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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