Worlds Within Worlds: The Story of Nuclear Energy, Volume 1 (of 3) - 1

Worlds Within Worlds:
The Story of Nuclear Energy
Volume 1
Atomic Weights · Energy · Electricity

by Isaac Asimov

U. S. Energy Research and Development Administration
Office of Public Affairs
Washington, D.C. 20545
Library of Congress Catalog Card Number: 75-189477
1972
_Nothing in the history of mankind has opened our eyes to the
possibilities of science as has the development of atomic power. In the
last 200 years, people have seen the coming of the steam engine, the
steamboat, the railroad locomotive, the automobile, the airplane, radio,
motion pictures, television, the machine age in general. Yet none of it
seemed quite so fantastic, quite so unbelievable, as what man has done
since 1939 with the atom ... there seem to be almost no limits to what
may lie ahead: inexhaustible energy, new worlds, ever-widening knowledge
of the physical universe._
Isaac Asimov
[Illustration: Photograph of night sky]


The U. S. Energy Research and Development Administration publishes a
series of booklets for the general public.
Please write to the following address for a title list or for
information on a specific subject:
USERDA—Technical Information Center
P. O. Box 62
Oak Ridge, Tennessee 37830
[Illustration: Isaac Asimov]


ISAAC ASIMOV received his academic degrees from Columbia University and
is Associate Professor of Biochemistry at the Boston University School
of Medicine. He is a prolific author who has written over 150 books in
the past 20 years, including about 20 science fiction works, and books
for children. His many excellent science books for the public cover
subjects in mathematics, physics, astronomy, chemistry, and biology,
such as _The Genetic Code_, _Inside the Atom_, _Building Blocks of the
Universe_, _Understanding Physics_, _The New Intelligent Man’s Guide to
Science_, and _Asimov’s Biographical Encyclopedia of Science and
Technology_.
In 1965 Dr. Asimov received the James T. Grady Award of the American
Chemical Society for his major contribution in reporting science
progress to the public.
[Illustration: Photograph of night sky]

VOLUME 1
Introduction 5
Atomic Weights 6
Electricity 11
Units of Electricity 11
Cathode Rays 13
Radioactivity 17
The Structure of the Atom 25
Atomic Numbers 30
Isotopes 35
Energy 47
The Law of Conservation of Energy 47
Chemical Energy 50
Electrons and Energy 54
The Energy of the Sun 55
The Energy of Radioactivity 57

VOLUME 2
Mass and Energy 69
The Structure of the Nucleus 75
The Proton 75
The Proton-Electron Theory 76
Protons in Nuclei 80
Nuclear Bombardment 82
Particle Accelerators 86
The Neutron 92
Nuclear Spin 92
Discovery of the Neutron 95
The Proton-Neutron Theory 98
The Nuclear Interaction 101
Neutron Bombardment 107

VOLUME 3
Nuclear Fission 117
New Elements 117
The Discovery of Fission 122
The Nuclear Chain Reaction 127
The Nuclear Bomb 131
Nuclear Reactors 141
Nuclear Fusion 147
The Energy of the Sun 147
Thermonuclear Bombs 149
Controlled Fusion 151
Beyond Fusion 159
Antimatter 159
The Unknown 164
Reading List 166
[Illustration: _A total eclipse of the sun._]


INTRODUCTION

In a way, nuclear energy has been serving man as long as he has existed.
It has served all of life; it has flooded the earth for billions of
years. The sun, you see, is a vast nuclear engine, and the warmth and
light that the sun radiates is the product of nuclear energy.
In order for man to learn to produce and control nuclear energy himself,
however (something that did not take place until this century), three
lines of investigation—atoms, electricity, and energy—had to develop and
meet.
We will begin with atoms.


ATOMIC WEIGHTS

As long ago as ancient Greek times, there were men who suspected that
all matter consisted of tiny particles which were far too small to see.
Under ordinary circumstances, they could not be divided into anything
smaller, and they were called “atoms” from a Greek word meaning
“indivisible”.
It was not until 1808, however, that this “atomic theory” was really put
on a firm foundation. In that year the English chemist John Dalton
(1766-1844) published a book in which he discussed atoms in detail.
Every element, he suggested, was made up of its own type of atoms. The
atoms of one element were different from the atoms of every other
element. The chief difference between the various atoms lay in their
mass, or weight.[1]
Dalton was the first to 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”.)
[Illustration: _John Dalton_]
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.
[Illustration: _Jöns Jakob Berzelius_]
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.


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.
[Illustration: _Michael Faraday_]
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 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.
[Illustration: _A Geissler tube._ labelled: Current source]
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.
[Illustration: _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 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.

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.
[Illustration: _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.
[Illustration: Roentgen’s laboratory]
[Illustration: _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) 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.
[Illustration: _Ernest Rutherford_]
[Illustration: _Marie Curie and her two daughters, Eve (left) and Irene,
in 1908._]
[Illustration: _Pierre Curie during a class lecture in 1906, the year of
his death._]
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.