Worlds Within Worlds: The Story of Nuclear Energy, Volume 2 (of 3) - 1

Worlds Within Worlds:
The Story of Nuclear Energy
Volume 2
Mass and Energy · The Neutron · The Structure of the Nucleus

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
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[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]
CONTENTS

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 field-ion microscope view of atoms in a crystal. Each
tiny white dot is a single atom, and each ring system is a crystal facet
or plane. The picture is magnified 1,500,000 times._]


MASS AND ENERGY

In 1900 it began to dawn on physicists that there was a vast store of
energy within the atom; a store no one earlier had imagined existed. The
sheer size of the energy store in the atom—millions of times that known
to exist in the form of chemical energy—seemed unbelievable at first.
Yet that size quickly came to make sense as a result of a line of
research that seemed, at the beginning, to have nothing to do with
energy.
Suppose a ball were thrown forward at a velocity of 20 kilometers per
hour by a man on top of a flatcar that is moving forward at 20
kilometers an hour. To someone watching from the roadside the ball would
appear to be travelling at 40 kilometers an hour. The velocity of the
thrower is added to the velocity of the ball.
If the ball were thrown forward at 20 kilometers an hour by a man on top
of a flatcar that is moving backward at 20 kilometers an hour, then the
ball (to someone watching from the roadside) would seem to be not moving
at all after it left the hand of the thrower. It would just drop to the
ground.
There seemed no reason in the 19th century to suppose that light didn’t
behave in the same fashion. It was known to travel at the enormous speed
of just a trifle under 300,000 kilometers per second, while earth moved
in its orbit about the sun at a speed of about 30 kilometers per second.
Surely if a beam of light beginning at some earth-bound source shone in
the direction of earth’s travel, it ought to move at a speed of 300,030
kilometers per second. If it shone in the opposite direction, against
earth’s motion, it ought to move at a speed of 299,970 kilometers per
second.
Could such a small difference in an enormous speed be detected?
[Illustration: _Albert A. Michelson_]
The German-American physicist Albert Abraham Michelson (1852-1931) had
invented a delicate instrument, the interferometer, that could compare
the velocities of different beams of light with great precision. In 1887
he and a co-worker, the American chemist Edward Williams Morley
(1838-1923), tried to measure the comparative speeds of light, using
beams headed in different directions. Some of this work was performed at
the U. S. Naval Academy and some at the Case Institute.
The results of the Michelson-Morley experiment were unexpected. It
showed no difference in the measured speed of light. No matter what the
direction of the beam—whether it went in the direction of the earth’s
movement, or against it, or at any angle to it—the speed of light always
appeared to be exactly the same.
To explain this, the German-Swiss-American scientist Albert Einstein
(1879-1955) advanced his “special theory of relativity” in 1905.
According to Einstein’s view, speeds could not merely be added. A ball
thrown forward at 20 kilometers an hour by a man moving at 20 kilometers
an hour in the same direction would not seem to be going 40 kilometers
an hour to an observer at the roadside. It would seem to be going very
slightly less than 40 kilometers an hour; so slightly less that the
difference couldn’t be measured.
However, as speeds grew higher and higher, the discrepancy in the
addition grew greater and greater (according to a formula Einstein
derived) until, at velocities of tens of thousands of kilometers per
hour, that discrepancy could be easily measured. At the speed of light,
which Einstein showed was a limiting velocity that an observer would
never reach, the discrepancy became so great that the speed of the light
source, however great, added or subtracted zero to or from the speed of
light.
Accompanying this were all sorts of other effects. It could be shown by
Einstein’s reasoning that no object possessing mass could move faster
than the speed of light. What’s more, as an object moved faster and
faster, its length in the direction of motion (as measured by a
stationary observer) grew shorter and shorter, while its mass grew
greater and greater. At 260,000 kilometers per second, its length in the
direction of movement was only half what it was at rest, and its mass
was twice what it was. As the speed of light was approached, its length
would approach zero in the direction of motion, while its mass would
approach the infinite.
Could this really be so? Ordinary objects never moved so fast as to make
their lengths and masses show any measurable change. What about
subatomic particles, however, which moved at tens of thousands of
kilometers per second? The German physicist Alfred Heinrich Bucherer
(1863-1927) reported in 1908 that speeding electrons did gain in mass
just the amount predicted by Einstein’s theory. The increased mass with
energy has been confirmed with great precision in recent years.
Einstein’s special theory of relativity has met many experimental tests
exactly ever since and it is generally accepted by physicists today.
Einstein’s theory gave rise to something else as well. Einstein deduced
that mass was a form of energy. He worked out a relationship (the
“mass-energy equivalence”) that is expressed as follows:
_E_ = _mc_²
where _E_ represents energy, _m_ is mass, and _c_ is the speed of light.
If mass is measured in grams and the speed of light is measured in
centimeters per second, then the equation will yield the energy in a
unit called “ergs”. It turns out that 1 gram of mass is equal to
900,000,000,000,000,000,000 (900 billion billion) ergs of energy. The
erg is a very small unit of energy, but 900 billion billion of them
mount up.
The energy equivalent of 1 gram of mass (and remember that a gram, in
ordinary units, is only ¹/₂₈ of an ounce) would keep a 100-watt light
bulb burning for 35,000 years.
[Illustration: ENERGY CREATED compared to MATTER (OR MASS) DESTROYED]
It is this vast difference between the tiny quantity of mass and the
huge amount of energy to which it is equivalent that obscured the
relationship over the years. When a chemical reaction liberates energy,
the mass of the materials undergoing the reaction decreases slightly—but
_very_ slightly.
Suppose, for instance, a gallon of gasoline is burned. The gallon of
gasoline has a mass of 2800 grams and combines with about 10,000 grams
of oxygen to form carbon dioxide and water, yielding 1.35 million
billion ergs. That’s a lot of energy and it will drive an automobile for
some 25 to 30 kilometers. But by Einstein’s equation all that energy is
equivalent to only a little over a millionth of a gram. You start with
12,800 grams of reacting materials and you end with 12,800 grams minus a
millionth of a gram or so that was given off as energy.
No instrument known to the chemists of the 19th century could have
detected so tiny a loss of mass in such a large total. No wonder, then,
that from Lavoisier on, scientists thought that the law of conservation
of mass held exactly.
Radioactive changes gave off much more energy per atom than chemical
changes did, and the percentage loss in mass was correspondingly
greater. The loss of mass in radioactive changes was found to match the
production of energy in just the way Einstein predicted.
It was no longer quite accurate to talk about the conservation of mass
after 1905 (even though mass was just about conserved in ordinary
chemical reactions so that the law could continue to be used by chemists
without trouble). Instead, it is more proper to speak of the
conservation of energy, and to remember that mass was one form of energy
and a very concentrated form.
The mass-energy equivalence fully explained why the atom should contain
so great a store of energy. Indeed, the surprise was that radioactive
changes gave off as little energy as they did. When a uranium atom broke
down through a series of steps to a lead atom, it produced a million
times as much energy as that same atom would release if it were involved
in even the most violent of chemical changes. Nevertheless, that
enormous energy change in the radioactive breakdown represented only
about one-half of 1% of the total energy to which the mass of the
uranium atom was equivalent.
Once Rutherford worked out the nuclear theory of the atom, it became
clear from the mass-energy equivalence that the source of the energy of
radioactivity was likely to be in the atomic nucleus where almost all
the mass of the atom was to be found.
The attention of physicists therefore turned to the nucleus.


THE STRUCTURE OF THE NUCLEUS

The Proton
As early as 1886 Eugen Goldstein, who was working with cathode rays,
also studied rays that moved in the opposite direction. Since the
cathode rays (electrons) were negatively charged, rays moving in the
opposite direction would have to be positively charged. In 1907 J. J.
Thomson called them “positive rays”.
Once Rutherford worked out the nuclear structure of the atom, it seemed
clear that the positive rays were atomic nuclei from which a number of
electrons had been knocked away. These nuclei came in different sizes.
Were the nuclei single particles—a different one for every isotope of
every element? Or were they all built up out of numbers of still smaller
particles of a very limited number of varieties? Might it be that the
nuclei owed their positive electrical charge to the fact that they
contained particles just like the electron, but ones that carried a
positive charge rather than a negative one?
All attempts to discover this “positive electron” in the nuclei failed,
however. The smallest nucleus found was that produced by knocking the
single electron off a hydrogen atom in one way or another. This hydrogen
nucleus had a single positive charge, one that was exactly equal in size
to the negative charge on the electron. The hydrogen nucleus, however,
was much more massive than an electron. The hydrogen nucleus with its
single positive charge was approximately 1837 times as massive as the
electron with its single negative charge.
Was it possible to knock the positive charge loose from the mass of the
hydrogen nucleus? Nothing physicists did could manage to do that. In
1914 Rutherford decided the attempt should be given up. He suggested
that the hydrogen nucleus, for all its high mass, should be considered
the unit of positive electrical charge, just as the electron was the
unit of negative electrical charge. He called the hydrogen nucleus a
“proton” from the Greek word for “first” because it was the nucleus of
the first element.
[Illustration: _One proton balances 1837 electrons._]
Why the proton should be so much more massive than the electron is still
one of the unanswered mysteries of physics.

The Proton-Electron Theory
What about the nuclei of elements other than hydrogen?
All the other elements had nuclei more massive than that of hydrogen and
the natural first guess was that these were made up of some appropriate
number of protons closely packed together. The helium nucleus, which had
a mass four times as great as that of hydrogen, might be made up of 4
protons; the oxygen nucleus with a mass number of 16 might be made up of
16 protons and so on.
This guess, however, ran into immediate difficulties. A helium nucleus
might have a mass number of 4 but it had an electric charge of +2. If it
were made up of 4 protons, it ought to have an electric charge of +4. In
the same way, an oxygen nucleus made up of 16 protons ought to have a
charge of +16, but in actual fact it had one of +8.
Could it be that something was cancelling part of the positive electric
charge? The only thing that could do so would be a negative electric
charge[1] and these were to be found only on electrons as far as anyone
knew in 1914. It seemed reasonable, then, to suppose that a nucleus
would contain about half as many electrons in addition to the protons.
The electrons were so light, they wouldn’t affect the mass much, and
they would succeed in cancelling some of the positive charge.
Thus, according to this early theory, _now known to be incorrect_, the
helium nucleus contained not only 4 protons, but 2 electrons in
addition. The helium nucleus would then have a mass number of 4 and an
electric charge (atomic number) of 4 - 2, or 2. This was in accordance
with observation.
This “proton-electron theory” of nuclear structure accounted for
isotopes very nicely. While oxygen-16 had a nucleus made up of 16
protons and 8 electrons, oxygen-17 had one of 17 protons and 9
electrons, and oxygen-18 had one of 18 protons and 10 electrons. The
mass numbers were 16, 17, and 18, respectively, but the atomic number
was 16 - 8, 17 - 9, and 18 - 10, or 8 in each case.
Again, uranium-238 has a nucleus built up, according to this theory, of
238 protons and 146 electrons, while uranium-235 has one built up of 235
protons and 143 electrons. In these cases the atomic number is,
respectively, 238 - 146 and 235 - 143, or 92 in each case. The nucleus
of the 2 isotopes is, however, of different structure and it is not
surprising therefore that the radioactive properties of the
two—properties that involve the nucleus—should be different and that the
half-life of uranium-238 should be six times as long as that of
uranium-235.
The presence of electrons in the nucleus not only explained the
existence of isotopes, but seemed justified by two further
considerations.
First, it is well known that similar charges repel each other and that
the repulsion is stronger the closer together the similar charges are
forced. Dozens of positively charged particles squeezed into the tiny
volume of an atomic nucleus couldn’t possibly remain together for more
than a tiny fraction of a second. Electrical repulsion would send them
flying apart at once.
On the other hand, opposite charges attract, and a proton and an
electron would attract each other as strongly as 2 protons (or 2
electrons) would repel each other. It was thought possible that the
presence of electrons in a collection of protons might somehow limit the
repulsive force and stabilize the nucleus.
Second, there are radioactive decays in which beta particles are sent
flying out of the atom. From the energy involved they could come only
out of the nucleus. Since beta particles are electrons and since they
come from the nucleus, it seemed to follow that there must be electrons
within the nucleus to begin with.
The proton-electron theory of nuclear structure also seemed to account
neatly for many of the facts of radioactivity.
Why radioactivity at all, for instance? The more complex a nucleus is,
the more protons must be squeezed together and the harder, it would
seem, it must be to keep them together. _More and more electrons seemed
to be required._ Finally, when the total number of protons was 84 or
more, no amount of electrons seemed sufficient to stabilize the nucleus.
The manner of breakup fits the theory, too. Suppose a nucleus gives off
an alpha particle. The alpha particle is a helium nucleus made up, by
this theory, of 4 protons and 2 electrons. If a nucleus loses an alpha
particle, its mass number should decline by 4 and its atomic number by 4
- 2, or 2. And, indeed, when uranium-238 (atomic number 92) gives off an
alpha particle, it becomes thorium-234 (atomic number 90).
Suppose a beta particle is emitted. A beta particle is an electron and
if a nucleus loses an electron, its mass number is almost unchanged. (An
electron is so light that in comparison with the nucleus, we can ignore
its mass.) On the other hand, a unit negative charge is gone. One of the
protons in the nucleus, which had previously been masked by an electron,
is now unmasked. Its positive charge is added to the rest and the atomic
number goes up by one. Thus, thorium-234 (atomic number 90) gives up a
beta particle and becomes protactinium-234 (atomic number 91).
If a gamma ray is given off, that gamma ray has no charge and the
equivalent of very little mass. That means that neither the mass number
nor the atomic number of the nucleus is changed, although its energy
content is altered.
Even more elaborate changes can be taken into account. In the long run,
uranium-238, having gone through many changes, becomes lead-206. Those
changes include the emission of 8 alpha particles and 6 beta particles.
The 8 alpha particles involve a loss of 8 × 4, or 32 in mass number,
while the 6 beta particles contribute nothing in this respect. And,
indeed, the mass number of uranium-238 declines by 32 in reaching
lead-206. On the other hand the 8 alpha particles involve a decrease in
atomic number of 8 × 2, or 16, while the 6 beta particles involve an
increase in atomic number of 6 × 1, or 6. The total change is a decrease
of 16 - 6, or 10. And indeed, uranium (atomic number 92) changes to lead
(atomic number 82).
It is useful to go into such detail concerning the proton-electron
theory of nuclear structure and to describe how attractive it seemed.
The theory appeared solid and unshakable and, indeed, physicists used it
with considerable satisfaction for 15 years.
—And yet, as we shall see, it was wrong; and that should point a moral.
Even the best seeming of theories may be wrong in some details and
require an overhaul.

Protons in Nuclei
Let us, nevertheless, go on to describe some of the progress made in the
1920s in terms of the proton-electron theory that was then accepted.
Since a nucleus is made up of a whole number of protons, its mass ought
to be a whole number if the mass of a single proton is considered 1.
(The presence of electrons would add some mass but in order to simplify
matters, let us ignore that.)
When isotopes were first discovered this indeed seemed to be so.
However, Aston and his mass spectrometer kept measuring the mass of
different nuclei more and more closely during the 1920s and found that
they differed very slightly from whole numbers. Yet a fixed number of
protons turned out to have different masses if they were first
considered separately and then as part of a nucleus.
Using modern standards, the mass of a proton is 1.007825. Twelve
separate protons would have a total mass of twelve times that, or
12.0939. On the other hand, if the 12 protons are packed together into a
carbon-12 nucleus, the mass is 12 so that the mass of the individual
protons is 1.000000 apiece. What happens to this difference of 0.007825
between the proton in isolation and the proton as part of a carbon-12
nucleus?
According to Einstein’s special theory of relativity, the missing mass
would have to appear in the form of energy. If 12 hydrogen nuclei
(protons) plus 6 electrons are packed together to form a carbon nucleus,
a considerable quantity of energy would have to be given off.
In general, Aston found that as one went on to more and more complicated
nuclei, a larger fraction of the mass would have to appear as energy
(though not in a perfectly regular way) until it reached a maximum in
the neighborhood of iron.
Iron-56, the most common of the iron isotopes, has a mass number of
55.9349. Each of its 56 protons, therefore, has a mass of 0.9988.
For nuclei more complicated than those of iron, the protons in the
nucleus begin to grow more massive again. Uranium-238 nuclei, for
instance, have a mass of 238.0506, so that each of the 238 protons they
contain has a mass of 1.0002.
By 1927 Aston had made it clear that it is the middle elements in the
neighborhood of iron that are most closely and economically packed. If a
very massive nucleus is broken up into somewhat lighter nuclei, the
proton packing would be tighter and some mass would be converted into
energy. Similarly, if very light nuclei were joined together into
somewhat more massive nuclei, some mass would be converted into energy.
This demonstration that energy was released in any shift away from
either extreme of the list of atoms according to atomic number fits the
case of radioactivity, where very massive nuclei break down to somewhat
less massive ones.
Consider that uranium-238 gives up 8 alpha particles and 6 beta
particles to become lead-206. The uranium-238 nucleus has a mass of
238.0506; each alpha particle has one of 4.0026 for a total of 32.0208;
each beta particle has a mass of 0.00154 for a total of 0.00924; and the
lead-206 nucleus has one of 205.9745.
This means that the uranium-238 nucleus (mass: 238.0506) changes into 8
alpha particles, 6 beta particles, and a lead-206 nucleus (total mass:
238.0045). The starting mass is 0.0461 greater than the final mass and
it is this missing mass that has been converted into energy and is
responsible for the gamma rays and for the velocity with which alpha
particles and beta particles are discharged.

Nuclear Bombardment
Once scientists realized that there was energy which became available
when one kind of nucleus was changed into another, an important question
arose as to whether such a change could be brought about and regulated
by man and whether this might not be made the source of useful power of
a kind and amount undreamed of earlier.
Chemical energy was easy to initiate and control, since that involved
the shifts of electrons on the outskirts of the atoms. Raising the
temperature of a system, for instance, caused atoms to move more quickly
and smash against each other harder, and that in itself was sufficient
to force electrons to shift and to initiate a chemical reaction that
would not take place at lower temperatures.
To shift the protons within the nucleus (“nuclear reactions”) and make
nuclear energy available was a harder problem by far. The particles
involved were much more massive than electrons and correspondingly
harder to move. What’s more, they were buried deep within the atom. No
temperatures available to the physicists of the 1920s could force atoms
to smash together hard enough to reach and shake the nucleus.
In fact, the only objects that were known to reach the nucleus were
speeding subatomic particles. As early as 1906, for instance, Rutherford