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April 13 2012 6 13 /04 /April /2012 14:20

3-20 - Use Control + F to find what you want in here

Beryllium reactions: Beryllium is placed in copper(II) chloride crystals and wet with a few drops of water. The reaction is extremely violent, so the solution is diluted. The reaction remains very violent. Large quantities of copper   are produced and a dirty solution is left behind. This solution is neutralized with ammonia to form an extremely gelatinous precipitate containing some occluded copper. Here are the initial reactants, the reaction (steam obscures the view), the precipitate under flash lighting and fluorescent lighting, and the copper residue.


Bismuth dissolution: Bismuth is dissolving in a 2:1 mixture of hydrogen peroxide and hydrochloric acid. The resulting bismuth piece lacks the original iridescent oxide coloration and is pitted and corroded. Addition of an excess of ascorbic acid followed by tincture of iodine formed no precipitate, only a light yellow solution. Therefore, it must be determined that I have failed to produce a precipitate of bismuth iodide, just like I did with tin iodide outside of my school lab. The resulting bismuth solution is diluted with water. Bismuth oxychloride precipitates. Left picture: Bismuth solution with bismuth metal in bottom. Right: Bismuth(III) chloride and water, hydrolyzed.

Iodine disproportionation and reactions: A piece of calcium is placed in tincture of iodine. The calcium vigorously dissolves, producing a large amount of heat. As the solution gets cloudy from the calcium hydroxide, the red-brown coloration of the triiodide ion disappears, replaced by the colorless iodide and iodate: 6 OH- + 3 I3- à 3 H2O + IO3-+ 8 I- As evidenced by this reaction, this is a highly inefficient way of producing calcium iodate, but the reaction itself is interesting for experimental purposes. When acetic acid was added, the solution became a shade yellow again, showing that some of the reverse reaction occurred, though in very small amounts. Addition of sodium hypochlorite did not cause any crystallization of calcium iodate or periodate, probably because the solution is mostly all water.


Bismuth dissolution: The remaining bismuth chunk from the above experiment is redissolved in the same solution. I then took some of the bismuth solution and placed a piece of zinc in it. The zinc immediately turned black, and the black coating of bismuth began thickening. An air bubble trapped against the zinc created a hole in the black coating, showing how thick it is. Some hydrogen was produced, and some of the bismuth hydrolyzed in the solution. Another portion of the bismuth solution was reacted with the previously produced calcium iodide-iodate-hydroxide mixture. Residual hydrogen peroxide, as well as the acidity of the solution, formed orange-brown triiodide, but when reduced with ascorbic acid, a yellow complex remained. This yellow complex turned white when diluted with water, showing that it is a bismuth iodide complex. Pictures are 1-4: bismuth reduction by zinc, 5: turbid bismuth(III) chloride solution, 6: bismuth iodo complex.

Iodine reactions: Tincture of iodine is acidified with hydrochloric acid and reacted with bleach. Some iodine forms initially, which reacts with nascent chlorine from bleach + hydrochloric acid to form iodine monochloride solution, which is yellow. Sodium bicarbonate is added, and iodine precipitates, as well as sodium chloride (due to the high concentration of ions in the solution). The iodine is filtered and dried a little. Some of the specks of iodine were placed in a vial and then placed in hot tap water. The purple color of iodine vapor formed, barely visible at that temperature. Then the vial is placed in boiling water. Almost all of the tiny iodine fleck evaporates, leaving the impurities behind and producing a bright purple coloration. When cooled, tiny iodine crystals are deposited on the walls of the vial. 1: Iodine crystals 2: Iodine vapor at 40 C 3: Iodine, deposited crystals.


Tantalum cleaning: I placed my tantalum chunk in hydrochloric acid in a vial and placed the vial in a boiling water bath .Most all of the manganese dioxide left dissolved and the solution was green because of dissolved chlorine gas. It reeked when I poured it down the drain.

Bismuth hydrolysis: Bismuth oxy compounds at pH 2 or 3 are white, while those at higher pH are light yellow.


Bismuth oxide production: The bismuth oxychloride is filtered and dried.

Silver dissolution: Silver is anodized and then placed in a 1:1 mixture of 3% hydrogen peroxide and 5% acetic acid.

Bismuth oxide reactions: Bismuth oxide is heated on aluminium foil. It melts and then rapidly contracts. When the resulting sphere is broken open, there is a small bead of bismuth metal along with a white fused oxide and a yellowish oxide. The molten liquid was splashed onto a cold surface to form a yellow bead that has a small bead of bismuth metal poking through one end. Since the melting point of bismuth trioxide is around 700 C, it easily melted in this hot flame. Some evidently reacted with either the iron wire loop or the aluminium foil to form the metal oxide and bismuth metal. The yellow surface dulled overnight; it is only barely visible in the second picture.


Silver reduction: Manganese dioxide is placed in the silver acetate solution. An abundance of hydrogen peroxide is decomposed. A copper wire is added. It gets coated with a thin layer of silver but no further reaction occurs. The solution turns green as the copper dissolves. Here is the copper wire at its initial and final silver level.

Further waiting produces a small amount of purplish micro-particulate silver, which collected on the bottom of the vial. Swishing around the precipitate produces a purplish suspension.

Reaction of ascorbic acid with some of the silver acetate solution forms the same purplish coloration without producing actual flecks of silver metal. The video that was taken was deleted because of its unimpressiveness.

The original silver acetate shows a little signs of disappearance as well; it is a little bluish, indicating some dispersed silver. Some of the silver acetate solution is placed in the calcium iodide/iodate mixture from a previous experiment. The residual hydrogen peroxide and acidity oxidized some of the iodide to triiodide again, leaving a yellow precipitate and brown solution. When ascorbic acid is placed in the solution, the triiodide is reduced to iodide, the hydrogen peroxide is reduced to water, and any residual iodate is reduced to iodide. This leaving no excess ascorbic acid to reduce the silver iodide, a yellow precipitate formed of silver iodide.

The remainder of the silver acetate solution is neutralized with sodium bicarbonate to form silver carbonate. Several reactions are occurring: baking soda is neutralizing vinegar, silver salts are catalytically decomposing hydrogen peroxide solution, sodium bicarbonate is precipitating silver bicarbonate, and silver bicarbonate is decomposing to silver carbonate and carbon dioxide gas. The resulting precipitate is pure white, showing no contamination of silver oxide. It gradually turns to an off-white color, however.

Nickel-copper sulfate reaction: A piece of nickel from a spark plug ground electrode is placed in moderately concentrated copper(II) sulfate solution. No immediate reaction is observed.

Nickel-copper chloride reaction: Just like above, nickel only reacts very very slowly with copper(II) chloride solution, hot or cold, concentrated or dilute.


Copper dissolution: A large amount of copper wire pieces were placed in 2:1 3% hydrogen peroxide and hydrochloric acid. Dissolution begins immediately. The smelly solution is placed outside.

Indium-copper sulfate reaction: Indium reacts rapidly with copper sulfate, decolorizing the solution and precipitating copper. Unlike aluminium, no difference between the reactions In + CuSO4 and In + CuCl2 is discernible. The resulting almost-colorless solution was reacted with ammonia. The formation of a light blue precipitate showed that a significant amount of copper remained and many ammonia washings will be necessary, possibly even a re-dissolution and precipitation, to get pure white indium hydroxide. It is also a light  and slow-settling precipitate, making the process slower than it was with bismuth hydroxide in the past. The resulting indium hydroxide is still slightly blue, showing that the last traces of copper are difficult to remove. If I ever do this again I will use hydrochloric acid instead.

Mischmetal burn: Mischmetal shavings burn with a sparkling flame. When glowing mischmetal embers are spit upon, they erupt in a manner similar to magnesium.

Pepto-Bismol reduction: Hydrochloric acid is diluted 5 to 1 and a few drops of Pepto-Bismol paste are placed in the solution. They turn white and the bismuth goes into solution. The solution is filtered and a piece of zinc is placed in it. A thin smear of bismuth metal forms on the surface of the zinc (picture). This is definitely a horrible way to obtain bismuth metal.

Silver dissolution with tracer: The residue of silver wire from the previous silver acetate production is placed in more hydrogen peroxide and acetic acid, with small amount of sodium chloride added as a tracer. The silver residue appears to be catalytically decomposing the hydrogen peroxide without forming any significant amount of silver acetate, as no murky streams are seem coming from the silver.


Indium hydroxide production: It is finished today (see above).

Copper(II) chloride formation: The copper dissolving solution is quite dark, and turns turbid when water is added, showing that some copper(I) chloride is also dissolved from comproportionation. I will have to just let it evaporate and reap the huge harvest of crystals.

Nickel dissolution: Nickel does not seem to be dissolving in acetic acid – hydrogen peroxide mixture. The passivation layer is really strong, on the scale of aluminium or higher.


Copper(II) chloride production: The resulting solution contains a significant amount of copper(I) chloride, which is what makes it so dark. I noticed that a white precipitate was collecting at the bottom of the container, as well as whitish substances on the top. To confirm my fears, I added hydrogen peroxide. A precipitate of copper(II) hydroxide chloride formed, turning the solution blue, and catalytic decomposition of the hydrogen peroxide commenced. All of the added hydrochloric acid was consumed in the formation of the coordination complex with copper(I)/copper(II), leaving none to dissolve the copper(II) oxychloride. I added more HCl to make a nice bluish solution, pretty much clear except for dust (I hate open top evaporation for that reason).

Indium hydroxide production: I took the small but reasonable amount of dried indium hydroxide and placed it on a piece of indium metal for photography, just for fun. The indium hydroxide has a very light blue tinge from residual copper; the dried and hydrolyzed copper(II) carbonate hydroxide mixture is less bluish than the ammine solution that was mixed with the precipitate in aqueous ammonia solution. Therefore, it hardly appears blue in this picture. The indium hydroxide is then placed in acetic acid. It appears to partially dissolve. However, no longer gelatinous to any degree, most of it settles on the bottom. IIRC, aluminium hydroxide dissolved in acetic acid when freshly precipitated. When it was evaporated, it left a crystalline mass that was wet and smelled very strongly of acetic acid, probably some form of basic aluminium acetate. Indium hydroxide probably behaves similarly, dissolving completely after an extended period of time.

Zinc reduction: Zinc is often the metal used to reduce other substances. However, magnesium is a strong enough reducing agent to reduce zinc. I dissolved some zinc acetate in water, forming a clear solution. (The zinc acetate was made 2 years ago by cutting up scratched up pennies and dissolving then very slowly in a yogurt cup of vinegar. Then a fan was blown on the solution to dry it. It took very long to dry. One day I knocked the solution over and about 80% of it spilled on the floor. I forgot about the 20% remainder and was surprised to see a mass of soft, damp crystals forming in the container a couple weeks later. I canned them and never had a use for such boring compounds as zinc compounds are until recently.) I then decided to go whole hog and chucked the entire end of my magnesium fire-starter into the solution. It began bubbling gently. After 12 hours, it was covered with bluish-gray zinc powder, some of which had fallen off to the bottom of the solution. After 24 more hours, the zinc had turned white from surface aerial oxidation.

Mischmetal oxidation: Mischmetal reacts with warm water to form a dark grayish film of metal hydroxides, along with hydrogen gas.


Lead-copper(II) chloride reaction: Lead reacts with concentrated copper(II) chloride solution to form, initially, a dark green mixed oxidation state copper complex. This is because only a small amount of copper was formed before the insoluble lead(II) chloride began forming a protective layer over the lead metal. The copper then dissolved in the excess copper(II) chloride. When the copper complex (which hydrolyzed to white copper(I) chloride and green copper(II) chloride when diluted with water) was washed away, an area of lead(II) chloride remained.



Vacuum tube getter: This getter could be either barium or caesium. However, it seems to be barium because 1) it does not react instantly with air – it takes a few seconds – and 2) its reaction with water is not instant either. The top picture shows the white coloration after the getter is exposed to air, as well as the basic pH of the white substance when damp. The left picture shows the mirror-like getter on the top end of the last remaining vacuum tube. Its golden color made me think that it was caesium, but barium could be a more realistic guess. Wikipedia says that barium is the most common getter material, so it is likely barium.



Copper(II) chloride dehydration: I heated the copper(II) chloride crystals in a test tube. They appeared to melt, but quickly solidified to a brown solid while releasing a mixture of HCl and water vapor. The remainder was a bland brown solid quite different from the dihydrate. When water is added, heat is released and a green substance is formed, which appears to be an insoluble basic copper chloride due to overheating of the test tube. The hydration is videotaped.

Nickel(II) chloride reduction: Nickel is dissolved in a mixture of hydrogen peroxide and hydrochloric acid. It is then neutralized with baking soda and the precipitate redissolved, making the solution only slightly acidic. After being split into three parts, three metals were added: magnesium, zinc, and iron. Based on magnetism, no precipitate of nickel formed in any container, showing that no divalent nickel was reduced at these concentrations.

Copper(II) oxide dissolution: After copper(II) oxide is calcined it is insoluble in 5% acetic acid, forming no coloration when mixed.




Lithium and oil: Lithium appears to react slowly with olive oil, probably because of impurities in the oil. But cooking oils can be used as a temporary storage medium for lithium metal. Here is the lithium in the oil, as well as the original lithium from the battery.


Cobalt(II) hydroxide: I reacted lithium with water to form a warm solution of concentrated lithium hydroxide. When cobalt(II) chloride crystals were added, they became coated with a layer of blue cobalt(II) hydroxide which turned red very quickly in the elevated temperature from the lithium-water reaction.

Triiodide disproportionation: The decoloration of tincture of iodine occurs very quickly when lithium is placed in it.

Lithium – mischmetal: Lithium mixed loosely with mischmetal shavings does not easily ignite, showing that increased reactivity in the alkali metals as compared to the alkaline earth metals and lanthanides does not mean increased flammability.


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April 12 2012 5 12 /04 /April /2012 18:58

3-8 - Use Control + F to find what you want

Copper-sulfate ascorbic acid reaction: When sodium bicarbonate is added to the above solution, a dark red precipitate forms, which soon turns a dirty yellow-green. When ammonia is added, a dark green solution is formed as the precipitate dissolves.

Titanium foil: I obtained my titanium foil from Gallium Source today. It is placed in concentrated copper(II) chloride solution and does not react, just like I expected.

Wet cell chemistry: I soaked a tissue in dilute copper(II) chloride solution and another tissue in saturated sodium chloride solution. I wrapped the copper(II) chloride solution tissue around a carbon rod, then wrapped the NaCl tissue around the copper(II) chloride tissue. I then placed a piece of magnesium foil (generously donated by GalliumSource) on the NaCl tissue and tied it tightly together using twist ties. The cell produced, at its peak, 1.95 V @ 120 mA. It could easily run a red LED. Since the tissues were quite dry, the soaking through of the copper(II) chloride to the magnesium was negligible. The magnesium began fizzing when the cell was short-circuited. Later the copper(II) chloride tissue and the magnesium foil were pressed together to demonstrate short-circuiting and the consequences. Pictures, please…

The top picture shows the disassembled cell. The copper(II) chloride tissue is wrapped around the carbon rod in the center. The center picture shows the removed magnesium foil and the sodium chloride tissue. The bottom picture shows the complete cell. This cell can be improved by 1) using filter paper instead of tissues 2) tying the cell together more tightly 3) using a more concentrated copper(II) chloride solution or 4) enclosing the cell to prevent evaporation. Although the voltage is unlikely to rise as a result of these procedures, the current should jump to around 500 mA.

Copper sulfate reactions: The dark green solution from above turns deep red-brown. When hydrogen peroxide is added, it turns blue, but not copper ammine blue. The bottom remains reddish colored. When the blue and red are remixed, the solution turns back to the original dark green coloration. The copper dust on the bottom of the solution has not entered the reaction. The solution is then disposed of.


Titanium burning: While wide pieces of titanium foil do not burn, pieces about 0.3 mm wide will burn with a white, dazzling flame which closely resembles magnesium. This reaction is videotaped. A blob of titanium oxides is formed on the end of the titanium “thread”.


Titanium vs. magnesium burning: The flames are compared in one video.



Titanium reactions: The titanium foil does not react with hydrochloric acid in the cold for several days, neither does not react with hot hydrochloric acid after 10 minutes. It appears that either the titanium is not finely divided enough or the acid is too dilute.


Solar heat: A magnifying glass directly applied to small chunks (100 mg) of various low-melting metals (zinc, indium, tin, lead, bismuth) only melted indium. When the metals are placed on a blackened wood chip and then heated, all of them melt. Magnesium, however, remains unchanged when heated, showing that the temperature is below 650 C but above 419 C. Maybe the magnesium was too large and reflective to melt. I need to try a smaller piece. Later, indium and bismuth are melted together, as well as indium and tin. Neither alloy melts in boiling water, showing that no eutectic was formed. The first picture is the metals after melting. Position is relative to the periodic table; zinc (top left), indium (center left) , tin and lead (center right), and bismuth (bottom right). The second picture shows the tin-indium (left) and indium-bismuth (right) alloys.

Titanium magnetism: Titanium appears to be paramagnetic by water boat method. In reality, it is paramagnetic. Good!

Sulfur heating: Sulfur melts to a light yellow liquid, then darkens to orange, then red. Ignition happens upon stronger heating, but the sulfur does not completely burn up; it seems that the aluminium foil conducts heat out of the sulfur, or the sulfur gets smothered in its own combustion products. Additional oxygen does not affect the burning because there is only a tiny bit of additional oxygen. I first tried a test tube of hydrogen peroxide with manganese dioxide catalyst, then a covered glass containing hydrogen peroxide with manganese dioxide catalyst.


Lithium-water reaction: Lithium reacts about twice as fast with near-boiling water than with cold water. The difference is seen by the time taken to dissolve, not by the speed of the bubbling.

Lithium heating: A lithium coin cell is dissected. The negative plate contains the lithium metal melted on (assumedly). I scraped some lithium off for the above experiment, then heated the plate containing the rest in a flame. The plate only reached about 200 C (approximately) before a bit of lithium hanging over the edge of the plate became incandescent. I thought that the lithium was igniting, but instead the bright white light spread over the entire piece of lithium and the mass grew almost white hot. The reaction was over in a few seconds and the plate returned to a much lower temperature. No lithium remained as evidenced by a water reaction.

Magnesium nitrides: I tried to form magnesium nitrides by burning magnesium wrapped in aluminium foil. Limited air should force the magnesium to begin burning nitrogen. However, the aluminium foil conducted the heat from the magnesium and extinguished it.

Magnesium-calcium burning: Magnesium shavings with calcium pellets burned beautifully, with orange flashes of calcium flame visible amidst the white magnesium flame.


Magnesium burning: When burning magnesium foil strip is immersed in water, it immediately is extinguished. Large amounts of water can be used to put out a small magnesium fire if the magnesium becomes submerged, not just splashed.

Copper(II) chloride flame test: I heated some copper(II) chloride on an iron (bad choice) wire in a flame. The copper(II) chloride soon changed from the blue-green dihydrate to the brown anhydrous form. Then the flame test began. The flame began green, then changed to a bright blue. Orange sparks were seen shooting out of the top. The copper(I) chloride (decomposition of CuCl2 at high temperature) then reacted with the iron, and the flame test was over. The videotape is poor quality as the bright blues of the flame are too bright for the camera to handle and they appear merely whitish.

Titanium-copper(II) chloride reaction. Copper(II) chloride crystals were placed on titanium metal and heated. As before, the copper(II) chloride dehydrated and the flame test was visible. However, when the titanium got hot enough, a reaction occurred, releasing a large reddish-orange flame and some white fumes. After this reaction is complete, no copper compounds remain as evidenced by the flame test. Was the reaction 2 CuCl2 + Ti à TiCl4 + 2 Cu? Of course, TiCl4 + H2O (present in atmosphere) à TiO2 (white smoke) + HCl, so there would not be any leftover titanium tetrachloride. Actually, it appears that there is very little titanium left over. Underneath the black surface is copper metal, so the black surface could be copper(II) oxide. Here is the aftermath of the reaction (the reaction was videotaped).


Calcium olive oil reaction: Olive oil seems to be a good storage medium for calcium. Well, air also seems just as good for Ca, as long as it remains dry.

Titanium anodization: Titanium is connected to the positive pole of a nine volt battery and a machine screw to the negative pole. A salt-water soaked tissue is wrapped around the head of the screw and this is touched on the titanium. It only takes an instant for a light golden yellow coloration to begin forming on the titanium metal. The coating is hardly visible with a nine volt battery, so a 24VDC power supply is tried. The golden coating becomes orange-tinted and much darker. Some more susceptible spots turn bluish. Here is the picture comparison. The left picture shows the ordinary titanium foil with some anodization on the edges. The right picture shows the nine volt anodization with blue spots (below) and 24VDC anodization (above the line). Patches of thicker oxide coating make the coating blotchy. A submersion technique may be tried next.

Titanium anodization: Titanium is anodized at 24VDC in a salt water bath (I later read that chlorides are not the ideal electrolytes). The golden color persists. A chart I downloaded shows the colors of anodized titanium at various voltages. Up to 140 volts is needed to generate certain colors. The golden color is at the lower end of the spectrum.

Titanium ignition: A thin piece of titanium shaving is easily ignited by the current from a 9 volt battery. Magnesium, however, seems to be a much better conductor, because although it ignites more easily than titanium it did not burn when the current was applied. The magnesium did seem a little thicker than the titanium (magnesium is much structurally weaker so a thin shaving easily breaks).

Lead dissolution: I placed a lead wheel weight in acetic acid and then added hydrogen peroxide. The lead turned yellow-orange. It almost appeared that lead(II) oxide was forming faster than the acetic acid could dissolve it. Unfortunately, when I went to take a picture, the acetic acid had caught up (maybe from sloshing around) and the yellow coloration was gone.


Lead iodide production: I reacted the lead solution (excess hydrogen peroxide present) with some sodium iodide solution (some excess ascorbic acid present). The lead iodide precipitated, and some of it began oxidizing to iodine as the peroxide oxidized both the ascorbic acid and the iodides. I then tried to recrystallize some lead iodide from boiling water solution. The lead iodide dissolved in the boiling water but failed to precipitate when the solution was cooled as the yellow snow. When placed in a saltwater-ice bath, the solution turned white and cloudy, not yellow, as if the lead had hydrolyzed. Maybe lead iodide is just much more soluble in water than I suspected. I took the entire residue of lead iodide, removed the supernatant liquid, and added some water. Heating it on a hot water bath dissolved just about all of the lead iodide, leaving a slightly yellow and cloudy solution behind. When this is cooled to around 40 C in air, a small amount of the iodide precipitates, leaving a yellowish cloudy suspension. When placed in a 10 C water bath to which ice was added, the whole of the iodide precipitates as an amorphous yellow solid in under a minute. This was videotaped through the distorting medium of a drinking glass.

Lead halides again: Lead chloride profusely precipitates as a heavy amorphous powder when hydrochloric acid is reacted with lead acetate. Lead bromide precipitates as an off-white amorphous powder, which soon turns crystalline and pure white. Lead iodide is of course a yellow amorphous powder when precipitated from a boiling water solution by ice.

Copper(I) iodide production: The leftover sodium iodide solution from the previous experiment is reacted with copper(II) chloride crystals. A strange mixture of colors and precipitates occurs, with 2 CuCl2 + 4 NaI à 2 CuI + I2 + 4 NaCl being the initial reaction. Also CuCl2 + ascorbic acid à CuCl and I2 + ascorbic acid àiodide occur. The resulting precipitate is a bland tan color, probably copper(I) iodide. Addition of hydrogen peroxide slightly darkens it, while addition of hydrochloric acid has no effect. The top left picture shows the solution just after the copper(II) chloride crystals were added. Brown iodine, along with light brown copper(I) iodide, is seen. The second picture shows the solution after some agitation. Some copper(I) chloride is visible on the top. The third picture shows the underside of the second picture, with some copper(II) chloride crystals still dissolving. The fourth picture shows the final result.

Lead dioxide formation and destruction: The lead bromide slurry is reacted with sodium hypochlorite. Initially, oxygen is released as the remainder of the hydrogen peroxide decomposes. Then, voluminous amounts of dark brown precipitate form. This precipitate seems to be lead dioxide. The headspace above the precipitate is orange with bromine.

This mixture is reacted with ascorbic acid. After some fizzing and heat, the bromine is reduced to bromide just as quickly as the lead dioxide precipitate is reduced to a light yellow solid, possibly containing some lead monoxide as well as lead bromide and lead hydroxide.

Lead chloride dissolution: Heating the lead chloride slurry did not dissolve it. The headspace was not water, though; it was mostly HCl.

Copper(I) bromide production: Copper sulfate crystals are placed in a mixture of ascorbic acid and sodium bromide. A white precipitate of copper(I) bromide forms.

Lead sulfate production: Lead sulfate is a white solid, just like many other lead compounds.

Lead dissolution: After the lead acetate supernatant was removed, hydrogen peroxide was added to the remaining precipitate. An extremely violent decomposition occurred, forming a large amount of insoluble gray powder (left). When hydrochloric acid is added, the precipitate turns white and a significant portion is insoluble.

More copper(I) iodide: Copper sulfate crystals are placed in a sodium iodide-ascorbic acid solution. A white precipitate forms. The supernatant liquid is bluish because there is excess copper sulfate. This precipitate is reacted with sodium hypochlorite solution. Chlorine was released, and a copper(II) solution was formed. The CuI turned brown on the edges before dissolving, but other than that, no sign of iodine was visible. It could have been found as iodate in the solution.

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April 12 2012 5 12 /04 /April /2012 18:47
I recently produced a tin oxide thermite.
 I took tin metal and dissolved it in hydrochloric acid, then neutralized the resulting solution with sodium bicarbonate. The tin(II) oxychloride was washed with acetic acid, then water, then dried. This light tan powder is then heated. Some combustion occurs, with white tin dioxide smoke being released and red-hot spots appearing and disappearing randomly. However, most of the powder is decomposed to a dark brown color (tin(II) oxide is blue-black but this is not pure). I then grind some magnesium with a file and grind this powder with the tin oxide powder. There is more magnesium than tin(II) oxide. I then grind some magnesium shavings and place them on top of the thermite for ignition. After a few minutes of inaccurate striking, the thermite ignites and does its thing. The thermite reacts noisily like a copper thermite, but without much dispersion. Among the heaps of magnesium oxide ash left behind is some tin metal, which did not fuse together because of the brevity of the reaction.
I really should try making larger thermites. These tiny micro-thermites are so hard to ignite and burn very quickly, without producing any useful products. If I ever make a chromium thermite, I really need to grind much magnesium.
P.S.: I tried grinding magnesium on a grinder to quickly form an ultra-fine powder. The magnesium gets blown all over and hardly any is able to be collected. It is a huge waste of magnesium to try to use such a device.
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April 12 2012 5 12 /04 /April /2012 18:44

Caesium does not have any household sources that I could find. Some exotic vacuum tubes use caesium metal as a getter, but barium is much more common. Caesium formate is used as a brine in fracking, but who has fracking fluids sitting around?



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April 11 2012 4 11 /04 /April /2012 15:59

2-2 - Use Control + F to find what you want in here

Antimony oxidation: Antimony powder is placed in hydrochloric acid. No reaction occurs since antimony is not reactive enough to dissolve in hydrochloric acid. When hydrogen peroxide is added, however, the antimony dissolves, forming either a colorless solution of antimony trichloride or a white precipitate of the trioxide. The reaction was recorded.


Calcium reactions: I recently obtained 10 grams of calcium for $13.00 from Gallium Source LLC (shipping included). The calcium is golden-colored and dry. Some of this calcium is placed in warm and cold water. The warm water reaction is faster than the cold. Addition of acetic acid rapidly speeds up reaction. A 1% solution of acetic acid reacts vigorously with calcium. Here are some small calcium chunks in the solution.


Indium properties: When indium is hammered, it easily picks up dirt. Granules of indium were hammered into a foil, and then folded to demonstrate flexibility. With one tap of a hammer, a granule completely flattens. By the way, the line in the indium granule is a gentle fingernail press.


Alkaline earth metals in acid: Beryllium does not react with 5% acetic acid, while calcium reacts vigorously. Beryllium, however, reacts vigorously with concentrated hydrochloric acid, just like calcium. The resulting piece of etched beryllium is blackish with pronounced crystal structure. It has a strange sweetish smell. I was sure to wash my hands after handling. Here is the piece of beryllium.

Galinstan alloying: This is a very haphazard experiment. Galinstan is first cleaned by a hydrochloric acid bath. Then the galinstan is washed with water and applied to a scratch in a wad of aluminium foil. Some bubbling is observed as the galinstan removes the passivation layer on the aluminium, allowing it to react with the residual water. Most of the water is pipetted away. Another bit of galinstan on unscratched aluminium foil did not absorb in the time interval of this experiment. The galinstan bead forms a thick gray coating containing a mix of galinstan and aluminium oxides. This floats to the top of the bead and can be scraped off, although it begins reforming instantly. Then the galinstan bead was mixed with the aluminium oxide paste. A golden-colored paste was formed, probably as the result of an oxide film. Then all the powder and paste was tossed in hydrochloric acid. A galinstan bead immediately reforms and the aluminium oxide, along with the galinstan content, dissolves. The galinstan bead is removed to another lid and zinc is tossed in the container. Violent fizzing erupts which quickly slows. For some reason, galinstan-impregnated zinc reacts about 10 times slower with hydrochloric acid than normal zinc, a good indication of the alloy’s formation. When the fizzing has sufficiently slowed, the zinc is placed on the lid where the galinstan bead is residing, hungrily awaiting the next bite of metal, be it zinc (I did not know that) or aluminium. When a large, unresisting, and recently cleaned (by the acid bath) piece of zinc already weakened by gallium reduction from the hydrochloric acid is added, the temptation to the galinstan is too great to resist. My first inkling that an invasion was occurring was the creeping of a silver-colored area on the opposite side of the zinc metal. This caught my attention and it was videotaped as it grew. It grew to about two-thirds of the zinc’s surface area before I lifted the zinc. Half of the galinstan was absorbed into the zinc, while the other half was flattened against the zinc in the process of being absorbed. However, the zinc was becoming saturated with galinstan. When the metal was tilted, a pocket of liquid galinstan formed on the lower side. By this time, the metal was being held by pliers in a location that was intentionally spared from the acid bath and was farthest from the site of the invasion by the next door neighbor. These various features of the metal were taped, using the last of my tiny memory card. I accidentally touched part of the zinc that did not appear to be amalgamated and the piece instantly fractured. I quickly placed the zinc in a Petri dish to prevent it from cracking out of the pliers and shattering on the floor. The end that was being held broke off. I can only imagine how brittle the silvery galinstan-saturated section is. Meanwhile, a thin oxide coating has begun forming on the zinc, turning it a milky yellow of Tyndale effect lore. Here are the pictures: Top is the zinc-galinstan alloy. In the flashbulb light, the shiny galinstan-soaked portion appears dark, probably because reflection is more specular (learn physics). Below, the difference between normal zinc and galinstan-impregnated zinc in reaction with hydrochloric acid is seen. Bottom left is the galinstan-aluminium oxide paste that was a golden color, along with the powder. Bottom right is the appearance of the galinstan in the cleaning bath.

Indium: I stepped on a stray indium bead. It prostrated itself so completely to the ground that it never rose again. No other stable metal is so submissive and shapeable.


Calcium burning: Small pieces of calcium are burnt in a wire loop and the results are videotaped.

Aluminium alloying: A bead of galinstan is placed on lightly scratched beverage can aluminium. The scratching was not enough to remove the thin transparent protective layer on the inside of the can. More vigorous scratching at a new location and moving the bead of galinstan to that location begins the alloy process. A shell of aluminium oxide forms around the liquid galinstan. The aluminium is manifesting evidence of corrosion on an edge about ¾ inch from the original introduction of the galinstan, showing that the aluminium is probably thoroughly impregnated.  The red dashed arrow shows the aluminium oxide shell. The blue circles show the galinstan bead in its shell, top view (bottom) and bottom view (top). The red solid arrow and red circle both show evidence of corrosion in areas of aluminium exposed to the air.


Indium dissolution: A small piece of indium foil is placed in hydrochloric acid. It is completely dissolved within 12 hours. The dissolution begins extremely slowly but evidently increases in speed as any oxide layers are removed from the metal. Two short videos are taken. Later, sodium bicarbonate is added to the solution. A white precipitate forms which easily passes through filter paper. This is most likely indium(III) hydroxide. Like aluminium(III) hydroxide it is very gelatinous.


Indium reactions: A piece of In is placed in CuCl2solution. The indium gets covered by a spongy Cu layer immediately. The Cu thickens quickly, and a dilute solution of CuCl2 is cleared of Cu in about 1 minute (when the indium is shaken, the reaction is much quicker). Here are some pictures of the reaction.    

The first picture shows the piece of indium foil. The second shows the indium foil floating on the copper(II) chloride solution. The solution has touched the edges and the underside, turning them both reddish-brown. The third shows the indium foil immediately after submersion. The copper layer is thin and dark. The fourth picture shows the indium foil about 10 seconds later. The copper layer has grown in thickness, although the solution has not experienced any significant de-coloration. (The difference in solution color is fluorescent vs. flashbulb light.) The fifth picture shows the thickness of the copper around the indium piece. The last picture shows the solution after the reaction has run to completion and been agitated. Pieces of spongy copper have broken from the indium piece, leaving the blackish-looking indium partially exposed. The solution has decolorized, turning into indium(III) chloride. The reaction that has occurred is 2 In + 3 CuCl2 à 2 InCl3 + 3 Cu. Indium is a quite reactive metal.

Indium reduction: The resulting indium chloride solution from the above experiment is reacted with zinc. Initially, no reaction is apparent, but a layer of indium later is seen to be forming on the zinc. It appears spongy, though not as spongy as the copper layer above. When the indium sponge is later compressed, it behaves like a metal, just like the lead sponge previously made, which also compresses to a solid metal. However, it was lost. These are pictures of the reaction. The first (left) is the zinc just after immersion in the indium chloride solution. The second shows the zinc about 8 hours afterwards. So with the third (the indium sponge is visible on the edges) and the fourth (compare to the first). The fifth shows the indium powder and the zinc outside of the solution. Some of this powder was pressed into the pellet which soon was lost.


Iron comproportionation: Iron(III) oxide is dissolved in an excess of hydrochloric acid to create a yellow solution. A machine bolt is then placed in the solution. A brief and swift dissolution of the electroplated zinc coat occurs, and the hydrogen production stops. It seems as if all of the iron is undergoing this reaction (Fe + 2 FeCl3 à 3 FeCl2) instead of the typical Fe + 2 HCl àH2 + FeCl2. The solution, after 30 minutes, has noticeably turned greener. The reaction will be left overnight. Here is the first set of pictures. The first picture shows the iron(III) oxide completing its dissolution in hydrochloric acid. The second picture shows the machine screw 30 seconds after starting. The third picture shows the solution after 30 minutes.

Beryllium copper(II) chloride reaction: Because of beryllium’s many similarities to aluminium (relatively high melting point, dissolution in alkalis, protective oxide coating), I wondered whether beryllium would undergo the same vigorous reaction that aluminium undergoes with copper(II) chloride. However, I did not want to exhaust, contaminate, or ruin my beryllium, so I only used a highly dilute copper(II) chloride solution (boring). The beryllium reacted more vigorously than the aluminium would under similar circumstances, forming hydrogen gas, beryllium chloride (somehow it stays in solution from the copper(II) chloride’s excess acidity), and copper metal. The beryllium was covered in a layer of dark brown copper, which smeared as it washed off, giving the impression that the beryllium was corroding. However, upon rubbing, the strange dark luster of corroded beryllium metal shone again. The brief (to prevent excess beryllium dissolution) reaction was videotaped.



Iron comproportionation: The reaction was deemed complete by morning, after about 10 hours. The solution was neutralized with sodium bicarbonate to precipitate the iron. A white precipitate formed. This white precipitate is iron(II) carbonate, which is white when pure and oxygen-free as a result of the vigorous carbon dioxide bubbling through the solution. Some of the solution not in contact with the iron was smeared near to the top of the vial, and it retains the brown color of iron(III) oxide (iron+++ doesn’t form a carbonate). This shows that the comproportionation reaction was completely successful and the resulting solution only contained ferrous, not ferric, ions. After a few minutes, the edges and top of the iron(II) carbonate had turned either brown or dark green, the products of aerial oxidation. Ordinary tap water, which contains dissolved oxygen, was added to the solution, and the iron(II) carbonate darkened to a greenish color. Addition of hydrogen peroxide turned it brownish amid fizzing. The left picture shows the resulting iron(II) chloride solution, containing excess hydrochloric acid. The center picture shows the white iron(II) carbonate precipitate. The brown spot at the top is the iron(III) oxide from the unreacted iron(III) chloride solution. The right picture shows the iron(II) carbonate after standing a few minutes.



Titanium reactions: Titanium foil piece from GalliumSource is placed in hydrochloric acid. A vigorous reaction occurs and the titanium dissolves as fast as a piece of aluminium would. Sodium bicarbonate is added. A precipitate occurs even in a strongly acidic solution (much fumes, pH < 1). Wasn’t titanium supposed to dissolve slowly in even boiling hydrochloric acid? Well, the heat produced by this reaction was enough to boil away the hydrochloric acid. The translucent gel on the side of the vial in the picture below is similar to aluminium hydroxide, which would not exist in such a strongly acid solution. The precipitate, however, is curdy and relatively heavy.





“Titanium” reactions: The foil reacts slowly with water, forming small bubbles of hydrogen gas and a white precipitate. This is definitely magnesium.





Metal and copper reactions: The reactions of magnesium, aluminium, and beryllium with copper sulfate and copper chloride solutions are compared. Aluminium does not react with copper sulfate, while it reacts vigorously with copper chloride. Magnesium and beryllium only exhibit slightly increased signs of reactivity with my excess-acid copper chloride; the excess acid is in all likelihood the only reason for the increased reaction rate. Therefore, only aluminium exhibits the strange phenomenon of being so much more reactive with the chloride than with the sulfate. The pictures are of Al-chloride (top left), Al-sulfate (top center), Be-chloride (top right), Be-sulfate (center left), Mg-chloride (center right), and Mg-sulfate (bottom).





Magnesium reactions: Magnesium reacts vigorously with concentrated copper sulfate solution, just like its reaction with copper chloride minus the acidity. Magnesium also reacts violently with a copper sulfate-ascorbic acid mixture, which turns green for some reason (complex formation?). Brown copper mud is produced in the reaction. Magnesium reacts moderately with plain ascorbic acid, showing a rare instance where ascorbic acid functions as an oxidizing agent. Top left is the magnesium-copper sulfate reaction. Top right is the magnesium-copper sulfate-ascorbic acid reaction. Bottom is the magnesium-ascorbic acid reaction, with small bubbles escaping along the sides.




Copper sulfate – ascorbic acid reaction: This try was not good. Too much ascorbic acid was present.





Copper sulfate – ascorbic acid reaction: Copper sulfate solution, which is sky blue, reacted with ascorbic acid to form a green solution. A tiny amount of fine copper metal dust has precipitated.



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April 11 2012 4 11 /04 /April /2012 15:47

I obtained iron-free manganese dioxide (probably mixed with carbon) from a tantalum capacitor. It is quite uniformly black, with some silver electrodes connecting it to the current. When placed in cold hydrochloric acid, it releases a small quantity of chlorine gas and produces a brown solution. I initially thought that this was an indication of iron, but the solution decolorized when heated, showing that the manganese dioxide was relatively pure. It seems that manganese(III) is relatively stable as a chloro complex in cold solutions, but decomposes to chlorine and manganese(II) in hot solutions.


If you have a hard time bringing a hot plate or any other heat source outside to perform the reaction, you may absorb the chlorine produced by a tissue soaked in either ascorbic acid or sodium metabisulfite mixed with sodium carbonate. No chlorine smell will be produced until the tissue is exhausted of its chemicals. After the dissolution, pure manganese(II) chloride solution will remain. If you desire to keep a source of manganese ions, precipitate the manganese as the carbonate, which is resistant to further aerial oxidation, unlike the hydroxide. Manganese(II) chloride can be evaporated but it needs to be stored in a tightly sealed container to avoid any deliquescence.


Manganese dioxide from an alkaline battery or a carbon-zinc battery is more problematic. In these, large quantities of iron impurities are generally present, which can be difficult to extract. Occasionally, you may find a battery that happens to be free of iron impurities. If so, a colorless solution will be formed upon heating the manganese dioxide with hydrochloric acid. Otherwise, a dirt brown solution will form. This may be remedied by dissolving the manganese dioxide in sodium metabisulfite solution. The sulfite will reduce the manganese(IV) to soluble manganese(II), while the iron(III) remains untouched.


You may get purer manganese dioxide from 3 volt lithium coin cells. I generally have had better experience with these than with alkaline batteries, but they are much more expensive than the aforementioned batteries.   Once I got a batch of pure manganese(II) chloride from an unknown alkaline battery.




Extracting pure manganese salts from batteries is difficult, but the fact that manganese(IV) is a stronger oxidizer than iron(III) is useful to know.

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April 11 2012 4 11 /04 /April /2012 15:32

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


Xenon is a heavy, colorless, odorless gas. It is quite inert and no compounds of it were made until the 1960s. Xenon produces a bright white flash when an electric discharge is passed through it. After the first xenon compounds were made, scientists produced a myriad of xenon compounds, including xenates, perxenates, xenic acid, xenon tetrafluoride, xenon trioxide, etc. Most of these are extremely strong oxidizing agents and highly reactive or unstable in some way.


In element form: Flash bulbs in cameras contain xenon gas. Some expensive headlights are filled with xenon. Xenon is also used in a few flashlight bulbs.


In compound form: No sources found.


Here is my sample of xenon. It is the core of a camera's flash bulb.



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April 10 2012 3 10 /04 /April /2012 15:42

Titanium is a rather unreactive metal, making it difficult to dissolve. There are several ways to oxidize titanium.


Hydrochloric acid: Boiling concentrated hydrochloric acid will slowly dissolve titanium, forming a purple solution of titanium trichloride. When I tried this, hardware store hydrochloric acid was unable to dissolve the titanium. It might work if it was boiled for an extended period of time, however.


Sulfuric, acetic, nitric acid: Titanium is quite resistant to all of these acids. It may dissolve extremely slowly in boiling sulfuric acid but that is a generally poor way of dissolution.


Bromide dissolution: I found somewhere on the internet that titanium has a weakness for bromine. Therefore, I decided to dissolve it by electrolyzing it in a sodium bromide solution with a nine volt battery. I was really amazed that absolutely no bromine was produced. All of the bromine reacted with the titanium, which appeared for form a soluble tetravalent complex with the bromine. When this contacted the sodium hydroxide produced at the cathode, a clumpy white precipitate began forming. The titanium gradually grew thinner until it was just a brittle fragment remaining that was no longer electrically connected. I noticed that the titanium hydroxide had a purplish tint to it, showing that some titanium was found in the trivalent state.




The titanium foil completely dissolved within a short period of time, showing the efficacy of this method. The product is also soluble in hydrochloric acid, forming a colorless solution which turns bright red upon addition of hydrogen peroxide. However, a smell of bromine is noticed, so is the red coloration the titanium peroxo complex or the bromine water? I will need to do more research into this.



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April 10 2012 3 10 /04 /April /2012 15:28

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


Iodine is a lustrous, bluish-black nonmetallic element. Its vapor pressure is relatively high, a strange thing for a solid. When heated, it emits purple fumes (sublimation) which condense as crystals on cold surfaces. Iodine is the last of the four common halogens and therefore the heaviest, highest melting, and least reactive one. It is found in small quantities in nitrate ores in the Atacama Desert, where strange water-soluble minerals can survive because of the minimal rainfall. It is also concentrated in seaweed, where it was discovered. Iodine forms iodides, which are similar to all the other halides. Iodates and periodates are strong oxidizing agents, just like chlorates and perchlorates.


Note: Purchasing large quantities of iodine or its compounds may make the DEA suspect you of being a methamphetamine synthesizer.


In element form: Iodine crystals used to be sold as a water disinfectant by Polar Pure, but the DEA might have put them out of business. React tincture of iodine with slightly acidified sodium hypochlorite. Iodine crystals will precipitate (some will remain dissolved in the alcohol). Tincture of iodine itself contains about 2% iodine, 2% sodium iodide, and some alcohol.


In compound form: Iodized salt contains a trace of iodine, as do most seaweeds. Radiation exposure tablets contain potassium iodide or iodate, made to prevent the radioactive iodine from accumulating in the thyroid.


Here was my sample of iodine. My iodine always evaporates on me since I have such small amounts. This is pure iodine crystals extracted from about 0.5 mL of tincture of iodine. The vapor reversibly stains the paper and my skin brown.



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April 9 2012 2 09 /04 /April /2012 15:32

What is the best way to store lithium?


I tried extra virgin cold-pressed heart-healthy olive oil. The lithium appears to react initially, producing a few bubbles. However, the lithium was covered with its foamy electrolyte and so it was not clear whether it was the lithium or the electrolyte reacting. The reaction appears to stop after a while. Olive oil primarily consists of oleic and palmitic acid esters, which do not seem to react with the lithium metal. So olive oil is a good, though not ideal, way to store lithium metal because of its lightness. Lithium is liable to get stuck on the side of the container after any disturbance because of its density. After the oil drains from the surface, it is susceptible to oxidation and therefore is destroyed.


Plastic wrap is the first method that I used to store lithium metal. However, the lithium's reaction with almost all atmospheric gases guaranteed that some gas would find its way into the plastic wrap. The lithium did have somewhat of a restriction on reacting; I detected a distinct smell of ammonia when the resulting white powder was placed in water, signifying that nitride was produced due to the lack of oxygen.


I then tried petroleum jelly. Lithium keeps very well in this medium. Since it does not float in petroleum jelly, there is no worry about lithium being isolated on the side of a container. A large piece of lithium can be covered with petroleum jelly and pressed against the side of a container to make an excellent lithium display for an element collection. A disadvantage is that the petroleum jelly coats the surface of the lithium, making it less likely to react and contaminating any reaction.


LIthium can remain in air for about a day before it starts significantly oxidizing (turning white). This is for normal air with about 40% humidity. With extra-humid air, however, the time period is much shorter. If you are opening a lithium coin cell for only one quick experiment, it is best to skip preservation and just use the lithium as-is. For a collection or display, use petroleum jelly. For further experiments, use petroleum jelly but be sure to completely wipe off the lithium before use.

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