Science Matters (12 page)

The energy needed to get the atoms into an excited state in the first place, and to get them to go back to it after they have emitted a photon, can be added to the system in many ways. Typically, scientists “pump” a laser by subjecting the material to heat, to a beam of energetic electrons, or to a bright light, from something like a flashbulb or even another laser.

Two precisely aligned mirrors at each end of the laser material cause the photons to move back and forth millions of times. Engineers design laser mirrors to allow a small fraction (maybe
5 percent) of the photons to escape on each bounce, and these leftover photons form the laser beam.

In the 1930s, Albert Einstein proposed a paradox that, to him, illustrated the fact that quantum mechanics, with its emphasis on probabilistic interpretations, could not possibly be the correct way to describe the subatomic world. His thought experiment was very simple: Suppose an atom emitted two particles back to back, and suppose further that we know that if one particle is spinning clockwise, the other must be spinning counterclockwise. If we write the wave function for either particle, it would be a combination of probabilities for these two directions of spin. The rules of quantum mechanics say that we have to describe the particles this way—each particle can have either spin until we make a measurement.

Einstein pointed out, however, that if we waited until the particles were far apart from each other—too far to allow a light signal to get from one to the other—and measured the spin of one particle, then we would know the spin of the other. In other words, measuring the spin of one particle would determine the spin of the other, even though no signal could pass between the two and no measurement was made on the second one. He regarded this as proof that quantum mechanics could not be the final theory of matter.

This proposal led to a strange series of events. In 1964, the Irish physicist John Bell proved a theorem that showed that there were certain experiments that would show unequivocally whether Einstein was right or not, and in the 1970s those experiments were done, showing that in spite of the apparent paradox, quantum
mechanics was right. It was about this time that the term “quantum weirdness” was born.

Today we understand that if two particles interact at some point in time (such as Einstein’s two emitted particles), then their wave functions are never really separate from each other. Scientists use the term “quantum entanglement” to describe this phenomenon. Because of quantum entanglement, we can’t measure only one of the particles in the pair—the measurement will always involve both. In the language of theoretical physics, we now understand that quantum mechanics, unlike ordinary Newtonian mechanics, is not a “local” theory, and that some quantum particles can never be separated from each other, no matter how far apart they are.

As strange as it may seem, quantum entanglement has not only been verified in the laboratory, but is being developed for a number of practical applications, of which quantum teleportation is one. It works this way: Imagine that you have a pair of entangled particles—call them photons, just for the sake of definiteness. Suppose further that you give each member of the pair to a different person. Traditionally, these people are called Alice and Bob, a joking reference to a 1970s movie titled
Bob and Carol and Ted and Alice
.

Now suppose that Alice takes another photon—call it the test photon—and lets it interact with her member of the entangled pair. She can then tell Bob what the result of that measurement was (a phone call will do), and Bob can then re-create the test photon in his laboratory using his member of the entangled pair.
Thus, the test photon was destroyed (or at least changed) in Alice’s lab and an identical photon was created in Bob’s. We say that the photon was teleported to emphasize that no photon ever traveled directly from Alice to Bob.

In 1997, physicist Anton Zeilinger of the University of Vienna used this technique to teleport the first visual image. (It was a photograph of the Venus of Willendorf, a stone age fertility statue found in Austria.)

Entanglement may also turn out to be an important phenomenon in the area of cryptography. If you think about the situation of Bob and Alice described above, you realize that it is absolutely secure. It would do an eavesdropper no good to intercept the call from Alice to Bob, because without one of the entangled photons he could not reconstruct Alice’s test photon. Furthermore, if he tried to intercept an entangled photon, the uncertainty principle guarantees that he would have to change it in the process, which would be readily detectable. Thus, quantum entanglement holds the prospect of serving as the basis for absolutely secure communications.

One way of thinking about a particle described by a wave function is to imagine that the particle is actually in all possible states simultaneously. (In technical jargon, we say that the wave function is a “superposition of all possible states.”) The key idea behind the new field of quantum computing is that if we let each of those superposed states perform a part of a calculation and add everything up at the end, we will have a system capable of doing calculations much more rapidly than they could be done in a conventional machine.

Think of this analogy: One way of reading a page of this
book would be to give one sentence each to a series of readers, then assemble the output from those readers at the end of the process. This would clearly take less time than having one reader go through the page sentence by sentence. In the same way, the hope of people trying to develop a quantum computer is that the states involved in the superposition will play the same role as the readers in our analogy, producing computers of unparalleled power.

We do not have a working quantum computer at this point, but scientists have succeeded in building parts of such a machine (what are called logic gates) and have hopes of proceeding farther in the future.

N
OTHING
Beats Margees Pennsylvania applesauce bread.

Combine all ingredients in large mixing bowl. Beat at medium speed until well blended. Grease a 9“x 5” loaf pan on the bottom only, and pour in mixture. Bake at 350°F for 55 to 60 minutes,
until a toothpick comes out clean. Loosen the edges with a spatula and remove from the pan. Cool before slicing.

Scientists like to cook up new recipes. Early in 1986 two scientists made history (and won the Nobel Prize) by mixing and grinding ordinary elements like copper, oxygen, and a couple of others in just the right proportions, and baking them at just the right temperature just long enough to produce a nondescript black wafer about the size of an aspirin tablet. That little black disk turned out to be the first of an entirely new kind of superconductor, a material with extraordinarily valuable electrical properties.

How is it possible that ingredients as different as flour, eggs, salt, and applesauce can combine to form a delicious loaf of bread? How can commonplace elements as different as copper and oxygen become a priceless black disk? The answer lies in the atoms—the building blocks of everything around us. Atoms combine in countless ways to form materials with every imaginable property. But in every case:

The properties of all materials, be they sticky, magnetic, brittle, or green, arise from the infinite different ways atoms can be arranged and linked together. Two atoms form a chemical bond when their electrons rearrange themselves so that each atom feels an attractive force. The attraction of positive and negative electrical charges holds everything together.

The distinctive materials that enrich our world can be described in architectural terms. Various atomic building blocks—the elements
—are joined in structures using chemical bonds. Elements can be combined with each other in countless ways to form chemical compounds that have properties and applications very different from their elemental raw materials. Nature and chemists have created millions of different chemicals. Some are very simple, like H
₂
O (water) with two hydrogen atoms for every oxygen, or NaCl (table salt) with a one-to-one ratio of sodium and chlorine. Other compounds are exceedingly complex, combining a dozen or more different elements.

With fewer than one hundred elements to play with, it might seem that the chemist’s job of making new chemical combinations would be quickly exhausted. Not so. A million chemists could each make a new sample every day for a million years and still not come close to running out of things to try. There are about 70,000,000 possible four-element combinations in one-to-one-to-one-to-one ratios, to say nothing of the countless variations in element ratios. For each possible composition there are hundreds of ways to mix and treat the chemicals, each resulting in a different arrangement of the same set of constituent atoms, and therefore in a material with different properties. Just as in cooking, the temperature and time of baking can make all the difference between a culinary masterpiece and an unmitigated disaster. The successful chemist, like a great chef, must apply skill and intuition in preparing new recipes.

Chemists make their living by trying to create new and useful mixtures of atoms, or by trying to manufacture established useful chemicals in new ways. This is not an idle pursuit. Almost every aspect of modern life—food and clothing, transportation and communications, sports and entertainment—depends on the discoveries of chemistry.

Chemistry is first and foremost the science of electrons and their interactions. That may seem an odd assertion, for chemistry is usually portrayed as the science of test tubes and bubbling beakers and odd mixtures of stuff that turns blue. But all those chemical reactions are the results of electrons shifting between atoms.

When two atoms approach each other, their outermost electrons come in close contact; the two negative electric charges repel each other. In most instances, as in the collision of gases in the atmosphere, the atoms just careen off to another chance rendezvous. Occasionally, however, colliding atoms stick together by an exchange or sharing of electrons.

The first step in understanding the extraordinary variety of materials that surround us is to look at the way that two or more atoms can link together: the chemical bond. There are several different kinds of bonds, each producing dramatically different properties. The general rule that dictates what kind of linkage, if any, will be formed when two atoms come near each other is simple: the two-atom system will try to reach a state of lowest possible energy Just as a ball rolls down a hill to get to a position of lowest gravitational potential energy, an atom’s electrons rearrange themselves to reach a state of lowest electrical potential and kinetic energy.

Three types of chemical bonds—ionic, covalent, and metallic—hold almost all materials together. Each type involves a different strategy to rearrange electrons among atoms, but the secret to forming most chemical bonds is the extreme stability of atoms with filled electron shells. Atoms with exactly 2 or 10 or 18 or 36 electrons (corresponding to 1, 2, 3, or 4 filled shells) are
much happier (if you will excuse the anthropomorphism) than those with a few electrons more or less—11 or 17 electrons, for example. We think of 2, 10, 18, and 36 as the magic numbers of chemical bonding. In fact, the only elements that don’t readily engage in any kind of chemical bonding are found in the extreme right-hand column of the periodic table—elements 2, 10, 18, and 36, with filled outer electron shells. These are the so-called inert gases, including lighter-than-air helium gas and the neon gas of colorful lights.

An ionic bond forms when one atom of a pair gives up an electron while the other acquires it on permanent loan, so that both atoms achieve a magic number. In common table salt, for example, a sodium atom with 11 electrons transfers its outermost electron to a chlorine atom, which has 17 electrons. When this shuffle occurs, the sodium atom winds up with only 10 electrons, while the chlorine atom increases to 18 electrons—both magic numbers. What’s more, by shifting an electron, sodium becomes a positive ion and chlorine becomes a negative ion, so the atoms form a strong electrostatic attraction—an ionic bond.

A glance at the periodic table (see page 75) reveals some elements engage in ionic bonding. All the elements in the two farthest left-hand columns of the table have only one or two electrons in their outer shell. These elements would be much happier to get rid of their outermost electrons and become positively charged ions. By contrast, elements in columns near the right-hand side of the table are one or more electrons shy of a full outer shell; they are eager to add a few more electrons and become negatively charged ions. Positive attracts negative and so, voilà, an ionic bond forms.

Hundreds of everyday materials are held together by ionic bonds. Table salt is the classic textbook example of ionic bonding, but compounds that combine silicon and oxygen, two of the Earth’s most abundant elements, are by far the most common ionic compounds. Most beach sand is a mineral called quartz, formed when silicon atoms, which yield four electrons, bond to pairs of oxygen atoms, each of which collects two electrons, leaving both silicon and oxygen with closed shells. The large electrostatic charges of +4 on every silicon and—2 on every oxygen results in a strong electrostatic force between adjacent atoms. That is why quartz, which can scratch steel, is such a tough and resistant mineral. Similar ionic silicon-oxygen bonds provide the strength and hardness of window glass, china and other ceramics, and most common rocks and minerals.