Vintage Sacks (2 page)

STINKS AND BANGS

Attracted by the sounds and flashes and smells coming from my lab, David and Marcus, now medical students, sometimes joined me in experiments—the nine- and ten-year age differences between us hardly mattered at these times. On one occasion, as I was experimenting with hydrogen and oxygen, there was a loud explosion, and an almost invisible sheet of flame, which blew off Marcus's eyebrows completely. But Marcus took this in good part, and he and David often suggested other experiments.

We mixed potassium perchlorate with sugar, put it on the back step, and banged it with a hammer. This caused a most satisfying explosion. It was trickier with nitrogen tri-iodide, easily made by adding concentrated ammonia to iodine, catching the nitrogen tri-iodide on filter paper, and drying it with ether. Nitrogen tri-iodide was incredibly touch-sensitive; one had only to touch it with a stick—a
long
stick (or even a feather)—and it would explode with surprising violence.

We made a “volcano” together with ammonium dichromate, setting fire to a pyramid of the orange crystals, which then flamed, furiously, becoming red-hot, throwing off showers of sparks in all directions, and swelling portentously, like a miniature volcano erupting. Finally, when it had died down, there was, in place of the neat pyramid of crystals, a huge fluffy pile of dark green chromic oxide.

Another experiment, suggested by David, involved pouring concentrated, oily sulfuric acid on a little sugar, which instantly turned black, heated, steamed, and expanded, forming a monstrous pillar of carbon rising high above the rim of the beaker. “Beware,” David said, as I gazed at this transformation. “
You'll
be turned into a pillar of carbon if you get the acid on yourself.” And then he told me horror stories, probably invented, of vitriol throwings in East London, and patients he had seen coming into the hospital with their entire faces all but burned off. (I was not quite sure whether to believe him, for when I was younger he had told me that if I looked at the Kohanim as they were blessing us in the shul—their heads were covered with a large shawl, a tallis, as they prayed, for they were irradiated, at this moment, by the blinding light of God—my eyes would melt in their sockets and run down my cheeks like fried eggs.)
¹

I spent a good deal of my time in the lab examining chemical colors and playing with them. There were certain colors that held a special, mysterious power for me—this was especially so of very deep and pure blues. As a child I had loved the strong, bright blue of the Fehling's solution in my father's dispensary, just as I had loved the cone of pure blue at the center of a candle flame. I found I could produce very intense blues with some cobalt compounds, with cuprammonium compounds, and with complex iron compounds like Prussian blue.

But the most mysterious and beautiful of all the blues for me was that produced by dissolving alkali metals in liquid ammonia (Uncle Dave showed me this). The fact that metals
could
be dissolved at all was startling at first, but the alkali metals were all soluble in liquid ammonia (some to an astounding degree—cesium would completely dissolve in a third its weight of ammonia). When the solutions became more concentrated, they suddenly changed character, turning into lustrous bronze-colored liquids that floated on the blue—and in this state they conducted electricity as well as liquid metal like mercury. The alkaline earth metals would work as well, and it did not matter whether the solute was sodium or potassium, calcium or barium—the ammoniacal solutions, in every case, were an identical deep blue, suggesting the presence of some substance, some structure, something common to them all. It was like the color of the azurite in the Geological Museum, the very color of heaven.

Many of the so-called transition elements infused their compounds with characteristic colors—most cobalt and manganese salts were pink; most copper salts deep blue or greenish blue; most iron salts pale green and nickel salts a deeper green. Similarly, in minute amounts, transition elements gave many gems their particular colors. Sapphires, chemically, were basically nothing but corundum, a colorless aluminum oxide, but they could take on every color in the spectrum—with a little bit of chromium replacing some of the aluminum, they would turn ruby red; with a little titanium, a deep blue; with ferrous iron, green; with ferric iron, yellow. And with a little vanadium, the corundum began to resemble alexandrite, alternating magically between red and green—red in incandescent light, green in daylight. With certain elements, at least, the merest smattering of atoms could produce a characteristic color. No chemist could have “flavored” corundum with such delicacy, a few atoms of this, a few ions of that, to produce an entire spectrum of colors.

There were only a handful of these “coloring” elements—titanium, vanadium, chromium, manganese, iron, cobalt, nickel, and copper, so far as I could see, being the main ones. They were, I could not help noticing, all bunched together in terms of atomic weight—though whether this meant anything, or was just a coincidence, I had no idea at the time. It was characteristic of all of these, I learned, that they had a number of possible valency states, unlike most of the other elements, which had only one. Sodium, for instance, would combine with chlorine in only one way, one atom of sodium to one of chlorine. But there were two combinations of iron and chlorine: an atom of iron could combine with two atoms of chlorine to form ferrous chloride (FeCl
₂
) or with three atoms of chlorine to form ferric chloride (FeCl
₃
). These two chlorides were very different in many ways, including color.

Because it had four strikingly different valencies or oxidation states, and it was easy to transform these into one another, vanadium was an ideal element to experiment with. The simplest way of reducing vanadium was to start with a test tube full of (pentavalent) ammonium vanadate in solution and add small lumps of zinc amalgam. The amalgam would immediately react, and the solution would turn from yellow to royal blue (the color of tetravalent vanadium). One could remove the amalgam at this point, or let it react further, till the solution turned green, the color of trivalent vanadium. If one waited still longer, the green would disappear and be replaced by a beautiful lilac, the color of divalent vanadium. The reverse experiment was even more beautiful, especially if one layered potassium permanganate, a deep purple layer, over the delicate lilac; this would be oxidized over a period of hours and form separate layers, one above the other, of lilac divalent vanadium on the bottom, then green trivalent vanadium, then blue tetravalent vanadium, then yellow pentavalent vanadium (and on top of this, a rich brown layer of the original permanganate, now brown because it was mixed with manganese dioxide).

These experiences with color convinced me that there was a very intimate (if unintelligible) relation between the atomic character of many elements and the color of their compounds or minerals. The same color would show itself whatever compound one looked at. It could be, for example, manganous carbonate, or nitrate, or sulfate, or whatever—all had the identical pink of the divalent manganous ion (the permanganates, by contrast, where the manganese ion was heptavalent, were all deep purple). And from this I got a vague feeling—it was certainly not one that I could formulate with any precision at the time—that the color of these metal ions, their chemical color, was related to the specific state of their atoms as they moved from one oxidation state to another. What was it about the transition elements, in particular, that gave them their characteristic colors? Were these substances, their atoms, in some way “tuned”?
²

A lot of chemistry seemed to be about heat—sometimes a demand for heat, sometimes the production of heat. Often one needed heat to start a reaction, but then it would go by itself, sometimes with a vengeance. If one simply mixed iron filings and sulfur, nothing happened—one could still pull out the iron filings from the mixture with a magnet. But if one started to heat the mixture, it suddenly glowed, became incandescent, and something totally new—iron sulfide—was created. This seemed a basic, almost primordial reaction, and I imagined that it occurred on a vast scale in the earth, where molten iron and sulfur came into contact.

One of my earliest memories (I was only two at the time) was of seeing the Crystal Palace burn. My brothers took me to see it from Parliament Hill, the highest point on Hampstead Heath, and all around the burning palace the night sky was lit up in a wild and beautiful way. And every November 5, in memory of Guy Fawkes, we would have fireworks in the garden—little sparklers full of iron dust; Bengal lights in red and green; and bangers, which made me whimper with fear and want to crawl, as our dog would, under the nearest shelter. Whether it was these experiences, or whether it was a primordial love of fire, it was flames and burnings, explosions and colors, which had such a special (and sometimes fearful) attraction for me.

I liked mixing iodine and zinc, or iodine and 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 particular heroes, had been nearly asphyxiated by nitrous oxide, poisoned himself with nitrogen peroxide, and severely inflamed his lungs with hydrofluoric acid. Davy also experimented with the first “high” explosive, nitrogen trichloride, which had cost many people fingers and eyes. He discovered several new ways of making the combination of nitrogen and chlorine, and caused a violent explosion on one occasion while he was visiting a friend. Davy himself was partially blinded, and did not recover fully for another four months. (We were not told what damage was done to his friend's house.)

The Discovery of the Elements
devoted an entire section to “The Fluorine Martyrs.” Although elemental chlorine had been isolated from hydrochloric acid in the 1770s, its far more active cousin, fluorine, was not so easily obtained. All the early experimenters, I read, “suffered the frightful torture of hydrofluoric acid poisoning,” and at least two of them died in the process. Fluorine was only isolated in 1886, after almost a century of dangerous trying.

I was fascinated by reading this history, and immediately, recklessly, wanted to obtain fluorine for myself. Hydrofluoric acid was easy to get: Uncle Tungsten used vast quantities of it to “pearl” his lightbulbs, and I had seen great carboys of it in his factory in Hoxton. But when I told my parents the story of the fluorine martyrs, they forbade me to experiment with it in the house. (I compromised by keeping a small gutta-percha bottle of hydrofluoric acid in my lab, but my own fear of it was such that I never actually opened the bottle.)