Sciencemadness Discussion Board - FePO4 blue !?

The trouble is that there really isn’t anything in there that can deprotonate HPO4(2-)

If FeNH4PO4 is sufficiently insoluble, then

Fe²⁺ + NH₄⁺ + 2 HPO₄^2- --> FeNH₄PO₄(s) + H₂PO₄^-

will be favourable.

Anyone know the Ksp of FeNH₄PO₄?

If we use K2 = 7e-8 and K3 = 4e-13 for phosphoric acid, then the disproportionation reaction would have a Keq of 5.7e-6. So if the Ksp for FeNH₄PO₄ is 6e-6 or less (and most Ksp values are), then this reaction will have a Keq of greater than one.

[Edited on 15-3-2013 by DraconicAcid]

[Edited on 15-3-2013 by DraconicAcid]

rstar - 15-3-2013 at 18:03

That means my BAP is contaminted and i made an iron complex. But what kind of iron complex could have bluish color ?

rstar - 15-3-2013 at 19:32

That thing now looks greyish black. Here's a photo

DraconicAcid - 15-3-2013 at 19:52

I would assume that your wet precipitate is oxidizing, and you're getting a mixed iron oxide/phosphate.

AJKOER - 16-3-2013 at 08:22

OK, some related research (see bottom of page 11 at http://click.infospace.com/ClickHandler.ashx?du=http%3a%2f%2... ). To quote:

≡FeOH0 + PO43- + 3H+ ↔ ≡FeH2PO4 + H2O log K1 = 32.08 (1)
≡FeOH0 + PO43- + 2H+ ↔ ≡FeHPO4- + H2O log K2 = 26.39 (2)
≡FeOH0 + PO43- + H+ ↔ ≡FePO42- + H2O log K3 = 20.73 (3)
12
How much ferric is needed to remove phosphorus?
Traditionally the ferricphosphate precipitant FePO4(s) was thought to be the primary precipitant in wastewater treatment. However, recent research (Smith et al., 2008) showed a complex structure, with several fundamental reactions occurring simultaneously as iron reacts to form ferric hydroxides and phosphates.
The amount of phosphate that binds to the metal hydroxide is still a topic of discussion. Current research suggest that the stoichiometric ratio of metal:phosphorus (Me: P) in the precipitant depends on many factors, including the phosphate concentration in the liquid, chemical dose, age of the hydroxyl complex, mixing, and many other factors (Szabo et al., 2008, de Haas et al., 2000, Yang et al., 2006)."

Now, this somewhat current (2008) 'topic of discussion' among researchers is before anyone has introduce NH4+. With respect to the precise chemistry involved in the cited procedure, I am even more indefinite without addressing impurities.

blogfast25 - 16-3-2013 at 10:38

I doubt strongly that the conditions decribed in that infospace paper can apply to what 'rstar' was doing.

These 'FeOH values' for instance are for surface reactions on hydrated ferric oxide. Not greatly relevant here, IMHO...

'rstar' should repeat his experiment using stoichiometric amounts of reagents and assuming FeNH4PO4 is formed. Then see what colour he gets.

Assuming FeNH4PO4 was precipitated and that there was an excess Fe(II) available, the latter would have started to oxidise to Fe(III). Having Fe in two different oxidation states could possibly lead to strange results (see e.g. magnetite and Berlin Blue).

[Edited on 16-3-2013 by blogfast25]

rstar - 17-3-2013 at 00:21

Hey ! now the colour actually looks greyish-green

blogfast25 - 17-3-2013 at 06:30

Hey ! now the colour actually looks greyish-green

You mean the blue has now changed colour to greyish-green? Did any drying take place?

rstar - 17-3-2013 at 07:22

Hey ! now the colour actually looks greyish-green

You mean the blue has now changed colour to greyish-green? Did any drying take place?

Oh yes...
There was drying...
I guess it may be a Fe - NH₄ complex.

DraconicAcid - 17-3-2013 at 10:50

I guess it may be a Fe - NH₄ complex.

Can't be. Ammonia can form complexes with transition metals because it has a lone pair it can donate to the metal. The ammonium ion doesn't have a lone pair, and can't act as a Lewis base (i.e., a ligand).

ElectroWin - 17-3-2013 at 11:11

any chance there is a bit of copper in that fertilizer?

rstar - 17-3-2013 at 21:34

What if this is FeNH4PO4, that also looks greenish, right?
is there any test for it ?

blogfast25 - 18-3-2013 at 06:37

It should dissolve in concentrated sulphuric acid, displacing the weaker phosphoric acid, into a greenish solution of FeSO4 and (NH4)2SO4. But that's not a definitive test...

rstar - 18-3-2013 at 08:43

I havnt stored it in a vessel yet, it's outside in the watch glass. Does it react with O2 in air ?

blogfast25 - 18-3-2013 at 10:22

Does it react with O2 in air ?

If it is an Fe(II) salt it probably will oxidise to Fe(III), especially when wet. But that oxidation is compound dependent: Mohr's Salt [(NH4)2Fe(SO4)2 hydrate] for instance keeps well without oxidation, in dry conditions.

platedish29 - 18-3-2013 at 15:20

Okay I got very "pure" mohr salt here, gonna try and mix it down with phosphoric and see what happens, reporting soon with a edit,.
EDIT: the precipitation of iron II out of mohr's salt takes a little while, so I'll refer to see photo that might be of interest: it is a sample from a local excavation which florished this amazing red clay which I believe is composed of iron, along this other green stuff inside the bigger bucket which I believed could be chromium but I'm more inclined for it be green rust. It came in larger chunks that were smashed into a powder!
(current reproduction of the original experiment coming..)
EDIT 2: yea ammonium iron phosphate is really grey-green. But when it was first prepared starting through iron hydroxide, thats green in color, and phosphoric acid, it first dissolved completely and turn red-brown? Then uppon addition of ammonium hydroxide it turned to the final color.
I think its just hydration, no big enigma.
Perhaps a roast should do for it to turn white?

[Edited on 19-3-2013 by platedish29]

blogfast25 - 19-3-2013 at 06:24

EDIT 2: yea ammonium iron phosphate is really grey-green. But when it was first prepared starting through iron hydroxide, thats green in color, and phosphoric acid, it first dissolved completely and turn red-brown? Then uppon addition of ammonium hydroxide it turned to the final color.

[Edited on 19-3-2013 by platedish29]

Take us through what you did precisely?

elementcollector1 - 19-3-2013 at 07:29

Iron(II) hydroxide is green in color... until it meets atmospheric oxygen. Then it turns to rust-brown Fe(OH)3.

rstar - 19-3-2013 at 09:17

Ok, when i scrapped off that stuff from my watchglass, i saw that it was the same ol' bluish color (seen in the first image) . I stored it in a amber coloured glass vessel. I don't think i will need it so i'm gonna discard it shortly

Well, i wanna make some FeCl3, can anyone tell me a fast method of oxidizing FeCl2, without the use of H2O2 or Cl2 ?

blogfast25 - 19-3-2013 at 10:05

Well, i wanna make some FeCl3, can anyone tell me a fast method of oxidizing FeCl2, without the use of H2O2 or Cl2 ?

There are plenty of oxidisers that will oxidise Fe(II) to Fe(III) but few leave no trace. Then you need to separate the oxidiser by-product (its reduced form) from the desired FeCl3. Not easy. That's why Cl2 and H2O2 are the most desirable. H2O2 is also cheap as chips and readily available.

Air will do it too but it's a bit of a slow boat to China for practical, 100 % conversions.

HNO3 also does it but NO/NO2 (and water) is released, so be careful. And for most, HNO3 is not OTC or easily available...

[Edited on 19-3-2013 by blogfast25]

platedish29 - 19-3-2013 at 10:29

Take us through what you did precisely?

Uh?! what? I successfully reproduced the experiment for a green variety of the phosphate, showed for that procedures. What I got after filtering indicates that the original poster have gone tricky in his affairs, and that the blue variety must be more of a valuable than we may dink + I added absolutely extra ammnonia so it may enhance the facts about stoichometry mentioned earlier. The solution smells ammonia but it still managed to yield the very green product at the same time. Interesting is also the gelly green hydroxide changing into a brown colloidal suspension in the rist acid addition showed how iron behaves in this particular situation...
nevethless practically showed everything about iron there is to be shown.
Thanks.

rstar - 19-3-2013 at 16:18

There are plenty of oxidisers that will oxidise Fe(II) to Fe(III) but few leave no trace.

Can Ca(ClO)₂ do it cleanly ?

platedish29 - 19-3-2013 at 16:41

There are plenty of oxidisers that will oxidise Fe(II) to Fe(III) but few leave no trace.

Can Ca(ClO)₂ do it cleanly ?

Do you mean calcium hypochlorite? I know calcium when added to iron III sulphate switches its colors from brown to red, an indicative it would somehow contribute iron III formation. Is it commonly believed Fe3+ ions would survive alone in the solution? I mean, not sticking to something else?
Fe(HOH)3 + Fe(HOH)3 + HCl + H2O --> Fe2(OH)(HOH)5Cl + H2OH+

*****Dear rstar,
not into the assurance (that posts that it is you the owner) producing it too are clearly demeriting...but if I posed a photo of them in the internet? I kicked a weird glitter which is so beautiful!
:cool:

[Edited on 20-3-2013 by platedish29]

elementcollector1 - 19-3-2013 at 18:01

Do you mean calcium hypochlorite? I know calcium when added to iron III sulphate switches its colors from brown to red, an indicative it would somehow contribute iron III formation. Is it commonly believed Fe3+ ions would survive alone in the solution? I mean, not sticking to something else?
Fe(HOH)3 + Fe(HOH)3 + HCl + H2O --> Fe2(OH)(HOH)5Cl + H2OH+

*****Dear rstar,
not into the assurance posts that its you the owner producing it too are clearly demeriting if I posed a photo of them in the internet? I kicked weird glitter which is so beautiful!
:cool:

This post is unacceptable.

Fe(OH)3, not Fe(HOH)3.

"Fe2(OH)(HOH)5Cl" and "H2OH" do not exist.

Compounds do not react with themselves, just put "2 Fe(OH)3" if that's what you meant.

Calcium ions, when added to iron sulfate, precipitate out *white* calcium sulfate, leaving an iron salt (likely also in the +3 oxidation state).

As for just Fe ions, not technically true, even if you used stoichiometry: some small amount of CaSO4 does dissolve in water.

As for the last paragraph... What? Are you high or something? Or just trolling?

platedish29 - 19-3-2013 at 18:28

Quote: Originally posted by elementcollector1

Do you mean calcium hypochlorite? I know calcium when added to iron III sulphate switches its colors from brown to red, an indicative it would somehow contribute iron III formation. Is it commonly believed Fe3+ ions would survive alone in the solution? I mean, not sticking to something else?
Fe(HOH)3 + Fe(HOH)3 + HCl + H2O --> Fe2(OH)(HOH)5Cl + H2OH+

*****Dear rstar,
not into the assurance posts that its you the owner producing it too are clearly demeriting if I posed a photo of them in the internet? I kicked weird glitter which is so beautiful!
:cool:

Nope sier!
I just want to post a photo of the glitter I hade made and steal the credits for the idea.
Sorry for that phrase it SOUNDS confusing, or I just have a very colloquial english!

yeah, Fe(HOH)3 Should be Fe(OH2)6(3+) or Fe(OH2)4(3+) not sure, sorry, just missed the notations.

[Edited on 20-3-2013 by platedish29]

elementcollector1 - 19-3-2013 at 18:52

Trying to understand here: You made ferric ammonium phosphate, or some other compound, and it was glittery. You want to sell this compound, or otherwise "steal the idea".
...I still don't get it.
Oh, you were referring to the hydrated Fe(H2O)_x +3 ions. I thought you were referring to Fe(OH)3, the compound. My bad.

rstar - 19-3-2013 at 21:14

@platedish : LOL

FeCl2 + Ca(ClO)2 = ?

blogfast25 - 20-3-2013 at 04:12

@platedish : LOL

FeCl2 + Ca(ClO)2 = ?

Calcium hypochlorite will oxidise Fe(II) very easily but you end up with an almost inseparable mixture of FeCl3 and CaCl2: fairly useless.

platedish29 - 20-3-2013 at 12:09

The cause of coloration was probably just caged iron hydroxides. High heat makes them dehydrate into a bid deal of orange/ black/ brown mass.
How's going dessication of your blue stuff rstar?

rstar - 21-3-2013 at 00:44

Ok, going to heat up that blue stuff today and post the results in this space .

[

]

[Edited on 21-3-2013 by rstar]

rstar - 21-3-2013 at 01:06

I guess FeCl2 can be easily oxidized to FeCl3 by KMnO4, but there will be a problem of separating the Mn++ ions

blogfast25 - 21-3-2013 at 09:08

I guess FeCl2 can be easily oxidized to FeCl3 by KMnO4, but there will be a problem of separating the Mn++ ions

Yes. That's why H2O2, HNO3 or Cl2 are preferred: they leave no cations that would have to be separated from the Fe3+ cations, a difficult task.

rstar - 22-3-2013 at 09:18

I guess FeCl2 can be easily oxidized to FeCl3 by KMnO4, but there will be a problem of separating the Mn++ ions

Yes. That's why H2O2, HNO3 or Cl2 are preferred: they leave no cations that would have to be separated from the Fe3+ cations, a difficult task.

HNO3 ? conc or dilute ? and can you give me the equation ?

blogfast25 - 22-3-2013 at 10:07