Mineralogy-101: Chemical Tests - 1

MINERALOGY-101: HOME LAB CHEMICAL TESTS - Part 1

SAFETY FIRST

Doing chemical testing in a home lab or workshop can be very helpful in whittling down the list of possible identities for a “mystery mineral.” But it does entail certain risks. Chemicals are dangerous! We each only have one set of eyes, one set of lungs – one life. And I can’t teach you lab safety protocols and procedures. Nor can you learn them from a book – any book! They need to be learned in the lab from someone who knows them well. Therefore I do NOT recommend that someone try to learn chemical testing on their own. I STRONGLY recommend that anyone new to doing chemical testing find someone who is experienced to apprentice under. See if there is someone in your local mineral club who does it – someone who still has all of his or her fingers, skin, and eyes – and isn’t on oxygen… Or stop by your local high school or a nearby state college and see if you can work something out with a chemistry instructor – develop a relationship. DON’T, though, dive in and hope for the best. You’ll get hurt. You could die…

GETTING STARTED

You will need to put together a lab area and stock it. A sturdy, stable, work table (a formica top is good, plain light color if possible, white!), a chair or stool the right height to save your back. At least a basic microscope that will give up to perhaps 50X. A cabinet for storing chemicals. LOCKABLE! Wood or plastic is probably best – metal ones tend to deteriorate if acids are stored in them. Test tube stands (a block of wood with holes drilled in it works just fine – so long as the holes are the right sizes), and the test tubes to go in them – lots and lots of small test tubes, semimicros if you are going to take the “minimalist” approach recommended below. Petri dishes (flat), glass rods, glass-stoppered bottles, beakers (50 ml are handy), glass slides (with depressions.) A wide flat glass container with low sides to work in under the ‘scope (save wear and tear on the ‘scope’s stage…) No bigger than the stage area, but not much smaller. An alcohol lamp (NEVER refill it while the wick is still hot!!!) or hotplate. (And a blowtorch for various flame tests that require more aggressive heat – but Al O. will cover that in his post on flame tests.) A sturdy mortar and pestle for grinding mineral chips to powder; preferably ceramic or granite – not marble, which is too soft. (You keep getting ground up calcite in with your powdered “mystery mineral”…) Keep it clean. Handling tools – lab spoons and small spatulas, tweezers, test tube tongs, etc. Needle nose pliers. (And probably a dozen other things I’m forgetting…)

GOGGLES, rubber apron, and rubber gloves (thin surgical “throw aways” are probably best – so long as you throw them away. Don’t reuse!)

Chemicals. I’m not going to try to give you a list. As I discuss the various tests, you can build your own list. They are largely acids, bases, and other reagents and fluxes. Some of them are restricted – difficult to obtain. (A good relationship with a high school or college chemistry instructor helps there.) They really need to be lab grade. Industrial grade hydrochloric acid (aka mureatic acid) might be fine for cleaning specimens. But it is not a testing lab chemical. Ditto the sulfuric acid that is used in car batteries. And other industrial chemicals that aren’t lab grade. You substitute such inferior chemicals at the risk of screwing up tests. If you stick to lab grade chemicals, you’ll be able to trust the results with a much higher degree of confidence. Honest!

Most of what you will need is available from chemical supply houses. If you do not live in an urban area where one of these is handy, some pharmacies will order things for you – though not restricted chemicals. Again, a local chemistry instructor is your best bet for those. Get to know one, gain their respect - and their understanding of what you are doing. They can help you in a lot of ways – safety, chemicals and labware, protocols and procedures, and tests that might not get mentioned here. (I can’t think of a good reason NOT to strike up a relationship with a chemistry instructor if you want to get into the chemical testing of minerals… Even if the nearest school is an hour away.)

A BASIC APPROACH TO TESTING

Think small! It doesn’t take a beaker full of HCl to dissolve a tiny pinch of powdered mineral. And you don’t need a pound or two of powdered mineral to do the testing. Work in small batches – pinches of powders and drops of liquids. Work in the flat glass dish under your microscope, placing the powders and liquids on a glass microscope slide (those with depressions ground in them work best – keeping the drops where you put them) for ease of cleanup and positioning, and use the ‘scope to observe reactions. While not all tests require this, most can be done this way. Even when you don’t work under the ‘scope, thinking small is still a good idea – minimal amounts in the bottom of a test tube. It’s cheaper, it should be safer, and it should be easier to mix the solutions you’ll need – get the strengths you’ll need. Besides, with a lot of mineral testing you probably won’t have much pure “mystery mineral” to work with. Maybe just a tiny crystal or two tweezed off its matrix. If you start out learning to work in small amounts, you won’t get caught short when you run into this situation – which will be quite often.

Be methodical. And take notes. Working helter-skelter gets helter-skelter results. And often a miss instead of a hit. And if you don’t take notes you may not remember exactly what you did to get the results you got – good or bad. It’s nice to be able to repeat the good tests – avoid the bad ones. Notes help guide you in that direction. Keep them to refer back to the next time you test something. It might be a year between doing certain tests for certain things – and those notes will be worth their weight in gold as reminders.

There are basically two broad categories into which tests fall: Qualitative and quantitative. Qualitative tests tell you whether or not a certain element, radical, or compound is present in your mystery mineral. Quantitative tests tell you how much of that constituent is present – weight percents or mole percents of various elements or oxides of elements, etc. Quantitative tests are rather tricky – or at least reliable ones are. They are “advanced mineral chemistry.” We’ll stick to qualitative tests for this basic introduction to home lab work, with maybe some hints on rough quantitative measurements now and then. You’re interested in seeing what the mystery mineral contains, not how much of it or how pure the sample is. At least that should be the focus when you start out. Any quantitative hints that get discussed will serve to point out that there is that next level of chemistry available – when you feel up to moving on.

In qualitative testing the goal is to find out whether or not certain constituents are in the mystery mineral. If you suspect a calcium-bearing mineral, you test to see if there really is calcium present. And so on. Most of the tests we’ll discuss here are of that type. Is it – or isn’t it – present? Negative results are as important as positive ones. They tell you what isn’t present. So you can move on to find out what is.

METHODOLOGY & THOROUGHNESS

You don’t need a home lab to identify minerals that can be identified from gross physical features – such as quartz, kyanite, and muscovite (usually…) Save the lab work for stuff that is defying identification by simpler means. Except it’s not a bad idea to experiment on “knowns” – see if you can “prove” a pyrite sample is pyrite by chemical testing. Same with other species you are familiar with. The practice will familiarize you with testing procedures, give you experience working with the chemicals, and give you some confidence once you get things worked out and are “proving” the samples are what you know they are. But, by-and-large, save the lab work for the tough ones.

As stated in the previous post, lab tests are all about finding out what’s in your mystery mineral(s). You narrow down the possibilities to two or a few species, and then you say “Okay – now what are the chemical differences between these species that I can test for to figure out which I have?” Let’s take a simple example: You have a sample that you’ve narrowed down to being one of the following species: calcite, magnesite, siderite, or smithsonite. Normally, figuring out the specific gravity of the mineral would tell you which you have – but your sample is small crystals, nothing big enough for you to work with on the triple beam balance (or whatever you use to do S.G. tests.) So you look these species up and find that calcite is CaCO3, magnesite is MgCO3, siderite is FeCO3, and smithsonite is ZnCO3. The only difference between these species is their metallic ion: Ca, Mg, Fe and Zn. So the way to find out which species you have is to test for those ions – find out which is present. You therefore need a chemical test that will identify Ca, another that will identify Mg, and so on…

A simple chemical test for Ca is to dissolve a pinch of the powdered sample in a drop or two of hydrochloric acid (HCl) and then add sodium sulfate (Na2SO4) to the solution. If you get a white precipitate (it’s CaSO4) there is Ca in the sample. It’s probably calcite. BUT – a single test is not conclusive. Any of these minerals may have Ca in them in small amounts as a “substitution” for the principal metallic ion. (We get the “ideal formulas” fed to us in books – not “actual formulas.”) So you would still need to test further. You can test for Mg by dissolving a pinch of the powdered mineral in HCl and adding several drops of a strong solution of ammonia to it. Then add a bit of Na2HPO4.12H2O to the solution. If you get a white precipitate (it’s NH4MgPO4) then there is Mg in the sample. To test for Fe you could use any of several test. One simple one is to dissolve a pinch of the powdered sample in a drop or two of HCl and then add several drops of ammonia hydroxide (NH4OH) solution to it to make it alkaline. If you get a tannish to white precipitate (it’s Fe(OH)3 ) then there is Fe in the sample. Then, for Zn, dissolve a pinch of powdered sample in HCl and add sodium disulfide (Na2S) solution – if you get a white precipitate (it’s ZnS) there’s Zn in the sample.

As you do each test, observe the volume of precipitate formed – keep adding the reagent to the solution (in tiny amounts) until no more reaction is taking place, no more precipitate forming. (If you keep the slides (or test tubes) of each solution lined up on the table, when all the tests are done you can compare the solutions with one another directly if necessary.) Note any negative results – such as no Mg found. Those eliminate species from your list of “likely suspects.” Let’s say you got negatives for Ca and Mg, positives for Fe and Zn. But the Zn test barely showed a trace, while the Fe test showed the sample is loaded with Fe. You’re “mystery mineral” is no longer a mystery – it’s siderite with a little bit of Zn substituting for the Fe. Actual formula:

(Fe,Zn)CO3. “Ideal” formula: FeCO3. [And here we’ve looked at rough quantitative testing: “How much?” As opposed to “Is it there?”]

“Aha!” Someone says. “But how do you know it is an iron *carbonate*? – not a sulfate or something else?” Well – I said you’d already decided it had to be on of the carbonates named. And you probably did that be testing a sample or two with HCl to see if it “fizzed.” And it did… Which is another chemical test – a simple test for CO2 (which is what comes bubbling out of a solution of carbonate mineral and acid.) See – you already knew one “chemical test” and didn’t even realize it! If you didn’t do that test, and decided it must be one of those carbonates by other means, then that test needs to be done to prove your guess is correct. Because without knowing for sure that the mineral is a carbonate, it may well be something else, like a sulfate. (Actually, if you think about it, you “proved” the mineral to be a carbonate as you dissolved the powder in acid – because it bubbled as it dissolved… PS: The sulfates of Fe, Mg and Zn are readily soluble in acids and likely to stay dissolved – not precipitate out of solution. Another reason you’d know the mineral was a carbonate – not a sulfate - in the above scenario.)

You have to test for EACH of the suspected constituents in a mineral – not just one. Which can be tricky if you only have a little bit of powdered mineral to work with and need to do a bunch of tests. Here the minimalist approach I recommend becomes critical – else you may not be able to do all the tests needed. It is always a good idea to figure out how many tests you are going to need to do and then divide up the available mineral powder solution into that many “batches.” In some cases this may mean that all you have to work with are a drop or two of solution for each test. A micro-eyedropper is handy for working with tiny amounts of solutions – and you’ll probably need to work under the microscope in order to see the reactions in the tiny drops of solutions you have. But it can be done. You can also divide up the mineral powder into tiny “piles” ahead of time and then make the solutions you need from them. This is handy if you need different powder solutions for different tests. But it can be tough to do with tiny quantities of powder – possibly ending up with just a few grains of powder to make solutions from.

Of course not all tests are as simple as the series described above. Or all results as certain as our hypothetical case. I picked an “easy” one to illustrate the methodology. And to point out that you can’t just test for one thing and – if its positive – say “Voila!” You can’t. You have to take other possibilities besides your “prime suspect” into consideration. In the above case, if you’d just tested for Zn you would have falsely concluded the mineral was smithsonite… Chemical testing is about “proving your case beyond a reasonable doubt.” Any single test is more than likely to leave plenty of room for doubt. Conclusive results come from being thorough – testing for

all possible culprits and finding out which you have, and which is the most likely “prime candidate” based on comparisons.

LAB PROCEDURES & TIPS

MAKING SOLUTIONS

Nearly all of the tests require that you first dissolve some of the powdered mystery mineral to make a solution that is then used for testing. This can be a problem in itself: What do you use to dissolve the powder? (And here we are talking about *dissolving* the powder – not just mixing it in liquid. In a mixture, the powder is not dissolved, just stirred in and floating around. In a solution the powder has been dissolved – broken up into its constituent ions.)

Only a few minerals are soluble in water – mostly halides and iodides. (Which can be a quick test to see if you have one of those species.) But the odds are that your powder won’t dissolve in water – you’ll have to try something else. Try hydrochloric acid (HCl) first. A great many minerals will dissolve in HCl. If that doesn’t work, try nitric acid (HNO3). If that doesn’t work, try sulfuric acid (H2SO4) last. If that doesn’t work the odds are your mystery mineral is a silicate, few of which will dissolve in acids. Zeolites do, and a few others, but by-and-large silicates are insoluble in anything short of hydrofluoric acid (HF) – and that’s one acid you can’t play with at home. (It’s extremely dangerous and requires special lab equipment – including a negative pressure ventilation hood – to work with safely.) If you run through the list of acids you can use at home and still can’t get the powder to dissolve, you’ve pretty much determined the mineral is a silicate – and probably not a Zeolite. Important information in itself.

A few other acids are also used in making solutions, such as acetic and oxalic. But these are rather weak acids, comparatively speaking (they can still be “strong,” for them, and are dangerous.) They are usually not used unless a specific test calls for them. And if a test does call for one and you can’t get the powdered mineral to dissolve in it, the odds are the mineral is not what you think it is and are testing for: You can scratch that possibility and move on to other tests.

ACIDITY vs ALKALINITY

Quite a few of the tests require you to start with an acid solution and then make it alkaline by adding a base, such as ammonium hydroxide (NH4OH). In some cases you simply overload the solution with the base to get it there. In others you need to work carefully to make the solution just slightly alkaline.

You use litmus paper for checking the pH of solutions. But since you are working with such small quantities of liquids, you don’t want to just stick the tip of the test strip in the solution – it might wick up your entire batch of liquid! Instead, dip a small diameter (approx. 2.5 mm) glass rod in the solution and then touch the wetted end of it to a fresh spot on the litmus paper. The wetted spot on the paper will then tell you where things stand alkalinity wise.

Also, sometimes (though rarely) you need a neutral solution – pH of 7. Which can be tricky to achieve. You pretty much need to use a micro-eyedropper to slowly add the base and – drop by drop – neutralize the acid, testing your solution with a glass rod and litmus paper after each drop has been added and stopping the first time the litmus paper doesn’t turn color.

MAKING A MICRO-EYEDROPPER

You can buy eyedroppers with small diameter ends if you have a good chemical supply house in your area. But otherwise they can be hard to come by – and they aren’t really “micro” eyedroppers. So you might want to try and make your own. Take an 8 inch length of glass tubing about the same diameter as a typical medicine eyedropper. Heat the center of it over a propane torch flame until it’s soft. Then remove it from the flame and pull on both ends so the middle stretches out. Snap the middle – so there is a roughly 3 inch long thin capillary tube coming off the end of the unstretched part of the tube. Once the glass has cooled, then heat the wide end of the tube until it’s soft and flare it slightly with the tang of a file – so it looks roughly like the top end of a regular medicine dropper. After it cools again, fit the flared end with a medicine dropper bulb.

If done right, you get an eyedropper that has a very thin capillary tube on its end – just right for creating micro-drops of solutions. Actually, one length of tubing can give you two droppers if you do the stretching and snapping carefully – each half having a capillary tube on its end. If you make several of them you’ll have enough to do testing without having to clean one between use in different solutions. Just make sure they – and all your glassware, etc. – get cleaned thoroughly between sessions. There’s nothing like evaporated residue to botch up a test…

HEATING THINGS

First off, where a test calls for heating liquids on a glass microscope slide you need to be very careful. If you just stick the slide over the flame the odds are it will shatter. And you can’t just try to hold it immediately above the flame – because that happens to be a pretty hot spot in the gas stream. You need to hold the slide “high” – two or more inches above the flame. There’s enough heat at that point to slowly warm the slide, but not so much that it will heat the slide quickly and shatter it. I suggest you practice (with a broken half of a slide if you have a couple that you’ve dropped) and use tongs or pliers to hold the slide in case it either gets too hot or breaks. But, when actually heating a slide during testing, I like to hold it in my fingers. If it starts getting too hot I know it immediately! :~} Basically such tests require that you *warm* the slide – not get it red-hot – in order to either promote evaporation of the liquid or hasten a chemical reaction. And you do so *gently* – not *strongly.*

Heating solutions in a test tube is less of a hassle – they are designed to be heated, even *strongly.* So you shouldn’t have to worry about them shattering when you put the heat to them. But you do need to worry about following the instructions for a test – not heating something strongly when the test calls for gentle heating. Again, you are usually trying to either evaporate the liquid or hasten a reaction in it.

Gentle heating calls for pretty much the same approach as handling slides. Again, I hold the tube in my hand, tipped at an angle, and don’t let it get too hot to hold – keeping my fingers below the rim so they aren’t in the path of any fumes issuing from the mouth of the tube.

Heating strongly requires a different approach. Here you want to put the heat to the tube – but without actually bringing the contents to a boil. If you experiment a little you’ll find that there is a curious phenomena which occurs when you heat a test tube: Just before the liquid comes to a boil the tube begins to vibrate. So what you do is heat the tube until you feel the vibration, then remove it from the heat for a few seconds. Then put it back over the heat until it vibrates again. And so on until you’ve achieved your purpose. After a while you’ll probably find yourself sort of waving the test tube over the flame, keeping it just shy of boiling by dipping it in and out of the flame with a continuous motion. Perfect!

DO NOT BRING ACIDS OR ACID SOLUTIONS TO A BOIL! You may only be working with a few drops of solution, the danger relatively low, but – believe me – you do not want to risk boiling off an acid all at once and getting a whiff of the fumes. Or getting them in your eyes! (While there is still some risk in slowly evaporating acid solutions, the fumes produced are much more dispersed and “diluted” with air.)

A few tests require heating the tube containing the test solution in a boiling water bath. This means heating a beaker of water to a boil and suspending the test tube in it for whatever time is required. A neat trick is to bend a medium or large sized paper clip so one part forms a test tube holder and the other part forms a “hanger” that can be hung on the edge of a beaker. Then you don’t need to hold the tube in the water by hand in a test tube holder.

SILICATE FUSIONS

As noted above, most silicate minerals will not dissolve in the acids you can use safely in a home lab. Some amateur mineral chemists avoid trying to test silicates for this reason. But there is a method for freeing up the metallic ions in silicates that can be done in the home lab. It is a bit tricky – and it won’t work for all silicates - but with a little practice most people should be able to become proficient at it and thus expand the list of minerals they can test. All it takes is sodium carbonate (Na2CO3), charcoal blocks, and a propane torch or blowtorch. It’s a bit messy, but you can use charcoal briquettes in place of charcoal blocks – you just have to scrape a flat on one side to provide a stable base, and then gouge a small hollow in the opposite side. Make sure the hollow is large enough to hold all the powder mixture you want to fuse. (It’s better to fuse several small batches than to try for one large one.)

Mix a pinch of powdered silicate mineral with about four or five times that amount of sodium carbonate powder. Place the powder mixture in the hollow on the charcoal block (or briquette). Then heat the mixture with the blue flame of the torch until it fuses into a solid globule on the charcoal. Once the globule cools, crush it back into a powder. This powder should be dissolvable in hydrochloric acid – and the resulting solution useable in a variety of tests for metallic ions. (If you find the powder does not dissolve, try another fusion – perhaps using a smaller amount. Sometimes it takes a couple tries to get a good fusion. If after several tries you can’t get a powder that dissolves, you’ve probably come upon one of the silicates that’s difficult to fuse or won’t fuse at all. That can be useful info in itself.)

QUANTITIES & MEASUREMENTS

I tend to think – and talk – in terms of “pinches” of powders and “drops” of liquids. But what’s a “pinch”? A “drop”? I have a tiny lab spoon that holds about ¼ of the volume of a Bayer aspirin. Maybe something like 5 to 7 milligrams. I call that a “pinch.” A “tiny pinch” is what I can pick up with a pair of small tweezers and tap off into the slide depression or test tube – maybe half a “pinch” or less. A “drop” is what drips out of the end of a lab eyedropper when you squeeze the bulb just enough to form a bead of liquid at the end that is big enough to drop off under its own weight. Don’t ask me what that is in milliliters… (The micro-eyedropper is used for

producing tiny drops when you only have tiny amounts of mystery mineral powder to work with, or when you need to slowly add solutions to achieve a desired degree of acidity or alkalinity.)

Just thought we ought to define these – ahem! – highly technical and precise terms before moving on. :~}

Also, this points out the fact that most qualitative testing for minerals is not really all that persnickety when it comes to measurements. You don’t usually need precisely controlled amounts of powders and liquids to get useable results. (Though some tests do require measured amounts.) In quantitative tests measurements become more of an issue. But we’re only talking rough quantities here – when we talk about them at all. So you don’t have to get all in a lather about making sure you use exactly the right amount of this to make that work.

SAFETY NOTE

Please read the previous posts on this topic before reading the information on tests below. There is important information in those posts that you need to know before you try doing any tests. If you try to short-cut things and dive into testing you could get hurt – or hurt someone else.

Also please note that each post is being posted in reverse order: The lowest one was posted first, and so on upwards to the last one. Read them in order; and if you print them out for future reference number them in the proper sequence.

………………………………….THINK SAFE – BE SAFE………………………………..

CHEMICAL TESTS, Part-1

[Note: In some cases two or more tests are offered for a given element, etc. This should help you in a couple ways: 1) You can try each test to confirm results, and 2) you might have chemicals for one of them, but not the others.]

Aluminum (Al): Dissolve a tiny pinch of powdered mineral in a few drops of acid. Add several drops of sodium hydroxide solution (NaOH) to the solution to make it alkaline. (Use litmus paper to determine alkalinity.) If Al is present a white precipitate of Al(OH)3 gel will form, and then re-dissolve. [WARNING: NaOH is extremely caustic. Use with all due precautions.]

Aluminum (Al): If the powdered mineral is pale in color or white, strongly heat a pinch of it in a test tube; then cool and add a few drops of cobalt nitrate solution (Co(NO3)2.6H2O) and reheat. If the powder turns blue there is Al present. (Same test shows Zn or Mg in some minerals.)

Antimony (Sb): Dissolve a pinch of powdered mineral in a few drops of hydrochloric acid (HCl). Slowly add micro-drops of ammonium hydroxide (NH4OH) until the solution is just slightly alkaline. Then add micro-drops of HCl until the solution becomes just slightly acidic again. (Use litmus paper to gauge alkalinity vs. acidity.) Then add one or two drops of freshly prepared sodium sulfide solution (Na2S). If an orange to reddish brown precipitate forms, there’s Sb in the mineral. (Always prepare fresh Na2S solution – it decomposes over time, old stuff isn’t any good.) [WARNING: NH4OH is extremely caustic. HCl is a strong acid. Use both with all due precautions.]

Antimony (Sb): Another test is to place a drop or two of the test solution prepared in the above test onto a small shaving of tin (pure tin – not an alloy) and see if Sb forms on the surface as a black powder.

Antimony (Sb): Yet another test is to dissolve a pinch of powdered mineral in a few drops of HCl, lower a shiny steel brad into the solution, heat gently, and see if the black powder of Sb forms on the shiny surface of the brad. [WARNING: HCl is a strong acid. Use with all due precautions.]

Aragonite vs. Calcite: Add a pinch of powdered mineral to a few drops of cobalt nitrate (Co(NO3)2), boil until liquid is nearly gone. Cool and rinse thoroughly in water. If the powder is stained pink it’s aragonite. If it isn’t, it’s calcite. (This assumes you have narrowed the choices down to one or the other.)

Arsenate (AsO4): Make a fresh solution of tin chloride (SnCl2) by adding shavings of tin to a few drops of HCl and heating gently until the tin is dissolved; then add several drops of a solution of fresh ammonium molybdate and powdered mineral. If the solution turns deep blue, there’s AsO4 present. (This test also works for phosphate – PO4.) To confirm AsO4 – as opposed to PO4 – first make a solution of distilled water and sodium carbonate (Na2CO3), slowly add the carbonate solution to the test solution until there is no effervescence. Check to make sure the solution is neutral to litmus paper. Then add one or two drops of silver nitrate solution. If AsO4 is present a reddish-brown silver arsenate precipitate will form. (Silver phosphate is yellow, as is silver arsenite.) (Note: Remember to always mix fresh ammonium molybdate solution and fresh tin chloride solution. Old stuff decomposes and isn’t any good.)

Barium (Ba): Dissolve a pinch of the powdered mineral in a few drops of HCl and add a drop or two of sodium sulfate (Na2SO4) to the mineral solution. A white precipitate of barium sulfate (BaSO4) will form if there is Ba in the mineral powder. (Same test will show calcium (Ca) as a white calcium sulfate precipitate (CaSO4) and strontium (Sr) as a white strontium sulfate precipitate (SrSO4).)

Barium (Ba): Dissolve a pinch of powdered mineral in a few drops of acetic acid (CH3CO2H), add several drops of potassium chromate (K2CrO4). If Ba is present a precipitate of yellow barium chromate will form. (Use approx. double the number of drops of potassium chromate solution as the number of drops of acid solution.) (This test will not work for all Ba minerals. If the powder will not dissolve in acetic acid, you need to use a different test.)

Bismuth (Bi): Make a slightly acid test solution of powdered mineral. Add several drops of NH4OH to a few drops of the test solution. If Bi is present it will form a white precipitate of bismuth hydroxide (Bi(OH)3). (This same test can be used for cadmium and tin – see below and next post.)

Bismuth (Bi): Use the same test solution described in the above test. Add freshly prepared Na2S solution. If Bi is present a dark brown precipitate of bismuth sulfide will form.

Boron (B)/Borate (BO3): Dissolve a pinch of powdered mineral in two or three drops of sulfuric acid, then add an equal amount of alcohol. The solution will burn with a bright green flame if B is present. (Really a “flame test” – sorry for snitching it, Al. :~} )

Cadmium (Cd): Dissolve a pinch of the powdered mineral in a few drops of acid (HCl, nitric, etc.) and add a few drops of sodium sulfide solution (Na2S). If Cd is present it will produce a bright yellow precipitate of CdS (which happens to be the mineral greenockite! :~} )

Cadmium (Cd): Use the same test solution described in the above test. Slowly add drops of NH4OH to make the solution slightly alkaline. If Cd is present it will form a white precipitate of cadmium hydroxide. The precipitate will re-dissolve if you add more ammonium hydroxide. (This same test can be used for lead and bismuth, but they will not re-dissolve in an excess of NH4OH.)

Calcium (Ca): Same test as first test for Ba, above, will produce a white precipitate of calcium sulfate (CaSO4.)

Calcium (Ca): Another test is to dissolve a tiny pinch of the powdered mineral in a drop or two of acid, then make the solution alkaline with a few drops of ammonium hydroxide solution; then add ammonium oxalate solution ((NH4)2C2O4.H2O.) If Ca is present in the mineral solution it will produce a white calcium oxalate precipitate (CaC2O4.) (Same test will show Ba as white barium oxalate precipitate, and Sr as white strontium oxalate precipitate.)

Carbonate (CO3) & Carbon dioxide (CO2): Most carbonate minerals bubble (“effervesce”) when acid is added to them – even un-powdered – due to the reaction giving off carbon dioxide. Calcite even effervesces in vinegar (a weak acid) a little bit. Powdered, it will effervesce in Coke, which has weak carbonic acid in it. Carbonates which do not effervesce in cold acids will do so if the acid is heated. [CAUTION: Never bring acids to a boil! The vapors can be deadly!] Some silicate minerals which contain carbonate will partially dissolve in heated acids, producing a gel or white powdery residue.

Chlorides/Chlorine (Cl): Most chloride minerals – such as halite (“rock salt”) - will dissolve in water. Adding silver nitrate solution (AgNO3) will precipitate white silver chloride (AgCl.) The dried precipitate will dissolve in ammonia solutions. Any mineral powder which can be dissolved in nitric acid (HNO3) can be tested the same way. [WARNING: NHO3 is a strong acid. Use all due precautions.]

Chromium (Cr): Dissolve a pinch of powdered mineral in a few drops of acid (HCl or HNO3) and add several drops of NH4OH to make the solution alkaline. If Cr is present it will form a green precipitate of chromium hydroxide, which will not re-dissolve if excess NH4OH is added. To confirm, add a few drops of hydrogen peroxide and heat in a boiling water bath for several minutes to oxidize the Cr ion to CrO4, then add several drops of silver nitrate solution; CrO4 will precipitate with Ag to form brick red silver chromate.

Cobalt (Co): Dissolve a pinch of powdered mineral in a few drops of HCl, add several drops of NH4OH (until solution is slightly alkaline). If Co is present a blue precipitate will form – which is soluble in excess NH4OH, producing a brownish yellow solution. On exposure to air the solution turns red. (Note: the bead test for Co is much better. Al O. will cover that in his post on flame tests.)

Copper (Cu): Most copper minerals (other than silicates – and even some of them) will color HCl green. Powder the mineral and add a pinch to a few drops of acid in a small test tube. Many copper minerals won’t bubble in the acid, but the acid will still turn green from the release of Cu ions.

Copper (Cu): Another test is to mix a tiny pinch of the powdered mineral in a few drops of HCl, then add several drops of NH4OH to make the solution strongly alkaline. If Cu is present it will precipitate as yellowish- to greenish-white Cu(OH)2 and then re-dissolve and color the solution bright royal blue.

Fluorine (F): Place a tiny pinch of the powdered mineral in the depression on a *new* (unscratched) glass slide. Add sulfuric acid (H2SO4) by the drop until the powder is just fully dissolved. Heat gently by holding the slide over a flame for several minutes. Cool the slide, then *thoroughly* wash it off and dry it. Check to see if the glass is etched – has fine little “scratches” or is “dissolved-looking.” Etching indicates that the mineral has F, which was converted to hydrofluoric acid (HF) which in turn etched the glass. [WARNING: Hydrofluoric acid is *extremely* dangerous! It’s fumes are deadly! Always do this test with the smallest possible amounts of the powdered mineral and sulfuric acid. Always do it in a well-ventilated area. Always use extreme caution not to spill the contents of the slide, or get it on you, your clothes, or the work area. Make sure the slide is thoroughly washed and dried before examining it closely for etching. (It is best to do the examination under the microscope – both to keep the slide away from you and to see the etching better if it’s there.)] Alternate method: Once the powder and acid are mixed, set the slide aside in a *safe* place for several days (at least three, maybe a week) instead of heating it. After several days proceed as if the slide had been heated. Still use the same extreme caution – if there was F in the powder, the HF is still there… [CAUTION: You have to thoroughly rinse and dry the slide before you put it under your microscope. If there is even a trace of HF left it can form fumes and etch the objective lens of your ‘scope. (That’s the lens that’s positioned right above the sample being examined…)]

Iron (Fe): There are several tests for Fe. One simple test is to dissolve the powdered mineral in either hydrochloric acid or nitric acid; Fe will turn the solution yellow. This will also work with lab grade oxalic acid (but may not with oxalic obtained from hardware stores, which is not lab grade and can cause yellow coloring due to impurities.)

Iron (Fe): Another test is to dissolve a tiny pinch of the powdered mineral in a drop or two of HCl or HNO3, then add several drops of NH4OH to make the solution alkaline, and see if reddish-brown iron hydroxide precipitates (Fe(OH)3.) (If you use HNO3 for this test and get positive results, the Fe is Fe+++ - see below for info on Fe of different valences.)

Iron (Fe): A third test (very sensitive) is to add a tiny pinch of the powdered mineral to a few drops of potassium thiocyanate solution (KCNS.) If Fe is present it will color the solution bright red. (This is used in dying agates red…)

[Iron has two “valance states” – Fe++ and Fe+++. There are tests which will determine which valence is present in a mineral. But be advised that some minerals have iron present in both valances.]

Iron (Fe++): Add a tiny pinch of the powdered mineral to a drop or two of potassium ferricyanide (K3Fe(CN)6) solution (water or acetone). If Fe++ is present it will yield a blue precipitate. (This is used for dying agates blue.) [WARNING: Any cyanide chemical is extremely poisonous! Use with all due precautions.]

Iron (Fe+++): Same test as immediately above, but using potassium ferrocyanide (K4Fe(CN)6.) If Fe+++ plus is present it will yield a blue precipitate. [See WARNING immediately above.]

[In minerals containing Fe of both valances, it may be possible to judge which is more prevalent by the volume of the precipitate. In each test, add drops of the solution slowly – one at a time with a wait in between - until the reaction is driven to completion – no more precipitate forms (all of the Fe present is used up). Compare the amount of precipitate in each test tube. The one with the greater amount should be the one which is more prevalent. (Another very rough quantitative test.)]

Lead (Pb): Dissolve a pinch of powdered mineral in a few drops of HNO3, add a drop or two of HCl. If Pb is present it will precipitate as white lead chloride (PbCl2). To confirm, draw off the solution, leaving only the moist precipitate in the test tube, then add 3 or 4 drops of distilled water to the precipitate, then heat gently while stirring with a glass rod. PbCl2 will dissolve in hot water. To confirm further, add 2 or 3 drops of potassium chromate – you should get a yellow precipitate of lead chromate (PbCrO4) if there is Pb in the solution. If you get the precipitate, you’ve confirmed the presence of lead. (Note: The test up to heating in distilled water will also show Ag and Hg, but neither will dissolve in the heated water.)

Magnesium (Mg): If the powdered mineral is pale in color or white, strongly heat a pinch of it in a test tube; then cool and add a few drops of cobalt nitrate solution (Co(NO3)2.6H2O) and reheat. If the powder turns blue there is Mg present. (Same test shows Al or Zn in some minerals.)

Magnesium (Mg): Another test is to dissolve a tiny pinch in a drop or two NH4OH to make a strongly alkaline solution, then add a few drops of sodium hypophosphate solution (Na4P2O6.10H20). If you get a white precipitate there’s Mg.

Manganese (Mn): Dissolve a pinch of the powdered mineral in a few drops of fresh sodium sulfide solution (Na2S.) If Mn is present a pink precipitate of manganese

sulfide (MnS) will form. When dried and exposed to air this precipitate will quickly turn brown and then black due to oxidation.

Mercury (Hg): Mix a pinch of powdered mineral with a pinch of sodium carbonate powder (NaCO3.) Heat strongly in the bottom of a test tube. If Hg is present it will form a silvery film on the wall of the test tube where the glass is still cool.

Nickel (Ni): Dissolve a tiny pinch of powdered mineral in a few drops of NH4OH to make a strongly alkaline solution, add a few drops of dimethylglyoxime and alcohol solution (let’s skip the formula for this one, it’s a bear…) The presence of Ni will cause a red precipitate. (This is a fairly sensitive test for Ni.)

Niobium (Nb): Mix a pinch of powdered mineral with a pinch of sodium carbonate powder on a slide and fuse over a blue flame. Crush the melt back into powder and dissolve in a few drops of HCl. Heat the solution to a boil and add tiny shavings of tin. If the solution turns first blue, then brown, there’s Nb present. [I know I said to never bring acids to a boil. Here’s a test that makes you break the rule. Remember – you’re doing it with a tiny amount of acid. So long as you don’t sniff the stuff – or spill it! – you should be okay. Just make sure you are in a well-ventilated place – like outdoors…]

Phosphate (PO4): Many phosphate minerals are rapidly etched by concentrated HCl – so a quick test for PO4 is to put a drop or two of concentrated acid on a chip of the mineral and set it aside for a while – maybe half an hour or so. Then examine the chip to see if it is etched. (Note: Don’t forget to examine the chip before you put the acid on it – so you know what it looked like and can compare that with how it looks after the acid has been applied.)

Phosphate (PO4): Another test is to dissolve a pinch of powdered mineral in a few drops of HNO3 and add a few drops of ammonium molybdate solution ((NH4)2MoO4.) If it yields a yellow precipitate there’s PO4 present – the precipitate is ammonium phosphomolybdate.

Phosphate (PO4): Another test is to make a fresh solution of tin chloride (SnCl2) by adding shavings of tin to HCl and heating, then add a solution of fresh ammonium molybdate and powdered mineral. If the solution turns deep blue, there’s PO4 present. (This test also works for arsenate – AsO4, see previous post. Since arsenates aren’t as readily etched by HCl, doing both the etch test and this test will usually distinguish between the two.) [Note: Remember to always mix fresh ammonium molybdate solution and fresh tin chloride solution. Old stuff decomposes and isn’t any good.]

Silver (Ag): Dissolve a pinch of powdered mineral in a drop or two of HNO3, add a drop of HCl. If Ag is present it will yield a white precipitate of silver chloride (AgCl) which turns dark on exposure to light. The precipitate will not dissolve in hot water (see test for Cr in the previous post) but will dissolve in NH4OH. (Same test shows Hg, but it won’t dissolve in NH4OH. To confirm, add 2 or 3 drops of potassium chromate – if you get a brick red precipitate it’s silver chromate (Ag2CrO4) and Ag is confirmed.

Sulfate (SO4): For water-soluble minerals, dissolve a pinch of powdered mineral in a few drops of distilled water, then add a few drops of barium chloride solution – BaCl2.2H2O. If SO4 is present it will yield a white precipitate of BaSO4.

Sulfate (SO4): For non water-soluble minerals, first dissolve the powdered mineral in HCl, then evaporate most of the acid by heating until the solution is almost gone. Then add a few drops of distilled water. This gets you to the same place as the start of the test described above, with the SO4 disassociated from the metallic ion(s) – proceed with that test.

Sulfur (S)/Hydrogen Sulfide (H2S): Suspected sulfide minerals can be easily tested by crushing to a powder and then adding a drop of HCl to a pinch of the powder. Sniff (gingerly!) – a “rotten egg” odor indicates the presence of hydrogen sulfide gas being released by the reaction. Sometimes this will even work with a crystal or chip that hasn’t been powdered. And sometimes the process of grinding the mineral to powder produces the odor – such as with pyrite. So, you may not even have to add the acid!

Tantalum (Ta): Same test as for niobium – see above. (Unfortunately, I don’t know of any “home lab chemistry” test that will distinguish between the two…)

Tin (Sn): Dissolve a pinch of powdered mineral in a few drops of HCl, add small shavings of zinc and heat over a flame; if Sn is present it will form a white precipitate on the Zn shavings.

Tin (Sn): Use the tests as for bismuth (see previous post). Sn will produce the same result, but the Sn-hydroxide will re-dissolve in excess NH4OH – and Sn-sulfide is a lighter brown color.

Uranium (U): Mix a pinch of the powdered mineral in a drop or two sulfuric acid (H2SO4), then heat the solution carefully to evaporate the liquid away until the remaining powder is almost dry. Add several drops of distilled water to make a yellow solution. Draw off the solution from the powder. Add tiny shavings of zinc to the solution. If U is present the solution will turn from yellow to green, and in time a yellow precipitate will form on the remaining zinc. (Note: If large amounts of iron or

vanadium are also present, they will interfere with this test.) [WARNING: Do not breath powdered U minerals or the fumes from the solution! Do this test in a well-ventilated area. U can cause radiation burns and cancer if breathed in. Use all due precautions. Exhaust all other possible means of identification – fluorescence, crystallographic, physical, geiger counter, etc. - before trying chemical ones.]

Vanadium (V)/ Vanadate (VO4): Dissolve a pinch of powdered mineral in a few drops of HCl, add zinc shavings and heat strongly for several minutes. If V is present the solution will turn first yellow, then green, and finally bluish-violet.

Vanadium (V): Another test is to slowly add NH4OH to the test solution until it is slightly basic (use litmus paper), then add two additional drops of NH4OH and one drop of fresh Na2S solution. If V is present the solution will turn a violet-red color. (Note: This test will not work if molybdenum is also present.)

Zinc (Zn): Dissolve a pinch of powdered mineral in a few drops of HCl, add several drops of fresh Na2S solution. If Zn is present it will form a white precipitate of ZnS.

GABBROS & BASALTS

Gabbro and basalt are two closely related rocks, distinguished only by their grain size – you can see the grains in a gabbro, but the grains in a basalt are so small that you can’t. Basalt is usually referred to as an *aphanitic * (very fine-grained) equivalent of gabbro. Gabbros are plutonic (deep-seated) rocks and crystallize from magmas that cool slowly, thus the grains can grow in size. Basalt is an extrusive or shallowly intrusive rock – formed from lavas that have reached the surface, or nearly so; thus the lava cools rapidly and the grains can’t grow.

Gabbros and basalts are actually a complex group of rocks with variations in chemistry and textures. The "Dictionary of Rocks” by Richard S. Mitchell has 28 entries for gabbro – and 50 entries for the various kinds of basalt! What makes them all similar is their essential minerals: calcic-plagioclase - usually bytownite or labradorite - with mafic (dark colored) pyroxenes - such as augite and hypersthene. In simple terms, gabbro is a common dark-colored coarse-grained plutonic igneous rock, and basalt is a very common dark-colored fine-grained extrusive igneous rock.

Minerals expected to be present in basalt include: copper, silver, prehnite, epidote, calcite, chlorite, quartz, and members of the Zeolite Group. In one type of basalt, called *amygdaloidal * because it has almond-shaped gas pockets in it. datolite, pumpellyite, and the adularia variety of orthoclase are often present. In pillow basalt, pectolite, babbingtonite and a few sulfides, such as pyrite, may be present. Gabbro has

much the same mineral content, but may also have magnetite, ilmenite, and nickel. Some zeolites likely to be present are: analcime, chabazite, heulandite, laumontite, mesolite, natrolite, scolecite, thomsonite, and stilbite. Apophyllite, prehnite, datolite and pectolite are closely associated with zeolites and should be looked for whenever zeolites are found in basalts or gabbros. (And vice-versa: If you find any of these "zeolite associates," look for the zeolites, too!)

On a recent two-day trip to northwestern Wisconsin, I was able to collect epidote, quartz, prehnite, calcite, chlorite, pyrite, copper -- just one tiny specimen - pumpellyite, magnetite, laumontite, and some unidentified pyroxenes and plagioclase feldspars at a basalt quarry and a gabbro quarry. I may have seen a few other zeolites, but did not try to collect or identify them because they lacked any crystal shape. There are many different mineral species found in basalt and gabbro, but the ones mentioned above are the most likely to be present. If an “unknown” mineral is not one of these minerals, then seriously consider some of the less common associated minerals such as albite, bornite, chalcopyrite, harmotone, hematite, phillipsite, and goethite.

Hydrothermal action may have invaded a basalt or gabbro, which can add lots of possible new minerals. Don't skip over a mineral species because it is seldom present in a certain mineral environment. Devote your attention to the likely suspects, but if they don't pan out, spread your search out. Alexander Falster, an authority on the Wausau and Stetton gabbro plutons in Wisconsin, collected many phenakite crystal specimens from them. Phenakite is a beryllium silicate present in granitic pegmatites and miarolitic pegmatites in granite. Beryl is also found in similar, but slightly different environments. Al found a beautiful aquamarine (beryl) crystal where phenakite should have formed. He explained that the crystal was on the fringe of the environmental area, so aquamarine, although not expected, could form instead of phenakite.

In the West Paterson, New Jersey area, there are several old “traprock” quarries in Triassic age basalt – the most famous being the New Street Quarry. These localities are noted for their specimens of prehnite, pectolite, chabazite, fluorapophyllite, heulandite, and other basalt species. Some of the materials from these quarries are considered “world-class” and grace the mineral rooms of major museums around the world.

The Connecticut River Valley in Connecticut and Massachusetts also has Triassic basalts that produce some fine specimens – prehnite, babbingtonite, calcite, amethyst quartz, datolite, and others. Babbingtonite is the State Mineral of Massachusetts because of the superb specimens that have been found at localities such as the Lane & Sons traprock quarries in Westfield and the Cheapside quarry in Deerfield. One of the oddities found in at least the Cheapside quarry is specular hematite. (Which reinforces

what Lloyd said above about widening the search for a mineral’s identity when the “usual suspects” don’t pan out…)

Gabbros and basalts provide a wide range of collecting opportunities. And knowing a little bit about them in advance of a trip to localities in these environments can go a long way towards making the trip productive – knowing how the rocks and minerals formed, what to expect, and what to look for at the site.

MASSIVE SULFIDES

Deposits of massive sulfides are found in several forms. Probably the two most prevalent types are *stratabound* deposits and *vein deposits.* The later are most often found along fault zones of one type or another. Both are usually associated with metamorphic rocks, though deposits in other settings are also known. I’m most familiar with the metamorphic stuff, so I’ll be concentrating on those here.

The stratabound deposits occur as layers sandwiched in between other layers of rock, and originated as sulfide-rich beds of volcanic materials – cooled lava flows, ash layers, and layers of sediments eroded off of volcanic island ranges or mountain chains. Volcanism often brings materials rich in sulfur and metals to the surface – laying them down in thick beds on the ground or sea floor. The ones I’m familiar with started out as rhyolitic tuffs and magnesium, iron and aluminum rich materials that metamorphism converts to *amphibolites* - schists and granulites composed largely of amphibole minerals such as hornblende and actinolite, usually with a fair amount of chlorite mixed in. The metamorphism of these rocks takes place after deep burial when tectonic forces exert high pressures and bring the rocks up to geologically moderate to high temperatures. Here in the Northeastern US such beds are found in Taconian age rocks created by volcanic island arcs that formed off the early coast of New England, and were then caught up in the “vise” of these islands and the seafloor colliding with the edge of the continent and being driven up onto the continental margin as the arcs and seafloor were driven land-ward by seafloor spreading further out, in the middle of the ancient Iapetus Ocean. During the height of this process there was a fair amount of hydrothermal activity – ocean water trapped in the rocks and brought up to high temperatures by the collision. These hydrothermal fluids served to enrich the stratabound massive sulfide beds by geochemically gathering in the sulfur and metal and depositing them along certain horizons or zones in the rocks.

There is evidence in New England which suggests that the stratabound massive sulfide “beds” actually occur in thrust fault zones that are regional in scale (see the post below on ultramafic tectonic slivers): The shearing forces of the island-continent collision created highly fractured and “mobile” zones at certain horizons in the rocks as they were metamorphosed, and these zones of weakness became the areas where

the hydrothermal fluids gathered in the metals and sulfur, depositing them. So technically – for the New England occurrences at least – the term “stratabound” is something of a misnomer. These deposits are closely related to the more typical fault zone deposits we see in the region as well.

Fault zone deposits of massive sulfides (whether along thrust faults or typical high angle faults, normal or reverse) are pretty much self-explanatory: Hydrothermal fluids rich in metals and sulfur coursing through the sheared rocks of fault zones deposit sulfide minerals as ore bodies along those faults. Here in New England quite a few of them occur due to magmatic doming during the Mesozoic era: Large bodies of magma – deep-seated liquid rock – slowly work their way towards the surface from the depths at which they are created. If they approach close enough to the surface their buoyancy actually lifts the overlying “veneer” of crustal rock; and since that rock is solid and somewhat brittle, it fractures from the stress and great blocks shift in relation to one another – faulting. [The magmas here, by the way, eventually cooled down and crystallized into plutons and a small batholith, called the “White Mountain Batholith.” These rocks contain pegmatites, and dikes and sills of pegmatite were forced out into the country rock – where we find a number of pegmatite minerals. See the post below on pegmatites for more info on these.]

Around where I live, we find quite a few fault zone deposits of massive sulfides, many of which have seen mining activity – people searching for ores of iron, sulfur, lead, zinc, copper, silver and gold. These deposits and the mines developed in them pretty much form a “halo” around the perimeter of the White Mountain Batholith. And it is my belief that the source of the metals and sulfur in these deposits are the older, Taconian age, *volcanogenic* stratabound massive sulfides discussed above. Just a little theory that I have – and that I am presently working on trying to prove or disprove. (Hey – the materials had to come from somewhere, and around here the stratabound massive sulfide “beds” are almost the only likely source…) Most of these deposits are rich in minerals such as pyrite (lots and lots of pyrite!), galena, sphalerite, and hematite - with some chalcopyrite and traces of arsenopyrite and other “primary” sulfide minerals. And massive sulfides usually carry a tiny percentage of silver and gold – though rarely (if ever…) in sufficient concentrations to be profitable on their own; as miners learned (to their dismay) when they tried to develop gold and silver mines on these deposits. If they didn’t capitalize on the base metal ores they were finding – turning the mines into lead, iron and zinc mines – they went bust. Those that read the writing on the wall (or in the rock, as the case may be) faired better, making a living for a space of time from base metals – with the occasional shipment of gold and silver as a bonus. Even most of the supposed “copper” mines around here really weren’t: Chalcopyrite – the principal copper ore – is also rather scarce. Some of the mines capitalized on the pyrite’s sulfur content, shipping ground pyrite to Boston for

processing to us the sulfur in the production of industrial sulfuric acid. In fact, the most profitable mines in the region’s massive sulfide deposits were those that went that route – considering all other ore production as subsidiary to sulfur production. The old Davis Pyrite Mine in Rowe, Mass., and the Ely Mine in VT are a couple of examples of such mines.

Whether we’re talking stratabound or faultbound, the mineral assemblages found in these deposits are largely a combination of sulfides (such as those mentioned above) and quartz. In places the sulfides are overshadowed by metallic oxides – particularly hematite, magnetite, and a bewildering array of manganese oxides that are damn near impossible to identify without high-tech lab equipment. These generally end up being called either “psilomelane” or “wad.” In one stratabound deposit down in the Berkshire Hills of western Massachusetts, the manganese was so rich it produced a deposit of rhodonite, rhodochrosite, and “wad” that was rich enough to be mined for the manganese – the old Betts Manganese Mine in Plainfield, Mass. A couple of places where the materials were particularly iron-rich are the Forge Hill Iron Mine in West Hawley, Mass. (specular hematite and magnetite), and at the Franconia Iron Mine on Ore Hill in the town of Sugar Hill, NH (magnetite.)

Often these deposits produce suites of secondary minerals – minerals that form after the deposit has basically been created. Percolating ground water – either geothermally heated to become hydrothermal fluid or simply cool water perking down through fractures and interstices as “epigene” or “meteoric” water – dissolve the primary minerals and redeposit them as crystals of the same minerals or chemically combine the constituents as new minerals. At one site where I am currently doing some study, this sort of activity has produced a major suite of unusual secondary minerals – unusual, at least, for New England: malachite, brochantite, linarite, native silver, acanthite, pyrite and chalcopyrite, cuprite sphalerite, cerussite and anglesite, posnjakite, schulenbergite, metazeunerite – and a few things we haven’t even identified yet. (We are up to “Unknown #37” – but some of the earlier ones have been identified since being listed as mystery minerals; so there are currently 15 to 20 which remain to be determined.) Anyway, good old Ma Nature seems to have had a field day with the fault zone sulfide deposit at the Mascot Lead Mine here in Gorham, NH - where I just happen to live! (Talk about the luck of the draw…) [PS: June 2001 update: We have determined that some of the secondaries found in one of the mine dumps formed in the dumps, not in the ore lode. Schulenbergite and posnjakite are two of them. "Modern mineralogy" :~} ]

More typical deposits tend to have fewer – if any – secondaries. The nearby Nay-Fogg Mine in West Milan, NH, hasn’t produced any good secondary minerals that I know of. The Silver Lake Lead Mine in Madison has produced a few – though nothing like

the plethora found at the Mascot Mine. The only other “rich” mine in New England that comes to mind is the old Manhan Lead Mine in Easthampton, Mass. – “Loudville” (the mining village near the mine that got its name because the miners were - well – a rather *loud* bunch…) That mine has produced secondaries such as wulfenite, pyromorphite, cerussite, and anglesite – the latter two in just about every conceivable crystal habit! There are even, I believe, still a few “mystery minerals” from there that await identification. (I have a shoe-box full of them that I inherited from an old friend who got them from another old friend of ours that studied the Manhan Mine minerals and wrote a paper on them for the “Mineralogical Record” back in the late ‘70s. Some day – if I ever find the time – I’m going to dig up that box and have some fun playing with Fred Lincks’ “Manhan Mystery Minerals.” Good hearted Fred – rest his soul – introduced me to the mineralogy of that mine back in the early ‘80s. Until I got to know him, I’d had no idea what a wealth of minerals the site produced. I thought it was just a place to go for quartz crystals, galena, and an occasional find of wulfenites. Little did I know…)

Finally, localities such as these often also produce “accessory” minerals – minerals that formed in the rocks around the deposit. Such as finds of almandine and epidote crystals at Ore Hill in Sugar Hill, NH, and finds of ferrohornblende crystals in the amphibolite containing the hematite and magnetite at the Forge Hill Mine in West Hawley, Mass. But here we are drifting afield from stratabound massive sulfides and getting into the metamorphic mineral environment – which I’ll discuss in another post to the Mineralogy-101 series.

1) Start out with simple observations and tests. Though they are simple, they can be telling:

...- Is it a metallic or submetallic mineral? Or is it nonmetallic?

...- If it is metallic or submetallic, will it readily mark a piece of paper? If it does, that leads you to one group of minerals. And these can be sorted out based on the color of their streak, hardness, and color of the mineral....-If it doesn’t mark paper, can you scratch it with a pocket knife? If it can, this leads you to another group of minerals that can be sorted out the same way as the softer group. If not, it leads you to a third group, also sortable the same way....The odds are in your favor that any metallic mineral at hand will fall into one of those three groups, and that once you are looking at species of the right group you will find the identity of your specimen in it.

...-If the mineral is non-metallic, does it have a definite colored streak? If it does, it will lead you to a group of minerals that can be sorted based on the color of the streak, the hardness, the color of the mineral, and other characteristics.

...-If the mineral does not have a colored streak, can you scratch it with your fingernail? If you can, you end up at another group of minerals that can be further subdivided based on whether or not the mineral has a prominent cleavage: If it does, you are lead to a group of minerals that can then be sorted based on the number of cleavages present, hardness, and color. If it doesn’t have a good cleavage you are lead to yet another group of minerals that are again sortable based on their characteristics.

And so on... (The methodology eventual takes you all the way up to very hard non-metallic minerals, including diamond. But I’m not going to lay it all out here – it is available to you - now - elsewhere here at Bob's.)

2) Mineral environments and associations are very important clues to a mineral’s identity. And this is a big reason why specimens of crystals on matrix are more “valuable” than loose crystals or clusters without any matrix. That rock, it’s mineral constituents, and any other crystals of different minerals with the one you are puzzling over, provide valuable information about what the mystery mineral might be – and often what it can’t be… You need to pay close attention to such info in field guides – and at mineral localities. Take your focus on the mystery mineral away from it long enough to consider what it occurs with. It will narrow down your list of “likely suspects” considerably.

3) If simple observations and tests, along with information provided by the matrix and associated minerals, don’t get you an identification, get a bit more “aggressive” in your search. Try things like specific gravity, crystal system and habit – maybe a few of the easiest chemical tests, perhaps some of the easier optical observations.

...- Specific gravity is a powerful tool in mineral identification. One which all too few collectors employ… Anyone who is at all serious about trying to identify minerals needs to get the equipment and learn this technique. Period!

...- Understanding at least the basics of crystallography – the differences between crystals in the seven crystal systems – is also a very powerful tool. With this under your belt you should be able to identify many of the crystal shapes you encounter – figure out at least which crystal system the mineral crystallizes in. Which effectively eliminates any mineral which crystallizes in any of the six other systems…

...- Basic optical techniques are also powerful tools. Along with physical characteristics, the optical properties of a mineral can narrow down the search for an identity by a wide margin – in your favor. (Optical observations are my short coming

in mineralogy. I now wish I’d paid closer attention and tried harder when the prof was trying to drill them into us...)

...- Flame and Bead tests are two more powerful tools. As Al O. has shown, they aren't that hard to do. A little practice and you're off and running. I've used them for years, a routine part of my mineral identification efforts. You should, too.

4) At some point you will undoubtedly run into something that defies identification using the above tools. At that point you have to decide how much you want to get into the nitty-gritty of mineral identification – go on to learn chemical testing, and advanced optical observations, or find a lab to do it for you. I don’t think any of us doing this seminar would say that chemical testing, or advanced optical work is for everyone. They aren’t. But I happen to think that anyone who really sets their mind to it can dope these out – become at least proficient enough to make good use of them. But you have to want to, and you have to be patient – take things a slow step at a time. If you do add these to your bag of tricks, there will be very few minerals you won’t be able to identify – some of the real “toughies,” like telling the various amphiboles apart, or the various zeolites, etc… Even some of the pros have a hard time with those – resorting to even more advanced techniques, such as SEM analysis and X-ray diffraction.

...- If you do decide to get into chemical testing, I strongly recommend that you take the “least is best” approach outlined in that topic below. And work at being as methodical in your testing using chemicals as you are (hopefully!) in your use of simpler observations and tests. This is the safest, most economical, and surest way to do chemical testing. The most practical by far... “Think Safe! – Be Safe!”