# Experiment precipitating gold out of solution 3 ways



## mls26cwru (Jul 29, 2013)

This weekend I tried an experiment with precipitating gold in a couple different ways. I have easy access to SMB, but was looking for a second way to more selectivity isolate the gold and get better purity. So, in my searching I found a couple different methods. I think the obvious one that comes to mind is oxalic acid precipitation. But after looking around a bit longer I decided on trying three different approaches... one of which i think was very interesting because there is not much info out there on the method. I will do my best to describe the conditions of each experiment

*Standard Oxalic Acid precipitation* (heating both solutions to near boiling and then mixing together)
_starting conditions:_ once refined gold dissolved in HCl/bleach
_method:_ auric chloride was produced, diluted, filtered. It was then heating on a hot plate until bubbles started to form in solution. A separate solution of oxalic acid was produced (using 1.5ounce of acid crystals to 1ounc AU ratio). The ph of this solution was adjusted to 6 using sodium hydroxide. This solution was then heated to near boiling as well.
_observations:_ upon adding the heated solutions together (AND SLOWLY) there was a lot of fizzing but no change to the solution. no gold was precipitating. I then prepared a sodium hydroxide solution by adding crystals to water (careful, this solution becomes hot). I added this to the hot auric chloride solution (VERY SLOWLY AS THE REACTION BETWEEN THE TWO IS VERY VIGOROUS!!!) until a reaction took place. It took a while, but when the gold finally started to precipitate, it did so very rapidly... i was rather amazed  
_results:_ happy!  actually, the gold powder that precipitated was the lightest tan I have seen so far. After washing the powder, the gold was melted and the button had no visible signs of tarnishing at all... definitely the purist gold I have produced so far. as an interesting side note, when the gold was molten, i noticed the mythical green hue that i have read about but never witnessed before 


*Oxalic Acid precipitation in conjunction with SMB*
_starting conditions:_ once refined gold dissolved in HCl/bleach
_method:_ after the auric chloride and oxalic acid solution was added in the above step, I set aside about 500 mL to precipitate with this method. No other ph adjustment was used, I simply added the standard ratio of SMB to the whole solution (1 to 1 by weight). 
_observations:_ precipitation was immediate. another observation was that the left over liquid was the clearest i have ever seen it... no hint of any color what-so-ever. also, the powder formed was the same tan color as the traditional oxalic acid drop. all rinses of the resulting powder were crystal clear as well.
_results:_ happy again!  The gold powder was indistinguishable from traditional oxalic drop. so was the gold button produced....and the green hue during melting was present once again. 


*Precipitation with Hydrogen Peroxide*
this was the most interesting in my opinion, and the one I was really looking forward to trying. I found the reaction on Wikipedia but almost no information on anyone actually doing it. here is the equation:
_(2)AuCl4[-] + (3)H2O2 + (6)OH[-] --> (2)Au + (8)Cl[-] + (6)H2O + (3)O2_
so, here is what i did:
_starting conditions:_ unrefined gold (clean foils, black powder)dissolved in HCl/bleach
_method:_ auric chloride was produced, diluted, and filtered. To this, I added the same volume of 3% H2O2. (i would be cautious about using strong h2o2). to this, i added sodium hydroxide in flake form. This was continued until the solution dropped the gold.
_observations:_ when the flakes were added, they would turn black immediately from a reaction taking place. as more was added, the solution would turn dark brown and gold would precipitate, but get dissolved back into solution. after continuing to add the sodium hydroxide, the solution eventually gave up the gold without putting it back into solution. the solution too a while to settle, probably due to how dilute it was.
_results:_ so-so. while the gold did precipitate, it ended up dark brown. (i can post pics if anyone is interested) This gold would have to be refined probably a couple times to get it as pure as the other methods. To me it seems as if SMB is much more selective as to what is reduced in the auric chloride solution. because of this and the fact that SMB is relatively cheap, I don't see any advantage to using peroxide over SMB. One plus I did observe was there was no nasty odors from this precipitation method...but this was still done under a hood. 


Just thought i would share with you guys... if you have any questions about the peroxide drop, please ask away and i will do my best to answer what i can 

Mike


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## solar_plasma (Jul 29, 2013)

Thanks for sharing! Really interesting. Somebody on the forum said, there is a reason, why there aren't many informations about some methods and you have showed, how true this is. Anyway, the chemistry itself is interesting. And who knows, if in future problems, materials and alternated settings maybe the worst method may become the only accessable.

I would like to see the pictures.


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## g_axelsson (Jul 29, 2013)

Interesting... number 3 was new for me, I hadn't heard about that before.

It feels a bit strange that hydrogen peroxide can both be used to dissolve gold and precipitate it, all depending on strength. That in it self would give a hint that it isn't that specific for gold.
The fact that you had to raise the pH to get the gold to drop could easily drop other hydroxides if you get the pH too high, to add flakes isn't the best way of doing it as it would raise the pH very high locally and some hydroxides isn't that easy to dissolve again. I think it would be less drag down if the NaOH is dissolved in water before added to the gold chloride.

Thanks for your description!

Göran


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## mls26cwru (Jul 29, 2013)

g_axelsson said:


> Interesting... number 3 was new for me, I hadn't heard about that before.
> 
> It feels a bit strange that hydrogen peroxide can both be used to dissolve gold and precipitate it, all depending on strength. That in it self would give a hint that it isn't that specific for gold.
> The fact that you had to raise the pH to get the gold to drop could easily drop other hydroxides if you get the pH too high, to add flakes isn't the best way of doing it as it would raise the pH very high locally and some hydroxides isn't that easy to dissolve again. I think it would be less drag down if the NaOH is dissolved in water before added to the gold chloride.
> ...



I was thinking along the same lines... I was also wondering if there is a specific ph 'target' that would help. I think I am going to try to do the second refining this way and see how it works, so there may be an addendum to this entry. I think another caveat is from my starting material... Since it was just gold fingers I worked on, i think that helps... if this was an AR solution with other metals dissolved, i think the result would be worse. thoughts?

Mike


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## solar_plasma (Jul 29, 2013)

Magnet stirring and adding the lye dropwise, - at all like if you would do a titration, could additionally help to prevent a locally too high pH.


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## Lino1406 (Jul 30, 2013)

The couples Au/Au+++, Cl-/Cl2, H2O2/(OH-)O2, have close oxidation potentials
therefore under different concentrations they may act as described above


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## lazersteve (Aug 4, 2013)

Dark colored Gold hydroxide is formed with the addition of NaOH to solutions of Chloroauric acid. The hydrogen peroxide is not required. The big drawback to this method is that nearly all metals are also precipitated as the metal hydroxide along with the gold at several different pH levels. 

If you have a very pure gold chloride solution and you add sodium hydroxide you will get dark gold hydroxide. If there are any other metals present, they will also report in the precipitate with the gold. The gold hydroxide can be melted directly to metallic gold.

I have experimented with hydrogen peroxide precipitation and was not pleased with the results. I did not add sodium hydroxide. I used the peroxide because it was listed as a precipitant for gold in DBC.

I have also used dry lye alone to precipitate gold hydroxide from a clean Chloroauric acid solution and the results was a nearly black powder that settled fast and melted into a shiny gold button. The source material (gold on quartz) had no other base metal contamination so I was not concerned with the formation of metal hydroxides along with the gold hydroxide.

Steve


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## mls26cwru (Aug 4, 2013)

lazersteve said:


> Dark colored Gold hydroxide is formed with the addition of NaOH to solutions of Chloroauric acid. The hydrogen peroxide is not required. The big drawback to this method is that nearly all metals are also precipitated as the metal hydroxide along with the gold at several different pH levels.
> 
> If you have a very pure gold chloride solution and you add sodium hydroxide you will get dark gold hydroxide. If there are any other metals present, they will also report in the precipitate with the gold. The gold hydroxide can be melted directly to metallic gold.
> 
> ...




I repeated my method tonight with the H2O2 and I have the same results as you Steve... The gold came out a little more pure, but the powder is still dark brown. maybe i will try to melt it in lite of what you mentioned about the gold hydroxide though. So far it has just been a lesson in chemistry more than anything else 

M


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## lazersteve (Aug 4, 2013)

Gold hydroxide is dark brown to black colored, hence the dark color. 

Steve


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## chlaurite (Aug 5, 2013)

Steve, relating to the need to further refine gold hydroxide - Everything I can find says that should dehydrate into a straight oxide at 140C, then decompose into plain ol' gold at 250C (making the melt no less pure than precipitating clean tan metallic gold).

I obviously defer to your experience, but have to ask - What contaminants _do_ you end up with if you try to melt gold hydroxide directly?


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## g_axelsson (Aug 5, 2013)

It all depends on the content of the original solution. Obviously the addition of hydroxide to a solution of metals will precipitate most other metals as hydroxides too, not only gold.

So whatever you start with together with the gold will end up in your melt.

Göran


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## GoldHappy (Aug 27, 2013)

I'm curious....did you use a standard flux with your gold powder or did you simply super heat the powder to bring it back to metallic form?


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## butcher (Aug 27, 2013)

Flux ( from the word flow) is normally used to change the chemistry of the melt, or the properties of the melt, like making it easier for the metal flow in the melt and combine together, different ingredients are used in a flux to give the desired results, these ingredients react with heat and can change the chemistry in the melt, such as oxidizing metals or reducing metals, or to create a glass slag that will absorb oxidized metals, some ingredients of the flux help to make the melt more liquid where precious metals can flow and collect together, sometimes the flux will have collector metals or metal oxides like litharge PbO to help collect the precious metals in a melt.

There are no "standard fluxes" that will work with any material,or in every situation, as when flux is used every situation will be different requiring a different recipe, or a different chemistry needed in the melt to get the desired result, some flux recipes are for a basic flux, or stock flux (these may be called a standard flux), that you add ingredients to so that you change the properties of the stock flux for each type of material you are melting.

With pure gold powders you need no flux to melt it or to change the chemical properties during the melt, with gold all we need is a sprinkle of borax to lightly coat the dish so the gold does not stick to the dish, and can roll around in the dish easier when melted or pour the molten gold from the dish.


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## GoldHappy (Aug 29, 2013)

butcher said:


> There are no "standard fluxes" that will work with any material,or in every situation, as when flux is used every situation will be different requiring a different recipe, or a different chemistry needed in the melt to get the desired result, some flux recipes are for a basic flux, or stock flux (these may be called a standard flux), that you add ingredients to so that you change the properties of the stock flux for each type of material you are melting.




Thank you, Butcher. Action Mining calls their 'stock flux' standard flux, so yes, that's what I was referring to. I'm still learning about changing the properties of the flux for the various types of material. I'm fascinated by the entire process and there's a lot to learn.


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## butcher (Aug 30, 2013)

Flux
from a few note I have:

Normally before fluxing we need to determine whether the ore we are melting is a neutral ore or if it is oxidizing or reducing in nature, here the chemistry of the ore itself can determine what flux recipe is needed, the ore itself can oxidize or reduce metals in the melt, to determine this normally a melt a test with a certain portion of litharge (lead oxide), it is used to determine the oxidizing or reducing nature of the flux, lead oxide itself can be an oxidizer for some metals or carbonous materials, this test melt with a predetermined amount of lead oxide to see whether the lead oxide reduces back to lead metal or how much is reduced to metal can help us determine what chemicals we may need to add to our flux, whether we need an oxidizing flux, a reducing flux or a neutral flux, this melt can also help us to determine the viscosity of the melt, remember the word flux came from the word flow, if we cannot make the melt fluid where the very fine particles of values like gold will flow around in the melt where the can join together into one molten button,if not they may end up as microscopic balls of gold locked up in the slag.

Certain metals , oxides, salts or ore, have a reducing or oxidizing power, some determination can be made from knowing these powers to help figure what flux to begin with as the flux itself has oxidizing or reducing powers.

Melting Metal powders, oxide hydroxide or salts of metals, Knowing how these metals react in a melt chemically can also help us to determine the flux needed, some examples:

Silver chloride will oxidize silver in the melt and much of our silver will go up in smoke, a flux high in soda ash (sodium carbonate can help to reduce the silver back to metal. 

When meting a sulfide where adding iron to the flux can help to convert sulfide to a form that will burn off as SO2 gas releasing the metal from its bond. 

Using litharge on a very carbonous ore to burn off the carbon as CO2 gas. 

Adding a carbon source to a metal oxide to reduce it back to metal.

Adding silica (crushed glass, sand or quartz) to a melt to offset the effects of lead oxide or other flux or ore from attacking the silica in our melting crucible, and also to help form a glass slag which will absorb the oxidized base metals into its bond.

Adding an oxidizing agent to oxidize base metals, to remove them in the slag while our more noble metals coagulate and sink to the bottom of the melt.

Adding a collector metal to the flux help collect the values, which would not form a button with this metal gathering them together in the melt (like using silver to collect the platinum group).

Borax glass Na2B4O7
viscosity, forms slag, (strong acid) to form oxides of base metals, lowers fusing point of slags.
can be hard on refractory.
(can be found in mule team laundry detergent but it contains moisture and will be very fluffy in melt) (anhydrous borax glass dehydrated form best)

Calcium fluoride CaF2
Used in small quantity.
Use more if aluminum in ore or melt.
increases fluidity.
It eats crucibles, or refractory.
(can be found in rocks)

sodium Carbonate Na2CO3 (soda ash, washing soda)
forms alkali sulfides in presence of air, can be considered a de-sulfurizing or oxidizing agent,
Emits free alkali, this with NaCO3 forms silicates and aluminates.
Can help to reduce silver chloride to metal in melt.
(can be found in automatic dish washing soaps)

Litharge PbO (lead oxide)
Reduces to lead metal.
Can act as a collector of values.
Can act as an oxidizing agent (oxidizing carbonous material).
Has a strong affinity for silicates, so it can dissolve some of the crucible if there is not enough sand glass or quartz (silica) added in the flux) so addition of silica may be needed to protect crucible.
(can be made)

Silica SiO2 (crushed glass, sand, quartz)
Strong acidic reagent in melt.
Combines with metal oxides to form silicates (slag).
Added when charge deficient in silica (to form slag or protect crucible).
Gives fluid to melt.
Can help to protect crucible from effects of litharge.

Flour 
Sugar
Charcoal 
wood chips 
Carbon source
reducing agent.
can help where oxides are high.
can also be used to reduce base metal oxides to metal.

potassium nitrate KNO3
Oxidizer, give oxygen.
Oxidizes sulfur.
Can oxidize base metals.
Gives oxygen to help oxidize sulfide ore.
(can be found in fertilizer)

A Basic flux is adjusted in recipe as needed to get desired results in the melt.

A Basic ore flux
15g sample
30g Na2CO3
35g PbO litharge
4g silica or crushed glass (vary depending on silica content of ore)
35g Na2B4O7
1g CaF2 (vary depending on aluminum content of ore)
3.2 g flour

Flux Quartz Ore
15g sample
29g Na2CO3
40g PbO
No silica needed
3g Na2B4O7
no CaF2
3g flour

There are many basic recipes, or house fluxes for ores or metal powders, these are normally just starting places to help determine what is needed to be able to adjust that basic recipe to get the desired result.
no one recipe will fit all chemistry going on in the melts, so in a way the melt chemistry determines the flux recipe.


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## Marcel (Aug 31, 2013)

I am uncertain about the mechanism that will precipitate gold from auric solution using H2O2 and OH.
My suspicion is that this is more or less a dillution that will affect the PH, which then will make the metals drop.
So the same effect should be achieved if using H2= and diluting the solution until the PH reaches a point, where it no longer can hold the metallions.
Following your equitation
(2)AuCl4[-] + (3)H2O2 + (6)OH[-] --> (2)Au + (8)Cl[-] + (6)H2O + (3)O2

We see that oxygen and chlorine excapes up in the air and a lot of water is formed.
Water is a very special liquid that can behave like and acid or caustic depending on whom it meets in the solution. It is neutral but will weaken any PH into the opposite direction. 
Just my thoughts on the H2O2 and H2O precipitation. No proof. Correct me, or add yours.


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## seriouscash (Aug 31, 2013)

I'm interested in seeing the pictures of your gold dust


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## butcher (Aug 31, 2013)

(2)AuCl4[-] + (3)H2O2 + (6)OH[-] --> (2)Au + (8)Cl[-] + (6)H2O + (3)O2

In the above equation, I do not see any chlorine Cl2 gas, I see the chloride Cl[-] ion in the solution of water, and the precipitated gold.

Diluting an acid will not change the pH of the acid.


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## g_axelsson (Sep 1, 2013)

butcher said:


> Diluting an acid will not change the pH of the acid.


Yes it will!

A highly concentrated acid has lower pH than the same acid diluted with water.
But you can't dilute enough to go from acidic to basic or the other way around as water has pH 7. Any dilution just brings it closer to pH 7.



Marcel said:


> I am uncertain about the mechanism that will precipitate gold from auric solution using H2O2 and OH.
> My suspicion is that this is more or less a dillution that will affect the PH, which then will make the metals drop.
> So the same effect should be achieved if using H2= and diluting the solution until the PH reaches a point, where it no longer can hold the metal ions.


Dilution with water to precipitate certain compounds happens in a few cases, silver chloride in strong aqua regia or copper chloride in HCl for example, but in both cases it is the compound that is precipitated out of solution, not a reduction of the compound to a metal.

I can't comment on the equation if it is feasible as I haven't tested it or calculated if it is possible. (Calculating the Gibbs free energy if my memory isn't failing me.) The main problem I see is that HCl + H2O2 dissolves gold (this I have tested) and the equation is creating HCl while reducing the gold. At least there should be an arrow towards both sides to show that this is a reversible reaction. The equilibrium point could be pushed towards either gold or gold chloride by changing the pH of the solution.
Hmmm... I think I managed to convince myself how hydrogen peroxide both can be used to dissolve gold and precipitate it again. It's a reversible reaction that can be pushed towards either side by changing the pH of the solution.

A problem with basic solutions though is that it will drop other metals as hydroxides and pollute our gold.

Göran


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## butcher (Sep 1, 2013)

Thanks Göran, for correcting me.

I was wrong in my statement; yes you can slightly change pH with large volumes of water.

Maybe I should have said that differently, dilution with water has very little effect on pH of an acidic solution (or an alkali solution), as it can only slightly change pH, and would take extremely large volumes of water to change the pH to any degree, and most of these metals will not precipitate from the acid until you get to the alkali side of the pH scale, and normally fairly high in that alkalinity side before they begin to precipitate as hydroxides, and you cannot add enough water to move an acid past neutral (pH 7) to the basic side of the scale where the metals would normally precipitate from solution as hydroxides, as water has a pH of about 7, most metals soluble in acidic solutions will not precipitate until on the alkali side of pH scale.

Yes you could slightly change a pH with water, but you will have to add ten times the amount of water to the acid to change the pH by one-pH unit, so to move from a strong acidic solution of a pH 1 to a strongly acidic solution of pH 3 we would have to add about twenty volumes of water, thus changing hydrogen H+ ion concentration and changing the pH with this large volume of water.

At this point I am not convinced the effects of AgCL2 forming AgCl upon dilution of an highly concentrated chloride solution is caused from the effects of a pH change. And likewise the precipitation of CuCl when we dilute the concentrated copper chloride solution I am not convinced the slight or minor pH change is the cause, I believed it was from the chloride ions being diluted in solution, (but I can also be wrong in this belief), than it was from the hydrogen ion being diluted.

Water itself can act as an acid or base in nature in solutions.

I assumed the Hydrogen peroxide worked as strong oxidizing agent in acidic solutions, and a strong reducing agent in alkaline solution,


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## g_axelsson (Sep 1, 2013)

butcher said:


> Yes you could slightly change a pH with water, but you will have to add ten times the amount of water to the acid to change the pH by one-pH unit, so to move from a strong acidic solution of a pH 1 to a strongly acidic solution of pH 3 we would have to add about twenty volumes of water, thus changing hydrogen H+ ion concentration and changing the pH with this large volume of water.


Butcher, I agree with you on all but one tiny detail... pH 1->2 takes 10 times volume (add 9 parts water), pH 2->3 also takes ten times (also add 9 parts) but now we're starting with the larger volume... 10*10 or 100 times the volume to go 2 steps (9/10 + 90/100 parts water, 1 part acid = 99 times the volume water added).

This is of course a simplification as we don't count the pH 7 of the water we are diluting with so actually we need even more than the 99 parts. But for strong acids or bases a dilution by ten moves the pH one step towards neutral is a good enough rule.
Anyone interested in the exact solution can look up the logarithmic laws and the definition of pH.

I know, I'm a besserwisser (a know-it-all for you English speakers)! :mrgreen: 

Göran


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## butcher (Sep 1, 2013)

Göran, 


Thank you again for the clarification and the math involved with the pH change and the volumes of water needed to make the change, I have not tried it but I have read where they take an acid and add ten volumes of water to raise the pH by one, then take a portion of that and add ten times water to this diluted acid raise that one more value in pH.

I guess I kind of generalize in my mind when I add water to a solution I am basically just diluting the ions in solution, and it has basically little to no effect on pH (although in reality there is a very slight effect on pH).

I have to admit you do keep me on my toes, and keep me thinking or rethinking my ideas, this also very good to help us both learn more or at least to help apply some of the things we learn.

I know I am not a besserwisser (Know it all), it is hard for me to see past what I think I know, until I can learn differently.

I also do not think you are a besserwisser, other wise you would not have learned as much as you have.


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