What do we mean by “the energy needed to probe it”? We have to get a little bit technical here and introduce you to the common currency in talking about energies of fundamental particles and atoms: the electron volt or “eV.”
4 Most of the science of chemistry, that is, the amount of energy involved in most chemical reactions, lies in the range from about 0.1 to 10 eV per atom. This means that when a given atom enters into a chemical bond with another atom (or an existing molecule) to form a new molecule, roughly 0.1 to 10 eV of energy is released. This is energy that comes from the forces that produce a chemical bond between two atoms, and it is typically released in the form of light, or the energy of motion, called kinetic energy.
The released energy is usually converted into heat (which is the aggregate random motion of atoms in a material), but it can also be seen as the light emitted from a fire or heard as the boom from a firecracker. You can usually see the released chemical energy with your eyes because a single visible particle of light, the photon, carries about 2 to 3 eV of energy—after all, the light entering our visual system that allows us to see is processed by various chemical reactions in our eyeballs and our brain, and so the perception of light entirely happens at the chemical energy scale.
If we can probe molecules with a source of energy of about 0.1 to 10 eV, we can often cause a chemical reaction to occur. For example, striking a match in a mixture of methane (CH4) and oxygen (O2) will provoke a rapid chemical reaction—a flame—yielding carbon dioxide (CO2) and water (H2O).5 The match is generating photons and kinetic energy of motion of atoms of about an eV each from its own burning (usually oxygen combining with phosphorous). These energetic particles strike the methane and oxygen and nudge them into reacting, which emits more photons. Then more and more energy is released in a chemical chain reaction, and VAVOOM, you might have an explosion.
The physics, that is, the motion and interactions of electrons and atoms in chemical reactions—the stratum of the chemical reactions—is very much independent of, or decoupled from, what is going on in other stratums of nature. For example, to analyze everyday chemical phenomena, one need not be bothered by such complications as the detailed motion of the protons and neutrons that comprise the inner atomic nucleus and that exist on much smaller distance scales than the overall size of the atom. Nuclear physics is a far different energy stratum, measured in millions of electron volts, or “MeV” (see note 4), compared to the lowly 0.1 to 10 eV stratum of chemistry. Nor need we, in studying chemistry, be bothered by the slow, lugubrious astronomical motion of the earth in its orbit around the sun. In fact, it is the relative motion of atoms and the electrons within the atoms that matters for chemistry. Thermal effects, the random motion and collisions of atoms due to heat, are typically about 0.1 eV per atom at room temperature, and they increase with temperature and therefore do have effects on chemical reactions (i.e., “cooking”). But the motion of the earth in its orbit is a common, uniform drift of the assemblage of all of the earth's atoms, producing no high-energy inter-atomic collisions. Of course, if an asteroid collides with the earth, the relative motion of the asteroid's atoms hitting the earth's atoms involves very high energies, and very serious chemistry, even nuclear physics, will occur!
While the triumph of Mendeleev's Periodic Table of the Elements formed a basis for understanding all chemical reactions, we learned in the early twentieth century that the atoms themselves are not truly elementary: they are composed of even smaller, more basic objects. To understand this we must probe into the atom. And, as it goes with probes, the probe we use to analyze something should preferably be smaller than the object we wish to probe. If the probe is bigger than the object probed, it becomes a bludgeon or battle-ax, and battle-axes don't work so well for dissecting tiny things or performing dental surgery, etc.
A simple home-brew experiment that you can perform, e.g., at a child's birthday party, will illustrate this point.
A SIMPLE HOME-BREW EXPERIMENT
Get a beach ball and a straw. Have someone blindfold you. Have your assistant take some randomly chosen small items unknown to you, like a peanut, an acorn, a coin, nuts and bolts, a few other small things, and place them on a table in front of you. Now, while still blindfolded, take hold of the beach ball with both hands. Holding only the ball, try to use it gently as a probe of the small objects on the table, the peanut, the acorn, etc. Can you discern which little object is which, while blindfolded, and coming into contact with them only through the very large beach ball? We would guess not, unless you peeked.
Next, take one end of the long straw and, while still blindfolded, use it to trace out the forms of the same small objects. Can you now discern what these objects are and which is which? When you trace out the objects’ shapes you must use a little thought and a little imagination to try to figure out what each of them is—you've become both an experimentalist and a theoretical physicist at the same time—like Enrico Fermi. With enough effort and thought you can probably figure out what little objects were placed before you on the table. Perhaps you can tell a dime from a nickel, and chunk of cauliflower from a golf ball. Go ahead—try it!
One thing is clear, if not obvious, from this experiment: a probe that is many times larger than the object to be probed does not work very well. Holding the beach ball, we doubt you can discern a nickel from a dime, if you can even detect either of them. On the other hand, probes that are much smaller than the object itself allow us to readily resolve the object's structure—even without seeing it with our eyes. This simple principle holds true in all of the strata of the onion of nature, including the stratum of subatomic particles. To explore the structure of the unknown “something,” we must construct a probe that is smaller than the “something” we seek to study.
This seems at first blush to pose what appears to be an insurmountable barrier to studying small objects, like the innards of an atom or a particle inside an atom. How can we study a particle's inner structure if all we have are other particles of the same size? Ah-ha! Here is where two of the greatest revolutions of science come to our aid: the quantum theory and Einstein's theory of relativity.
Essentially, we learn from quantum theory that all particles in nature are also waves. This seems to be a ridiculous and nonsensical paradox, but it is the mysterious and jarring reality of quantum theory. The effects of waves vs. particles for most things don't show up until we reach atomic dimensions, but they can be seen readily for ordinary light. But to be precise, a quantum state is neither a particle nor a wave—it is both at once!
Small objects can be described by quantum mechanical waves that are associated with the probability of detecting a point-like particle at any point in space and time. That is a mouthful, and the interested reader should grab a copy of our book Quantum Physics for Poets (Amherst, NY: Prometheus Books, 2011). However, if you can just “ride the wave” with us for a few more paragraphs, you need only accept that a wave always has a characteristic wavelength. The wavelength is just the familiar distance between two crests or two troughs of a wave, like a water wave. It is the quantum wavelength that tells us how big an object is when it used as a probe.
Now here is a second relevant fact about quantum physics: as we increase the energy of any particle, its quantum wavelength becomes smaller and smaller. When the wave motion approaches the speed of light, then Einstein's theory of relativity kicks in. If you double the energy of a particle moving near the speed of light, you will halve its quantum wavelength. So, investing a lot of energy in a particle makes its quantum wavelength smaller. This, in principle, allows us to make an arbitrarily tiny probe simply by accelerating a particle to arbitrarily high energies. This is the most important principle underlying microscopes and particle accelerators. The more energy in a particle, the smaller it becomes. And, by the way, you now understand why today's particle accelerators are very large: it takes a very large accelerator to put a lot of energy into a particle to make it become a smaller probe.
The
wavelength of ordinary visible light ranges from, approximately, higher-energy blue light, 0.00004 centimeters (4 × 10-5 cm, about 3 eV per photon; recall that a centimeter is about a half an inch) to lower-energy red light, 0.00007 centimeters (7 × 10-5 cm; about 2 eV per photon). A typical visible particle of a light, a photon, has a quantum wavelength in this range, with an energy of approximately 2 to 3 eV. Objects larger than about 0.0001 centimeters (10-4 cm) can be readily probed with visible light because they are smaller than the wavelength of the light wave. You need only make a precise optical microscope to do this, and you can see little things that your eye cannot resolve.
Figure 2.1. Wave and Wavelength. (A) A traveling wave. The wave moves to the right at a speed of v and has a wavelength (the length of one full cycle, from trough to trough, or crest to crest). An observer watching the wave travel by would see a frequency of v/(wavelength) crests, or troughs, passing by per second. The amplitude is the height of a crest above zero. (B) As we increase the energy of a quantum particle-wave, the wavelength becomes smaller.
However, visible light falters when it is used to study structures smaller than this size scale, such as the tiny components found inside the living cell of a biological organism. Visible light is unable to resolve two objects of much less than 0.00001 centimeters (10-5 cm) or smaller. You now know the reason: these objects are smaller than the wavelength of the visible photons of light—visible photons are as useless as beach balls to probe such small distance scales. No improvement in your microscope optics can ever improve the image. You could spend hundreds of thousands of dollars on a top-of-the-line Bausch and Lomb microscope, and still the tiniest denizens deep inside living organisms will only appear fuzzy or will not appear at all. Crank up the magnifying power of your microscope, and all you'll get is a bigger fuzzier image. You absolutely cannot see DNA in an optical (light-based) microscope. Visible light is hopeless to use as a probe of an atom at the atomic-size scale of 0.00000001 centimeter (10-8 cm) or less.
Fortunately, at shorter distance scales, even down to the atomic stratum and beyond, electrons become excellent probes. Electrons can be accelerated in a small type of particle accelerator called an electron microscope, giving them more energy. Electrons, too, have a quantum wavelength, as do all particles. Electrons can easily be endowed with kinetic energies of about 20,000 eV (that's 10,000 times more than a visible photon). This is the energy of acceleration of the electrons in the old TV picture tubes that could at one time have been found in any household but that have now gone the way of the horse and buggy. At this energy the electrons have a quantum wavelength of about 0.000000001 centimeters (10-9 cm), considerably smaller than that of visible light, and they can be used to make images of DNA, a virus, and even resolve individual atoms.
PEERING INSIDE THE ATOM
The first peek into the internal structure of the atom, the “atomic onion layer of nature,” began about 50 years after Mendeleev with a British scientist named J. J. Thomson.6 Thomson cleverly demonstrated in 1897 that certain “rays” that could be provoked out of atoms in an electric discharge tube (something like a fluorescent light tube) were actually particles. In particular, Thompson established that these particles lurked within all atoms, and he called the new particle the electron. Thomson is deservedly heralded as the “father of modern particle physics” for this discovery. He proved that electrons were extremely low-mass particles, weighing two thousand times less than the atom itself, and that they each carried a negative electric charge.
Thanks to the work of J. J. Thomson, atoms were now known to be full of these very lightweight, negatively charged electrons. But atoms themselves are normally electrically neutral, i.e., they have no electric charge; they can be “ionized” and lose an electron, hence they acquire the opposite, positive charge. Obviously, then, there would have to be an equal amount of positive electric charge inside the atom, balancing the negative electric charges of the electrons. But where this positive charge resided within the atom remained a mystery. In 1905, Thomson had proposed a theoretical model of the atom in which the positive electric charge is a medium that is uniformly dispersed throughout the entire atom, with the electrons embedded within it like raisins in a loaf of raisin bread.
Throughout this era a gruff walrus of a young man, Ernest Rutherford, played a prominent role in unraveling the inner properties of matter. Rutherford was a skilled and masterful craftsman and won the Nobel Prize for his work in elaborating the properties of radioactivity (which we'll describe later), and he had now become the director of the famous Cavendish Laboratory in Cambridge, England.7
Rutherford had grown up as one of a dozen children in a farm family in rough-and-tumble New Zealand, where he had learned hard work, thrift, and tinkering with technological innovation. As a child he'd played with clocks and made models of his father's waterwheels, and by the time he was a graduate student he was investigating the physics of electromagnetism. He had managed to devise a detector of wireless (radio) signals even before Marconi began his famous experiments that led to the wireless telegraph. When a scholarship brought Rutherford to the Cavendish Laboratory, he hauled his wireless device along to England and was soon sending and receiving signals over half a mile, a feat that impressed the Cambridge dons, including J. J. Thomson, the Cavendish director at the time. Later Thomson would declare, “I have never had a student with more enthusiasm or ability for original research than Mr. Rutherford.”8
By 1909, Rutherford and his students were shooting tiny subatomic probes, called “alpha particles,” at a piece of thin gold foil. They were carefully measuring the way in which the particles were slightly deflected as they scattered off of the heavy gold atoms in the target foil. The alpha particles were produced from the radioactive decay of the element radium, which naturally accelerated them to high energies—there were no man-made particle accelerators in those days. Alpha particles were now known, thanks to Rutherford, to be very heavy compared to an electron, and with a little energy they have a very tiny quantum wavelength and are therefore capable of probing deep inside the atom.
One day something utterly unexpected happened. The alpha particles were usually deflected only slightly by their passage through the gold foil, and the scientists were measuring in detail this “forward scattering.” They decided, simply as a sanity check, to see if there was any “backward scattering.” To their astonishment they found that one in 8,000 alpha particles bounced back toward the source. As Rutherford remembered it, “It was as if you fired a fifteen inch artillery shell at a piece of tissue paper and the shell came back and hit you.”9 What was happening here? What kind of thing inside the atom was repelling the positively charged and massive alpha particle?
No one before Rutherford had devised any way of mapping the inner shape of the atom. According to the “raisin bread” model of the atom of J. J. Thomson, the alpha particles should always have bullied their way straight through the atom—always! The atom was like a big glob of shaving cream, and the alpha particles were rifle bullets. Rifle bullets would tear straight through globs of shaving cream. Imagine seeing a rifle bullet occasionally deflected and ricocheting backward upon colliding with a blob of shaving cream. Such was the observation of Rutherford and his students.
Rutherford devoted his full energies to understanding this remarkable discovery. According to his detailed calculations there was only one way that any alpha particles could ever be deflected backward. That could only happen if the entire mass, and positive charge, of the atom was concentrated into a tiny volume in the center of the atom—the “atomic nucleus” was discovered! The nucleus's hefty mass and large positive charge could repel the positively charged alpha particles that came within its range and deflect them through a large angle, even kick them backward. It was as if within the glob of shaving cream there were dense, hard ball bearings that could cause bullets to collide and deflect. The electrons were orbiting this dense central charge of the atom. The raisin bread theory of the atom of J. J. Tho
mson was tossed in the trash-bin. An atom now resembled a tiny solar system, with miniature “planets” (electrons) orbiting a dense “dark star” at the center (nucleus), and it was all held together by electromagnetism.
Further experiments indicated that the nucleus was indeed tiny—one-trillionth of the volume of the atom—even though most of the mass of the atom, more than 99.98 percent of it, resided in the nucleus. At the time of this discovery, within this tiny solar system model of an atom, all the classical laws of physics were still thought to be rock solid, just as in the macroscopic solar system with the sun and its planets. The same laws of classical physics were believed to work in the atom just as they did everywhere else—until Niels Bohr showed up.
THINKING THROUGH THE ATOM
Niels Bohr10 was a young theoretical physicist from Denmark who was studying at the Cavendish Lab, and he happened to attend a lecture by Rutherford. He was immediately captivated by this new atomic theory of electrons orbiting nuclei. Bohr arranged to visit the great man for four months in 1912, while Rutherford, at the time, was working in Manchester. Sitting down and thinking about the new data, Bohr quickly perceived something profoundly significant about Rutherford's planetary model of the atom. It was a complete disaster, according to the known laws of physics!
Beyond the God Particle Page 5