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

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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
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