Despite being a staple demonstration in many introductory chemistry classes, the classic explanation for the explosive reaction between alkali metals and water has long been incorrect.

Many middle and high school students are familiar with the demonstration. Almost immediately following contact with water, alkali metals such as sodium and potassium produce a brilliant and highly energetic explosive pop. Instructors the world over would often confidently follow by explaining that the  reaction produces hydrogen gas, whose subsequent ignition is responsible for the theatrics.

However, recent research published in Nature Chemistry shows that things are not actually so simple. Although the hydrogen gas may indeed eventually ignite, the initial rapid explosion is caused by something almost entirely unrelated.

In retrospect, it seems obvious that there was something wrong with the orthodox explanation. In order for a reaction to produce an explosion, the reactants would have to mix very effectively in order to react rapidly and release energy suddenly. This is why flour mills are so susceptible to explosive outbreaks of fire; a build-up of finely ground flammable particles suspended in the oxygen-rich air allows any spark to consume an enormous amount of fuel virtually instantaneously.

Alkali metals, on the other hand, are solids. The water can only come into contact with the outer surface, which should result in a brief layer of products preventing it from reaching deeper layers right away. Water isn’t immediately in contact with every metal atom, so at the very least the reaction should proceed more slowly than it does.

In order to investigate this further, researcher Pavel Jungwirth and others set out to scrutinize the reaction with the use of high-speed cameras. Because pure alkali metals tend to accumulate an oxidized layer on their outer surfaces, causing them to be less reactive in water, he used an alloy of sodium and potassium that is liquid at room temperature.

The images captured by the cameras were very telling. The reaction begins less than a millisecond after the droplet contacts the water. At 0.4 milliseconds, spike-like tendrils of metal shoot outward, much too quickly to have been produced by heat. Most interestingly, this spiked droplet develops a never-before-seen aura of dark bluish purple color in the surrounding solution between 0.3 and 0.5 seconds (see supplemental video). This blue color turned out to be the key to understanding what was really going on.

The origin of this mysterious color was confirmed when Jungwirth’s colleage Frank Uhlig recreated the reaction in a quantum-mechanical simulation. This digital analysis revealed that atoms at the surface of the cluster were each stripped of an electron within just a few picoseconds. The electrons then rapidly shoot away from one another and become solvated in the surrounding solution. Free electrons in solution, as many chemists know, appear blue to the naked eye. The loss of these electrons leaves the atoms positively charged, resulting in an incredibly strong repulsive force blowing the cluster apart.

This research represents a feature of science that keeps so many people fascinated by it. Although it may seem like the basics are well-understood, surprises like this frequently come from the most unexpected of places. Scientific knowledge is highly dynamic and constantly evolving, as nature proves time and time again that the richness and complexity of reality rivals the limits of human imagination.

Written by: Aisling Williams


Mason, Philip E., Frank Uhlig, Vaclav Vanek, Tillmann Buttersack, Sigurd Bauerecker, and Pavel Jungwirth. "Coulomb Explosion during the Early Stages of the Reaction of Alkali Metals with Water." Nature Publishing Group, 26 Jan. 2015. Web. 8 Feb. 2015.