![]() ![]() The reason for these two observations became clear when Jungwirth’s postgraduate student Frank Uhlig carried out quantum-mechanical computer simulations of the process with clusters of just 19 sodium atoms. What’s more, between 0.3 and 0.5 ms, this “spiking” droplet becomes surrounded by a dark blue/purple colour in the solution. After just 0.4 ms, “spikes” of metal shoot out from the droplet, too fast to be expelled by heating. The reaction starts less than a millisecond after the metal droplet, released from a syringe, enters the water. The movies revealed a vital clue to what was fueling the violent reaction in the early stages. “Phil took it off once to blow out a small fire and a tiny piece of metal exploded into his face: luckily lower part of it, so he only had a few scratches on his cheek.” ![]() But getting a reliable explosion has its hazards. “The basic trick Phil came up with is to use liquid metal – a sodium/potassium alloy that is liquid at room temperature”, says Jungwirth. The experiments were conducted by his colleague Philip Mason, and he says that “an equally important part is Phil's love for exciting experimentation and the easy availability of our balcony, where the first experiments were carried out.” There Mason set up a high-speed video camera to film the process, although the final movies were shot in the lab of coauthor Sigurd Bauerecker at the Technical University of Braunschweig in Germany.ĭespite its notoriously explosive nature, the reaction of sodium with water is in fact extremely erratic: sometimes it explodes and sometimes it doesn’t, largely because of surface oxidation of the metal. This, Jungwirth admits, was only a part of the original motivation for looking more deeply into the reaction. But the hydrogen gas and steam released at the surface of the metal should impede the further access of water and quench the reaction. “In order to have a runaway explosive behaviour of a chemical reaction, very good mixing of the reactants needs to be ensured,” he says. The process seems so straightforward and understandable that no one previously seems to have felt there was anything else to explain.īut as Jungwirth says, there is a fundamental problem with the conventional explanation. Highly reactive sodium and potassium react with water to form sodium hydroxide and hydrogen, and the reaction liberates so much heat that the hydrogen may ignite spontaneously. We all exulted – this was chemistry with a vengeance.” ![]() Neurologist and chemical enthusiast Oliver Sacks offers a vivid account of how, as a boy, he and his friends carried out the reaction on Highgate Ponds in North London with a lump of sodium bought from the local chemicals supplier : “It took fire instantly and sped around and around on the surface like a demented meteor, with a huge sheet of yellow flame above it. That may happen eventually, but it begins as something far stranger: a rapid exodus of electrons followed by explosion of the metal driven by electrical repulsion. The explosion, say Pavel Jungwirth and his collaborators at the Czech Academy of Sciences in Prague, is not merely a consequence of the ignition of the hydrogen gas that the alkali metals release from water. Yet a paper in Nature Chemistry reveals that this familiar piece of pyrotechnics has not previously been understood. It’s the classic piece of chemical tomfoolery: take a lump of sodium or potassium metal, toss it into water, and watch the explosion. There’s more than exploding hydrogen in the violence of the reaction of alkali metals with water. Could you possibly need any more evidence that chemistry rocks? I saw the experiments being done by Phil M when I visited Pavel a couple of years ago, and have been waiting for the work to come together ever since. Here’s the longer version of my latest news story for Nature. ![]()
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