The Nucleus
The nucleus is an extraordinary thing. It contains 99.9 percent of an atom’s mass, but occupies only a trillionth of its volume. Atomic nuclei are tightly packed clusters composed primarily of protons and neutrons. These two atomic building blocks have nearly the same weight, and each weighs about 1,860 times more than an electron. Don’t be taken in by the fact that protons and neutrons are “massive” on the atomic scale—it still takes about 600,000,000,000,000,000,000,000 of them to balance a thimbleful of water.
Protons determine how an atom will behave. Each proton has a positive electrical charge of +1, so the number of protons in the nucleus dictates the electrical characteristics of an atom. Each chemical element is defined exclusively by its number of protons—the so-called atomic number. Every gold atom has exactly 79 protons. Helium, carbon, oxygen, and iron are element names for atoms with exactly 2, 6, 8, and 26 protons respectively The number of other particles is irrelevant for the purposes of assigning names.
All the naturally occurring elements, from number 1 (hydrogen) to 94 (plutonium), are found in the rock, water, or air of Earth. Of these ninety-odd elements, about fifty form almost everything that you are likely to see or use in a lifetime. Elements beyond number 94 can be created in specially equipped physics laboratories, although these “heavy atoms” are highly unstable and do not survive long. Elements heavier than 94 have names that honor prominent people and places of twentieth-century physics as in berkelium (97), einsteinium (99), and fermium (100).
Neutrons weigh roughly the same as protons, but lacking an electric charge, they have little effect on the structure of the atom or on the way one atom interacts with another. They play an important role, however, in holding the nucleus together, and they are as important as the proton in giving the atom mass. In fact, scientists weigh atoms and subtract the known weight of all the protons to determine how many neutrons are present.
The number of neutrons in a nucleus can vary from zero, in most hydrogen atoms, to more than 140 in the heaviest atoms. In most familiar atoms, protons and neutrons are present in roughly equal numbers. The commonest kind of carbon atom, for example, has 6 protons and 6 neutrons in its nucleus, while oxygen usually has 8 of each. In heavier elements, like iron (26 protons usually coupled with 30 neutrons) or platinum (78 protons and 117 neutrons), the neutrons outnumber protons by a modest margin.
The nuclei of a given chemical element must all have the same number of protons, but may have different numbers of neutrons. The nucleus of carbon, for example, usually has 6 of each, but occasionally there is a nucleus with 6 protons and 7 or 8 neutrons. An atom with this nucleus is still carbon, since it has six protons, but it weighs more than ordinary carbon. Atoms with nuclei that contain different numbers of neutrons are called isotopes of the element. Scientists customarily denote an isotope by giving the combined total of protons and neutrons. Thus ordinary carbon is called carbon-12, and the other kinds carbon-13 and carbon-14.
Electrons
We have already met electrons, the mobile carriers of negative electric charge. Electrons are tiny things, weighing only about 1/1860th as much as a proton or neutron. If you imagine an atomic nucleus as a basketball-sized 5-pound weight, then electrons are something like bumblebees flying around several miles away while everything in between is a void. Atoms, which form all “solid” matter, are themselves almost entirely empty space.
In order to keep the atom as a whole electrically neutral, the number of electrons in orbit must equal the number of protons in the nucleus exactly. Occasionally, because of collisions or some other mechanism, a particular atom may lose or gain electrons. An atom in which the number of electrons is not the same as the number of protons (and which is not, therefore, electrically neutral) is called an ion. The atoms in the gas in an ordinary fluorescent lightbulb, for example, are often ionized when electrons are torn off normal atoms in collisions.
What You Need to Know
We’ve tried to stay away from jargon in this book, but there really are a few terms that a scientifically literate American should know. Here’s a brief summary of key words relating to atoms—words that appear again and again in news stories about nuclear waste, fusion power, superconductivity, and advances in electronics.
ATOM: The basic building block of everything around you.
NUCLEUS: The heavy central part of every atom; the nucleus contains protons and neutrons.
PROTON: A positively charged particle in the nucleus; the number of protons distinguishes one element from another.
NEUTRON: An electrically neutral particle in the nucleus.
ELECTRON: An electrically negative particle that orbits the nucleus.
ELEMENT: A substance that can’t be broken down by chemical means; an atom for which you know the exact number of protons—the element carbon always has 6 protons.
ISOTOPE: An atom for which you know both the exact number of neutrons and the number of protons; the isotope carbon-14 always has 6 protons plus 8 neutrons. Different isotopes of a given element have the same chemical properties.
ION: An atom that has gained or lost electrons, and hence has an electrical charge.
THE BOHR ATOM
Although negative electrons circle the positive nucleus much as planets circle the sun, electron orbits differ from those of planets in one significant way. A planet like Earth does not have to be any specific distance from the sun—if its orbit were 10 feet closer to the sun or 10,000 miles farther out, no laws of physics would be violated. Electrons in an atom, however, can “orbit” only in certain well-defined paths, and can never be found anywhere else. Each of these so-called allowed energy levels corresponds to a different energy, so the energy of atomic electrons can only have certain exact, specified values.
Electrons move from one energy state to another by means of a “quantum leap,” a phenomenon that is impossible to picture. The electron simply disappears from one energy level and reappears in another without traversing the space in between. It’s as if you moved up and down a staircase by simply vanishing from one step and popping up on the next. You may think that sounds like nonsense, but when we look at things in the universe that are extremely small, extremely massive, or traveling extremely fast, they just don’t behave in familiar ways. In fact, many of the phenomena we encounter inside the atom have no analogs in our everyday experience.
Since each electron state has a different energy, an electron moving to a new state must either give up or take in energy when it makes the leap. If the new state is closer to the nucleus than the old, the electron emits energy in the form of electromagnetic radiation. This is how atoms emit visible light. The light energy is precisely equal to the difference in energy between the old and new energy states. Similarly, if an atom takes in energy (by absorbing light, for example, or by collision with another atom), that energy can be given to the electron, which will then move to a higher orbit. This picture, in which electrons shuttle back and forth between allowed energy states as they absorb and emit energy, is called the “Bohr atom,” after the Danish physicist Niels Bohr (1885–1962), who first suggested it as a young postdoctoral fellow in 1912.
THE PERIODIC TABLE
OF CHEMICAL ELEMENTS
When chemists first investigated how materials were put together, they made an important discovery. It turned out that you could take anything—a piece of wood, for example, or a stone—and break it down into other things. Wood, when burned, yields carbon dioxide, water, and some mineral ash. Each of these substances in turn can be broken down by the appropriate procedures into still other substances—carbon, hydrogen, and water, for example. But try as they would, chemists could not find a way to break things like carbon and oxygen down further. These substances, which seemed to make up the pieces of more complex materials but which were themselves irreducible, scientists termed “chemical elements.” At the end of the eighteenth century chemists knew of about twenty-six elements, and today that list has expanded to ov
er a hundred.
Every element has a name. “Hydrogen,” for example, comes from the French for “generator of water,” a name that tells the story of how the element was discovered. We customarily represent each element by a simple one-or two-letter abbreviation—H for hydrogen, O for oxygen, Ca for calcium, and so on.
Chemists systematize all the elements in one simple table—the periodic table. Every element has its own box, with atomic number increasing as you read from left to right. Boxes are arranged so that elements in a vertical grouping have similar chemical behavior—that is, elements in the same column enter into similar reactions and combine to form similar compounds.
The periodic table of the elements, a fixture in chemistry lecture halls, was first written down by the Russian scientist Dmitri Mendeleyev (1834–1907) in 1869. He didn’t understand why a ranking of the elements in increasing order of weight had something to do with their chemical properties, but it seemed that it did. In addition, when Mendeleyev first wrote down the table there were two holes, corresponding to the elements we now call germanium and scandium. When these elements were discovered (in Germany and Sweden, as the names imply), it was considered an important piece of evidence that the organization of Mendeleyev’s table has a deep reason behind it.
Today we understand that elements can be grouped into rows and columns because of the way electrons fit into shells of electrons, corresponding to different allowed Bohr energy levels. It turns out that electrons cannot be crowded too closely together. Like two cars in a parking lot, two electrons cannot share the same space. We call this the exclusion principle (because the presence of one electron excludes others).
Thus, there is space for only two electrons in the lowest Bohr electron shell. The atoms of hydrogen (one electron) and helium (two electrons) respectively fill that shell, so the inner electron shell becomes “closed,” to use physicists’ jargon. The next element, lithium, has three electrons, however, so the third electron must go into the next higher shell. Thus lithium, like hydrogen, has a single electron in its outermost shell, which explains why the two have similar chemical properties.
The second and third shells have space for eight electrons each, so we can squeeze in seven electrons after lithium before we start on another shell, and thus find another atom with one electron in its outermost shell. This element is sodium, with eleven electrons. Not surprisingly, you will discover hydrogen, lithium, and sodium in the same column of the periodic table, along with potassium, rubidium, cesium, and francium—all of which have one electron in their outer Bohr shells.
Using quantum mechanics, the subject of the next chapter, we can predict the number of spaces available for electrons in each atomic shell. It is this calculation that ultimately justifies and explains the periodic table.
FRONTIERS
One of the most striking areas of modern research is the rapid development of “microscopes” capable of producing photographs of individual atoms in a material. The best developed of these instruments, the scanning tunneling microscope (or STM), works by measuring the electric current that flows between a tiny, precisely positioned point and the atoms on the surface of a material. The closer the point of the atom, the greater the current. A typical STM photograph shows the presence of individual atoms on a surface.
A scanning tunneling microscope reveals the locations of individual iodine atoms coated on a surface of platinum. The distance between atoms is only about a billionth of an inch. PHOTO COURTESY OF BRUCE SCHARDT, PURDUE UNIVERSITY.
As scientists improve the resolution of these microscopes, and as they learn more and more about the surfaces of materials all around us—metals, plastics, paper, and skin, to name just a few—we can expect to see more spectacular pictures of atoms.
SUPERHEAVY ELEMENTS
Although natural elements occur only up to atomic number 94 (the radioactive element plutonium), artificial “superheavy” elements up to 114 (with the inelegant but temporary name “ununquadium”), and as high as 118, have been created in the laboratory by colliding less massive atoms together and hoping that some of them stick together. For example, researchers at Russia’s Joint Institute for Nuclear Research at Dubna fired a beam of calcium-48 atoms into a target of plutonium-244 to produce a few ephemeral atoms of element 114. These atoms, with their massive collections of protons and neutrons, have unstable nuclei and split apart in a few seconds into groups of smaller atoms.
Some theorists speculate that if and when we get to atomic numbers in the vicinity of 120 to 130 we may find a new group of stable elements. If they are correct, then the periodic table of the late twenty-first century may be a lot longer than the one we now have.
CHAPTER FIVE
The World of the Quantum
THE WORD “QUANTUM” is familiar to lay people—splashed on cars and bandied about by Madison Avenue. Pundits and newscasters talk about “quantum leaps,” mostly in a context that has nothing to do with physics.
Most scientists do not use quantum mechanics directly in their work, and those who do usually regard it as a mathematical tool to help them understand the subatomic world, without worrying too much about how that tool squares with our ideas about how the world ought to behave.
Keep two important points in mind as you think about quantum mechanics: (1) no matter how weird it seems, it works—indeed, it has been estimated that a third of the GDP of the United States is ultimately based on quantum mechanics; and (2) the world of the quantum is not anything like the familiar Newtonian world in which we live. Perhaps the human brain is simply not wired to understand this world. Nevertheless, the two basic ideas behind quantum mechanics—what you need to know to be scientifically literate—are quite simple:
Everything—particles, energy, the rate of electron
spin—comes in discrete units.
and
You can’t measure anything without changing it.
Together, these two basic facts explain the operation of atoms, things inside atoms, and things inside things inside atoms.
THE WORLD OF THE VERY SMALL
Quantum mechanics is the branch of science devoted to the study of the behavior of atoms and their constituents. Quantum is the Latin word for “so much” or “bundle,” and “mechanics” is the old term for the study of motion. Thus, quantum mechanics is the study of the motion of things that come in little bundles.
A particle like the electron must come in a “quantized” form. You can have one electron, or two, or three, but never 1.5 or 2.7. It’s not so obvious that something we usually think of as continuous, like light, comes in this form as well. In fact, the quantum or bundle of light is called the “photon” (you may find it useful to remember the “photon torpedoes” of Star Trek fame). It is even less obvious that quantities such as energy and how fast electrons spin come only in discrete bundles as well, but they do. In the quantum world, everything is quantized and can be increased or decreased only in discrete steps.
The behavior of quanta is puzzling at first. The obvious expectation is that when we look at things like electrons, we should find that they behave like microscopic billiard balls—that the world of the very small should behave in pretty much the same way as the ordinary world we experience every day. But an expectation is not the same as a commandment. We can expect the quantum world to be familiar to us, but if it turns out not to be, that doesn’t mean nature is somehow weird or mystical. It just means that things are arranged in such a way that what is “normal” for us at the scale of billiard balls is not “normal” for the universe at the scale of the atom.
THE UNCERTAINTY PRINCIPLE
The strangeness of the quantum world is especially evident in the operation of the uncertainty principle, sometimes called the Heisenberg uncertainty principle after its discoverer, the German physicist Werner Heisenberg (1901–76). The easiest way to understand the uncertainty principle is to think about what it means to say that you “see” something. In order for you to see these words, for exa
mple, light from some source (the sun or a lamp) must strike the book and then travel to your eye. A complex chemical process in your retina converts the energy of the light into a signal that travels to your brain.
Think about the interaction of light with the book. When you look at the book, you do not see it recoil when the light hits it, despite the fact that floods of photons must be bouncing off every second in order for you to see the page. This is the classic Newtonian way of thinking about measurement. It is assumed that the act of measurement (in this case, the act of bouncing light off the book) does not affect the object being measured in any way. Given the infinitesimal energy of the light compared to the energy required to move the book, this is certainly a reasonable way to look at things. After all, baseballs do not jitter around in the air because photographers are using flashbulbs, nor does the furniture in your living room jump every time you turn on the light.
But does this comfortable, reasonable, Newtonian viewpoint apply in the ultrasmall world of the atom? Can you “see” an electron in the same way that you see this book?
If you think about this question for a moment, you will realize that there is a fundamental difference between “seeing” these two objects. You see a book by bouncing light off it, and the light has a negligible effect on the book. You “see” an electron, on the other hand, by bouncing another electron (or some other comparable bundle) off it. In this case, the thing being probed and the thing doing the probing are comparable in every way, and the interaction cannot leave the original electron unchanged. It’s as if the only way you could see a billiard ball was to hit it with another billiard ball.
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