Sun in a Bottle_The Strange History of Fusion and the Science of Wishful Thinking

by Charles Seife

  There was only a handful of known elements, and they each had different properties. For example, the atoms of some elements, such as hydrogen, oxygen, and carbon, were very light. Other elements’ atoms, like those of mercury, lead, and uranium, were very heavy. And these particles—these atoms—were fixed in their properties; it was impossible to transmute an atom of hydrogen, say, into an atom of lead.

  This picture explained the nature of matter extremely well. Within a century, atomic theory changed the subject of chemistry from a quasi-mystical hodgepodge of contradictory ideas into a real science. Physicists soon joined the chemists in their support of atomic theory; they began to provide evidence for the existence of tiny atomic particles. Theorists like Ludwig Boltzmann realized that you could explain the properties of gases simply by imagining matter as a collection of atoms madly bouncing around. Observers even saw the random motion of atoms indirectly: the jostling of water molecules makes a tiny pollen grain swim erratically about. (Albert Einstein helped explain this phenomenon—Brownian motion—in 1905.) Though a few stubborn holdouts absolutely refused to believe in atomic theory,14 by the beginning of the twentieth century the scientific community was convinced. Matter was made of invisible atoms of various kinds: hydrogen atoms, oxygen atoms, carbon atoms, iron atoms, gold atoms, uranium atoms, and a few dozen others. But, as scientists were soon to find out, atoms are not quite as uncuttable as the ancient Greeks thought. Indeed, to figure out why different elements have different properties, physicists had to slice the atom into pieces.

  The first piece came off in 1898. The Cambridge physicist J. J. Thomson was studying a mysterious phenomenon known as cathode rays. He used electric and magnetic fields to deflect the rays and came to the correct conclusion that the rays were made up of negatively charged particles that had been stripped away from atoms. These very, very light particles came to be known as electrons.

  Since an atom is, on balance, neither positively nor negatively charged, the positive and negative charges in the atom must be equal and opposite; the charges in the atom have to cancel each other out. This means that for every electron in an atom, there has to be something else in the atom that carries the equivalent positive charge. About a decade after the discovery of the electron, the physicist Ernest Rutherford found out where that equal and opposite charge sits. It resides in tiny, but extremely solid, nucleus at the very center of the atom. This nucleus is quite heavy, thousands of times heavier than an electron, so the nucleus of an atom had to be made of stuff very different from electrons. Rutherford soon figured out what that positively charged stuff was: he realized that the positive charge is cloistered inside a heavy particle known as a proton.

  For every electron zipping around in the outer regions of the atom, a proton had to be sitting in the nucleus. Since positively charged objects attract negatively charged ones, the nucleus attracts the electrons through electrical forces, in roughly the same way that the sun attracts its planets with gravitational forces. Rutherford took this analogy fairly literally; he imagined the atom to be like a miniature solar system. At the center is a heavy, dense, positively charged nucleus. Quite a distance away, lighter, quick-moving, negatively charged electrons are in “orbit” around it.15 In between, there is empty space—lots of it.

  When physicists discovered the proton and electron, they sparked a revolution in the scientific understanding of matter. Two subatomic particles suddenly explained the properties of the elements. No longer were atoms of different elements considered to be fundamentally different objects; an atom of gold need not be thought of as a different sort of creature compared with an atom of lead. Gold and lead were essentially the same kind of object: bundles of protons surrounded by bundles of electrons. Gold has properties different from those of lead—and they both have properties different from the other elements—because they have different numbers of protons in their nuclei (and, hence, different numbers of electrons). A hydrogen atom has one proton per nucleus, helium has two. Oxygen has eight; gold, forty-seven; lead, eighty-two; uranium, ninety-two. In each case, the number of protons in a nucleus—known as an atom’s atomic number—determines how the atom behaves chemically. It tells you which atoms it will react with and which it won’t; it tells you whether a collection of atoms is likely to be a gas or a metal, whether it will burn in oxygen or explode in water or refuse to react with anything at all. This theory was a tremendous success for science. The uncuttable atom had been dissected into its component parts. But one piece was still missing.

  The discovery of the electron had come from Thomson’s investigations into cathode rays. Cathode rays come from a fairly simple piece of laboratory equipment: put a couple of pieces of metal in a vacuum tube, hook them to a battery, and radiation streams from one end to the other.

  The concept of radiation was a new phenomenon at the turn of the twentieth century. Scientists knew little about it, but they were beginning to detect it everywhere. Marie Curie’s radium emitted a substance—particles or rays or something as yet unknown—that carried energy; something was fogging a photographic plate. That was one kind of radiation. The German scientist William Roentgen discovered another kind in 1895. When he sent electrical current through an evacuated tube, he noticed it would generate invisible rays that could make fluorescent screens glow. Like the rays coming from radioactive elements like radium and uranium, Roentgen’s x-rays could expose a photographic plate. X-ray radiation, too, carries energy. (It turned out that x-rays are beams of light so energetic that they pass right through flesh.) Then there were the mysterious rays coming from Thomson’s cathode. By the turn of the century, scientists across the world were finding all sorts of rays in strange places. The scientific world was going radiation crazy.

  We now know that these “radiations” are not all the same thing. Some, like x-rays, are varieties of light. (Gamma rays, too, are light beams even more energetic, and more penetrating, than x-rays.) These high-energy light rays penetrate matter relatively easily. Not all the radiations had this property. Thomson’s cathode rays couldn’t penetrate very far into an object before being absorbed. Neither could beta radiation, another type of emanation that streams from certain kinds of unstable atoms. Alpha radiation, which comes from yet other varieties of unstable atoms, penetrates even less than beta rays. It turns out that cathode rays, beta rays, and alpha rays are all subatomic fragments. Cathode rays and beta rays are both made up of electrons; alpha rays are made up of heavier, positively charged pieces of large atoms.16

  Not surprisingly, researchers were so excited about finding new kinds of radiation that some of their discoveries were entirely fictional. In 1903, the French physicist René Blondlot thought he had discovered a new type, which he dubbed the “N-ray.” But Blondlot had deceived himself; his desire to believe in N-rays made him ignore the evidence against them. When a skeptical researcher removed a crucial component of the experimental apparatus and the unsuspecting Blondlot continued to observe the N-rays, N-rays were exposed as a fiction. Blondlot was made a laughingstock.

  There was one type of radiation, though, that did not fit neatly into the pattern scientists had been seeing. High-energy light, such as x-rays or gamma rays, penetrates matter easily; its beams are hard to block. Fragments of atoms—charged particles like protons and electrons and alpha particles—tend not to penetrate matter much at all. Because of their charge, they get tangled in the electrons and protons in a given hunk of matter and quickly slow to a halt. But a new type of radiation, discovered in the 1930s, seemed like a weird cross between light and atom fragment. Scientists generated this bizarre radiation by shining a beam of alpha particles upon certain kinds of atoms (such as beryllium atoms). This new kind of radiation did not have an electric charge: it was unaffected by electric or magnetic fields. It penetrated matter as readily as gamma rays did, but it did not behave as a light beam should. It behaved like a heavy particle: it would hit a block of paraffin and knock protons out; mere light couldn’t do that so easily.
In 1932, the British physicist James Chadwick concluded, correctly, that this new type of radiation consisted of particles almost identical to protons but for one major difference: they had no electric charge. Chadwick won the Nobel Prize for his discovery: the neutron.

  The neutron is just a tiny fraction of a percent heavier than a proton, so it has quite a bit of oomph. But because it is electrically neutral, it doesn’t “feel” the electrical charges of the electrons and protons in a material. It is only affected by an atom when it slams directly into the nucleus. However, since atoms are mostly empty space and atomic nuclei are very small, a neutron can zoom straight through a chunk of matter without ever encountering something that deflects it. Neutrons penetrate matter extremely well, going through lead bricks almost as if they didn’t exist. But when a neutron does, by chance, hit an atomic nucleus, it packs a punch. A light atom (such as hydrogen) might be kicked out of the substance altogether. A heavy atom (such as uranium) might shiver and break apart when struck with the right amount of force. (As described in chapter 1, neutrons doing just this is what causes the chain reaction at the heart of the atom bomb.)

  The discovery of the neutron also solved a puzzle that was beginning to vex physicists. When chemists and physicists used their newfound knowledge of protons and electrons to understand the nature of the elements, they were surprised by a strange inconsistency. They discovered that an atom of a given element did not have a fixed weight. For example, in 1932 scientists found that hydrogen came in several flavors. There was ordinary hydrogen—which was thought to be made up of one proton (and one very light electron, whose weight is negligible). Then there was a heavier form of hydrogen that weighed twice as much. They called it deuterium. Soon, they realized there was yet another version that weighed almost exactly three times as much as hydrogen: tritium. All three of these varieties had the same chemistry as hydrogen, but they all had different weights. (And tritium, as it turned out, was radioactive.) Until the neutron was discovered, nothing could explain why a single element could have multiple weights.

  Once Chadwick discovered the neutron, though, the answer to the puzzle was obvious. Scientists already knew that the number of protons determined the chemical properties of an atom; hydrogen, deuterium, and tritium each had a single proton in the nucleus, so they were almost identical, chemically speaking. But neutrons can also sit in an atom’s nucleus. Because neutrons don’t have a charge (and don’t attract extra electrons), they don’t affect an atom’s chemical behavior; an extra neutron doesn’t turn hydrogen into a different element. But an extra neutron makes that hydrogen weigh more than before.

  Ordinary hydrogen’s nucleus is simply one proton. It weighs as much as one proton, so it is known as hydrogen-1, or 1H. Deuterium’s nucleus, too, has one proton. But it also has a neutron that weighs roughly the same as the proton; the mass of the nucleus (hence, the mass of the atom) is doubled. Deuterium is thus known as hydrogen-2, 2H. Tritium has a single proton in its nucleus, but in addition it has two neutrons, making it three times as heavy as ordinary hydrogen. Tritium is therefore designated hydrogen-3, 3H. All these atoms are considered to be varieties, or isotopes, of hydrogen. In a chemical reaction, all three behave more or less the same way. But they have slightly different physical properties by virtue of their nuclei’s different weights.

  Scientists were thrilled when they discovered the neutron because it gave them a complete model to explain an atom’s chemical behavior. Just figure out how many protons and neutrons are in a given atom and you can predict its properties extremely well.

  Despite the spectacular success of atomic theory, scientists, in some sense, were astonished that atoms could exist at all. Nuclei are finicky things, and it is amazing that any of them are stable. By rights, they should fly apart instantly. They are filled with positively charged protons, and positively charged things repel one another. If the protons in a nucleus were to obey their electrical urges, they would flee each other’s presence, and the nucleus would explode in all different directions. But something forces the protons to stay put and in close proximity to one another. A very strong force—stronger than gravity, stronger than electromagnetism—glues nuclei together, trapping protons inside. In a great burst of creativity, scientists dubbed this strong force . . . the strong force. This force holds the secret to nuclear fusion.

  The strong force is powerful enough to overcome the natural repulsion that protons have for other protons. However, it can do so only under a fairly narrow range of conditions. If there is the right balance of particles in the nucleus—the correct number of protons and neutrons—the strong force keeps the nucleus stable (or nearly so), preventing the nucleus from exploding. If there are too many neutrons or too few, the nucleus will be unstable. An unstable atom will destroy itself somehow, changing the balance of particles in its nucleus until the nucleus reaches a more stable state. A nucleus can break apart, spit a particle out, or swallow one to get closer to an ideal, stable balance of protons and neutrons.

  For example, hydrogen (one proton) and deuterium (one proton and one neutron) are stable. Left to their own devices, they would not change at all. But add a second neutron to the mix, making tritium, and the atom has too many neutrons for comfort. It is no longer stable. Eventually, a tritium atom will, spontaneously, transmute one of its neutrons into a proton (and spit out an electron in the process). The substance left behind is no longer tritium; it has become helium-3, a stable if rare isotope of helium that has two protons and one neutron. (Most helium is helium-4, which has two protons and two neutrons.) It takes an average of twelve years or so for any given tritium atom to undergo this decay process, but over time, if you have a jar full of hydrogen-3, you will find that it slowly transforms itself into helium-3.

  This transformation process releases energy, because the helium- 3 does not weigh exactly the same as the tritium did. The neutron that disintegrated weighed more than the proton and the electron that it turned into.17 There is mass missing: it was converted into energy, just as E = mc2 says. As the unstable tritium changed itself into the stable helium-3, it lost a little bit of mass and released a bunch of energy.

  This is an example of a general rule. When a nucleus converts itself from a less-stable variety to a more-stable one, it releases a little bit of energy because some of its mass disappears. And nuclei always “want” to become more stable, just as a ball perched on a hill “wants” to roll down to the bottom. In the process of getting more stable, an atom releases energy, just as a ball rolling down a hill picks up more and more speed as it goes.

  Marie Curie was seeing this process with radium. Radium-226 is a heavy atom with 88 protons and 138 neutrons. It is almost stable ... but not quite. On average, after 1,600 years, a radium-226 nucleus spits out an alpha particle (a helium-4 nucleus: two protons and two neutrons), leaving behind 86 protons and 136 neutrons—radon-222—and releasing a bunch of energy. This energy heats the hunk of radium, and it is why Curie observed that chilled radium would warm itself. It is also why a hunk of radium emits radon and helium. It so happens that radon, itself, is unstable; it decays into thorium, releasing energy, which, in turn, decays into another species and another and another, emitting energy at each step. Like a ball rolling down a bumpy hill, it keeps rolling and rolling until it reaches a stable place to rest: in this case, lead-206, which is much more stable than radium-226. The ball has rolled a long way down the hill. But it didn’t roll all the way down. There are, in fact, nuclei more stable than lead-206. At the very bottom of the valley are the most stable atoms of them all: the iron group.

  Iron-56 (26 protons, 30 neutrons), nickel-62 (28 protons, 34 neutrons), and a few other nearby iron and nickel isotopes are the ne plus ultra of the nucleus world. They are the most stable elements of them all. They are at the very bottom of the valley. All other atoms “want” to be iron, just as a ball anywhere on the slope of a hill “wants” to be at the very bottom.

  The landscape of nuclei is very much like a val
ley with a steep hill on one side and a shallow hill on the other (see the graph on page 47). At the very bottom of the valley is iron. Heavy nuclei, like radium and uranium, have more protons and neutrons than iron—they are high up on the shallow hill. To roll down to iron, they have to get lighter, shedding those extra protons and neutrons. Sometimes they do it in small steps, like radium does. Sometimes they do it violently, by breaking into two or more parts: fission. (Indeed, fission is little more than a way for heavy atoms to roll quickly down the shallow hill, losing mass and releasing energy in the process.) Light elements, on the other hand, are on the steep hill. They have to get bigger if they want to get into the valley of iron. This is what fusion is all about. Two light nuclei, if they slam together, can stick to one another to create a larger nucleus.

  CURVE OF BINDING ENERGY: Iron is at the bottom of the valley. Light elements on the left and heavy elements on the right release energy when they move down the valley of iron by fusing or fissioning.

  Fusion is a way for light atoms to roll down the steep hill toward iron. Since the fusion hill is much steeper than the fission one, a fusion reaction yields much more energy than an equivalent fission reaction. Fusion and fission are two sides of the same coin, but as Teller well knew, fusion is more powerful than fission.

  In the 1930s, fusion was soon to solve the puzzle that had so vexed Darwin and Kelvin, and it would answer a question that had bothered humans for millennia: Why does the sun shine? Hans Bethe, then a physicist at Cornell University, would uncover the answer.

  In retrospect, William Thomson was fundamentally correct. If the sun were, in fact, an incandescent ball of liquid, the energy released by infalling matter would only power it for a few tens of millions of years, far short of the time that Darwin’s theory needed to explain the diversity of life on Earth (and far short of the time that other scientists needed to explain geological processes). But Thomson’s work preceded E = mc2 by decades; nobody had yet puzzled over the nature of radioactivity or understood how fission and fusion turn matter into energy. Fusion held the solution to Thomson’s puzzle and vindicated Darwin. Fusion is the source of the energy that has powered the sun for billions of years.