Volume 1 — ELECTRICITY

Units of Electricity

Through the 18th century, scientists had been fascinated by the properties of electricity. Electricity seemed, at the time, to be a very fine fluid that could extend through ordinary matter without taking up any room.

Electricity did more than radiate through matter, however. It also produced important changes in matter. In the first years of the 19th century, it was found that a current of electricity could cause different atoms or different groups of atoms to move in opposite directions through a liquid in which they were dissolved.

The English scientist Michael Faraday (1791-1867) noted in 1832 that a given quantity of electricity seemed to liberate the same number of atoms of a variety of different elements. In some cases, though, it liberated just half the expected number of atoms; or even, in a few cases, just a third.

Scientists began to speculate that electricity, like matter, might consist of tiny units. When electricity broke up a molecule, perhaps a unit of electricity attached itself to each atom. In that case, the same quantity of electricity, containing the same number of units, would liberate the same number of atoms.

In the case of some elements, each atom could attach 2 units of electricity to itself, or perhaps even 3. When that happened a given quantity of electricity would liberate only one-half, or only one-third, the usual number of atoms. (Thus, 18 units of electricity would liberate 18 atoms if distributed 1 to an atom; only 9 atoms if distributed 2 to an atom; and only 6 atoms if distributed 3 to an atom.)

It was understood at the time that electricity existed in two varieties, which were called positive and negative. It appeared that if an atom attached a positive unit of electricity to itself it would be pulled in one direction through the solution by the voltage. If it attached a negative unit of electricity to itself it would be pulled in the other direction.

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Michael Faraday

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The units of electricity were a great deal more difficult to study than the atomic units of matter, and throughout the 19th century they remained elusive. In 1891, though, the Irish physicist George Johnstone Stoney (1826-1911) suggested that the supposed unit of electricity be given a name at least. He called the unit an “electron”.

Cathode Rays

An electric current flows through a closed circuit of some conducting material, such as metal wires. It starts at one pole of a battery, or of some other electricity generating device, and ends at the other. The two poles are the positive pole or “anode” and the negative pole or “cathode”.

If there is a break in the circuit, the current will usually not flow at all. If, however, the break is not a large one, and the current is under a high driving force (which is called the “voltage”), then the current may leap across the break. If two ends of a wire, making up part of a broken circuit, are brought close to each other with nothing but air between, a spark may leap across the narrowing gap before they actually meet and, while it persists, the current will flow despite the break.

The light of the spark, and the crackling sound it makes, are the results of the electric current interacting with molecules of air and heating them. Neither the light nor the sound is the electricity itself. In order to detect the electricity, the current ought to be forced across a gap containing nothing, not even air.

In order to do that, wires would have to be sealed into a glass tube from which all (or almost all) the air was withdrawn. This was not easy to do and it was not until 1854 that Heinrich Geissler (1814-1879), a German glass-blower and inventor, accomplished this feat. The wires sealed 14 into such a “Geissler tube” could be attached to the poles of an electric generator, and if enough voltage was built up, the current would leap across the vacuum.

A Geissler tube.

Such experiments were first performed by the German physicist Julius Plücker (1801-1868). In 1858 he noticed that when the current flowed across the vacuum there was a greenish glow about the wire that was attached to the cathode of the generator. Others studied this glow and finally the German physicist Eugen Goldstein (1850-1931) decided in 1876 that there were rays of some sort beginning at the wire attached to the negatively charged cathode and ending at the part of the tube opposite the cathode. He called them “cathode rays”.

These cathode rays, it seemed, might well be the electric current itself, freed from the metal wires that usually carried it. If so, determining the nature of the cathode rays might reveal a great deal about the nature of the electric current. Were cathode rays something like light and were they made up of tiny waves? Or were they a stream of particles possessing mass?

There were physicists on each side of the question. By 1885, however, the English physicist William Crookes 15 (1832-1919) showed that cathode rays could be made to turn a small wheel when they struck that wheel on one side. This seemed to show that the cathode rays possessed mass and were a stream of atom-like particles, rather than a beam of mass-less light. Furthermore, Crookes showed that the cathode rays could be pushed sideways in the presence of a magnet. (This effect, when current flows in a wire, is what makes a motor work.) This meant that, unlike either light or ordinary atoms, the cathode rays carried an electric charge.

J. J. Thomson in his laboratory. On his right are early X-ray pictures.

This view of the cathode rays as consisting of a stream of electrically charged particles was confirmed by another English physicist, Joseph John Thomson (1856-1940). In 1897 he showed that the cathode rays could also be made to take a curved path in the presence of electrically charged 16 objects. The particles making up the cathode rays were charged with negative electricity, judging from the direction in which they were made to curve by electrically charged objects.

Thomson had no hesitation in maintaining that these particles carried the units of electricity that Faraday’s work had hinted at. Eventually, Stoney’s name for the units of electricity was applied to the particles that carried those units. The cathode rays, in other words, were considered to be made up of streams of electrons and Thomson is usually given credit for having discovered the electron.

The extent to which cathode rays curved in the presence of a magnet or electrically charged objects depended on the size of the electric charge on the electrons and on the mass of the electrons. Ordinary atoms could be made to carry an electric charge and by comparing their behavior with those of electrons, some of the properties of electrons could be determined.

There were, for instance, good reasons to suppose that the electron carried a charge of the same size as one that a hydrogen atom could be made to carry. The electrons, however, were much easier to pull out of their straight-line path than the charged hydrogen atom was. The conclusion drawn from this was that the electron had much less mass than the hydrogen atom.

Thomson was able to show, indeed, that the electron was much lighter than the hydrogen atom, which was the lightest of all the atoms. Nowadays we know the relationship quite exactly. We know that it would take 1837.11 electrons to possess the mass of a single hydrogen atom. The electron is therefore a “subatomic particle”; the first of this sort to be discovered.

In 1897, then, two types of mass-containing particles were known. There were the atoms, which made up ordinary matter, and the electrons, which made up electric current.

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Radioactivity

Was there a connection between these two sets of particles—atoms and electrons? In 1897, when the electron was discovered, a line of research that was to tie the two kinds of particles together had already begun.

In 1895 the German physicist Wilhelm Konrad Roentgen (1845-1923) was working with cathode rays. He found that if he made the cathode rays strike the glass at the other end of the tube, a kind of radiation was produced. This radiation was capable of penetrating glass and other matter. Roentgen had no idea as to the nature of the radiation, and so called it “X rays”. This name, containing “X” for “unknown”, was retained even after physicists worked out the nature of X rays and found them to be light-like radiation made up of waves much shorter than those of ordinary light.

Antoine Henri Becquerel.

At once, physicists became fascinated with X rays and began searching for them everywhere. One of those involved in the search was the French physicist Antoine Henri Becquerel (1852-1908). A certain compound, potassium uranyl sulfate, glowed after being exposed to sunlight and Becquerel wondered if this glow, like the glow on the glass in Roentgen’s X-ray tube, contained X rays.

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Wilhelm Roentgen and his laboratory at the University of Würzburg.

It did, but while investigating the problem in 1896, Becquerel found that the compound was giving off invisible penetrating X-ray-like radiation continually, whether it was exposed to sunlight or not. The radiation was detected because it would fog a photographic plate just as light would. What’s more, the radiation would fog the plate, even if the plate were wrapped in black paper, so that it could penetrate matter just as X rays could.

Others, in addition to Becquerel, were soon investigating the new phenomenon. In 1898 the Polish (later French) 20 physicist Marie Sklodowska Curie (1867-1934) showed that it was the uranium atom that was the source of the radiation, and that any compound containing the uranium atom would give off these penetrating rays.

Until then, uranium had not been of much interest to chemists. It was a comparatively rare metal that was first discovered in 1789 by the German chemist Martin Heinrich Klaproth (1743-1817). It had no particular uses and remained an obscure element. As chemists learned to work out the atomic weights of the various elements, they found, however, that, of the elements then known, uranium had the highest atomic weight of all—238.

Once uranium was discovered to be an endless source of radiation, it gained interest that has risen ever since. Madame Curie gave the name “radioactivity” to this phenomenon of continuously giving off rays. Uranium was the first element found to be radioactive.

It did not remain alone, however. It was soon shown that thorium was also radioactive. Thorium, which had been discovered in 1829 by Berzelius, was made up of atoms that were the second most massive known at the time. Thorium’s atomic weight is 232.

But what was the mysterious radiation emitted by uranium and thorium?

Almost at once it was learned that whatever the radiation was, it was not uniform in properties. In 1899 Becquerel (and others) showed that, in the presence of a magnet, some of the radiation swerved in a particular direction. Later it was found that a portion of it swerved in the opposite direction. Still another part didn’t swerve at all but moved on in a straight line.

The conclusion was that uranium and thorium gave off three kinds of radiation. One carried a positive charge of electricity, one a negative charge, and one no charge at all. The New Zealand-born physicist Ernest Rutherford (1871-1937) called the first two kinds of radiation “alpha rays” and “beta rays”, after the first two letters of the Greek alphabet. The third was soon called “gamma rays” after the third letter.

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Ernest Rutherford

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Marie Curie and her two daughters, Eve (left) and Irene, in 1908.

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Pierre Curie during a class lecture in 1906, the year of his death.

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The gamma rays eventually turned out to be another light-like form of radiation, with waves even shorter than those of X rays. The alpha rays and beta rays, which carried electric charges, seemed to be streams of charged particles (“alpha particles” and “beta particles”) just as the cathode rays had turned out to be.

In 1900, indeed, Becquerel studied the beta particles and found them to be identical in mass and charge with electrons. They were electrons.

By 1906 Rutherford had worked out the nature of the alpha particles. They carried a positive electric charge that was twice as great as the electron’s negative charge. If an electron carried a charge that could be symbolized as -, then the charge of the alpha particle was ++. Furthermore, the alpha particle was much more massive than the electron. It was, indeed, as massive as a helium atom (the second lightest known atom) and four times as massive as a hydrogen atom. Nevertheless, the alpha particle can penetrate matter in a way in which atoms cannot, so that it seems much smaller in diameter than atoms are. The alpha particle, despite its mass, is another subatomic particle.

Here, then, is the meeting point of electrons and of atoms—the particles of electricity and of matter.

Ever since Dalton had first advanced the atomic theory over a century earlier, chemists had assumed that atoms were the fundamental units of matter. They had assumed atoms were as small as anything could be and that they could not possibly be broken up into anything smaller. The discovery of the electron, however, had shown that some particles, at least, might be far smaller than any atom. Then, the investigations into radioactivity had shown that atoms of uranium and thorium spontaneously broke up into smaller particles, including electrons and alpha particles.

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It would seem, then, that atoms of these elements and, presumably, of all elements, were made up of still smaller particles and that among these particles were electrons. The atom had a structure and physicists became interested in discovering exactly what that structure was.

The Structure of the Atom

Since radioactive atoms gave off either positively charged particles or negatively charged particles, it seemed reasonable to assume that atoms generally were made up of both types of electricity. Furthermore, since the atoms in matter generally carried no charge at all, the normal “neutral atom” must be made up of equal quantities of positive charge and negative charge.

It turned out that only radioactive atoms, such as those of uranium and thorium, gave off positively charged alpha particles. Many atoms, however, that were not radioactive, could be made to give off electrons. In 1899 Thomson showed that certain perfectly normal metals with no trace of radioactivity gave off electrons when exposed to ultraviolet light. (This is called the “photoelectric effect”.)

It was possible to suppose, then, that the main structure of the atom was positively charged and generally immovable, and that there were also present light electrons, which could easily be detached. Thomson had suggested, as early as 1898, that the atom was a ball of matter carrying a positive charge and that individual electrons were stuck throughout its substance, like raisins in pound cake.

If something like the Thomson view were correct then the number of electrons, each with one unit of negative electricity, would depend on the total size of the positive charge carried by the atom. If the charge were +5, there would have to be 5 electrons present to balance that. The total charge would then be 0 and the atom as a whole would be electrically neutral.

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If, in such a case, an electron were removed, the atomic charge of +5 would be balanced by only 4 electrons with a total charge of -4. In that case, the net charge of the atom as a whole would be +1. On the other hand, if an extra electron were forced onto the atom, the charge of +5 would be balanced by 6 electrons with a total charge of -6, and the net charge of the atom as a whole would be -1.

Such electrically charged atoms were called “ions” and their existence had been suspected since Faraday’s day. Faraday had known that atoms had to travel through a solution under the influence of an electric field to account for the way in which metals and gases appeared at the cathode and anode. It was he who first used the term, ion, from a Greek word meaning “traveller”. The word had been suggested to him by the English scholar, William Whewell (1794-1866). In 1884 the Swedish chemist Svante August Arrhenius (1859-1927) had first worked out a detailed theory based on the suggestion that these ions were atoms or groups of atoms that carried an electric charge.

Svante A. Arrhenius

By the close of the 19th century, then, Arrhenius’s suggestion seemed correct. There were positive ions made up of atoms or groups of atoms, from which one or more of the electrons within the atoms had been removed. There were negative ions made up of single atoms or of groups of atoms, to which one or more extra electrons had been added.

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Although Thomson’s model of the atom explained the existence of ions and the fact that atoms could give off electrons or absorb them, it was not satisfactory in all ways. Further investigations yielded results not compatible with the raisins-in-the-pound-cake notion.

In 1906 Rutherford began to study what happened when massive subatomic particles, such as alpha particles, passed through matter. When alpha particles passed through a thin film of gold, for instance, they raced through, for the most part, as though nothing were there. The alpha particles seemed to push the light electrons aside and to act as though the positively charged main body of the atom that Thomson had pictured was not solid, but was soft and spongy.

The only trouble was that every once in a while an alpha particle seemed to strike something in the gold film and bounce to one side. Sometimes it even bounced directly backward. It was as though somewhere in each atom there was something at least as massive as the alpha particle.

How large was this massive portion of the atom? It couldn’t be very large for if it were the alpha particles would hit it frequently. Instead, the alpha particles made very few hits. This meant the massive portion was very small and that most alpha particles tore through the atom without coming anywhere near it.

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Rutherford’s alpha particle bombardment apparatus. A piece of radium in the lead box (B) emits alpha particles that go through the gold foil (F). These particles are scattered at different angles onto the fluorescent screen (S), where the flashes caused by each impact are seen through the microscope (M). Below, alpha particles are shown bouncing off a nucleus in the gold foil.

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By 1911 Rutherford announced his results to the world. He suggested that just about all the mass of the atom was concentrated into a very tiny, positively charged “nucleus” at its center. The diameter of the nucleus was only about 1/10,000 the diameter of the atom. All the rest of the atom was filled with the very light electrons.

Hans Geiger (left) and Ernest Rutherford at Manchester University about 1910.

According to Rutherford’s notion, the atom consisted of a single tiny positively charged lead shot at the center of a foam of electrons. It was Thomson’s notion in reverse. Still, the nucleus carried a positive charge of a particular size and was balanced by negatively charged electrons. Rutherford’s 30 model of the atom explained the existence of ions just as easily as Thomson’s did and it explained more besides.

For instance, if all the electrons are removed so that only the nucleus remains, this nucleus is as massive as an atom but is so tiny in size that it can penetrate matter. The alpha particle would be a bare atomic nucleus from this point of view.

Rutherford’s model of the “nuclear atom” is still accepted today.

Atomic Numbers

Since the atom consisted of a positively charged nucleus at the center, and a number of negatively charged electrons outside, the next step was to find the exact size of the nuclear charge and the exact number of electrons for the different varieties of atoms.

The answer came through a line of research that began with the English physicist Charles Glover Barkla (1877-1944). In 1911 he noted that when X rays passed through atoms, some were absorbed and some bounced back. Those that bounced back had a certain ability to penetrate other matter. When the X rays struck atoms of high atomic weight, the X rays that bounced back were particularly penetrating. In fact, each different type of atom seemed associated with reflected X rays of a particular penetrating power, so Barkla called these “characteristic X rays”.

In 1913 another English physicist, Henry Gwyn-Jeffreys Moseley (1887-1915), went into the matter more thoroughly. He measured the exact wavelength of the characteristic X rays by reflecting them from certain crystals. In crystals, atoms are arranged in regular order and at known distances from each other. X rays reflecting from (or more accurately, diffracting from) crystals are bent out of their path by the rows of atoms. The longer their waves, the more they are bent. From the degree of bending the wavelength of the waves can be determined.

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Charles Glover Barkla

Henry Gwyn-Jeffreys Moseley

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Moseley found that the greater the atomic weight of an atom, the shorter the waves of the characteristic X rays associated with it and the more penetrating those X rays were. There was such a close connection, in fact, that Moseley could arrange the elements in order according to the wavelength of the characteristic X rays.

For some 40 years prior to this, the elements had been listed in order of atomic weight. This was useful especially since the Russian chemist Dmitri I. Mendeléev (1834-1907) had arranged them in a “periodic table” based on the atomic weight order in such a way that elements of similar properties were grouped together. The elements in this table were sometimes numbered consecutively (“atomic number”) but this was inconvenient since, when new elements were discovered, the list of atomic numbers might have to be reorganized.

Dmitri Mendeléev and Bohuslav Brauner in Prague in 1900. Brauner was a professor of chemistry at the Bohemian University in Prague.

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The Danish physicist Niels Bohr (1885-1962) had just advanced a theory of atomic structure that made it reasonable to suppose that the wavelength of the characteristic X rays depended on the size of the nuclear charge of the atoms making up a particular element. Moseley therefore suggested that these X rays be used to determine the size of the positive charge on its nucleus. The atomic number could then be set equal to that charge and be made independent of new discoveries of elements.

Hydrogen, for instance, has an atomic number of 1. Its nucleus carries a unit positive charge, +1, and the hydrogen atom possesses 1 electron to balance this. Helium, with an atomic number of 2, has a nuclear charge of +2 and 2 electrons, with a total charge of -2, to balance it. (The alpha particle released by radioactive atoms is identical with a helium nucleus.)

The atomic number increases as one goes up the line of atoms. Oxygen atoms, for instance, have an atomic number of 8 and iron atoms have one of 26. At the upper end, thorium is 90 and uranium is 92. Each uranium atom has a nucleus bearing a charge of +92 and contains 92 electrons to balance this.

Once the notion of the atomic number was worked out, it became possible to tell for certain whether any elements remained as yet undiscovered and, if so, where in the list they might be.

Thus, when Moseley first presented scientists with the atomic number it turned out that there were still 7 elements that were not discovered. At least elements with atomic numbers of 43, 61, 72, 75, 85, 87, and 91 were still not known. By 1945, all seven had been discovered.

It quickly turned out that the atomic number was more fundamental and more characteristic of a particular element than was the atomic weight.

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Niels Bohr

Bohr’s study.

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Since Dalton’s time it had been assumed that all the atoms of a particular element were of equal atomic weight and that atoms of two different elements were always of different atomic weight. The first inkling and the first proof that this might not be so came through the study of radioactivity.

Isotopes

In 1902 Rutherford and his co-worker Frederick Soddy (1877-1956) showed that when uranium atoms gave off alpha particles, a new kind of atom was formed that was not uranium at all. It was this new atom that was eventually found to give off a beta particle, and then another atom of still another element was formed. This work of Rutherford and Soddy began a line of investigation that by 1907 had shown that there was a whole radioactive chain of elements, each one breaking down to the next in line by giving off either an alpha particle or a beta particle, until finally a lead atom was formed that was not radioactive.

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Frederick Soddy

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There was, in short, a “radioactive series” beginning with uranium (atomic number 92) and ending with lead (atomic number 82). The same was true of thorium (atomic number 90), which began a series that also ended with lead. Still a third element, actinium (atomic number 89) was, at that time, the first known member of a series that also ended in lead.

The various atoms formed in these three radioactive series were not all different in every way. When the uranium atom gives off an alpha particle, it forms an atom originally called “uranium X₁”. On close investigation, it turned out that this uranium X₁ had the chemical properties of thorium. Uranium X₁, had, however, radioactive properties different from ordinary thorium.

Uranium X₁ broke down so rapidly, giving off beta particles as it did so, that half of any given quantity would have broken down in 24 days. Another way of saying this (which was introduced by Rutherford) was that the “half-life” of uranium X₁, is 24 days. Ordinary thorium, however, gives off alpha particles, not beta particles, and does so at such a slow rate, that its half-life is 14 billion years!

Uranium X₁, and ordinary thorium were in the same place in the list of elements by chemical standards, and yet there was clearly something different about the two.

Here is another case. In 1913 the British chemist Alexander Fleck (1889- ) studied “radium B” and “radium D”, the names given to two different kinds of atoms in the uranium radioactive series. He also studied “thorium B” in the thorium radioactive series and “actinium B” in the actinium radioactive series. All four are chemically the same as ordinary lead; all four are in the same place in the list of elements. Yet each is different from the radioactive standpoint. Though all give off beta particles, radium B has a 38 half-life of 27 minutes, radium D one of 19 years, thorium B one of 11 hours, and actinium B one of 36 minutes.

In 1913 Soddy called atoms that were in the same place in the list of elements, but which had different radioactive properties, “isotopes”, from Greek words meaning “same place”.

At first, it seemed that the only difference between isotopes might be in their radioactive properties and that only radioactive atoms were involved. Quickly that proved not to be so.

It proved that it was possible to have several forms of the same element that were all different even though none of them were radioactive. The uranium series, the thorium series, and the actinium series all ended in lead. In each case the lead formed was stable (not radioactive). Were the lead atoms identical in every case? Soddy had worked out the way in which atomic weights altered every time an alpha particle or a beta particle was given off by an atom. Working through the three radioactive series he decided that the lead atoms had different atomic weights in each case.

The uranium series ought to end with lead atoms that had an atomic weight of 206. The thorium series ought to end in lead atoms with an atomic weight of 208 and the actinium series in lead atoms with an atomic weight of 207.

If this were so, there would be 3 lead isotopes that would differ not in radioactive properties, but in atomic weight. The isotopes could be referred to as lead-206, lead-207, and lead-208. If we use the chemical symbol for lead (Pb), we could write the isotopes, ²⁰⁶Pb, ²⁰⁷Pb, and ²⁰⁸Pb. (We read the symbol ²⁰⁶Pb as lead-206.) Atomic weight measurements made in 1914 by Soddy and others supported that theory.

All 3 lead isotopes had the same atomic number of 82. The atoms of all 3 isotopes had nuclei with an electric charge of +82 and all 3 had 82 electrons in the atom to balance that positive nuclear charge. The difference was in the mass of the nucleus only.

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Isotopes of two elements.

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But what of ordinary lead that existed in the rocks far removed from any radioactive substances and that had presumably been stable through all the history of earth? Its atomic weight was 207.2.

Was the stable lead that had no connection with radioactivity made up of atoms of still another isotope, one with a fractional atomic weight? Or could stable lead be made up of a mixture of isotopes, each of a different whole-number atomic weight and was the overall atomic weight a fraction only because it was an average?

It was at the moment difficult to tell in the case of lead, but an answer came in connection with another element, the rare gas neon (atomic symbol Ne), which has an atomic weight of 20.2.

Was that fractional atomic weight something that was possessed by all neon atoms without exception or was it the average of some lightweight atoms and some heavyweight ones? It would be a matter of crucial importance if isotopes of neon could be found, for neon had nothing to do with any of the radioactive series. If neon had isotopes then any element might have them.

In 1912 Thomson was working on neon. He sent a stream of cathode-ray electrons through neon gas. The electrons smashed into the neon atoms and knocked an electron off some of them. That left a neon ion carrying a single positive charge—an ion that could be written Ne⁺.

The neon ions move in the electric field as electrons do, but in the opposite direction since they have an opposite charge. In the combined presence of a magnet and of an electric field, the neon ions move in a curved path. If all the neon ions had the same mass, all would follow the same curve. If some were more massive than others, the more massive ones would curve less.

The neon ions ended on a photographic plate, which was darkened at the point of landing. There were two regions of 41 darkening, because there were neon ions of two different masses that curved in two different degrees and ended in two different places. Thomson showed, from the amount of curving, that there was a neon isotope with an atomic weight of 20 and one with an atomic weight of 22—²⁰Ne and ²²Ne.

What’s more, from the intensity of darkening, it could be seen that ordinary neon was made up of atoms that were roughly 90% ²⁰Ne and 10% ²²Ne. The overall atomic weight of neon, 20.2, was the average atomic weight of these 2 isotopes.

Thomson’s instrument was the first one capable of separating isotopes and such instruments came to be called “mass spectrometers”. The first to use the name was the English physicist Francis William Aston (1877-1945), who built the first efficient instrument of this type in 1919.

He used it to study as many elements as he could. He and those who followed him located many isotopes and determined the frequency of their occurrence with considerable precision. It turned out, for instance, that neon is actually 90.9% ²⁰Ne, and 8.8% ²²Ne. Very small quantities of still a third isotope, ²¹Ne, are also present, making up 0.3%.

As for ordinary lead in nonradioactive rocks, it is made up of 23.6% ²⁰⁶Pb, 22.6% ²⁰⁷Pb, and 52.3% ²⁰⁸Pb. There is still a fourth isotope, ²⁰⁴Pb, which makes up the remaining 1.5% and which is not the product of any radioactive series at all.

The isotopes always have atomic weights that are close to, but not quite, whole numbers. Any atomic weight of an element that departs appreciably from an integer does so only because it is an average of different isotopes. For instance, the atomic weight of chlorine (chemical symbol Cl) is 35.5, but this is because it is made up of a mixture of 2 isotopes. About one quarter of chlorine’s atoms are ³⁷Cl and about three-quarters are ³⁵Cl.

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Francis W. Aston

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Mass spectrograph as used by Thomson and Aston to measure the atomic weight of neon.

To avoid confusion, the average mass of the isotopes that make up a particular element is still called the atomic weight of that element. The integer closest to the mass of the individual isotope is spoken of as the “mass number” of that isotope. Thus, chlorine is made up of isotopes with mass numbers 35 and 37, but the atomic weight of chlorine as it is found in nature is 35.5 (or, to be more accurate, 35.453).

In the same way, ordinary lead is made up of isotopes with mass numbers 204, 206, 207, and 208, and its atomic weight is 207.19; neon is made up of isotopes with mass numbers 20, 21, and 22, and its atomic weight is 20.183, and so on.

If the atomic weight of some element happens to be very close to a whole number to begin with, it may consist of a single kind of atom. For instance, the gas fluorine (chemical symbol F) has an atomic weight of nearly 19, while that of the metal sodium (chemical symbol Na) is nearly 23. As it turns out, all the atoms of fluorine are of the single variety ¹⁹F, while all the atoms of sodium are ²³Na.

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Sometimes the atomic weight of an element, as it occurs in nature, is nearly a whole number and yet it is made up of more than 1 isotope. In that case, one of the isotopes makes up very nearly all of it, while the others are present in such minor quantities that the average is hardly affected.

Helium, for instance (atomic symbol He) has an atomic weight of just about 4 and, indeed, almost all the atoms making it up are ⁴He. However, 0.0001% of the atoms, or one out of a million, are ³He. Again, 99.6% of all the nitrogen atoms (atomic symbol N) are ¹⁴N, but 0.4% are ¹⁵N. Then, 98.9% of all carbon atoms (atomic symbol C) are ¹²C, but 1.1% are ¹³C. It is not surprising that the atomic weights of nitrogen and carbon are just about 14 and 12, respectively.

Harold Urey

Even hydrogen does not escape. Its atomic weight is just about 1 and most of its atoms are ¹H. The American chemist Harold Clayton Urey (1893- ) detected the existence of a 45 more massive isotope, ²H. This isotope has almost twice the mass of the lighter one. No other isotopes of a particular atom differ in mass by so large a factor. For that reason ²H and ¹H differ in ordinary chemical properties more than isotopes usually do and Urey therefore gave ²H the special name of “deuterium” from a Greek word meaning “second”.

W. F. Giauque

In 1929 the American chemist William Francis Giauque (1895- ) found that oxygen was composed of more than 1 isotope. Its atomic weight had been set arbitrarily at 16.0000 so it was a relief that 99.76% of its atoms were ¹⁶O. However, 0.20% were ¹⁸O, and 0.04% were ¹⁷O.

As you see, ¹⁶O must have a mass number of slightly less than 16.0000 and it must be the more massive isotopes ¹⁷O and ¹⁸O that pull the average up to 16.0000. Disregarding this, chemists clung to a standard atomic weight of 16.000 for oxygen as it appeared in nature, preferring not to concern themselves with the separate isotopes.

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Physicists, however, felt uneasy at using an average as standard for they were more interested in working with individual isotopes. They preferred to set ¹⁶O at 16.0000 so that the average atomic weight of oxygen was 16.0044 and all other atomic weights rose in proportion. Atomic weights determined by this system were “physical atomic weights”.

Finally, in 1961, a compromise was struck. Chemists and physicists alike decided to consider the atomic weight of ¹²C as exactly 12 and to use that as a standard. By this system, the atomic weight of oxygen became 15.9994, which is only very slightly less than 16.

The radioactive elements did not escape this new view either. The atomic weight of uranium (chemical symbol U) is just about 238 and, indeed, most of its atoms are ²³⁸U. In 1935, however, the Canadian-American physicist, Arthur Jeffrey Dempster (1886-1950), found that 0.7% of its atoms were a lighter isotope, ²³⁵U.

These differed considerably in radioactive properties. The common uranium isotope, ²³⁸U, had a half-life of 4500 million years, while ²³⁵U had a half-life of only 700 million years. Furthermore ²³⁵U broke down in three stages to actinium. It was ²³⁵U, not actinium itself, that was the beginning of the actinium radioactive series.

As for thorium (atomic symbol Th) with an atomic weight of 232, it did indeed turn out that in the naturally occurring element virtually all the atoms were ²³²Th.

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