#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.