Urea, its use in refining

[4metals]

At the request of Freechemist, I am starting this thread for the discussion of the use of urea in refining.

Please note that the use of excessive nitric acid is wasteful and requires either time consuming multiple evaporations of the acid to drive off the excess nitric acid or the use of urea to consume the excess nitric and drive it from the solution.

Urea or carbamide is an organic compound with the chemical formula CO(NH2)2. It is a white solid and comes in a form called prills, or little beads of the reagent. It is extremely soluble in water.

I pasted the following descriptions on the use of urea to kill the nitric in aqua regia solutions before precipitating the gold.

Urea prills are added to the acid, start slowly and add a small scoop, if there was excessive free nitric the reaction will be noticeable and can double the reaction volume as the solution rises with gas formation and foams up. Wait for the foaming to subside and add a bit more urea. You will notice less and less foaming as you add more urea and finally you will see the prills begin to float to the surface reacting very slowly if at all. That is the sign that the nitric has been consumed and the solution is ready for precipitation.

I really do not like urea but many refiners use it regularly. Some of the issues are that all of platinum in solution will not be dropped by using ammonium chloride if urea was used and this is true, however the platinum group metals will come down with cementation on copper so it is often a moot point unless you are looking for a lot of platinum in the solution.


Another point about urea, it is always important to leave room for the solution to rise when reacting. As a minimum never fill a vessel more than half full. I've seen many large tanks of acid foaming up right to the top before the big bubbles form and collapse to drop the level. That's nerve wracking! I always like to keep a hose handy to give a fine spray to those bubbles when they get close to the top.

In normal aqua regia digestions there is often some platinum or palladium. There are many scenarios for dropping the PGM's but all involve solutions with no nitric. Using urea in a solution where you expect to drop out PGM's will form some ammonia complexes with platinum which will inhibit their drop with ammonium chloride.

In another thread Freechemist commented the following;

I dislike urea, too, especially for destroying/neutralizing excess nitric acid, simply because I think it doesn't do this job. But still, urea has become a valuable tool for me as an easy to handle, precisely weighable source of ammonia, be it for neutralization of excess acid as a base, or to prepare some important PGM-ammine compounds in a simple, well controlled manner. To me, use of urea is an important theme, and so I propose, to discuss it in a new, separate thread

Of the forum members with extensive refining experience, I think it is unanimous that we all dislike urea and go to lengths to eliminate its use for killing excess nitric. When a member with a background like freechemist mentions PGM ammine synthesis using urea, it peaks my interest and warrants further discussion.

I invite posts concerning experiences where urea has either inhibited complete precipitation of platinum from solutions or, as freechemist implied, aided in synthesis of specific products.

 

[HAuCl4]

I've never used it for PM processing, but I have heard it's easy to make explosives with it, even unadvertently.

That's all I will say.

 

[freechemist]

To enter the discussion, my first point is definitely the fact, that urea doesn't do the job of neutralizing excess nitric acid left after dissolution of precious metals in aqua regia. What can be "neutralized" with urea, are the resulting reduction products of nitric acid, namely nitrous acid, HNO2, and NOx.

Reaction of urea with nitrous acid gives the gaseous products nitrogen (N2) and carbon dioxide (CO2), according to the following reaction:

H2N-CO-NH2 + 2 HNO2 + ==> CO2 + 2 N2 + 3 H2O

A similar reaction of HNO2, which has also been mentioned in this forum, occurs with sulfamic acid, H2NSO3H, leading to sulfuric acid and elemental nitrogen (N2):

NH2SO3H + HNO2 ==> H2SO4 + N2 + H2O

With nitric acid urea reacts as a weak base, forming the salt urea nitrate, which apparently can be crystallized easily:

(H2N)2CO + HNO3 ==> (H2N)2COH(+) + NO3(-)

Like HAuCl4 says in his post, urea nitrate can be used as an explosive. Urea nitrate is stable in acid. As a solid on a filter, it can be washed even with concentrated nitric acid. It can be prepared in aqueous media from HNO3 and urea. It seems to be quite soluble in water, and thus an old patent proposes it's synthesis in more or less non aqueous media.

Here some links leading to properties and preparation of urea nitrate:

http://en.wikipedia.org/wiki/Urea_nitrate

Sorry, but we will take this one out 4metals

http://www.google.com/patents/US3141039

http://www.google.com/patents/US3330864

In his entering post 4metals says:

"Please note that the use of excessive nitric acid is wasteful and requires either time consuming multiple evaporations of the acid to drive off the excess nitric acid or the use of urea to consume the excess nitric and drive it from the solution."

To me this statement is most important. In my practice I never used premixed aqua regia. Instead I did some maths, to calculate the amounts of acid (HCl and HNO3) needed, to dissolve the material, based on it's weight and composition. To calculate the amount of HNO3 needed, I always assumed an overall-reduction of HNO3 to HNO2, corresponding to a 1:1-mixture of NO and NO2, escaping the reaction mixture, thus avoiding the need to destroy an excess of nitric acid successfully.

 

[4metals]

Freechemist,

I thank you for your input.

Since you have mentioned the use of sulfamic acid, I would like to explore that with you a bit further.

The product of the sulfamic acid additions yields sulfuric acid in the process. Have you ever, after the precipitation of the gold, tried to convert the copper in solution from the chloride to the sulfate? My reason for asking is that a copper sulfate waste stream is a natural for electrodeposition of copper with the recovery of valuable copper and the lessening of the load on waste treatment.

While we are at it, and I may check this out the next time I have some bench time in the lab, would sulfamic elimination of the nitric eliminate the need for boiling with PGM drops?

 

[butcher]

Just my two centavos.

I do not want to add fertilizer to my clean gold solution, I worked so hard to get.

So I prefer not to use urea to remove excess nitric acid from solution, I also do not want to make urea nitrate, as many times I take powders from my stock pot, where these salts can mix with my metal powders, many times these powders are incinerated, and reprocessed, I am not saying that it would explode but these metal powders are dangerous enough, and why would I want to take the chance of making a primary explosive I am standing over in a glass or ceramic vessel?
I just prefer not to take a chance.

I have also read other sources that say urea nitrate does not have to be dry and have a detonator, to react explosively, but can also in a liquid state under some conditions, again I do not know all of these conditions but I will not take that chance.

Use minimum amount of nitric needed when dissolving the gold, (You can also leave some gold metal un-dissolved with heat, and get it next time to be sure the nitric is used up), then you do not need urea.

If you do use excess you can use the evaporation method to rid excess nitric acid from solution, for me this makes more sense than adding fertilizer, it is not hard if done right and once you understand how to do it.

Or if you do not like the evaporation method, and use more nitric than needed to dissolve gold, then using the trick of Harold's, adding a gold button some heating and nitric is used by the gold you get back later, you will only dissolve the amount of gold from this added button in proportion to the amount of excess nitric you added (that you did not need) to dissolve the gold in your original aqua regia solution, if you calculated the nitric well, none of this added gold button would dissolve.

 

[freechemist]

4metals,

I never have tried to convert copper in solution after precipitation of gold to the sulfate, because I didn't see the need to do so. The copper containing mother liquor after precipitation of gold and/or PGMs is always a mixture of aqueous HCl (main component) and H2SO4, regardless of the origin of the H2SO4-part, be it from oxidation of SO2, used in gold dropping or from pregoing treatment with sulfamic acid to destroy excess nitrous products. In my opinion it shoudn't matter much, if copper is electrodeposited from chloride containing acid solution, as long as cathode and anode compartments are separated. Sure, I see the problem of anodic chlorine generation in an undivided cell, but I think, that this can be overcome by a combination of electrodialysis with electrolysis of water with the aid of bipolar membranes, using the fact, that dissolved remnants of precious metals are complex anions, and thus wander together with the chloride- and sulfate-anions in an appropiately compartmented cell-stack, so, that in an ideal case in one compartment an acid CuSO4-solution, and in an other compartment a concentrated HCl-solution containing also the anionic precious metal complexes in a more concentrated form will be obtained.

http://www.fumatech.com/EN/Membrane-technology/Membrane-processes/Bipolar-membranes/

Your second question about elimination of nitric acid before PGM-precipitation, I cannot answer, because I never had problems in dropping Pt and/or Pd out of HCl-solutions containing small amounts of remaining HNO3.

Have a nice bench time in the lab, regards,

freechemist

 

[freechemist]

To continue the discussion, please let me comment on a few topics in 4metals introductory post.

1.) "Another point about urea, it is always important to leave room for the solution to rise when reacting. As a minimum never fill a vessel more than half full....."

A very important point, and also true, when using sulfamic acid in place of urea.

2.) "In normal aqua regia digestions there is often some platinum or palladium. There are many scenarios for dropping the PGM's but all involve solutions with no nitric...."

Having worked up a lot of dental alloys, quite frequently I had to deal with charges containing appreciable quantities of Pt and/or Pd, and I never encountered any problems in precipitating them, when small amounts of nitric acid were present.

3.)"...Using urea in a solution where you expect to drop out PGM's will form some ammonia complexes with platinum which will inhibit their drop with ammonium chloride."

This is primarily pH-dependent. If there is enough excessive HCl, corresponding to a solution pH of 0 or even less, not even traces of ammonia, and therefor, no ammine complexes are formed. The reason for inhibited PGM-precipitation may be a different one. The PGM's form complexes with sulfite and/or SO2, generally more stable, than their chloro-complexes, even in strongly acidic solution. In addition Pt(IV) in PtCl6(2-) is rapidly reduced by SO2 (or sulfite) to Pt(II), which remains dissolved as PtCl4(2-), not being precipitated with ammonium chloride. Thus, to me it seems not unreasonable, that in the foregoing gold-reduction-process at least some platinum is reduced to the two-valent state and thus not precipitated with ammonium chloride.

 

[freechemist]

Reaction of urea with water, a useful reaction

Urea reacts with water, especially at high temperatures, according to the following stoichiometry:

(H2N)2CO + H2O => 2 NH3 + CO2

The final reaction products are ammonia and carbon dioxide. Ammonia at room temperature is a gas, very well soluble in water, reacting as a weak base. The formed ammonia can serve two purposes: 1.) It can be used as a base, to neutralize acid, itself being transformed to the ammonium-ion, NH4(+). 2.) It can be used as a ligand, to form ammine complexes with the PGMs.

As an example prolonged boiling under reflux of an aqueous H2PdCl4-solution with the appropriate amount of urea yields an aqueous solution of tetrammine-palladium-II-chloride, while gaseous CO2 continuously escapes the reaction mixture:

H2PdCl4 + 3 (H2N)2CO + 3 H2O => [(NH3)4Pd]Cl2 + 2 NH4Cl + 3 CO2

This solution can be treated with the calculated amount of acid, to precipitate the yellow complex dichloro-diammine-palladium(II), (NH3)2PdCl2:

[(NH3)4Pd]Cl2 + 2 HCl => (NH3)2PdCl2 + 2 NH4Cl

If necessary, this compound can again be redissolved by treating it with the calculated amount of urea in aquous solution like described above:

(NH3)2PdCl2 + (H2N)2CO + H2O => [(NH3)4Pd]Cl2 + CO2

and reprecipitated with the appropriate quantity of aqueous HCl.

 

[samuel-a]

a useful reaction

Indeed it is ! Thank you freechemist for all of that info and thank you 4metals for starting that thread.


I have been scrubbing NH3 gas through my palladium nitrate solutions on a regular basis, then reduce volume by evaporation and precipitate with HCl in an ice bath.
I will definatlly give Urea a chance as i don't have bottled NH3 therefore i'm producing it on demand. Which is no fun.

 

[4metals]

I have always assumed the pH of the system effected the formation of the ammonia. When acidic, I never noticed the ammonaical smell, but when alkaline, for waste treatment, the ammonia smell is strong and the ammonia redissolves the copper hydroxide and causes issues with the discharge limits for copper. Usually locking up the copper hydroxide with dithiocarbamate works well in the presence of the ammonia.

At what pH are the ammine PGM complexes forming?

 

[freechemist]

4metals asks: "At what pH are the ammine PGM complexes forming?"

I never have measured the pH of the reacting mixtures. I assume, that ammine-complex formation occurs in a practically neutral or very weakly acidic medium (pH 6-7, may be even lower). During reaction gaseous CO2, a weak Lewis acid escapes, as long, as NH3 also formed is bound by dissolved PGM-ions. A pH-test paper always indicated a pH of about 6-7 after termination of the reaction.

 

[HAuCl4]

Hi freechemist and 4metals:

Let's say one has a solution of mainly silver nitrate, palladium nitrate, platinum nitrate, with some copper nitrate and minor amounts of other base metal nitrates. This solution was obtained by alloying silver and PGM compounds in a proportion of 7:1. There could be some minor amounts of gold in the solution as well, but mostly it has been filtered out after digestion in nitric acid.

What is the best plan ahead for recovery of platinum, palladium and silver, and does urea serve any useful purpose at some point in the process?.

Cheers.

 

[Lou]

I hate urea and it hasn't much use. Sulfamic works well at removing the nitric acid because a.) it's very keen on forming stable sulfate (which won't participate in redox), b.) nitrous oxide is a rather non polar gas that readily disperse from the solution and is removed from the system's equilibria.
The reaction of H3NSO3 with HNO3 is also the premier method to make N2O in the laboratory and has been used as such for years, as it makes very clean gas (much better than the decomposition of ammonium nitrate).

To answer 4metals's question "at what pH" are the ammine PGM complexes forming?":

this depends on the metal and its relative solution concentration and the answers may be derived from ligand field theory.
It also depends what else is attached (i.e. chloride, bromide, etc.).

In the case of palladium, the ammine ligand will add quite quickly to give pretty pink Pd(NH3)4PdCl4 even at low pH of 1-2. This binuclear complex shares +2 charge. This can be isolated from pH 3-8 and is most easily reduced.
As you increase the pH, it further solubilizes--this is when (over pH 7-8) the ammine ligand gets tacked on and forms (NH3)4PdCl2, which is wholly soluble and rather base stable.

for chloroauric acid:

If you had silver nitrate, palladium nitrate, platinum nitrate and copper nitrate all in solution...

add concentrated HCl (equivalent to silver content) to precipitate the silver, filter cold, rinse with dilute 0.5 M HCl until colorless filtrate.
at this point, the remaining metals are +2, including palladium. That is the opportune time to remove palladium with DMG.
If you're not interested in that, consider adding more HCl, syrupizing, precipitating the platinum classically (Pt (IV) is rather stable, Pd(IV) isn't).
Then go after the Pd as per normal.

 

[HAuCl4]

for chloroauric acid:

If you had silver nitrate, palladium nitrate, platinum nitrate and copper nitrate all in solution...

add concentrated HCl (equivalent to silver content) to precipitate the silver, filter cold, rinse with dilute 0.5 M HCl until colorless filtrate.
at this point, the remaining metals are +2, including palladium. That is the opportune time to remove palladium with DMG.

This leaves platinum nitrate, arguably the largest value in $$ of the 3 metals (the one I'd like to get out the door fast, if possible), in solution with a large amount of concentrated nitric acid. How do I get the platinum from this very acid solution?.

I have a few ideas, but none that I really like.

 

[Lou]

Can simply add an 2X excess of HCl to it and either kill the nitric whichever way you like (boil down, sulfamic, sparging, sulfite, more platinum...etc.)

 

[freechemist]

Hi Lou and HAuCl4,

For treating chloroauric acid's hypothetical solution, Lou makes the following proposition:

"If you had silver nitrate, palladium nitrate, platinum nitrate and copper nitrate all in solution...

add concentrated HCl (equivalent to silver content) to precipitate the silver, filter cold, rinse with dilute 0.5 M HCl until colorless filtrate. at this point, the remaining metals are +2, including palladium. That is the opportune time to remove palladium with DMG.
If you're not interested in that, consider adding more HCl, syrupizing, precipitating the platinum classically (Pt (IV) is rather stable, Pd(IV) isn't).
Then go after the Pd as per normal."

I think, that this procedure looks much simpler, than it actually is. The main reason for this can be found in the fact, that not only Ag, but Pd and Pt as well form chloro-complexes, probably more stable than their Ag-congeners. Thus, adding the stoichiometrically calculated amount of HCl to a solution of all three metals in HNO3 in order to precipitate silver quantitatively as AgCl, will form chloro-complexes of Pd and Pt, too, thus leading to a solution containing Pd- and Pt-chloro-complexes, together with silver and some already precipitated AgCl. The only practicable remaining solution is, to add the all three metals bearing HNO3-solution to a big excess of HCl, separate AgCl by filtration and continue in treating the filtrate, like described above.

Knowing the exact quantities of dissolved Pd and Pt, an alternative procedure could be their coprecipitation with dimethylglyoxime, leaving silver in solution, by using a small excess (10-20%) over the stoichiometrically calculated amount. To do so, the dimethylglyoxime should be dissolved in 1-2 M NaOH, not in an organic solvent, to prevent reduction of Pt(II) to the metal. It is obvious, that urea never is serving any useful purpose in all these procedures.

 

[HAuCl4]

Thanks guys. It looks like that alloy is best avoided, like Hoke said 100 years or so ago...

 

[Lou]

I wonder what the thermodynamics have to say about silver chloride vs the chloroanion complexes of the PGMs?

To be truthful, I've never tried adding silver nitrate to a PdCl2 solution to see if it can displace.

 

[4metals]

I think it will come down in theory but with an eV difference of 0.15 and all of that fluffy silver chloride with its large surface area a lot may get hung up mechanically.

Ag(+) + e(-) ----> Ag(0) eV= 0.80

Pd(2+) + 2 e(-) ----> Pd(0) eV= 0.95

 

[HAuCl4]

Is there a simple procedure to make silver "blacks"?. (i.e. super fine powders of pure silver).