[Date Prev][Date Next][Thread Prev][Thread Next][Date Index][Thread Index]

Re: PO4 & Fe




> From: "Daniel Larsson" <defdac at hotmail_com>
> Subject: Re: PO4 and traces
> 
>> I just tried adding to 500ml of DI water, 5mls of Flourish, 1/8th teaspoon
>> of KH2PO4 and got nothing, no precipitate.
>> So I tried the tap water in case the mix reacts with the KH/GH, again
>> nothing. So I tried it Plantex trace mix , again same result.
> 
> That's weird. The PO4-precipitation is actually used to
> clean lakes:
> http://www.clearwatertechnologies.net/how_it_works/how.htm
> and it's easy to find more pages with google that explains
> the chemistry behind it.

And it's used for wastewater but you have to collect the precipitate to
remove it from the system. It doesn't go away without some form of removal.
You'll need a PO4 limited lake also for this to be effective.

Otherwise it is simply cycled back into the system during spring/fall
turnover.
In high aerobic conditions, you'll get precipitation, in anoxic conditions
along the substrate interface that occur commonly during summer
stratification, the reverse happens and much of Fe and PO4 is re released
back into the water column.

During summer stratification then following occurs:
Ferric iron is reduced to ferrous which is more soluble, freeing PO4.
Reduction usually occurs at substrate-water interface.
Fe concentration in hypolimnion(bottom layer) increases throughout
stratification(summer etc) but is practically undetectable in
epilimnion(upper water layers). The hypolimnion operates as Fe trap/sink.

At overturn:
The water column is oxygenated oxidizing Fe2+ to Fe3+ liberating PO4 which
is mixed throughout water column causing algal bloom The PO4 rapidly forms
FePO4 or becomes adsorbed to Fe(OH)3

We do this when we remove and uproot plants to some degree, not much I'd
imagine unless you do a very large hack/pruning. Also, DOC and mulm likely
binds a fair amount of PO4 and this is brought up when uprooting occurs.
Then there's Ca precipitation and good ol rust, Fe(OH)3.

So what about O2 then?
It'll precipitate out to form Fe(OH)3
There's also more sulfur in the form of SO4 than PO4 in most waters to
compete for the Fe.
What about Ca precipitation? Lots of Ca floating around.
There's more than simply one thing removing PO4 etc.
Roots of plants can and do use FePO4, Fe3(PO4)2.
They can add H+ ions till the the Fe is reduced. Rust, same thing.
Recycling of P: excreted by fish and zooplankton among phytoplankton.
Excreted P consists of about 50% orthophosphate PO4-P and rest as organic P.
In some cases zooplankton excretion may supply most all of daily
phytoplankton demand.

Rooted Macrophytes uptake of P occurs through leaves when water is rich in
PO4. Main method macrophytes obtain P is absorption of PO4 directly from
interstitial soil water(in that rich reduced place but the roots still add
O2 and H+ ions). Losses occur by direct excretion or death and
decomposition.
This is a pH dependent driven system and there are other competing reactions
besides Fe for PO4.

> Both iron (III) and phosphate are pH dependent. Phosphate chemistry
> derives from dissociation of phosphoric acid, a triprotic acid.
> FePO4 (S) _ Fe3+ + PO43-    KSO = 10 -26 = [Fe3+][PO4 3-]
> Iron (III) complexation:
> 1.    Fe3+ + H2O _ FeOH2+ + H+    K1 = 10 -2.2
> 
> 2.    FeOH2+ + H2O _ Fe(OH)2+ + H+    K2 = 10 -4.56
> 
> 3.    Fe(OH)2+ + H2O _ Fe(OH)3 (S) + H+    K3 = 10 2.76
> 
> 4.    Fe(OH)3 (S) + H2O _ Fe(OH)4- + H+    K4 = 10 -19
> Phosphate chemistry:
> 
> 5.    H3PO4 _ H2PO4 - + H+    K5 = 10 -2.2
> 
> 6.    H2PO4 - _ HPO4 2- + H+    K6 = 10 -7
> 
> 7.    HPO4 2- _ PO4 3- + H+    K7 = 10 -12
> Putting it all together:
> KSO = [Fe3+][PO4 3-] = (a0 [Fe (III)T]) (b3 PT)
 
> Therefore, the amount of phosphorus that can be present in equilibrium with
> FePO4 (S) at various pH values is:
>  a0 and b3 are functions of pH
>  Fe (III)T is either the maximum soluble concentration or the amount that was
> put into the system, whichever is less.
> Application:
> Phosphates are often a problem in lakes.  When phosphates are introduced,
> eutrophication (runaway algae growth), can occur because phosphorous is often
> the limiting nutrient in natural systems.  Therefore, there is increased focus
> on phosphorus removal in wastewater treatment processes.  One chemical method
> for phosphorus removal is iron addition to form iron-phosphate precipitant,
> which can be removed from the treated wastewater prior to discharge.

"The precipitant is removed prior to discharge." Other wise it'll be
recycled next turnover event. That is a key point.

> So I don't understand how you can't get any precipitation.
> More people that have tried this experiment?
> I have GH 6-8 and KH 2-3.

Same here. I'm just reporting I did the same thing and found nothing. Why?
Let's ask an inorganic chemist familiar with ETDA, and other binding agents.

Water chemistry of Fe is very complex based just on the inorganic component
alone.

Roger Miller and I went after this from two different angles and discussed
this in the past. I found that's very tough to say what is bioavailable
iron.
There's no test for that other than controlling everything/all else and
manipulating this one parameter. Keep DOC and detritus down(water
changes/vacuuming). Some 20 complexes form when Fe is add to water. A
difficult and complex issue.

Maybe an inorganic chemist, good with Fe, can go into the gory details.

Regards, 
Tom Barr

 
> // Daniel.