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July 1 2012 1 01 /07 /July /2012 02:31

There are a large variety of chemicals that can be obtained for use in cleaning and disinfecting pools and spas. These can often provide a good source of chemicals to the home experimenter.

 

Chlorine shock or chlorine pool oxidizer is calcium hypochlorite, typically from 48 to 65%. This can be used for generation of chlorine gas. Household bleach, a dilute solution of sodium hypochlorite, tends to dissolve most of the chlorine formed when hydrochloric acid is reacted with it. The higher concentration in the calcium hypochlorite allows for the more efficient production of chlorine gas. The remainder is probably inert calcium chloride which does not significantly enter the reaction. Household bleach also has sodium chloride in it due to the similarities of the reaction.

 

Several chemicals are used to increase the pH or alkalinity of a pool. The most common is pure sodium carbonate. This chemical is useful for precipitating metals as their carbonates. (Sodium bicarbonate may be a better choice for neutralizations and some reactions.) It might be better, though, to buy sodium carbonate as "washing soda", or to simply heat sodium bicarbonate in an oven at 300 degrees Fahrenheit to form the carbonate. To increase alkalinity (and not pH), sodium bicarbonate is used.

 

To decrease pH, sodium bisulfate is the most common modern chemical. This is slightly more expensive than sodium carbonate, but is quite useful as an alternative to sulfuric acid in beginner chemistry. Hydrochloric (muriatic) acid is also used, although it has fallen into disfavor due to the fumes it releases. This might be easier to find in a hardware store for experimental purposes. Buffered forms of muriatic acid are impure and unsuitable for most experiments. However, it is still possible to find traditional 31.45% muriatic acid to decrease the pH of pool water. Sulfuric acid is also used to decrease pH.

 

Potassium persulfate (Oxone) is used as a non-chlorine pool shock. The purity is disappointing, only about 33%. However, this is still fine when qualitative oxidation is only necessary. Purer forms of persulfates may be available as etchants.

 

Sodium bromide is available from two sources. It is sold as a spa bromide reserve in the form of pure crystals. Some algaecides also use sodium bromide.

 

Copper sulfate is occasionally used to kill algae in pools.

 

Alum (potassium aluminium sulfate) is used to clear up the water in pools.

 

Calcium chloride crystals (ice melt) is used to increase hardness in pool water.

 

Sodium thiosulfate is used to decrease chlorine level in pools. 

 

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July 1 2012 1 01 /07 /July /2012 02:15

Yesterday, I went to see some fireworks that my town was having. As always, I thought of the chemical compositions of the fireworks while watching the display.

 

Fireworks are colored using several elements. A green coloration is obtained using barium or some forms of copper. A rarer blue coloration can also be obtained by copper. A purple coloration is produced with potassium, while an orange coloration is produced with calcium. Either lithium or strontium can be used to create a red color. Combinations of colors are produced by layering these various elements in the pyrotechnic mixture. As it burns up, the different colors are shown.

 

Some fireworks leave streams of glowing particles behind or produce an array of sparks. If the particles are a dull orange, then they consist of iron filings. If they are bright white, they are most likely magnesium. If they are slightly orange, they could be either aluminium or titanium powder.

 

My favorite fireworks were the flash powder ones. While I am not definite about the composition, bright flash powders are often made using magnesium powder and potassium perchlorate. This mixture is quite stable until it is ignited, after which it burns rapidly. Enclosing this easily makes a bright spot in the sky and a loud bang as the enclosing material is torn apart, similar to a flash-bang grenade.

 

Plain and simple chemistry is not the only element necessary to create a dazzling fireworks display. The physical arrangement of fireworks can provide a variety of effects. Here are several pages detailing the physical aspects of fireworks.

 

Enjoy the fireworks this holiday season (for US readers) and remember the intricacies of chemistry that produce the dazzling effects.

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June 23 2012 7 23 /06 /June /2012 21:10

My ultimate goal was to form the dioxide of both elements. With that goal in mind, I performed a wide range of experiments but failed to obtain the goal in either case.

 

My tellurium came in crystalline chunks. These were extremely brittle and somewhat conductive of electricity.  My selenium came in glassy beads. These were also brittle but did not conduct electricity. They are probably the black allotrope of selenium.

 

When either element is heated, it volatilizes before it burns. This is because neither of these elements are metals. Tellurium has a similar appearance to a metal but different properties. It melts at 450 C, well below the temperature necessary to burn it. Selenium is even worse. It has a melting point of about 220 C, and the fumes produced by the volatilization are very toxic. Therefore, it is practically impossible to create the oxide using the heat of a propane torch in an open atmosphere.

 

Here is the video of a piece of tellurium being heated on a piece of brick to prevent the highly mobile liquid tellurium from falling out of anything. Volatilization occurs, tellurium is deposited, but no oxidation occurs. The flame test of tellurium vapor is visible as a green coloration.

 

 

When tellurium is heated on an iron loop, it melts and quickly falls through the loop. Again, very little oxidation occurred.  The hot tellurium only left the vapor behind, not any white oxide.

 

 

Tellurium, though conductive to some degree, does not form the dioxide when connected to the anode of an electrolysis apparatus. It does form an interesting purple ion when connected to the cathode. This reacts with water to produce a black water-insoluble solid. This solid is soluble in hydrogen peroxide. I thought that this might be a finely divided form of tellurium, but I do not see any insoluble tellurium dioxide precipitating out as the substance is oxidized. Selenium does not conduct electricity in this state so electrolysis is out of the picture. It is very difficult for tellurium due to its brittleness. Any gentle pressure on the metal causes it to fracture and crumble, making electrolysis a painstaking procedure and a waste of tellurium, of which I only have 1 gram.

 

Here is a video of the electrolysis of tellurium. The tellurium is in the solution and is being contacted with a carbon rod. Uniquely, most of the current flows into the tellurium. A faint smell of hydrogen telluride was noticeable. There was a film of elemental tellurium floating on the surface of the solution after the experiment was complete, indicative of the oxidation of the gas.

 

 

I next tried simple dissolutions. Tellurium did not appear to dissolve in a mixture of hydrochloric acid and hydrogen peroxide, nor did it dissolve in sodium hypochlorite, and hypochlorous acid. Selenium dissolved in sodium hypochlorite but did not appear to dissolve in hydrogen peroxide, or the hydrogen peroxide hydrochloric acid mixture. The red allotrope of selenium is supposed to dissolve in hydrogen peroxide, forming the dioxide. Because of this, I tried to form this allotrope. Blowing hot dilute selenium vapor onto a cold surface forms the red allotrope of selenium. However, this process consumes a tremendous amount of fuel and produces a large amount of highly toxic vapors. Because I am not equipped to do this, I wrote off this process as unattainable. So I tried to melt black selenium in a dish and pour it into water. This formed a small amount of red selenium on the surface of the re-formed black selenium beads. It dissolved in added hydrogen peroxide but there was no residue. This was expected due to the extremely small amount of this form of selenium.

 

After a few more experiments and repetitions of previous experiments, I gave up on finding an easy route to selenium dioxide with my limited equipment and materials. Until I obtain some barium nitrate and sulfuric acid to make nitric acid, I will hold off on any anticipated experiments with these metals. This shows the error in trying to use transition metal chemical pathways (acid dissolution, burning in air, etc.) to perform the same results on semimetals and nonmetals.

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June 17 2012 1 17 /06 /June /2012 00:41
I heard from NurdRage's excellent video (http://www.youtube.com/watch?v=BliWUHSOalU) that one can easily extract lithium metal from commonly available Energizer lithium - iron disulfide batteries. So I decided to go and purchase some of these batteries. There are actually two grades of these batteries. The Advanced grade has significantly greater power than an alkaline battery and increased shelf life (10 years). The Ultimate grade has slightly greater power than the Advanced grade and has a 15 year shelf life. Because my goal was extracting the lithium and I wanted to purchase the largest amount of lithium for the money, I purchased 4 Advanced AA batteries for 7.46USD at Wal-mart. The batteries were quite a bit lighter than alkaline batteries (lithium being the lightest metal known).
 
The next day, I decided to tear open one of the fully charged batteries. These batteries contain lithium metal and iron disulfide (pyrite) as electrodes. Due to my past experience with ripping open batteries, the battery came apart like a charm. It hardly required 5 minutes to open it. (By comparison, the extremely tough sub-C nickel-cadmium batteries in cordless drill battery packs took me about 30 minutes to open.) The lithium almost immediately fell out from between the electrolyte membranes. It was silvery gray but quickly turned golden, then greenish, and gradually began darkening from there. I quickly rolled it up (it was very soft, malleable, and flexible and warm from reactivity with the air) and placed it in some mineral oil I stole from my strontium storage container. It floated to the top, as I expected, but still very little of the lithium was above the liquid due to the container's design.
 
The iron disulfide stunk like hydrogen sulfide, but did not dissolve to any significant extent in hydrochloric acid. Therefore, it was regarded as useless and discarded.
 
Soon after I decided to try my first experiment with the lithium. I took a large piece of it and placed it on the bottom of a burning place. I had previously melted selenium on the opposite side in an attempt to form red selenium. Some selenium adhered still to the plate. Despite that, I placed the lithium on the plate and ignited it with a propane torch.
 
 
The mineral oil ignited first, which soon ignited the lithium. The lithium shrunk, melting, in the flame. (And this is the highest melting of the alkali metals, the hardest, and the toughest. ) The white glowing spots finally manifested themselves into a real flame. In the video, the bright bottom part of the flame is the burning lithium. The long dim flames are the mineral oil combustion, which ceases about half-way into the video. The lithium burns up, leaving behind glowing formations of white lithium oxide ash. (By the way, this ash absorbs carbon dioxide from the air quite well, to the extent that a hydrated form is used in some spaceships to absorb the astronaut's exhalations.)
 
However, on the dark underside of the plate, sinister events were occuring. The selenium melted due to the heat of the combustion (selenium has a low melting point of 221 Celsius). Selenium has a high vapor pressure, causing evaporation from the dish and subsequent deposition on the floor. High concentrations of selenium vapor resulted in a black deposit, while low concentrations resulted in a red deposit. One piece of selenium actually dripped from the dish's underside onto the floor. Selenium vapors are also very toxic, making it difficult to work with. Fortunately, this was done in a well-ventilated place, making the fumes less of a concern than the stain. I took hydrogen peroxide and a wash cloth and easily scrubbed all of the red stains off the floor, but the black metallic looking spot which was the real concern remained. I used a torch to try to spread it out, which was a bad idea due to the amount of toxic selenium vapor released. This made it necessary to wait until the selenium fumes were completely ventilated. Again the hydrogen peroxide was applied and the floor scrubbed, without much mitigation of the black spot. A scrub brush and a scraper were both applied, but without effect. Just in time, however, I recalled an earlier experiment where the size of a selenium bead shrunk significantly when immersed in bleach, hinting that selenium is soluble in sodium hypochlorite. Therefore, I decided to apply bleach. It worked wonderfully. The selenium stain disappeared in a minute, and nothing remained except for a slightly whiter spot on the concrete.
 
I then dissected another battery which I almost completely depleted by running a 200 mA flashlight lamp on it for 24 hours. The battery had a nominal voltage of 0.12 V and a maximum current of 20 mA. I opened it up. Just about all of the lithium foil was completely reacted. The black iron looked the same (the iron disulfide is reduced to iron). The lithium sulfide was only a smeary residue on the electrolyte papers. I expected a large amount of lithium sulfide to be present, so I placed it in water. I was greeted with the hydrogen release from all of the left-over lithium inside the electrolyte "paper". The solution turned black and has remained black ever since, even after filtering. It looks like my lithium sulfide is irreversibly contaminated and useless. Two dollars almost went down the drain. However, there was an extra piece of lithium in the battery that had escaped the depletion reaction. I decided to react it with water to obtain some lithium hydroxide. To do this, I took a plastic paint mixing container and placed a significantly sized piece of lithium in it. I then took a garden hose sprayer and sprayed a fine stream of water into it. Unexpectedly, the lithium ignited, spewing out a long red flame and melting a hole in the bottom of the container as all of the water was consumed and the flame's heat had its effect on the container. It was fortunate that this experiment was done outdoors on a concrete patio instead of on some flammable object. The lithium ignition occurred in several steps. My hypothesis is that a piece of the lithium in the upper part of the container was temporarily covered in water. The water reacted as it ran off the lithium, producing heat and steam. Despite lithium's high heat capacity, it does not exceed that of water, so the lithium easily heated to a high temperature in the absence of the water. Steam-laden hydrogen passing over the metal from a reaction in the bottom of the container heated the metal so much more. Eventually it reached the ignition point. The lithium ignited, igniting in turn the hydrogen gas. The hydrogen gas was ignited in a steady stream, enabling it to ignite more of the lithium. A runaway reaction occurred, during which the lithium melted through the container.
 
Excited by the prospect of the lithium's reactivity, I took another piece of lithium, placed it on a brick, and emptied an eyedropper full of water on it. No ignition occurred. This was because of the lack of the specific circumstances stated beforehand which were used to get ignition. No hydrogen laden steam passed over any portion of the lithium, and there was way too little water.
 
 
Not to be discouraged, I tried again two more times. Each time was a failure. Finally, I decided to reproduce original conditions. Original conditions produced (nearly) original results. Take a watch. The steam hid much of the reaction from view.
 
 
I then hit upon a novel idea. Why not combine my two previous experiments and drop lighted lithium into water? This should produce a more vigorous reaction. I was right. The first experiment involved a small piece of lithium. It was lighted and dropped into the water. The hydrogen immediately caught fire and kept burning throughout most of the video. The lithium reacted completely, leaving only a red-hot sphere of lithium oxide floating on the water by the Leidenfrost effect. When the temperature dropped to a certain extent, it fell down to the bottom with a fizz. This reaction was quite similar to the reaction of a similar piece of sodium with this amount of water.
 
 
 However, I was not satisfied with such a small amount of metal. I decided to step it up a bit. I used a larger piece of lithium and got it burning more thoroughly before dropping it into the water. The lithium ignited the hydrogen with a pop, burning brilliantly for a few seconds. However, due to the reaction thinning out the foil, the vigorous bubbling of hydrogen, and the heat of the flame, the lithium foil broke into several pieces and exploded, shooting gorgeous bits of burning lithium up to six feet into the air. This reaction is more worthy of potassium metal than lithium.
 
I was done with the bangs and flashes of the alkali metals for now but wanted to catch a more close-up video of the combustion of lithium. I did so with a small piece of foil.
 
The mineral oil first caught fire, melting the lithium. Unlike the alkaline earth metals, the alkali metals have a low melting point and generally melt before burning. The lithium solidified into a blob and began burning. The white flame gradually became brighter as the lithium was completely converted to the oxide. Then it turned red due to the flame spectrum of lithium oxide and faded as the lithium burnt up. The lithium oxide ash appeared in an interesting lump, which I immediately preserved from atmospheric attack in a glass vial, though in a crumbled state. The formation shown here is very delicate.
Lithium oxide
Since I wanted to keep some lithium for further experiments without resorting to breaking open another expensive battery, I ceased experimenting with the metal after this.
If you have any ideas about what do with lithium, drop a note in the comments section below.     
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May 21 2012 2 21 /05 /May /2012 16:30

The flammability of magnesium is well-known. Magnesium ignites around 500 degrees Celsius, burning with a temperature of 3100 degrees Celsius. It is highly reactive, burning in oxygen, nitrogen, and carbon dioxide. It also reacts with water and can burn in small amounts of it. (Large amounts of water still extinguish small amounts of magnesium.)

 

Curls of magnesium can be ignited by a 3 inch magnifying glass on a piece of wood. They can also be ignited by sparks from burning ferrocerium, the material used to ignite cigarette lighters. This is the principle behind the camp firestarter.

 

Magnesium burns vigorously when mixed with most metal oxides. Even sodium hydroxide burns with magnesium, forming impure sodium metal. Metal oxides of less reactive metals such as copper forms violent mixtures with magnesium powder. These mixtures burn best when finely powdered and ground together.

 

Magnesium is extinguished quite rapidly in an enclosed location, but not before it burns some nitrogen, forming magnesium nitride. This produces an ammoniacal smell when reacted with water.

 

Burning magnesium powder explodes when struck by a fine stream of water. This can be dangerous, especially if the original intent was extinguishing the fire.

 

When a magnesium strip is placed in water and ignited at the top, it burns down until it hits the water, where it continues burning for a while, forming hydrogen gas from the water. Eventually, the water pulls enough heat from the magnesium to stop the combustion.

 

Magnesium does not readily form sparks when ground in air, but I did get a few sparks. They could have been from impurities (e.g. iron) on the grinding wheel.

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May 14 2012 2 14 /05 /May /2012 14:52

Copper is a first-row transition metal that produces a wide variety of colors in its compounds and complexes. Copper chemistry is interesting and easy, making it optimal for chemistry beginners. One such exercise involves the creation of every color in the rainbow using copper. In my example below, the only starting copper chemical was copper(II) chloride.

 

Copper-rainbow.JPG

 

Red: Copper metal is reddish. In the photo, I took copper(II) chloride solution and reacted it with zinc metal to produce the reddish copper. I also added some ascorbic acid to the solution after the reaction was completed to absorb dissolved oxygen and prevent the copper from oxidizing to black copper(II) oxide. A red coloration is also obtained through copper(I) oxide. When glucose is heated with Benedict's reagent, red copper(I) oxide is produced. Another copper complex is deep red, and preparation instructions are found here.

 

Orange: Copper(I) oxide has the unique ability of forming a range of colors. The exact coloration depends on the rapidity of the compound's formation. When slowly prepared such as in the above reaction of glucose and Benedict's reagent, copper(I) oxide is red. The orange coloration in the picture was obtained by reacting ascorbic acid with copper(II) chloride to precipitate white copper(I) chloride. Sodium carbonate is added, and the orange compound precipitates.

 

Yellow: Tetrachlorocuprate(II) ion has a yellow coloration. When a small amount of copper(II) chloride is reacted with hydrochloric acid, a yellow solution (in the photo) is formed. Reaction of sodium hydroxide with copper(I) chloride precipitate forms yellow copper(I) oxide.

 

Green: A mixture of blue copper(II) aqua complex and yellow copper(II) chloro complex produces this deep green solution. I made it by dissolving a significant amount of copper(II) chloride in hydrochloric acid. When copper(II) chloride is crystallized from this solution, it is a green solid.

 

Blue: A dilute solution of copper(II) chloride (or sulfate or acetate) in water produces the sky-blue coloration of copper(II) aqua ions. Copper sulfate has a blue coloration when solid.

 

Violet: Tetramminecopper(II) solution is often violet. Copper(II) chloride (or carbonate or sulfate or acetate or oxide, provided  that they are not strongly heated) is dissolved in aqueous ammonia. The precipitate of copper(II) hydroxide dissolves in the ammonia, forming a deep purplish solution of cuprammonium chloride.

 

White: Copper(I) chloride, produced when copper(II) chloride is reacted with ascorbic acid, is white until it gets oxidized by air.

 

Black: When copper(II) chloride is reacted with sodium hydroxide and heated to boiling, the unstable blue copper(II) hydroxide decomposes to black copper(II) oxide in solution, despite being surrounded by water.

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May 6 2012 1 06 /05 /May /2012 03:04

 

Hydrolysis is the reaction of a compound or ion with water. Most hydrolysis reactions in inorganic chemistry result in the donation or removal of a proton by water, forming either hydroxide or hydronium ions. Two sample hydrolysis reactions are:

NH4(+) + H2O à NH3 + H3O(+) forms an acidic solution

 

CN(-) + H2O à HCN + OH(-) forms an alkaline solution

Here, ammonium ion reacts with water, generating ammonia and hydronium, which makes an acidic solution. Below, cyanide ion reacts with water as well, generating hydrogen cyanide and hydroxide. Learn more by taking a General Chemistry II course.

These reactions happen with salts of weak acids or weak bases, which is what gives the salts their pH. Many other chemical compounds form their unique hydrolysis products which can be interesting. Most metal oxides are weak bases, making their salts prone to hydrolysis. The less basic the oxide is, the stronger the hydrolysis of the salt and the more acidic the resulting solution. Below several metal and semimetal ions are listed, along with their hydrolysis properties and any uses of such properties.

 

Aluminium: Aluminium forms a stable complex with water, but solutions of aluminium with strong acids are highly acidic, showing the weakness of aluminium(III) hydroxide as a base.

 

Antimony(III) and arsenic: This ion is stable in strongly acidic solutions. When diluted with water, it produces a white precipitate of antimony oxychloride in a similar manner to bismuth. However, the hydrolysis happens at an even lower pH. Arsenic(III) hydrolyzes even more easily than antimony(III) in water. Arsenic pentoxide, however, is insoluble in concentrated hydrochloric acid, preferring instead to dissolve in water, forming its own acidic solution. When the nonmetals are reached, the oxides no longer keep up any pretense of basicity.

 

Bismuth(III): When a colorless acidic (pH 1 or so) solution is bismuth(III) chloride is diluted with water, a white precipitate of bismuth oxychloride is instantly formed, even though the pH remains around 2. Bismuth(III) oxide is a very weak base and so its salts are highly acidic and prone to hydrolysis. This is how bismuth oxychloride, a substance used in cosmetics, is created.

 

Monovalent ions: Monovalent ions hardly even have an acidic pH in most cases, showing that they do not hydrolyze at all. The alkali metal salts with strong acids are completely neutral. Other monovalent ions (thallium, silver, copper) can have different properties because of their greater proximity to the right side of the periodic table, but none are easily hydrolyzed AFAIK.

 

Most divalent ions:  Copper shows a very small tendency toward hydrolysis. The production of an acidic pH of 3 and higher is the most hydrolysis that occurs. Therefore, aqueous solutions of metals which form divalent ions (cobalt, copper, manganese, iron, nickel, zinc, cadmium, mercury, lead, alkaline earth metals) are often the easiest to study in amateur chemistry due to ease of dissolution in water.

 

Niobium(V): Niobium pentachloride is a yellow solid that hydrolyzes completely in water and in moist air. Even in the most acidic solutions it slowly hydrolyzes, depositing white niobium pentoxide.

 

Phosphorus: Although not technically an ion, simple phosphorus compounds are highly prone to hydrolysis. Phosphorus trichloride, a colorless liquid, fumes upon contact with water, forming phosphorous acid and hydrochloric acid. Phosphorus pentachloride also has a vigorous reaction. The reaction between phosphorus triiodide and water (which creates hydriodic acid) is used in illegal drug manufacture.

 

Silicon compounds: Silicon tetrachloride, a typical binary silicon compound, hydrolyzes rapidly with water, releasing silicic acid (hydrated sand, essentially) as a colorless gel along with hydrogen chloride fumes. Because of silicon tetrachloride's unique properties, it is used to produce high-purity silica gel by mixing with water.

 

Sulfuryl chloride: This compound hydrolyzes much more slowly in water. I have heard that it can take weeks for a layer of sulfuryl chloride to completely react with water.

 

Thionyl chloride: This compound reacts vigorously with water, releasing sulfur dioxide gas and hydrogen chloride fumes in large amounts, as I accidentally experienced upon opening a lithium battery containing the substance. The high acidity of this complex is used to prevent sensitive metal chlorides (e.g. rare earth chlorides) from hydrolyzing as they are dehydrated. The thionyl chloride reacts with any released water, releasing a cloud of acidic gases that keep the anhydrous chloride stable.

 

Tin(II): A solution of tin(II) chloride is stable in hydrochloric acid, but slowly hydrolyzes when diluted, forming a white precipitate of tin(II) oxychloride. This is not desired in most circumstances, and is prevented by using hydrochloric acid.

 

Tin(IV): Solutions of tin tetrachloride in water are generally turbid to some degree. Because tin dioxide is a weak base, the solution hydrolyzes easily, forming white insoluble tin dioxide. The anhydrous form of tin tetrachloride fumes upon contact with air or water. This was used in the past as a naval smokescreen.

 

Titanium(IV): Titanium tetrachloride, one of the common tetravalent titanium compounds, hydrolyzes very strongly. When the anhydrous compound is sprayed into air, it forms a dense white smoke as a result of reacting with the water vapor in the air. This is used to determine air flow in a room or to test smoke detectors for effectiveness.

 

 

 

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May 2 2012 4 02 /05 /May /2012 13:04
Hydrogen peroxide is a rather unstable material. It has a strong tendency to exothermically decompose to water and oxygen. This process is limited in drug store hydrogen peroxide by the addition of a stabilizer such as phenacetin. However, certain substances vastly accelerate the decomposition of hydrogen peroxide. One of them is blood, where the enzyme catalase breaks down hydrogen peroxide very quickly and efficiently. Other inorganic compounds break down hydrogen peroxide too.
 
Iodide: This is used in the famous "elephant's toothpaste" demonstration. 30% hydrogen peroxide is mixed with detergent and some potassium iodide is added, after which it vigorously foams up to 30 times its original size. This is not my video. Iodide's decomposition is quite fast, as is evident by these pictures.
 
Copper(II) chloride: This is another strong catalyst in the decomposition of hydrogen peroxide. The green solution visible at the end of this video decomposes the hydrogen peroxide quite rapidly.
Alkali and alkaline earth metal salts: These decompose hydrogen peroxide slowly. Of all of the salts that I tried the carbonates and bicarbonates caused the most decomposition, undoubtedly because of the basic solution produced.
 
Manganese dioxide: One of the most well - known catalysts for hydrogen peroxide, manganese dioxide vigorously and completely decomposes hydrogen peroxide, causing a dirty brownish-black solution to be left behind. The other lower manganese oxides also decompose hydrogen peroxide but to a lesser extent.
 
Cobalt(II) chloride: Cobalt(II) chloride, like most transition metal salts, decomposes hydrogen peroxide relatively fast, but not in a spectacular way.
 
Silver compounds: Soluble silver compounds are also excellent decomposition catalysts. When silver covered in a chloride crust is immersed in hydrogen peroxide, vigorous fizzing begins which only slows a little when acidified by acetic acid. Catalysis continues until all of the peroxide is depleted.
 
Other precious metals: Platinum sponge and other precious metal catalysts likely have a strong effect on peroxide decomposition, but I could not obtain any of these for experimentation.
 
Iron(II) compounds: These are used to make Fenton's reagent by reaction with hydrogen peroxide. While they destabilize hydrogen peroxide to some extent they are not as strong a catalyst as the other transition metal compounds.
 
Chromium(VI) compounds: These actually react with hydrogen peroxide, forming a metastable coordination complex. It soon decomposes, giving off oxygen and returning to the original state in some cases. If it does so (which it does more in basic solution), it is a catalyst, although a slow catalyst in comparison with normal ones.
 
 
Titanium compounds: These form a bright red and surprisingly stable peroxo complex with titanium.
Titanium-peroxo-complex.JPG
Most other metals only decompose hydrogen peroxide to a fair extent or are actually oxidized by it.
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April 30 2012 2 30 /04 /April /2012 13:12

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

 

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

 

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

 

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

 

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

 

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

 

DSCF0295.JPG

 

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

 

Copper-II--oxide.JPG

 

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

 

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

 

Manganese-II-III--oxide.JPG

 

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

 

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

 

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

 

DSCF0304.JPG

 

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

 

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

 

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

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Published by LanthanumK - in Experiments
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April 27 2012 6 27 /04 /April /2012 13:20

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

 

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

 

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

 

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

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