Asimov's New Guide to Science

by Isaac Asimov

  Rutherford himself came upon the beginning of the answer. Between 1906 and 1908, he kept firing alpha particles at a thin foil of metal (such as gold or platinum) to probe its atoms. Most of the projectiles passed right through undeflected (as bullets might pass through the leaves of a tree). But not all did: Rutherford found that, on the photographic plate that served as his target behind the metal, there was an unexpected scattering of hits around the central spot, and some particles bounced back! It was as if some of the bullets had not passed through leaves alone but had ricocheted off something more substantial.

  Rutherford decided that they had hit some sort of dense core, which occupied only a very small part of the volume of the atom. Most of an atom’s volume, it seemed, must be occupied by electrons. As alpha particles charged through the foil of metal, they usually encountered only electrons, and they brushed aside this froth of light particles without being deflected. But once in a while an alpha particle might happen to hit an atom’s denser core, and then it was deflected. That this happened only very occasionally showed that the atomic cores must be very small indeed, because a projectile passing through the metal foil must encounter many thousands of atoms.

  It was logical to suppose that the hard core was made up of protons. Rutherford pictured the protons of an atom as crowded into a tiny atomic nucleus at the center. (It has since been demonstrated that this nucleus has a diameter of little more than 1/100,000 that of the whole atom.)

  This, then, is the basic model of the atom: a positively charged nucleus taking up very little room, but containing almost all the mass of the atom, surrounded by a froth of electrons taking up nearly all the volume of the atom, but containing practically none of its mass. For his extraordinary pioneering work on the ultimate nature of matter, Rutherford received the Nobel Prize in chemistry in 1908.

  It now became possible to describe specific atoms and their behavior in more definite terms. For instance, the hydrogen atom possesses but a single electron. If this is removed, the proton that remains immediately attaches itself to some neighboring molecule. But when the bare hydrogen nucleus does not find an electron to share in this fashion, it acts as a proton—that is to say, a subatomic particle—and in that form it can penetrate matter and react with other nuclei if it has enough energy.

  Helium, with two electrons, does not give one up so easily. As Imentioned in the preceding chapter, its two electrons form a closed shell, and the atom is therefore inert. If helium is stripped of both electrons, however, it becomes an alpha particle—that is, a subatomic particle carrying two units of positive charge.

  The third element, lithium, has three electrons in its atom. Stripped of one or two, it is an ion. If all three of its electrons are removed, it, too, becomes a bare nucleus, carrying a three-unit positive charge.

  The number of units of positive charge in the nucleus of an atom has to be exactly equal to the number of electrons it normally contains, for the atom as a whole is ordinarily neutral. And, in fact, the atomic numbers of the elements are based on their units of positive, rather than negative, charge, because the number of an atom’s electrons may easily be made to vary in ion formation, whereas the number of its protons can be altered only with great difficulty.

  This scheme of the construction of atoms had hardly been worked out when a new conundrum arose. The number of units of positive charge on a nucleus did not balance at all with the nucleus’s mass, except in the case of the hydrogen atom. The helium nucleus, for instance, had a positive charge of two but was known to have four times the mass of the hydrogen nucleus. And the situation got worse and worse as one went down the table of elements, until one reached uranium with a mass equal to 238 protons but a charge equal only to 92.

  How could a nucleus containing four protons (as the helium nucleus was supposed to) have only two units of positive charge? The first, and simplest, guess was that two units of its charge were neutralized by the presence in the nucleus of negatively charged particles of negligible weight. Naturally the electron sprang to mind. The puzzle might be straightened out if one assumed the helium nucleus to consist of four protons and two neutralizing electrons, leaving a net positive charge of two—and so on all the way to uranium, whose nucleus would have 238 protons and 146 electrons, netting 92 units of positive charge. The whole idea was given encouragement by the fact that radioactive nuclei were actually known to emit electrons—that is, beta particles.

  This view of matter prevailed for more than a decade, until a better answer came in a roundabout way from other investigations. But, in the meantime, some serious objections to the hypothesis arose. For one thing, if the nucleus was built essentially of protons, with the light electrons contributing practically nothing to the mass, how was it that the relative masses of the various nuclei did not come to whole numbers? According to the measured atomic weights, the nucleus of the chlorine atom, for instance, had a mass of 35½ times that of the hydrogen nucleus. Did it, then, contain 35½ protons? No scientist (then or now) could accept the idea of half a proton.

  Actually, this particular question has an answer that was discovered even before the main issue was solved. It makes an interesting story in itself.

  Isotopes

  UNIFORM BUILDING BLOCKS

  As early as 1816, an English physician named William Prout had suggested that all atoms were built up from the hydrogen atom. As time went on and the atomic weights were worked out, Prout’s theory fell by the wayside, because it developed that many elements had fractional weights (taking oxygen as the standard at 16). Chlorine has an atomic weight of 35.453. Other examples are antimony, 121.75; barium, 137.34; boron, 10.811; cadmium, 112.40.

  Around the turn of the century there came a series of puzzling observations that was to lead to the explanation. The Englishman William Crookes (he of the Crookes tube) separated from uranium a small quantity of a substance that proved much more radioactive than uranium itself. He suggested that uranium was not radioactive at all—only this impurity, which he called uranium X. Henri Becquerel, on the other hand, discovered that the purified, feebly radioactive uranium somehow increased in radioactivity with time. After it was left standing for a while, the active uranium X could be extracted from it again and again. In other words, uranium was converted by its own radioactivity to the still more active uranium X.

  Then Rutherford similarly separated a strongly radioactive thorium X from thorium and found that thorium, too, went on producing more thorium X. It was already known that the most famous radioactive element of all, radium, broke down to the radioactive gas radon. So Rutherford and his assistant, the chemist Frederick Soddy, concluded that radioactive atoms, in the process of emitting their particles, generally transformed themselves into other varieties of radioactive atoms.

  Chemists began searching for such transformations and came up with an assortment of new substances, giving them such names as radium A, radium B, mesothorium I, mesothorium II, and actinium C. All of them were grouped into three series, depending on their atomic ancestry. One series arose from the breakdown of uranium; another, from that of thorium; and a third, from that of actinium (later it turned out that actinium itself had a predecessor, named protactinium). Altogether, some forty members of these series were identified, each distinguished by its own peculiar pattern of radiation. But the end product of all three series was the same: each chain of substances eventually broke down to the same stable element—lead.

  Now obviously these forty substances could not all be separate elements; between uranium (92) and lead (82) there were only ten places in the periodic table, and all but two of these belonged to known elements. The chemists found, in fact, that though the substances differed in radioactivity, some of them were identical with one another in chemical properties. For instance, as early as 1907, the American chemists Herbert Newby McCoy and William Horace Ross showed that radiothorium, one of the disintegration products of thorium, showed precisely the same chemical behavior as thorium. Radium D behaved chem
ically exactly like lead; in fact, it was often called radiolead. All this suggested that the substances in question were actually varieties of the same element: radiothorium, a form of thorium; radiolead, a member of a family of leads; and so on.

  In 1913, Soddy gave clear expression of this idea and developed it further. He showed that when an atom emits an alpha particle, it changes into an element two places lower in the list of elements; when it emits a beta particle, it changes into an element one place higher. On this basis, radiothorium would indeed fall in thorium’s place in the table, and so would the substances called uranium X1 and uranium Y: all three would be varieties of element 90.

  Likewise, radium D, radium B, thorium B, and actinium B would all share lead’s place as varieties of element 82.

  To the members of a family of substances sharing the same position in the periodic table, Soddy gave the name isotope (from Greek words meaning “same position”). He received the Nobel Prize in chemistry in 1921.

  The proton-electron model of the nucleus (which, nevertheless, eventually proved to be wrong), fitted in beautifully with Soddy’s isotope theory. Removal of an alpha particle from a nucleus would reduce the positive charge of that nucleus by two—exactly what was needed to move it two places down in the periodic table. On the other hand, the ejection of an electron (beta particle) from a nucleus would leave an additional proton unneutralized and thus increase the nucleus’s positive charge by one unit. The effect was to raise the atomic number by one, so the element would move to the next higher position in the periodic table.

  How is it that when thorium breaks down to radiothorium, after going through not one but three disintegrations, the product is still thorium? Well, in the process the thorium atom loses an alpha particle, then a beta particle, then a second beta particle. If we accept the proton building-block idea, the thorium atom has lost four electrons (two supposedly contained in the alpha particle) and four protons. (The actual situation differs from this picture but in a way that does not affect the result.) The thorium nucleus started with 232 protons and 142 electrons (supposedly). Having lost four protons and four electrons, it is reduced to 228 protons and 138 electrons. In either case, the number of unbalanced protons—232 – 142, or 228 – 138—is 90. This still leaves the atomic number 90, the same as before, therefore. So radiothorium, like thorium, has ninety planetary electrons circling around the nucleus. Since the chemical properties of an atom are controlled by the number of its planetary electrons, thorium and radiothorium behave the same chemically, regardless of their difference in atomic weight (232 against 228).

  The isotopes of an element are identified by their atomic weight, or mass number. Thus, ordinary thorium is called thorium 232, while radiothorium is thorium 228. Similarly, the radioactive isotopes of lead are known as lead 210 (radium D), lead 214 (radium B), lead 212 (thorium B), and lead 211 (actinium B).

  The notion of isotopes was found to apply to stable elements as well as to radioactive ones. For instance, it turned out that the three radioactive series I have mentioned ended in three different forms of lead. The uranium series ended in lead 206; the thorium series, in lead 208; and the actinium series, in lead 207. Each of these was an “ordinary,” stable isotope of lead, but the three leads differed in atomic weight.

  Proof of the existence of stable isotopes came from a device invented by an assistant of J. J. Thomson named Francis William Aston. It was an arrangement that separated isotopes very sensitively by virtue of the difference in deflection of their ions by a magnetic field; Aston called it a mass spectrograph. In 1919, using an early version of this instrument, Thomson showed that neon was made up of two varieties of atom: one with a mass number of 20, the other with a mass number of 22. Neon 20 was the common isotope; neon 22 came with it in the ratio of 1 atom in 10. (Later a third isotope, neon 21, was discovered, amounting to only 1 atom in 400 in the neon of the atmosphere.)

  Now the reason for the fractional atomic weights of the elements at least became clear. Neon’s atomic weight of 20.183 represented the composite mass of the three different isotopes making up the element as it was found in nature. Each individual atom had an integral mass number, but the average mass number—the atomic weight—was fractional.

  Aston proceeded to show that several common stable elements were indeed mixtures of isotopes. He found that chlorine, with a fractional atomic weight of 35.453, was made up of chlorine 35 and chlorine 37, in the abundance ratio of 3 to 1. Aston was awarded the Nobel Prize in chemistry in 1922.

  In his address accepting the prize, Aston clearly forecast the possibility of making use of the energy bound in the atomic nucleus, foreseeing both nuclear power plants and nuclear bombs (see chapter 10). In 1935, the Canadian-American physicist Arthur Jeffrey Dempster used Aston’s instrument to take a long step in that direction. He showed that, although 993 of every 1,000 uranium atoms were uranium 238, the remaining seven were uranium 235. This was a discovery fraught with a significance soon to be realized.

  Thus, after a century of false trails, Prout’s idea was finally vindicated. The elements are built of uniform building blocks—if not of hydrogen atoms, at least of units with hydrogen’s mass. The reason the elements do not bear this out in their weights is that they are mixtures of isotopes containing different numbers of building blocks. In fact, even oxygen, whose atomic weight of 16 was used as the standard for measuring the relative weights of the elements, is not a completely pure case. For every 10,000 atoms of common oxygen 16, there are twenty atoms of an isotope with a weight equal to 18 units and four with the mass number 17.

  Actually there are a few elements consisting of a single isotope. (This is a misnomer: to speak of an element as having only one isotope is like saying a woman has given birth to a “single twin.”) The elements of this kind include beryllium, all of whose atoms have the mass number 9; fluorine, made up solely of fluorine 19; aluminum, solely aluminum 27; and a number of others. A nucleus with a particular structure is now called a nuclide, following the suggestion made in 1947 by the American chemist Truman Paul Kohman. One can properly say that an element such as aluminum is made up of a single nuclide.

  TRACKING PARTICLES

  Ever since Rutherford identified the first nuclear particle (the alpha particle), physicists have busied themselves poking around in the nucleus, trying either to change one atom into another or to break it up to see what it is made of. At first they had only the alpha particle to work with. Rutherford made excellent use of it.

  One of the fruitful experiments Rutherford and his assistants carried out involved firing alpha particles at a screen coated with zinc sulfide. Each hit produced a tiny scintillation (an effect first discovered by Crookes in 1903), so that the arrival of single particles could be witnessed and counted with the naked eye. Pursuing this technique, the experimenters put up a metal disk that would block the alpha particles from reaching the zinc sulfide screen so that the scintillations stopped. When hydrogen was introduced into the apparatus, scintillations appeared on the screen despite the blocking metal disk. Moreover, these new scintillations differed in appearance from those produced by alpha particles. Since the metal disk stopped alpha particles, some other radiation must be penetrating it to reach the screen. The radiation, it was decided, must consist of fast protons. In other words, the alpha particles would now and then make a square hit on the nucleus of a hydrogen atom (which consists of a proton, remember) and send it careening forward, as one billiard ball might send another forward on striking it. The struck protons, being relatively light, would shoot forward at great velocity and so could penetrate the metal disk and strike the zinc sulfide screen.

  This detection of single particles by scintillation is an example of a scintillation counter. To make such counts, Rutherford and his assistants first had to sit in the dark for 15 minutes in order to sensitize their eyes and then make their painstaking counts. Modern scintillation counters do not depend on the human eye and mind. Instead, the scintillations are conve
rted to electric pulses that are then counted electronically. The final result need merely be read off from appropriate dials. The counting may be made more practical where scintillations are numerous, by using electric circuits that allow only one in two or in four (or even more) scintillations to be recorded. Such scalers (which scale down the counting, so to speak) were first devised by the English physicist Charles Eryl Wynn-Williams in 1931. Since the Second World War, organic substances have substituted for zinc sulfide and have proved preferable.

  In Rutherford’s original scintillation experiments, there came an unexpected development. When his experiment was performed with nitrogen instead of hydrogen as the target for the alpha-particle bombardment, the zinc sulfide screen still showed scintillations exactly like those produced by protons. Rutherford could only conclude that the bombardment had knocked protons out of the nitrogen nucleus.

  To try to find out just what had happened, Rutherford turned to the Wilson cloud chamber, a device invented in 1895 by the Scottish physicist Charles Thomson Rees Wilson. A glass container fitted with a piston is filled with moisture-saturated air. When the piston is pulled outward, the air abruptly expands and therefore cools. At the reduced temperature. it is supersaturated with the moisture. Under such conditions, any charged particle will cause the water vapor to condense on it. If a particle dashes through the chamber, ionizing atoms in it, a foggy line of droplets will therefore mark its wake.

  The nature of this track can tell a great deal about the particle. The light beta particle leaves a faint, wavering path; the particle is knocked about even in passing near electrons. The much more massive alpha particle makes a straight, thick track. If it strikes a nucleus and rebounds, the path has a sharp bend in it. If it picks up two electrons and becomes a neutral helium atom, its track ends. Aside from the size and character of its track, there are other ways of identifying a particle in the cloud chamber. Its response to an applied magnetic field tells whether it is positively or negatively charged, and the amount of curve indicates its mass and energy. By now physicists are so familiar with photographs of all sorts of tracks that they can read them off as if they were primer print. For the development of his cloud chamber, Wilson shared the Nobel Prize in physics in 1927.