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                         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.

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.

[Illustration: _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.

                            [Illustration: ]

  Neutral atom
    Each unit of positive charge is balanced by a unit of negative
          charge
    In this case, total charge = +2 -2 = 0
  Ionized atom:
    If an electron is removed, the balance is destroyed
    In this case, total charge = +2 -1 = +1

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.

[Illustration: _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._]

[Illustration: ]

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.

[Illustration: _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 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.

[Illustration: _Charles Glover Barkla_]

[Illustration: _Henry Gwyn-Jeffreys Moseley_]

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.

[Illustration: _Dmitri Mendeléev and Bohuslav Brauner in Prague in 1900.
Brauner was a professor of chemistry at the Bohemian University in
Prague._]

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.

[Illustration: _Niels Bohr_]

[Illustration: _Bohr’s study._]

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.

  [Illustration: showing Helium atom, Hydrogen atom; Nucleus, Proton,
                      Neutron, Electron labelled]


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.

[Illustration: _Frederick Soddy_]

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 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.

[Illustration: _Isotopes of two elements._]

  Atomic number, 1
    Hydrogen-1: Mass number, 1; 1 Proton, 1 Electron
    Hydrogen-2: Mass number, 2; 1 Proton, 1 Neutron, 1 Electron
  Atomic number, 2
    Helium-3: Mass number, 3; 2 Protons, 1 Neutron, 2 Electrons
    Helium-4: Mass number, 4; 2 Protons, 2 Neutrons, 2 Electrons

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 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.

[Illustration: _Francis W. Aston_]

[Illustration: _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.

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.

[Illustration: _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 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”.

[Illustration: _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.

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.




                                 ENERGY


The Law of Conservation of Energy

We have now gone as far as we conveniently can in considering the
intertwining strands of the atom and of electricity. It is time to turn
to the third strand—energy.

To physicists the concept of “work” is that of exerting a force on a
body and making it move through some distance. To lift a weight against
the pull of gravity is work. To drive a nail into wood against the
friction of its fibers is work.

Anything capable of performing work is said to possess “energy” from
Greek words meaning “work within”. There are various forms of energy.
Any moving mass possesses energy by virtue of its motion. That is, a
moving hammer will drive a nail into wood, while the same hammer held
motionlessly against the nailhead will not do so. Heat is a form of
energy, since it will expand steam that will force wheels into motion
that can then do work. Electricity, magnetism, sound, and light can be
made to perform work and are forms of energy.

The forms of energy are so many and so various that scientists were
eager to find some rule that covered them all and would therefore serve
as a unifying bond. It did not seem impossible that such a rule might
exist, since one had been found in connection with matter that appeared
in even greater variety than energy did.

All matter, whatever its form and shape, possessed mass, and in the
1770s, the French chemist Antoine Laurent Lavoisier (1743-1794)
discovered that the quantity of mass was constant. If a system of matter
were isolated and made to undergo complicated chemical reactions,
everything about it might change, but not its mass. A solid might turn
into a gas; a single substance might change into two or three different
substances, but whatever happened, the total mass at the end was exactly
the same (as nearly as chemists could tell) as at the beginning. None
was either created or destroyed, however, the nature of the matter might
change. This was called the “law of conservation of mass”.

[Illustration: _Lavoisier in his laboratory during his studies on
respiration. From a sketch made by Madame Lavoisier._]

[Illustration: _Antoine Lavoisier and his wife._]

Naturally, it would occur to scientists to wonder if a similar law might
hold for energy. The answer wasn’t easy to get. It wasn’t as simple to
measure the quantity of energy as it was to measure the quantity of
mass. Nor was it as simple to pen up a quantity of energy and keep it
from escaping or from gaining additional quantity from outside, as it
was in the case of mass.

Beginning in 1840, however, the English physicist James Prescott Joule
(1818-1889) began a series of experiments in which he made use of every
form of energy he could think of. In each case he turned it into heat
and allowed the heat to raise the temperature of a given quantity of
water. He used the rise in temperature as a measure of the energy. By
1847 he was convinced that any form of energy could be turned into fixed
and predictable amounts of heat; that a certain amount of work was
equivalent to a certain amount of heat.

In that same year, the German physicist Hermann Ludwig Ferdinand von
Helmholtz (1821-1894) advanced the general notion that a fixed amount of
energy in one form was equal to the same amount of energy in any other
form. Energy might change its form over and over, but not change its
amount. None could either be destroyed or created. This is the “law of
conservation of energy”.


Chemical Energy

There is energy in a piece of wood. Left quietly to itself, it seems
completely incapable of bringing about any kind of work. Set it on fire,
however, and the wood plus the oxygen in the air will give off heat and
light that are clearly forms of energy. The heat could help boil water
and run a steam engine.

The amount of energy in burning wood could be measured if it were mixed
with air and allowed to burn in a closed container that was immersed in
a known quantity of water. From the rise in temperature of the water,
the quantity of energy produced could be measured in units called
“calories” (from a Latin word for “heat”). The instrument was therefore
called a “calorimeter”.

In the 1860s the French chemist Pierre Eugène Marcelin Berthelot
(1827-1907) carried through hundreds of such determinations. His work
and similar work by others made it clear that such “chemical energy”—the
energy derived from chemical changes in matter—fit the law of
conservation of energy.

Here’s how it looked in the last decades of the 19th century.

Molecules are composed of combinations of atoms. Within the molecules,
the atoms stick together more or less tightly. It takes a certain amount
of energy to pull a molecule apart into separate atoms against the
resistance of the forces holding them together.

If, after being pulled apart, the atoms are allowed to come together
again, they give off energy. The amount of energy they give off in
coming together is exactly equal to the amount of energy they had to
gain before they could separate.

This is true of all substances. For instance, hydrogen gas, as it is
found on earth, is made up of molecules containing 2 hydrogen atoms each
(H₂). Add a certain amount of energy and you pull the atoms apart; allow
the atoms to come back together into paired molecules, and the added
energy is given back again. The same is true for the oxygen molecule,
which is made up of 2 oxygen atoms (O₂) and of the water molecule (H₂O).
Always the amount of energy absorbed in one change is given off in the
opposite change. The amount absorbed and the amount given off are always
exactly equal.

However, the amount of energy involved differs from molecule to
molecule. It is quite hard to pull hydrogen molecules apart, and it is
even harder to pull oxygen molecules apart. You have to supply about 12%
more energy to pull an oxygen molecule apart than to pull a hydrogen
molecule apart. Naturally, if you let 2 oxygen atoms come together to
form an oxygen molecule, you get back 12% more energy than if you allow
2 hydrogen atoms to come together to form a hydrogen molecule.

It takes a considerably larger amount of energy to pull apart a water
molecule into separate atoms than to pull apart either hydrogen or
oxygen molecules. Naturally, that greater energy is also returned once
the hydrogen and oxygen atoms are allowed to come back together into
water molecules.

Next, imagine pulling apart hydrogen and oxygen molecules into hydrogen
and oxygen atoms and then having those atoms come together to form
_water_ molecules. A certain amount of energy is put into the system to
break up the hydrogen and oxygen molecules, but then a much greater
amount of energy is given off when the water molecules form.

It is for that reason that a great deal of energy (mostly in the form of
heat) is given off if a jet of hydrogen gas and a jet of oxygen gas are
allowed to mix in such a way as to form water.

Just mixing the hydrogen and oxygen isn’t enough. The molecules of
hydrogen and oxygen must be separated and that takes a little energy.
The energy in a match flame is enough to raise the temperature of the
mixture and to make the hydrogen and oxygen molecules move about more
rapidly and more energetically. This increases the chance that some
molecules will be broken up into separate atoms (though the actual
process is rather complicated). An oxygen atom might then strike a
hydrogen molecule to form water (O + H₂ → H₂O) and more energy is given
off than was absorbed from the match flame. The temperature goes up
still higher so that further breakup among the oxygen and hydrogen
molecules is encouraged.

[Illustration: _The formation of a sodium chloride molecule._]

This happens over and over again so that in very little time, the
temperature is very high and the hydrogen and oxygen are combining to
form water at an enormous rate. If a great deal of hydrogen and oxygen
are well-mixed to begin with, the rate of reaction is so great that an
explosion occurs.

Such a situation, in which each reacting bit of the system adds energy
to the system by its reaction and brings about more reactions like
itself, is called a “chain reaction”. Thus, a match flame put to one
corner of a large sheet of paper will set that corner burning. The heat
of the burning will ignite a neighboring portion of the sheet and so on
till the entire sheet is burned. For that matter a single smoldering
cigarette end can serve to burn down an entire forest in a vastly
destructive chain reaction.


Electrons and Energy

The discovery of the structure of the atom sharpened the understanding
of chemical energy.

In 1904 the German chemist Richard Abegg (1869-1910) first suggested
that atoms were held together through the transfer of electrons from one
atom to another.

To see how this worked, one began by noting that electrons in an atom
existed in a series of shells. The innermost shell could hold only 2
electrons, the next 8, the next 18 and so on. It turned out that some
electron arrangements were more stable than others. If only the
innermost shell contained electrons and it were filled with the 2
electrons that were all it could hold, then that was a stable
arrangement. If an atom contained electrons in more than one shell and
the outermost shell that held electrons held 8, that was a stable
arrangement, too.

Thus, the helium atom has 2 electrons only, filling the innermost shell,
and that is so stable an arrangement that helium undergoes no chemical
reactions at all. The neon atom has 10 electrons—2 in the innermost
shell, and 8 in the next—and it does not react. The argon atom has 18
electrons—2, 8, and 8—and it too is very stable.

But what if an atom did not have its electron shell so neatly filled.
The sodium atom has 11 electrons—2, 8, and 1—while the fluorine atom has
9 electrons—2 and 7. If the sodium atom passed one of its electrons to a
fluorine atom, both would have the stable configuration of neon—2 and 8.
This, therefore, ought to have a great tendency to happen.

If it did happen, though, the sodium atom, minus 1 electron, would have
a unit positive charge and would be Na⁺, a positively charged ion.
Fluorine with 1 electron in excess would become F⁻, a negatively charged
ion. The 2 ions, with opposite charges, would cling together, since
opposite charges attract, and thus the molecule of sodium fluoride (NaF)
would be formed.

In 1916 the American chemist Gilbert Newton Lewis (1875-1946) carried
this notion farther. Atoms could cling together not only as a result of
the outright transfer of 1 or more electrons, but through sharing pairs
of electrons. This sharing could only take place if the atoms remained
close neighbors, and it would take energy to pull them apart and break
up the shared pool, just as it would take energy to pull 2 ions apart
against the attraction of opposite charges.

In this way the vague notions of atoms clinging together in molecules
and being forced apart gave way to a much more precise picture of
electrons being transferred or shared. The electron shifts could be
dealt with mathematically by a system that came to be called “quantum
mechanics” and chemistry was thus made a more exact science than it had
ever been before.


The Energy of the Sun

The most serious problem raised by the law of conservation of energy
involved the sun. Until 1847, scientists did not question sunlight. The
sun radiated vast quantities of energy but that apparently was its
nature and was no more to be puzzled over than the fact that the earth
rotated on its axis.

Once Helmholtz had stated that energy could neither be created nor
destroyed, however, he was bound to ask where the sun’s energy came
from. It had, to man’s best knowledge, been radiating heat and light,
with no perceptible change, throughout the history of civilization and,
from what biologists and geologists could deduce, for countless ages
earlier. Where, then, did that energy come from?

The sun gave the appearance of being a huge globe of fire. Could it
actually be that—a large heap of burning fuel, turning chemical energy
into heat and light?

The sun’s mass was known and its rate of energy production was known.
Suppose the sun’s mass were a mixture of hydrogen and oxygen and it were
burning at a rate sufficient to produce the energy at the rate it was
giving it off. If that were so, all the hydrogen and oxygen in its mass
would be consumed in 1500 years. No chemical reaction in the sun could
account for its having given us heat and light since the days of the
pyramids, let alone since the days of the dinosaurs.

Was there some source of energy greater than chemical energy? What about
the energy of motion? Helmholtz suggested that meteors might be falling
into the sun at a steady rate. The energy of their collisions might then
be converted into heat and light and this could keep the sun shining for
as long as the supply of meteors held out—even millions of years.

This, however, would mean that the sun’s mass would be increasing
steadily, and so would the force of its gravitational pull. With the
sun’s gravitational field increasing steadily, the length of earth’s
year would be decreasing at a measurable rate—but it wasn’t.

In 1854 Helmholtz came up with something better. He suggested that the
sun was contracting. Its outermost layers were falling inward, and the
energy of this fall was converted into heat and light. What’s more, this
energy would be obtained without any change in the mass of the sun
whatever.

Helmholtz calculated that the sun’s contraction over the 6000 years of
recorded history would have reduced its diameter only 560 miles—a change
that would not have been noticeable to the unaided eye. Since the
development of the telescope, two and a half centuries earlier, the
decrease in diameter would have been only 23 miles and that was not
measurable by the best techniques of Helmholtz’s day.

Working backward, however, it seemed that 25 million years ago, the sun
must have been so large as to fill the earth’s orbit. Clearly the earth
could not then have existed. In that case, the maximum age of the earth
was only 25 million years.

Geologists and biologists found themselves disturbed by this. The slow
changes in the earth’s crust and in the evolution of life made it seem
very likely that the earth must have been in existence—with the sun
delivering heat and light very much in the present fashion—for many
hundreds of millions of years.

Yet there seemed absolutely no other way of accounting for the sun’s
energy supply. Either the law of conservation of energy was wrong (which
seemed unlikely), or the painfully collected evidence of geologists and
biologists was wrong (which seemed unlikely),—or there was some source
of energy greater than any known in the 19th century, whose existence
had somehow escaped mankind (which also seemed unlikely).

Yet one of those unlikely alternatives would have to be true. And then
in 1896 came the discovery of radioactivity.


The Energy of Radioactivity

It eventually became clear that radioactivity involved the giving off of
energy. Uranium emitted gamma rays that we now know to be a hundred
thousand times as energetic as ordinary light rays. What’s more, alpha
particles were being emitted at velocities of perhaps 30,000 kilometers
per second, while the lighter beta particles were being shot off at
velocities of up to 250,000 kilometers per second (about 0.8 times the
velocity of light).

At first, the total energy given off by radioactive substances seemed so
small that there was no use worrying about it. The amount of energy
liberated by a gram of uranium in 1 second of radioactivity was an
insignificant fraction of the energy released by a burning candle.

In a few years, however, something became apparent. A lump of uranium
might give off very little energy in a second, but it kept on for second
after second, day after day, month after month, and year after year with
no perceptible decrease. The energy released by the uranium over a very
long time grew to be enormous. It eventually turned out that while the
rate at which uranium delivered energy did decline, it did so with such
unbelievable slowness that it took 4.5 billion years (!) for that rate
to decrease to half what it was to begin with.

If _all_ the energy delivered by a gram of uranium in the course of its
radioactivity over many billions of years was totalled, it was
enormously greater than the energy produced by the burning of a candle
with a mass equal to that of uranium.

Let’s put it another way. We might think of a single uranium atom
breaking down and shooting off an alpha particle. We might also think of
a single carbon atom combining with 2 oxygen atoms to form carbon
dioxide. The uranium atom would give off 2,000,000 times as much energy
in breaking down, as the carbon atom would in combining.

The energy of radioactivity is millions of times as intense as the
energy released by chemical reactions. The reason mankind had remained
unaware of radioactivity and very aware of chemical reactions was,
first, that the most common radioactive processes are so slow that their
great energies were stretched over such enormous blocks of time as to be
insignificant on a per second basis.

Secondly, chemical reactions are easily controlled by changing
quantities, concentrations, temperatures, pressures, states of mixtures,
and so on, and this makes them easy to take note of and to study. The
rate of radioactive changes, however, could not apparently be altered.
The early investigators quickly found that the breakdown of uranium-238,
for instance, could not be hastened by heat, pressure, changes in
chemical combination, or, indeed, anything else they could think of. It
remained incredibly slow.

But despite all this, radioactivity had at last been discovered and the
intensity of its energies was recognized and pointed out in 1902 by
Marie Curie and her husband Pierre Curie (1859-1906).

Where, then, did the energy come from? Could it come from the outside?
Could the radioactive atoms somehow collect energy from their
surroundings, concentrate it several million-fold, and then let it out
all at once?

To concentrate energy in this fashion would violate something called
“the second law of thermodynamics”. This was first proposed in 1850 by
the German physicist Rudolf Julius Emmanuel Clausius (1822-1888) and had
proved so useful that physicists did not like to abandon it unless they
absolutely had to.

Another possibility was that radioactive atoms were creating energy out
of nothing. This, of course, violated the law of conservation of energy
(also called “the first law of thermodynamics”) and physicists preferred
not to do that either.

The only thing that seemed to remain was to suppose that somewhere
within the atom was a source of energy that had never made itself
evident to humanity until the discovery of radioactivity. Becquerel was
one of the first to suggest this.

It might have seemed at first that only radioactive elements had this
supply of energy somewhere within the atom, but in 1903 Rutherford
suggested that all atoms had a vast energy supply hidden within
themselves. The supply in uranium and thorium leaked slightly, so to
speak, and that was all that made them different.

[Illustration: _The room in which the Curies discovered radium. Pierre
Curie’s writing is on the blackboard._]

But if a vast supply of energy existed in atoms, it was possible that
the solution to the puzzle of the sun’s energy might rest there. As
early as 1899 the American geologist Thomas Chrowder Chamberlin
(1843-1928) was already speculating about a possible connection between
radioactivity and the sun’s energy.

If it were some variety of this newly discovered source of energy (not
necessarily ordinary radioactivity, of course) that powered the
sun—millions of times as intense as chemical energy—then the sun might
be pouring out energy for hundreds of millions of years without
perceptible physical change—just as uranium would show scarcely any
change even in so mighty a time span. The sun would not have to be
contracting; it would not have had to fill the earth’s orbit 25,000,000
years ago.

This was all exciting, but in 1900 the structure of the atom had not yet
been worked out and this new energy was just a vague supposition. No one
had any idea of what it actually might be or where in the atom it might
be located. It could only be spoken of as existing “within the atom” and
was therefore called “atomic energy”. Through long habit, it is still
called that much of the time. And yet “atomic energy” is not a good
name. In the first couple of decades of the 20th century, it became
apparent that ordinary chemical energy involved electron shifts and
those electrons were certainly components of atoms. This meant that a
wood fire was a kind of atomic energy.

The electrons, however, existed only in the outer regions of the atom.
Once Rutherford worked out the theory of the nuclear atom, it became
apparent that the energy involved in radioactivity and in solar
radiation had to involve components of the atom that were more massive
and more energetic than the light electrons. The energy had to come,
somehow, from the atomic nucleus.

What is involved then in radioactivity and in the sun is “nuclear
energy”. That is the proper name for it and in the next section we will
consider the subsequent history of the nuclear energy that broke upon
the startled consciousness of scientists as the 20th century opened and
which, less than half a century later, was to face mankind with untold
consequences for good and for evil.




                               FOOTNOTES


[1]“Mass” is the correct term, but “weight”, which is a somewhat
    different thing, is so commonly used instead that in this book I
    won’t try to make any distinction.




                            QUOTATION CREDIT


  Inside front cover    Copyright © by Abelard-Shuman, Ltd., New York.
                        Reprinted by permission from _Inside the Atom_,
                        Isaac Asimov, 1966.




                             PHOTO CREDITS


  Cover                 The Metropolitan Museum of Art
  Page facing inside    The “Horsehead” Nebula in Orion. Hale
  cover                 Observatories.
  Author’s Photo        Jay K. Klein
  Contents page &       Lick Observatory
  page 4
  Page
  7                     New York Public Library
  9                     From _Discovery of the Elements_, Mary E. Weeks,
                        Chemical Education Publishing Company, 1968.
  12                    Library of Congress
  15                    Sir George Thomson
  18                    Burndy Library
  19                    New York Public Library
  21                    Copyright © 1965 by Barbara Lovett Cline,
                        reprinted from her volume _The Questioners:
                        Physicists and the Quantum Theory_ by permission
                        of Thomas Y. Crowell Company, Inc., New York.
  22 & 23               Curie Foundation, Institute of Radium
  26                    Academic Press, Inc.
  29                    Van Nostrand Reinhold Company
  31                    Top, Nobel Institute; bottom, from _Discovery of
                        the Elements_, Mary E. Weeks, Chemical Education
                        Publishing Company, 1968.
  32                    From _Discovery of the Elements_, Mary E. Weeks,
                        Chemical Education Publishing Company, 1968.
  34                    Top, Nobel Institute; bottom, Niels Bohr
                        Institute.
  36, 42, 44, & 45      Nobel Institute
  48                    Academic Press, Inc.
  49                    From _Discovery of the Elements_, Mary E. Weeks,
                        Chemical Education Publishing Company, 1968.
  60                    Curie Foundation, Institute of Radium

                      ★ U.S. GOVERNMENT PRINTING OFFICE: 1975—640—285/13


A word about ERDA....

The mission of the U. S. Energy Research & Development Administration
(ERDA) is to develop all energy sources, to make the Nation basically
self-sufficient in energy, and to protect public health and welfare and
the environment. ERDA programs are divided into six major categories:

· CONSERVATION OF ENERGY—More efficient use of existing energy sources,
development of alternate fuels and engines for automobiles to reduce
dependence on petroleum, and elimination of wasteful habits of energy
consumption.

· FOSSIL ENERGY—Expansion of coal production and the development of
technologies for converting coal to synthetic gas and liquid fuels,
improvement of oil drilling methods and of techniques for converting
shale deposits to usable oil.

· SOLAR, GEOTHERMAL, AND ADVANCED ENERGY SYSTEMS—Research on solar
energy to heat, cool, and eventually electrify buildings, on conversion
of underground heat sources to gas and electricity, and on fusion
reactors for the generation of electricity.

· ENVIRONMENT AND SAFETY—Investigation of health, safety, and
environmental effects of the development of energy technologies, and
research on management of wastes from energy production.

· NUCLEAR ENERGY—Expanding medical, industrial and research applications
and upgrading reactor technologies for the generation of electricity,
particularly using the breeder concept.

· NATIONAL SECURITY—Production and administration of nuclear materials
serving both civilian and military needs.

ERDA programs are carried out by contract and cooperation with industry,
university communities, and other government agencies. For more
information, write to USERDA—Technical Information Center, P. O. Box 62,
Oak Ridge, Tennessee 37830.

    [Illustration: ENERGY RESEARCH & DEVELOPMENT ADMINISTRATION USA]

                             United States
             Energy Research and Development Administration
                        Office of Public Affairs
                         Washington, D.C. 20545




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