# "de-NOxing" of questionable need in Pt purification



## tensor9 (May 23, 2012)

I've been browsing the forum and have seen quite a bit about de-NOxing in Pt purification. I don't think it's necessary. Possibly it's needed from some reason with a mixture of noble metals?

According to Inorganic Syntheses, this is not really a problem. The dissolution of platinum in aqua regia yields hexachloroplatinic acid (H3O)2PtCl6. Some (really old) literature claims that there is considerable nitrosyl contamination as a result, but it turns out that this is a very minor problem. The compound (NO)2PtCl6 is the the culprit, and it is fact formed in very small amounts under the aforementioned reaction conditions, but it is quite unstable in aqueous solution. (Considerable effort is required to obtain this compound in good yield.) So, if your aim is to synthesize hexachloroplatinic acid, you just might be inclined to worry about this since the two can co-precipitate. However, in Pt recovery, you're going for the very poorly soluble ammonium salt, (NH4)2PtCl6. So, the recommended procedure is to evaporate the aqua regia/HCPA solution JUST to dryness to get rid of the strong oxidizers, then redissolve in pure water, then add ammonium chloride. The resulting precipitate will be the ammonium hexachloroplatinate salt, so the small nitrosyl impurity, if present, will remain in solution and just not matter.


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## Oz (May 23, 2012)

It almost sounds as though you are speaking of re-purification of platinum with only traces of base metals.

The typical reason for the removal of HNO3 in most refining is that palladium is often present and will co-precipitate with the platinum with the classic addition of ammonium chloride to a degree. If HNO3 is still present.

To be clear, I have never had a chemistry course in my life, but I will enjoy seeing this discussed in detail by those that have.

Edit in red for clarity.


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## tensor9 (May 23, 2012)

Yeah, that's kind of what I'm wondering. Is there some other metals that a little bit of NOx will interfere with?


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## Oz (May 23, 2012)

I just do not know all that much as to other elements in this case and the use of AR. 

I typically avoid the problem of HNO3 in this situation by using Cl2 gas with my HCl. For mixed platinum group metals it has the added benefit of doing a better job of digestion when Rh is present. Free Cl2 is also easily removed compared to HNO3 with no interference.


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## lazersteve (May 23, 2012)

Denoxxing is important when other PMs are present. Check out this chart I made based upon several common reducing agents and precipitants listed in literature:

Precipitant Chart

The chart does not contain all possible combinations of solutions and reagents, but it helps one understand why certain processes and precipitants are required for the highest purity metals.

I agree with Oz, use chlorine instead, it just does a better job on powders and foils in my opinion. If your metals (especially PGMs) are not in a finely divided state AR with careful control of the nitric additions is the better choice. 

Steve


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## tensor9 (May 23, 2012)

Thanks! For my purposes I shouldn't have to worry about this too much, but it's good to know.


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## Lou (May 23, 2012)

It's not so much the nitroysl complexes as it is the fact that AR is a great solvent for ammonium hexachloroplatinate. Precipitate from a strong hot aqua regia solution and see if your precipitate doesn't redissolve. I had a customer and his chemists leave tens of thousands in solution not knowing that it interferes with the completeness of the reaction.


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## tensor9 (May 23, 2012)

Yeah, even still. One evaporation to dryness is all that's required. Multiple evaporations with HCl I don't think is necessary if you're only worried about Pt.


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## Lino1406 (May 24, 2012)

If you call "killing" of HNO3 "de-noxing"
then the reason is preventing new chlorine
from forming and intefering with precipitation.
If you refer to NO2 removal it is just for neatly
working


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## freechemist (May 30, 2012)

Hi tensor9,

What is written in your beloved "Inorganic Synthesis", I can only confirm. In my practice I never had problems with remnants of some HNO3 or NOx in H2PtCl6-solutions, prepared by careful dissolution of the metal in half-concentrated, hot HCl, by slow oxidation of the metal with a calculated amount of HNO3, using the following overall-redox-stoichiometry:

Pt(0) + 2 N(V) => Pt(IV) + 2 N(III)

A good method, to prepare H2PtBr6-solutions is, to use about 10 moles 48% HBr (azeotrope) per gram-equivalent of Pt and a small excess of 35% H2O2 (ca. 2.1-2.2 moles per gram-equivalent of Pt). The metal (preferrably powder or sponge) is suspended in 48% HBr with constant stirring, and 35% H2O2 is added slowly. The reaction, especially at the beginning, is fairly exothermic, so means to cool down eventually should be available. In order to minimize bromine-losses, the reaction apparatus should be equipped with a reflux-condenser. After complete addition the mixture is boiled under reflux until all metal is dissolved and no more dark colored heavy droplets of elemental bromine are falling down from the reflux-condenser.

The resulting, intensely red colored solution of H2PtBr6 is diluted to ca. 50 g Pt/liter, filtered to remove potential insoluble impurities (or traces of unreacted metal) and can be used directly to prepare either pure Pt-metal by reduction or to prepare pure salts, (NH4)2PtBr6 and/or K2PtBr6 by precipitating with NH4Br or KBr. 

Regards, freechemist


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## Lou (May 30, 2012)

About how fast is that HBr/H2O2 method?


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## tensor9 (May 31, 2012)

Inorganic Syntheses (plural) is the journal. Inorganic chemists like myself LOVE it because if it's in there, it works. (Syntheses are submitted and actually confirmed to work by checkers.)

In my Pt recovery, I tend to stay away from the platinic acids of any sort because they are not convenient to handle or store.

Along the lines of what Lou asked, is there an advantage to HBr/H2O2? Having to deal with the bromine is kind of a pain. Speaking of the hydrogen peroxide method in general, the literature I've found suggests 10:1 concetrated HCl to 35% hydrogen peroxide (v:v). Is this what you've found as well?

I do use 48% HBr to dissolve PtO2 residues. Interestingly, HBr will dissolve it when nothing else will. I use potassium bromide to precipitate the potassium hexabromoplatinate salt. I was pretty sure ammonium bromide would do it too, but hadn't actually found any hard data on the solubility of the ammonium salt. Thanks for confirming it.


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## lazersteve (May 31, 2012)

I can confirm the HCl and peroxide ratios you quoted from literature. They have worked perfectly for me on numerous occasions. My starting material is typically foils, thin wires, or sponge. Dissolution proceeds slowly at first, then more rapidly as the temperature of the reacition increases (exothermic).

Be very careful when working with finely divided Pt and foils as the reaction takes off with lots of chlorine off gassing, foaming, and heating. Your reaction vessel needs three to four volumes of the reaction solution in head room. I always keep an overflow basin handy in case of rapid boil overs. I tend to be a little heavy handed on the peroxide on occasion as well. Filter the solution until transparent,:







preciptate with your favorite precipitant (I like potassium chloride as seen in the photo):






filter canary powder as seen here:






, and reduce with your chosen reducing method (hydrazine+base = fast, zinc+ dilute acid = slower, but safer to handle, calcine, etc.).

The hydrazine method gives a beautiful sand like product, while the zinc method produces a 'brain-like' sponge. 

Here's the result of a hydrazine reduction:






Both are easier to handle once they have been vaccum filtered and clinkered up with a hydrogen-oxy torch.











I stopped part way through the melting process and snapped this photo:








Steve


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## freechemist (Jun 1, 2012)

> *Lou:* _"About how fast is that HBr/H2O2 method?"_


I never paid a lot of attention on how fast the metal is dissolved. Usually, after a few hours of refluxing, dissolution was complete. However, dissolution of Pt-powder or sponge in my hands was always faster in HBr/H2O2, than in HCl/NaClO3.

*tensor9's remarks:*

_"Along the lines of what Lou asked, is there an advantage to HBr/H2O2?"_

Yes, there is definitely. Br2, at room temperature, is a liquid and is much less volatile, than Cl2. It condenses in a reflux-condenser and flows back into the reactor. Secondly Br2 forms an ionic addition compound with Br(-), Br3(-), the tribromide anion. Thus much more Br2 dissolves in (nearly) azeotropic HBr, than in water. The system HCl/Cl2 behaves differently, especially on heating, where Cl2 is volatilized rapidly and not retained by a reflux condenser.

_"Having to deal with the bromine is kind of a pain."_

You don't handle bromine actually. You generate it in the reactor directly, where it is slowly used up in a safe containement. You don't have to isolate and measure it as liquid Br2.

_"Speaking of the hydrogen peroxide method in general, the literature I've found suggests 10:1 concetrated HCl to 35% hydrogen peroxide (v:v). Is this what you've found as well?"_

In my practice I always used calculated, measured (excess included) amounts of hydrohalic acid and oxidant to dissolve precious metals (Au and PGMs). This led me to controlled conditions, especially on working with HCl/HNO3, where I never encountered problems, caused by excessive remnants of HNO3. Excessive hydrohalic acid is never a problem. Working in the system HCl/Cl2, I seldom used H2O2 as an oxidant. I preferred oxidation with NaClO3, despite the fact that resulting PM-solutions contain some sodium.

_"I do use 48% HBr to dissolve PtO2 residues. Interestingly, HBr will dissolve it when nothing else will."_

Your observation shows, that Pt(IV) as a center in complex hexahalo-anions has a greater affinity for Br(-) obviously, than for Cl(-) as a ligand. It could also explain the faster dissolution of Pt in HBr/H2O2, than in HCl/NaClO3.

_"Inorganic Syntheses (plural) is the journal. Inorganic chemists like myself LOVE it because if it's in there, it works."_

I love it too, and I fully accept, to see me corrected.


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## Lou (Jun 1, 2012)

Too bad HBr is so darn expensive.


Tensor, refer to your knowledge of HSAB theory to understand why bromine is better than chlorine. In general, Cl<Br<I for precious metals insofar as affinity for complexation reactions.


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## tensor9 (Jun 1, 2012)

freechemist,

I wasn't under the impression that isolation of bromine was necessary. Though I did forget that bromine was so volatile. (I just let the hood take care of it.)

As for the nature of the chloride and bromide complexes, I am well aware of all of this. Although, a formally 4+ Pt center with such a high charge is not considered to be soft. : ) The degree of covalency in the Pt-halogen bond is strikingly high and of course increases going down the the halogen group.

Wow, I remember TM chemistry. My actual research is with lanthanides, so it's kinda fun to work with transition metals, even if it's quite ancillary to my project.


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## Lou (Jun 1, 2012)

Lanthanides? and you're asking questions of us PGM chemists? Surely something is amiss in the world.


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## tensor9 (Jun 1, 2012)

Yeah, need Pt for my ligand synthesis, and I need my ligand for my lanthanide cluster synthesis.


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## MichaelB (Mar 29, 2014)

For a moment I was afraid you were going to want someone here to explain how to separate all of the rare earth elements from each other. As you know that would probably be a little much for a forum.


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