Another experiment, suggested by David, involved pouring concentrated, oily sulphuric 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.)[9]
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 a 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 aluminium oxide, but they could take on every color in the spectrum – with a little bit of chromium replacing some of the aluminium, 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 (FeCl2) or with three atoms of chlorine to form ferric chloride (FeCl3). 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’?[10]
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 sulphur, 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 sulphide – 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 sulphur 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.
9
I read John Hersey’s
When he had penetrated the bushes, he saw there were about twenty men, and they were all in exactly the same nightmarish state: their faces were wholly burned, their eyesockets were hollow, the fluid from their melted eyes had run down their cheeks. (They must have had their faces upturned when the bomb went off…)
10
Such thoughts about ‘tuning’, I was later to read, had first been raised in the eighteenth century by the mathematician Euler, who had ascribed the color of objects to their having ‘little particles’ on their surface – atoms – tuned to respond to light of different frequencies. Thus an object would look red because its ‘particles’ were tuned to vibrate, resonate, to the red rays in the light that fell on it:
The nature of the radiation by which we see an opaque object does not depend on the source of light but on the vibratory motion of the very small particles [atoms] of the object’s surface. These little particles are like stretched strings, tuned to a certain frequency, which vibrate in response to a similar vibration of the air even if no one plucks them. Just as the stretched string is excited by the same sound that it emits, the particles of the surface begin to vibrate in tune with the incident radiation and to emit their own waves in every direction.
David Park, in
I think this was the first time anyone who believed in atoms ever suggested that they have a vibrating internal structure. The atoms of Newton and Boyle are clusters of hard little balls, Euler’s atoms are like musical instruments. His clairvoyant insight was rediscovered much later, and when it was, nobody remembered who had it first.