The Knowledge: How to Rebuild Our World From Scratch (26 page)

Photography is a wonderful technique—a way of harnessing light to record an image, capturing an instant in time and preserving it for eternity. A holiday picture can trigger vivid reminiscences even decades later, but photography records the visual world with far greater fidelity than memory can ever offer. Yet beyond drunken party snapshots, family portraits, or breathtaking landscapes, the incomparable value of photography over the past two hundred years has been in presenting what the eye cannot see. It represents a key enabling technology across numerous fields of science, and will be vital in accelerating the reboot. Photography allows investigators to record events and processes that are exceedingly faint or occur over timescales too rapid or too slow for us to perceive, or at wavelengths invisible to us. For example, photography offers extended exposure times to soak up feeble light over far longer periods than the human eye can offer, allowing astronomers to study a multitude of dim stars and resolve faint smudges into detailed galaxies and nebulae.
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Photographic emulsions are also sensitive to X-rays and so allow you to create medical images for examination of the body’s interior.

The crucial chemistry behind photography is simple enough: certain compounds of silver darken in sunlight and so can be employed to record a black-and-white image. The trick is to create a soluble form of silver that can be spread evenly in a thin film, but then convert it into an insoluble salt that sticks on the outside surface of your photographic medium and doesn’t get washed away again.

First, coat a sheet of paper with egg whites containing some dissolved salt, and allow it to dry. Now dissolve some silver in nitric acid, which will oxidize the metal to soluble silver nitrate,
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and spread the solution over your prepared paper. The sodium chloride will react to create silver chloride, which is both light-sensitive and insoluble, and the egg albumin will prevent the photographic emulsion from soaking into the paper fibers. If you scavenge in the post-apocalyptic world, a single solid-silver teaspoon will contain enough of the pure element to produce over 1,500 photographic prints.

When light rays hit this sensitized paper, they provide the energy to liberate electrons in the grains and so reduce the silver chloride back to metallic silver. Large lumps of silver, such as a polished platter, have a bright luster, but a speckle of tiny metallic crystals scatters the light instead and so looks dark. On the other hand, areas of the sensitized sheet not exposed to light remain the white of the paper behind. The key follow-up step after the exposure is to kill this photochemical reaction and so stabilize the captured shadows. Sodium thiosulfate is the fixing agent still used today and is relatively easy to prepare. Bubble sulfur dioxide gas (
here
) through a solution of soda or caustic soda,
then boil with powdered sulfur and dry for crystals of “hypo” (a nickname derived from its old name, hyposulfite of soda).

Using a lens set into the front of a light-tight box to project an image onto sensitive paper on the back wall produces a photographic camera, but even in bright sunshine it can take many hours for this rudimentary silver chemistry to take a “photo.” Luckily, you can increase the sensitivity of your camera enormously with a developer—a chemical treatment that completes the transformation of partially exposed grains, reducing them entirely to metallic silver. Ferrous sulfate works well, and can be synthesized easily enough by dissolving iron in sulfuric acid. And as the chemical proficiency of the post-apocalyptic society improves, in place of chlorine salt you can substitute that of one of its atomic siblings, iodine or bromine, which produce far more light-sensitive photographic emulsions.

However, the fact that light exposure turns the photosensitive grains dark with silver metal, whereas shadows in the scene remain pale, means that the photo comes out tonally reversed from what your eye sees—you’ll get a “negative.” There is no fast-acting chemical reaction that produces a permanent positive image—no initially black substance that rapidly bleaches in sunlight—and so photography is burdened with this negative outcome. The necessary conceptual leap is to realize that if this reversed, negative image is created in the camera on a transparent medium, all that is needed is a second stage of printing out using the negative as a mask on top of sensitized paper, so that the pattern of highlights and shadows reverses again back to normal. The wet collodion process uses guncotton dissolved in a mixture of ether and ethanol solvents—all substances we have already come across in this book—to produce a syrupy, transparent fluid. It’s perfect for coating a glass plate with photochemicals, and then exposing and developing the image before it dries into a tough, waterproof film. And if instead you use gelatin (boiled out of animal bones, as we saw
in Chapter 5), it is possible to create a dry plate that is even more photosensitive and allows much longer exposure times.

Photography is a fantastic example of a novel application created by the fusion of several preexisting technologies, and using relatively straightforward materials and substances. Build a fireclay-lined kiln to produce your own glass by melting silica sand or quartz with a soda ash flux. Take one dollop to grind into a focusing lens and another to flatten into a rectangular pane for a negative plate, and draw on your papermaking skills to produce smooth prints. The chemistry underlying photography uses the same acids and solvents marshaled time and time again in this book, and you can take a primitive photo using substances derived from a silver spoon, a dung heap, and common salt. Indeed, if you fell through a time warp back to the 1500s you’d readily be able to source all the chemicals and optical components you needed to build a rudimentary camera, so you could show Holbein how to take a photograph of King Henry VIII rather than create an oil painting.

Filling in the periodic table of the elements, exploiting explosives, and using photography as a tool for rediscovery will all be important activities for a civilization restarting after the Fall. But as the post-apocalyptic society recovers and begins to flourish, it will need ever-increasing amounts of the basic substances we’ve discussed throughout this book. And to meet these demands, civilization will need to develop the more advanced processes of
industrial chemistry.

We often hear about the Industrial Revolution and the innovation of ingenious mechanical contraptions for alleviating the toil of humankind, thereby greatly accelerating the pace of progress and transforming eighteenth-century society. But the transition to an advanced civilization is as much about the invention of chemical processes for
the large-scale synthesis of the necessary acids, alkalis, solvents, and other substances critical to the running of society as it is about machinery for automating spinning and weaving and building roaring steam engines.

Many of the vital necessities of civilization that we’ve covered in this book rely on the same reagents to drive the transformations from raw materials gathered from the environment to the required commodities or products. And as the post-apocalyptic population swells over the generations of recovery, your ability to meet the demand for these critical substances by using the rudimentary methods we’ve looked at so far will become limiting, threatening to stall further progression.

We’ll focus here on the creation of two substances that became grievous choke points in our own history: soda in the late 1700s and nitrate in the late 1800s. Securing an adequate supply of both of these will also inevitably become critical for a rebooting post-apocalyptic society. So how can a recovering civilization release itself from the restrictions of relying on ashes for soda or manure for nitrates? Let’s start with the large-scale synthesis of soda, which formed the very beginnings of industrial chemistry in our history.

As we’ve seen, soda ash—sodium carbonate—is a vitally important compound used in a vast number of different activities throughout society. It’s indispensable as a flux for melting sand to produce glass (more than half of the global production of sodium carbonate today is used for glassmaking), and when converted to caustic soda (sodium hydroxide) it’s the best agent for driving the central chemical reactions in creating soap and separating the plant fibers needed for making paper. Glass, soap, and paper are central pillars of civilization, and since the Middle Ages we have relied upon a constant, cheap supply of alkali.

Traditionally, alkali was provided by potash from burned timber, and by the eighteenth century, widespread deforestation of Europe meant that potash came to be imported from North America, Russia,
and Scandinavia. For many applications, however, soda ash is preferred (the caustic soda made from it is a far more potent hydrolyzing agent than caustic potash), which was produced in Spain by burning the native saltwort plant, and along the Scottish and Irish coastlines from kelp washed ashore by storms. Sodium carbonate was also mined from dried lake-bed deposits of the mineral natron in Egypt. But by the second half of the eighteenth century, as the population and economy grew, the demand for soda began to outstrip the supply from these natural sources, as will inevitably happen again with a recovering post-apocalyptic society. Common sea salt and soda ash are chemical cousins,
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so can you convert an essentially limitless substance into an economically crucial commodity?

One simple two-step operation, developed by the eighteenth-century French chemist Nicolas Leblanc, is to react the salt with sulfuric acid and then roast the product with crushed limestone and charcoal or coal in a furnace at around 1,000°C to form a black ashy substance. The sodium carbonate you are interested in is soluble in water, so you can soak it out using exactly the same technique you would use when working with seaweed ashes. However, while this Leblanc process is a readily achievable way of morphing salt into soda, releasing you from the limitations of burned plants or mineral deposits, it is grossly inefficient and churns out noxious waste products.
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So
ideally the rebooting civilization would want to be able to leapfrog straight over the easy but wasteful LeBlanc process to a more efficient system.

The Solvay process is slightly more intricate, but ingeniously employs ammonia to close the loop: the reagents it uses are recycled back within the system, minimizing wasteful byproducts and thus also pollution. The chemical reaction sitting at the core of the Solvay process is this: when a compound known as ammonium bicarbonate is added to strong brine, the bicarbonate ion swaps over onto the sodium to form sodium bicarbonate (identical to the rising agent used in baking), which can then simply be heated to change to soda ash. The first step in achieving this is to pass the strong brine through two towers, with first ammonia gas and then carbon dioxide bubbling up through them to dissolve in the salty water and combine to make the crucial ammonium bicarbonate. The swapping reaction occurs with the salt, creating sodium bicarbonate, which doesn’t dissolve and so settles as a sediment to be collected. The ammonia is the key ingredient for this stage, as it keeps the brine nicely alkaline and so ensures that the bicarbonate of soda can’t dissolve, neatly separating these two salts.

The carbon dioxide needed for this initial step is baked out of limestone in a furnace (in exactly the same way we saw in Chapter 6 for burning lime for mortar and concrete production). The quicklime left behind is itself added to the brine solution after the soda has been extracted and regenerates the ammonia bubbled in originally, ready to be used again. So overall, the Solvay processes consumes only sodium chloride salt and limestone, and alongside the valuable soda produces just calcium chloride as a byproduct, which itself is used for spreading as a de-icing salt on wintry roads. This elegantly self-contained system, cleverly recycling the key ammonia as it goes and built using only fairly rudimentary chemical steps, is still the major source of soda worldwide today (except in the United States, where a large deposit of the sodium carbonate mineral trona was discovered in Wyoming in the 1930s). And for a recovering civilization the Solvay process presents a marvelous
opportunity to leapfrog over less efficient and noxiously polluting alternatives for producing vital soda to form the foundation of a post-apocalyptic chemical industry.

A SODA PLANT IN NEW YORK IN THE LATE NINETEENTH CENTURY, OWNED BY THE SOLVAY PROCESS COMPANY (TOP). THE FOUR STEPS OF THE SOLVAY PROCESS FOR ARTIFICIALLY SYNTHESIZING SODA (BOTTOM). THE RECYCLING OF AMMONIA SITS AT THE HEART OF THIS CRITICAL CHEMICAL PROCESS.

The Solvay process converts an abundant source of the element sodium, common salt, into the crucial alkaline compound soda. But before too long an advancing civilization will run into problems of supply limitation of another critical commodity. One of the most fundamental chemical processes for all of us alive today involves the element nitrogen—and another miraculous transformation of a common base substance into something vitally valuable.

In terms of the sheer number of people it directly affects every day, the most profound technological advance of the twentieth century was not the invention of flight, antibiotics, electronic computers, or nuclear power, but the means to synthesize a humble, foul-smelling chemical: ammonia. As we’ve seen throughout this book, ammonia and the related (and thus chemically interconvertible) nitrogenous compounds nitric acid and nitrates are foundation stones of the chemistry underpinning civilization. Nitrates are absolutely vital for the production of both fertilizers and explosives, but by the closing years of the nineteenth century, the industrialized world was running out: demand was starting to exceed supply, and America and European nations became concerned not just about ensuring munitions for their armies, but about fundamentally providing enough food to keep their citizens alive.

For millennia, the response to an expanding population had been simply to clear more ground for cultivation. Once you’ve hit the limit of available land, though, the only solution to feed the multiplying mouths is to increase the yield of crops from the same cultivable area. As we saw in Chapter 3, returning manure to the soil and planting legumes are both effective. But when a population reaches a certain limit—a capacity crowd, if you like—civilization inescapably hits a snag. You cannot produce more manure from livestock, as animals would need to be fed plants grown on the land in the first place; and
you can’t sow more fields with legumes, as that decreases the land available for growing cereal crops. You’ve hit the carrying capacity of organic agriculture.

The only recourse is to inject external nitrogen from outside of the agricultural loop. Through the nineteenth century, Western agriculture relied heavily upon imported guano and on saltpeter mined from the Chilean desert.
But these sources rapidly became depleted, and the president of the British Association for the Advancement of Science, Sir William Crookes, warned in 1898 that “we are drawing on the Earth’s capital, and our drafts will not perpetually be honored” (a caution we would be wise to heed today as our civilization’s voracious appetite for crude oil and other natural resources threatens to exhaust them). The world we leave behind will already have been stripped of these natural deposits of nitrate, and a maturing post-apocalyptic civilization will slam into this wall soon enough.

The planet’s atmosphere is rich with nitrogen gas—it makes up almost four-fifths of every breath you take—but it’s recalcitrantly unreactive. The two atoms of nitrogen are locked firmly together by a triple bond; in fact,
nitrogen gas is the least reactive diatomic substance known. This makes it very difficult to convert nitrogen gas into accessible forms—to “fix” it. By the end of the nineteenth century it had become clear that figuring out how to fix nitrogen was critical to the progress of civilization itself—that chemistry had to come to the rescue of humanity.

The solution, discovered in 1909 and still used today, is known as the
Haber-Bosch process. On the face of it, the process appears beguilingly simple. Nitrogen, the most common gas in the Earth’s atmosphere, and hydrogen, the most abundant element in the entire universe, are the only raw materials and are mixed together in a one-to-three-parts recipe within a reactor and combined to form NH
₃
—ammonia. Nitrogen can be simply sucked out of the air, and today
hydrogen is made from methane natural gas, but it can also be gathered from the electrolysis of water. Coaxing nitrogen into participating requires severing the sturdy bonds clamping the twin atoms together, and that requires a catalyst. A porous form of iron, with potassium hydroxide (the caustic potash we covered
here
) as a promoter to increase its effectiveness, works well for hustling this reaction along. The reaction is never complete, so the gases are cooled for the desired product to condense as ammonia rain to be drained off and stored, and the as-yet-unreacted gases are recycled back through the reactor repeatedly until virtually everything is successfully transformed. But, as with many things, the devil is in the details, and the Haber-Bosch process is actually pretty tricky to pull off.

Many chemical reactions are essentially unidirectional: they are a one-way street of reactants recombining into products. For example, in a burning candle the waxy hydrocarbon molecules are oxidized by the combustion process into water and carbon dioxide, but the reverse transformation would never occur spontaneously. Other chemical processes, however, are reversible reactions, and the two opposing conversions run in both directions simultaneously. “Reactants” are transformed into “products,” but these are being converted back again at the same time. The conversion between a nitrogen-hydrogen mixture and ammonia is one such reversible process, and to tip the balance toward the desired compound, you need to carefully arrange the conditions within the reactor. For producing ammonia, this means running at a high temperature (around 450°C) and crushingly high pressure (around 200 atmospheres). And these extreme conditions for the reactor and pipework are why the Haber-Bosch process is so troublesome to run. Far more so than the other crucial processes we’ve looked at that require the heat of a furnace—such as glassmaking or metal smelting—pulling off nitrogen fixation is a feat of accomplished engineering. If your post-apocalyptic society can’t salvage a suitable reactor vessel, you’re
going to need to learn how to construct your own industrial pressure cooker.