Worlds Within Worlds: The Story of Nuclear Energy, Volume 1 (of 3) Atomic Weights; Energy; Electricity — Edition Insights

o try to determine what these masses might be. He could not work out the actual masses in ounces or grams, for atoms were far too tiny to weigh with any of his instruments. He could, however, determine their relative weights; that is, how much more massive one kind of atom might be than another.

For instance, he found that a quantity of hydrogen gas invariably combined with eight times its own mass of oxygen gas to form water. He guessed that water consisted of combinations of 1 atom of hydrogen with 1 atom of oxygen. (A combination of atoms is called a “molecule” from a Greek word meaning “a small mass”, and so hydrogen and oxygen atoms can be said to combine to form a “water molecule”.)

To account for the difference in the masses of the combining gases, Dalton decided that the oxygen atom was eight times as massive as the hydrogen atom. If he set the mass of the hydrogen atom at 1 (just for convenience) then the mass of the oxygen atom ought to be set at 8. These comparative, or relative, numbers were said to be “atomic weights”, so that what Dalton was suggesting was that the atomic weight of hydrogen was 1 and the atomic weight of oxygen was 8. By noting the quantity of other elements that combined with a fixed mass of oxygen or of hydrogen, Dalton could work out the atomic weights of these elements as well.

Dalton’s idea was right, but his details were wrong in some cases. For instance, on closer examination it turned out that the water molecule was composed of 1 oxygen atom and 2 hydrogen atoms. For this reason, the water molecule may be written H₂O, where H is the chemical symbol for a hydrogen atom, and O for an oxygen atom.

It is still a fact that a quantity of hydrogen combines with eight times its mass of oxygen, so the single oxygen atom must be eight times as massive as the 2 hydrogen atoms taken together. The oxygen atom must therefore be sixteen times as massive as a single hydrogen atom. If the atomic weight of hydrogen is 1, then the atomic weight of oxygen is 16.

At first it seemed that the atomic weights of the various elements were whole numbers and that hydrogen was the lightest one. It made particular sense, then, to consider the atomic weight of hydrogen as 1, because that made all the other atomic weights as small as possible and therefore easy to handle.

The Swedish chemist Jöns Jakob Berzelius (1779-1848) continued Dalton’s work and found that elements did not combine in quite such simple ratios. A given quantity of hydrogen actually combined with a little bit less than eight times its mass of oxygen. Therefore if the atomic weight of hydrogen were considered to be 1, the atomic weight of oxygen would have to be not 16, but 15.87.

As it happens, oxygen combines with more elements (and more easily) than hydrogen does. The ratio of its atomic weight to that of other elements is also more often a whole number. In working out the atomic weight of elements it was therefore more convenient to set the atomic weight of oxygen at a whole number than that of hydrogen. Berzelius did this, for instance, in the table of atomic weights he published in 1828. At first he called the atomic weight of oxygen 100. Then he decided to make the atomic weights as small as possible, without allowing any atomic weight to be less than 1. For that reason, he set the atomic weight of oxygen at exactly 16 and in that case, the atomic weight of hydrogen had to be placed just a trifle higher than 1. The atomic weight of hydrogen became 1.008. This system was retained for nearly a century and a half.

Throughout the 19th century, chemists kept on working out atomic weights more and more carefully. By the start of the 20th century, most elements had their atomic weights worked out to two decimal places, sometimes three.

A number of elements had atomic weights that were nearly whole numbers on the “oxygen = 16” standard. The atomic weight of aluminum was just about 27, that of calcium almost 40, that of carbon almost 12, that of gold almost 197, and so on.

On the other hand, some elements had atomic weights that were far removed from whole numbers. The atomic weight of chlorine was close to 35.5, that of copper to 63.5, that of iron to 55.8, that of silver to 107.9, and so on.

Throughout the 19th century, chemists did not know why so many atomic weights were whole numbers, while others weren’t. They simply made their measurements and recorded what they found. For an explanation, they had to wait for a line of investigation into electricity to come to fruition.

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.

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”.

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 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.

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 (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.

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 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.

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.

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.

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) 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.

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.

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.

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.