Tag Archives: Chemistry

Colored Fire for a Happy Halloween!

Note: This involves combining fire which can always be dangerous with chemicals that can irritate skin or badly hurt eyes! Only do this with proper supervision and safety equipment including goggles, gloves, and a fire extinguisher!

The classic Jack-o-lantern has been around for hundreds of years. The tradition was brought to America by the Irish, who had originally started carving the spooky faces into turnips based on a folk tale. Moving to pumpkins was undoubtedly an upgrade, and we think it’s about time for another one!


The lighting of a jack-o-lantern is classically done by candles. Candles burn with a yellowish flame giving off a special glow, but there are ways to make flame in different colors! To do this, the fuel simply needs to have different kinds of salts added. When you think of salt, it probably conjures images of a table seasoning known as sodium chloride, but this is not the only possibility.

Salt actually a scientific term referring to the category of molecule that is left over from pouring acidic and basic solutions together. Fortunately, to acquire these salts, you don’t need to risk working with potentially dangerous acids that could leave you with a costume you couldn’t take off!

Most of these are the result of mixing something with hydrochloric acid, which means that they are often chlorine combined with another molecule, just like sodium chloride! Adding energy in the form of fire causes these molecules to give off unique energy in the form of differently colored light!


Caption: Different salts give off different colors of light. This can come in many forms including from red strontium chloride and green from cupric chloride.

jackolantern-colorOur fire source was rubbing alcohol, with the chosen salt stirred in. To keep it from going everywhere, we used a small glass jar instead of the dishes above. To make sure the liquid burned easily and evenly, we used a cotton ball as a wick like a candle, and put it in the jack-o-lantern!.

The Science of Fireworks

It’s easy to change the color of a flame. Just add salt! A substance’s chemistry determines the hue of light that will be released by its combustion. Copper-based salts burn green. Strontium turns flames bright red. Sodium salts, such as table salt, burn yellow. If you’ve ever seen a multicolored fireworks show, you’ve experienced this science firsthand.


Left to right: lithium chloride, boric acid, calcium chloride, and potassium chloride burn in a hand sanitizer fuel base.

When energy is added to an atom, its outermost electrons can become excited, or jump to energized orbitals. When these electrons eventually relax back into their default minimum-energy or ground state, the excess is re-released in a form dictated by the atom’s physical structure– often a specific wavelength (color) of visible light.


The visible spectrum of the sun as observed with the Fourier Transform Spectrograph at Kitt Peak National Observatory. Credit: N.A.Sharp, NOAO/NSO/Kitt Peak FTS/AURA/NSF. Scientists use emission spectra to study the chemical composition of stars, including our own sun.

Color isn’t the only property that distinguishes burning substances. Some elements and compounds also have noticeably characteristic combustion behavior. For instance, iron filings sparkle when ignited, and magnesium flashes blindingly.


Igniting a strip of magnesium.

Pyrotechnic engineers fill firework shells with carefully selected chemical arrangements. Small packages of individual compounds, called stars, are the building blocks of their designs. By packing stars together with strategically placed explosive powder, technicians lay the groundwork for everything from simple starbursts to multi-colored tropical scenes.

Written By: Caela Barry

Fire & the History of Matches


Humans have been creating and controlling fire for almost a million years! Our early ancestors used friction – essentially rubbing sticks together – to create their first fires for cooking food and making tools. Today we can carry fire making tools around in our pockets.  Every year 500 billion matches are used in the United States alone. Even though we’ve come a long way from rubbing sticks together, matches today work on a very similar principle – friction. For an example of this, you can try rubbing your hands together. You should feel them get warm. It should not start a fire.

Match bodies are made of wood or stiff paper, to provide fuel for the fire. Match heads are coated in phosphorous based compounds that catch fire when heated up. The heat that lights a match generally comes from friction when you rub or “strike” a match on a rough surface. Early “strike anywhere” matches were coated in white phosphorous, but the white phosphorus was too easy to light. This made them rather dangerous, as they tended to ignite accidentally. Great if you needed to light a fire in a hurry…but not so great if you need to ship them long distances, or keep them for a long time.


To get a combustion reaction like this to start requires something in chemistry called “activation energy”. This simply refers the amount of energy needed to start the reaction. These matches lit so easily because the white phosphorous in the match head needed very little energy to light. Simply rocking around in a crate could cause enough friction to ignite them. Today’s strike anywhere matches use a less dangerous form of phosphorous (phosphorous sesquisulfide). They can still be lit on any surface rough enough to create the right amount of heat from friction, but anyone who has tried (and failed) to light a match multiple times can tell you it takes a bit more effort.

In the video, the bright parts you see as the fire moves along consuming all of the matches is this phosphorus beginning to ignite. It doesn’t last very long, and its job is just to keep the match burning long enough for it to get hot and start burning the wood, which is what sustains most of the flame that you see in the end!


Liquid Nitrogen & Glowsticks

Note: Opening glowsticks is not a particularly great idea. They contain bits of broken glass and some unpleasant chemicals that can be hazardous to your health! Read full article to learn more.

Glowsticks are incredibly fun, but how do they work!? Naturally because of some very cool science, notable chemistry. Glowsticks are made up of an outer plastic casing with a smaller glass casing inside. The plastic tube is filled with a dye which determines the color of the glowstick and a chemical called diphenyl oxalate. The glass tube contains hydrogen peroxide, the same thing you might use to clean out a cut or scrape.


When you crack the glowstick you break the glass, the hydrogen peroxide is released into the mixture. This causes a series of chemical reactions to take place. The main end products of this chemistry are carbon dioxide and energy, as well as another molecule we will talk about later. The energy that is released goes into the dye, which converts the chemical energy into light energy! The reaction happens slowly, so that the glow lasts for a long time. Companies can vary the amount of each chemical to have glowsticks that glow brightly for a short time or more dimly for a very long time!


It also forms something called phenol that is a somewhat toxic chemical. Repeated exposure to this chemical, and even its vapors can be dangerous. This is why breaking open a glow stick without proper protection is not advisable. There will be phenol and broken glass in the mixture, neither of which are good things to have around! We were very careful, and used protective equipment on our skin as well as working in a well ventilated area to keep ourselves safe!

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One interesting thing is that this reaction can be stopped by extreme cold. Check out what happened when we left the glowsticks in liquid nitrogen for a while! Science!

The Power of Air!


How did the can get crushed? You could see in the video it wasn’t pushed in by the tongs, so what did it!? This very simple experiment works because of something called Charles’s Law. Charles’s Law says that a gas will get bigger if it gets hotter, or smaller if it gets colder, as long as the pressure doesn’t change.

One thing that you can’t see in the video is that the water in the can is boiling. This means that the can is full of water vapor that is around 200℉! Next, the can is placed open-side down into a container of cool water, probably about 50℉. Note that we aren’t changing the pressure, so Charles’s Law tells us what happens next. The cold water cools down the water vapor, causing it to contract (and even condense!), but this is not what really crushes the can. The real culprit….is air.


Air doesn’t seem to weigh anything. We can’t see it, or pick it up and hold it in our hands very well. However, that doesn’t mean it is light! The atmosphere weighs a whopping 6,000,000,000,000,000 tons! The earth is pretty big, but that means that at sea level, there is about 15 pounds of air pushing down on every single square inch!

However, not everything gets crushed by the atmosphere. Your body effortlessly pushes back on the air to not get squished, just like the hot air in the can pushed out to keep the can from imploding. However, when the cold water cooled and contracted the air, there was nothing to push out against the atmosphere, and no way for the atmosphere to get in. So yes, the air just crushed it!

This is a great DIY experiment to do at home or try in class! It requires few materials, and can teach a lot of science! Charles’s Law is a very powerful idea, and is half of the Ideal Gas Law, which is seen in both chemistry and physics classes!

Note that this is the same principle that we used to get our egg into the bottle experiment!

DIY Science Experiment: Egg in a Bottle

You might be asking yourself why you would want to put an egg in a bottle. The answer is, of course, for science! This is a great experiment for explaining the basics of the ideal gas law. Mainly, that gases expand and contract when they change temperature. Here we will explore how to actually do the experiment, and what the science is behind it.


  • Peeled hard boiled eggs. Note that if they crack during peeling, they will likely not survive the ordeal!

  • Bottle with a neck smaller than the egg. Erlenmeyer flasks work great!

  • Matches, or a small piece of flammable material

  • Workspace clear of burning hazards

That’s it!

Doing the experiment:

  • Read the steps first. They have to happen quickly!

  • Light a match or something small and flammable. A small piece of paper works great!. In the video, we use four strike-anywhere matches.

  • Drop your flaming object of choice into the flask. With matches, do it quickly! If you wait too long, there wont be enough fire to heat the air!

  • Put the egg on top of the flask so it completely covers the opening. This must be done quickly

  • Watch!

You should see the fire go out and the egg get sucked into the bottle shortly thereafter.

What happened?

The fire rapidly heats the air in the flash. Then, the fire should quickly become starved of oxygen and go out. Once the fire stops and the air begins to cool. The molecules in the gas slow down as it cools, decreasing the pressure inside the flask. The air pressure outside is then greater, and pushes the egg down the seemingly-too-small neck of the bottle. Because the air pushes equally from all sides, the egg stays intact, unlike if you had done it with your hand!

How do I get it out?!

There are three good ways to do this.

  1. Get something pokey, like a butter knife, and chop the egg into bits and dump it out. Messy. Not my favorite.

  2. Blow in behind the egg (like in the video–note, we shook the matches out first for safety). The hot air on your breath should be enough to push it out.

  3. Flip the flask over so the egg covers the opening from within. Run hot tap water over the base of the flask. As the air inside heats up, it pushes the egg out, simply doing the experiment in reverse!


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