by Oliver Sacks
And so I set up a little lab of my own at home. There was an unused back room I took over, originally a laundry room, which had running water and a sink and drain and various cupboards and shelves. Conveniently, this room led out to the garden, so that if I concocted something that caught fire, or boiled over, or emitted noxious fumes, I could rush outside with it and fling it on the lawn. The lawn soon developed charred and discolored patches, but this, my parents felt, was a small price to pay for my safety – their own, too, perhaps. But seeing occasional flaming globules rushing through the air, and the general turbulence and abandon with which I did things, they were alarmed, and urged me to plan experiments and to be prepared to deal with fires and explosions.
Uncle Dave advised me closely on the choice of apparatus – test tubes, flasks, graduated cylinders, funnels, pipettes, a Bunsen burner, crucibles, watch glasses, a platinum loop, a desiccator, a blowpipe, a retort, a range of spatulas, a balance. He advised me too on basic reagents – acids and alkalis, some of which he gave me from his own lab, along with a supply of stoppered bottles of all sizes, bottles of varied shapes and colors (dark green or brown for light-sensitive chemicals), with perfectly fitting ground-glass stoppers.
Every month or so, I stocked my lab with visits to a chemical supply house far out in Finchley, housed in a large shed set at a distance from its neighbors (who viewed it, I imagined, with a certain trepidation, as a place that might explode or exhale poisonous fumes at any moment). I would hoard my pocket money for weeks – occasionally one of my uncles, approving my secret passion, would slip me a half crown or so – and then take a succession of trains and buses to the shop.
I loved to browse through Griffin & Tatlock as one would browse through a bookshop. The cheaper chemicals were kept in huge stoppered urns of glass; the rarer, more costly substances were kept in smaller bottles behind the counter. Hydrofluoric acid – dangerous stuff, used for etching glass – could not be kept in glass, so it was sold in special small bottles made of gummy brown gutta-percha. Beneath the serried urns and bottles on the shelves were great carboys of acid – sulphuric, nitric, aqua regia; globular china bottles of mercury (seven pounds of this would fit into a bottle the size of a fist), and slabs and ingots of the commoner metals. The shopkeepers soon got to know me – an intense and rather undersized schoolboy, clutching his pocket money, spending hours amid the jars and bottles – and though they would warn me now and then, ‘Go easy with that one!’ they always let me have what I wished.
My first taste was for the spectacular – the frothings, the incandescences, the stinks and the bangs, which almost define a first entry into chemistry. One of my guides was J.J. Griffin’s Chemical Recreations, an 1850ish book I had found in a secondhand bookshop. Griffin had an easy, practical, and above all playful style; chemistry was clearly fun for him, and he made it fun for his readers, readers who must often have been, I decided, boys like myself, for he had sections like ‘Chemistry for the Holidays’ – this included the ‘Volatile Plum Pudding’ (‘when the cover is removed…it leaves its dish and rises to the ceiling’), ‘A Fountain of Fire’ (using phosphorus – ’the operator must take care not to burn himself’), and ‘Brilliant Deflagration’ (here, too, one was warned to ‘remove your hand instantly’). I was amused by the mention of a special formula (sodium tungstate) to render ladies’ dresses and curtains incombustible – were fires that common in Victorian times? – and used it to fireproof a handkerchief for myself.
The book opened with ‘Elementary Experiments,’ experiments first with vegetable dyes, seeing their color changes with acids and alkalis. The most common vegetable dye was litmus – it came from a lichen, Griffin said. I used some of the litmus papers that my father kept in his dispensary, and saw how they turned red with different acids or blue with alkaline ammonia.
Griffin suggested experiments with bleaching – here I used my mother’s bleaching powder in place of the chlorine water he suggested, and with this I bleached litmus paper, cabbage juice, and a red handkerchief of my father’s. Griffin also suggested holding a red rose over burning sulphur, so that the sulphur dioxide produced would bleach it. Dipping it into water, miraculously, restored its color.
From here Griffin moved (and I with him) to ‘sympathetic inks,’ which became visible only when heated or specially treated. I played with a number of these – lead salts, which turned black with hydrogen sulphide; silver salts, which blackened when exposed to light; cobalt salts, which became visible when dried or heated. All this was fun, but it was chemistry, too.
There were other old chemistry books lying around the house, some of which had been my parents’ when they were medical students, and some, more recent, belonging to my older brothers Marcus and David. One such was Valentin’s Practical Chemistry, a workhorse of a book – straight, uninspired, pedestrian in tone, designed as a practical manual, but nevertheless, for me, filled with wonders. Inside its cover, corroded, discolored, and stained (for it had done time in the lab in its day), it bore the words ‘Best wishes and congratulations 21⁄1⁄13 – Mick’ – it had been given to my mother on her eighteenth birthday by her twenty-five-year-old brother Mick, already a research chemist himself. Uncle Mick, a younger brother of Dave, had gone to South Africa with his brothers, and then worked in a tin mine on his return. He loved tin, I was told, as much as Uncle Dave loved tungsten, and he was sometimes referred to in the family as Uncle Tin. I never knew Uncle Mick, for he died of a malignancy the year I was born – he was only forty-five – a victim, his family thought, of the high levels of radioactivity in the uranium mines in Africa. But my mother had been very close to him, and his memory and image stayed vividly in her mind. The notion that this was my mother’s own chemistry book, and of the never-known, young chemist uncle who gave it to her, made the book especially precious to me.
There was a great popular interest in chemistry in the Victorian era, and many households had their own labs, as they had their ferneries and stereoscopes. Griffin’s Chemical Recreations had originally been published around 1830 and was so popular that it was continually revised and brought out in new editions; I had the tenth, published in 1860.«8»
A companion volume to Griffin’s, published at much the same time and in the same green and gilt binding, was The Science of Home Life, by A.J. Bernays, which focused on coal, coal gas, candles, soap, glass, china, earthenware, disinfectants – everything that might be contained in a Victorian home (and much of which was still contained in houses a century later).
Very different in style and content, though equally designed to awake the sense of wonder (‘The common life of man is full of Wonders, Chemical and Physiological. Most of us pass through this life without seeing or being sensible of them…’) was The Chemistry of Common Life, by J.F.W. Johnston, written in 1859. This had fascinating chapters on ‘The Odours We Enjoy’, ‘The Smells We Dislike’, ‘The Colours We Admire’, ‘The Body We Cherish’, ‘The Plants We Rear,’ and no less than eight chapters on ‘The Narcotics We Indulge In.’ This introduced me not only to chemistry, but to a panorama of exotic human behaviors and cultures.
A much earlier book, of which I was able to get a battered copy for sixpence – it had no covers, and a few pages missing – was The Chemical Pocket-Book or Memoranda Chemica, written in 1803. The author was a James Parkinson, of Hoxton, whom I would reencounter in my biology days as the founder of paleontology, and then again, when I was a medical student, as the author of the famous Essay on the Shaking Palsy – which came to be known as Parkinson’s disease. But for me, at eleven, he was just the author of this delightful little pocket book of chemistry. I got a strong sense, from his book, of how chemistry was expanding, almost explosively, at the beginning of the nineteenth century; thus Parkinson spoke of ten new metals – uranium, tellurium, chromium, columbium (niobium), tantalum, cerium, palladium, rhodium, osmium, iridium – all having been discovered in the preceding few years.
It was from Griffin that I first gained a clear
idea of what was meant by ‘acids’ and ‘alkalis’ and how they combined to produce ‘salts.’ Uncle Dave demonstrated the opposition of acids and bases by measuring out precise quantities of hydrochloric acid and caustic soda, which he mixed in a beaker. The mixture became extremely hot, but when it cooled, he said, ‘Now try it, drink it.’ Drink it – was he mad? But I did so, and tasted nothing but salt. ‘You see,’ he explained, ‘an acid and a base come together, and they neutralize each other; they combine and make a salt.’
Could this miracle happen in reverse, I asked? Could salty water be made to produce the acid and the base all over again? ‘No,’ Uncle said, ‘that would require too much energy. You saw how hot it got when the acid and base reacted – the same amount of heat would be needed to reverse the reaction. And salt,’ he added, ‘is very stable. The sodium and chloride hold each other tightly, and no ordinary chemical process will break them apart. To break them apart you have to use an electric current.’
He showed me this more dramatically one day by putting a piece of sodium in a jar full of chlorine. There was a violent conflagration, the sodium caught fire and burned, weirdly, in the yellowish green chlorine – but when it was over, the result was nothing more than common salt. I had a heightened respect for salt, I think, after that, having seen the violent opposites that came together in its making and the strength of the energies, the elemental forces, that were now locked in the compound.
Here, too, Uncle Dave showed me, the proportions had to be exact: 23 parts of sodium, by weight, to 35.5 of chlorine. I was struck by these numbers, for they were already familiar: I had seen them in lists in my books; they were the ‘atomic weights’ of these elements. I had learned these numbers by rote, in the same mindless way one learns multiplication tables. But when Uncle Dave brought up these selfsame numbers in relation to the chemical combination of two elements, a slow, underground questioning started in my head.
In addition to my collection of mineral samples, I had a collection of coins, housed in a small wooden cabinet of highly polished mahogany, with doors that opened like the doors of a toy theater, revealing a series of slim trays with velvet-covered circles for the coins – some as small as a quarter-inch across (this for groats, for silver threepenny pieces, and for Maundy money, tiny silver coins given on Easter to the poor), others almost two inches across (for crowns, which I loved, and even larger than these, the gigantic twopenny pieces made at the end of the eighteenth century).
There were also stamp albums, and the stamps I most loved were those of remote islands, with pictures of local scenes and plants, stamps which could themselves provide a vicarious voyage. I adored stamps showing different minerals, and peculiar stamps of various sorts – triangular ones, imperforate ones, stamps with inverted watermarks or missing letters or advertisements printed on the back. One of my favorites was a strange Serbo-Croat stamp from 1914 which was said to show the features of the murdered Archduke Ferdinand when viewed from a certain angle. But the collection closest to my heart was a singular collection of bus tickets. Whenever one took a bus in London in those days, one got a colored oblong of cardboard bearing letters and numbers. It was after getting an O 16 and an S 32 (my initials, also the symbols of oxygen and sulphur – and added to these, by a happy chance, their atomic weights, too) that I decided to make a collection of ‘chemical’ bus tickets, to see how many of the ninety-two elements I could get. I was extraordinarily lucky, so it seemed to me (though there was nothing but chance involved), for the tickets accrued rapidly, and I soon had a whole collection (W 184, tungsten, gave me particular pleasure, partly because it provided my missing middle initial). There were, to be sure, some difficult ones: chlorine, irritatingly, had an atomic weight of 35.5, which was not a whole number, but, undismayed, I collected a CI 355 and inked in a tiny decimal point. The single letters were easier to get – I soon had an H 1, a B 11, a C 12, an N 14, and an F 19, besides the original O 16. When I realized that atomic numbers were even more important than atomic weights, I started to collect these as well. Eventually, I had all the known elements, from H 1 to U 92. Every element became indissolubly associated with a number for me, and every number with an element. I loved carrying my little collection of chemical bus tickets with me; it gave me the sense that I had, in the space of a single cubic inch, the whole universe, its building blocks, in my pocket.
CHAPTER EIGHT
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 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.