Periodic Tales

by Hugh Aldersey-Williams

  Wright’s fanciful image could not be taken as reliable evidence, of course, but it seemed to offer encouragement. The set-up was simple and brutal, which would make it easier for me to replicate the experiment. And I now knew what I should expect to see if it worked. But the starting materials concealed within the furnace remained as mysterious as ever. How did one get from liquid urine to anything that could be put in a furnace?

  Fortunately, as Brand and his competitors traipsed round the courts of Europe with specimens of the noctiluca, or ‘night-light’, stoppered up in their pockets, some of the leading scientists of the day were there to take notes and make their own investigations, from which more coherent recipes did begin to appear. One of the clearest is to be found in the papers of Robert Hooke, one of the original fellows of the Royal Society, published twenty-three years after his death, in 1726:

  Take a Quantity of Urine (not less for one Experiment than 50 or 60 Pails full); let it lie steeping in one or more tubs, or an Hogshead of oaken Wood, till it putrify and breed Worms, as it will do in 14 or 15 days. Then, in a large kettle, set some of it to boil on a strong Fire, and, as it consumes and evaporates, pour in more, and so on, till, at last, the whole Quantity be reduced to a Paste, or rather a hard Coal, or Crust, which it will resemble; and this may be done in two or three Days, if the Fire be well tended, but else it may be doing a Fortnight or more. Then take the said Paste, or Coal; powder it, and add thereto some fair Water, about 15 Fingers high, or four Times as high as the Powder; and boil them together for 1/4 of an Hour. Then strain the Liquor and all through a woollen Cloth; that which sticks behind, may be thrown away, but the Liquor that passes must be taken and boil’d till it come to a Salt, which it will be in a few Hours.

  After that, it was simply a matter of adding some Caput Mortuum (or ‘death’s head’, which, apparently, ‘you have at any Apothecary’s’) to the salt and soaking the resulting mess in alcohol ‘so as it will become a kind of Pap’:

  Then evaporate all in warm Sand, and there will remain a red, or reddish, Salt. Take this Salt, put it into a Retort, and, for the first Hour, begin with a small Fire; more the next, a greater the 3d, and more the 4th; and then continue it, as high as you can for 24 Hours. Sometimes by the Force of the Fire, 12 Hours proves enough; for when you see the Recipient white, and shining with Fire, and that there are no more Flashes, or, as it were, Blasts of Wind, coming from Time to Time from the Retort, then the Work is finished. And you may, with Feather, gather the Fire together, or scrape it off with a Knife, where it sticks.

  The Fire is best preserved in a Vessel of Lead, closed up from the Air. But to be seen, ’tis also put into a Glass, in Water, where it will shine in the Dark…

  This was beginning to sound epic. Fifty or sixty pails of urine was an awful lot for a start. How long would it take me to produce that much? In fact, I learnt, I should be able to take some short cuts and stand a chance of replicating the experiment on a smaller scale. One bucket of urine–about three days’ worth–should contain around four grammes of phosphorus. This, if only I could separate it, would be more than enough to ignite ‘the Fire’.

  The first question was which urine to collect. The health guides always say it should be ‘straw-coloured’, as if everybody is intimately familiar with the colour of straw. Should I take this typical sauvignon blanc effluence? I decide it would be better to use the over-oaked chardonnay of the morning’s first pee. This strikes me as more likely to be rich in dissolved substances. I collect four litres and allow it to evaporate in an open container which I put out in the garden. It reeks at first, but gradually the disgusting smell disperses, and the liquor turns a rich ale brown. I am relieved to see that it shows no sign of breeding worms, not only because I have no especial wish to pick them out of the putrifying concentrate, but also because it implies that my sample is uncontaminated with stray organic matter, and that I might therefore be able to omit some of the repetitious purifying stages found necessary in the seventeenth century. After several weeks baking in the sun, all the liquid has evaporated, and I am left with twenty-two grammes of an almost odourless crystalline residue the colour of sawdust. This, I hope, is the reddish salt observed by Hooke.

  Now I am ready to begin the long roasting process. For this I need some more professional laboratory apparatus and advice. I enlist the help of Andrew Szydlo, one of my former chemistry teachers. Andrew is a man of many talents, and I remember him as always liable to whip out his gipsy violin in mid-lesson, or pass on some piece of lore about bee-keeping or car maintenance. More pertinently, he is an authority on alchemical history and the author of a treatise on Michael Sendivogius, the Polish alchemist who may have discovered oxygen in the early seventeenth century and contributed to its use in a pioneering manned submarine crossing of the Thames by the Dutchman Cornelis Drebbel in 1621. Andrew speaks with an English orotundity and the trace of a Polish accent and is wont to greet his former pupils as ‘professor!’ He is enthusiastic about the attempt to replicate this first isolation of a chemical element, and has set out various ingredients that may prove helpful in our quest, not least some best-quality gunpowder charcoal he has made from willow wood.

  We grind up some of my urine residue in a pestle and tip it into a test tube for heating. This tube is connected to an apparatus that will allow us to collect any distillate and test any gases that come off. Volatile material, including any phosphorus, should condense in a second test tube, while the gases escape through a vent. We aim two bunsen burners at the base of the loaded test tube, turn them up to maximum heat and wait. At first, a little water vapour comes off, which is followed by thick yellow curls that look and smell a little like burning tobacco. ‘Very curious,’ says Andrew in his demoniacal way. ‘It’s the most bizarre experiment, I must say.’ This vapour condenses as a tarry brown oil much like that which is produced when many forms of organic matter are burnt in this controlled fashion. At the vent, wisps of a white vapour appear. Could this be phosphorus pentoxide, the acidic combustion product of phosphorus? Litmus paper shows that unfortunately it is alkaline; another quick test with hydrochloric acid confirms that it is merely ammonia. We allow the solid remaining in the test tube to cool. It is now a dark slate grey. A flame test–a little of this solid dabbed on to a platinum wire and placed in a hot blue flame–reveals the characteristic yellow light of sodium and a fainter carmine red due to calcium. Andrew is now giving me a masterclass in analytical chemistry, which he intersperses with tirades against the parlous state of chemical education: how the school janitors are always seeking to clear away miscellaneous apparatus that they take to be scrap; how students are hardly allowed to do experiments for themselves these days; or, if they are, how the experiments must be designed so as to deliver a result before the end of the lesson–a constraint that slams the door on any slow-roasting chemistry such as this.

  Is the sodium just ordinary salt–sodium chloride–or is it perhaps, as we hope, a phosphate or phosphite salt, which would mean we’re getting closer to our goal? Dissolving a little of the grey residue in water and adding a drop of silver nitrate rapidly produces a dirty precipitate. This separates into a milky white sludge–the standard proof of chloride–and a mysterious brown residue, which does not dissolve either in acid or alkali, suggesting that it is rich in inorganic substances. This is probably where the phosphorus still lurks. We decide to reheat the residue mixed with the special charcoal Andrew has brought in order to reduce the phosphate or phosphite to elemental phosphorus. We grind the two materials together–the grey burnt urine residue and the black willow charcoal–and subject the mixture to the bunsen burners. ‘We’re roasting the pants off this now,’ he says gleefully.

  I am surprised when the residue, which has already spent an hour or so at the highest temperatures we can produce in a school laboratory, begins to react anew. Andrew explains that by grinding it with the charcoal we have greatly increased the surface of contact between the two materials, so boosting the chances of a
reaction. More ammonia comes off, followed by a gas that burns with a low blue flame when a taper is put to it. It is now dusk, and we turn out the laboratory lights in order to study the flame more closely. Could this be our phosphorus? It cannot be, because it would produce a thick white smoke of phosphorus pentoxide. It is presumably carbon monoxide burning to the invisible carbon dioxide. As the flame dwindles away in the darkened laboratory, it seems to betray a faint white edge in its dying moments. ‘We may just be beginning to get something,’ Andrew tells me. Limited now by the temperature–five or six hundred degrees Celsius–achievable with bunsen burners, we are brought up against the knowledge that Brand and his imitators used far hotter furnaces and ran the experiment for hours or days. We resolve to meet again, armed with quartz test tubes and an oxyacetylene torch that will enable us to turn up the heat.

  This time, it is immediately clear we are reaching a much higher temperature. The sequence of observations that we previously noted over an hour or more is repeated within minutes. Very soon, the roasted residue in the quartz tube begins to glow with a dazzling white light. Excitedly, we assume this may be our phosphorus, but the glow stays resolutely at the tip of the turquoise oxyacetylene flame where the heat is greatest. If it was really phosphorus, it would flow out of the tube in a vapour which would condense in the cooler second tube, as in Wright’s picture. It seems it is merely an incandescence produced by the extreme heat as it vaporizes the very substance of the quartz tube. We are forced to concede that, whatever his delusions, Brand was clearly a formidable experimental scientist.

  Joseph Wright of Derby painted The Alchymist in 1771. It was one of a number of scientific demonstrations he committed to canvas: his most famous work is probably An Experiment upon a Bird in the Air Pump done a few years earlier, in which a well-to-do household gathers in various states of wonder, horror and pity around a glass bulb from which the natural philosopher who stares unflinchingly out at us from the centre of the composition has evacuated all the air, extinguishing the life, or at least the temporary consciousness, of the bird inside.

  Wright was well connected with the Lunar Society in nearby Birmingham, whose members, including James Watt, the inventor of the steam engine, the physiologist and poet Erasmus Darwin and the chemist Joseph Priestley, met mostly on the full moon so that they could see their way home after evenings of ‘dinner and a little philosophical laughing’ that on occasion also included an experimental demonstration. Inspired by Robert Boyle’s work on vacuums in the 1650s, the painting seems also to anticipate Priestley’s experiments on the life-affecting properties of the new gases oxygen and carbon dioxide that lay a few years in the future–and the full moon shines in through the window. Other members of the society, such as the industrialists Josiah Wedgwood and Richard Arkwright, bought his work. With these paintings, Wright made his name as a recorder of the scientific Enlightenment.

  Like The Air Pump, The Alchymist is reimagined history. It purports to show the first making of phosphorus, which also happened more than a century before. Interpreted as an allegory, the painting seems to represent modern science shining its light into the alchemical darkness, a message that would naturally go down well with Wright’s patrons. However, the work appealed neither to them nor to contemporary viewers; it was unsold at Wright’s death in 1797. An astute analysis by the art and science historian Janet Vertesi attempts to explain its ‘curious failure’, and to account for the weird get-up of the protagonist. The painting balances three sources of light–once again the full moon outside, the radiant phosphorus pouring into the flask, and, on a bench in the background, the dimmer light of an oil lamp by which two laboratory assistants are going about their own business apparently oblivious of the miraculous scene unfolding before them. This trinity of lights may have religious significance, but it also symbolizes a contest between nature (the moon), the Enlightenment (the oil lamp) and some mysterious, more powerful, third force. The rational students of nature (they are in modern dress and use modern apparatus in contrast to their druidical master) toil by lamplight, but they are outshone, literally, by the light of the ignorant alchemist’s accidental discovery. Recall Wright’s carefully phrased title: ‘The Alchymist, In Search of the Philosopher’s Stone, Discovers Phosphorus…’ In other words, the alchemist, while doing whatever it is that alchemists are supposed to do, inadvertently makes a genuine contribution to science, a contribution moreover that the rationalists have failed to make for themselves. What kind of message did this send to the Enlightenment progressives of the Lunar Society in the fast-industrializing English Midlands?

  Science had the last laugh, however. Brand and the few rivals who eventually managed to repeat his experiment toured the courts of Europe with their precious luminous cargo. In England, Charles II attended a demonstration, as did Samuel Pepys and his fellow members of the Royal Society. John Evelyn wrote of how, while dining with Pepys in 1685, they witnessed ‘a very noble experiment’ in which two liquids were mixed to produce ‘fixed divers suns and stars of real fire, perfectly globular, on the sides of the glass, and which there stuck like so many constellations, burning most vehement’. But for a long time phosphorus remained little more than a high-end party trick. Obtaining it was arduous and obscure, and its elemental status was far from agreed, with chemical dictionaries sometimes listing it as no more than a ‘species of sulphur’.

  Exactly 100 years after Hennig Brand isolated phosphorus from urine, the Swedes Carl Scheele and Johan Gahn showed that it was a major constituent of bone. This richer source of the element made it possible at last to consider how it might be put to practical use. For more compelling than a mysterious light in nature is, as Keats observed, a light that may be captured by man. By the time that Keats was writing ‘Lamia’ in 1819, phosphorus lamps such as he describes were the latest thing, inventors having found a way to prevent the outright combustion of the phosphorus by diluting it in a suitable inert medium and regulating the admission of air. In this way, they were able to obtain a lamp that could give a steady glow on demand over a period of weeks. The discovery and application of phosphorus was well timed for the element to become a symbol of the taming of nature, of progress, and, literally, of enlightenment.

  The British returned Hamburg’s chemical gift to the world with vengeful interest during the last week of July 1943. In nightly raids, hundreds of aircraft dropped 1,900 tonnes of white phosphorus incendiary bombs on the city, the culmination of a strategy of ‘morale bombing’ authorized in 1941 by the prime minister, Winston Churchill, and Arthur Harris, the chief of the Royal Air Force Bomber Command, who sought to direct the aerial assault on locations most likely to weaken the spirit of the enemy. Increasingly, the manner of the bombing became a factor too, so that by the summer of 1943 the Allies’ objective was to destroy cities not only of historic and industrial importance, but also those densely populated with key workers, and to use means specifically designed to terrify the Germans into submission. This led to an unprecedented emphasis on incendiary bombs, and especially on phosphorus.

  On 27 July, the third night of the onslaught, incendiary bombing combined with the hot, still weather to produce a fire-storm, a phenomenon where the intensity of the conflagration sucks in air from all directions, so feeding the flames and creating a ferociously hot vortex of fire. In the words of a recent analysis by a German historian:

  The combination of the climate, the incendiary ratio, the collapsed defenses, and the structure of the city blocks created what Harris’s code-name ‘Gomorrah’ predicted: Like Abraham in Genesis 19:28, Harris looked toward the sinful city ‘and beheld, and lo, the smoke of the country went up as the smoke of a furnace.’ It melted between forty thousand and fifty thousand people.

  Many others were asphyxiated as the upward rush of flame simply sucked the air out of their underground shelters. Although the old town survived, the fires devastated much of the rest of Hamburg-Mitte, the central district where Brand had first isolated phosphorus nearly 300
years before. More than a quarter of a million dwellings were destroyed, along with factories, shipping and the all-important U-boat docks. Fifty-eight churches were reduced to rubble, but although its neighbourhood was badly hit, St Michaelis survived for another year until it was badly damaged in an American bombing raid. That autumn the trees of Hamburg flowered again as if it were spring.

  ‘Dropping phosphorus bombs on the homes of innocent civilians is never likely to happen again,’ writes John Emsley, while explaining that the element is nevertheless bound to remain a part of modern armouries because of its sheer versatility, used to illuminate targets, to create smokescreens or to ignite and clear vegetation. Yet as I write, in January 2009, Israel has admitted using white phosphorus during its offensive in Gaza. Israeli fire first hit a United Nations school, and a week later officials at the United Nations Relief and Works Agency for Palestinian Refugees in the Near East claimed that its Gaza City compound was set ablaze by phosphorus shells. In this conflict, as in others since the First World War, phosphorus is regarded as a legitimate agent of warfare, but its use is confined by convention to the open battlefield, and it is not allowed to be used against civilian populations. In Gaza, it happened that the ‘battlefield’ was densely populated: the smokescreen that phosphorus produces remains moral as well as literal.