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May 5 2012 7 05 /05 /May /2012 12:38

I have compiled a list of sources for the elements that are available to the amateur chemist. Osmium will be discussed here. 

 

Osmium is an interesting element. It is named after osme, which means stink. This refers to the tendency of powdered osmium to oxidize in air to form colorless, volatile, smelly, and highly toxic osmium tetroxide. Osmium itself is the densest element on the periodic table and one of the few colorful metals; it is light bluish, shiny, and brittle. Osmium is quite inert in bulk form but in powdered form it is quite reactive. It was discovered from platinum ores in the 1800s.

 

In element form: Some old photographic needles were supposedly made of osmium. Fountain pen nibs were sometimes made of osmiridium ore.

 

In compound form: No sources found.

 

I do not have any osmium at the time of this writing.

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May 4 2012 6 04 /05 /May /2012 12:56

I have compiled a list of sources for the elements that are available to the amateur chemist. Rhenium will be discussed here.

 

Rhenium was the last non-radioactive element discovered. The metal itself is extremely dense and lustrous. Rhenium has the highest boiling point of any metal. Rhenium forms perrhenate compounds, which are the most commonly traded form of rhenium. These are much weaker oxidizing agents than the permanganates. Rhenium metal is made by heating the ammonium perrhenate. The rhenium powder is then sintered to form the desirable shape. Rhenium is one of the rarest elements in the earth's crust, with an average concentration of 1 milligram per metric ton in the earth's crust. Rhenium forms a wide variety of colorful and not well-defined compounds in many oxidation states.

 

In element form: Some thermocouples have tungsten/rhenium alloy in them. Antique GE rhenium flashbulbs have a tungsten/rhenium alloy filament used to trigger the flash.

 

In compound form: No sources found.

 

I do not have any rhenium collected at the time of this writing.

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May 3 2012 5 03 /05 /May /2012 17:43

I have compiled a list of sources for the elements that are available to the amateur chemist. Tungsten will be discussed here.

 

Tungsten comes from the name "tung sten", which means "heavy stone" in Swedish. Tungsten, true to its name, is a dense, gray metal of Group 6. Tungsten, when pure, is quite flexible, but becomes brittle from impurities. It is quite inert at room temperature to most acids and alkalis. However, it burns when heated, forming the yellow-green compound tungsten trioxide. This is why the filament in an incandescent bulb must be protected from oxygen by filling the bulb with argon and nitrogen. Tungsten forms colorless tungstates, which precipitate the hydrated trioxide upon acidification. Tungsten is found as calcium tungstate, which is acified to precipitate tungsten trioxide and reduced with hydrogen to form tungsten powder. This is sintered into various shapes. Tungsten hexafluoride, a highly corrosive substance, is the heaviest known gas.

 

In element form: Open an incandescent or halogen bulb (the higher wattage the better) and remove the tungsten filament. Tungsten is used in rod form in TIG (tungsten inert gas) welding.

 

In compound form: When a tungsten filament is burned in air, it produces yellowish tungsten trioxide. Tungsten carbide is both used in jewelry and as an abrasive for blades and grout cutters.

 

Here is my sample of tungsten. It is a filament from a 500 watt halogen lamp. This is a significant amount of tungsten.

 

Tungsten-ring.JPG

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May 2 2012 4 02 /05 /May /2012 13:26

I have compiled a list of sources for the elements that are available to the amateur chemist. Tantalum will be discussed here.

 

Tantalum is an extremely dense, bluish-gray metal. It is named after the king Tantalus, who was submerged in water he could not ever drink and was near fruit he could never eat. Tantalum's behavior of being submerged in acids but never able to dissolve was the reason that it was given the name. Tantalum is a rare and expensive metal, found in niobium ores. It was mistakenly confused with niobium, despite its much higher density, and is difficult to remove from the aforementioned metal. When tantalum is pure, it has excellent physical strength and properties, making it ideal for artificial limbs and fine medical instruments. Tantalum, despite its seeming invincibility, has its weak points; it can be dissolved in hydrofluoric acid or molten hydroxides. Tantalum is primarily pentavalent, which is part of the reason for its inertness. The oxide is a white unreactive substance, difficult to reduce to the metal. It can be made to form compounds to dissolving in molten metal hydroxides or electric arc fusion with other metals, forming a range of tantalates. 

 

In element form: High capacity (30 mfd or so) tantalum capacitors have a lump of anodized tantalum in the middle. Use hot hydrochloric acid to dissolve any manganese dioxide which may be present. Some expensive medical instruments use tantalum.

 

In compound form: Motion sensors may contain lithium tantalate as the sensing material.

 

Here is my sample of tantalum. It is a nugget of tantalum metal with tantalum connector wire, taken from a 22 mfd tantalum capacitor. The green anodized color is visible in certain places.

 

Tantalum-capacitor-core.JPG

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May 2 2012 4 02 /05 /May /2012 13:04
Hydrogen peroxide is a rather unstable material. It has a strong tendency to exothermically decompose to water and oxygen. This process is limited in drug store hydrogen peroxide by the addition of a stabilizer such as phenacetin. However, certain substances vastly accelerate the decomposition of hydrogen peroxide. One of them is blood, where the enzyme catalase breaks down hydrogen peroxide very quickly and efficiently. Other inorganic compounds break down hydrogen peroxide too.
 
Iodide: This is used in the famous "elephant's toothpaste" demonstration. 30% hydrogen peroxide is mixed with detergent and some potassium iodide is added, after which it vigorously foams up to 30 times its original size. This is not my video. Iodide's decomposition is quite fast, as is evident by these pictures.
 
Copper(II) chloride: This is another strong catalyst in the decomposition of hydrogen peroxide. The green solution visible at the end of this video decomposes the hydrogen peroxide quite rapidly.
Alkali and alkaline earth metal salts: These decompose hydrogen peroxide slowly. Of all of the salts that I tried the carbonates and bicarbonates caused the most decomposition, undoubtedly because of the basic solution produced.
 
Manganese dioxide: One of the most well - known catalysts for hydrogen peroxide, manganese dioxide vigorously and completely decomposes hydrogen peroxide, causing a dirty brownish-black solution to be left behind. The other lower manganese oxides also decompose hydrogen peroxide but to a lesser extent.
 
Cobalt(II) chloride: Cobalt(II) chloride, like most transition metal salts, decomposes hydrogen peroxide relatively fast, but not in a spectacular way.
 
Silver compounds: Soluble silver compounds are also excellent decomposition catalysts. When silver covered in a chloride crust is immersed in hydrogen peroxide, vigorous fizzing begins which only slows a little when acidified by acetic acid. Catalysis continues until all of the peroxide is depleted.
 
Other precious metals: Platinum sponge and other precious metal catalysts likely have a strong effect on peroxide decomposition, but I could not obtain any of these for experimentation.
 
Iron(II) compounds: These are used to make Fenton's reagent by reaction with hydrogen peroxide. While they destabilize hydrogen peroxide to some extent they are not as strong a catalyst as the other transition metal compounds.
 
Chromium(VI) compounds: These actually react with hydrogen peroxide, forming a metastable coordination complex. It soon decomposes, giving off oxygen and returning to the original state in some cases. If it does so (which it does more in basic solution), it is a catalyst, although a slow catalyst in comparison with normal ones.
 
 
Titanium compounds: These form a bright red and surprisingly stable peroxo complex with titanium.
Titanium-peroxo-complex.JPG
Most other metals only decompose hydrogen peroxide to a fair extent or are actually oxidized by it.
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May 1 2012 3 01 /05 /May /2012 13:15

I have compiled a list of sources for the elements that are available to the amateur chemist. Hafnium will be discussed here.

 

Hafnium is the first element after the long and boring series of rare earth metals are completed. (Whether you wish to purchase all of them is a matter of your decision, of course, but some of the very similar ones may be skipped.) Anyway, hafnium is a dense (more so than lead), silvery gray transition metal. It absorbs neutrons, making it a harmful impurity in zirconium nuclear fuel casings which must allow the radiation to escape. Therefore, many pains are taken to remove hafnium from zirconium. In most applications, however, most zirconium is mixed with hafnium and vice versa. Therefore, your cubic zirconia is probably also cubic hafnia. Hafnium forms rather low-melting tetravalent compounds; its chloride hydrolyzes in water to form an insoluble hafnium oxychloride, just like most tetravalent compounds. Hafnium is extremely corrosion resistant and does not dissolve in acids or bases. It can burn in air when finely powdered, though, just like titanium. Hafnium finds use in very high-melting alloys, such as those used in rocket nozzles.

 

In element form: Find a plasma cutter tip at a welding shop and cut off the copper end. You will see a small cylinder of hafnium metal in the center.

 

In compound form: Cubic zirconia undoubtedly contains some hafnium, as well as most other sources of zirconium.

 

Here is my sample of hafnium. It is a plasma cutter tip tip.

 

Hafnium.JPG

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April 30 2012 2 30 /04 /April /2012 13:12

Most metal carbonates decompose when heated. Their decomposition products vary between metals.

 

Sodium and other alkali metal carbonates: I heated sodium carbonate to a red heat and it did not decompose. These carbonates are extremely stable and do not decompose at an ordinary temperature.

 

Calcium and other alkaline earth metal carbonates (except beryllium): When heated strongly (840 C for calcium), magnesium, calcium, and strontium carbonates decompose into the oxide. The production of lime (calcium oxide) from limestone (calcium carbonate) has been a well known process for many years. Barium carbonate may form the peroxide when heated strongly. Beryllium does not form a carbonate.

 

Trivalent carbonates: Aluminium and its neighbors do not form carbonates. Scandium and the rare earth metals form carbonates. When heated vigorously, they decompose to the oxide. 

 

Tetravalent or higher carbonates: The carbonates of Group 4, 5, and 7 do not exist.

 

Cobalt(II) carbonate: This easily decomposes to black  cobalt(II,III) oxide, glowing red-hot as it is calcined. The resulting product is slightly soluble in hydrochloric acid, releasing no significant quantity of chlorine gas. Upon evaporation, a tiny crop of cobalt(II) chloride crystals is visible.

 

DSCF0295.JPG

 

Copper(II) carbonate: This easily decomposes to black copper(II) oxide. Calcining does not occur at the low decomposition temperature, making the resulting oxide quite soluble in hydrochloric acid and useful for thermites and other metal oxide reactions. 

 

Copper-II--oxide.JPG

 

Zinc carbonate: This white substance decomposes to white zinc oxide, which reversibly turns light yellow when heated. The change in color is due to a loss of a few oxygen atoms.

 

Manganese carbonate: White manganese(II) carbonate decomposes when heated in air to a higher oxide of manganese such as manganese(III) oxide or manganosic oxide. This oxide forms a mud brown suspension in hydrochloric acid which slowly gives off chlorine gas at room temperature to form manganese(II) chloride. A little bit of the white carbonate is still visible in the picture below. In a vacuum, manganese(II) carbonate decomposes to green manganese(II) oxide.

 

Manganese-II-III--oxide.JPG

 

Iron(II) carbonate: White iron(II) carbonate decomposes without heating in air to iron(III) oxide and carbon dioxide gas. Since it does not have any excess acid like Mohr's salt, it is very sensitive to aerial oxidation. Addition of hydrochloric acid to a completely dry (and orange brown) precipitate of "iron(II) carbonate" produces no carbon dioxide gas, showing that a chemical reaction has occurred. 

 

Iron-III--oxide--2-.JPG

 

Nickel(II) carbonate: Nickel(II) carbonate, a green solid, decomposes at a moderate temperature to green stoichiometric nickel(II) oxide, which is soluble in HCl. When heated strongly (600 C), it turns black as it oxidizes further to non-stoichiometric nickel oxide. This nickel oxide has a formula (this is just armchair speculation based on experimental results) of NiO1+x where x is about 0.3. When heated strongly, the surface of the nickel oxide particles become further oxidized, reaching a formula of either Ni2O3 or NiO2. When placed in hydrochloric acid, the surface instantly dissolves, producing chlorine gas much more vigorously than with manganese dioxide. The inside portion dissolves slowly in hydrochloric acid, just like an ordinary calcined nickel oxide. Here is the black nickel oxide.

 

DSCF0304.JPG

 

Lead(II) carbonate: Lead carbonate decomposes to lead(II) oxide, which is light orange - brown. The white lead carbonate is visible beside the orange lead oxide. 

 

DSCF0359.JPG

 

Silver(I) carbonate: This light yellow compound decomposes to dark brown silver(I) oxide at 210 C, which ultimately decomposes to black silver metal at 280 C.  

 

Gadolinium(III) carbonate: This pale yellow solid decomposes to white gadolinium(III) oxide when heated.

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April 30 2012 2 30 /04 /April /2012 13:03

I have compiled a list of sources for the elements that are available to the amateur chemist. Thulium, ytterbium, and

lutetium will be discussed here. 

 

Thulium is a typical rare earth metal. Its compounds are pale green to yellow. No sources of it are available to the general public.

 

Ytterbium is another rare earth metal. Its liquid range from 824-1196 C is the smallest of all metals. Most ytterbium compounds are white and boring. It does form divalent compounds, however, of which the chloride is green. Earthquake meters may contain ytterbium metal to measure the quake, but it is unlikely that they do.

 

Lutetium is the last of the rare earth metals. It is the rarest and most expensive one. Its compounds are boring and white. Lutetium tantalate is the densest white substance besides thorium dioxide, which is slightly radioactive. There  are no household sources.

 

Here ends the sequence of little articles about rare earth metals.

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April 28 2012 7 28 /04 /April /2012 13:18

I have compiled a list of sources for the elements that are available to the amateur chemist. Erbium will be discussed here.

 

Erbium is a typical not-so-reactive rare earth metal. Erbium forms pink compounds that fluoresce in UV light. It is solely trivalent. These heavy rare earth metals are quite boring as they are so similar to each other.

 

In element form: No sources found.

 

In compound form: Some glass is colored pink by the addition of erbium oxide.

 

I do not have any erbium at the time of this writing.

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April 27 2012 6 27 /04 /April /2012 13:20

Chlorine is a light yellow gas, second from the top in the halogen group (Group 17) on the periodic table. Although it is highly toxic (it was the first poison gas used in World War I), its extremely pungent smell is fair warning. Breathing in small amounts of chlorine is not bad, but larger amounts cause burning of the mucus membranes and lungs.

 

Bromine, on the other hand, is a reddish liquid that smells similarly, although some say that the smell is less repulsive. Bromine has the additional hazard of being spillable, and liquid bromine evaporates twice as fast as water AFAIK.

 

Both of these poisonous elements have one property that enables them to be easily, quickly, and completely neutralized. They are strong oxidizing agents. The chloride and bromide ions are harmless, ionic substances in most cases. Therefore, reduction can be used to remove chlorine gas.

 

For example, if you want to produce some pure manganese(II) solution by dissolving some pure manganese dioxide from a tantalum capacitor in hydrochloric acid, you will produce a large amount of chlorine gas. However, the reaction needs to be heated, so placing it outside on a normal day will slow the reaction very considerably. To prevent gassing yourself in chlorine or bromine gas, you may use a variety of reducing agents. Sodium or potassium metabisulfite, readily available from brewer's shops or chemical supply stores, is a good reducing agent. Mix this with sodium carbonate solution to form a basic reducing solution. Chlorine dissolves easily in a basic solution, forming hypochlorite and chloride. The hypochlorite is easily reduced by the sulfite ions. For a beaker, soak a rag in this solution and place it on top of the beaker. Use the bottom of another beaker if needed to hold the rag down, making sure that all of the chlorine contacts the rag. For a test tube or flask, use a scrap piece of hose (chlorine tends to damage hoses) and lead to a beaker with carbonate-sulfite solution. For a vial, wrap a tissue soaked in the solution around the cap to absorb any produced gases.

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