Acetic Acid (Ethanoic Acid, Vinegar, Glacial Acetic Acid)By Barry Flannery (edited by Dr. Stephen Moreton) Acetic acid,CH3COOH, is a simple carboxylic acid. It is a weak organic acid better known as the chemical which gives vinegar its distinctive odour. Its weak acidity makes it a useful chemical for dissolving calcite from sensitive minerals such as galena and fluorite where the much stronger hydrochloric acid may cause dulling and damage to lu..

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Acetic Acid (Ethanoic Acid, Vinegar, Glacial Acetic Acid)

By Barry Flannery (edited by Dr. Stephen Moreton)

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Acetic acid, CH3COOH, is a simple carboxylic acid. It is a weak organic acid better known as the chemical which gives vinegar its distinctive odour. Its weak acidity makes it a useful chemical for dissolving calcite from sensitive minerals such as galena and fluorite where the much stronger hydrochloric acid may cause dulling and damage to lustre. It is very slow to act and specimens will usually require several days of treatment to produce an acceptable result. Dilute (<10%) acetic acid is preferred for particularly sensitive minerals, or by beginners due to its safety and availability. If obtainable, orthophosphoric acid is far superior in terms of speed, and is similar to acetic acid with regard to its effect on sensitive minerals. It is not, however, as widespread and easy to obtain. 

Acetic acid will dissolve calcite by the following reaction producing the soluble salt calcium acetate, water and carbon dioxide (the fizzing).

2CH3COOH + CaCO3 -> Ca(CH3COO)2 + H2O + CO2

In an ideal system 2 moles of acetic acid will dissolve 1 mole of calcite. In layman's terms, given enough time one litre of 5% solution (50g/L) will be able to dissolve ~40g of calcite. In physical terms, 40g is roughly equivalent to a 2.5cm (1'') cube of calcite. 

Directions for Use

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Acetic acid is commonly found as some form of household vinegar, which will usually have a concentration of 3-7% acetic acid. Ordinary vinegar for food use can be used for dissolving calcite straight out of the bottle. However the price, added sugars (to improve taste) and slow reaction rate make it less than ideal for dissolving large amounts of calcite. It should be noted that increasing the temperature greatly increases the rate of reaction, however it also causes the acid to fume and a pungent smell of vinegar will soon fill the room, so good ventilation is essential. Higher concentrations of acetic acid do result in faster reaction times, however the dangerous nature of the concentrated acid offsets these gains (see Safety Precautions) and so it should NOT be used.

Safety Precautions

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At low concentrations (<5%) acetic acid is relatively harmless and used as the food additive vinegar. However as the concentration increases, acetic acid becomes far more dangerous and has a powerful pungent odour. At high concentrations (>25%) acetic acid is corrosive and should be handled with the same care and safety precautions as stronger, more dangerous chemicals, such as concentrated hydrochloric acid. At high concentrations it should be used in VERY well ventilated areas, ideally outside or in a fume hood as the vapours are corrosive, EXTREMELY pungent and harmful if inhaled. Wear gloves and goggles if handling it.

EU Hazard Classification

Concentration by weight Molarity Classification R-Phrases 10–25% 1.67–4.16 mol/L Irritant (Xi) R36/38 25–90% 4.16–14.99 mol/L Corrosive (C) R34 >90% >14.99 mol/L Corrosive (C) Flammable (F) R10/R35

Examples

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This specimen of fluorite was treated with dilute HCl(~10%). There are clear signs of dulling and damage to the lustre of the fluorite.
 

This specimen was treated with vinegar (~5% acetic acid) which has left the lustre completely intact. 
 

Cleaning Galena with Acetic Acid by Peter Haas

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There is no way to restore the lustre of galena. The tarnish usually consists of cerussite, anglesite, hydrocerussite or a mixture thereof. That is, the tarnish forms by corrosion of galena and therefore, the surface won't become brilliant even if the tarnish is removed. 

Though, the specimen will probably look a bit better when you manage to remove the tarnish, but this is also a problem. Anglesite is almost insoluble even in concentrated mineral acids. The only exception is concentrated sulfuric acid, but this acid has oxidizing properties and will corrode your specimen further. A dilute sulfuric acid is not oxidizing, but it also doesn't dissolve anglesite. As for the carbonates, they are dissolved by acids, but many acids (the mineral acids in particular) form lead salts which are even less soluble than the carbonates. Again, there are some exceptions, one of which is nitric acid. Unfortunately, nitric acid is an oxidizing acid as well and will convert your specimen to anglesite, so it will look even worse after the treatment. 

In my opinion, the only chance to make the specimen look better will be a treatment with acetic acid. You should use commercial white vinegar (a 20% w/w aqueous solution of acetic acid) for that purpose. The treatment will take some time, since acetic acid dissolves lead carbonates rather slowly, so you'll have to leave the specimen in the solution for at least several hours. Make sure that the matrix (it might consist of carbonates as well ...) survives the treatment. Afterwards, rinse the specimen with cold water. You might also be able then to remove residues of the tarnish with a toothbrush. Do not expect too much: the crystals will never become brilliant for the reasons explained above. 

DISPOSE OF THE USED ACETIC ACID SOLUTION PROPERLY: lead compounds, like all heavy metals, are VERY toxic, and even more so, if they are soluble. 

DO NOT HEAT THE CLEANING SOLUTION. The effect of heat is negligible with this type of reaction. The cleaning process will not become any faster by heating. Acetic acid is volatile and the inhalation of its vapours will harm your respiratory tract considerably.

Addition Information Links

 

General information on acetic acid http://en.wikipedia.org/wiki/Acetic_acid

 

Formic Acid (Methanoic Acid)

By Barry Flannery (edited by Dr. Stephen Moreton)


General Discussion

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Formic Acid, HCOOH, is the simplest carboxylic acid and commonly found as the venom in ant and bee stings. Its chemistry and uses to the mineral collector are effectively the same as acetic acid and it can be used to dissolve calcite. Formic acid is a stronger acid than acetic acid and so it reacts more readily with calcite, dissolving it faster. Nonetheless it is still a weak acid and slow working, usually taking several days to produce an acceptable result. However formic acid is not widely available to the average collector and so is not usually used.

Formic acid will dissolve calcite by the following reaction producing the soluble salt calcium formate, water and carbon dioxide (the fizzing).

2 HCOOH + CaCO3 -> Ca(HCOO)2+ CO2+ H2O

In an ideal system 2 moles of formic acid will dissolve 1 mole of calcite. In layman's terms, given enough time one litre of 5% solution (50g/L) will be able to dissolve ~52g of calcite which is slightly more than acetic acid. In physical terms, 52g is roughly equivalent to a 2.7cm (~1'') cube of calcite. 

Directions for Use

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A solution of <10% (100g/L) should be used as the acid becomes more dangerous and corrosive at higher concentrations. Heating will increase the rate of reaction but vapours are pungent and irritant so it must be done in a WELL ventilated area, ideally a fume hood. 

Safety

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As with dilute acetic acid, dilute formic acid is used as a food additive but at high concentrations it becomes both hazardous and corrosive and should be handled with the same care and precautions as concentrated hydrochloric acid. At high concentrations it should be used in VERY well ventilated areas, ideally outside or in a fume hood as the vapours are corrosive, EXTREMELY pungent and toxic if inhaled.
Concentrated formic acid is extremely toxic if ingested and is partly responsible for optic nerve damage caused by drinking methanol (‘’blind drunks’’) as it is one of the metabolites of methanol. 

EU Hazard Classification

Concentration by weight Classification R-Phrases 2%–10%% Irritant (Xi) R36/38 10%–90% Corrosive (C) R34 >90% Corrosive (C) R35

Additional Links

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General Information on formic acid http://en.wikipedia.org/wiki/Formic_acid



 

EDTA (Chelaton III)

by Dr. Peter Haas

General Discussion

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The active chemical belongs to the group of aminopolycarboxylic acids. That is, the backbone is a low-molecular weight aliphatic amine or polyamine that has been exhaustively carboxymethylated. Since the synthesis is relatively simple to perform at an industrial scale, there is a vast variety of compounds available. Though structurally similar, their complexing properties vary widely, so it is important to know which compound you're dealing with. 

Second, the chelants themselves (i.e. the free acids) are very weak acids and they're almost insoluble in water. Therefore, they're traded in form of their sodium salts. Depending on their chemical nature, there may be different salts available of one complexant, and they have different trade names. If my memory serves me correctly, Chelaton III is the disodium salt of "ethylene diamine tetraacetic acid" (a pseudo-chemical trivial name), also known as EDTA. This is not well suited for the purpose of mineral cleaning, since its solubility in water is still rather low. The tetrasodium salt is more easily handled. 

When the salts are dissolved in water, hydrolysis occurs, because water is a stronger acid than the complexant, and one ends up with a strongly alkaline solution. For instance, the pH of a 1% (10 g/l) solution of the EDTA tetrasodium salt (M = 380 g/mol) is well above 12 ! In these conditions, many minerals are unstable, among them many silicates, most secondaries of metal ions with amphoteric hydroxides (copper, chromium and others precipitate as hydroxides at pH values around 10, but re-dissolve as stable hydroxo complexes when the pH is further increased), but also some oxides and sulphides. For instance, chalcopyrite already starts to decompose at pH values around 9. The latter is due Fe2+ being oxidized to Fe3+ by dissolved oxygen and formation of an iron hydroxide precipitate. The driving force in this reaction is the particularly low solubility of iron(III) hydroxide, reducing the equilibrium concentration of Fe3+ tremendously and thereby shifting all preceding equlibria (reduction of the Fe2+/Fe3+ oxidation potential, virtual increase of the solubility of chalcopyrite (!) etc.). This holds for many other iron-bearing minerals as well, and it is also the reason why iron hydroxide stains are so difficult to remove in general: the acid alone is not of much help, because iron(III) hydroxides start to precipitate in acidic solutions (due to the low solubility product, the hydroxide concentration in acidic solution is still high enough to exceed the saturation limit). Removing iron stains requires to reduce most of the Fe3+ to iron Fe2+, whose hydroxide is much better soluble. 

Aminopolycarboxylic acids are not redox active. Also, although they form very stable complexes with alkaline earth and transition metal ions, their Fe3+ complexes are still not stable enough to prevent iron(III) hydroxides from precipitating, even in weakly alkaline conditions. Hence, they DO NOT WORK ON IRON STAINS at all. Oxalic acid, on the other hand, does, becaues it can reduce Fe3+ to Fe2+. This is just one of the many differences between these compounds. 

The need to work in strongly alkaline solutions is a severe drawback of the aminopolycarboxylic acids, and limits their use to particular applications. One may acidify the solutions, but this results in the precipitation of most of the complexant in form of the free acid. Solutions saturated in this way are still useful for some applications, but they are weak. On the other hand, they are also very selective (as a rule of thumb in mineral cleaning, low reactivity means high selectivity) and this may present an advantage. I should add that heating the solutions does not provide an advantage. At elevated temperature, the solubility does not markedly increase, but the stability constants of some complexes do somewhat decrease, so this is even counter-productive. 

However, the alkaline environment, combined with the low acidity of EDTA, may also turn out as an advantage. For instance, the solution does not affect calcite, although EDTA forms a very stable complex with Ca2+, and the same holds for other minerals that are otherwise sensible to acids. This is because EDTA does not assist the dissolution process. It is a strong complexant, but has otherwise almost no reactivity. This is another important difference between oxalic acid and EDTA. 

I use EDTA quite often, because I don't like specimens that are depraved of any signs that still relate them to the environment they were found in. EDTA enables me to improve the appearance of metallic minerals without doing any harm to secondary minerals, but this requires profound chemical knowledge. Also, there are very few minerals for which it produces outstanding results: these are just native copper and some copper sulphides and sulphosalts. In fact, if you do not specifically collect copper, chalcocite, bornite and tennantite, there is actually no need for you to deal with it. Of course, it may also be used on a variety of other minerals, but it does not provide better results for these than other chemicals do. 

Safety Precautions

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I have to add some notes on handling and disposal. Commercial aminopolycarboxylic acid salts come as very fine powders and they are usually of technical grade. That is, their purity is only around 90%. The rest are synthesis by-products, mostly sodium acetate, sodium formiate and ammonium chloride. The latter severely harm your lungs upon inhalation. Since it is almost impossible to handle fine powders without generating more or less of an aerosol, this should be done in a well ventilated place. Consumed EDTA solutions will be enriched with heavy metals and, consequently, are very toxic. The complexants themselves are toxic to the aquatic environment. Not all of them are readily biodegradable: they may become enriched in sediments, and they also may mobilise heavy metals in surface waters. Therefore, they HAVE to be declared as chemical waste and disposed of appropriately. In this point of view, they are by no means comparable to citric acid, oxalic acid, etc., which have a very short mean lifetime in the environment and are fully degraded into entirely harmless products. 

Cleaning Process

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Before cleaning, the specimen should be placed in tap water for at least two hours, for the matrix becomes soaked with liquid. As a general advice, this should be taken care of in every cleaning process. 

For most applications (i.e. native copper, chalcocite, bornite, fahlore, enargite), a 2% (20 g/l) solution of EDTA tetrasodium salt can be used. Cleaning takes a couple of minutes. Only specimens with thick, spongy alteration crusts may need to stay in the solution somewhat longer. Though, anything that not dissolves within a short period of time, will never dissolve, even after hours or days. Carbonates, including malachite, are not affected. Copper halides and many copper sulphates, however, will completely dissolve. In this case, EDTA cannot be applied. 

The solution also works well for some tarnishes on galena, and for Alston Moor specimens of sphalerite on ankerite. Fahlores, especially tennantite, will come out very bright, but the tarnish often re-appears within a short period of time. This depends on their chemical composition, which can be largely variable. Chalcocites may come out bright silvery, bright black or bright blue-metallic, depending on whether they are true chalcocites or pseudomorphs. This is not always obvious underneath a black, sooty layer. 

If there is cuprite or chalcopyrite present, add acetic acid to the solution until it is slightly acidic (pH around 6, check with a test paper). In these conditions, chalcopyrite is stable, and it will also be cleaned to some extent. However, carbonates won't survive. Hence, if there are both carbonates and chalcopyrite present, EDTA cannot be applied. Cuprite does withstand a couple of minutes in an alkaline solution, but it may become dull after a longer period of time. 

Examples

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This specimen was found in debris, and it had obviously been exposed to the atmosphere for quite some time, since it was completely black when I noticed it:

The black layer did come off easily, within several minutes. The specimen is clean, but not overly improved: the copper remains dull, as it should, and the cuprite was left unaffected. 

This one was completely black also:

There was a black spongy mass that covered the whole matrix and also filled the cavity with the crystals. Cleaning time was about two minutes. 

This one was cleaned, but without much improvement:

There is strong superficial alteration evident, and in some places, a thin layer of bornite has formed. Stable alteration products are not affected by the treatment. The specimen came out a tad lighter than before, because there was thin, dark tarnish that readily dissolved. 

This one also had a dark tarnish, but did not look much worse before:

The dolomite was not affected by the treatment, nor were the iron stains. I should add that this specimen was of superior quality even before the treatment; it only had this typical dark layer that is common with many specimens that have been kept in cabinets over decades. The lustre of galena specimens is impossible to restore when there are only faint signs of superficial alteration.

 

No. It would not work. Chelate formation is almost entirely entropy driven. There are no naked ions in an aqueous solutions (except at very high salt concentrations). Metal ions bind several water molecules to form a complex; this may be simply an electrostatic cluster (such as with alkali and alkaline earth ions) or a compound with some degree of covalent bonding (such as in the case of most of the transition metal ions). These complexes are different species - they have different physical and chemical properties than the naked ions. For a metal ion with a coordination number of 6 (coordination numbers here are not the same as those encountered in solids), the reaction is as follows: 

[M(OH2)6]n+ + EDTA -> [M(EDTA)]n-4 + 6H2O 

We start with two molecules and produce seven. This increase in particle number of the system corresponds to a tremendous entropy increase. When the coordination number is only 4 (e.g. Li+), the reaction is less favourable; when it is only 2, as is the case with Ag+, the complex will be much less stable. Solubilities are also not compatible: EDTA is almost insoluble in acidic solution, whereas in alkaline solution, Ag2O will precipitate. EDTA is generally not applicable, when the respective metals form low soluble oxides or hydroxides. 

Lead minerals do not improve; this is because EDTA is unable to dissolve the lead secondaries.

Addition Information Links

 

Additional information can be found in the original post here http://www.mindat.org/mesg-19-65376.html
General information on the compound EDTA http://en.wikipedia.org/wiki/EDTA

 

Pyrite Rot

Based on correspondence with Dr Stephen Moreton, edited by Barry Flannery.

General Discussion

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The common iron sulphide minerals marcasite, pyrrhotite and pyrite are all rather unstable and tend to oxidise when exposed to the atmosphere. They turn into ferrous sulphate then ferric sulphate and sulphuric acid. The acid accelerates the reaction, so it becomes autocatalytic, with the result that once it gets started it speeds up and there's no stopping it. As the reaction products have a larger volume than the starting material the whole specimen eventually disintegrates and ends up as a handful of crumbs. Meanwhile the acid generated eats through everything around it.

There has been much speculation about microbes. Certainly Thiobacilus ferrooxidans eats pyrite, and turns it into ferrous sulphate. The consensus now seems to be that microbes are unlikely to be involved in the decay of specimens in cabinets. Instead it is a function of surface area and humidity. The larger the surface area, the more reactive the sample. Thus, the framboidal/granular pyrites that are so typical of, for example, the big Irish zinc mines - Lisheen, Galmoy, Mogul and Tara are very unstable. They are composed of many fine grains, and have pores or fissures penetrating deep inside them, so they have a large reactive area available. They also tend to contain some marcasite, which only worsens things. Large well-formed pyrite crystals, like the Spanish ones, have a very low surface area to volume ratio, and so last indefinitely.

So reactive is finely divided pyrite it has even been known to catch fire, as happened once in Tara mine, when a pile of such produced by drilling was left over the Christmas vacation, and ignited. To extinguish it that section of the mine was sealed to starve it of air. When everything had cooled down and it was re-entered the walls were coated with sulphur crystals distilled out of the pyrite.

Keeping the pyrite/marcasite dry is key to its survival. Water is essential for the reaction. As the reaction proceeds sulphuric acid is generated as a by-product. This is hygroscopic (absorbs moisture from the atmosphere) thus drawing in more water and causing more reaction. The reaction is therefore self-perpetuating and keeps on going until the pyrite is all consumed.

Treating The Rot

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As a research chemist for what used to be Europe's biggest desiccant gel manufacturer I am, perhaps, in a good position to figure out how best to keep my pyrites dry! Storing in a desiccator with a suitable desiccant, is my preferred method, but laboratory desiccators are big and expensive. Recently I have begun experimenting with cheap, plastic, air-tight food containers, which seem to be OK, at least for the short term. Fill them to about a fifth or a quarter of their volume (the more the better) with dry silica gel (can be in a sachet for convenience) and place the specimen next to it. There are many other desiccants available too. For very low humidities molecular sieves (zeolites) available as little grey pellets, are particularly effective, although I've found silica gel to suffice (at least my pyrites stopped falling apart once I began using it). Clay based ones are cheap, but generally not as good as silica gel. These can all be regenerated by heating, although zeolites need a good cooking at about 400°C to get them really dry. Silica gel can be freshened in the kitchen oven at about 110-150°C for an hour. I recommend an indicating gel that changes colour when used up. That way you will know if your container is not sufficiently air-tight. A few indicating granules mixed in with standard white gel will suffice. The traditional one uses cobalt chloride and turns from blue to pink. Turns blue again when baked in an oven to refresh it and can be used many times over. As cobalt salts are slightly carcinogenic I invented one based on iron, which goes from amber to near colourless (United States Patent Application 20040209372). A variant I came up with later uses a Fe-Bromo complex and gives a better colour change at even lower humidity and is manufactured under licence by Engelhard (as "Sorbead Orange Chameleon"), although I am not sure if it is available to non-industrial consumers, but would be the best one for the job. Both can be refreshed in the oven many times over. Some people use a microwave oven, although temperature control is harder. There are other indicating silica gels, but those based on organic indicators (commonly blue/yellow colour change, although there are others) tend not to withstand many drying/redrying cycles. I have no qualms about using the traditional blue/pink cobalt-based one for my samples however, and it is probably still the most readily available. Just don't eat it as it is toxic!

Calcium oxide is an even more powerful desiccant (drying agent), however, unlike silica gel, it cannot be regenerated (except in a furnace). It is also very alkaline. Where I have found it useful, however, is in treating specimens in the early stages of pyrite rot. For this one needs a large air-tight container (a desiccator is ideal, although I guess a particularly large food box will do as long as it is air-tight). Fill a small beaker or pot (100ml) with calcium oxide ("quicklime", best as lumps) and place it in the container. Place the specimens nearby in the container. Add a few ml of strong ammonia solution to the calcium oxide and put the container lid on fast. The calcium oxide reacts with the water in the ammonia solution, setting the ammonia free as gas. Leave for several days. Being a small molecule ammonia can penetrate into the tiniest spaces in the specimens, neutralising any sulphuric acid within them. It will thereby bring to a halt any pyrite rot, but will not repair the damage already done. Do this treatment in a well ventilated space (NOT IN THE HOUSE) as ammonia stinks. I use my garden shed. When done take the lid off (outside) and leave for 20 minutes for the smell to go then get the specimens into an air-tight container with desiccant as discussed above. If the specimen is already badly affected you may see some brownish colouration (hydrated iron oxides) where the products of the pyrite rot have reacted with the ammonia. In that form they will do no more harm, but rot will resume if the specimen is ever exposed to atmospheric moisture again, so keep it in its container and only bring it out briefly for ceremonial occasions.

Once rot is well underway there is little you can do to stop it. I have tried soaking in water to remove the acid and ferrous sulphate, then in dilute sodium or potassium hydroxide solution to make it alkaline. This seems to slow it down, but still they fall apart eventually.

Example Of Unstable Pyrite

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All fine grained framboidal pyrite specimens should be periodically monitored for signs of rot as a high relative surface area makes them extremely prone to decomposition. An example of framboidal pyrite from an Irish zinc mine:
 

Conclusion

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There is no true and absolute solution for preventing pyrite rot but the problem is readily treatable and progression of the rot can be effectively halted but only under controlled conditions. One should be aware that many treatments out there regarding ''pyrite rot'' are akin to old wives tales and either don't work or are outdated. Only the most current and up to date information should be taken as valid. The most active areas of research into pyrite rot are in the field of Palaeontology where the extremely valuable type fossils are often replaced by pyrite or marcasite and their preservation is of great importance to the wider scientific community. When researching treatments it is always very wise to check the fossil world as it will most likely have been tried extensively here. Pyrite rot requires BOTH oxygen and water to occur and is catalysed by acidity. Neutralization of acidity and the removal of either oxygen or water will severely reduce the rate of rot and many techniques can be used to achieve this. 

- Neutralize acidity in the specimen (by using some form of alkaline solution or substance)
- Remove ferrous sulphates and other oxidation products from the specimen (generally by reducing them or other complexing agents)
- Eliminate water from the specimen and environment (use of desiccants, drying in volatile water loving solvents such as acetone)
- Remove oxygen from the specimen and environment (use of anoxic or oxygen absorbing substances)