Vintage Sacks Read Free Page B

antimony—no added heat was needed here—and seeing how they heated up spontaneously, sending a cloud of purple iodine vapor above them. The reaction was more violent if one used aluminum rather than zinc or antimony. If I added two or three drops of water to the mixture, it would catch fire and burn with a violet flame, spreading fine brown iodide powder over everything.
    Magnesium, like aluminum, was a metal whose paradoxes intrigued me: strong and stable enough in its massive form to be used in airplane and bridge construction, but almost terrifyingly active once oxidation, combustion, got started. One could put magnesium in cold water, and nothing would happen. If one put it in hot water, it would start to bubble hydrogen; but if one lit a length of magnesium ribbon, it would continue to burn with dazzling brilliance
under
the water, or even in normally flame-suffocating carbon dioxide. This reminded me of the incendiary bombs used during the war, and how they could not be quenched by carbon dioxide or water, or even by sand. Indeed, if one heated magnesium with sand, silicon dioxide—and what could be more inert than sand?—the magnesium would burn brilliantly, pulling the oxygen out of the sand, producing elemental silicon or a mixture of silicon with magnesium silicide. (Nonetheless, sand was used to suffocate ordinary fires that had been started by incendiary bombs, even if it was useless against burning magnesium itself, and one saw sand buckets everywhere in London during the war; every house had its own.) If one then tipped the silicide into dilute hydrochloric acid, it would react to form a spontaneously inflammable gas, hydrogen silicide, or silane—bubbles of this would rise through the solution, forming smoke rings, and ignite with little explosions as they reached the surface.
    For burning, one used a very long-stemmed “deflagrating” spoon, which one could lower gingerly, with its thimbleful of combustible, into a cylinder of air, or oxygen, or chlorine, or whatever. The flames were all better and brighter if one used oxygen. If one melted sulfur and then lowered it into the oxygen, it took fire and burned with a bright blue flame, producing pungent, titillating, but suffocating sulfur dioxide. Steel wool, purloined from the kitchen, was surprisingly inflammable—this, too, burned brilliantly in oxygen, producing showers of sparks like the sparklers on Guy Fawkes night, and a dirty brown dust of iron oxide.
    With chemistry such as this, one was playing with fire, in the literal as well as the metaphorical sense. Huge energies, plutonic forces, were being unleashed, and I had a thrilling but precarious sense of being in control—sometimes just. This was especially so with the intensely exothermic reactions of aluminum and magnesium; they could be used to reduce metallic ores, or even to produce elemental silicon from sand, but a little carelessness, a miscalculation, and one had a bomb on one’s hands.
    Chemical exploration, chemical discovery, was all the more romantic for its dangers. I felt a certain boyish glee in playing with these dangerous substances, and I was struck, in my reading, by the range of accidents that had befallen the pioneers. Few naturalists had been devoured by wild animals or stung to death by noxious plants or insects; few physicists had lost their eyesight gazing at the heavens, or broken a leg on an inclined plane; but many chemists had lost their eyes, limbs, and even their lives, usually through producing inadvertent explosions or toxins. All the early investigators of phosphorus had burned themselves severely. Bunsen, investigating cacodyl cyanide, lost his right eye in an explosion, and very nearly his life. Several later experimenters, like Moissan, trying to make diamond from graphite in intensely heated, high-pressure “bombs,” threatened to blow themselves and their fellow workers to kingdom come. Humphry Davy, one of my