#12: Burning Magnesium in Carbon Dioxide

If you have ever stared into a fireplace or campfire chances are you have pondered a particular question. That being, what is fire? Technically, most fire is a plasma-like portion of semi-ionized gas. However, what most people really want to know is what it means for something to burn? The answer is somewhat simple. Burning, or combustion, happens when a substance has enough energy to oxidize. In other words, materials burn because they want to bond with oxygen. I explored this interesting concept in an experiment that I performed recently.

To start, I filled two identical flasks with vinegar (dilute acetic acid [C2H4O2]) and baking soda (sodium bicarbonate [NaHCO3]). The reaction that took place produced a nice amount of carbon dioxide (CO2). After corking the containers, I was ready to perform two tests. For the first part, I simply lit a large match and held it over one of the flasks. Instantly, it was snuffed out. In my second test, I ignited a strip of magnesium metal and suspended it within the other flask of gas. Unlike the wooden match, the magnesium continued to blaze brightly until it was completely consumed. To explain these differing results, I looked at the science.

In the case of the lit match, I knew that it ran out of oxygen upon entering the CO2. The question was, why didn't the same thing happen to the magnesium? The answer lies in the fact that magnesium burns with an extremely high amount of energy. In other words, the metal has a very powerful attraction to oxygen and will try incredibly hard to bond with it. It is because of this extreme desire that magnesium is able to rip oxygen atoms straight off of other molecules. Thus, I concluded that, when placed in the flask, the burning magnesium obtained all of its oxygen from the carbon dioxide. I found that to be particularly remarkable.

Even while in carbon dioxide, the magnesium metal burned in a spectacularly bright fashion.

To watch the reaction between magnesium and solid carbon dioxide (dry ice)...

http://www.youtube.com/watch?v=wqErrNvns4o#t=61

If you want to learn more about the science behind fire...

http://en.wikipedia.org/wiki/Fire

#11: Candy Cane Incineration

In the spirit of the holiday season, I decided that I ought to perform a festive chemistry experiment. The question was, what should I do? My answer: incinerate a peppermint candy cane in a test tube of molten potassium chlorate (KClO3). What else? To start, I unwrapped my candy cane and ate the top portion. Next, I scooped a small amount of potassium chlorate powder into a test tube. This would act as my oxidizer. After holding the tube over my butane burner, I was able to melt the powder into its more reactive liquid form. With the test tube securely in place, I dropped the half-eaten candy cane inside. 

As soon as the spiraled segment entered the liquid, there was a miniature explosion that caused it to shoot out like a mortar shell. Even though the piece of candy was on fire, I still had my potassium chlorate and wasted no time adding another candy cane. This time, the clump of peppermint fuel stayed in place long enough to be completely consumed. At its peak, the sugary inferno resembled a propane torch on full blast. While the pyrotechnics were pretty cool, the best part of the reaction was its wonderful smell of burnt caramel. Merry Christmas!

The vigorous firestorm produced by this reaction shows how much energy is in a simple candy cane.

#10: Shapeshifting Gallium

Of the many fascinating chemicals that I have recently obtained, I have a favorite. It is a shiny low-melting metal called gallium. While you may not have heard of it, gallium is actually the 31st element on the periodic table, located just below aluminum. It has quite a few interesting traits, some of which I plan to explore in future experiments. Gallium is best known for its relatively low melting point of 86 degrees Fahrenheit. This, along with the fact that it is also non-toxic, makes it a great substitute for mercury. In general, liquid metals have many unique properties. Unfortunately, because most metals have extremely hot liquid states, they can't react with other fluids without instantly solidifying. This makes gallium useful in examples of aqueous chemistry such as my most recent experiment. 

I started by placing a small piece of my gallium into a beaker with some water. After heating the beaker for a short time, the metal melted into an irregular blob. Next, I added a few drops of concentrated sulfuric acid (H2SO4). Within seconds, the liquid gallium became a perfect sphere. This was because the newly formed gallium sulfate had a higher surface tension than the metal by itself. In the second part of the experiment, I reversed this effect by adding a small amount potassium permanganate (KMnO4). Just like the sulfuric acid, the permanganate changed the chemical composition of the gallium. However, instead of increasing the metal's surface tension, it lowered it. This caused the gallium to spread out like oil slick. The interesting part came when I poured in more sulfuric acid. This reversed the reaction once again as it caused the layer of gallium to form beads.

Ideally, this is supposed to be an oscillating reaction in which both chemicals are balanced. In this situation, the gallium would constantly contract and relax as the chemicals on its surface changed places. This unique behavior of the gallium is why the demonstration is often called the "gallium beating heart experiment." The process is significant because it shows that by changing a liquid's chemical makeup, you can change its surface tension and thus, the way it behaves.

While the experiment didn't go entirely as planned, the gallium behaved in some interesting ways.

To see the true gallium beating heart experiment...

http://www.youtube.com/watch?v=N6ccRvKKwZQ

#9: Spontaneous Combustion

I recently received a ton of awesome chemistry supplies for my 16th birthday. Along with items such as a butane burner and periodic table poster, I was given my first batch of lab-grade chemicals. Until then, I obtained most substances from the pharmacy or hardware store which greatly limited what I could do. Thanks to my new materials, I have many fascinating projects planned for the upcoming months. For my first experiment, I wanted to start things off with a bang. Literally.

I began by pouring out a small amount of potassium permanganate (KMnO4), which is an oxidizer. I then added a little bit of glycerin (C3H8O3) as my fuel. While the two chemicals would have reacted eventually, I decided to speed things up by adding a few drops of water. Within seconds, the tiny pile went from a smolder to a miniature flash-bang explosion. It was pretty crazy.

Because the reaction involved the burning of glycerin, which contains hydrogen and carbon, it can sort of be considered a form of combustion. Thus, it is no surprise that with the presence of the oxygen-rich potassium permanganate, the result would be explosive. Caution: this reaction is unpredictable in nature and therefore should only be attempted on a small scale with extreme care.

Considering the small quantity of chemicals used, the reaction was quite vigorous.

#8: Hydrochloric Kisses

I was eating a handful of half-metled Hershey Kisses when I found myself struggling to fully unwrap them. Then a thought occurred to me: what if I could remove the wrapping-foil using science? After pondering the question, I came up with an idea. I grabbed a small beaker and headed outside. There I filled the beaker with some hydrochloric acid (HCl) and a little bit of water. With my camera ready, I dropped a fully-foiled Hershey Kiss into the beaker and began to wait. It wasn't long before the shiny candy was writhing and bubbling. Eventually, all of the aluminum surrounding the chocolate had disintegrated into almost nothing.

The science behind what happened was simple. The hydrochloric acid ripped off the aluminum atoms forming aluminum chloride (AlCl3) and leftover hydrogen gas (H2). The hydrogen is what formed the bubbles which means that they were actually quite flammable. In the end, I was left with what seemed to be a perfectly unwrapped piece of milk chocolate. However, the HCl and aluminum that it absorbed probably made it extremely toxic. 

After a few minutes, the wrapping-foil surrounding the chocolate completely dissolved in the HCl.

#7: Pepto-Bismuth

I was in the pharmacy the other day when I saw something neat that I didn't think still existed: Pepto-Bismol. For those who don't know, Pepto-Bismol is a digestive medicine that has been around for quite a long time. The reason I was interested in it was not because I had a terrible stomach ache, but because of the drug's chemical makeup. I had heard a while back that the medicine contained bismuth as part of its active ingredient bismuth subsalicylate (C7H5BiO4). Bismuth is a non-toxic heavy metal that resembles lead in a lot of ways. If you saw my first post about diamagnetic levitation, you'll also know that it weakly repels strong magnets. Anyway, I decided to buy a small pack of Pepto-Bismol tablets with the hope of somehow extracting bismuth from them. After going online and finding the correct procedure, I learned that my hope could easily become a reality.

I started by crushing up the majority of the bright pink tablets in my mortar and pestle. I then placed the neon powder in a beaker and added some diluted hydrochloric acid (HCl). It took a while, but I eventually dissolved the Pepto-Bismol into the hydrochloric acid with a good amount of stirring. After filtering out any leftover solids, I placed a strip of aluminum into the pale pink solution and let it sit for a while. What happened was that the chlorine ions from the HCl let go of the bismuth ions and bonded with the aluminum ones instead. This left the bismuth to simply clump together and form a gray precipitate. I was then able to collect these bismuth particles using a coffee filter. After allowing the powder to dry and putting it into a steel ladle, I attempted to melt it over our stovetop.

Sadly, I never got the shiny silver liquid that I was looking for. Instead the dark dust simply fused together and turned mustard yellow. After thinking about it, I would say that there were most likely too many impurities. As far as the yellow color goes, I suspect that this was due to the oxidation of either the bismuth or the leftover aluminum. I may never know.

Bright pink Pepto-Bismol digestive medicine containing minor amounts of the element bismuth.

Precipitation of solid bismuth
out of HCl using aluminum. 

Pure bismuth metal from my "Diamagnetic Levitation" project.

Impure bismuth sludge.

To perform this experiment yourself...

http://www.popsci.com/diy/article/2012-07/gray-matter-extracting-bismuth-pepto-bismol-tablets

#6: Burning Sulfur

In my last post I mentioned how I sometimes break off chunks of my sulfur specimen to use in experiments. So this time I am going to do just that. While I have done this experiment before, I thought it would be neat to try it again and record it this time. To start, I simply chiseled off a few pieces of elemental sulfur and ground them into a fine powder. After placing this powder on a piece of wood, all I had to do was ignite it. Due to its low melting point, the sulfur actually melts upon contact with a lit match and burns quite steadily. While it may seem in the video below that there is no flame, it is merely too dim to see during the day. Luckily, you don't get to smell the horrible stench of rotten eggs that is produced.

What happened was that the heated sulfur bonded with the surrounding oxygen to produce sulfur dioxide (SO2) and sulfur trioxide (SO3). Both of these products are poisonous and here is the interesting reason why. These gases bond with water molecules in the air, or on the surface of your body, to form both sulfurous acid (H2SO3) and sulfuric acid (H2SO4) on the spot! While this reaction isn't extremely dangerous, it should obviously be performed outdoors or in a fume hood. In my opinion, sulfur is a fascinating element with many unique aspects.

As you can see, sulfur burns somewhat strangely.

If you would like to learn a bit more about sulfur...

http://en.wikipedia.org/wiki/Sulfur

#5: Rocks & Minerals

Even though I stopped collecting them a few years ago, I continue to be fascinated by the many rocks and minerals that I own. I started acquiring rocks on a small scale when I was ten or eleven until I eventually became a serious rock hound. Sadly, due to the area that I lived in, all I ended up doing was paying lots of money to order them online or buy them at stores. I eventually halted my spending spree and moved on to other nerdy things.

Recently, I have taken a new interest in my mineral collection, this time, from a chemistry perspective. I find it neat that most rocks are really just large, dense samples of simple compounds. For example, in the picture below, the big yellow rock on the bottom left is composed entirely of the element sulfur. In fact, on a few occasions I have actually chipped off a piece in order to smash it up and use it in an experiment. While grinding stones in a mortar and pestle may not seem like the most efficient way to find substances, it is, when you think about it, how most materials are obtained in the real world. The truth is, rocks and minerals are really just chemicals in their raw, sometimes pretty, but usually ugly form.

Through the years I accumulated quite a few remarkable specimens.

I placed my smaller rocks in tackle boxes.

It's clear to see that minerals come in a variety of shapes and colors.

#4: Uranium Glass

Last week I visited my grandparents in Frederick Maryland where my grandma owns a large collection of very old glassware. As I was looking through the many pieces she owned, I found something that sparked my interest. I read a book recently about the many ways that dangerous chemicals have been used throughout the past hundred years. The book talked specifically about the use of toxic and radioactive substances in dyes for glassware. With this in mind, I noticed one piece that I thought might be example of this.

It was a small dirty-green pitcher made in Czechoslovakia sometime before the 1950s. According to my book and wikipedia, there is a high likelihood that it was dyed using uranium ore (uranium oxide). This so-called "uranium glass" was very popular during the 20th century and was said to contain large amounts of radioactive uranium. Due to the fact that I don't own a Geiger counter, all I can do is speculate. While the possibility seems wild, the use of uranium as a glass dye was not uncommon and really wasn't dangerous as radiation levels were low.

Green glassware from Czechoslovakia possibly containing radioactive uranium.

To read more about uranium glass...
http://en.wikipedia.org/wiki/Uranium_glass

#3: Synthesizing Ferrofluid

I successfully created a liquid, otherwise known as a ferrofluid, that responds to magnetic fields. Through a couple shopping trips and an online order, I was able to find hydrochloric acid (HCl), steel wool, hydrogen peroxide (H2O2), ammonia (NH3), oleic acid (C18H34O2), and kerosene. I started by dissolving the steel wool in a beaker of hydrochloric acid until the solution turned a beautiful lime green color. I poured this newly-created solution of what is called ferrous chloride (FeCl2), into two flasks, one of which I corked and the other I left open and added a few drops of hydrogen peroxide. By doing this I introduced oxygen directly to the second flask forming what is known as ferric chloride (FeCl3).

Next, I combined both flasks of ferrous and ferric chloride and added them to a larger flask containing ammonia. After letting the mixture react for a little while, I placed it on my alcohol burner and began heating it. Once it became relatively hot, I added a small amount of oleic acid that would act as a surfactant. I continued to heat the solution for about two hours until most of the excess ammonia and hydrochloric acid had boiled off as a gas. 

After letting it cool down for a while, I poured the jet-black solution into a beaker and added some kerosene. It took a couple hours, but eventually my ferrofluid, which had dissolved in the kerosene thanks to my surfactant, separated from the leftover water. By decanting this oily layer, I had my magnetic liquid.

Here is my ferrofluid suspended in sugar water; it isn't top notch, but it's still pretty neat.

If you want to learn more about ferrofluid...
http://en.wikipedia.org/wiki/Ferrofluid

#2: Hydrogen Combustion

(September 2013) I reacted lye (sodium hydroxide (NaOH) with aluminum foil in a flask full of water to produce flammable hydrogen gas. I then placed a large balloon over the head of the flask to collect the gas. After tying the balloon and setting it in the middle of our patio, I lit the end of a long bamboo pole and extended it toward the balloon. The balloon instantly exploded into a bright fireball and created a loud boom as the hydrogen bonded with the surrounding oxygen to make dihydrogen monoxide (H2O).

It was quite the explosion.

#1: Diamagnetic Levitation

(September 2013) I built a real levitation device that relies on the complex diamagnetic properties of bismuth metal. I ordered a bismuth ingot online and, thanks to its melting point of 521 degrees Fahrenheit, I was able to melt it down on our stovetop. It took me a couple tries, but eventually I casted the molten metal into the pieces I wanted.

After sanding and polishing them, I positioned one over the other using Legos. I then built a wood structure to hold up an adjustable bolt to which I attached a lifter magnet. After adjusting this to the right height, I was able freely levitate a small neodymium magnet that I placed in-between the pieces of bismuth.

From the right angle it looks super awesome.

The full setup is pretty simple.

If you want to read more about diamagnetism...
http://en.wikipedia.org/wiki/Diamagnetism

If you would like to order some bismuth for yourself...
http://www.rotometals.com/Bismuth-s/4.htm