Isotopes are particles which have the same position in the
Periodic Table, that is, they are atoms of the same chemical element but their nucleon
numbers are different. Isotopes of an element have nuclei with the same number of protons
but different numbers of neutrons. Neon, for instance, has three isotopes with nucleon
numbers of 20, 21 and 22, corresponding respectively to 10, 11 and 12 neutrons in the
nucleus. The most common isotopes of uranium are uranium-235 and uranium-238 (143 and
146 neutrons respectively.)
It is important to realise that since the number of
electrons is identical for all isotopes of the same element, the chemical properties of isotopes
of the same element are identical. Since the structure of the nuclei are different, however,
their nuclear properties will be different and, since their relative atomic masses are different,
some of their physical properties are different as well. For example, the boiling point of
'heavy water' (water containing the isotope of hydrogen with a neutron in the nucleus) is 104
oC.
In 1906 isotopes were discovered in radioactive elements (although
their nature was not understood) and in 1912 Thomson discovered the three isotopes of
neon with the nucleon numbers shown above.
Two isotopes of carbon are shown in the following diagram
| Element | Nucleon numbers of isotopes |
| Hydrogen | 1,2,3 |
| Helium | 3,4 |
| Carbon | 12,14 |
| Oxygen | 16,17,18 |
| Neon | 20,21,22 |
| Calcium | 40,42,44 |
| Iron | 56,57 |
| Mercury | 198,199,200,201,202 |
| Lead | 206,207,208 |
| Uranium | 235,238 |
There are
several methods for separating isotopes.
(a) Centrifuge
Due to the difference in
the masses of the two isotopes of uranium, a centrifuge method can be used to separate
them. The mass difference is three neutron masses and this is sufficient to make this method
effective.
(b) Gaseous diffusion
This method is used when the difference in
mass is small, for example one neutron mass as in the case of hydrogen (1p) and deuterium
(1p, 1n). It is also used to enrich uranium.
(c) Electromagnetic/electrostatic
deflection
Where a very high purity is required the sample may be built up particle by
particle, by deflection and collection in a mass spectrometer.
An ion current of 0.1 mA
gives 6.21x1014 particles per second, assuming that the ions are singly charged. To produce
10-4 mole of the sample by this method would take 1.33
days!
The
following in an extract from an article by Dr F.W.Aston first published in Nature in
1920.
In the atomic theory put forward by John Dalton in 1801 the second postulate was
'Atoms of the same element are similar to one another and equal in weight'. For more than a
century this was regarded by chemists and physicists alike as an article of scientific faith. The
only item among the immense quantities of knowledge acquired during that productive period
which offered the faintest suggestion against its validity was the inexplicable mixture of order
and disorder among the elementary atomic weights [relative atomic masses].The general
state of opinion at the end of the last century may be gathered from the two following
quotations from Sir William Ramsay's -address to the British Association at Toronto in 1897:
'There have been almost innumerable attempts to reduce the differences between
atomic weights to regularity by contriving some formula which will express the numbers
which represent the atomic weights with all their irregularities. Needless to say such attempts
have in no case been successful. The idea has been advanced that what we call the atomic
weight is a mean; that when we say the atomic weight of oxygen is 16 we merely state that
the average atomic weight of oxygen is 16; and it is not inconceivable that a certain number
of oxygen molecules have a weight somewhat higher than 32 and a certain number have a
lower weight.'
This idea was placed on an altogether different footing some ten
years later by the work of Lord Rutherford and his colleagues on radioactive transformations.
The results of these led inevitably to the conclusion that there must exist elements which
have chemical properties identical for all practical purposes, hut the atoms have different
weights. This conclusion has been recently confirmed in a most convincing manner by the
production in quantity of specimens of lead from radioactive and other sources, which,
although perfectly pure and chemically indistinguishable, give atomic weights differing by
amounts quite outside the possible experimental error. Elements differing in mass but
chemically identical have been called isotopes by Professor Soddy.
The work of Sir
J.J.Thomson before the war led to the belief that neon also existed as a mixture of two
isotopes with atomic weights of 20 and 22, the accepted atomic weight being 20.2. The
methods available were not accurate enough to distinguish between 20 and 20.2 with
certainty but in 1913 a diffusion experiment gave positive results, an apparent change in
density of 0.7 per cent between the lightest and heaviest fractions being obtained after many
thousands of operations.
By the time work was started again after the war the
isotope theory had been generally accepted so far as the radioactive elements were
concerned and a good deal of theoretical speculation had been made as to its applicability to
the elements generally. As separation by diffusion is at best extremely slow and laborious
attention was again turned to positive rays in the hope of increasing the accuracy of
measurements to the required degree.
[A description of the Aston mass
spectrograph then follows. I will continue the extract at the point where the results are
considered.]
By far the most important result obtained from this work is the
generalisation that, with the exception of hydrogen, all atomic weights so far measured are
exactly whole numbers on the scale 0 = 16. Hydrogen is found to be 1.008, which agrees
with the value accepted by the chemists. This exception from the whole number rule is not
unexpected, as on the Rutherford 'nucleus' theory the hydrogen atom is the only one not
containing any negative electricity in its nucleus.
The results which have been so far
obtained with eighteen elements make it possible that the higher the atomic weight of an
element, the more complex it is likely to be, and that there are more complex elements than
simple. It must be noticed that, though the whole number rule asserts that a pure element
must have a whole number atomic weight, there is no reason to suppose that all elements
having atomic weights approximating to integers are therefore pure.'
1. Why was the
existence of isotopes initially thought to be unlikely?
2. Write an account of
Rutherford's work on radioactive decay.
3. Assuming that neon occurs as two
isotopes of atomic weights 20 and 22, what proportion of naturally occuring neon is neon-20
if the atomic weight of neon is 20.2? (Ignore other isotopes)
4. What is the modern
unit for 'atomic weights', and why was it chosen?
5. Why did hydrogen appear to be
an exception to the whole number rule?