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Halide Leaching

The group of chemicals known as halides is composed of the elements fluorine, chlorine, bromine, and iodine. We will ignore fluorine because it does not dissolve gold. The elements listed above are in order of their molecular weights, density, cost, and reactivity. Chlorine is the lightest, cheapest and least reactive while iodine is heaviest, most expensive, and fastest reacting. Iodine is classified as a rare element. There are no ores containing iodine. It is obtained primarily from extracting seaweed, hence, the high cost.

Chlorine was the first commercial gold extracting agent. In the late 1800’s it was introduced on the market as " DuPont Mining Salts". Although I have no direct proof, I’m sure that this was nothing more than sodium hypochlorite, the stuff that Clorox is made from. It was used by taking ore and chlorine and agitating them in a "Pachuca Tank" or rotating the mixture in lead-lined barrels. This method fell into disfavor with the discovery that cyanide extraction was cheaper and easier for the untrained miner to use. The usefulness of various extraction techniques for the production of gold is attested to by the fact that at the present more than 80% of the world’s gold production is obtained by chemical extraction.

In recent years there has been a resurgence of interest in halide extraction of gold. Partly because of a better understanding of the process but mostly because of the bad rap that cyanide has received at the hand of the do-gooders, zealots and ignorant crusaders. I understand that there is at least one commercial companyare selling a sort of "magic" proprietary chemical that only they can provide at a cost that knocks my socks off. Let me give you a truism about chemistry, MERLIN THE MAGICIAN DIED. THERE AIN’T BEEN NO MAGIC SINCE!!! You can do the same thing in your carport or out in the mountains. All you need is a little information that average Joe Beerdrinker just doesn’t have ready access to. I am going to give you (I hope) all the information you need to do it. So, stay with me.

If you will read through all this stuff which, I know, is more than you ever wanted to know I think you will eventually agree that you will have learned some principles that you can use to great advantage. I really hate to see folks trying to follow a cookbook recipe to the point that when something goes a little wrong it wrecks everything. I'm going to try to give you enough information so that if something unexpected happens you will be able to deal with it rather than panic.

Leaching
O.K., I hate to try to explain the principle of oxidation/reduction or "redox" because I have never been able to put it into words that anyone can understand. But, we have to try to get at least a rudimentary understanding of this or you will never really understand what is going on in gold extraction or any form of metallurgy.

There are three ways to look at redox. Oxidation is:

  • The addition of oxygen.
  • The removal of hydrogen.
  • The loss of an electron.

Now I know that doesn’t mean much to you but it is important. For example when you recover mercury from a solution of nitric acid which you have used to clean some amalgamated gold you usually simply hang a piece of copper or aluminum in the clear solution and mercury plates out on the copper and drips to the bottom of the vessel. In this case the copper is reducing the mercury in mercuric nitrate to mercury metal and oxidizing the nitrate to copper nitrate. When something is reduced something also has to be oxidized, they are simply the opposite of each other. This is how a battery works. In the case of metals, which are cations, that is they carry positive charges, the pure metal is the reduced form. That is, iron metal is the reduced form of iron. The rust that is on it is it’s oxidized form because it has combined with oxygen to make iron oxide. If you take iron oxide and put some rust remover (reducing agent) on it you remove the oxygen and reduce the metal back to pure iron. Now, halides are not metals. In fact they are the opposite. They carry negative charges and are anions. So, In their elemental form they are oxidized. Just the reverse of metals.

Why is all of this important to you? Because if you are going to use these chemicals you must be able to use them. If you want to store iodine, it is much better to store it in its reduced form so that it won’t evaporate. And, if you want to dissolve gold with it, it must be in its elemental or oxidized form. To summarize, you must be able to shift the iodine (or other halide) from its oxidized to reduced form. I’m going to tell you how to do it.

When you want your halide in it’s elemental or oxidized form, simply add a little Clorox bleach. If you add an excess, Iodine will precipitate from the water solution and settle to the bottom as pure elemental, blue, iodine. If you want to dissolve the iodine, you must reduce it. This you can do by adding a little Red Devil Lye dissolved in water. The blue iodine will begin to go in solution so that the solution will first turn bright blue and upon addition of more Lye it will become colorless as all of the iodine is reduced to sodium iodide. Hey. Great! Now we know how to convert/handle it. As my son tells me re computers, you can make it do what you want it to do. Now you have control.

Now for a couple of facts re iodine. Elemental (oxidized) iodine will not dissolve in water. So if it won’t dissolve, how are we going to use it? It has another characteristic. It will dissolve in solutions of sodium iodide. I didn’t do it, it just works that way. That fact gives us our way to dissolve gold. We will pursue that a little later.

I should say that there are differences in the way the halides behave but they are very minor differences. For example, chlorine and bromine will not precipitate from solution when oxidized. That makes them a bit more difficult to handle and recover, as we will see later. I am simply using iodine as my example.

Recovery of the Gold
Now we have our ore with a nice red-brown solution of iodine and, hopefully, a whole bunch of dissolved gold in the solution. We can’t sell it that way. Got to get it out and make it look like gold. First we have to get rid of all the material that we have been extracting. This can be done in a number of ways but probably the most Basement way is to filter it. Depending on the scale that you are working you can use a coffee filter in a funnel or a piece of canvas in a 20 ft dia tank. Anyway, filter it and try to get a rather clear solution. Remember that the solution must still be red-brown. If it isn’t, you have left your gold behind in the filter. This is why I insisted on your being able to oxidize or reduce at will. KEEP IT RED-BROWN.

Once you have your red-brown solution free of material, now, you can let it go colorless or make it go colorless by addition of the Lye solution. Your gold will now slowly settle to the bottom as a black powder. Or, you can filter the solution through a fine filter to recover the gold.

Great, we got a whole bunch of gold but if we lose that iodine we are still going to be in the hole. Remember, at this point it don’t look like iodine but it is still there. By this time you probably have a pretty good volume of liquid with the iodine in it. If its more liquid than you want to deal with simply dump in an excess of Clorox , let the iodine settle to the bottom, pour off most of the water, add some Lye solution and you have your sodium iodide in a concentrated solution which is the way you want to store it anyway. Ready for another extraction.

Of course, no one is going to buy that black powder from you. " Sure it’s gold, anyone can see that! Take it somewhere else"! To make this stuff look like gold again simply smelt it. I assume most of you know how to smelt. If not, please contact me.

Hey guys, It ain’t brain surgery. You can do it as well as me. Please let me know how it comes out or if you have problems.

The other two Halides
Now, I have talked almost entirely about iodine. 99% of what I said about the use of iodine applies to both chlorine and bromine. The significant differences are;

· The price as discussed previously.

· Solutions of chlorine are colorless so you cannot rely on the color to tip you off as to the eH of the solution. With chlorine you will have to use an eH indicator dye which will not be readily available, or (would I leave you with that problem?), you can use your nose instead of your eyes. You will have to add oxidizer to maintain a little chlorine smell coming off the extraction.

· So, what oxidizer do I use? Same as with iodine. You use chlorox. However, chlorox is very weak and you would have to use great volumes of it to maintain some chlorine in solution. You should use the stuff that chlorox is made of, sodium hypochlorite. This is a solid, white powder which in a 5.0% solution is chlorox. This chemical can be obtained at any chemical supply. Now if you can’t or won’t go to that effort there is still another alternative. You see, we never leave you hanging without another way to skin the cat. Almost as good as sodium hypochlorite is our old friend, swimming pool chlorine. This is calcium hypochlorite. It works very well. It just doesn’t dissolve very well. Again, we never leave you without a way to escape. Just add a little lye solution and it will go right in. You have converted it to sodium hypochlorite. Don’t add too much, just enough to put it in solution.

· Remember chlorine is very slow. You will have to maintain this system for several hours to a day or more.

· Like all of the halides, if you allow the solution to go acid, the chlorine will rapidly boil off and if you are close by will be extremely uncomfortable. Bromine and iodine are not so bad. They are not nearly so volatile and will give you a little more time to rectify the situation by addition of a little lye water.

· If you are extracting with halides, in particular, chlorine where you have no visual reference as to what is happening, you should have some hydrochloric (muriatic) or sulfuric acid at hand. If your reaction should start to slow down and you are sure you have an excess of halide in solution you might have to add a little acid in order to liberate the halide from it’s salt form in order to keep sufficient free halide to ensure a good extraction. If you can keep the pH at say 8.5 you will be about right.

· If chlorine production should get out of hand, you probably should have some solution of sodium thiosulfate on hand. It can be bought from any chemical supply. This is the stuff that tropical fish freaks use to treat tap water to destroy chlorine. You will need much more than they use however. This is a good way to neutralize any solution you wish to dump. Your neighbors will probably appreciate your thoughtfulness.

· Just a little about bromine. It’s a bit expensive so we are not going to throw it away, right? Bromine is rather nasty stuff to handle in its elemental form. It is a liquid, which is sold sealed in 1-LB glass ampoules. When you break the container you have to use it or contain it some other way. It has a way of fuming right through the tightest bottle-cap type seals. The only way you might do it successfully is to seal around the cap with hot paraffin. Let’s don’t do it that way. Always an option, right? Since we are rather knowledgeable about oxidation/reduction we can handle it in a more sophisticated way. We can buy it in its reduced and cheaper form, sodium bromide. After all, we know how to convert it to elemental bromine; it just takes a little clorox, right? After all, that’s what we are going to store it as anyway. When Clorox is added and the bromine oxidized to its active form, like iodine; it exhibits a distinctive color, deep red. We are now a bit smarter than the average bear

Now that we can jump around in redox reactions like a monkey in a fig tree let me add just one final "consejo". Whatever you decide to do with these extractions, please. Read the above several times until you are certain that you understand it well enough to control the reactions. It’s not hard, use common sense, and if thing start to go sour don’t panic, fix it.