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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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