# Palladium in various states of process



## Harold_V

I've threatened for some time to post a picture of samples of palladium that I kept from my refining years. Below, you will see five sample bottles, along with one of the new US $1 bronze coins. 

The first specimen, left side, contains palladium that was precipitated with ammonium chloride and sodium chlorate. The salt was then dissolved in ammonium hydroxide and re-precipitated with HCL. The color changes from red to pale yellow, as you can see. 

The second specimen is palladium precipitated with ammonium chloride and sodium chlorate.

The third specimen is palladium precipitated with dimethylglyoxime (DMG). Note the bright yellow color. 

Fourth specimen is calcined palladium salt----in this case precipitated with ammonium chloride and sodium chlorate.

Fifth specimen is a small bead of palladium. 

Harold


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## lazersteve

Great post Harold,

Now I have a good idea of what to look for when I precipitate my Palladium batch!  

Steve


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## blueduck

Is there difference's in the weight to volume ratio of the different precipitations? 

If that is not quite clear what I am trying to ask is when the different methods are used, like the DMG, I believe you or someone posted elsewhere that it gives volumes of precipitation with minute amounts of Pd.....

In my reading, I see that there is always more than one way to skin a cat [or in this case precipitate Pd out of solution] I guess I want to know what is easiest, and then which would be the better to use when trying to bring it back to a button form [if you have a guess even] more than likely since I do not have an induction furnace access [though I have not checked up to the university of Idaho yet] I will more than likely use a torch to reduce the powders to metal form.... 

btw very nice and clear photo!

William


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## Harold_V

blueduck said:


> Is there difference's in the weight to volume ratio of the different precipitations?



I'm unable to answer that. I never bothered to make any determinations, for I never had to account for the platinum group so recovered. As a result, I wasn't concerned with ratios. The salt reduces considerably in volume when it is calcined, so it's quite light in weight as compared to the sponge. 



> If that is not quite clear what I am trying to ask is when the different methods are used, like the DMG, I believe you or someone posted elsewhere that it gives volumes of precipitation with minute amounts of Pd.....



It's clear for me----and I'd comment that using DMG for precipitation isn't the smartest thing you can do. While it will precipitate the most dilute of solutions, the volume of precipitant is huge----and as I've mentioned, not easy to handle. When dry it tends to stick to everything it touches. The other methods yield tiny crystals that are more like sugar---so they handle quite well---but are water soluble, so they must be handled with dry utensils. Using a washing bottle, traces left behind in handling are easily washed away to the stock pot , where they are stored for future recovery. 

If the concentration level isn't high enough to precipitate with ammonium chloride and sodium chlorate, I highly recommend it be either evaporated, or poured to the stock pot, where it can accumulate and increase the volume for future recovery en mass. The process of recovering palladium is slow and tedious---and not worth taking the time to recover in small volumes. You'll understand this better after having done it a few times. 



> In my reading, I see that there is always more than one way to skin a cat [or in this case precipitate Pd out of solution] I guess I want to know what is easiest, and then which would be the better to use when trying to bring it back to a button form [if you have a guess even] more than likely since I do not have an induction furnace access [though I have not checked up to the university of Idaho yet] I will more than likely use a torch to reduce the powders to metal form....



Please read Hoke's book on reducing the salts. Doing it with a torch can be a mistake. It should be heated slowly, so as to not drive off values when calcining. Further, a direct flame would be inclined to blow some of the salt away. It's a slow process, one that would take too long to achieve with a hand held torch. I used a Fisher type burner, with a Vicor or quartz melting dish held in a ring. A batch might take as much as a half hour. The calcining operation is best accomplished in a hood. The smell leaves a great deal to be desired, and is likely not good to breath. I avoided doing so by doing the operation in the mouth of my hood.

An induction furnace is not necessary for melting palladium, and can even be avoided with platinum for small volumes. A large Hoke torch, or even an oxy/acet torch with a rosebud can be used with success. The tip should be well cleaned before application, to avoid contamination from scale. 

If you're not familiar with induction furnaces, I'd encourage you to do some research. Just having one may not be the solution. As volume increases, frequency of the induction furnace can be reduced. Large induction furnaces (capable of melting many tons of metal) can run line frequency, but as volumes are reduced, the required frequency to melt must increase. To give you an idea, I own a 50kw Ajax Magnethermic power supply--which is capable of melting up to 200 pounds of metal, assuming you have the proper furnace. It runs @ 3,000 Hz, but they also made one of the same capacity that operated @ 10,000 Hz. As the frequency increases, they are capable of starting with smaller feed material. You might not be successful achieving a molten mass starting with fine particles at 3,000 Hz, where with a 10,000 Hz machine you could, as an example. 

There is a published chart that shows optimum frequency for given loads, which would be a good reference for those that might be interested in exploring induction furnaces. 

A microwave oven is nothing more than a high frequency induction furnace. Anything can be heated by induction if the frequency is high enough. 



> btw very nice and clear photo!



Thanks. I use a Sony FD-97 that's about 5 years old now. It's proven to be a good camera---although somewhat large in size. It records to a floppy disc, or a memory stick, which I do not have. Very easy to take pics, then remove the disk and insert it in the floppy drive of the computer. No cables required. 

Harold


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## Lou

http://webpages.charter.net/dawill/tmoranwms/Electronics.html

Sorry to resurrect and old thread, but this fellow has done a lot of work on the small scale (<10 kW) with inductions furnaces. As Harold's already mentioned, they're not required, but they are very nice to have as they melt quickly, cleanly, and efficiently. You can also melt your metals in a vacuum.




Also, I wouldn't use DMG as it will also drop out any nickel in there as well. Dimethylglyoxime is also pretty expensive, and fairly toxic as well. 
Ammonium chloride + sodium chlorate works very well. If you don't have chlorate and don't want to make it from salt (easy to do), you can make chlorine gas with manganese dioxide and HCl and bubble that through a warm solution. Then chill it all and most of your precipitate crashes out as a nice brick red slimy precipitate.

For any of you that care to make your own chlorate for a buck a pound...
http://webpages.charter.net/dawill/tmoranwms/Chem_Chlorate.html



Lou


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## Irons

This was presented in earlier threads to dissolve Gold by generating Chlorine gas in situ. It might also make a good Chlorine gas generator for the Precipitation of Palladium complex.


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## Lino1406

Dear Harold,
Is it possible that the subject material
is Palladousammine dichloride? the 
color about fits.
Lino1406


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## Harold_V

Lino1406 said:


> Dear Harold,
> Is it possible that the subject material
> is Palladousammine dichloride? the
> color about fits.
> Lino1406


Chuckle!

You're asking a guy that struggled to graduate from high school, having taken no chemistry classes, a question like that?  

Certainly could be------but I'd have no way of knowing. I welcome your input.

Harold


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## Lou

The brick red is your (NH4)2PdCl6, ammonium hexachloropalladate in its +4 oxidation state and that is the start of your pyrolysis pathway for palladium. At about 350C it goes to palladite, the +2 oxidation state.


The way you're naming sounds like it’s complexed with NH3. Nope, it's actually an ammonia salt. Don’t even get me started with why it’s a salt not a complex and not vice versa…

I’ve been meaning to post this up for a while. I won’t mention Ru, Os, Ir, and Re unless y’all ask me to, since not many people deal with it. Same with some of the other weird compounds that you rarely ever see, like platinum oxide (sounds like an oxymoron doesn’t it  )To the CRC (my Hoke  ):
The properties of Palladium, platinum, rhodium and their common salts:

Palladium
Pd-palladium, 106.42 g/mol, melts at 1554.9oC. Soluble in aqua regia, nitric acid, conc. sulfuric, and Cl2+ HCl to give a deeply yellow solution that appears brown.

(NH4)2PdCl6—ammonium hexachloropalladate (IV)—355.21 g/mol. Brick red hygroscopic crystals. Soluble in hot water, concentrated ammonia. Reduce to sponge with heat or hydrazine, Pd black with formic acid or borohydride.

PdCl2—palladium (II) chloride—177.33 g/mol. Ruby red hygroscopic crystals. Soluble in acetone, (m)ethanol, and water. Reduce with hydrogen, 500oC.

PdO—palladium (II) oxide -- 122.42 g/mol. Greenish black crystals. Insoluble in water, readily soluble in aqua regia. Palladium hydroxide is a hydrated form of this compound. Reduce with ammonium formate or heat til dissociation of the elements at 750oC.

Pd(NO3)2—palladium (II) nitrate—230.43 g/mol. Brown hygroscopic crystals. Slightly soluble in water, very soluble in dilute nitric acid. Reduce with pyrolysis.

Platinum
Pt—platinum, 195.08g/mol. Melts at 1768.4 oC. Soluble in hot aqua regia, practically nothing else. Gives a red orange solution.

(NH4)2PtCl6 – ammonium hexachloroplatinate (IV) – 443.87 g/mol. Orange-red cubic crystals. Slightly soluble in hot water, conc. ammonia, insoluble in ethanol. Reduce to sponge with heat or hydrazine, Pt black with formic acid or borohydride. Alkali metal analogues.

(NH4)2PtCl4—ammonium tetrachloroplatinate (II)—372.97 g/mol. Dark red crystals. Soluble in water. Reduce to sponge with heat.

H2PtCl6*6H2O—hexachloroplatinic acid hexahydrate—517.9g/mol. Brown-yellow hygroscopic crystals. Soluble in HCl, very much so in water and (m)ethanol.

Rhodium
Rh, rhodium 102.91 g/mol, melts at 1964 oC, really damn hot. Slightly soluble in aqua regia and fuming sulfuric acid. Best dissolved by fusion.

RhCl3—rhodium (III) chloride—209.264 g/mol. Red to chocolate brown crystals. Soluble in alkalis, strong acids; insoluble in water. Reduction with hydrogen at 500 oC gives clean Rh sponge.

Rh2(SO4)3—rhodium (III) sulfate—494.002 g/mol. Red-yellow solid, brown solution. Slightly soluble in water, more so with sulfuric acid. Decomposes to the oxide at 500 oC. Cement with merrilite zinc or magnesium turnings. Or plate, or reduce the oxide with hydrogen.

Rh4(CO)12--Rhodium [dodeca]carbonyl—747.747g/mol. Red crystals. Decomposes with water or heat to yield rhodium metal.


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## Lino1406

Thanks Lou and Harold 
You gave me the answer, since:
The subject material, Pd(NH3)2Cl2
is received from PdCl2 by the method
Harold mentionned, that is NH3, followed
by HCl - no need for chlorate in between
and the Pd there is +2. Now, if Pd+4 gives
the same color to Pd+2 (which personally
I doubt), that could be interesting, or - some
traces of Pd+2 have been also present there
despite the oxidation by chlorate...
(words words words but how interesting)
Lino1406


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## Lou

ok Pd(NH3)2Cl2, diamminopalladous chloride is a complex salt. Acidification with peroxide will probably make (NH4)2PdCl6, or skip the peroxide--the hydronium from the HCl will break the complex and make the ammonia salt.

Most complexes are really soluble, did you crystallize it out, or did it precipitate? If so, from what temperature solution?


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## Lino1406

Hello Lou,
This is very resistant precipitate
at room temperature. The exact
procedure: 1)Dissolve PdCl2 in
HCl, or take ready Pd solution
2) Add conc. NH4OH (or NH3)
to get some scattered yellow
red brownish particles. 
3)Add fluently NH4OH, this should
dissolve most of the particles
4)This part is a bit nasty, add
HCl carefully (much boiling and
smoking), and you get it at once.
You may want to wash it with alcohol.
Lino1406


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## lazersteve

Lino,

That is the exact same procedure I use to redissolve and purify my red Pd salts and complexes from ammonium chloride + sodium chlorate precipitations. 

Hoke outlines it in her book.

By the way Lou,

I tried Irons idea to simply add sodium hypochlorite 5% (bleach) to the ammonium chloride, HCl, and palladium solution. It worked like a charm. Looks promising for eliminating the the sodium chlorate as a precipitation reagent with Pd salts and complexes. There is still a very fine balance of ammoinum chloride and chlorine required. 

Steve


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## Irons

Interesting. I had envisioned using the Hypochlorite/HCl as a separate Chlorine generator to minimize the addition of more water but adding it directly scavenges excess HCL. I'm a little leery about adding any more spectaror ions that would increase contaminants.

In an earlier thread, I had proposed to bubble Ammonia gas directly to scavenge the excess HCL and convert it to Ammonium Chloride in situ. This would reduce the water even more. The Ammonia comes from the Nitric acid production from Ammonium Nitrate and Calcium Hydroxide to produce Calcium Nitrate which is reacted with Sulfuric Acid togive HN03 and Calcium Sulfate as a precipitant.

Ammonia is also good for converting any excess Nitric Acid to Ammonium Nitrate which could be recycled back. This might be useful where Palladium is digested in Nitric alone.


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## lazersteve

Irons,

I agree with you about adding too many ions to the solution. I was just doing a quick small scale test of the idea and it seemed to work. A separate generator setup would be much cleaner, but would involve more labware. 

Since the resultant Pd complex is only slightly soluble in dilute (15% >)ammonum chloride solution the extra water doesn't seem to be a factor. The key is to produce ammonium chloride (NH4OH + HCl ) in situ or add in enough solid NH4Cl to prevent the resulant complex from redissolving when the Cl2 is added.

Steve


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## Irons

Having a Chlorine generator is handy when working With PMs. HCl+Cl makes for a clean dissolution. I's like the analog to using SO2 to precipitate Gold rather than SMB. They both work but the gas gives a cleaner product.

Calcium Hypochlorite is available cheaply at pool chemical suppliers and all you need is a HCl drip to generaate the gas. It's also available in cylinders, as is anhydrous Ammonia.

I wasn't trying to be critical, just a nitpicker. :

BTW, Anhydrous Ammonia is also great for getting rid of Gophers. Stick the hose down the hole and pack dirt around it to keep it from blowing out. Trust me, a face full of anhydrous Ammonia is something you won't forget.


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## Lou

I'm with you Irons, the less spectator ions in solution, the better! I usually use SO2 for my first Au precip, then oxalic acid the second time round. The sulfur dioxide is in a lecture bottle, and so is my chlorine. It's kind of hard/expensive to get chlorine and sulfur dioxide cylinders nowadays.



Hypochlorite will work instead of chlorate, and so will hydrogen peroxide and HCl (forms your chlorine in situ thru hypochlorite). You need an oxidant to get free chlorine and your solution must be strongly acidic. Irons, it will scavenge only so much HCl in dilute solution because there's an equilibrium between the chlorine formed from OCl- and OCl- formed from the chlorine! That is how I make strong 12% bleach, using chlorine gas and cold sodium hydroxide solution.

Awhile back I did a good sized precipitation of the ammonium hexachloropalladate salt in a 4L beaker. Probably 500g worth. I did the precipitation with excess HCl and a little bit of sodium chlorate to make free chlorine (that's owing to the fact that I couldn't get the damn chlorine cylinder open, it's really corroded). I also added about 130% of theoretical amount of ammonium chloride in several portions with magnetic stirring. 

The precipitation was done hot, then the whole lot of it was chilled in an ice bath, volume decanted, and then vacuum filtered. I held back some of my ammonium chloride solution and acidified it slightly, chilled it, and used that for rinsing my precipitate in the buchner. Extra ammonium chloride has no effect on the reduction of your palladate salt.

I'd use bleach, but it's a real pain to make it up (I don't use store bleach) and I've got plenty of chlorate laying around.


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## Irons

It's good to have different options to do the same process depending on what you have available or is easy and inexpensive to get. Many of us live in the boonies where chemical supply houses are few and far between and shipping and HAZMAT charges can really cut into profit.

I'm sure we've confused a lot of people with all of the different variations and the techniques to adapt chemicals becomes more of an art form than cookbook chemistry. 

The best thing for newbies is to stick with the tried and true procedures and leave the experimentation to those with those who understand the pitfalls.

Buy the Hoke book if you're not sure what to do. Mistakes in processing can be very expensive and sometimes hazardous.

I learn something new every day, especially since joining here.


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## Lou

I seem to hear a lot of good about this Hoke book. I think I'm going to have to look around for it. Is it more on the practical/cookbook side than it is on chemistry?


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## aflacglobal

Written in 1940 by a lady no less. The tried and true methods of every want to be refiner for over 67 years. The one the only C.M. Hoke :arrow: 
http://www.lmine.com/Merchant2/merchant.mvc?Screen=PROD&Product_Code=17205&Category_Code=assaying&Product_Count=3


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## aflacglobal

Irons said:


> It's good to have different options to do the same process depending on what you have available or is easy and inexpensive to get. Many of us live in the boonies where chemical supply houses are few and far between and shipping and HAZMAT charges can really cut into profit.
> I'm sure we've confused a lot of people with all of the different variations and the techniques to adapt chemicals becomes more of an art form than cookbook chemistry.



That's why i love this guy. :wink: 

Question :?: 

Now i understand about ions + and - . And most of the other good stuff involved. Sometimes i just can't see the obvious.

Someone please break down for me the how's and why's of what happens in the reaction. In other words what make the acid attack the metals and exactly does the roles of oxidizers and other forms of chemicals play in it. 
I understand the swapping of electrons because oxygen needs 8 electrons and blah, blah, blah. but for some reason i can't just quiet picture it all. 
Is it the chemical energy that is held by the electrons of the acid that react with the lower density electrons of the metal to reduce it to ions.

I have been bouncing from chemistry to physics and all over the place lately. Some things even escape the cat. :shock:


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## lazersteve

Ralph said:


> Is it the chemical energy that is held by the electrons of the acid that react with the lower density electrons of the metal to reduce it to ions.



Reactions are driven by several factors including:


 Temperature
 pH (H+ concentrations)
 Light (photons)
 Electronegativity (atomic charge)
 Catalytic activity 
 Population (species concentrations)
 Solubility
 Pressure

I'm sure I missed some and welcome some of the other chemical guys to chime in.

In a nutshell the conditions of the reaction are what drives any reaction. 

The insoluble gold becomes soluble when it combines with another species like chlorine.

Combination is catalyzed by chlorine oxidizing the gold to become positively charged (+1 and +3 states) and the negatively charged chlorine attaches to the positively charged gold (opposites attract). This process repeats three times for each atom of gold.

Once combined the newly formed auric chloride (AuCl3), which is now neutrally charged with full electron orbitals, is soluble in the aqueous acid medium.

Various literature state some reactions for gold as :








And the reduction of gold by FeSO4 (copperas) as:

FeSo4 * 7H2O + AuCl3 = Fe2(SO4)3 +FeCl3 + Au + 7H2O

via SO2:

3SO2 + 2AuCl3 + 3H2O = 3SO3 + 6HCl + 2Au

via SMB in water:

Na2SO3 + H2O = 2NaHSO3

3NaHSO3 + 2AuCl3 + 3H2O = 3NaHSO4 + 6 HCl + 2Au

Sodium Meta Bisulfite becomes Sodium Meta Bisulfate (this is why bisulfate won't precipitate gold)

via Hydrazine:

3N2H4 + 4AuCl3 = 12HCl + 3N2 + 4Au

New: Via Oxalic Acid:
3H2C2O4 + 2AuCl3 = 6HCl + 2Au + 6CO2

etc.

I hope answer clears up your question.

Steve


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## Lou

For what we're doing, it's all an electrons game. We're just pushing them from one nucleus to another. What we're concerned with are the valence electrons.

PM me aflac! I'd like to hear what's on your mind and have a try at answering some of your questions.


I think light (or any irradiation/energy input--even ultra sound) falls under catalyst :wink: Though that's mainly in organic where you have to worry about radicalization, but there are some things in inorganic that are light dependent... BUT the big one you forgot is equilibrium of the reaction!

I suppose surface area is an obvious one to include with population, since they are hand in hand.

Anyway, nice reactions...saves me the bothersome minute it takes to balance!! Thank you, Steve! Most of them seem like straightforward redox. But looking at the equations will give you many questions to think about if you've not had some chemistry courses...

Take what Steve said about gold III chloride, it's happy with those three chlorines around it, so why does it want to go react with HCl to make chloroauric acid? Why does it suck up moisture to the extent of it being hygroscopic. After all, doesn't it seem counterintuitive and against the common-ion effect? Shouldn't gold chloride be not so soluble in HCl like silver chloride (a population/equilibrium problem here). So why does HCl protonate the gold salt and make a very soluble compound, a weak acid but not do the same to silver? Why are some things soluble and others not a gram in thousands of gallons?

Remember, any aqueous AuCl3 you have isn't quite AuCl3 

It's not just an electron thing anymore. It then becomes more physical and there are physical propensities that each compound has. To me, a lot of chemistry is intuitive: if this is this, and this happens like that, well then shucks, this sure as hell better mean that this would act like this. <-- I can see where you get confused :shock: ) Still, there's an exception to every rule. But that's the fun of it  :wink: 

Sometimes I still want to :roll: when I think about it all.


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## Lino1406

If Mrs. Hoke uses the HCl/NH3 procedure
as purification method for Pd, then (a guess),
then undissolved particles by exess ammonia
are not Pd, but a contaminant, e.g. Ni
and should be discarded?
Lino1406


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## Harold_V

Reprocessed for potentail contained values would be a far better idea. If nothing else, I'd place such residues in the stock pot, where they'd be reduced to the elements and recovered (and purified) another day. It's not good practice to discard anything that isn't well known to be valueless. Good chance the remains would contain platinum, or other platinum group metals. 

Harold


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## aflacglobal

Combination is catalyzed by chlorine oxidizing the gold ( by oxidizing the gold you mean the chlorine excepted or stripped electrons from the gold ) Or is this promoted thru the use of an oxidizer. Or is the chlorine acting as the oxidizer. See this is where i get lost.to become positively charged (+1 and +3 states) and the negatively charged chlorine attaches to the positively charged gold (opposites attract). This process repeats three times for each atom of gold. 

Once combined the newly formed auric chloride (AuCl3), which is now neutrally charged with full electron orbitals, is soluble in the aqueous acid medium.


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## hungry

Hello everyone,
I am going to ask a question and I'm not going to say it is a stupid one because as you guys say there is only one stupid question, the one not asked.
I was wondering if Muriatic Acid dissolves gold and if it could what are the conditions that would need to be met to do so?
I am specifically referring to pins in a crockpot. I experimented with one pound of descent pins. It got to where it was hotter than I actually wanted it to be. I tested the solution with stannous chloride. It showed no color but I do not really trust my detection solution. As of now I do not have any gold in solution to test against. 
I am stuggling to get get a handle on all these processes and since I have no chemistry background at all, not even the prehistoric schools I attended I am in doubt about a lot. Although, I have absorbed a great amount of Knowledge this forum, I am lurking around religiously.
I have become totally committed, I have bought chemicals galore including my 70% nitric acid and every time I turn around I read here of another must have. I am enjoying even this part. I have bought or found a large amount of computer parts (the good stuff). I have karat gold, sterling jewelry and coins. Of course now I am PM rich (lol) and dollar poor so I am putting spending on hold for a while and try to get some results. I did not let it sink in soon enough that I should have CM Hoke's book at my side at all times, now it is going to have to wait for a while.
Well, I hope I have not bored everyone with my rambles and I also hope that it was alright to post this here. I saw a lot of talk about chemical reaction and it reminded me of my question.
That is all. Thank you for your time.
ED


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## Harold_V

If you have used HCl exclusively, and have not added any nitrates, you are safe to assume that no gold is in solution. Please consider that, even if some had been dissolved by some means, as long as there's any base metal present, it will have precipitated on the base metal as a dark colored mud, often not resembling gold in the least. From this you can conclude that it's wise to discard only known substances when you process the material coming from the material you process. 

One of the procedures for washing precipitated gold revolves around boiling the finely divided particles in HCl-----which can be safely accomplished using any level of concentration one desires, along with any amount of tap water. I used that procedure routinely for years with no negative consequences. 

You can safely assume that you have not dissolved any of your gold, but DO NOT DISCARD ANYTHING UNTIL YOU HAVE MADE A RELIABLE TEST. That should include making a standard solution. Remeber, some gold will lurk in the fine mud that you find----it, too, should be processed. 

Do yourself a favor. Buy Hoke's book, which will detail methods to make all of the solutions that would prove useful for you as a refiner. Operating without being able to test with certainty is akin to driving with a blindfold. Don't do it. 

Remember-----if you work without a known test, you have no idea what to save and what to discard. You'll quickly find yourself up to your hips in solutions and have no clue what to do with them. Testing by Hoke's suggested methods is 100% safe and sure. 

You might be better off to explore stripping clean pins instead of processing with HCl, or even investigate using nitric instead. Both methods are likely better suited to such material. 

BUY THE BOOK. 

Good luck,

Harold


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## hungry

Thanks Harold,I appreciate your input greatly You are right about being up to the hips with solutions as I have been doing a lot of experiments lately. Most of it is has had steel aluminum or zinc added to drop out the copper chloride and then neutralized with baking soda. I was wondering if I could some way refine it over again just to see if I could get different results as I am sure there is some PMs there.
Thanks
ED


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## Harold_V

hungry said:


> I was wondering if I could some way refine it over again just to see if I could get different results as I am sure there is some PMs there.


Hey Ed,

Well, there's certainly not much standing in your way, and it could prove to be a valuable learning tool to refine the wastes, but I think I'd give the idea some careful thought before jumping in with both feet. 

What you've done up to this point is lower the value of the materials, assuming you've extracted any of the gold. Traces now may be worth far less than the value of the time and acid, but perhaps that's what you're trying to determine. 

If I was to pursue this adventure, I'd approach it from the standpoint of refining the spoils from the stock pot. What you've done is put back in the elemental state, any metals that were present. You have also contaminated the material with substances that can be troublesome when you attempt to process chemically. It would be in your best interest to eliminate them, so you are sure that when you apply acid again, that the base metals alone go into solution. I'd start with several hot water rinses, to eliminate any chemical compounds that would be soluble, then I'd incinerate, heating each lot to a dull redness, which would eliminate carbonaceous materials. 

I'd then take a small sample and digest it with a few drops of water and nitric acid. Don't use any more acid than necessary to dissolve anything that's willing. When fuming stops, if the addition of a drop or two more acid makes no difference, it's work is done. The entire lot should then be taken up with water, allowed to settle, then decanted. Your target would be the remaining, undigested material. That would then be dissolved with very small amount of AR, then the solution tested, possibly after evaporating to limit free nitric acid. That often interferes with the tests you must perform. Hoke's book discusses evaporation 

By running a tiny sample, if there are no values to recover, you won't have wasted a lot of time and acid. 

I'd suggest testing the nitric solution for silver before discarding. If it's green, test for gold as well. If it's a distinct blue color, almost no chance it would have any values. 

One thing to consider. Once you've converted values to tiny particles, they often will be suspended in solution as tiny specs, usually black, or very dark brownish/purple color. Allow your solutioins to settle once digested, and look for a black or dark covering of material on the bottom. It may well be mixed with various white substances, depending on the nature of the material you've processed. 

You could also reduce the material in a furnace, fluxing properly. You'd still face the same problem of re-dissolving an amount of metal that may well not have any value at all, however, and furnace reduction tends to be quite hard on crucibles and the furnace lining due to the aggressive nature of the required flux. It's a good method, but costly-----best reserved to high grade wastes if possible. 

If each of your questionable samples started with a slight amount of material, so even if you had 100% recovery, you wouldn't get much gold, I'm of the opinion you'd be better off to simply discard the questionable material and chock up the experience to learning. Avoid recovering base metals along with traces of values from future lots by testing properly before hand, then discard things as quickly as possible. Our memories have a way of making us think garbage may well be worth saving, and it's usually wrong. When you come back to a solution, you'll always question your previous decision------and it will drive you nuts. If you test with fresh or proven test solution, and the solution in question is barren, it's time to get rid of it. You'll get more comfortable with this as you progress, gaining the necessary confidence not only in the testing solution, but your own judgment. I fully understand your reluctance to do so at this point. 

I mentioned buying the book. It will teach you what to look for regards the platinum group, plus how to properly test. It's the cheapest and best advice you can get from any source. I've mentioned time and again, I am nothing more than a high school graduate that took no chemistry classes------but I learned to refine, and ran a (very successful) commercial refining business for a little more than 20 years, all with knowledge I gleaned from Hoke's book. It's good-----very good------and written so the common man can make sense of everything mentioned. 

Keep us all posted.

Harold


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## Lino1406

Hello Steve,
Maybe the following reaction is
also of interest:

Au(+1) + ascorbic acid = Au(0) + dehydroascorbic acid

Lino1406
(N.B. ascorbic acid is called also: Vitamin C)


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## Irons

Here's a cheap source of Manganese Dioxide to make Chlorine gas as per Lou's suggestion.

Buy some Carbon zinc D batteries and take them apart. They contain Zinc metal (good), Ammonium Chloride as the electrolyte (good) and the black material that fills the battery is powdered Manganese Dioxide (good) and you have a useful Carbon rod for experiments.

Cut open the Zinc outer container and mix the Manganese Dioxide depolarizer with water to extract the Ammoniun Chloride.
The steel outer case goes into the stock pot.

It's all good, Baby!

Everything but the squeal.


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## lazersteve

Lino said:


> If Mrs. Hoke uses the HCl/NH3 procedure
> as purification method for Pd, then (a guess),
> then undissolved particles by exess ammonia
> are not Pd, but a contaminant, e.g. Ni
> and should be discarded?



According to Hoke, as Harold pointed out, the Higher PGMs will remain as a colored powder.

Most all of the base metal amines are soluble in the NH4OH solution and will remain dissolved (like dissolves like). You may recall the test for Nickel in solution involves basification with NH4OH followed by DMG addition to form the pink Nickel precipitate of DMG. So it follows that Nickel amines are soluble in NH4OH until the DMG is added.

Steve


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## Lino1406

Thanks Steve,
You opened my eyes.
I give here a more exact
equation for gold reduction
by ascorbic acid:

2Au(+1) + C6H6O4(OH)2 = 2Au(0) + C6H6O6 + 2H(+1)

Lino1406


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## lazersteve

Lino,

I see two possible minor drawbacks to using Asorbic acid to precipitate gold as you have demonstrated:


 Cost
 The indication in your reaction equation that it reduces Au +1

How much does Asorbic cost vs. other precipitants per gram? If you have plenty on hand already this is not an issue.

The second point above may be of more concern. Is asorbic acid equally effective at reducing Au +3 such as found in AuCl3 ? If so what is the molar ratio required to reduce a single mole of AuCl3?


Steve


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## Lino1406

Since I get my ascorbic acid (or VitaminC)
from pharmacy expiries the cost is not an
issue. The molar ratio is Gold:Ascorbic=2:1
And one doesn't need heating as against
oxalic acid. For Au(+3) I use copperas/bisulphite, 
so one still has the chance to be pioneer on that, 
my guess, it will work since Au(+1) and Au(+3)
have very near redox potentials. The expected
molar ratio is G:A=2:3
Lino1406


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## Rhodium

I had forgot about this thread. Thanks Steve.


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## goldnugget77

On page 106 of the Hoke book
she says when you add ammonium chloride to the palladium chloride
Nothing happens
I know that she is doing this to make a point 
But why does Harold do the same thing and Steve


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## lazersteve

What you are missing is that this is the first step of precipitating the palladium from the solution.

The next step is to add small crystals of sodium chlorate to the hot solution and stir. The palladium ammonium chloride precipitates if the solution if of the right concentration, temperature, and enough ammonium chloride is present in the solution.

Steve


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## golddie

> I tried Irons idea to simply add sodium hypochlorite 5% (bleach) to the ammonium chloride, HCl, and palladium solution. It worked like a charm. Looks promising for eliminating the the sodium chlorate as a precipitation reagent with Pd salts and complexes. There is still a very fine balance of ammoinum chloride and chlorine required.



Hi Steve!
Can you be a bit more clear.
Thanks


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## wicky

Harold_V said:


> I've threatened for some time to post a picture of samples of palladium that I kept from my refining years. Below, you will see five sample bottles, along with one of the new US $1 bronze coins.
> 
> The first specimen, left side, contains palladium that was precipitated with ammonium chloride and sodium chlorate. The salt was then dissolved in ammonium hydroxide and re-precipitated with HCL. The color changes from red to pale yellow, as you can see.
> 
> The second specimen is palladium precipitated with ammonium chloride and sodium chlorate.
> 
> The third specimen is palladium precipitated with dimethylglyoxime (DMG). Note the bright yellow color.
> 
> Fourth specimen is calcined palladium salt----in this case precipitated with ammonium chloride and sodium chlorate.
> 
> Fifth specimen is a small bead of palladium.
> 
> Harold



Hi 

Thanks for sharing all the knowledge you have...
I am not that much familiar with moler mass or molecular weight...
I been searching online to find the calculation or theory of Palladium present in hot air dried Palladium Dimethylglyoxime complex...

Or in simple way i want information about...
How much Palladium is present in 100 grams of Palladium Dimethylglyoxime? (70⁰ C hot air dried... Same as you shared 3rd sample)


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## rucito

Мultiply the weight by 0.3167


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