Fireworks

If you haven’t seen a video of London’s NYE 2012 Explosive Extravaganza, brace yourself:

http://www.youtube.com/watch?v=q1yLRK2M8YQ


My viewing experience went something like this:

Minute One: Enchanted. Fireworks are shooting out of Big Ben!! The timing is perfect. Ooh, pretty sparkles.

Minute Five: Saturated. Mind begins to drift… What’s this singer’s name again? That was a cool music video, the one with the water glasses.

Minute Seven: Awakened from my daze by a Bhangra remix of Jump Around. The extravaganza is STILL going on?! Becoming a little bit angry. As an aspiring atmospheric chemist, I can’t help but fixate on the sulfurous plume masking the colors of the explosions.

Minute Ten: Check the timing-bar on the YouTube video. What could the grand finale possibly be at this point? Is Big Ben going to rocket into space while spewing a rainbow of sparks? (Alas, the clock tower remains grounded and the ending is decidedly anticlimactic).

Before I go into atmospheric science mode, here’s a quick lesson about the chemistry of those pretty pyrotechnics: A firework is a metal capsule consisting of several compartments stacked on top of one another. Each compartment contains the ingredients for a single explosion; a time-delay fuse runs through the center of the capsule so that the compartments explode one by one. Initially, the capsule sits inside a launch tube. The bottom-most compartment is ignited first, and the rapid buildup of hot gas in the base of the tube propels the firework high into the air. The fuse continues to burn, and the other compartments ignite into colorful bursts. The magic ingredients are gunpowder (for the explosions), and metal-containing salts (for color).

As you’ve probably learned at some point, combustion is the process of breaking chemical bonds (usually carbon-carbon and carbon-hydrogen bonds) and re-forming stronger bonds with oxygen, thus releasing energy. Gunpowder is a mixture of chemicals that undergoes particularly rapid and violent combustion. It contains 75% potassium nitrate (KNO₃), which provides the oxygen, 15% charcoal, which provides the energy-containing bonds (the fuel), and 10% sulfur, which ignites relatively easily, causing the explosion to happen more quickly and at a lower temperature.


What makes these explosions appealing is, of course, the colorful light. Any hot object gives off light. The random motion of sub-atomic particles produces electromagnetic radiation, and the wavelength range of this radiation depends on the object’s temperature (a very hot flame looks blue, while a less-hot flame looks orange). So, controlling the temperature of the combustion by adding different chemicals to the firework will change its hue. However, this kind of light contains a broad range of wavelengths. To get those brilliant reds and blues, a purer source of radiation is needed. This is why salts like lithium carbonate and sodium chloride are added to the mix. The metal atoms absorb energy during the blast, and their electrons are excited to higher energy levels. When an electron falls back toward the nucleus, it emits a photon (a little packet of light). The energy of the photon is characteristic of the atom. For example, copper emits blue light, lithium emits red light, and calcium emits orange light.


So, we’ve got metal-encased bundles of potassium nitrate, charcoal, sulfur, salts, plus the binding material to hold everything together, exploding into the air. Quite a lot of interesting chemistry going on. And quite a lot of stuff being released into the atmosphere.

Like any other combustion source, fireworks emit carbon dioxide, carbon monoxide, and myriad organic molecules. During the blasts, the nitrate and sulfur from the gunpowder are converted into sulfur dioxide (SO₂), which is poisonous, and nitrogen oxides (NO_x), which can then react to produce ozone. SO₂ and NO_x also contribute to the formation of atmospheric aerosols and acid rain. This is all bad news for local air quality, especially during the wintertime when temperature inversions trap pollutants by preventing the vertical mixing of air.

To be fair, the total pollution emitted by fireworks is probably much less than that emitted by transportation and power plants. What is really concerning, though, is that “sulfurous plume” I mentioned. The hazy appearance is a result of light being scattered by millions of tiny particles: little bits of salt and carbon drifting above the heads of 250,000 spectators. When inhaled, those aerosols can damage the respiratory and cardiovascular systems. Several scientific studies have shown that fireworks produce high levels of sub-micron particles (those with a diameter of less than one-millionth of a meter). These super-tiny bits are particularly hazardous, as they can easily penetrate the lungs. One group of scientists measured a sub-micron aerosol concentration of 600 micrograms per cubic meter (μg/m³) during a fireworks show in 2005 in Germany. That’s nearly 20 times the U.S. EPA’s current standard for fine aerosol particles!

One more disconcerting thought: firework aerosols contain unusually high amounts of trace metals. Increased airborne concentrations of potassium, barium, magnesium, strontium, copper, arsenic, and lead have been measured during pyrotechnical displays. Imagining an army of little atomic symbols parachuting into the Thames and infiltrating people’s bloodstreams detracted from my enchantment…

… which is why a fireworks show should not overload the attentive capacities of its viewers.

Like most good things, dazzling explosives are best enjoyed in moderation. Over-indulge, and the consequences begin to outweigh the benefits. The total happiness derived from London’s massive display could not have been worth the sum of the public health costs and the ₤1.9 million production bill. Perhaps the national self-esteem boost made up Britain’s deficit? Hopefully so, because I doubt the extravagant show gained them much admiration from abroad.

References:

Bushby, H. London 2012: How new year fireworks dazzled. BBC News [online] January 3, 2012. http://www.bbc.co.uk/news/uk-16391068.
Drewnick et al. Atmospheric Environment, 2006, 40, 4316.
Gondhia, Reema. The Chemistry of Fireworks. http://www.ch.ic.ac.uk/local/projects/gondhia/.
Moreno et al. Journal of Hazardous Materials, 2010, 183, 945.
U.S. Environmental Protection Agency. PM Standards Revision-2006. http://epa.gov/pm/naaqsrev2006.html.
Vecchi et al. Atmospheric Environment, 2008, 42, 121.