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Why this popular chemistry experiment is a blast

Explosions are fun—ask any budding chemistry student. That’s why dropping a chunk of pure sodium into water is a classic classroom demonstration. The resulting violent reaction can produce impressive flames and a loud bang. Although the basic chemistry of the popular experiment has long been understood, the details were not. Now, scientists have captured the action using high-speed video cameras and discovered an unexpected trigger. Less than a millisecond after sodium and water meet, the sodium contorts into a sea urchin–like shape, growing spikes that shoot out into the water and initiate a runaway reaction.

Sodium is one of the alkali metals, a group of famously reactive elements. Because sodium is so reactive, it is rare in its pure form, which is solid at room temperature but soft enough to cut with a knife. The sodium we’re most familiar with, in table salt, is stable because it’s bound to chloride. Pure alkali metals are a different story:  When they get together with water, they erupt furiously. A flurry of electrons abandons the metal, interacting with the water to form hydrogen gas and other byproducts. The reaction generates heat, which melts the sodium, and because hydrogen gas is flammable, it ignites. Or so the story goes.

But scientists led by chemist Pavel Jungwirth of the Academy of Sciences of the Czech Republic in Prague weren’t satisfied. “There was a missing piece in this explanation,” Jungwirth says. To produce an explosion, the chemicals involved need to be well mixed so that the reaction snowballs instead of sputtering out. But because only the surface of the sodium chunk contacts water, only atoms in its outer layer can react. Moreover, the production of hydrogen gas creates a layer that separates the sodium and the water, which should further slow the reaction, resulting in slow bubbling rather than a kaboom.

To study the details of the detonation, Jungwirth and colleagues first had to tame the reaction. The results of the demonstration are unpredictable—sometimes it flashes and sometimes it fizzles, depending on small variations in the size and shape of the chunks of sodium used. So the scientists used a liquid mixture of sodium and another alkali metal, potassium, which they could slowly drip into water in drops of a uniform size and shape. They arranged high-speed video cameras to capture the reaction at thousands of frames per second. Then they donned protective gear and stood back.

The cameras captured a never-before-seen effect. Less than a millisecond after the reaction begins, tens to hundreds of spiky metal protrusions pierce the water, the researchers report online today in Nature Chemistry. The spikes appear, the researchers deduced, because when electrons flee the metal for water, an intense positive charge builds up. The mutual repulsion of those positive charges rips the metal apart, and it blasts outward in tiny needles. This increases the surface area of the metal in contact with water, generating a vigorous reaction. Computer simulations performed by the researchers confirmed this effect, although for much smaller quantities of sodium due to the limits of computing power.  

Because the basic chemistry of the reaction is 19th century science, it’s an unconventional subject for a modern study. “I haven’t thought much about explosions like this,” says chemist Benjamin Schwartz of the University of California, Los Angeles. But, he says, “the idea makes a lot of sense once you see it.”

“The surprising thing is that we didn’t think of it before,” says chemist David Bartels of the University of Notre Dame in Indiana. “We just [took] this for granted without thinking it through.”

The detailed explanation of this reaction could be useful for preventing such explosions in industrial applications that use alkali metals, including certain nuclear reactors that are cooled with liquid metal. But the applications are an afterthought, Jungwirth says. The true motivation for the research? “We like to play with explosives.”

(Video credit: Mason et al.)