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.