Fire, now in Technicolor!
All right, I think I’ve played with enough fire to last me for a good long time so with this post I bring Flammable Week to a close. (Let me just point out that this has been the longest week ever.) It’s going out in style though, as today’s topic is how to make your fire burn in a full spectrum of brilliant colors. After all, what’s better than fire? Fire that looks cool, that’s what.
Making colored fire is easy. In fact, it’s basically throwing chemicals onto the flames. Not just any chemicals, mind you, as random chemicals will more likely than not create horribly toxic vapors without any noticeable change in the fire’s color. (Even worse than sustaining liver damage is doing so and not even getting a good lightshow out of it.) However, there are a whole host of chemicals, mostly metal salts, that are well-known in the field of pyrotechnics for the dramatic colors they create when burned. Some are hard to come by, like strontium chloride (SrCl2) and potassium nitrate (KNO3), but many are common household chemicals like calcium chloride (CaCl2, road salt) and boric acid (H3BO3, a disinfectant).
Before I start listing them off though, there’s science to be done. Let’s talk about why burning metal salts make pretty colors.
The phenomenon at work here is a chemical property called the ‘bright line emission spectrum’ (sometimes just ’emission spectrum’), that is created by activity within an atom’s electron cloud.
We know that molecules are made up of atoms, and that individual atoms are made up of (usually) equal numbers of protons, neutrons, and electrons. (Unequal amount of the three cause things like isotopes and ions, but that’s a story for another day.) The protons and neutrons from the nucleus of the atom and account for most of its mass while the electrons orbit them in semi-determined areas/paths (it’s wibbly) called, surprise surprise, electron orbitals. These orbitals are what we’re interested in at the mo.
Electrons, in addition to their negative charge, carry a certain amount of energy that determines what orbital they’ll stay in. The closer to the nucleus an orbital is, the lower energy it is. The further an orbital is from the nucleus, the more energy an electron needs to stay in it.
The order of electrons expanding out from the nucleus is: 1s, 2s, 2p, 2p, 2p, 3s, 3p, 3p, 3p, 3d, 3d, 3d, 3d, 3d, 4s, 4p, 4p, 4p, 4d, 4d, 4d, 4d, 4d, 4f, 4f, 4f, 4f, 4f, 4f, 4f, 5s… ad nauseum.
The letter s, p, d, or f refers to the shape of the orbital (spherical, hourglass, clover leaf, or ‘Oh god, what is that?!’ respectively) while the number denotes what electron shell the orbital is located in. The higher the number, the higher the energy.
Because low-energy states are more stable, an atom’s electrons will fill its innermost orbitals first. This means that smaller elements like oxygen, having 8 electrons, will fill its most central s and p orbitals and stop when it runs out of electrons instead of storing them in the higher-energy d and f orbitals far out from the nucleus.
A more massive atom like, say, good old uranium which has 92 electrons will fill up far more orbitals simply because it has more electrons that need to be stored somewhere. For example, oxygen fills only the 1s to second 2p orbitals while uranium fills all the orbitals from 1s to 7s.
The thing is, even if an orbital isn’t occupied by electrons at the moment it’s still there, just vacant. Oxygen may not use its 8f orbitals, but it does still have seven of them. This is an important concept because, if given enough energy, electrons can jump between their home orbitals and higher energy orbitals.
They’ll stay in the new orbital for a short time (fractions of seconds) before falling back to their original position. When they make that return trip, the extra energy they absorbed in order to jump up in the first place is emitted in the form of a photon. The wavelength, and thus color, of the photon emitted depends on how far the electron jumped between orbitals. (Did it jump a single orbital? A shell? Two shells?) A greater distance requires more energy and will release a photon with a shorter wavelength.
When viewed through diffraction grating the photons emitted by all of an atom’s electrons jumping and falling at once produce a specific pattern of colored lines, called a bright line emission spectrum, which is unique to each element and can be used as a form of identification when trying to analyze chemical compounds.
Right about now all this info should be starting to coalesce into a nice, big, wibbly-wobbly ball of science in your mind.
When you burn metal salts in a campfire, you’re providing those chemicals’ electrons with massive amount of energy, causing them to begin wildly jumping and falling between electron orbitals. All that activity unleashes burst after burst of photons which are the lovely colors we see. The different colors produced by each chemical are derived from the brightest band in their emission spectrum, which is determined by the unique way their electrons jump. Makes sense, no?
Now, I’m obliged to say that because this project involves burning metals, which can create very nasty fumes, do this outside. Ventilation is important here because damage to one’s major organs is not fun. You should be fine as long as you’re outdoors and don’t go out of your way to inhale large quantities of smoke. (Even with a plain old wood fire that’s bad.)
This concludes the educational segment of today’s program. Let’s burn things.
The most commonly used chemicals (called ‘colorants’ in this application) are salts, which are most commonly compounds of a metal ion and halogen. Since compounds of this sort are usually very biologically active (read: toxic) you can find many of them among your normal household chemicals as disinfectants, anti-fungal agents, herbicides, ect. (Can you see why you don’t want to breathe the fumes?)
Taking it from the top of the spectrum (or bottom, depending on whether you’re talking about frequency or wavelength), here are the most common, effective, and least toxic colorants needed to make…
Strontium-containing chemicals like strontium chloride (SrCl) and strontium nitrate (Sr(NO3)2) produce a bright red flame. SrCl2 is a little difficult to come by, but Sr(NO3)2 can be obtained from road flares if you’re careful.
Calcium chloride (CaCl2), aka road salt, is common, cheap, and will produce a strong orange flame. Just be sure to read the label on your salt carefully and make sure it’s CaCl2 and not some other chemical before casting it in.
Normal wood fires are already yellow, but I’d adding this for completion’s sake. Every day table salt, sodium chloride (NaCl), will make your soup saltier and your yellow fires yellower.
A lighter green fire can be made by adding borax to the mix.
To make a vivid green flame, use boric acid, a common insecticide and disinfectant that can be found at the pharmacy. (I like this one best.)
This is in some ways the easiest color to produce and the hardest in others. One way to make a strong blue fire is simply to burn a fuel other than wood. Alcohols of all sorts (methanol, ethanol, isopropyl, ect.) all burn with a blue flame on their own. No need for colorants.
The harder way is using copper chloride (CuCl). The main problem with this is that CuCl isn’t widely available. Your options in getting it essentially boil down to ordering some from a chemical supplier or making your own by dissolving copper bits, likely wire, in muriatic (hydrochloric) acid (HCl), which is a messy and dangerous reaction that you should probably avoid all together.
Your next best bet is just throwing copper wire or powder into the fire. You could also use copper (II) sulfate (CuSO4), which is a combination fungicide/pesticide and can sometimes be found at gardening and pool supply stores. The only thing is that it burns more blue-green than straight blue.
Sorry, no indigo fire. Yet.
To get purple fire, just toss on potassium compounds. A good choice is potassium chloride (KCl), which is sold as a substitute for regular table salt for people who are on a low-sodium diet and is readily available at the grocery store. Note, it’s not ‘low-sodium’ salt. That’s something else entirely. Failing that, you can always throw on bananas. They’re chock full of potassium.
Now, the only question left is that of a delivery system. How exactly are you going to get your colorant into the fire to be burned?
Sure, you could just throw the raw chemicals on in a powdered or chunked form, that’s obvious (If you do, wear gloves. A lot of those compounds are toxic and can be absorbed through the skin.), but there are other options available as well.
One is you can soak porous materials (pinecones, sawdust, natural fiber fabrics, paper, wood, ect.) in a solution of the colorant then allow the items to dry. When they burn in a fire they’ll carry the colorant they’re saturated with along for the ride.
Another, slightly more hazardous method involves dissolving the colorants in alcohol and using a spray bottle to spritz the mix onto the fire. (Careful with this one.)
However you decide to do it, have fun, enjoy the show, and don’t light yourself on fire. Have at it!