Al Bruno says every fireworks show ends with “the last word.”
Bruno, a scientific interpreter at The Franklin Institute, says that’s the term pyrotechnicians use to refer to a compression wave. The wave is that “boom” released after the initial combustion causes the gases to expand and collide. If you’re close enough, you can feel as well as hear it. Bruno says every kid in the audience knows exactly what that means.
“That is usually the end of the show,” he said. “We end it with a bang.”
Bruno’s job is to stage science demonstrations, sometimes in the theater, but other times throughout the Franklin, using a cart full of materials labeled “random acts of science.” This month’s “Science of Fireworks” show is one of them.
“I love the fireworks show,” Bruno said. “I look forward to the fireworks show every summer.”
The anatomy of a firework
Since the first anniversary of Independence Day in 1777, America’s love affair with fireworks has become one of our most iconic traditions. But the earliest use of fireworks dates back to 7th century China.
Bruno explains that fireworks come in a neat package. Each package includes a bursting charge and pellets of elements that produce different colors when the burn. They’re packed into a tube and fired out of mortars into the air to create the “aerial display show” we have all come to expect on the the holiday.
The bursting charge is what sets the whole thing off. “Traditionally the bursting charge was made with an oxidizer known as saltpeter, that’s potassium nitrate, and with a fuel, usually carbon, sulfur and maybe some phosphorus,” Bruno said, adding that the ingredients in fireworks have changed little over the course of two centuries.
“But the real part is that oxidizer, that potassium nitrate, the saltpeter,” he said. “That was used as a food preservative thousands of years ago, just like salt, and it wasn’t until old Chinese chefs found that if you take some of that saltpeter and sprinkle it on the coals, on the carbon, that the flame would burn more intensely, [it] would burn hotter.”
Yellows, reds and blues
Next comes the colors. Anyone remotely familiar with the periodic table of elements will recognize the names responsible for fireworks colors.
For yellows, there’s sodium. “Now, everyone’s seen this sodium color before in old street lamps,” Bruno said. “If you ever notice how old street lamps have a yellowish orangish color. Those are sodium bulbs.”
To get vibrant reds, technicians use an metal called strontium. “If you’ve ever seen a road flare. Road flares have strontium in them,” he said.
Blues or greens can usually be achieved with copper, but it’s a very tricky color to achieve. A lot depends on very carefully on the temperature of the flame and the mixture of copper with other elements.
Potassium, an element found in bananas, makes for a light, lavender color, while lithium makes for a bright fuchsia or pink.
To get white, either magnesium or aluminum is used. “That’s also traditionally what’s in flash powder that old photographers used to use,” Bruno said.
When the bursting charge goes off, the the expanding gases ignite the pellets and scatter them into the air.
Light and heat
To demonstrate the colors on a small scale for audiences, Bruno sprays an open flame with different metals dissolved in water. He picks up a bottle labeled “copper” and sprays it on a candle to create a fan of bright green flames.
Other demonstrations aim to explain the two most important outcomes of fireworks: light and heat. Bruno invites a volunteer to ignite a balloon filled with hydrogen (from a safe distance), which explodes briefly into a ball of flames.
“It’s wonderful science,” Bruno said. “It’s wonderful, understandable, accessible science.”