No other rust battle in America has been fought so visibly, contentiously, or been celebrated so grandly. On July 4, 1986, millions of people showed up for the centennial celebration, as did 40,000 boats—including an aircraft carrier, the Queen Elizabeth II, and more tall ships than had ever gathered together before. There were so many boats in the harbor that the Staten Island Ferry took twice as long as normal to weave through them. Queens and Staten Island made available 10,000 camping spots. Bleachers were erected. Governors Island became the VIP island, where the Secretary of the Navy sat. On Liberty Island, Walter Cronkite performed as the master of ceremonies, and Nancy Reagan cut a ribbon. Moffitt sat there. Baboian sat there with his wife. Long-haired Ed Drummond was not invited, even though he had forced Moffitt to take a close look at the statue with binoculars, then announced, in bold red letters that needed little poetic interpretation, that liberty was framed. Nobody ever thanked him.
The evening before the big celebration, Cardinal John O’Connor held an ecumenical mass at St. Patrick’s Cathedral. That day, Chief Justice Warren Burger swore in 250 new US citizens on Ellis Island. The Boston Pops played in New Jersey. The New York Philharmonic Orchestra played in Central Park. Events were held at the Meadowlands. Sinatra didn’t show up. The weekend alone cost just under $40 million. So many visitors flocked to the statue, and were forced to stand in such long lines, that a riot almost ensued. Almost a third of the world’s population saw the ceremony on TV.
As Grover Cleveland had shown up on a boat for the dedication a hundred years earlier, that’s how President Reagan planned to arrive. He wanted to show up aboard the USS John F. Kennedy and “relight” the torch with a laser beam. Instead, he performed the ceremony from Governors Island. In any case, the achievement was no less triumphant than it had been a century before; in both cases, the Statue of Liberty exemplified the greatest fusion of engineering and art of the day. The occasion was followed by the largest fireworks show ever—twenty tons of fireworks, launched from forty barges, by a pyrotechnics partnership called the All-American Fireworks Team.
That may be the best symbol of all: planned oxidation, commemorating defeated oxidation. Which makes you wonder: had people known about the rust battle, would the celebration have been even grander? What’s more impressive: liberty or engineering? Philosophy or power? Belief or might? History or science? From the metal’s point of view, there was nothing democratic going on. The metal had become part of a totalitarian regime, planned, controlled, observed, denied the opportunity to do what it yearned most to do. It would be a strange thing to celebrate the metal’s fate. Better to focus on the fireworks.
Today, a plaque, installed by the National Association of Corrosion Engineers, marks the site. Below the NACE logo—a triangle within a circle, with two fig leaves around it—there’s this text:
THE STATUE OF LIBERTY
HAS BEEN SELECTED BY THE
NATIONAL ASSOCIATION OF
CORROSION ENGINEERS
AS A
NATIONAL CORROSION RESTORATION SITE
AS AN EXAMPLE OF MAN’S TECHNOLOGICAL
ACCOMPLISHMENTS TO CONTROL CORROSION
APPLIED TO A HISTORIC STRUCTURE SO
THAT FUTURE GENERATIONS CAN BENEFIT
FROM THE SYMBOLIC HISTORY OF THE
STATUE AS THE WORLD’S BEST-KNOWN
MONUMENT TO MAN’S SEARCH FOR FREEDOM
AND LIBERTY.
PRESENTED TO THE
NATIONAL PARK SERVICE
OCTOBER 28, 1986
IN COMMEMORATION
OF THE STATUE’S 100TH BIRTHDAY
It is a unique plaque, marking the only National Corrosion Restoration Site in the country. But it won’t be the last.
2
SPOILED IRON
The first recorded words about rust express what you might expect: exasperation. The words belong to a Roman army general. Two thousand years ago, during the Blue Nile campaign, he complained about corrosion in his giant catapults. “The pentle hooks on the onagers are weakened so badly by corrosion,” he wrote, “that the arbalests are causing more casualties in our own army than to the enemy.” A generation later, at a loss to explain how rust worked, Pliny the Elder figured, metaphysically, that the benevolence of nature had inflicted the penalty of rust to limit the power of iron, thus making nothing in the world more mortal than that which is most hostile to mortality. He called rust ferrum corrumpitur, or “spoiled iron.” Only in myths was rust less troublesome. It helped Iphicles father a son, and it healed the stubborn wound in Telephus’s thigh. Rust confounded the rest of us.
Robert Boyle, the tall, wealthy English “father of chemistry,” took up an investigation of rust during the seventeenth-century reign of King Charles II. He began by insulting Pliny: “I have not found among the Aristotelians,” he wrote, “so much as an offer at an intelligible account.” Born a year after the death of Francis Bacon, Boyle talked metaphysics with Sir Isaac Newton, hung out with the founders of the scientific group known as the Royal Society, and was in Florence, Italy, when Galileo died. He taught himself Hebrew, Greek, and Arabic so that he could read original sources. He conducted medical experiments on himself, and tasted his own urine. By the end of his life, he’d written more than two and a half million words. In his 1675 treatise Experiments and Notes About the Mechanical Origine or Production of Corrosiveness and Corrosibility, Boyle summed up rust as a mechanical phenomenon resulting from the “congruity between the agent and the patient.” For a metal to be corroded, it had to be “furnish’d with pores of such bigness and figure, that the corpuscles of the solvent may enter them.” These “pores” also explained why light penetrated glass. (That same year, Boyle also published an account of the “transmutation” of mercury into gold.) By his own intelligible standard, Boyle failed, too, only more discursively than Pliny, and with a higher word count.
His twenty experiments were not for naught, though. He found that salt, by itself, didn’t corrode lead nearly as fast as saltwater did. By pouring saltwater, lemon juice, vinegar (which had “edges like blades of swords”), urine, turpentine, lye, and various acids onto lead, iron, mercury, copper, antimony, and tin, he showed that all metals were vulnerable. Silver, for example, fell victim to nitric acid. Even gold corroded when subjected to a mixture of nitric and hydrochloric acids called aqua regia.
Today we know that only a handful of rare metals don’t corrode: tantalum, niobium, iridium, and osmium. The others—all of them—can be invited, urged, or forced to react with oxygen. Some react spontaneously in air or water. Some, like aluminum, chrome, nickel, and titanium, form a thin outer layer of protective metal oxide, and then call it quits. Many of the corrosion-resistant metals are named in honor of Greek gods or kings, for no other entity could have created such marvelous stuff. Nevertheless, most metals met oxygen long ago, which explains why precious few metals present their naked selves anywhere on Earth. (This also explains why oxygen did not accumulate in the atmosphere for billions of years, until rocks on the surface had reached their fill.) Three-quarters of the universe’s elements are metals, and nature, apparently, abhors almost all of them.
A few lucky ancient wanderers found rare but marvelously resilient chunks of metal. Shiny and strong, the stuff was perfect for Inuit spears, or Sumerian shields, or Tibetan jewelry. It was a nickel-iron alloy, not unlike stainless steel, and it came from the sky, in the form of meteorites. The world’s largest meteorite, in fact, in Namibia, is of this variety. Today we call it “meteoritic iron,” but long before stainless steel was invented, one person called this “celestial metal” an advantage “providentially placed before us.” The disadvantage placed providentially before everyone else was rust.
Rust, of course, is the corrosion of iron, while corrosion is the gnawing away, thanks to oxygen, of any metal. To the horror of engineers, I use the word colloquially.
As a result of rust, oxygen commands love and hate. It contributed to the filling of the oceans, got life beyond a
green slime, and assisted with the evolution of two sexes, earning one modern biochemist’s nickname “the molecule that made the world.” Just two centuries earlier, though, on account of oxygen’s potency and ubiquity, a less impressed chemist called it “the fire that burns up all things slowly.” It does not get along with metals, or rather, the way it gets along with metals, as Boyle demonstrated, does not agree with us.
But oxygen wasn’t discovered until a couple of years before the United States declared independence from Britain, and wasn’t recognized as the culprit in rust for another fifty years. Until then, experiments relating to rust caused confusion and enlightenment in equal proportions.
In Bologna, Italy, Luigi Galvani figured out that dissimilar metals could be used to create sparks—enough to make the legs of a frog twitch—and rust. Two hundred miles northwest, in Como, a physics professor named Alessandro Volta figured out how to make a battery by stacking sixty different metals in a tower, but didn’t know what the metals had to do with it, or why the current eventually ebbed. Using such a battery—a “voltaic pile”—the frantic British chemist Sir Humphry Davy demonstrated a technique named after his Italian colleague. By galvanizing steel—coating it in a thin layer of zinc—he could protect it from rust.
At the turn of the century, though, chemistry was still muddled with alchemy and philosophy. There were about a dozen known elements, and three recognized processes: combustion, respiration, and oxidation. Combustion and oxidation were explained by a mysterious agent called phlogiston. Heat was understood as a weightless fluid called caloric. An experimentalist at heart, Davy didn’t like either theory. In 1806, after more experiments with batteries, he coined the term electrochemistry. The next year, he discovered sodium and potassium, new metals of the reactive variety. Colleagues compared him to Boyle, and the government saw him as society’s savior. He was asked to investigate explosions in mines, and after studying the combustion of methane, designed a miners’ safety lamp that, because it was surrounded by iron mesh, dissipated heat and eliminated explosions. It made him famous.
In 1823 the Royal Navy commissioned Davy to solve the problem of corrosion on their warships. Copper sheathing on ships protected wooden hulls from destruction by worms and rot, and prevented the adhesion of barnacles, weeds, and other sea life, which slowed down ships. By then, Davy had been knighted, and had successfully campaigned for the presidency of the Royal Society. He began by putting copper in beakers of seawater and investigating the greenish precipitate. Deducing that it formed with oxygen, he figured the oxygen had to be coming from the water or the air, and since no hydrogen was produced, it couldn’t be coming from the water. It was the air. To prove it, he showed that very salty water—brine—didn’t dissolve the copper as fast as regular seawater, because it contained less dissolved oxygen. (He’d already shown years before that deoxygenated water killed fish.)
Then, with the principles of galvanizing in mind, he attached nails made of zinc and iron—“more oxidable metals”—to sheets of copper, and put them in the sea. Davy spent months experimenting at Chatham and Portsmouth, studying the ratios of the metals. Finally, in January 1824, he announced his findings. “A piece of zinc as large as a pea, or the point of a small iron nail, were found fully adequate to preserve forty or fifty square meters of copper,” he wrote. A zinc-to-copper ratio anywhere between 1:40 and 1:150 prevented corrosion. Between 1:200 and 1:400, corrosion was slowed down, but not prevented. Less than that had little effect. It didn’t matter where the zinc was placed, as long as it was connected to the copper. He called his results “beautiful and unequivocal.”
The navy tested his method on the HMS Comet. Months into the trial, the ship wasn’t corroding, but was fouling with barnacles far worse than before. The navy was annoyed, and the Royal Society was embarrassed. Letters in newspapers tarnished Davy’s reputation. But Davy had been right. His sacrificial anodes—still widely used on boats today—prevented corrosion.
Michael Faraday, the renowned experimentalist and father of electromagnetism, whom Davy had hired as an assistant in 1813, took Davy’s work even further and, in the 1840s, determined that electrical current—not a fluid at all—could be used advantageously to prevent corrosion. “All chemical phenomena,” he wrote, “are but exhibitions of electrical attractions.”
For much of the nineteenth century, though, chemists focused elsewhere—on the composition of molecules, on using spectroscopy to identify and isolate and eventually describe new elements. At the end of the century, most believed that acids, carbonic acid in particular, were responsible for corrosion. (Acids were only part of the story.) Others thought hydrogen peroxide was somehow involved. Some thought that imperfections in metal were to blame, and that perfectly pure metals wouldn’t corrode.
Not until the first half of the twentieth century did corrosion theory take shape. It started with Swiss chemist Julius Tafel, a handsomely bearded insomniac, who in 1905 related current and voltage to the rate of a chemical reaction. He killed himself before chemists Johannes Bronsted, Martin Lowry, and Gilbert Lewis, in 1923, each came up with the notion that chemical bonding resulted from the pairing up of acids and bases, which, depending on how you looked at it, either donated/received protons, or donated/received electrons. Three years later, Linus Pauling and Robert Mulliken—both future Nobel Prize winners—began quantifying the tendencies of elements to attract electrons, a property called electronegativity. Because every element is structured differently with its electrons orbiting its neutrons, each has a unique electronegativity, though many are similar. Little known francium, on the lower left corner of the periodic table, is the least electronegative, while fluorine, on the opposite corner, is the most electronegative. The scale runs from 0.7 to 4, in Pauling units. Fluorine, reacting with everything, steals electrons with fury. Oxygen, a gas five hundred times more abundant on earth than fluorine, is the second most electronegative element—and this explains why life relies on it. For transporting energy, it is the best thing going. (Aerobic metabolisms, which rely on oxygen, are fifteen times more efficient than anaerobic metabolisms.) The third most electronegative element is chlorine. Considering that two-thirds of the world is covered by a liquid containing a great deal of dissolved chlorine, it’s almost as if God stacked the cards against admirals intent on employing metals in their service. God didn’t make calm air or seas.
In sort of a mirror of electronegativity, metals can be ranked in nobility. Noble metals don’t give up their electrons, no matter how electronegative the other elements may be. The most noble metals—gold, platinum, iridium, palladium, osmium, silver, rhodium, and ruthenium—are also the most valuable, and this is no coincidence. They’re valuable because they’re reliable. They don’t corrode. The nobility of a metal is measured in volts, from 1.18 (platinum) to -1.6 (magnesium).
Therein lay the source of the current in Volta’s piles of metals: the unnoble metals were giving up electrons to the noble ones. A quarter-volt difference is enough to compel an electron to migrate. There’s a quarter volt between lead and titanium, and there’s a quarter volt between tin and silver. The quarter volt between iron and copper is what saved the copper on the bottom of the HMS Comet, and what saved the Statue of Liberty’s skin at the expense of its frame. The quarter volt between aluminum and steel is what causes seat-posts to rust firmly into place on bicycles, and what caused the white powder around the rivets in Syzygy’s mast.
In other words, the unnoble metals are anodic. Paired with more cathodic (in other words, more noble) metals, they sacrifice themselves at the altar of physics. Evans, of Wimbledon, diagrammed this back in Pauling’s day. That’s why galvanizing works: you’re giving nature something to chew on for a while.
Oxygen wants to steal electrons, and so get reduced. Unnoble metals give them up, and so get oxidized, especially when compelled by more noble metals. Water’s the convenient pathway for the particles and elements to zip around in. In 1938 Carl Wagner and Wilhelm Traud described this
phenomenon, putting together Tafel’s observation. They said that the sum of the charges lost and gained in a corrosion reaction was zero, and that each metal would corrode at a new rate. It became known as the mixed-potential theory of galvanic corrosion.
By then chemists also recognized that most of one volt was enough to compel electrons to stay put. In other words, if a pipeline operator pushed 0.85 volts into his buried pipeline, he could convince the electrons in the steel not to be lured elsewhere.
Together, this cathodic protection and anodic protection form half of the techniques available to combatting corrosion. The third arm of defense, far blunter, precedes them both. It’s paint. If you can stop oxygen (and water, which contains oxygen) from getting to metal, you can stop corrosion. The fourth arm is sort of a modern version of paint. Inhibitors, binding to metal before oxygen has a chance to, work just as well in abetting a brown outcome. Many are synthetic, but they’ve been made from mangos, Egyptian honey, and Kentucky tobacco. Anodizing—intentionally oxidizing the surface of aluminum by dipping it in acid and applying current—works because the thick oxide layer is then sealed with an inhibitor. Electroplating with a metal more durable than zinc—cadmium, chromium, nickel, or gold—is sort of the rich-man’s galvanizing.
Of course, subtleties in the theory of corrosion abound. Poorly mixed alloys may become anodic and cathodic to each other in opposite corners of the same piece. Once electrons begin to flow, the disparity only increases, and corrosion accelerates. Metallurgists using X-ray crystallography nearly a hundred years ago figured that out.
In the same year that Wagner and Traud described galvanic corrosion, Marcel Pourbaix, of Brussels, Belgium, came up with thermodynamic diagrams of corrosion reactions. He broke down the oxidation and reduction reactions in a metal across the full range of electrochemical conditions, from super acidic to super basic. The resulting graph showed zones of corrosion, passivity, stability. It showed where a metal was safe and where it was under threat—revealing why the acids that Boyle toyed with were bad news for metals. Pourbaix did this for each element, publishing his results in his Atlas of Electrochemical Equilibria in Aqueous Solutions in French in 1963 and in English three years later. Then this pioneer of the field traveled around the world talking about the new science of rust.
Rust: The Longest War Page 5
