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January 14 2013 2 14 /01 /January /2013 20:17



Pepper spray is a solvent containing extracts of hot peppers which is placed in a pressurized canister. It is applied to the face of an attacker as a non-lethal self-defense weapon. Pepper spray causes "immediate closing of the eyes, difficulty breathing, runny nose, and coughing." (Wikipedia) Pepper sprays are often rated by their OC content, but all OC is not equal. Generally, inexpensive pepper sprays will advertise a high percentage of OC but will have little strenth in reality.


I was pepper sprayed with the spray from Harbor Freight Tools listed above. The temperature was 30 Fahrenheit (0 Celsius) and there was a slight breeze.


Initially, there was no effect, so it is evident that this no-name brand of pepper spray from HFT did not have the expected potency. So I decided to rub it into my face and eyes. The pain started when I opened my eyes. The spray made me close my eyes immediately, and I was unable to open them again. After one minute, the skin on my face began burning and my nose began running as pepper spray ran through the tear ducts into the nasal passages. No coughing was observed, though. I used paper towels to sponge the majority of the pepper spray from my face. I used a little water but it was so cold outdoors that my face froze. It definitely felt strange to have one's skin numb from cold yet burning up at the same time. After 10 minutes of rubbing with paper towels (something that I was not supposed to do), I entered a slightly warmer room (45 degrees F). Breathing was becoming a little difficult and my nose was running rapidly. After 30 minutes, some sporadic coughing began, and my face continued burning. I continued wiping pepper spray from my face with paper towels. Milk helped relieve the pain but only for a few seconds. After 45 minutes, my eyes were able to be opened again. My face was bright red where I rubbed with the paper towels, and my eyes were quite swollen. After 1.5 hours, most of the symptoms were gone, although when I took a shower, pepper spray ran from my hair into my eyes and made them burn again for 10 minutes or so.


I do not think a bargain brand of pepper spray would stop an attacker due to its low potency, so if you want a spray for self-defense, get a brand-name spray such as Sabre Red.


After the experience, I decided to purchase Police Magnum pepper spray, which seems to be much more potent, although I have not sprayed myself with it yet.

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January 10 2013 5 10 /01 /January /2013 23:22

Right now on the shelf in my laboratory I have the following substances: CuCl2(s), KBrO3(s), NiCl2(aq), SnCl2(aq), MnCO3(s), PbCO3(s), Zn(Ac)2(s), Na2HPO4(s), Li2CO3(s), Na2CO3(s), SrCO3(s), CaCO3(s). Below are the methods I used to make each chemical. Be notified that the method that I used is not always the most efficient synthesis method.


CuCl2: Muriatic acid and 3% H2O2 are mixed in about equal concentrations. (A moderate excess of either ingredient is harmless.) Household copper wire is chopped up and placed in the solution. The solution will begin to turn green as the copper dissolves. Once the solution gets concentrated enough, remove the solution and allow it to evaporate in a dish at room temperature. The dish can be covered with a tissue to prevent dust contamination. Once it turns blue-green (fully dry), put it in a container.


KBrO3: Dissolve NaBr in water to make a moderately concentrated solution and place it in a beaker in a well-ventilated location. Clip a carbon rod and a nail to opposite sides of the glass, with both items in the solution. Find a 5 volt power supply and connect the ground to the nail and the positive wire to the carbon rod. When the power supply is turned on, hydrogen and sodium hydroxide solution will be released from the nail, and smelly red bromine will be released from the carbon rod. Periodically swirl the solution. The red color will disappear as the bromine reacts with the sodium hydroxide to form sodium hypobromite, the bromine equivalent of bleach. When bromine production slows, disconnect the electrode. Filter the black carbon powder from the solution. Heat the solution until it boils for five minutes. This will disproportionate the hypobromite to bromate and bromide. Add potassium nitrate and allow it to evaporate in air. The potassium bromate should crystallize first due to its lower solubility. Scrape it out, pat it dry with a filter paper, and put it in a container.


NiCl2: Muriatic acid and 3% H2O2 are mixed in about equal concentrations. A piece of nickel (Canadian nickel coin that is magnetic) is placed in the solution. The solution begins to turn green as the nickel dissolves. When the nickel stops dissolving, remove the metal, add sodium carbonate, and filter the nickel carbonate precipitate. Wash it and let it dry, then redissolve it in as little muriatic acid as possible.


SnCl2: Muriatic acid is placed in a vial and pieces of pewter is added. Black antimony powder gets left behind and tin(II) chloride is formed. Remove the antimony periodically. Once the pewter stops dissolving, remove all the antimony, add a new piece of antimony to prevent aerial oxidation, and keep the solution out of air.


MnCO3: Muriatic acid is placed in a vial and pieces of cathode material from a carbon-zinc battery are added. (Most battery brands are contaminated with iron oxide, so if the solution is dark brown after filtering, discard it.) Chlorine gas is produced, and the almost colorless manganese(II) chloride is formed. After the reaction stops, let it sit overnight to help the carbon settle, then slowly filter the solution through filter paper. A relatively clear liquid should come through the bottom. Add sodium bicarbonate to this liquid. A light brown precipitate will form and fall to the bottom of the solution. Filter, wash, and dry this precipitate. While manganese(II) hydroxide oxidizes in air to Mn3O4, manganese(II) carbonate does not suffer from such an effect, so it is a good source of manganese(II) for chemical reactions.


PbCO3: White vinegar and 3% H2O2 are mixed in about equal concentrations. A bicycle trip is taken, during which wheel weights (some of them lead) are found and taken off the shoulders of the local roads. The lead ones are sorted out by softness and higher density, and they are chopped up and placed in the solution. Bubbling begins, and a hazy black mixture of antimony, tin, arsenic, bismuth, and other lead alloying materials begins forming as the lead dissolves. Wait until the bubbling stops, then filter the solution. Add sodium bicarbonate to the resulting lead acetate solution. An extremely dense white precipitate of lead carbonate will form and sink to the bottom. Filter, wash, and dry the precipitate.


Zn(Ac)2: White vinegar is placed in a dish. Some new-fangled pennies (which do not contain copper except on the skin) are chopped up and placed in the vinegar. They slowly dissolve over the course of a week or two. The solution is forgotten about for a month and when it is checked, zinc acetate crystals have grown. They are soft, vinegar smelling, but quite crystalline. They are placed in a vial.


Na2HPO4: Baking powder is added to water and the resulting goop is allowed to evaporate. For some reason, when water is added again, the cornstarch did not gum up for me. It was filtered and the filtrate was allowed to evaporate in air. I believe these are reasonably pure disodium hydrogen phosphate crystals.


Li2CO3: Lithium (I know, very expensive) is added to water. The resulting lithium hydroxide is allowed to evaporate in air, furnishing crystals of lithium carbonate.


Na2CO3: Baking soda is heated on a baking sheet at 300 F in an oven for an hour. The carbon dioxide and water escapes and sodium carbonate remains.


SrCO3: Strontium (I know, very expensive) is added to water. The resulting strontium hydroxide mix is allowed to evaporate in air, furnishing strontium carbonate powder.


CaCO3: Calcium oxide is allowed to sit in air. It absorbs carbon dioxide, transforming into calcium carbonate over the span of several years.

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December 11 2012 3 11 /12 /December /2012 20:33
During the course of amateur experimentation, there is always the occasional experiment which has an unexpected problem. Sometimes this is due to poor lab technique and lack of safety precautions, while other times it is one of those one-in-a-thousand coincidences that seem so unreal. In this article I will detail some of my experiences as well as include videos and excerpts of the errors of other amateur scientists. Proper attribution will be given for all work that is not mine.
Hydrogen torch fail: My younger brother occasionally performs chemistry experiments, which do go wrong due to lack of proper planning. In one of these experiments, aluminium foil was placed in muriatic acid in a plastic shampoo bottle. A special nozzle / valve apparatus constructed of spare parts was connected to the modified cap. After the normal waiting time, the reaction began and the small amount of hydrochloric acid in the bottle began heating up as the reaction progressed. Fortunately, he waited a while before igniting the expelled hydrogen, and did not have a gas explosion inside the bottle. However, the bottle was not designed for the kind of heat produced by this reaction, and began to soften and deform in his hands. He panicked and finally decided to drop the bottle in the bathroom sink. He then spilled the hot hydrochloric acid right down the drain, eating the thin chrome plating off the drain plug.
Pyrophoricity: I was exploring battery chemistry and dissecting common household batteries to determine if their contents held any useful chemicals. I opened a nickel-cadmium battery expecting to see a shiny lump of soft cadmium metal in it, but was disappointed. I decided to try opening a similar battery, the nickel metal-hydride battery. Fortunately, I correlated the high power capacity and jelly-roll construction of the NiMH battery with highly reactive chemicals, and placed all of the components on a metal surface. Sure enough, a minute later I come back and the electrolyte paper is flaming. I beat out the small flames with another metal object and tell my younger brother about it. He decides to open a NiMH battery and see whether it could really catch on fire. He dissects it over a metal dish but unwittingly (as I was videotaping) dumped some powder from the battery into his garbage can, which was full of flammables. Half a minute later he smelled smoke  coming from under his table and looked down to see a garbage can full of burning tissues. I tell him to take it to the bathroom sink, where he quenches the flames with water. I do not need to tell you what would happen if these chemicals got casually dumped in the trash.
Tert-Butyllithium fire death: A researcher at the University of California, Los Angeles, was removing tert-butyllithium from a bottle into a syringe when the syringe fell apart and the dangerous chemical became exposed to air. It instantly ignited (due to moisture in the air), causing burns that caused her death. Link_to_article
Sciencemadness member death: User myfanwy94, it appears, passed away a couple of years ago shortly after announcing that he was going to synthesize phosgene. This is not proven, but this link will show the thread on sciencemadness.org about his disappearance. By the way, here is an example of his careless experimentation, this time with sodium metal.
Bottle bomb explosion: Shaking is often not necessary. Holding these dangerous items for too long is the primary reason for premature explosions.
Bottle bombs are notoriously unpredictable, and are made with some nasty chemicals. A toilet bowl cleaner bomb, for instance, will splatter you with boiling hot acid, whereas a chlorine bomb produces toxic gases. Never use these as a prank because of this.
Periodic videos: Chromium trioxide reacts with alcohol, causing a violent flaming reaction. The Professor forgets to shut off the smoke detectors, and there is a problem. Even professional chemists make mistakes sometimes.
Chlorine accident: I mixed manganese dioxide from alkaline batteries with muriatic acid in a small container and left it in an unventilated room. When my mother unwittingly walked into the room 20 minutes later she was not happy with the smell. Of course, there was splatter all over the nice white surface I placed the container on as well. No one was harmed as the chlorine concentration was not dangerously high.
Gasoline flames:
Gasoline has a high vapor pressure, which means it evaporates swiftly. The vapor flows along the ground for quite a distance, mixing with the air to form a highly-flammable air zone. Any spark or flame in that zone causes an instant flare-up. This is why gasoline should not be used to start fires.
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November 6 2012 3 06 /11 /November /2012 19:47
The creation of an artificial atmosphere is necessary for manned underwater and space travel. Humans breathe in oxygen and exhale carbon dioxide and water. Air contains 78% nitrogen, 21% oxygen, 0.04% carbon dioxide, and a trace amount of water vapor. Humans exhale about 4% carbon dioxide and 17% oxygen. 15% oxygen is the minimum that can be breathed safely. Mental impairment starts at 14%, and unconsciousness starts at 10%. Oxygen deprivation causes disorientiation but is not particularly painful. A pig can show that well. (Video clip from the BBC documentary "How to Kill a Human Being")
Carbon dioxide levels below 2% are not harmful, while above that level causes an increased rate of breathing and a sense of suffocation, as carbon dioxide concentration and not oxygen concentration is used to produce the feeling of asphyxiation. Greater than 10% causes "a strangling sensation". (ref) Excess water vapor is not harmful to humans, but it can damage equipment.
Ideally, an artificial atmosphere will add oxygen and remove carbon dioxide and water, keeping them at a normal level. There are several ways to add oxygen. One is a simple oxygen tank. Oxygen is separated from liquefied air and placed in a tank under high pressure. These tanks are readily available in welding supply stores and medical supply companies. However, they are not compact and require storage of a compressed gas. This is not an option for nuclear submarines which stay underwater for months. Instead, electrolysis of water is used. This process runs off the nuclear power from the submarine  and produces pure oxygen from a solution of an electrolyte in water. This is much more compact and can produce a huge amount of oxygen. Sea water cannot be used because electrolysis of salt water produces a toxic mixture of chlorine and oxygen. Another method of oxygen generation is using dry chemicals. A mixture of sodium chlorate with a little iron powder generates a reliable stream of oxygen, while leaving only non-toxic byproducts behind. Some of the sodium chlorate helps burn the iron, while the rest is decomposed by the heat of the burning iron, releasing oxygen and sodium chloride. Other chlorates and perchlorates can be used similarly. Commercial airliners use barium peroxide and sodium chlorate, which react when ignited by a percussive igniter, producing a stream of oxygen gas.
An all-around chemical for creating an artificial atmosphere is potassium superoxide. This chemical reacts with water vapor, releasing oxygen and potassium hydroxide. The potassium hydroxide reacts with exhaled carbon dioxide, forming potassium carbonate and releasing water. The water starts the cycle again, and the reaction goes on until all of the chemicals are exhausted. However, not enough carbon dioxide is absorbed by this reaction, and it is dangerously reactive with liquid water.
Carbon dioxide absorbers are primarily alkalis. The most common is soda lime, which consists of a mixture of sodium and calcium hydroxides. The sodium hydroxide readily absorbs carbon dioxide, forming sodium carbonate. This then reacts with calcium hydroxide, precipitating insoluble calcium carbonate (which drives the reaction forward) and producing sodium hydroxide. On a space station, the more expensive but lightweight lithium hydroxide is used. Due to its smaller atomic mass, a smaller mass of lithium hydroxide can absorb an equal amount of carbon dioxide (compared to soda lime).
Water absorbers can be such common substances as silica gel. Calcium chloride and magnesium sulfate are excellent water absorbers and are quite cheap. Dehumidifiers condense water on cold coils, removing it from the atmosphere by electricity without the need for a chemical treatment. These are better for large artificial atmospheres.
Soda lime and oxygen canisters are not difficult to get hold of, but producing an apparatus with enough surface area and circulation to absorb exhaled carbon dioxide at home can be challenging. Still, it would be an interesting project to create an artificial atmosphere. Firefighters, hazmat crews, and space technicians use them all of the time to get pure breathing air in an unbreathable atmosphere.
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September 25 2012 3 25 /09 /September /2012 19:37

The batteries in common use today have a small variety of chemicals that are used to create power. The electricity is created by separating one chemical reaction into two parts and harnessing the flow of electrons between those two parts. Many chemical reactions can be used, but only a few have the properties required by modern batteries.


The first primary cells (non rechargeable batteries) used chemical reaction between zinc and manganese dioxide. Ammonium chloride is used as an electrolyte (electrically conductive solution). The zinc is on the outside (often behaving as the case of the battery) while the manganese dioxide is mixed with carbon to improve conductivity and is placed inside the cell. A piece of paper is placed between the two chemicals to prevent a short circuit. A carbon rod in the center is used to collect the current from the manganese dioxide. These batteries are the cheapest batteries sold, as well as the lowest quality. Due to the acidic nature of the electrolyte, the zinc case will wear away and eventually will begin leaking the corrosive electrolyte. These batteries have a shelf life of about 1.5 years at most. The small surface area between the zinc and the manganese dioxide also gives a low capacity as well as a low discharge rate. A typical AA carbon-zinc battery can cost 0.15 USD and provide 500 mAh of current. They are becoming increasingly harder to find as a modified version which costs about the same but has more capacity and shelf life is becoming the new bargain battery.


The "heavy duty" or "super heavy duty" batteries contain zinc chloride in their electrolyte in addition to ammonium chloride. Many of the super heavy duty batteries also have an iron case surrounding the zinc can, so the battery does not become soft as it discharges and the zinc dissolves. These batteries cost about USD 0.20 per AA cell but have a capacity of about 900 mAh. There are an abundance of "heavy duty" battery brands, with many manufacturers of cheap devices using them in those devices. These contain much the same chemistry and construction as general purpose Leclanche cells.


Alkaline batteries, although they use a very similar chemical reaction, have different properties and a different construction. The electrolyte is an alkaline solution of KOH. This is better because basic solutions do not corrode zinc, prolonging the shelf life. The case is also made of iron instead of zinc; the zinc is stored inside the battery in a powdered form, which enables the battery to produce a higher discharge rate. Alkaline batteries cost about USD 0.50 per AA cell and have a capacity of about 3000 mAh at low current draw (more like 1200 mAh at high current draw). To obtain manganese dioxide or zinc, it is cheaper and better to get heavy duty batteries. Therefore alkaline batteries are not a good source of chemicals.


Lithium iron disulfide batteries have much better quality. Lithium is highly energetic and lightweight, making for a high power battery. The electrolyte is an organic solvent because lithium reacts with water. The electrode materials are spiraled within the shell to make a jelly roll construction, which allows high current draw (much higher than alkalines). This is their main advantage. The durable nickel-plated steel case allows for a shelf life of 10-15 years. Lithium batteries cost about USD 1.80 per cell and have a capacity of 3000 mAh at any current draw. They are an excellent source of small quantities of lithium for an element collection. The iron disulfide is full of carbon impurities and does not function as a good source of sulfides. It could probably be thrown into HCl to get hydrogen sulfide gas, but there are cheaper sources of sulfides available.


Lithium manganese dioxide batteries are more common in smaller sizes. They can take high loads but they function much better under small loads. With a shelf life of 10 years, they are quite cheap and perfom well in watches and CMOS units. They are also a good source of lithium but small coin cells can be an expensive way to acquire large amounts of the metal. The electrolyte is the same as in the Li - FeS2 batteries. They cost about USD 1.50 and have a capacity of 1500 mAh. They have a nominal voltage of 3 volts, so their capacity is equivalent to 3000 mAh at 1.5 volts. They have a displacement of 7.83 cm3, while AA cells have 8.34 cm3 displacement. This gives them a slightly higher capacity than most other batteries despite their smaller size.


Lithium thionyl chloride batteries have the highest capacity of all common batteries in use today, as well as the largest tolerable temperature range. They can only take low loads but some of them have been proven to last over 25 years while still retaining charge. Lithium is the active metal, making these batteries very lightweight. Thionyl chloride is mixed with a lithium salt and carbon powder to form the cathode. These batteries are not a good source of thionyl chloride or lithium because of this. Besides, they are very expensive. They cost about USD 3.00 for a AA cell and have a capacity of 2400 mAh. They have a nominal voltage of 3.6 volts, so their capacity is equivalent to 5760 mAh at 1.5 volts.


Zinc air batteries are unique in that they get their cathode from the oxygen in the air. They only contain zinc metal and come in very tiny sizes, so they are only useful for hearing aids and not for chemistry.


Silver oxide batteries have some of the highest energy densities outside of lithium batteries. They have a silver oxide cathode (mixed with carbon powder) and a powdered zinc anode. They are very expensive but can be a source of silver. However, they only come in very small sizes.


Lead acid batteries, often seen in large packs of 3 or 6, rely on a chemical reaction between lead and lead dioxide in the presence of sulfuric acid to generate large quantities of electricity. A lead acid battery is not the ideal source of lead (fishing sinkers are much cheaper), but lead, its dioxide (which is a powerful oxidizer) and sulfuric acid can be extracted from a fully charged lead acid battery. Discharged ones only contain white lead sulfate crystals, hardly soluble in anything. Lead acid batteries produce 2 volts per cell, a little higher than the 1.5 volts produced by most other batteries. A sealed lead acid battery available on Amazon costs USD 50.00 for a 2300 cm3 battery with a capacity of 18000 mAh at 12 volts, equivalent to 144000 mAh at 1.5 volts. If it were AA size, rough calculations show that it would have a capacity of 520 mAh at 1.5 volts while costing $0.18. This shows that lead acid batteries are quite powerful and economic, which is why they are so prevalent in automobiles. Considering that the technology was invented in 1859 (before any other consumer battery), lead acid batteries are surprisingly good when compared to a modern heavy duty carbon zinc battery which costs just as much but is not rechargeable, has a shorter working life, and has a much lower maximum current draw.


Nickel-cadmium batteries were the next rechargeable batteries to become widespread. They use an anode of cadmium, a toxic heavy metal which is difficult to extract from the battery. The cathode is nickel oxide, which also does not come out in a pure form. The electrolyte is potassium hydroxide, which can be economically dissolved out of old batteries. Other than that, these batteries do not offer much in the way of chemistry. They can take large currents and are quite tough but generally have a low capacity and high price. A typical Nicad (as they are called) costs about USD 0.80 for a AA size battery and produces 1000 mAh at 1.2 volts, equivalent to 800 mAh at 1.5 volts.


Nickel-metal hydride batteries dispense with the toxic cadmium and use hydrogen stored in a chemical form on metal as the anode. These anodes contain a large mixture of various compounds and are not useful for chemistry. However, due to their high energy capacity, they often catch fire in air and can startle an unsuspecting battery dissection crew. A typical AA cell costs USD 1.80 but produces 2500 mAh at 1.2 volts, equivalent to 2000 mAh at 1.5 volts. Advances in technology means that these batteries can hold current for many months and complete many discharge cycles, definitely justifying their high price.


Lithium-ion batteries, instead of using elemental lithium, use a lithium ion - containing material to produce electricity. I have not found any useful chemicals in lithium ion batteries, but they are good for electronics due to their light weight and high power. A typical AA lithium-ion battery costs about USD 3.00 and produces 3.7 V at 900 mAh, equivalent to 1.5 volts at 2220 mAh. Another battery I tried got only 1400 mAh for a AA 1.5 volt equivalent. Oh well, I though lithium ion had the highest capacity of all batteries. It must be their extremely light weight that gets them their high energy density. And they are rechargeable. It is understandable that rechargeable batteries have a lower capacity than non-rechargeable ones, as well as a higher cost.


Thus ends my overview of common consumer batteries. I may post pictures in the future.


For further learning:


  • Battery Energy Density
  • Comparing rechargeable batteries, must keep price in mind
  • A similar battery comparison site, I didn't source anything from here
  • Company selling high-end thionyl chloride batteries
  • Energizer lithium battery technical "data sheet"
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September 13 2012 5 13 /09 /September /2012 18:42


Thermite generally refers to a mixture of a finely powdered metal and metal oxide in various ratios. Depending on the components, these mixtures can burn smoothly and slowly or explode with a bang (well, almost). Getting the components of thermite is quite easy, and there are several ways to do this.


First of all, pick your metal. Metals that have been used include aluminium, magnesium, zinc, and even ferrocerium (think camping firesteel). The size of the particles can range from coarse shavings (2 mm long) to a powder 100 times finer than white flour. The former are much less reactive and therefore much safer, while the latter have much more power and can be more interesting. Magnesium is the most reactive of these metals, followed closely by ferrocerium, then followed by aluminium and finally by zinc.


Then, obtain the metal. Aluminium powder is easily available in many sizes from online sellers. Here is 1 pound of 30 micron (very fine) aluminium powder: http://alphachemicals.com/inc/sdetail/165. For a smaller amount, here is 1 ounce of 200 mesh aluminium powder: http://www.elementalscientific.net/store/scripts/prodView.asp?idproduct=1363. Course magnesium turnings are easily obtained: http://www.mcssl.com/store/gallium-source/magnesium-metal/044---magnesium-shavings-454-grams. Magnesium powder is a bit more difficult to obtain (due to shipping regulations, I believe) but still is easily sourced. Here is 40 mesh magnesium powder: http://www.elementalscientific.net/store/scripts/prodView.asp?idproduct=2173.  Zinc powder is also sold around the interwebz: http://www.ebay.com/itm/Zinc-Powder-1-lb-325-mesh-Purity-99-10-/190353593264. Ferrocerium shavings are too dangerous to sell, so they can be easily made. Find a knife, file, or similar object and preferably a low-quality Chinese firesteel, such as the ones found on bargain magnesium firestarters. Good quality firesteels like Swedish firesteel are extremely flammable and light at a very low scraping speed. This is bad because 1) it contaminates your metal powder mixture with metal oxides or 2) worse, it ignites the entire pile of shavings you are making, leaving you nothing to show for your work. Be always prepared for spontaneous ignition, and do not try to put out the resulting flame with water. To get zinc powder, open a fresh alkaline battery and get the center grayish part out. Wash the paste with water, then filter it and allow it to dry. This is impure zinc powder. To get zinc filings at home, take a piece of zinc casing from a carbon-zinc battery and apply a file or sandpaper to it. Zinc filings will be produced by the action. Since zinc is a harder metal than magnesium, it will be more work to create a pile of metal which is much less reactive than magnesium anyway. So magnesium is probably your best bet for tiny homemade thermites.


To get magnesium, get a preferably medium-quality magnesium fire starter (I use Coleman’s brand) and begin filing or sanding down one end. Ultra lightweight magnesium filings will begin to collect on the tissue below. Periodically dump them into a small storage container.  Once you have reached your desired amount (which is about 3 grams for most thermites), cap off the container to prevent dispersal and store in a dry area. An expensive fire starter may make a nicer thermite but I have not tried due to price concerns, while some YouTube videos have shown some of the bargain fire starters being unable to light up.


Then we need an oxidizer. Thermites all use some form of metal oxide to act as the oxidizer in the reaction, thus negating the need for atmospheric oxygen. Many metal oxides can be used, including iron(II,III) oxide, iron(III) oxide, copper oxide, manganese dioxide, zinc oxide, titanium dioxide, and even finely powdered sand (silicon dioxide). Exotic thermites using vanadium pentoxide (which is very toxic) as well as molybdenum trioxide also work, but they are very expensive and performed for the sole purpose of isolating the resulting metal in an ultrapure state. Other thermites, such as lead dioxide (very toxic) and silver oxide (astronomical price: http://www.ebay.com/itm/Lab-Chemical-Purified-Silver-Oxide-Ag2O-10-grams-/320976242642?pt=LH_DefaultDomain_0&hash=item4abbacc7d2), are extremely violent and powerful. The iron oxides are probably your best thermite oxidizers. They have a high reduction potential, meaning that they make a reliable and energetic thermite. They are also dirt cheap (literally, they are some of the most common minerals on earth) so obtaining them can be as simple as dragging a strong NdFeB magnet (wrapped in a tissue) through beach sand and collecting the magnetite particles. These must be thoroughly crushed and mixed between two metal plates for best use. If you do not have a source of magnetite around, you can buy some easily (http://alphachemicals.com/inc/sdetail/147), or you may want to use hematite, another form of iron oxide. Artificial hematite, which is more reactive than magnetite, can be easily made using salt, electricity, and any small iron object. Dissolve salt in water and attach two iron objects to the sides of the container. Run any voltage higher than 3 volts DC and lower than 24 volts DC through the iron objects, with the larger object preferably connected to the positive and the smaller one connected to ground. All kinds of green gunk will begin to collect at the positive object, some of it turning to brown after a while. The negative object will just bubble. Run this for as long as you like and then filter the slop through a coffee filter. When it is completely dried, it will be completely red. Put the crumbly chunks of red iron oxide in an aluminium foil sheet and heat them in a kitchen oven at 350 F for half an hour to dehydrate them, then pour them into a container. Red iron oxide can be easily purchased, too (http://alphachemicals.com/inc/sdetail/225


To make copper oxide, replace the negative iron object with a copper wire and the positive copper object with a large number of copper wires tied together at the top or a piece of copper. Run the electricity through as long as you desire. A yellow-orange solid will form throughout the container, turning green at times. Some blue will also be noticed. Then filter all of the copper pieces and copper oxide paste. Let it dry and heat it in the oven at the maximum temperature until it is completely black through and through. If it does not get completely black, make an aluminium foil dish and carefully place the powder in it, then heat it on the stovetop. Pour the black powder into a container. Copper oxide can also be purchased (http://alphachemicals.com/inc/sdetail/175).


To get manganese dioxide in an impure form, rip open a fresh alkaline battery (look on YouTube for instructions). Keep as much of the black gunk as you can. The gray inside stuff is not used in this step. The other metal oxides are not practical for low-price high-energy thermites, but they can be purchased [website]. Pure manganese dioxide can be found at http://www.ebay.com/itm/MANGANESE-DIOXIDE-1-lb-Pound-Lab-Chemical-MnO2-Ceramic-Technical-Grade-Pigment-/190634439641?pt=LH_DefaultDomain_0&hash=item2c62b2b3d9. Other manganese oxides are also used in thermites, but they are less reactive (http://alphachemicals.com/inc/sdetail/211). 


Mix the thermite by grinding, where you can get creative. I sometimes use the tail of a pen to grind together the finely powdered metal oxide with the magnesium shavings, thus getting an even mix. Remember, the more finely ground and thoroughly mixed the thermite is, the better, hotter, and brighter it will burn.


To light your thermite, pile your thermite on a brick or similar surface and find an ignition method. You need to use a flame source hotter than butane (cigarette lighter). It is recommended to use short (1.5 inch) lengths of narrow magnesium ribbon to light the thermite. If the thermite is composed of an unreactive oxide (titanium dioxide, zinc oxide), you can consider placing some potassium nitrate crystals (Spectracide stump remover @ Home Depot) around the base of the magnesium ribbon. Bend the ribbon at the base so it can stand upright, then pile the thermite mix around it. The magnesium ribbon can be ignited by a butane lighter and a quick evacuation of the immediate area can be performed. The thermite should ignite as soon as the magnesium burns down to base level. Another common ignition system is potassium permanganate / glycerin. Both of these chemicals can be purchased easily online. When they are mixed, spontaneous combustion occurs after a waiting period of a few seconds to a few minutes. The thermite should ignite after the glycerol begins burning. Firework sparklers can also be used for ignition. Another method is a magnesium fire starter. Make a big pile of coarse magnesium shavings on top of the thermite mix. Use the provided ferrocerium to set the shavings on fire, and quickly move back. It is recommended to use heat-resistant gloves when doing this process. Another novel method is to place the thermite in a dry brown leaf (if you live in a deciduous area) and place some magnesium filings alongside the thermite (and overlapping it) on top of the leaf. Char the leaf – magnesium junction with a high-power magnifying glass. The black leaf char will absorb the heat, igniting the magnesium filings which will ignite the thermite. Use eye, face, and hand protection and back quickly away once ignited.


Warning: Iron thermites are very nontoxic to the environment and to people but they burn very hot and produce ultraviolet light. Do not look directly into a large number of thermites. Copper and manganese thermites are a little more toxic but can be disposed of normally. Beware of residual embers which may flare up when the ash is disturbed. Large thermites can explode if sprayed with water in an attempt to extinguish them.


Have fun with these fiery little reactions!




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August 19 2012 1 19 /08 /August /2012 02:28
There are several ways to make invisible ink, but I decided to try a dye that responds to changes in pH. Phenolphthalein is a white crystalline substance which is soluble in alcohol. Solutions of it are colorless at a pH below 8.2 and magenta or pink at a higher pH. (Wikipedia says that it turns colorless above pH 13 as well.) Sodium carbonate solutions have a pH of around 11.6 for a dilute solution, making it the perfect trigger for phenolphthalein. (It is easily made by heating baking soda or sodium bicarbonate in an oven at 300 degrees F.)
To make the invisible ink, I was originally planning to paint the phenolphthalein solution on with an eyedropper. Revealing it would be done by a spray mist bottle filled with sodium carbonate solution. However, my eyedroppers were dirty and we had no excess spray bottles, so I decided to use paint brushes. I used a fine tipped paint brush for the phenolphthalein solution and a 1 inch brush for the sodium carbonate solution. Phenolphthalein dissolved in isopropyl alcohol is delicately painted onto paper in the form of writing. It dries in about 30 seconds. Then sodium carbonate solution is swiped over the paper until the invisible writing shows itself. The writing stays visible after drying.
Here is a video demonstration:
Only tiny amounts of chemicals are used, making this a good invisible ink for those already possessing phenolphthalein. Those who do not possess it should look to other compositions, some of which may be discussed here at a later date.
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August 12 2012 1 12 /08 /August /2012 00:20

Warning: Bromates are suspected carcinogens. Limit exposure to them. They are also strong oxidizers. Keep away from combustible materials. Bromine is smelly and toxic in large quantities. Perform this reaction outdoors.


You will need: Source of potassium (No Salt salt substitute, Spectracide stump remover, etc.), bromide reserve for spas (sodium bromide), a carbon rod, a wire of any sort, a five volt power supply, and some connecting wires.


Boil some water and dissolve quite a bit of sodium bromide in it, though not enough to make a supersaturated solution when cooled. Anchor the carbon rod and the bit of wire to the sides of the chosen electrolysis container and pour the hot sodium bromide solution in. Attach the positive electrode of the power supply to the carbon electrode and the ground to the wire. Begin electrolysis.


A brown solution of elemental bromine will begin forming at the carbon electrode (anode), while a colorless solution of sodium hydroxide will begin forming at the wire (cathode) (net reaction 2 NaBr + 2 H2O --> 2 NaOH + H2 + Br2). The hydrogen escapes as a gas from the cathode. When sloshed around, these two solutions will react to form sodium hypobromite and sodium bromide (2 NaOH + Br2 --> NaBrO + NaBr + H2O). Because of the heat, the sodium hypobromite will disproportionate into sodium bromide and sodium bromate (3 NaBrO --> NaBrO3 + 2 NaBr). All of the sodium bromide will reenter the reaction at the beginning.


When bubbles begin forming at the anode at a significant rate, this means that much of the sodium bromide is exhausted. Stop the electrolysis and filter the solution to remove the carbon particles. What remains is a light yellow solution of sodium bromate, bromide, and hypobromite. It smells just like household bleach, just with chlorine instead of bromine. Heat it to disproportionate the majority of the remaining sodium hypobromite, then let it cool. When crystals begin forming (it may take several days of evaporation), add some potassium salt solution and stir to dissolve the crystals. They should dissolve but new crystals should take their place over time. These are crystals of not-so-soluble potassium bromate. Wait until no more crystals form and then remove them. They should be quite pure and white. The remaining solution contains some bromate in solution as well, but it is quite impure. When I tried this, the yield was not spectacular, but the resulting crystals were good. You can use them as an oxidizer or a source of reactions involving bromine oxyanions.



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August 6 2012 2 06 /08 /August /2012 21:33




In this article I will describe the history of my personal experiences with chemistry. My extracurricular study of chemistry has covered much ground since I started performing experiments about four years ago. Like many people, my very first experiments involved kitchen chemistry. Strangely, I never performed any kitchen chemistry reactions until my teenage years. Before that, I only read about chemistry. Obtaining chemicals of any form always seemed like a distant fantasy, so I contented myself with reading science books and dreaming about myself as a scientist. Using these books, I learned about the periodic table and memorized many of the elements and their properties and uses. In one of the books, I noticed a description for a homemade hydrogen generator and became fascinated by it. Producing my own flammable “helium” balloons from scrap metal and a common acid had to be magic. However, we did not have chemicals such as muriatic acid and zinc strips, so I could only theorize. This was tacitly encouraged by my parents, who like all parents did not want me hurting myself with caustic chemicals but did not mind me spending all of my free time studying and learning. However, I had become curious, and it is hard to stop a curious child.


I accidentally discovered a chemical reaction; baking powder fizzes when mixed with water. This fascinated me, as I had believed that chemical reactions can only be performed in labs and that they were all dangerous. I wanted to step it up a little, though. I had heard much about the reaction between baking soda and vinegar from science books and decided to try that. I was even more surprised at the reaction than I was at the previous reaction, but, illustrating my current chemophobia, I still believed that the resulting product, sodium acetate, was toxic because it did not have a household name. I assumed at that time that I had completed all of the reactions possible with household substances as my parents told me that everything else was too dangerous to play with.


To increase my chemical knowledge, I began searching Encarta® Encyclopedia for chemicals. The small amount of information present about chemistry only made me search harder than ever, so I then decided to search online for random chemical compounds that I saw in textbooks. I soon found that Wikipedia was almost always the top result in these searches. I decided to check out the extent of Wikipedia’s chemical information and was dumbfounded by the number and size of the articles on chemicals and chemistry. For several months after this, most of my free time was spent taking painstaking notes from Wikipedia chemical articles and learning as much as I could about the few chemicals and elements I knew. I filled several notebooks with chemical information, drawings, and data about basic inorganic compounds and elements. Despite being able to find excellent data on the chemicals, these articles only described laboratory methods of synthesis, so I still did not have an idea on how to make any of these mysterious substances which were so different from the materials I had seen all of my life. Because I was planning to take chemistry in my junior year, I had purchased a chemistry textbook. To further increase my knowledge, I thoroughly perused that book, reading through it about a dozen times. Because I read the book so many times before taking the course to glean every bit of chemical information I could get, I got almost a perfect A+ with practically zero effort. That was only the beginning of the rewards I got for pursuing amateur chemistry as a hobby.


My introduction to genuine chemistry experimentation itself was gradual. I became fascinated with electrolysis over time. Since I heard only water and electricity were necessary to obtain an explosive mixture of gases (including hydrogen, my earlier dream), I decided to give it a try. I was disappointed with the performance of pure water. Even at 24 volts DC, the electrolysis occurred so slowly that I could hardly notice anything. I then read that electrolytes could be used to speed up the reaction. I added vinegar and used longer copper electrodes to provide more surface area. A green coloration appeared in the solution and bubbling was happening quite rapidly. The solution heated up and the smell of vinegar became strong. I added salt and the reaction improved even more. However, I noticed that all of the gas was being produced at the negative electrode while nothing but green slime was being made at the positive electrode. This colorful slime got the attention of my sister, who began a course of electrolysis experiments on her own. She used salt and copper and got “yellow water” (microcrystalline copper(I) oxide), while baking soda and copper got “green water” (basic copper carbonate), and Epsom salts and copper got “blue water,” (copper sulfate) named because of the colors produced. I became curious as to the composition of the precipitates/solutions and how they were formed. I then discovered that iron makes a green-gray precipitate which turns brown over time, and this metastable substance only made me more interested. A science kit I received as a gift required copper sulfate for most of the experiments in the chemistry section, and I was desperately looking for a source of this material. When I tried to form “blue water” and obtained a white precipitate at the cathode, I decided to research this reaction. I discovered that magnesium hydroxide was being formed, and the magnesium was being replaced in solution by copper. The solution turned bright blue and I stopped the electrolysis only when copper was falling out at the cathode as fast as it was dissolving at the anode. Upon evaporation, I got a blue solid that turned green on exposure to air (it lost some of its pentahydrate). I performed some experiments with this copper sulfate but was somewhat disappointed that it was only about 15% pure. Another more successful experiment I did was the heating of sodium bicarbonate to obtain sodium carbonate, also in the science kit’s chemistry section. This was done in a kitchen pot which was blackened on the outside much to my mother’s disapproval, although I kept trying to persuade her that the pot was fine for cooking. This began my synthesis craze.


I continued expanding my Wikipedia research and performing new syntheses of compounds. I electrolyzed iron in Epsom salts to form iron(III) sulfate. I electrolyzed lime in water and nothing happened. I also began finding more household sources of chemicals. I extracted potassium hydroxide from an alkaline battery and (unknowingly) converted it to potassium carbonate by long standing in air. I created sodium hydroxide from the chloralkali process and zinc acetate with modern pennies and vinegar. I formed copper(II) oxide by electrolyzing copper in sodium bicarbonate solution and heating the precipitate formed, although I did not know what to use it for. My room was constantly filled with new experiments almost exclusively in the aqueous state. Because I had no source of fire, my only way to form crystals was to evaporate in an open pan. This made my bedroom a very humid place, filled with all kinds of strange metallic smells. This caused my parents to relocate my chemistry lab in the basement, where I realized that storing chemicals in the individual compartments of an egg carton (my portable chemical shelf at the time) is inferior to storing them in closed paint mixing containers. However, I still had only a small amount of starting point chemicals. Then came my first glimpse of a genuine chemistry kit.


My pastor had an old chemistry kit belonging to his son and kindly donated it to me. The kit was designed for a middle-school scientist and being a junior in high school I found the experiments themselves quite uninteresting. The chemicals, however, were a completely different story. Chemicals that I have never heard of – phenolphthalein, protease, and copper(II) chloride – as well as those that I have read about but never expected to see until college – cobalt chloride, copper sulfate, zinc sulfate, silver nitrate, ammonium chloride – I now had in my possession. I began to do a huge array of experiments involving these new chemicals, wasting most of them. Suddenly, formerly abstract reactions from Wikipedia became intriguing realities as I continued to explore the wonderful world of chemistry. A significant number of glassware also came with the kit, which turned out to be the most useful portion of it. I still lacked many basic chemicals, but finally a school project gave an excuse to purchase them.


Due to my chemistry craze, it is no wonder that I decided to use chemistry for my science project, which was entitled “The Catalytic Decomposition of Hydrogen Peroxide.” I prepared a huge list of catalysts that would be mixed in equal weights (if possible) with a fixed volume of hydrogen peroxide. The oxygen production after a certain length of time is measured and recorded and the results compared. I intentionally added certain catalysts such as tincture of iodine, ammonia, and muriatic acid for the dual purpose of observing their catalytic effects and adding such valuable reagents to my chemical stockpile. The science project was a success, with several general trends visible on the periodic table and a whole new host of experiments opened to me. I was finally able to perform the cobalt chloride equilibrium with hydrochloric acid in my textbook and numerous other experiments, as well as restock critically low chemicals like copper(II) chloride (by dissolving CuO in HCl). These colorful and wonderful reactions made me think of using my camera to begin recording reactions. My first pictures were poor quality, but they steadily increased in quality as I learned how to take macro pictures. I then got a new camera and began taking videos as well as high-resolution pictures. Some of my first photographs were of flame tests, which were done on my parents’ old but only kitchen stove. I remember promising myself never to burn lead powder on the kitchen stove anymore as it produces a spectacular grayish flame but toxic red smoke. However, the only reactions I knew were only the ones that I stumbled upon or found on Wikipedia. Therefore, I did not know how to dissolve metals like lead, as I never suspected a 1 M solution of a very weak acid like vinegar to dissolve a relatively inert metal like lead with any reasonable speed. I heard the Romans had to wait months for their lead pots full of vinegar to accumulate enough atmospheric oxygen to produce a significant amount of lead acetate, which they used as a calorie-free sweetener. However, there were many metals that I could dissolve, many precipitates and complexes to be formed, and many mysteries to be sorted out with the unknown compositions of impure household products like magnets.


Ever since, I have expanded my range of experiments steadily. I found the wonderful forum Sciencemadness, and discovered a whole world of amateur experimentalists who have created processes for synthesis that were designed for chemists like me. For example, using some of this newly found information, I purchased bismuth and dissolved it in hydrochloric acid and hydrogen peroxide. I was quite happy to finally have bismuth chloride from which I could do the hydrolysis experiment featured in my chemistry textbook. I also extracted lithium (my first sight of a metal more reactive than aluminium) from a lithium battery and reacted it with water. I also found W. Oelen’s great site, which has a large number of riveting experiments, complete with high-quality videos and stunning photographs. I discovered how to dissolve lead (hydrogen peroxide and acetic acid) and did several experiments with it. The iodide of lead was an especially intriguing compound. (You can see some of my experiments on the YouTube channel LanthanumK or on this blog.) I also found out that placing a magnifying glass in front of my camera lens while taking photos drastically increases the quality and zoom level of the picture or video. I use this tactic for most all of my macro shots now. But all of my experiments have not been good.


I have gotten several temporary bans and restrictions on my experimentation due to chemistry epic fails. I electrolyzed tincture of iodine in my room, and the heat of the electrolysis evaporated much of the iodine and filled my room with fumes of iodine when I was sleeping. I woke up the next morning with a headache and a foggy feeling. That resulted in my chemistry being relocated again to the lower level of the house as it had gradually shifted back into my bedroom. Halogen strike one. Once I placed the manganese dioxide from a dead battery in hydrochloric acid. Very little chlorine was released, and it was barely noticeable in the indoor environment. Incorrectly deducing that a dead alkaline battery equals a charged one, I placed the manganese dioxide from a charged battery in hydrochloric acid in my mother’s hobby room (where I did my chemical experiments at the time) and filled the room with chlorine gas. Unfortunately it was my mother who walked in before me and got an unpleasantly strong but not suffocating whiff of chlorine gas. The solution was placed on top of a white surface, which was covered with black manganese dioxide stains after that due to splattering during the reaction. This resulted in a several-month break, my halogen strike two. Then I made some bromine water with bleach, sodium bromide, and hydrochloric acid, and poured it down a sink drain without dilution after some experiments. My sister bent over the sink and got a whiff of bromine vapors that allegedly gave her a headache. This was halogen strike three. I now respected the halogens much more than I previously did. Fortunately, my parents did not know that I made the same error, releasing halogen vapors indoors, three times in a row. I hope I get to work with fluorine someday, as this is the most fascinating of the halogens. Another time I drilled a hole in a lithium battery extracted from an EZ-Pass device. I should have researched why the battery voltage was 3.6 V instead of 3.0 V (like a normal lithium – manganese dioxide battery) before opening. Anyway, I drilled out the center carbon electrode and removed a smoking drill bit along with a smell of burning matches. I thought that I had caused a short and that the battery would catch on fire and explode or some bad thing, so I dunked it in water (bad move with any lithium battery). Fortunately, only a small amount of thionyl chloride – yes, this was a lithium thionyl chloride battery – could leak out at a time, but I still ended up stinking up my parents’ kitchen with a sulfur dioxide / hydrogen sulfide smell. After this, I decided to dismantle a partially charged nickel metal-hydride battery. I peeled back the coating without shorting out the jelly roll and took out the electrode materials, placing them on a metallic surface. This metallic surface was both my saving and my undoing. What happened was partly due to the metallic surface forming a short, aerial oxidation of the hydride material, and improper contact of the electrodes. I was busy somewhere else in the room when I noticed a light coming from the battery and looked up just in time to notice the battery igniting. Not knowing what would happen if the battery was allowed to burn, I immediately smothered it with a metal case of drill bits. (Out of the two more NiMH batteries I opened, two of them ignited upon opening as well.) Due to these types of experiments, my chemical lab ended up in the garage, where I accidentally vaporized some selenium and filled the garage with a fecal smell that I was sure was toxic. Heating the deposited selenium on the floor in an attempt to vaporize it did not help matters. This did not result in a ban as I was able to clean it up with bleach, a selenium oxidizer that can handle gray as well as red selenium. Another time I wanted to see whether the sodium atoms from celery would be visible in a flame test and filled the kitchen with a horrible smell of burning celery, which caused my parents to strongly discourage my use of the kitchen stove.


During one of these bans I wondered how many pure elements I had. I counted quite a few that I had on hand and quite a few that I was sure I could make. Thus began my element collection. I rigorously searched household materials for elements until I accumulated about 50 of them, then began purchasing and isolating a few. I have spent about 80 dollars on elements so far, mostly from Gallium Source. Contact me and I will give you my 90-page description of my element collection. You could also check in the archives of this blog for articles on how to obtain elements. I will give you tips and tricks on how to obtain individual elements if you ask me.


I decided to begin posting my experiments online on this blog to gauge reader’s responses and interest. I got a significant number of page views (over ten thousand so far), much more than I got for a generic blog that I also worked on. For a while I was posting two articles a day, but declined lately due to work and school. I am trying to work on posts when I get the time and hope to be coming out with more posts in the future.


I learned several lessons over my years in chemistry. One is “research before reacting”. A well-informed chemist can make split-second decisions regarding a runaway reaction or a smelly one without having to worry about uncertainty. When my sodium bromate reacted with hydrochloric acid to produce a cloud of brown bromine fumes, I immediately took it outdoors because I followed this rule. Another would be to regard other people’s noses. That acrid smell that you may not find annoying may irritate other people, and that is not an option. A third would be to use small quantities of reagents at first. Using large amounts of chemicals and accidentally contaminating or hydrolyzing it can be one of the biggest disappointments, especially if the chemicals were expensive.


I hope my readers learn something from this rehearsal of my experimental history. Be sure to stay safe but not sanitized. Experimental chemistry needs a touch of the unknown, puzzling, beautiful, dazzling, and even the noisy to be an interesting hobby for most ages. It cannot be a computer simulation, no matter how good they are. For me, it has guided me in the easy choice of a major, aced me through several classes, saved much study time, given me inspiration for my blog and YouTube channel, and decided the way I have spent much of my free time. Enjoy studying the building blocks of the universe at your desired level.


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July 14 2012 7 14 /07 /July /2012 21:15
Flash powder is a pyrotechnic mixture that, when ignited, burns with a bright flash. The reaction happens so rapidly that the powder seems to disappear after ignition. This mixture is often used in fireworks due to its ability to explode when confined. The most common flash powder is aluminium – potassium perchlorate mixture. However, perchlorates are not readily accessible. A less spectacular but more available composition is magnesium – potassium nitrate.
Course magnesium powder can be easily made from a file, muscles, and a magnesium bar. I obtained my magnesium bar from a camp fire starter, which generally costs about 10 dollars. There may be a cheaper source for large magnesium chunks. The shavings are approximately 100 microns across when ground properly. Potassium nitrate, also known as saltpeter, is also readily available. Some fertilizers use this ingredient to supply nitrogen and potassium, two essential plant materials. It is also an important ingredient in black powder. The best source of pure potassium nitrate that I have found is Spectracide® stump remover. This comes in fine white crystals which may be ground further if desired but still make a good flash when whole. A one pound bottle of the stuff costs about six US dollars.
I mixed the potassium nitrate with the magnesium powder in a 1:2 volume ratio. There is an excess of magnesium in this reaction, but that is fine. (If you want to make a more powerful flash powder, use stoichiometric quantities and fine powders.) The reaction of the two chemicals is probably this one:
5 Mg + 2 KNO3 => K2O + 5 MgO + N2
My first flash powder experiment was a success. I ignited it with a magnesium fuse and it burnt with an extremely bright flame, producing a cloud of white smoke. The nitrogen gas blows the reaction mixture apart, making the chance of embers quite small. The heat is intense, however, and it could easily ignite a flammable base, e.g. grass. Caution must be observed because of the heat.
Then I tried igniting flash powder in a test tube. I was curious whether the containment of a test tube would cause the flash powder to explode. To do this, I created an extremely long-winded and fragile ignition setup. First of all, I made a titanium igniter. Titanium has a higher resistance than quite a few other metals and so heats up easier when current is applied. Even better, a thin titanium strip will ignite when heated. Wires are tied to two ends of a thin piece of titanium foil and the wires are hooked up to two nine-volt batteries in series. The load resistance is about 1 ohm at first. This means a huge current spike as the nine-volt batteries heat the titanium, which decreases as the titanium becomes red-hot and is cut off as the titanium ignites and burns to the non-conductive titanium dioxide.
I used this titanium to ignite a pile of coarse magnesium powder and calcium pieces. The calcium was added in the hope of increasing the strength of the burning magnesium powder. Then a wooden splint, one half of which was soaked in potassium nitrate solution, was placed in the bottom of the powder pile. On top of it laid a magnesium ribbon which ran directly into a test tube filled with the flash powder mixture. The test tube was on its side.
It did not work as I expected. The ignition pile ignited when the current was applied, but I used the file to make the majority of the pile and it blew away before igniting the magnesium ribbon. The calcium did not have any significant effect. The wooden splint ignited however, and the nitrate – soaked section burnt up. The magnesium appears to have oxidized, although the bright white flame of magnesium is not seen. Once the splint passed into the unsoaked portion, the flame extinguished due to the wind and the heat removal by the brick. The reducing fumes produced by the extinguishing put out the magnesium for good, about 2 inches from the flash powder.
I decided to remove everything except the magnesium strip and light it using a propane torch. The test tube was standing upright in a holder. The force of the flame from the propane torch melted and bent the strip over the edge of the test tube, invisible to me. When I applied more heat, it ignited in the middle and instead of burning toward the flash powder, it burnt along the outside edge of the test tube, making a black charred spot. I took video of this but since it is boring the video was deleted.
I then laid the test tube on its side and tried ignition. The magnesium burnt up and hit the flash powder inside the test tube. The test tube was immediately melted and ripped open on the top, leaving just a small chunk of black plastic behind. The flash powder burnt brilliantly, even in the broad daylight, throwing flame and sparks around for an entire second. If the powders were fine it would have been gone almost instantly and the test tube would have exploded. However, due to the large number of fails, I forgot to turn my camera on and my biggest flash powder combustion event went unrecorded.
My brother got interested in flash powder and decided to make some himself. He ignited it with a wooden splint. He said that it was hard to see for a minute after ignition.
I then assembled a flash powder module. A small amount of flash powder is placed on some duct tape and a titanium igniter is placed inside. It is ignited using an 18V nickel – cadmium battery pack. The duct tape melts by the igniter, but the flash powder does not go off. I make a slit in the duct tape and the flash powder ignites. Therefore, the ignition of the titanium is necessary for the flash powder to deflagrate. To help this out, I wrap the flash powder in filter paper. The wires are connected to a long power cable which is hooked up to the 18V battery pack. A water bottle cap is placed on top of the paper construction to determine the effects of the blast. Ignition happens instantly, and the burning water cap is thrown 5 feet.  This shows the danger of being too close to the deflagrating mixture. Small amounts of smoldering paper are left behind.
At this stage, my flash powder modules are becoming too close to fireworks, which are illegal to use in NJ. Therefore, I will quit my experimentation with Mg – KNO3 flash powder for now.
After a couple of months, I decided to make Zn - KNO3 flash powder. I ground up a piece of zinc from a carbon-zinc battery on a file and mixed the fine powder with potassium nitrate. The mixture ignited with difficulty. The zinc tended to melt at first, then suddenly combined with the potassium nitrate, making a dim green flash. This is expected because zinc is a much less flammable and much less reactive metal than magnesium.
Then I tried to coat a wooden splint in glue and magnesium - potassium nitrate flash powder. I hoped that when the splint was ignited, the flash powder would behave as a sparkler. Unfortunately, the nitrogen produced during the flash powder reaction tended to immediately extinguish the wooden splint, making reignition necessary. I hope to try using potassium bromate/chlorate as the oxidizer in the future as it seems to produce much less gas.
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