are insoluble in water.

EDTA stabilizes metal ions in solution, but it is in equilibrium with unbound ions and metal-free EDTA. The small amount of free iron reacts with phosphate in a very short period of time. As iron ions are removed from the solution, the Fe-EDTA equilibrium shifts farther and farther in the direction of free ions, continuously losing them to the phosphate.

Ask me how I know. (I'm a chemistry teacher; I should have known better...)
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post #9 of 12 (permalink) Old 06-13-2015, 10:22 PM Thread Starter
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Yes this is the ppm in my resevoir. Phosphate ppm in tank is 1ppm. I'm not sure why the iron is so high. I used https://sites.google.com/site/aquati...r/home/pps-pro which gave me 5g of KH2PO4 and 64g of TE. It doesn't say what TE to use, perhaps 64g is too high for CSM+B? I can take it easy on the trace.

Evaporation consistently consumes about 50% of my reservoir every week, so I just refill it all the way and add a week's dose of PPS Pro solutions.

Quote:
Originally Posted by jasonpatterson View Post
Yes, you will precipitate out the iron phosphate. The Ksp for FePO4 is something like 10^-22, so your phosphate will be consumed (to a concentration on the order of a few ppb) and your iron will be at roughly 10ppm. Precipitation reactions are generally fairly fast, and here you'll wind up with some brownish-gray sediment in whatever container you put this in. It will be a mix of ferric phosphate and phosphates of whatever other metals you put in there. Most phosphates are insoluble in water.

EDTA stabilizes metal ions in solution, but it is in equilibrium with unbound ions and metal-free EDTA. The small amount of free iron reacts with phosphate in a very short period of time. As iron ions are removed from the solution, the Fe-EDTA equilibrium shifts farther and farther in the direction of free ions, continuously losing them to the phosphate.

Ask me how I know. (I'm a chemistry teacher; I should have known better...)
Thanks, this is pretty valuable information - I couldn't find it anywhere on the internet. It seems like what you're saying is the phosphate will be completely depleted and the iron will be depleted by however much phosphate there is.

You mentioned the "Ksp for FePO4 is something like 10^-22". Does this mean if your phosphate is over 1 ppb, it will start binding with iron?? Doesn't this mean I will get FePO4 even if I dose the iron separately, as my tank is consistently at 1 ppm of phosphate?

Thanks!


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post #10 of 12 (permalink) Old 06-13-2015, 11:23 PM
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The ppm of iron and PO4 are the same for PPS-Pro, 0.1 ppm each day.

To raise Fe and PO4 to 0.1 ppm in 10 gallons use the amounts below.

Plantex 57.97 mg
KH2PO4 5.42 mg

A full weeks worth of PPS for 10 gallons...

405.79 mg Plantex
37.94 mg of KH2PO4

If you give me the tank volume your dosing I can give you the exact amount for one week or you can do the math from the amounts above.
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post #11 of 12 (permalink) Old 06-14-2015, 12:36 AM
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Quote:
Originally Posted by RisingSun View Post
You mentioned the "Ksp for FePO4 is something like 10^-22". Does this mean if your phosphate is over 1 ppb, it will start binding with iron?? Doesn't this mean I will get FePO4 even if I dose the iron separately, as my tank is consistently at 1 ppm of phosphate?

Thanks!
Mini chemistry lesson: Ksp is called the solubility product. It describes the extent to which a substance dissolves and is used primarily for substances that are barely soluble. In this case, FePO4(s) <--> Fe(aq) + PO4(aq). For this system,

Ksp = [Fe][PO4] = 1.3x10^-22

Where [Fe] is the concentration of iron(III) ions in units of molarity. I guessed at the concentration of free phosphate earlier, and I think I overestimated substantially. Your 3.8ppm is 0.00004M PO4, and 0.00025M Fe. If those two solutions are combined, you would expect them to react almost to completion. With some math you find that the final concentration of iron would essentially be 0.00025M and the concentration of phosphate would essentially be zero (5 parts in 10^17, by my math.)

Phosphate and iron will combine at any concentration, and the reaction will get really, really close to completion every time. There is no minimum concentration; they'll just keep on sticking to one another.

The way that we can get around that is by using dilute solutions of both and dosing at different times. The former slows down the precipitation reaction dramatically (and more so with an EDTA complex), so that what might take minutes in concentrate takes days in the tank. The latter gives the plants time to consume the ions before they can react.

Finally, and I'm not sure whether this is true for FePO4 or not, many compounds are difficult to crystallize because the tiny seed crystals they're based on are unstable. These crystals are only stable when there are enough of them to bond into an extended lattice, and with the tiny amounts we're using, that is not likely to occur. Again, I'm not sure whether this is the case for iron(III) phosphate or not. What I do know is that in the concentrations used in aquarium dosing solutions, the concentrations are more than high enough for iron(III) phosphate to precipitate.
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post #12 of 12 (permalink) Old 06-14-2015, 03:20 AM Thread Starter
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Amazing explanation! Thanks!

Quote:
Originally Posted by jasonpatterson View Post
Mini chemistry lesson: Ksp is called the solubility product. It describes the extent to which a substance dissolves and is used primarily for substances that are barely soluble. In this case, FePO4(s) <--> Fe(aq) + PO4(aq). For this system,

Ksp = [Fe][PO4] = 1.3x10^-22

Where [Fe] is the concentration of iron(III) ions in units of molarity. I guessed at the concentration of free phosphate earlier, and I think I overestimated substantially. Your 3.8ppm is 0.00004M PO4, and 0.00025M Fe. If those two solutions are combined, you would expect them to react almost to completion. With some math you find that the final concentration of iron would essentially be 0.00025M and the concentration of phosphate would essentially be zero (5 parts in 10^17, by my math.)

Phosphate and iron will combine at any concentration, and the reaction will get really, really close to completion every time. There is no minimum concentration; they'll just keep on sticking to one another.

The way that we can get around that is by using dilute solutions of both and dosing at different times. The former slows down the precipitation reaction dramatically (and more so with an EDTA complex), so that what might take minutes in concentrate takes days in the tank. The latter gives the plants time to consume the ions before they can react.

Finally, and I'm not sure whether this is true for FePO4 or not, many compounds are difficult to crystallize because the tiny seed crystals they're based on are unstable. These crystals are only stable when there are enough of them to bond into an extended lattice, and with the tiny amounts we're using, that is not likely to occur. Again, I'm not sure whether this is the case for iron(III) phosphate or not. What I do know is that in the concentrations used in aquarium dosing solutions, the concentrations are more than high enough for iron(III) phosphate to precipitate.


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