CHEMICAL TESTS FOR SMALL SPECIMENS
By Jesse Crawford

INTRODUCTION

This file is available in .pdf format. Please click on the following link:
Chemical Tests.

This is a work in progress. The objective is to collect together the chemical tests that are useful for identifying minerals and that are also within the reach of the typical hobbyist with the typical hobbyistís budget. If anyone would like to suggest a test for the collection, please email me (jesse@lava.net) with the details. Put ìMineral testî in the title so the message can get through my spam filter. Tests should be easy to perform, and should use materials that are reasonably easy to obtain.

The tests described here are intended for small samples. For most tests, a piece that's 1 or 2 millimeters across is adequate for at least 3 or 4 tests. That's about 5 to 20 milligrams.

Scientists have directed a lot of effort toward developing ways to make chemical tests on tiny samples using a microscope to interpret the results. Most of these tests have become obsolete in recent years, but they still offer useful and fairly low cost methods that amateur scientists can use to test minerals.

What follows is a description of some chemical tests that work reasonably well when scaled down to a size appropriate for testing tiny samples. The author has tried most, but not all, of the tests included. Not much detail is included about what positive or negative test results look like because it is assumed that the tests will be performed on both the unknown sample, and a sample that is known to contain the element being tested for. It's also a good idea to run the test on a blank sample known to not contain the element being tested for.

No test is perfect. Most of these tests offer a level of confidence probably no more than the 80 to 90 percent level, which is pretty good in a world as uncertain as this one. As I see it, it's the uncertainty that keeps things interesting. Remember, if youíre not having fun, then youíre not doing it right.

Chemists (at least old chemists) form the habit early in their careers of treating all chemicals as if theyíre dangerous.

BE SAFE! Respect the fact that chemicals can be hazardous. Scientists donít know all the ways that chemicals can injure people. You donít want to be the first to discover a new one. Don't let chemicals stay on your skin, and don't breathe them. If you can smell them, then you probably need to improve the ventilation.




SAMPLE PREPARATION

After we have a sample to test, the next thing to do is dissolve it, or at least dissolve enough of it to test. Ideally, the objective is to dissolve as much as possible of the sample so that the result is a drop of clear liquid about 10 to 50 microliters in volume containing all of the ions of interest. (in other words, a small drop). Most of the time, it's not necessary to go for complete dissolution. Often there will be a part of the sample that remains as pulverized fragments or sometimes a gelatinous mass of silica. If the grinding of the specimen is done with a mortar and pestle, the acid can be added while the grinding is being done. Then touching the pestle to a slide makes a drop that's sufficient for a test.

Almost all samples are prepared by dissolving them in some kind of acid. The following is a list in the suggested order to use in trying to dissolve the sample. If nothing in the list attacks the sample, then that's a lot of information already. The list of minerals that are impervious to all acids is a comparatively short one. A table of solubilities of some minerals is included at the end of this paper. Acids should be full strength. When one is found that attacks the sample, the solution can be diluted with a drop of water before beginning the tests. Some of the tests need to be carried out in a neutral or basic environment. Ammonia is handy for neutralizing acids.

THESE ACIDS ARE DANGEROUS! Handle them carefully in a well ventilated environment. Don't breathe the fumes. Be especially careful with fluoride minerals. Hydrofluoric acid and sometimes elemental fluorine is evolved when fluorides are treated with some acids. Itís very nasty stuff.

Water
Hydrochloric Acid
Nitric Acid
Sulfuric Acid
Aqua Regia (3 parts Hydrochloric 1 part Nitric) CAUTION! Chlorine is evolved from aqua regia.


Hydrofluoric acid, if it were less dangerous, would certainly belong on this list. It neatly solves the problem of dissolving silicate minerals by converting silicon to a gas, silicon tetrafluoride. With that goal in mind, there is an alternative to using a strong solution of hydrofluoric acid. Small samples of silicate minerals can be digested in platinum or teflon dishes with a mixture of sulfuric acid and calcium fluoride. Hydrofluoric acid is thereby generated ìin situî and immediately reacts with the silica in the mineral. The technique is not without dangers, but with proper precautions can be used when necessary. The hydrogen fluoride generated is still dangerous, and must be respected, but the risk is more manageable.

There is a method that can be employed to dissolve even the minerals that resist all the above acids. Heating the sample with a flux to a high temperature until it is thoroughly fused alters the composition of most minerals so that they can be dissolved in water or one of the above acids. The usual flux used is sodium or potassium carbonate, or for some minerals sodium or potassium bisulfate.

Fusing the mineral sample at red heat with a flux can induce almost any mineral to dissolve either in water or in one of the acids. These are extreme measures, and because they involve a lot more handling of the sample than simply treating it with acid it's usually good to start with a larger piece. 50 to 100 milligrams is good. Carbonate fusions can be carried out in a platinum crucible or piece of platinum foil, but bisulfate fusions should not be made on platinum, as the platinum will be attacked. Fusions with carbonate can be done in a small ceramic crucible, or on a block of charcoal, or a loop of platinum or nichrome wire using pretty much any small torch.

To do a carbonate fusion, start by grinding the sample as fine as possible. Add about twice the volume of dry sodium carbonate, and mix them. If you have a platinum crucible, then put in the sample mixed with flux, cover with a little more pure flux and support the crucible for heating. Begin heating the side of the crucible and as the mass begins to fuse, regulate the heat so as to avoid any loss of sample. The melt will evolve carbon dioxide and water vapor and possibly other gasses, and it will probably do a lot of bubbling. After the bubbles stop, raise the heat to redness and continue heating for 10 or 15 minutes, until itís thoroughly melted. Let everything cool down and add a few drops of nitric acid and a little water, and let it sit for a while. The melt will loosen and dissolve. Put the contents into a beaker, rinse the crucible with water, and add the washings to the beaker. Then set the beaker on a low source of heat so that the water and nitric acid can evaporate. It should not boil at any time. A double boiler arrangement is desirable for this phase of the operation. When the contents of the beaker are dry, add the minimum amount of water necessary to dissolve the soluble part. There may be an insoluble residue of silica. If the dried sample doesn't dissolve in water, one of the acids may be necessary. The fusion can also be done using a wire loop. Start with a hot loop, pick up as much sample and flux mixture as will stick to it, and fuse it. Then, touch the fused bead to the sample mixture to pick up a little more, and continue. Repeat the process until enough of the sample is fused.

The procedure for a bisulfate fusion is similar, but should be carried out in a porcelain crucible. It's messier, and the fumes are more toxic, so BE CAREFUL. During the bisulfate fusion, there's a lot of bubbling at first. After the bubbling stops, there comes a point where the melt solidifies, and a higher heat is needed to get it to fuse again. This is the point at which the generation of sulfur trioxide and other corrosive sulfur oxides begins, which is the objective of the procedure. If the melt is allowed to cool at this point, the process will not be complete. The heat should continue until the mass fuses again, and no further changes are in evidence. At this point, cool the melt, add a drop of concentrated sulfuric acid (carefully) and resume heating. This is repeated two times. Then the melt is cooled and removed from the crucible as above using a little sulfuric acid and water. Sulfuric acid gives off dense clouds of white fumes when it is heated to dryness. DON'T BREATHE ANY OF IT. This procedure is not for the faint hearted. It's noisy and hot and frightening and suitable only for a well ventilated garage or lab. Have a fire extinguisher close by and an escape route cleared in case of emergency. Other than that, it's kind of fun.

Vycor labware works well for fusions.

Sodium peroxide also makes a good flux. One author asserts that any mineral can be brought into solution by sodium peroxide fusion. Peroxide fusions are ordinarily carried out in a zirconium crucible.

DECIDING WHAT TESTS TO PERFORM

In deciding what tests to make, it's sometimes handy to remember that it can be just as valuable to know what isn't present in a sample as what is.

Once we know what will dissolve the sample, tables of the solubility of minerals can be consulted to help in selecting which further tests to undertake. Try to find a test that will split the list of possibilities in half. This has been called the half-split technique.

Whenever we read about a test, it usually starts out with a list of needed equipment and reagents, then a description of the procedure, and somewhere near the end will be a list of ions that interfere with the test. That's always the catch. There are very few tests that respond only to one element. Usually there's a list of them.

A lot of the difficulty with interfering ions can be sidestepped by careful selection of the sample. Picking a well formed crystal of the mineral of interest improves the chances that there won't be a lot of interfering ions. Naturally, those are always the prettiest crystals.

MATERIALS for SPOT TESTS

A small mortar and pestle for grinding samples.
A box of microscope slides.
A box of cover slips.
A glass or plastic ring about 15 or 20 millimeters in diameter ( it must be smaller than the cover slips) and about 2 to 3 millimeters thick
A glass rod 1 to 2 millimeters in diameter.
A small bulb type pipet (eyedropper).

REAGENTS for SPOT TESTS

Acetic Acid (Glacial)
Acetylsalicycilic Acid (Aspirin)
Ethyl Alcohol
Aluminon 0.1 percent solution
Ammonium Acetate Solution 3N
Ammonium Chloride
Ammonium Hydroxide
Ammonium Molybdate
Ammonium Oxalate
Ammonium Phosphate (Dibasic)
Aniline hydrochloride
Barium Chloride
Cesium Chloride
Chloroplatinic Acid
Citric Acid
Curcumin
Dimethylglyoxime
Hydrogen Peroxide
Hydroquinone
Lead Acetate
Oxalic Acid or sodium or potassium oxalate
m-phenylenediamine hydrochloride or sulfate
Potassium Dichromate
Potassium Iodide
Potassium Mercuric Thiocyanate (This reagent is made by combining mercuric nitrate with potassium thiocyanate in molar proportions of 1 part mercuric nitrate to 4 parts potassium thiocyanate. Tabular and needle-like crystals separate easily from acidic aqueous solution).
Potassium Nitrite
Potassium Phosphate (Dibasic)
Potassium or Sodium Sulfite
Rubidium Chloride
Silica sand
Sodium Acetate
Sodium Chloride
Sodium Fluoride
Sodium Phosphate (Dibasic)
Silver Nitrate
Starch
Tartaric Acid
Thiourea
Uranyl Acetate

The following are spot tests that are carried out on microscope slides and viewed through the microscope. Some authors recommend coating the microscope slides with wax or some other hydrophobic material to make it easier to control the drops. Some manufacturers make microscope slides with small wells that prevent solutions from running off the slide or to use with the "hanging drop" method (to be described below). They're all good ideas, yet just a plain microscope slide works fine for most tests. It's also a good idea to have a piece of black paper and a piece of white paper handy to put under the slide for contrast when viewing crystalline precipitates.

TECHNIQUES

The most general method for carrying out tests is to place a drop of the solution of the sample on a slide and put a drop of a reagent solution near it. Then a thin glass rod is used to bring the two drops together. The entire process is observed under the microscope.

Another important technique that's used is the hanging drop method. It's used to trap gaseous reaction products that are evolved from the sample as it reacts with a test reagent. For this technique a glass or plastic ring supports a cover glass with a drop of reagent or water hanging from the underside. The hanging drop is positioned over the sample, so it's close but not touching.

It is occasionally desirable to separate a drop of a solution from solid material, such as a precipitate or the fragments that remain after grinding the sample. It's often possible to precipitate an interfering ion and then move the clear sample solution to another slide for further tests. To remove the iron, for example, from a drop of solution, the pH of the sample solution can be raised by adding a drop of ammonia. At high values of pH, iron forms a dark gelatinous precipitate. To separate the sample from the iron precipitate a small piece of filter paper, an eighth of an inch or so in diameter, is placed on the slide near the sample. Then a dropper tube with an opening a little smaller than the diameter of the paper is pressed against the paper. The bulb of the dropper should be squeezed so that a small amount of suction will be supplied when the bulb is released. The tip of the tube with the filter paper is slid across into the sample drop, and the pressure on the bulb is released. If the dropper tube is not pressing too hard on the filter paper, the fluid will be drawn up into it through the filter paper, and it can be picked up and moved to another slide. As described, this procedure for separating iron is not selective, and would also leave behind other elements that precipitate at high values of pH, notably aluminum. Something else is needed to separate iron and aluminum (See "Aluminon Test" below). It takes a little practice to get this technique just right, but it opens a lot of possibilities when mixtures of ions interfere with one another. It helps to roughen the end of the tip of the dropper tube with fine sandpaper to prevent it from slipping off the filter paper when sliding it along the glass.


It is sometimes necessary to protect a glass slide or cover slip from the action of hydrogen fluoride. Plastic slides can often be used in these situations, or the glass can be coated with a hydrophobic material. Smearing grease on the glass works, but itís difficult to get a uniform thickness, and the irregularity of the coating can interfere with visibility. It works well to keep on hand a thin solution of microscope grease dissolved in xylene for this purpose. A drop is spread easily over the slide, and the xylene evaporates quickly, leaving a thin film of grease that prevents the hydrogen fluoride from attacking the glass.


TESTS for SPECIFIC IONS

In the descriptions of the tests to follow, references in parentheses following the name of each test are to the sources listed in the bibliography.

CATIONS

Aluminum

Aluminon Test: (Welcher) (Lange) If the sample is not already acidic, acidify it with dilute hydrochloric acid (1 drop). Some authors also recommend adding a drop of ammonium acetate buffer (3N). Place a drop of 0.1 percent Aluminon nearby and combine the two drops using a glass rod. A Red precipitate develops in the presence of aluminum and a number of other ions. If the precipitate persists after adding a drop of ammonium hydroxide, aluminum is indicated. This is the simplest form of the test and unless the material being tested is reasonably free of other ions, it is likely to produce a false positive. Aluminon is a very versatile reagent that can be used to detect very small amounts of aluminum, but in order to be confident of the results it must be recognized that it forms colored precipitates with a number of other ions. It forms a purple precipitate with iron, and red to brown precipitates with aluminum, actinium, barium, beryllium, calcium, cerium, chromium, europium, gadolinium, hafnium, indium, lanthanum, magnesium and neodymium, and white precipitates with antimony, bismuth, lead, mercury, and titanium. In order to interpret the result of this test it's necessary to separate the aluminum from at least some of these other ions, particularly iron.

Iron may be separated from aluminum by adding tartaric acid or citric acid to the sample solution. These will bind with the iron in such a way as to make it soluble in alkaline solution from which aluminum can be separated as a gelatinous hydrous oxide. Add a little of the tartaric or citric acid, and stir the drop until it dissolves, then add a small drop of ammonia to form the aluminum precipitate. It's important to realize that if the pH will go too high, the precipitate of aluminum will re-dissolve. This does not ordinarily happen with ammonium hydroxide, but if it does, warm the slide a little to drive off some of the ammonia. When the ph of the drop is in the right range, the aluminum will be in the form of a white translucent gelatinous precipitate and the iron will still be in solution. The aluminum precipitate will be translucent white. If it takes on a dark color then it probably means that the iron (or something else) is also precipitating. Add more tartaric acid and adjust the amounts until it looks right. Then use the eyedropper to filter off the liquid phase, or decant it carefully. Add a drop of water to wash the precipitate and draw that off through the filter too. Repeat the wash a couple of times. Then dissolve the white precipitate containing the aluminum and test it with the aluminon reagent as described above. It's good to practice on some fake "unknowns" until you have a feel for the amounts needed to make it go right.

Aluminon works better as a reagent for use with the enhanced spot tests described later.

Ammonium Molybdate Test: (Chamot and Mason) This test should be carried out at a fairly neutral pH. To the drop of sample, if it is not already neutral, add a small drop of saturated sodium acetate buffer. Then add a small pinch of ammonium molybdate. Watch for the formation of transparent four sided plates indicating the presence of aluminum. At first the crystals look square, but on closer examination they prove to have a more interesting shape. Too much buffer tends to inhibit the formation of the crystals. These crystals show symmetrical extinction when viewed between crossed polarizers. The presence of some ions can inhibit their formation. Warming the slide and/or adding more water to the drop sometimes helps. These should always be tried before deciding that the test is negative for aluminum. Nickel and iron both form similar crystals. Mercury forms six sided slightly elongated crystals under these conditions.

Barium

Sodium Bicarbonate Test: (Chamot and Mason) See the sodium bicarbonate test under ìCalciumî below. Barium carbonate crystals form more slowly than calcium or strontium carbonates, and the crystal are larger and more well formed.


Beryllium

Aluminon Test: (Welcher) See the test for aluminum. A red precipitate develops in the presence of aluminum or beryllium. The red color from the beryllium looks much like the color from aluminum, but when ammonia is added, the red precipitate dissolves if it's beryllium. This is not a very good test for beryllium, because several of the other ions that form red precipitates with aluminon behave the same way. Aluminum is the only one that does not dissolve when the ammonia is added.



Potassium Oxalate Test: (Chamot and Mason) This test produces characteristic crystals of a double salt of potassium and beryllium oxalate. A large drop of potassium oxalate solution is placed near the sample and the two drops are joined using a glass rod. As the water evaporates, rhombs and prisms will become evident if beryllium is present. It is easy to mistake crystals of potassium oxalate for the double salt, so care should be exercised in interpreting the result of this test. The double salt is strongly birefringent, and exhibits an extinction angle of 39 degrees. Prisms are sometimes formed with difficulty, but if the solution is heated so the crystals re-dissolve and a tiny drop of a solution containing mercuric ions is added, the prisms will have more of a tendency to form as the solution cools.

Both of the photos show the results of a positive test for beryllium, made on a known sample of phenakite. The well developed crystals in the upper one appeared only after recrystallization. These are crystals of the double salt of potassium and beryllium oxalate. The extinction angle is 39 degrees from the direction of elongation of the prisms. The colors are due to the birefringence of the crystals viewed between crossed polarizers. The lower photo includes crystals of other compounds as well, and is more difficult to interpret.
Curcumin Test: (Chamot and Mason) (Smith) See Curcumin Test for borate under ìAnionsî below.
Bismuth
Thiourea Test: (Chamot and Mason) (Lange) Thiourea added to a solution containing bismuth in nitric acid makes a strong yellow colored solution.
Dimethylglyoxime Test: (Budavari) A sample containing bismuth forms a bright yellow color and precipitate with this reagent.
Boron
See borate under ìAnionsî below.
Cadmium
Potassium Mercuric Thiocyanate Test: (Chamot and Mason) (Schaeffer) Put a small drop of potassium mercuric thiocyanate solution near the sample drop. Combine the two drops with a thin glass rod and watch for the characteristic crystals that indicate cadmium. This test is also used for other ions. Crystal shapes are distinctive for each type of ion.

Oxalic Acid Test: (Chamot and Mason) Cadmium oxalate crystallizes as long prisms with oblique ends, or as Xís or radiating groups. From concentrated solutions it forms octahedrons. Calcium, zinc and strontium interfere with this test.


Calcium, Barium, Strontium
These three elements are difficult to separate because they have very similar chemistries. Getting a good identification is possible by using two of the following tests in combination. The difference in solubilities of the sulfates makes a good way to tell the difference between them.
Ammonium Oxalate Test: (Chamot and Mason) Add a small drop of sodium acetate buffer to bring the pH to neutral. Place a drop of the reagent near the test sample and join the two drops with a glass rod. A white precipitate indicates calcium or strontium or barium. The test can also be made using a crystal of oxalic acid. The crystals typical of calcium are quite small, squarish tablets. Strontium oxalate looks much the same. The crystals are larger, and some are elongated but differentiating the two types of crystals in a mixture of both is not practicable. Barium makes distinctive crystals with oxalic acid. They assume the form of branching tree-like structures. The presence of calcium or strontium will suppress the formation of the barium crystals. Barium oxalate is very soluble in acids. If just a trace of nitric acid is present, the crystals will not form.

Sodium Bicarbonate Test: (Chamot and Mason) To the sample drop add a small drop of saturated sodium acetate to adjust the pH to a value near neutral. Then add a pinch of sodium bicarbonate. If calcium is present, small crystals of calcium carbonate will begin to separate out, floating on the surface and adhering to the slide. After some of the water has evaporated, larger crystals of the double salt of sodium and calcium carbonate will form beginning at the edge of the test drop. Itís not always obvious which is the calcium carbonate and which is the double salt, however the double salt is more soluble in water. Calcium carbonate, once it has precipitated from neutral solution, will not redissolve on the addition of water. The double salt of calcium and sodium carbonate can be redissolved by adding more water to the drop. This is a good test for calcium. Strontium and barium, which also precipitate as insoluble carbonates, do not form a double salt under these conditions.

Sulfuric Acid Test: (Chamot and Mason) Add a small drop of sodium acetate buffer to make the pH neutral and join the sample drop with a drop of dilute sulfuric acid. In the presence of calcium, prisms of calcium sulfate separate gradually from the solution. The ends of the prisms are terminated at an angle of 66 degrees, which serves to confirm their identity. Twinning is common. The precipitation with barium and strontium is too finely divided to recognize crystal forms. There are several other elements that react to form insoluble sulfates, so itís best to do a preliminary separation with oxalic acid or sodium bicarbonate, so that the other elements will not interfere. The oxalates and carbonates of calcium, barium and strontium tend to adhere to the surface of the microscope slide, so the precipitate containing these can be carefully washed and redissolved prior to making the sulfate test. The precipitate containing strontium can be recrystallized from hydrochloric acid to produce recognizable crystals, but they look a lot like the calcium oxalate crystals, so itís not really worth the effort. The solubility of the sulfates of calcium, barium, and strontium differ widely. Calcium sulfate dissolves readily in hydrochloric acid. Strontium sulfate dissolves slightly, while barium sulfate is virtually insoluble.

Cobalt


Potassium Mercuric Thiocyanate Test: (Chamot and Mason) Adding potassium mercuric thiocyanate solution to a sample containing cobalt results in deep blue crystals like the ones shown in the photo. These crystals are somewhat more soluble than the ones that develop in the presence of other ions, and sometimes do not appear until the drop has been allowed to sit for a while so that some of the water has evaporated.

Quinoline - Ammonium Thiocyanate Test: This test responds to Co, Fe, Mo, Ti, U, V, and Zr. For this test, the sample should be dissolved in strong hydrochloric acid. A small drop of quinoline mixed with an equal volume of 6N hydrochloric acid is mixed with the test drop. It is then connected in the usual way with a drop of saturated ammonium thiocyanate. An oil phase separates out quickly and over time crystals develop from the droplets of oil. After a half hour or so, the drop becomes filled with ammonium chloride crystals from the reaction between the hydrochloric acid and the ammonium thiocyanate. Antimony and bismuth must be absent for this test to work, as these cause an immediate precipitation when the quinoline reagent is added to the sample. In the case of Mn, Cd, Sn, and Hg (and possibly others), crystals may be formed immediately before the addition of the ammonium thiocyanate solution.

Cobalt develops light blue dendrites and blue crystalline blades from blue oil droplets. In the case of titanium, the oil is yellow to orange and the crystals, if they appear, are small thin yellow discs, scales and elongated hexagons or prisms. Zirconium yields a colorless oil and thin scales, plates, and rosettes from yellow to orange in color. Vanadium causes a colorless oil, and crystals are difficult to form. Uranium causes a yellow oil with rectangular plates and prisms of a light yellow color. Molybdenum produces a reddish oil but seldom produces crystals. Nickel and copper both yield dark colored oils but rarely produce crystals.

Quinaldine ñ Ammonium Thiocyanate Test: Quinaldine is a compound almost identical to quinoline. Itís comprised of the same heterocyclic ring system with a single methyl substituent. Itís chemistry is much the same as quinoline, and when substuted for quinoline in the above test protocol, the results are similar. The crystal shapes and colors produced are a little different however, presumably because of steric effects due to the methyl group. If the test solution is not already in hydrochloric acid, it should be evaporated to dryness and dissolved in a drop of hydrochloric acid before adding the quinaldine. If crystals do not develop readily, sometimes it helps to also add a drop of water to the sample. In the case of some ions (Co, Cu, Ni, U) crystals can take up to two or three hours to develop. The slide must be covered in those cases with an inverted petri dish or watch glass in order to retard evaporation.

The presence of cobalt in the test drop is indicated by a blue oil separating out. Crystals do not form immediately. If conditions are not perfect, they do not form at all. If the crystals donít make their appearance before the slide becomes covered with a mixture of ammonium chloride and unreacted quinaldine hydrochloride crystals, then itís too late. These metastable states are common with a number of ions, making this test a little frustrating at times. Iím still working on improving it because crystals, when they can be coaxed to appear, can be quite distinctive.

Copper

Ammonia Test: (Chamot and Mason) A dilute nitric acid solution of copper ions will turn a strong characteristic blue color with the addition of a drop of ammonia. This is not a precipitate, tetraamine copper ions are soluble but strongly colored.

Triple Nitrite Test: (Schaeffer)(Chamot and Mason) Evaporate the sample to dryness and then just cover the residue with a small drop of 30 percent acetic acid. Add a small crystal of sodium acetate. Wait for the crystal to dissolve and add a small crystal of lead acetate. When that has dissolved, add a crystal of potassium nitrite. Characteristic crystals will form if copper is present. The triple nitrite test has several variations and is used to test for a number of ions. It requires some practice, and even then the results can be confusing. Thereís a good discussion in ìHandbook of Chemical Microscopyî by Chamot and Mason.

Quinaldine Ammonium Thiocyanate Test: See the notes for this test under ìCobaltî above. The coordination compound made by copper and quinaldine produces clusters of long slender crystals arranged in branching structures, dark reddish brown in color. These crystals only appear when conditions are perfect, and they take a long time to form. Ordinarily they donít put in an appearance before patience runs out. The appearance of the oil is similar to the oil that separates when the sample contains nickel.





Potassium Mercuric Thiocyanate Test: (Chamot and Mason) Copper causes long green needle shaped crystals to form when combined with potassium mercuric thiocyanate reagent. These crystals, like the ones that develop with cobalt, are fairly soluble. Allowing the slide to sit undisturbed for a period of time while the water evaporates produces crystals like the ones in the photos above.

Gold

Potassium Mercuric Thiocyanate Test: (Chamot and Mason) Adding a drop of potassium mercuric thiocyanate reagent and joining it to a drop of sample solution containing gold causes the immediate separation of a densely branched structure of finely divided crystals. The crystals have a reddish hue and are unmistakable, making this an easy test for the presence of gold.


Potassium mercuric thiocyanate is a versatile reagent, but it must be realized that with mixtures of ions, the results can be variable and confusing. It works best when one ion dominates the sample mixture. Chamot and Mason give an extensive discussion of the behavior of this reagent under various conditions in ìHandbook of Chemical Microscopy.î




Iridium
Thiourea Test: (Chamot and Mason) In a solution of the sample in concentrated hydrochloric acid, a few crystals of thiourea are added. If iridium is present, the reddish color of the sample drop will decolorize, and become water clear. Iridium does not cause the formation of crystals.



Iron

A quick assessment of iron can be made by adding a little sodium hydroxide solution, or ammonia to the unknown. Iron will cause a precipitate immediately of a dirty green color if itís ferrous, or brown if itís ferric. This is a quick test for iron, and can easily lead to false conclusions unless followed up with more specific tests, because there are several elements that yield gelatinous precipitates with bases. Specifically, the following ions all yield gelatinous hydrous oxides under these conditions: aluminum, chromium, tin, titanium, zirconium, hafnium, thorium, bismuth, and uranium. There are probably others.

Iron is a common element in the earthís crust. In minerals, it assumes one of two oxidation states, either the +2 or ferrous state, or +3, ferric. Itís sometimes important to be able to determine which of the two states are present. Often both are, and if both, then itís good to get some idea of the ratio. This ratio is destroyed if the sample is dissolved in nitric acid, since nitric acid is a strongly oxidizing acid, all iron in nitric acid solution is in the ferric state, even if it was originally ferrous. This difficulty does not apply if the solution is made in hydrochloric acid. Ferric iron can be changed into ferrous iron by the addition of a reducing agent, such as sodium sulfite. Ferrous iron can be changed back to the ferric state by adding an oxidizing agent such as hydrogen peroxide. Because some of the reagents respond only to iron in one state and not the other, advantage can be taken of these facts to design a sequence of operations that will give a pretty good idea of the proportions of each in an unknown sample. Thiocyanate produces a red color in the presence of ferric iron, but remains colorless if only ferrous iron is present. If a solution in hydrochloric acid is first treated with ammonium or potassium thiocyanate (potassium works best)the red color, if there is any, gives an estimation of the ferric iron present. If the color increases significantly after adding hydrogen peroxide, then the change in the color gives an idea of how much ferrous iron was in the sample.

Quinoline Test: (Chamot and Mason) See ìQuinoline Testî under ìCobaltî above.



Quinaldine Test: See ìQuinaldine Testî under ìCobaltî above. Iron causes the immediate separation of a dark red oily phase and the prompt separation of dark red, almost black rectangular tabular and skeletal crystals, some elongated, as well as elongated blades and rhombohedral forms. Dendritic clusters predominate. A concentrated solution of iron turn opaque very quickly. Because of the strong color, this is a very sensitive test for iron.

Iron (Ferric)

Potassium Ferrocyanide Test: (Chamot and Mason) Potassium Ferrocyanide and ferric (Fe3+) iron produce Prussian blue.

Ammonium or Potassium Thiocyanate Test: (Chamot and Mason) Either of these reagents react with a solution of ferric ions to produce a red color.

The ferrocyanide and thiocyanate tests for iron may fail in the case of minerals that contain phosphate, fluoride, or borate. Also cobalt, chromium, nickel and copper interfere.

Iron (Ferrous)

Potassium Ferricyanide Test: (Chamot and Mason) Potassium Ferricyanide and ferrous (Fe2+) iron produce Turnballís blue.
Orthophenanthroline is an excellent reagent for detecting ferrous iron.

Lead


Thiourea Test: (Schaeffer) This test must be carried out on a solution of the sample in nitric acid. Add a small drop of nitric acid to the sample and a small lump of thiourea. Characteristic crystals will slowly form in the presence of small amounts of lead. It's important to observe the form of the crystals. Other ions may form other kinds of crystals and if large quantities of some impurities are present, the crystals may not form at all. The form of the crystals varies depending on the acidity and concentration of lead present. Thiourea makes distinctive crystals with several elements, including gold, platinum, ruthenium, palladium, rhodium, and osmium. Unfortunately, most if not all of these elements can produce crystals of several different habits, depending on the concentration, pH, and the nature and amount of interfering ions. Consequently, interpreting the results of a thiourea test is something of an art. More so than with most tests, running standards and blanks in parallel with the test sample is necessary in order to have much confidence in the results. Itís worthwhile to take some extra precautions to make sure the thiourea is pure. Thiourea can be purified by recrystallization from alcohol.

Hydrogen Chloride Ammonia Test: (Chamot and Mason) This test for lead is also a test for silver and for mercury. In a dilute nitric acid solution of the sample, add a small drop of hydrochloric acid. A white precipitate indicates lead, silver, or mercury. Adding a drop of ammonia will dissolve the precipitate if it's silver. If not, then lead or mercury is indicated. Remove the liquid by touching the edge of the drop with a piece of filter paper. Then add two drops of water and heat, but not to boiling. Put a small drop of potassium dichromate solution near the hot solution and combine the two drops with a glass rod. A yellow precipitate indicates lead. These tests are sometimes ambiguous because silver and lead often occur together and sometimes all three can be present. Lead and mercury chloride are more soluble than silver chloride. This fact can be exploited by washing the precipitate several times with warm water to separate the silver from the other two.

Potassium Iodide Test: (Chamot and Mason) Drop a few small crystals of potassium iodide into the sample solution. Lead will cause a yellow precipitate.




Magnesium

Sodium Phosphate (dibasic) Test: (Chamot and Mason) To the sample drop add a few crystals of ammonium chloride, stir and add a few crystals of citric acid. Warm and stir until dissolved. Add a crystal of disodium phosphate, warm gently and stir. Put a drop of strong ammonium hydroxide near the sample and cause the two drops to join using a glass rod. Ammonium magnesium phosphate slowly develops as dendritic forms, feathery stars, and Xís turning into plates and tabular forms. The precipitate can be recrystallized by decanting, then dissolving the crystals in dilute hydrochloric acid and precipitating with ammonium hydroxide. This should be done in order to reduce the chance of false results from interfering ions. Similar double ammonium phosphates are formed with Fe2+, Mn2+, Co2+, and Ni2+. Of these only Mn2+ precipitates (partly) in the presence of citric acid, and then only if the Mn is in high concentration. If in doubt, decant the crystals, wash with distilled water, and add hydrogen peroxide. If manganese is present the crystals will turn brown.

In ammoniacal citrate solution, disodium phosphate will completely precipitate Mg, Ca, Sc, Pb, Au, and the rare earth elements. In addition, Be, Sr, Ba, Hg, In, U, Zr, and Mn are partially precipitated.

Manganese

Sodium Phosphate (dibasic) Test: (Chamot and Mason) See sodium phosphate test under magnesium.

Sodium Bismuthate Test: Manganese dissolved in dilute nitric acid gives a purple color with sodium bismuthate.

Acetylsalicylic acid test: (Welcher) The reagent must be freshly prepared by dissolving a 15 grain aspirin tablet in 1 ml of 10 percent ammonia. Add 0.5 ml of hydrogen peroxide solution. The color developed is red to reddish brown in the presence of manganese. Iron also produces a strong color which may mask the results. The color produced with iron is dark brown at high concentrations and brown to yellow at low concentrations.

Ammonium Molybdate Test: (Chamot and Mason) Evaporate the sample drop to dry without overheating. Place a very small crystal of ammonium molybdate on the spot and put a drop of water on it. Set aside for a half hour. The orange crystals produced in the presence of manganese are markedly dichroic, going from red-orange to a pale yellow color as the polarization of the light is rotated. Itís important not to use too much ammonium molybdate. If too much is used, the result will be a white crystalline mass that covers the entire spot, making any red-orange crystals difficult to see. This is a good test, but not very sensitive. Manganese concentration needs to be at least two parts per thousand. The presence of significant amounts of copper, chromium, strontium, titanium or tungsten can prevent the crystals from developing.

Mercury

Hydrogen Chloride - Ammonia Test: (Chamot and Mason) To a nitric acid solution of the sample, add a small drop of hydrochloric acid. A white precipitate forms if silver or mercury or lead are present. Adding a drop of ammonium hydroxide will cause the precipitate to go back into solution if it is silver. See notes for the lead test above.

Potassium Iodide Test: (Chamot and Mason) Add a tiny crystal of copper sulfate to the sample drop. Put a drop of potassium iodide solution nearby and bring the two drops together in the usual way. Red mercuric iodide indicates the presence of mercury.

Ammonium Molybdate Test: (Chamot and Mason) See the discussion of this test under ìaluminumî above.

Molybdenum
Dipotassium Phosphate Test: (Chamot and Mason) The test sample must be strongly acidified with nitric acid, A solution of dipotassium phosphate is combined in the usual way, and, if no precipitation occurs, warm the slide gently. Then set the slide aside to cool. Examine at high magnification. If molybdenum is present, small yellow isotropic octahedral crystals are formed. If the principal element is tungsten the crystals will be white. A negative test result does not mean that molybdenum or tungsten are not present, only that they are not present in the form of molybdate or tungstate ions. If diammonium phosphate is used instead of the dipotassium salt, the test is more sensitive, but itís harder to read.


Nickel

Quinaldine Ammonium Thiocyanate Test: See the notes for this test above under ìCobalt.î Crystals of the nickel quinaldine coordination compound are deep garnet red. They are rarely seen however, because conditions must be perfect, and even then theyíre very slow to form. Dark, almost black drops of oil separate out on addition of the reagent. The initial appearance is similar to the oil droplets seen when the sample contains copper.
Dimethylglyoxime Test: (Schaeffer) (Lange) This reagent forms a bright red precipitate in the presence of nickel. Make the sample alkaline with a drop of ammonium hydroxide. Put a drop of saturated dimethylglyoxime in water near the sample, and combine the two drops with a glass rod. A deep pink or magenta precipitate indicates nickel. It might be necessary to warm the sample.

Ammonium Molybdate Test: (Chamot and Mason) See the discussion of this test under ìaluminumî above.


Osmium
Thiourea Test: (Chamot and Mason) Add a few crystals of thiourea to the sample dissolved in concentrated hydrochloric acid. In the presence of osmium, a red color develops immediately and, over time, red crystals form.
Palladium

Dimethylglyoxime test: (Smith) Dimethylglyoxime forms a yellow precipitate with palladium under acid conditions which is soluble in a solution made basic by ammonia.

Thiourea Test: (Chamot and Mason) Adding a few crystals of thiourea to a drop of the sample in concentrated hydrochloric acid causes an orange or yellow region to develop around the crystals, the outer edge of which is crystalline. In concentrated solutions of palladium, the orange crystals form closely around the reagent crystals and prevent them from dissolving.



Platinum

Potassium Chloride Test: (Schaeffer) Octahedral crystals of potassium chloroplatinate form when a drop of potassium chloride solution is joined to a drop of a sample solution containing chloroplatinic acid. This is the form produced by the action of aqua regia on platinum. Rubidium chloride can also be used in place of the potassium chloride. The rubidium chloride test is more strongly colored.

Thiourea Test: (Chamot and Mason) Adding a crystal of thiourea to a solution containing chloroplatinic acid causes a yellow reaction followed by reddish brown feathery dendrites.
Potassium

Uranyl Acetate Test: (Schaeffer) Characteristic crystals of potassium uranyl acetate are formed in the presence of potassium.

Tartaric Acid Test: (Schaeffer) Tartaric acid causes crystals typical of potassium acid tartarate to precipitate if potassium is present. Ammonia must not be present for this test to work. Add a little sodium hydroxide solution and warm the slide first to remove it. Then make the test for potassium.

Chloroplatinic acid Test: (Schaeffer) To use this reagent, ammonium must not be present. Add a little sodium hydroxide solution and warm the slide first to remove it. Then place a drop of chloroplatinic acid solution near the sample, and proceed in the usual way. In the presence of potassium, characteristic crystals will form. The test can also be used in the same way to test for the presence of ammonium. If a mixture of ammonium and potassium is suspected, use the hanging drop method to trap the ammonia in a drop of water. Then test the water drop separately.

Ruthenium
Thiourea Test: (Chamot and Mason) This test works only on a sample dissolved in concentrated hydrochloric acid. The sample solution should not be too darkly colored, if it is, dilute it with concentrated hydrochloric acid. Add several small crystals of thiourea to the test drop. Warm the slide gently. Over time, a blue color will develop in the presence of ruthenium.
Silver

Hydrogen Chloride ñ Ammonium Hydroxide Test: (Schaeffer) To a nitric acid solution of the sample, add a small drop of hydrochloric acid. A white precipitate forms if silver, lead or mercury are present. The silver precipitate will dissolve in ammonium hydroxide.

Sodium

Uranyl Acetate Test: (Schaeffer) This test is conducted in the usual way. Characteristic crystals of sodium uranyl acetate will form in the presence of sodium.



Fluosilicic Acid Test: The initial step of this procedure is to create a drop of water containing fluosilicic acid. Mix powdered silica sand with powdered calcium fluoride and put it in a small lead dish. Add a drop of concentrated sulfuric acid. Then put a hanging drop of water over it. It works well to use a cover slip coated with a very thin film of something hydrophobic such as stopcock grease. Silicon hexafluoride is evolved from the acid mixture and is trapped in the drop of water where it breaks down forming silicic acid which separates out, and fluosilicic acid which remains dissolved in the water. After a few minutes, lift the cover slip and touch the drop to a microscope slide. Then put a drop of the sample solution nearby, and cause the two drops to join using a thin glass rod. Set the slide aside for several minutes while the water evaporates. If sodium is present, hexagonal crystals of sodium fluosilicate, some looking like little flowers, will appear beginning near the edges of the drop. These crystals have a very low index of refraction so they may be difficult to see. If necessary, let the slide dry completely and examine it with a high powered objective. Often the crystals appear as small prisms lying on their sides with irregular terminations.

Strontium

Sodium Bicarbonate Test: (Chamot and Mason) See the sodium bicarbonate test under ìCalciumî above. Strontium carbonate looks much the same as calcium carbonate at first, but does not form the double salt, and after standing for a while it forms small acicular tufts of crystals attached to the slide near the test reagent. These are easy to overlook.

Tin
Potassium Iodide Test: (Chamot and Mason) A yellow to reddish orange precipitate is formed with the addition of a solution of potassium iodide to a sample containing stannic tin (+4). If the sample contains stannous ions (+2) the precipitate is a lighter yellowish white which changes to orange in the presence of an excess of potassium iodide.

Stannic Tin (Sn+4)

Cesium Chloride or Rubidium Chloride Test: (Schaeffer) (Chamot and Mason) These tests are conducted in the usual way. Small crystals characteristic of tin can be recognized. It is difficult to have much confidence in this test, since the reagents form insoluble crystals with a number of other ions. If the sample is first treated with nitric acid and evaporated to dryness on a double boiler several times before making the test, all the tin will be converted to an insoluble hydrous oxide, which can be washed several times again with dilute nitric acid. This removes many of the interfering ions. Then the insoluble oxide can be dissolved in hydrochloric acid and tested for tin as described above.

Stannous Tin (Sn+2)

Oxalic Acid or Alkali Oxalate Test. (Chamot and Mason) The addition of oxalic acid or a solution of an alkali oxalate causes a precipitate of irregularly shaped crystals. Prisms, if formed, exhibit either parallel extinction or, if twinned, an extinction angle of approximately 15 degrees to the direction of elongation.

Titanium
Quinoline Test: (Chamot and Mason) See ìQuinoline Testî under ìCobaltî above.

Tungsten

Dipotassium Phosphate Test: (Chamot and Mason) See ìDipotassium Phosphate Testî under ìMolybdenumî above.

Vanadium
Quinoline Test: (Chamot and Mason) See ìQuinoline Testî under ìCobaltî above.

Quinaldine Test: See ìQuinaldine Testî under ìCobaltî above.

Uranium
Sodium Fluoride Bead Test. This test is incredibly sensitive, however it is not very specific to uranium. A sodium fluoride bead is made by heating a loop of platinum wire until it is red hot and touching it to some sodium fluoride so that a little of it adheres to the loop. This is re-heated until fused, and the bead is built up in increments until it reaches the desired size. Then the hot bead is used to pick up a bit of the pulverized sample and heated thoroughly until itís completely fused. Let the bead cool, and examine it under an ultraviolet light source. If uranium is present, the bead will glow brightly with green fluorescence. Because there are other elements that can also cause fluorescence, this test should be followed up with a confirmatory test. There is a lot of old literature devoted to bead tests. They are still some of the most useful of field tests. With some practice itís possible to glean a great deal of information from them.
Quinoline Test: (Chamot and Mason) See ìQuinoline Testî under ìCobaltî above.




Potassium Oxalate Test: This is not an especially sensitive test for uranium, but because the crystals produced are dichroic, itís fairly definitive. A large drop of sample containing something on the order of a half a milligram of uranium in dilute nitric acid is combined with a similarly sized drop of saturated potassium oxalate solution. Crystals of oxalic acid are immediately precipitated. Over time these re-dissolve leaving small pale yellow rectangular prisms and tablets of the uranium compound. These are dichroic, going from pale yellow to colorless as the polarization of the light is turned through 90 degrees. The crystals are small and require an hour or two to develop. After three or four hours they will be large enough to easily determine their dichroic character. The best crystals seem to develop near the edges where the two drops come together.

Zinc



Potassium Mercuric Thiocyanate Test: (Schaeffer) (Chamot and Mason) Put a small drop of potassium mercuric thiocyanate solution near the sample drop. Bring the two drops together with a thin glass rod and watch for the characteristic crystals that indicate Zinc.
The photos show crystals of zinc mercuric thiocyanate. These are typical of the crystals that form after adding a solution of potassium mercuric thiocyanate to a sample containing zinc. The graceful branching is typical of a solution with a high concentration of zinc.
Sodium Bicarbonate Test: (Schaeffer) Expose the test drop to ammonia fumes long enough to make it alkaline, or add a small drop of sodium hydroxide solution, then join with a drop of saturated sodium bicarbonate solution. Watch for the formation of characteristic crystals that indicate zinc. The reaction begins as a slightly milky area where the two drops join, and the crystals grow slowly. Avoid stirring the drop when adding the baking soda. If it is agitated, the crystals that form may be too small to be seen even at maximum magnification.

Zirconium
Quinoline Test: (Chamot and Mason) See ìQuinoline Testî under ìCobaltî above.

Quinaldine Test: See ìQuinaldine Testî under ìCobaltî above. Red crystals of the quinaldine zirconium complex look like little red footballs. They tend to be a little slow in forming and develop from a light colored oil that separates on addition of the reagent. This is another one thatís a little tempermental. Conditions must be just right for the crystals to develop, and often they donít.




ANIONS

Borate

Curcumin Test: (Smith) (Chamot and Mason) Use an alcoholic solution of curcumin (0.5 percent). This test works well on paper. First put the spot of unknown on the paper and let it dry. Then put a spot of sodium fluoride solution over the spot and let it dry. Then add a drop of dilute hydrochloric acid. Let the spot evaporate until almost dry. Then add a drop of alcoholic solution of curcumin (0.5 percent). A red color indicates boron or beryllium. Then hold the test strip over a bottle of ammonia so the fumes can reach it. If the spot turns blue, boron is indicated. Titanium, columbium, molybdenum, tantalum, and zirconium interfere.

Bromide

m-phenylenediamine or aniline Test: (Schaeffer) Either of these reagents can be used to demonstrate the presence of bromide ions. The procedure involves the use of the hanging drop technique to trap the volatile bromine as it is released from the sample. First, place a small ring about 2 mm in thickness around the sample drop. Add several crystals of potassium dichromate to the sample and warm it until it's dry. Then add a drop of concentrated sulfuric acid. If bromide ions are present in the sample, free elemental bromine will be evolved from the reaction. This bromine must be trapped in a solution of the reagent by placing a cover glass with a droplet of a solution of m-phenylenediamine (sulfate or hydrochloride) in the middle over the reaction mixture, supported on the ring. After a few minutes, small crystals characteristic of the tribromo derivative of the test reagent will appear on the underside of the cover glass. Aniline works the same way.

Chloride (in the absence of fluoride)

Chromyl Chloride Test: (Schaeffer) This is an indirect test for chloride ions. Add potassium dichromate to the test solution and evaporate to dryness. Then add a small drop of concentrated sulfuric Acid to the sample. Set up a hanging drop of water to catch any gas evolved by the reaction. Allow to stand for a few minutes, and retrieve the hanging drop. Evaporate it to dryness and add a small drop of water to the dry residue. Add a small crystal of silver nitrate. A red precipitate of silver chromate establishes the presence of chloride ions in the sample. For this test to work, fluoride ions must not be present. Bromide and iodide ions do not interfere. The reason this test works to identify chloride is that the mixture of sample containing chloride mixed with potassium dichromate and treated with sulfuric acid produces chromyl chloride which is trapped by the hanging drop, where it decomposes to chromic acid and hydrochloric acid. Evaporation to dryness leaves only the chromic acid which reacts with the silver nitrate to produce red silver chromate. if there are no chloride ions in the original sample, then there will be no chromic acid, and thus no silver chromate.



Chromate or Dichromate

Silver Nitrate Test: (Chamot and Mason) (Schaeffer) Acidify the sample with nitric acid. Place a drop of 2 percent silver nitrate solution close by and combine the two drops. Dark red crystals of silver chromate form if the sample contains chromate or dichromate ions.

Lead Acetate Test: (Smith) The same procedure using lead acetate instead of silver nitrate produces a yellow precipitate in the presence of chromate or dichromate.

Fluoride

Sulfuric acid and silica Test: (Schaeffer) This is another test that involves the hanging drop technique to catch the reaction product in a drop of water. The sample is mixed with some pulverized silica sand and a drop of concentrated sulfuric acid is added. Silicon tetrafluoride gas is evolved if a fluoride is present. This is collected in a hanging drop of water where the silicon tetrafluoride breaks down into silicic acid and fluosilicic acid. The silicic acid is insoluble and forms a precipitate, while the fluosilicic acid remains in solution. It can be detected by converting it into its insoluble sodium salt by the addition of a few crystals of sodium chloride. Compare the fluosilicic acid test for sodium.

Halides (other than Fluoride)

Silver Nitrate Test: (Schaeffer) The presence of a precipitate when the sample is combined with a drop of silver nitrate solution indicates chloride, bromide or iodide.

Iodide

Potassium nitrite and starch Test: (Schaeffer) Potassium (or sodium) nitrite is an oxidizing agent that releases free iodine from a mixture containing the iodide ion. Put a few crystals of potassium nitrite in the test drop together with a few grains of starch. The starch grains will turn blue if iodine is present. A drop of hydrogen peroxide to which a little hydrochloric acid has been added can be used instead of the potassium nitrite. Bleach also works well for this test. As confirmation, adding potassium or sodium sulfite reduces the iodine back to iodide, causing the blue color to disappear.

Phosphate or Arsenate

Ammonium Molybdate Test: (Smith) Add a drop of ammonium molybdate reagent, and a 1 drop of concentrated nitric acid. Warm the slide. A yellow precipitate indicates phosphate or arsenate.

Selenium

Hydroquinone Test: (Chamot and Mason) Hydroquinone is a reducing agent that works well to detect the presence of selenium and tellurium. The sample, dissolved in nitric acid is placed on a slide and evaporated to dryness without overheating. The spot is covered with sulfuric acid and heated until dense fumes of sulfur trioxide begin to come off. The slide is cooled and another drop of sulfuric acid is added, then separate the clear solution from any insoluble material. This can be problematic. Glass fiber microfilters work well, if you have them, otherwise decanting the drop is about the best that can be done. Let the drop settle for a long time and then decant very slowly. The clear drop is then heated again until sulfur trioxide fumes begin to come off. Let the drop cool, and then combine it in the usual way with a saturated drop of hydroquinone dissolved in sulfuric acid. Warm the slide gently. Selenium will separate as a brown or red precipitate. After a few minutes, decant the clear liquid from the selenium precipitate, and put the drop on a fresh slide. Heat the drop again until the dense fumes of sulfur trioxide begin to come off. Tellurium, if present will precipitate as black bundles and aggregates. Hot sulfuric acid is dangerous. Drops tend to spread out when hot and itís a little difficult to keep things together. For this reason, this test would probably work better carried out on a small watch glass, or something that has a shape that helps to keep the drop in one place.

Silicate

The same reaction that is used to detect fluoride can be used to test for silicon. The hanging drop setup is used. Concentrated sulfuric acid is added to a mixture of the unknown material with calcium fluoride. Any gas that is evolved is trapped in a hanging drop of water. After a few minutes, if silicon is a major component of the sample, a precipitate of silicic acid will be visible in the water drop. If the water is subsequently treated with a few crystals of sodium chloride, insoluble hexagonal crystals of sodium fluosilicate will confirm the presence of silicon. This test gives a false positive if carried out in the presence of glass. A small lead dish works well. A glass cover slip can be coated with a film of stopcock grease to prevent the hydrogen fluoride from attacking it, or a plastic cover slip can be used.

Sulfate

Barium Chloride Test: (Chamot and Mason) Use the normal procedure. A white precipitate indicates sulfate.

Tellurium

Hydroquinone Test: (Chamot and Mason) See Hydroquinone Test under Selenium.




ENHANCED SPOT TESTS


CHEMICALS

70% Isopropyl Alcohol
8 Hydroxyquinoline
Alizarin
Aluminon
Curcumin
Dimethylaminobenzyledene Rhodanine
Dimethylglyoxime
Sodium Sulfide
Quercetin
Rhodamine B

You don't need all of these, just the Isopropyl alcohol (available at any drug store) and one or two of the other reagents for making the spots visible. 8-hydroxyquinoline and alizarin are the best. They are both very versatile reagents. The colors that develop are often unique for a given ion, and some ions produce spots that show fluorescence under ultraviolet light. Curcumen and quercetin are also pretty good and they're a lot cheaper, since they're both available at health food stores. They also show fluorescence with certain ions.

There are a number of reagents that react to a wide variety of ions with colors that, in many cases, are diagnostic. They can be used as simple spot test reagents, but the chance for success can be increased by using the following method to spread the sample over a wider area and separate the different ions somewhat using a technique borrowed from chromatography.

Chromatography is a method that has evolved into one of the main technologies both for detection and for the separation of compounds that are mixed together. There are many variations on the method, but the one that seems most useful for the basement scientist uses paper as the support. There are a lot of possibilities for the mobile phase. The best solvent system for a given application has been the subject of a lot of research. Isopropyl alcohol is not optimum, but it has the virtue of being easily available and seems to work reasonably well. It's also fairly non-toxic, which is always an important consideration.

It may be stretching the definition a little to call this method chromatography. It's really just a spot test, with a slight enhancement borrowed from chromatography. Before applying the reagent that develops the color, the sample spot is caused to spread across a region of the paper by allowing the isopropyl alcohol to climb the length of the paper by capillarity. Since different ions in a mixture in the sample have different solubilities in the isopropyl alcohol (and differing affinities for the paper), they will move along the paper at different rates. A particular ion then can be recognized by how far along it moves, and by the color it develops with a sprayed on reagent. Spreading the test spot across a region of the paper overcomes many of the problems of interfering ions by moving each ion to a different part of the paper before adding the reagent to detect it. Resolving two ions that may be eclipsing one another can sometimes be accomplished by using a longer strip of paper, or by trying a different solvent. Arthur Ritchie's book "Chromatography in Geology" contains a lot of helpful information.

You'll need a a stock of chromatography paper. It comes in various sizes. Tests of the method were made with pieces 1x4 inches cut from larger sheets. Precondition the strips to be used by soaking them in 70 percent isopropyl alcohol (or whatever solvent youíre using) for several hours and then let them dry before use. Cleanliness is very important. Don't let fingerprints get on the paper. About a half inch from one end of the paper, place a spot of the sample solution made by dissolving a small crystal (5 to 20 milligrams) of your unknown mineral. The spot should be approximately a quarter inch in diameter. Let the spot dry. Prepare a jar that is large enough for the paper to stand in without touching the sides. Fabricate a way to hang the paper in it so that the bottom end of the paper is about a half inch above the bottom of the jar, and the sides of the paper do not touch anywhere else. Then put about a half inch of 70 percent isopropyl alcohol in the jar and hang the paper that you prepared with the spot down so that the end just dips into the alcohol. It must not go deeply enough that the spot is beneath the surface of the alcohol, or the test will not work. Cover the jar and watch as the alcohol rises up the paper. The front of the solvent should rise fairly evenly up the paper over a time of several minutes until it reaches somewhere near the top. The time required will depend on the kind of paper used. Some papers are very fast, others may take an hour or more. Don't let the solvent front reach all the way to the top. When it's ready, remove the paper and hang it up to dry. After it's dry, spray the paper with a developer made from one of the reagents described below. Let it dry again. The positions and colors that will be on the paper will depend on what ions were present in the spot that you applied, and the type of spray reagent used. The possibilities are many, and this is both the power and the weakness of the method. The interpretation of the result is strictly empirical. To determine whether a sample contains, for example, gold, compare it with a reference strip that was made with a sample known to contain gold. Ideally, the reference strip should contain about the same amount of gold that the unknown has, in order for them to look the same. Even when they don't look exactly the same though, the colors and the distance moved, expressed as a percentage of the distance moved by the solvent front, will be the same, or nearly so. This can be an extremely sensitive test. There is no right or wrong way to do it. The important thing is to keep the paper clean, and do enough of them that you develop a system that works for you.


There are a lot of chemicals that can work as developers. Below is a list of several, and the ions that they are sensitive to. Reagents can also be applied by dipping the paper in them, but spraying works better. There are some really cute little chromatography sprayers available on ebay from time to time.

8 Hydroxyquinoline 0.5 percent in ethyl alcohol Al, Ag, Au, Ba, Be, Bi, Ca, Cd, Co, Cr, Cs, Cu, Fe, Ga, Ge, Hf, Hg, In, K, La, Li, Mg, Mn, Mo, Nb, Ni, Pb, Pd, Pt, Rb, Sb, Sc, Sn, Sr, Ta, Th, Ti, Tl, U, V, W, Zn, Zr

Alizarin: A saturated solution in ethyl alcohol Al, As, Bi, Ce, Cr, Cs, Cu, Fe, Hg, In, Li, Mg, Mn, Pb, Sb, Ta, Th, Ti, V, W, Y, Zn, Zr

Aluminon: 0.1 percent in 1 percent ammonium acetate in water Ac, Ag, Al, Ba, Be, Ca, Ce, Cr, Cu, Eu, Ga, Ha, In, La, Li, Mg, Mn, Nd, Ni, Ti

Curcumin: 0.1 percent in ethyl alcohol Ag, Al, Au, B, Be, Cr, Cu, Fe, Li, Ni, Pt, Ta, Ti, V, W, Zr

Dimethylaminobenzyledene Rhodanine 1 percent in ethyl alcohol Al, Au, Ag, Co, Cu, Fe, Hg, Li, Mn, Ni, Pb, Pd, Pt, Ta, Ti, V, W, Zn

Dimethylglyoxime 1 percent in ethyl alcohol Al, Cu, Fe, Co, Li, Ni, V, W, Zn

Sodium Sulfide 0.5 percent in water Au, Cd, Co, Cu, Fe, Hg, Mn, Ni, Pb, Pd, Pt, Zn

Quercetin 0.2 percent in ethyl alcohol Ag, Al, Bi, Ca, Cd, Co, Cr, Cu, Fe, Hg, Mg, Mn, Ni, Pb, Sb, Sn, U, Zn

Rhodamine B 0.1 percent in ethyl alcohol Ag, Au, Cu, Fe, Ni, Pt, Sb, V, W


After development, exposing the paper to ammonia fumes will sometimes enhance the picture and sometimes not. Also, viewing them under ultraviolet light can reveal features that otherwise are not visible.




BIBLIOGRAPHY

1. "Identification and Qualitative Chemical Analysis of Minerals" by Orsino C. Smith 1953
2. "Microscopy for Chemists" by Harold F. Schaeffer 1953
3. "Chromatography in Geology" by Arthur S. Ritchie 1964
4. ìHandbook of Chemical Microscopyî Chamot and Mason 1940
5. ìHandbook of Chemistryî (9th Edition) Norbert Adoplph Lange Ph.D. 1956
6. ìOrganic Analytical Reagentsî (Volume 2) Frank J. Welcher Ph.D. 1947
7. ìThe Merck Indexî (11th Edition) Susan Budavari Editor 1989

EXPERIMENTAL

SAMPLE: Approx 20 mg sample of heulandite. PROCEDURE: Pulverized sample in a small mortar. Added 1 drop of concentrated nitric acid. Continued grinding for a minute or two. Added 1 drop of water. Grind some more and touch the pestle to a microscope slide, leaving a small drop with quite a bit of undissolved material suspended in it. Placed a small drop of 0.1 percent aqueous aluminon beside it. Used a glass rod to cause the two drops to touch. Over a period of several minutes the point where the two drops meet developed a pink color that spread across to eventually cover the entire reagent drop. The color persists after adding a drop of ammonia. CONCLUSION: This test is positive for aluminum. DISCUSSION: Heulandite is approximately 9 percent aluminum. This experiment demonstrates the sensitivity of aluminon as a reagent for the detection of aluminum.

SAMPLE: Approx 25 mg crystal of heulandite. PROCEDURE: Ground the sample with about twice the volume of calcium fluoride. Placed in a small lead dish with a drop of sulfuric acid. Placed a cover slip with a hanging drop of water over the sample, supported on a plastic ring about 2 mm thick. The hanging drop spread out and ran under the plastic ring, but did not run down into the acid solution because of the hydrophobic nature of the plastic ring. Over several minutes gas bubbles evolved from the sulfuric acid and a residue accumulated on the cover slip in a ring around the inner edge of the plastic ring. It appears to be salicic acid. CONCLUSION: This test demonstrates a positive test for a silicate mineral.

SAMPLE: Approximately 50 mg piece of phosphate rock containing mostly strengite with some rockbridgeite. PROCEDURE: Pulverized the sample with a drop of concentrated nitric acid. After a minute or so, added a drop of water. Grind more, and allow to stand for a few minutes. Touch pestle to a slide, leaving a small drop with a little solid material. Placed a drop of 1 percent ammonium molybdate nearby and let the two drops flow together. There was no reaction immediately. Placed the slide on a low heat source (a small transformer). After a few minutes the drop, still wet, shows a border of yellow material visible against a white paper background. CONCLUSION: Test is positive for phosphate.

SAMPLE: Approximately 50 mg piece of phosphate rock containing mostly strengite with some rockbridgeite. PROCEDURE: Pulverized the sample with a drop of concentrated nitric acid. After a minute or so, added a drop of water. Grind more, and allow to stand for a few minutes. Sample stands for 5 to 10 minutes. Touch pestle to a slide, leaving a small drop with a little solid material. Placed a drop of 9 percent ammonium molybdate nearby and let the two drops flow together. A yellow color develops immediately at the interface between the two. Put the slide on a low heat source (a small transformer). A yellow crystalline mass develops as the drops proceed to dryness. CONCLUSION: Test is positive for phosphate. DISCUSSION: The stronger ammonium molybdate reagent develops a stronger yellow residue, as expected, however, it was not difficult to conclude that the test was positive, even with the 1 percent reagent.

SAMPLE: Blank. PROCEDURE: Placed about 50 mg of calcium fluoride in a small lead dish with a drop of concentrated sulfuric acid. A cover slip with a hanging drop of water was placed over the sample. No bubbles were observed coming from the sulfuric acid. The hanging drop spread out and ran under the edges, but did not run down into the acid. Over time the cover glass appeared to have a white film on the underside where the water drop was. The cover slip was removed and inverted on a piece of black paper. There appears to be a residue of salicic acid where the water drop was. The amount is much less that it was in a similar experiment in which the sample contained heulandite. CONCLUSION: This test could be interpreted as a false positive for silicate in a sample. It is significant that there was no visible evolution of gas bubbles from the acid. When silicate is present, Silicon tetrafluoride gas bubbles are distinctly visible, and probably should be part of the criterion for a positive result. The salicic acid on the cover slip in this case was probably from the hydrofluoric acid attacking the glass of the cover slip. A plastic cover slip might work better for this test or a glass one covered with a thin film of stopcock grease.

SAMPLE: Approximately 25 mg piece of aurorite. Procedure: The sample was ground together with a small drop of concentrated nitric acid for several minutes. The sample dissolved almost completely. The reagent solution was prepared by dissolving an aspirin tablet in a milliliter of ammonia solution and adding a half milliliter of strong hydrogen peroxide. A small drop of the sample solution was transferred to a microscope slide by touching the pestle to the slide. A reagent drop of similar size was placed nearby using a small glass rod, and the drop was carefully moved until it just barely touched the sample drop. A dense light gray precipitate developed immediately at the interface between the two drops and faint but clearly visible red streamers developed along the reagent side of the precipitate. The streamers slowly spread and intensified over time lending a pink cast to the solution. After a few minutes the precipitate turned a dirty brown. CONCLUSION: Test is positive for manganese. The red color was weak but unmistakable against a white paper background at first, but faded to a light brown as it spread.

SAMPLE: About 100 mg piece of aurorite. PROCEDURE: Sample is ground with a drop of nitric acid, and the solution is diluted with a drop of distilled water. A small drop is placed on the slide near a drop of a solution of oxalic acid, made by mixing a few crystals with a drop of water on the slide. The oxalic acid solution is then connected to the sample solution using a glass rod. A white precipitate develops immediately at the interface between the drops. CONCLUSION: The test is positive for calcium. DISCUSSION: The results might be due to the presence of strontium or barium. PROCEDURE: Another drop of sample solution was placed on a fresh slide, and tested with sodium chloride solution. The solution remained clear of any precipitate. CONCLUSION: Negative for silver, mercury and lead. DISCUSSION: Silver sometimes occurs in aurorite, but is absent from this specimen.

PROCEDURE: Another drop of test solution was placed on a slide and caused to join with a small drop of ammonium thiocyanate solution. No color was observed. CONCLUSION: Negative for iron. DISCUSSION: At this point, it has been established that the sample of aurorite probably contains manganese and calcium, and that it does not contain significant amounts of silver or iron.

SAMPLE: Approx 50 mg of picotite. PROCEDURE: The steps taken were similar to the ones described above. The sample was macerated in a drop of nitric acid and a drop placed on a microscope slide. A similar drop of acetylsalicylic acid reagent was placed nearby and carefully encouraged to join. There was a heavy white precipitate at the point where the two drops joined, and a yellow hue developed in the reagent drop over the following minute or so. CONCLUSION: Negative for manganese, positive for iron. DISCUSSION The absence of a red color is consistent with the fact that picotite does not contain manganese. This yellow color appears to be due to iron in the sample.

SAMPLE: Approx 50 mg of picotite. PROCEDURE: The sample was ground together with a drop of nitric acid for one minute. A small drop was transferred to a microscope slide. Another small drop of aluminon reagent was placed next to it, and the two drops caused to touch. A red streamer reached immediately into the reagent drop. Color persists after adding a drop of ammonia. CONCLUSION: Test is positive for aluminum.

SAMPLE: Approx 50 mg of picotite. PROCEDURE: The sample was ground together with a drop of nitric acid and allowed to stand for some time. A small drop of the nitric acid solution is placed on a microscope slide and a small lump of thiourea is placed into it. A reddish brown color spreads out from the thiourea as it begins to dissolve. Over time, the red color fades. No crystals are observed to form. CONCLUSION: Test is negative for lead. DISCUSSION: This test can also be taken to imply that thallium is absent, since it is also known to form crystals with thiourea.

SAMPLE: A drop of solution known to contain zinc ions. PROCEDURE: A drop of sample solution was placed on a slide and joined to a nearby drop of saturated potassium mercuric thiocyanate solution. Within a minute, small crosses were beginning to become visible. These grew and developed a fine branching structure, taking on a feathery appearance, eventually becoming like round fluffy snowflakes. CONCLUSION: Test is positive for zinc. DISCUSSION: The published description for this test is confirmed in this experiment. This test is also sensitive to cadmium, and produces crystals of a different form in the presence of cadmium, or a cadmium and zinc mixtures.


SAMPLE: A drop of solution made from dissolving a crystal of bertrandite in dilute hydrochloric acid. PROCEDURE: A drop of potassium oxalate solution was made on the slide by the following method. A small quantity of oxalic acid was placed on the slide and several small drops of potassium hydroxide solution were mixed into it and stirred. After a few minutes, most of the oxalic acid dissolved. The pH of the drop was measured to be about 6 to 7. A drop of the solution was decanted from the solid oxalic acid remaining and this was caused to join the drop of sample. Then a glass rod was touched to a solution of mercuric nitrate, and most of the adhering drop of mercury ions was transferred from the glass rod to an unused part of the slide, so that very little of the mercury solution remained on the glass rod. Then the glass rod was lightly touched to the point where the sample drop joined the potassium oxalate drop, and removed without stirring the drop. Over the next minute or two, crystals separated from the solution, mostly rhombs, and some prisms. The slide was placed under a polarizing microscope, and the extinction angle of several crystals was measured. Some of the crystals measured around 44 degrees, and some measured 38 degrees. CONCLUSION: The test is positive for beryllium, based on the presence of crystals with a measured extinction angle of 38 degrees. DISCUSSION: It is apparent that not all of the crystals are the double oxalate salt of potassium and beryllium, but some of the observed crystals fit the criterion for this salt, and therefore the conclusion that beryllium is present is supported. This test is a little difficult, and the first attempt to perform it failed. Making the potassium oxalate ìin situî in the manner described is probably not the best way to perform the test, but was necessitated by the fact that no potassium oxalate was available.







 

SOLUBILITIES OF SOME COMMON MINERALS

Information for the following table was taken, for the most part, from Orsino Smith’s book “Identification and Qualitative Analysis of Minerals.” In Smith’s tables, minerals soluble in hydrochloric acid were not tested for solubility in other acids, so subsequent entries for that mineral under nitric and sulfuric acids will indicate “No”, even though the mineral might in fact be soluble in those acids. Likewise, minerals not soluble in hydrochloric acid, if soluble in nitric acid, were not tested in sulfuric acid, and so may show an erroneous “No” under sulfuric acid. These errors are regrettable, but hopefully the table will nevertheless be useful if this caveat is kept in mind. Most, but not all, minerals that are soluble in hydrochloric acid are also soluble in nitric acid, and many are soluble in all three. Clearly, the reasoning for this order of presentation in Smith’s book was to bring the sample into solution in the mildest solvent possible, or perhaps the least dangerous. Using the sequence beginning with water and progressing through to sulfuric acid is a good method to use when experimenting to discover what acid will dissolve an unknown sample.

SOLUBILITY OF COMMON MINERALS

Name Composition Water HCl HNO3 H2SO4 Insoluble Note Acanthite Ag2S 0 0 1 0 0 Alabandite MnS 0 1 0 0 0 Albite NaAlSi3O8 0 0 0 0 1 Allanite (Ce Ca Y)2(Al Fe3+)3(SiO4)3(OH) 0 1 0 0 0 Silicate remains insoluble Alunite KAl3(SO4)2(OH)6 0 0 0 1 0 Amblygonite (Li Na)Al(PO4)(F OH) 0 0 0 0 1 Analcite NaAlSi2O6·(H2O) 0 1 0 0 0 Silicate remains insoluble Anatase TiO2 0 0 0 1 0 Soluble in Hot Sulfuric acid or potassium bisulfate fusion or alkalai hydroxides or carbonates Andalusite Al2SiO5 = Al[6]Al[5]OSiO4 0 0 0 0 1 Andradite Ca3Fe3+2(SiO4)3 0 1 0 0 0 Silicate remains insoluble Anglesite PbSO4 0 0 1 0 0 Anhydrite CaSO4 0 1 0 0 0 Ankerite Ca(Fe++ Mg Mn)(CO3)2 0 1 0 0 0 Evolves Carbon Dioxide Annabergite Ni3(AsO4)2-8(H2O) 0 1 0 0 0 Anthophyllite (Mg Fe)7Si8O22(OH)2 0 0 0 0 1 Antigorite (Mg Fe++)3Si2O5(OH)4 0 1 0 0 0 Silicate remains insoluble Antlerite Cu3SO4(OH)4 0 1 0 0 0 Apatite Ca5(PO4)3(OH F Cl) 0 1 0 0 0 Apophyllite KCa4(Si4O10)2F·8(H2O) 0 1 0 0 0 Silicate remains insoluble Aragonite CaCO3 0 1 0 0 0 Evolves Carbon Dioxide Arfvedsonite Na3(Fe Mg)4FeSi8O22(OH)2 0 0 0 0 1 Arsenopyrite FeAsS 0 0 1 0 0 Precipitates Sulfur Atacamite Cu2Cl(OH)3 0 1 0 0 0 Augite (Ca Na)(Mg Fe Al Ti)(Si Al)2O6 0 0 0 0 1 Aurichalcite (Zn Cu)5(CO3)2(OH)6 0 1 0 0 0 Evolves Carbon Dioxide Avicennite Tl2O3 0 1 0 0 0 Azurite Cu3(CO3)2(OH)2 0 1 0 0 0 Evolves Carbon Dioxide Barite BaSO4 0 0 0 1 0 Sol in Hot Sulfuric acid See Smith p. 148 Bertrandite Be4Si2O7(OH)2 0 0 0 0 1 Fuse with sodium carbonate filter. Beryllium is in the solid phase. Soluble in HCl. Beryl Be3Al2Si6O18 0 0 0 0 1 Biotite K(Mg Fe++)3[AlSi3O10(OH F)2 0 0 0 1 0 Silica separates out Boracite Mg3B7O13Cl 0 1 0 0 0 Bornite Cu5FeS4 0 0 1 0 0 Precipitates Sulfur Boulangerite Pb5Sb4S11 0 1 0 0 0 PbCl2 precipitates. Soluble hot. Sb separates on dilution Bournonite PbCuSb3 0 0 1 0 0 Brochantite Cu4(SO4)(OH)6 0 1 0 0 0 Calaverite AuTe2 0 0 0 1 0 Calcite CaCO3 0 1 0 0 0 Evolves Carbon Dioxide Carnotite K2(UO2)2(VO4)2- 1-3(H2O) 0 1 0 0 0 Cassiterite SnO2 0 0 0 0 1 Celestite SrSO4 0 0 0 0 1 Cerussite PbCO3 0 0 1 0 0 Evolves Carbon Dioxide Chabazite CaAl2Si4O12-6H2O 0 1 0 0 0 Silicate remains insoluble Chalcanthite CuSO4-5(H2O) 1 0 0 0 0 Chalcocite Cu2S 0 0 1 0 0 Evolves Nitric Oxide precipitates Sulfur Chalcopyrite CuFeS2 0 0 1 0 0 Precipitates Sulfur Chlorite 0 0 0 1 0 Silica separates out Chromite FeCr2O4 0 0 0 0 1 Chrysocolla (Cu Al)2H2Si2O5(OH)4·n(H2O) 0 1 0 0 0 Chrysotile Mg3Si2O5(OH)4 0 1 0 0 0 Silicate remains insoluble Cinnabar HgS 0 0 1 0 0 Clinochlore (Mg Fe++)5Al(Si3Al)O10(OH)8 0 0 0 1 0 Silica separates out Clinozoisite Ca2Al3(SiO4)3(OH) = Ca2AlAl2(SiO4)(Si2O7)O(OH) 0 1 0 0 0 Partially insoluble Cobaltite CoAsS 0 0 1 0 0 Precipitates Sulfur Colemanite Ca2B6O11·5(H2O) 0 1 0 0 0 Conichalcite CaCu(AsO4)(OH) 0 1 0 0 0 Copiapite Fe++Fe3+4(SO4)6(OH)2·20(H2O) 1 0 0 0 0 Copper Cu 0 0 1 0 0 Evolves Nitric Oxide Cordierite Mg2Al4Si5O18 0 1 0 0 0 Silicate remains insoluble Corundum Al2O3 0 0 0 0 1 Covellite CuS 0 0 1 0 0 Cryolite Na3AlF6 0 0 0 1 0 Cuprite Cu2O 0 1 0 0 0 Danburite CaB2(SiO4)2 0 0 0 0 1 Datolite CaBSiO4(OH) 0 1 0 0 0 Silicate remains insoluble Diaspore AlO(OH) 0 0 0 0 1 Diopside CaMgSi2 O6 0 0 0 0 1 Dolomite CaMg(CO3)2 0 1 0 0 0 Evolves Carbon Dioxide Dyscrasite Ag3Sb 0 0 1 0 0 Enargite Cu3AsS4 0 0 1 0 0 Enstatite MgSiO3 0 0 0 0 1 Epidote Ca2(Fe3+ Al)3(SiO4)3(OH)=Ca2(Fe Al)Al2(SiO4)(Si2O 0 1 0 0 0 Silicate remains insoluble Erythrite Co3(AsO4)2-8(H2O) 0 1 0 0 0 Ferberite Fe++WO4 0 0 0 0 1 Sodium Carbonate fusion dissolves tungsten in water Fluorapatite Ca5(PO4)3F 0 1 0 0 0 Fluorite CaF2 0 0 0 1 0 Forsterite Mg2SiO4 0 1 0 0 0 Silicate remains insoluble Franklinite (Zn Fe Mn)(Fe Mn)2O4 0 1 0 0 0 Gahnite ZnAl2O4 0 0 0 1 0 Galena PbS 0 1 1 0 0 PbCl2 separates but is sol in Hot water Glauberite Na2Ca(SO4)2 0 1 0 0 0 Glauconite (K Na)(Fe3+ Al Mg)2(Si Al)4O10(OH)2 0 0 0 0 1 Glaucophane Na2(Mg3Al2)Si8O22(OH)2 0 0 0 0 1 Gmelinite (Na2 Ca)Al2Si4O12·6(H2O) 0 1 0 0 0 Silicate remains insoluble Goethite Fe3+O(OH) 0 1 0 0 0 Gold Au 0 0 0 0 1 Graphite C 0 0 0 0 1 Gypsum CaSO4 0 1 0 0 0 Halite NaCl 1 0 0 0 0 Harmotome (Ba Na K)1-2(Si Al)8O16·6(H2O) 0 1 0 0 0 Hausmannite (Mn+2)(Mn+3)2O 4 0 1 0 0 0 Hematite Fe2O3 0 1 0 0 0 Heulandite Ca Na)2-3Al3(Al Si)2Si13O36·12(H2O) 0 1 0 0 0 Hornblende (Ca Na)2-3(Mg Fe Al)5(Al Si)8O22(OH F)2 0 0 0 0 1 Hydromagnesite Mg5(CO3)4(OH)2·4(H2O) 0 1 0 0 0 Evolves Carbon Dioxide Hydrozincite Zn5(CO3)2(OH)6 0 1 0 0 0 Evolves Carbon Dioxide Hypersthene (Mg Fe)SiO3 0 1 0 0 0 Ilmenite Fe++TiO3 0 1 0 0 0 Ilvaite CaFe++2Fe3+Si2O7O(OH) 0 1 0 0 0 Silicate remains insoluble Jamesonite Pb4FeSb6S14 0 1 0 0 0 PbCl2 precipitates. Soluble hot. Sb separates on dilution Jarosite KFe3(SO4)2(OH)6 0 1 0 0 0 Kaolinite Al2Si2O5(OH)4 0 0 0 0 1 Kernite Na2B4O6(OH)2·3(H2O) 1 0 0 0 0 Kyanite Al2SiO5 0 0 0 0 1 Latrappite (Ca Na)(Nb Ti Fe)O3 0 0 0 0 1 try Potassium Bisulfate fusion Laumontite CaAl2Si4O12·4(H2O) 0 1 0 0 0 Silicate remains insoluble Lazulite MgAl2(PO4)2(OH)2 0 0 0 0 1 Lepidolite K(Li Al)3(Si Al)4O10(F OH)2 0 1 0 0 0 Incomplete dissolution. Silicate remains insoluble Leucite KAlSi2O6 0 1 0 0 0 Limonite Mixed Iron Oxides 0 1 0 0 0 Magnesite MgCO3 0 1 0 0 0 Evolves Carbon Dioxide Magnetite Fe3O4 0 1 0 0 0 Malachite Cu2CO3)OH)2 0 1 0 0 0 Evolves Carbon Dioxide Manganite MnO(OH) 0 1 0 0 0 Marcasite FeS2 0 0 1 0 0 Precipitates Sulfur Melanterite Fe++SO4·7(H2O) 1 0 0 0 0 Mercury Hg 0 0 1 0 0 Evolves Nitric Oxide Microcline KAlSi3O8 0 0 0 0 1 Microlite (Ca Na)2Ta2O6(O OH F) 0 0 0 1 0 Mimetite Pb5(AsO4)3Cl 0 0 1 0 0 Molybdenite MoS2 0 0 1 0 0 Monazite (Ce La Nd Th)PO4 0 1 0 0 0 Muscovite KAl2(Si3Al)O10(OH F)2 0 0 0 0 1 Natrolite Na2Al2Si3O10-2H2O 0 1 0 0 0 Silicate remains insoluble Olivenite Cu2AsO4(OH) 0 0 1 0 0 Olivine (Mg Fe)2SiO4 0 1 0 0 0 Silicate remains insoluble Opal SiO2·n(H2O) 0 0 0 0 1 Orpiment As2S3 0 0 1 0 0 Precipitates Sulfur Orthoclase KAlSi3O8 0 0 0 0 1 Pectolite NaCa2Si3O8(OH) 0 1 0 0 0 Silicate remains insoluble Perovskite CaTiO3 0 0 0 1 0 Silica separates out Phillipsite (K Na Ca)1-2(Si Al)8O16·6(H2O) 0 1 0 0 0 Silicate remains insoluble Phlogopite KMg3(Si3Al)O10(F OH)2 0 0 0 1 0 Silica separates out Phosgenite Pb2(CO3)Cl2 0 0 1 0 0 Evolves Carbon Dioxide Polybasite (Ag Cu)16Sb2S11 0 0 1 0 0 Prehnite Ca2 Al2 Si3 O10(OH)2 0 1 0 0 0 Proustite Ag3AsS3 0 0 1 0 0 Pyrargyrite Ag3SbS3 0 0 1 0 0 Pyrite FeS2 0 0 1 0 0 Precipitates Sulfur Pyrolusite MnO2 0 1 0 0 0 Pyromorphite Pb5(PO4)3Cl 0 0 1 0 0 Pyrophyllite AlSi2O5OH 0 0 0 1 0 Silica separates out Pyroxene (Mg Fe)SiO3 0 0 0 0 1 Pyrrhotite FeS 0 1 0 0 0 Quartz SiO2 0 0 0 0 1 Realgar As4S4 0 0 1 0 0 Precipitates Sulfur Rhodochrosite MnCO3 0 1 0 0 0 Evolves Carbon Dioxide Rhodonite (Mn++ Fe++ Mg Ca)SiO3 0 1 0 0 0 Rutile TiO2 0 0 0 0 1 Scheelite Ca(WO4) 0 1 0 0 0 Yellow Tungstate precipitates Scolecite CaAl2Si3O10·3(H2O) 0 1 0 0 0 Silicate remains insoluble Scorodite Fe3+AsO4·2(H2O) 0 1 0 0 0 Serpentine (Mg Fe)3Si2O5(OH)4 0 1 0 0 0 Silicate remains insoluble Siderite Fe++CO3 0 1 0 0 0 Evolves Carbon Dioxide Silver Ag 0 0 1 0 0 Evolves Nitric Oxide Smithsonite ZnCO3 0 1 0 0 0 Evolves Carbon Dioxide Sodalite Na4Al3(SiO4)3Cl 0 1 0 0 0 Sphalerite (Zn Fe)S 0 1 0 0 0 Spinel MgAl2O4 0 0 0 1 0 Spodumene LiAlSi2O6 0 0 0 0 1 Stannite Cu2FeSnS4 0 0 1 0 0 SnO2 separates out Stephanite Ag5SbS4 0 0 1 0 0 Stibnite Sb2S3 0 1 1 0 0 Dilution of Nitric acid solution may precipitate sulfur Stilbite NaCa4[Al8Si28O72]·n(H2O) (n=28-32) 0 1 0 0 0 Strontianite SrCO3 0 1 0 0 0 Evolves Carbon Dioxide Sulfur S 0 0 0 0 1 Sylvanite AuAgTe4 0 0 1 0 0 Gold separates out Talc Mg3Si4O10(OH)2 0 0 0 0 1 Tetrahedrite (Cu Fe)12Sb4S13 0 0 1 0 0 Thenardite Na2SO4 1 0 0 0 0 Thomsonite NaCa2Al5Si5O20·6(H2O) 0 1 0 0 0 Silicate remains insoluble Titanite CaTiSiO5 0 1 0 0 0 Silicate remains insoluble Topaz Al2SiO4(F OH)3 0 0 0 1 0 Silica separates out Tourmaline Variable Composition 0 0 0 0 1 Tremolite Ca2Mg5Si8O22(OH)2 0 0 0 0 1 Turquoise CuAl6(PO4)4(OH)8*5(H2O) 0 1 0 0 0 Ulexite NaCaB5O6(OH)6·5(H2O) 1 0 0 0 0 Uraninite UO2 0 0 1 0 0 Vanadinite Pb5(VO4)3Cl 0 1 1 0 0 PbCl2 precipitates. Soluble hot. Vesuvianite Ca10Mg2Al4(SiO4)5(Si2O7)2(OH)4 0 1 0 0 0 Silicate remains insoluble Vivianite Fe3(PO4)2-(H2O)8 0 1 0 0 0 Wavellite Al3(PO4)2(OH)3-5(H2O) 0 1 0 0 0 Weloganite Sr3Na2Zr(CO3)6·3(H2O) 0 1 0 0 0 Strontium sodium zirconium Carbonate… should evolve CO2 Wernerite Na4 [Cl|(AlSi3O8)3 ] - Ca4 [CO3|(Al2Si2O8)3] 0 1 0 0 0 Silicate remains insoluble Willemite Zn2SiO4 0 1 0 0 0 Wollastonite CaSiO3 0 1 0 0 0 Silicate remains insoluble Wulfenite PbMoO4 0 1 0 0 0 PbCl2 separates but is sol in Hot water Wurtzite (Zn Fe)S 0 1 0 0 0 Zincite ZnO 0 1 0 0 0 Zinkenite Pb9Sb22S42 0 1 0 0 0 PbCl2 precipitates. Soluble hot. Sb separates on dilution Zircon ZrSiO4 0 0 0 1 0 Silica separates out Zoisite Ca2Al3(SiO4)3(OH) 0 0 0 0 1

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