Supernovas: Making Astronomical History

Nuclear Energy

Nook-yuh-lur. It's pronounced Nook-yuh-lur.
--Homer Simpson

Our Sun

TRACE image of the Sun (September 2000)

A typical supernova can glow as brightly as an entire galaxy of a hundred billion stars, pouring out as much energy as our Sun will in its entire lifetime in a few months. To understand where this fantastic amount of energy comes from, it helps to understand where stars derive their normal energy supply, during the many millions of years their lives can last. Before the year 1896, scientists did not even have a candidate for the Sun's energy source, for the simple reason that no phenomenon then known could provide that much energy for that long. Were the Sun composed entirely of a chemical fuel like coal, for example--and if it also had the oxygen necessary to burn the coal, of course--it would go dark in a few thousand years. Hermann von Helmholtz (1821-1894) proposed the idea that the Sun was contracting under its own gravitational pull, heating up as its own weight pulled its gases into a tighter and tighter ball. When he calculated how long this process would last, he discovered that the Sun could only have begun glowing at its present intensity some 70,000 years ago. This was worrisome, and not just because it implied that the Sun would die out comparatively soon, too: geologists studying the Earth's past had already decided that the planet was at least tens of millions of years old. Fossils of extinct animals like dinosaurs, already known to be millions of years old, clearly showed that the Earth was habitable at that time period, with the Sun performing then much as it does today. What could possibly power the Sun for those countless epochs, possibly upwards of a billion years?

Radioactivity

Marie Curie

Marie Curie (Nobel Chemistry Prize photograph)

In 1896, Antoine-Henri Becquerel (1852-1908) discovered that a mineral containing the element uranium released energy, at a steady rate, for prolonged periods of time. A crystal of uranium mineral could fog a photographic plate, even if the plate were wrapped in dark cloth to block all visible light. In the following years, other elements were discovered which had the same effect, some even more strongly than uranium itself. The Polish physicist and chemist Marie Slodowska Curie (1867-1934) coined the term radioactivity to describe these elements' peculiar property. Working with a uranium ore known as pitchblende (mined, it happens, in the mountains of Bohemia), Marie and her husband Pierre Curie discovered that the ore was more radioactive than the uranium within it could account for. They deduced the presence of additional radioactive elements, more intensely radioactive than uranium. In July 1898, they isolated one of these elements, which they named polonium; in December of the same year, they isolated a still more radioactive element named radium. This work won them the 1903 Nobel Prize (the third ever given) in Physics, which they shared with Becquerel. For her work on polonium and radium, Marie won a second Nobel, this time in Chemistry, in 1911. (Her husband, Pierre, had by that time sadly been killed in a traffic accident.)

The radioactivity of these minerals was the first indication that an energy source existed which was more powerful than chemical reactions. Radium, for example, glows in the dark, with practically undiminished intensity, for much longer than a human lifetime. Pierre Curie measured radium's rate of heat release to be 590 joules (140 calories) per hour for every gram of radium, enough to warm 1 cc of water 2.3 degrees Celsius every minute. What's more, radium maintains this heat rate not just for hours or days, but centuries: after 1600 years, the heat-emission rate is still half of its original value.

When this discovery became known, it rapidly found an application in industry. The hands of watches, and the numbers on their dials, were soon painted with radium compounds, which glowed a greenish color bright enough to be read at night. Sadly, the same radioactivity which produced this useful illumination--and which fogged photographic plates, as Becquerel had shown--also produced deleterious effects upon the human body. At the same time the Curies were working, the Dutch botanist Hugo Marie De Vries (1848-1935) described the phenomenon of mutations: he observed that the offspring of American evening primrose flowers (which had been introduced into the Netherlands in the 1880s) were sometimes noticeably different from the parent plants that had produced the seeds. De Vries observed traits in the flowers which appeared suddenly in one generation without having been visible in earlier ones, and he concluded that random factors must be producing changes in the flowers' genetic material. By 1927, the American biologist Hermann Muller (1890-1967) had shown conclusively that radiation, including X-rays and the emissions produced by radioactive minerals, sharply increased the rate of mutations. Most mutations, scientists discovered, were harmful to the organism they affected, frequently robbing the organism of the ability to conduct some necessary chemical function. Frequently, mutations are carcinogenic; that is, the genetic changes they involve lead to the uncontrolled cellular growth we call cancer. Muller, having discovered this harmful consequence of radiation, publicized his knowledge widely and fought for strong precautions on the handling of radiation--but his efforts were too late for the workers who had already developed serious health problems licking the brushes they'd used to paint radium onto watch dials.

Portrait of Pierre Curie

Pierre Curie (1903 Nobel Physics Prize photograph)

Now, even when Pierre Curie made his measurements in 1901, it was fairly clear that neither the Sun nor any other visible star obtained its energy from radium. (For one thing, the Sun doesn't glow green like a painted watch dial!) However, it was still evident that somewhere within the atoms of these elements there lurked a potent energy source, one which could pour out heat and light for centuries without noticeable diminishment. The study of radioactivity led to many astonishing breakthroughs in our understanding of the atom; in terms of pure knowledge, it is certainly one of the most significant discoveries of the last several centuries. (The fact that the last sixty years have seen us use this knowledge to develop weapons capable of annihilating our own species and much besides may make the discovery of radioactivity more significant still.)

Atomic Structure

Alpha Particle Drawing

Helium-4 nucleus, with two protons (arbitrarily colored red to denote positive charge) and two neutrons. (Art by Blake Stacey.)

The picture of the atom we have today, developed largely thanks to intensive study of radioactivity, is as follows. The center of each atom is a lump of material, known as the atomic nucleus, which contains almost all of the atom's mass and carries a positive electric charge. By the 1930s, it was known that the nucleus is composed of protons, which are positively charged and give the nucleus its overall electric properties, and neutrons, which are electrically neutral and therefore do not "feel" the effects of electric fields. (This made neutrons difficult to detect.) The outer regions of each atom--in fact the vast majority of the atom's volume-- are occupied by electrons, particles much lighter than either protons or neutrons, which carry a negative electric charge. In the normal state of affairs, the number of electrons in the outer regions equals the number of protons in the nucleus; since the size of the charge each particle carries is equal, the overall charge of the atom balances out to zero.

It is frequently said that electrons "orbit" the nucleus, in much the same way that planets orbit the central star of their star system. This is a useful image for studying certain aspects of atomic behavior, but it is by no means a complete picture. By the late 1910s, for example, it was deduced that electrons could only "orbit" the nucleus at certain fixed distances, whereas (in principle) a planet could orbit the Sun in a circle of any radius. By 1928, it was well established that objects as small as protons, neutrons and electrons--in fact, anything on the sub-atomic scale--behave in a fundamentally different way than the entities with which we are familiar, which are made of many trillions of atoms. (A liter of something as insubstantial as ordinary air, for example, contains roughly 2.68x1022 or 26,800,000,000,000,000,000,000 molecules, most of which are the two-atom nitrogen molecule N2.) Thanks principally to Erwin Schrödinger (1887-1961) and Werner Heisenberg (1901-1976), we have a mathematical way of describing this kind of subatomic behavior, a system called quantum mechanics. Countless experiments have shown that the world really is quantum-mechanical, when observed closely enough, but that doesn't make quantum behavior any more intuitive. (The physicist Richard Feynman (1918-1988) was fond of observing that, though the newspapers once claimed "only twelve men understand Einstein's Theory of Relativity", it wasn't really true: once people read Einstein's work, they probably understood it to some degree. Quantum mechanics, on the other hand, nobody understands, even though scientists are able to use it to make calculations.)

Electron wavefunction visualization

Rough visualization of an electron "wavefunction". (Art by Blake Stacey.)

Schrödinger's prescription for doing quantum mechanics, which lets us compute numerical answers for questions we pose, involves a quantity we call the "wave function". Essentially, if we want to know where in an atom we might find an electron, we imagine a sort of "cloud" which is denser in some regions than in others. This "cloud"--the wave function--tells us that we are most likely to observe the electron in the regions where the cloud is thickest. Before we actually make an observation, we can only know the probability that the electron is in one place or another. The shape of the wave function is influenced by electric fields or other forces; the familiar notion that like charges repel and opposite charges attract becomes, in quantum mechanics, a formula for the shape of the wave function, which we call Schrödinger's Equation.

Hydrogen in quantum mechanics

How quantum mechanics sees a hydrogen atom: one electron "inhabiting" the space around one proton. (Art by Blake Stacey.)

Schrödinger was able to calculate the different ways electrons can "inhabit" the space around an atom. With quantum-mechanical calculations as a background, it is possible to show that ordinary chemical reactions (combustion, rusting, or even the biochemical interactions inside living organisms) result from interactions among the electrons in the outer reaches of atoms. Atoms may "share" electrons, for example, forming a bond between them. Alternatively, one or more electrons may be removed entirely from one atom and donated to another. The donor atom, stripped of its electrons, has a net positive charge, while the recipient gains a net negative charge. Because opposite charges attract, these atoms can be bound together quite tightly. Large numbers of atoms, held together in such a fashion, can form structures difficult to disrupt. Ordinary table salt is an example: it is a crystal made of negative chloride ions (chlorine atoms with one extra electron) and positive sodium ions (which have lost one electron). The "electrostatic" force holding these ions together makes table salt ("sodium chloride") a crystal which is difficult to melt.

The bonds formed by these electron interactions are, however, relatively weak. The shared-electron bonds which hold together the hydrogen and oxygen in water (H2O), for example, can be broken by an electric current supplied by a small battery. Put another way, it doesn't take much energy to disrupt a shared-electron bond. (Bonds produced by the transfer of electrons from atom to atom, such as those in salt crystals, are stronger, but can still be broken by electricity or chemical reactions.)

Shared-Electron bond image

Qualitative image of a shared-electron bond. (Art by Blake Stacey.)

There is not much energy involved in rearrangeing the pattern of electrons around the nucleus, particularly when compared to the amount involved with altering the nuclear structure itself. To grasp the difference, note that ordinary chemical explosives like dynamite or TNT involve breaking shared-electron bonds to release energy. Nuclear weapons work by rearranging the protons and neutrons inside the nucleus (more details will follow, but this is the key notion). How much is the difference? Well, the first time we tried to blow something up using a nuclear "device", a few kilograms of material produced a blast equivalent to fifteen thousand tons of TNT. We have become more clever since then; technological advances have led to what Feynman (himself a group leader at Los Alamos) liked to call a "worse bomb".

Fission and Fusion

There are two primary ways nuclear reactions can produce energy. The first, which is familiar to us from Earthly radioactivity, involves large atomic nuclei splitting apart. The second, in which small nuclei are put together, is the mechanism which powers the stars. We call the joining of small nuclei into larger ones fusion, and we use the term fission to cover cases where large nuclei split into pieces.

Of the two processes, fusion produces most of the energy we see in the Universe, and is also responsible for the variety of chemical elements we find in the world around us. However, fission was the first nuclear process human beings learned how to control.

You may have noticed something odd about the picture of the atom described earlier: if like charges repel each other, and opposite charges attract, what keeps the protons in the nucleus--who all have the same charge--from flying apart? This is a real difficulty, and it puzzled scientists for many years. At first, they believed that the nucleus contained a few electrons, whose negative charge helped neutralize the protons' positive charge, making a kind of "cement" that held the nucleus together. Eventually, though, further experiments showed that the nucleus contained neutrons instead; because the neutron carries no electric charge, it cannot hold the nucleus together by electrical forces.

This still leaves the possibility that some other force exists, some force which operates between protons and neutrons and makes them attract each other. Because this force must be strong enough to overcome the protons' electrical repulsion, scientists termed it the strong nuclear force. (There is also a weak nuclear force, which behaves in a different way, but we needn't worry about it just yet.) We now have a fairly good understanding of the way the strong force operates, though why such a force exists in the first place is still an unanswered question.

The strong force can hold a small nucleus together very tightly. For example, the helium-4 nucleus, consisting of two protons and two neutrons, is a very stable unit. Larger nuclei, however, can be too big for the strong force to hold together. Sometimes, a large nucleus may "shed" small pieces. Because it is the number of protons in the atom which indicates what element the atom is, a nucleus emitting particles in this fashion changes the atom's chemical identity. Often, a heavy nucleus will emit a unit of two protons and two neutrons, all bound together, which is identical to the nucleus found in helium-4 (the most common variety of helium). This group is often called an alpha particle, a name invented before the particle's nature was known. (Early investigations of radioactivity led to the discovery of several kinds of radiation, which scientists named after the first letters in the Greek alphabet. This is kind of a running pattern: we find scientists talking about alpha and beta subunits in proteins, alpha and beta carbon atoms in molecules, and so on. As it turned out, beta rays are speeding electrons, released from the nucleus itself, and gamma rays are photons, or particles of light--but light of an extremely short wavelength, on the scale of the atomic nucleus itself.)

Uranium Nucleus

Nucleus of U-235: protons in red, neutrons in grey. (Art by Blake Stacey.)

Sometimes, instead of merely emitting smaller particles, a heavy nucleus will react in a more radical way. If we take an atom of uranium-238, for example, and "whack" it with a neutron, the neutron may impact the U-238 nucleus, pushing it into an "excited" state, which then splits apart, releasing extra neutrons and extra energy. The extra neutrons released in this reaction are not themselves energetic enough to cause other U-238 atoms to fission. However, natural uranium consists of a mixture of isotopes, including one with three fewer neutrons per nucleus known as U-235. If we strike a U-235 nucleus with a neutron, it may also fission, but unlike U-238, nuclei of U-235 can be broken apart by the neutrons fission releases. Therefore, if the U-235 atoms are packed together tightly enough, so that too many neutrons do not escape to the outside, we may witness a chain reaction, in which one fission event causes several more, which in turn cause many more, building up to a fantastic release of energy.

Uranium is not the only "fissile" element. Plutonium, discovered in nuclear-reactor products in 1940, also works quite well. (The elements uranium, neptunium and plutonium were named after the planets Uranus, Neptune and Pluto, which were in turn named after ancient Greco-Roman gods.) Many elements, even if they aren't actually suitable for chain reactions (and therefore nuclear fission bombs), break up into smaller pieces. Generally speaking, the more unstable an isotope, the less likely one will find it occuring naturally "in the wild". This is why, for example, technetium (element number 43) is so rare that it was produced in a lab before it was ever seen in nature: no isotope of technetium is stable enough to have survived the 4.5 billion years since the Earth formed. It can be extracted, however, from used nuclear fuel rods. (See this, fairly comprehensive article and the references therein.)

How did human beings first manufacture technetium? The answer will lead us to the second type of nuclear reaction, since it involves the putting-together of smaller nuclei. In 1937, Carlo Perrier and Emilio Segre (1905-1989) bombarded molybdenum atoms (element number 42) with deuterium nuclei--that is, with "heavy hydrogen" nuclei having one proton and one neutron apiece. When a "deuteron" approaches a molybdenum nucleus, the positive charge of the molybdenum nucleus makes the deuteron split apart. The proton gets pushed away, but the uncharged neutron feels no electric repulsion, so it impacts the Mo nucleus and causes a reaction. (Why not bombard with neutrons in the first place? The positive charge of the deuteron provides a convenient "handle" for accelerating the particles, which is much more difficult to arrange with uncharged neutrons.) The neutron impact causes a nuclear reaction, which produced two different technetium isotopes, Tc-95 and Tc-97.

By 1929, spectral analysis had shown that the Sun is composed mostly of hydrogen and helium, in a 3:1 ratio by mass. In that year, the physicist George Gamow (1904-1968) suggested that, in the intense pressures and temperatures within the Sun's core, four hydrogen nuclei could fuse together to make one helium nucleus. The excess energy of this reaction would be, ultimately, what makes the Sun shine. In later years, Gamow's ideas were refined, particularly by Carl Weizsacker (b. 1912) and Hans Bethe (1906-2005). It is this theory of nuclear fusion which answered Helmholtz's original question, providing a mechanism which could power the Sun's radiation for literally billions of years.

References and Further Reading

For information on the topics this page treats, we recommend the following books:

Researched, written and maintained by Blake Stacey.