I tried to extract pure chromium metal from a stainless steel spoon. It turned out to be much more difficult than I expected.
I had two extraction theories. One was to oxidize the electrolyzed chromium to chromate(VI), reduce it to chromium(III) in aqueous solution, precipitate it as chromium hydroxide, dry it, and react it with magnesium powder in a thermite-type reaction to make chromium metal.
Another was to use the chromium(III) in aqueous solution and electrolyze it to produce chromium metal at the cathode.
Trying it was a different matter.
First, I anodized the spoon in salt water solution. A huge amount of metal hydroxides were produced. I filtered them and placed the paste in sodium hypochlorite to oxidize the chromium(III) to chromate(VI). However, I did not realize that the inside of the huge lump of hydroxides still had a large amount of unoxidized iron(II) hydroxide. The oxidation from iron(II) to iron(III) used up most of the oxidative power of the bleach, severely diluting the solution.
Then, I reduced the light yellow chromate solution with ascorbic acid. It formed, as before, a light purple solution of some chromium complex. However, the resulting solution was by now so dilute that nothing could really be done with it.
I took a small portion, added hydrochloric acid and reacted it with zinc metal. No chromium(II) ions were produced in a measurable quantity.
I then took another portion, larger this time, and added some ammonia. Apparently, chromium(III) hydroxide forms a complex with ammonia, preventing it from precipitating anything.
I then reacted the previous ammonia mixture with sodium carbonate. Nothing precipitated.
I then took some more chromium(III) solution and reacted it with sodium carbonate. Nothing precipitate. Apparently, chromium(III) forms a complex with ascorbate.
I then electrolyzed some chromium(III) solution. Nothing happened.
I reduced the remaining chromate to chromium(III) to avoid environmental catastrophe and dumped down the drain.
This page will detail all of the ammine complexes that can be formed with various metals.
Copper(II): Copper(II) ammine complex shows more information about this complex. This complex may be dried after being washed with an organic solvent, forming an extremely dark crystalline mass that tends to turn whitish in air. I have never done this without hydrolysis to insoluble copper hydroxide.
Chromium(III): Chromium forms an ammine complex as well. It forms when chromium(III) chloride is reacted with excess ammonia.
Cobalt(III): Cobalt(II) forms several ammine complexes which easily oxidize in air to cobalt(III). I have yet to make some of these.
Nickel(II): Nickel(II) forms an ammine complex when nickel(II) chloride is reacted with excess ammonia. It is much bluer than other nickel solutions.
Silver(I): Silver forms a colorless ammine complex when silver chloride, oxide, carbonate, or nitrate is dissolved in ammonia. Upon standing, it can precipitate dangerously explosive silver nitride. It was used in the past to make silver mirrors by reacting with KOH and glucose.
Zinc: Zinc forms a colorless ammine complex when zinc hydroxide reacts with excess ammonia.
There are several super-strong acids that exist, with highly interesting properties.
Sulfuric acid: The most common of the strong acids, it is the acid that all other strong acids are compared to. Sulfuric acid does not have an extreme ability to protonate substances, but it is a highly reactive substance.
Magic acid: Containing antimony pentafluoride and fluorosulfonic acid, it was called this because of its ability to dissolve hydrocarbons such as paraffin wax. Dissolving a candle in this acid was a magic trick.
Fluoroantimonic acid: This acid contains antimony pentafluoride with hydrofluoric acid. This forms the extremely stable cation SbF6(-), leaving the hydronium ion quite bare.
Triflic acid: one of the earliest discovered superacids, this acid is relatively tame and in common use, but is still much stronger than sulfuric acid.
This is nasty stuff.
What temperature can be reached with a three-inch magnifying glass in direct sunlight?
1. Indium melts when directly focused on it. Indium has a melting point of 156.6 degrees Celsius. At least this temperature can be reached by focusing the light onto a shiny piece of metal.
To get a higher temperature, I needed to use a dark surface. I charred the surface of a piece of wood with the magnifying glass and placed my metal sample on the dark surface. Upon focusing the light, the metal melted.
2. Tin, bismuth, and lead easily melted. Their respective melting points are 232, 271.5, and 327.5 degrees Celsius. All of these temperature are easily reached by using this method.
3. Zinc melts with difficulty, requiring an extended period of time to warm. Zinc has a melting point of 419.4 degrees Celsius.
4. Magnesium burns with difficulty. This will not work with a large piece of magnesium; to get the magnesium to burn I cut a thin strip of the foil and focused the light on the super-thin and fragile end. Magnesium autoignites at 473 degrees Celsius. Hotter temperatures may be reachable but not very easily by such a primitive method.
5. Titanium does not burn using a dark charred wood base. The wood turns white due to ash formation and disappears from under the titanium, never allowing it to reach the maximum temperature. Titanium autoignites at 1200 degrees Celsius. A different method will need to be tried to determine if this temperature can be attainable.
I have finished my second semester of general chemistry labs. I have had a generally positive experience with them academically, although there is some to be desired scientifically. Here is my opinion on ways to improve the chemistry labs in general chemistry.
1) There should be more questions on the lab reports that make students do some research about chemistry. This should not be inevery question, but once in a while, there should be an unknown result which will have to be researched. Also, the practical aspects of the labs and their use in a real-life setting can be mentioned, to help connect the student's learning to their career (if they are planning on a chemistry-related career).
2) Providing problems which will have to be solved by existing chemistry knowledge will help students apply all that they have learned about chemistry to synthesize a solution or a compound, for instance.
3) Do not do all of the preparation work for the students. Allow them to make their own solutions, fill their own burets, etc. This will help enforce the basic concepts such as molarity and molality much more than having everything done already. Maybe even leave some leeway in the instructions instead of providing a point-by-point process that can be followed without thinking or learning.
4) Perform some interesting and some perplexing demonstrations occasionally to get students interested in the course more.
5) Lab write-ups should be more thorough for those labs which are more complex. Obviously, there is not much to write about in a titration lab, for example, but other labs will benefit from a longer write-up.
Here are some things that are already good in the labs.
1) The labs are not too long, which helps the student to keep the goal in mind.
2) The instructors were very helpful and ready to fix any problems that the students might have.
3) Some of the experiments were quite thought-provoking for me as they did not produce the expected result or the results were just plain interested (mercury amide is just a weird and interesting-sounding compound).
I may expand this later.
Trivalent iron has a significant amount of oxidative power to it. Here I will list some of the reactions of this common yet interesting substance.
Thermites: Iron(III) oxide (or red iron oxide) is one of the most common ingredients in thermites. Iron thermites tend to burn extremely hot, yet without much violence, allowing a prolonged heating of the supporting materials. See Magnesium - iron(III) oxide thermite for my experience with an iron(III) oxide thermite.
Most of the other experiments involve trivalent iron in aqueous solution as a chloro complex. To make this, iron(III) oxide is dissolved in hydrochloric acid and the resulting brown and highly acidic solution is the starting point.
Iron: Iron comproportionates with iron(III) chloride, forming a greenish solution of iron(II) chloride. This reaction goes to completion, as evidenced by the complete whiteness when the solution is neutralized with sodium carbonate. See Comproportionation Reaction for other comproportionation reactions including the iron one.
Nickel: Nickel reacts with iron(III) chloride, producing a green solution of iron(II) and nickel(II) ions. While nickel dissolves very slowly with hydrogen peroxide/hydrochloric acid and dissolution is almost nonexistent with pure HCl, nickel dissolves quite rapidly in trivalent iron solution.
Copper: Copper reacts with iron(III) chloride, forming copper(II) and iron(II) ions. However, copper(II) reacts further with copper, producing white and insoluble copper(I) chloride. This then further reacts with more copper(II) chloride, producing a brown solution. This breaks up when diluted with water to white copper(I) chloride and green copper(II) chloride again.
Bismuth: Bismuth dissolves extremely slowly and incompletely in iron(III) chloride solution. When the resulting solution is diluted, some bismuth hydrolysis is noticed.
Tin: Tin reacts with iron(III) chloride solution, forming tin(IV) ions and iron metal. Since tin and its divalent ion are both relatively strong reducing agents, they reduce iron(III) all the way to iron metal.
Lead: Lead reacts with iron(III) chloride, forming insoluble lead(II) chloride and iron(II) ions. However, the reaction does not run to completion because the lead(II) chloride forms an impermeable layer.
As you can see, iron(III) is quite a reactive substance, reacting with and reducing most metals. That is why it is used as an etchant.
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.
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.