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					Isotopes
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 0C.

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.

The following table shows some of the more common isotopes of a few elements:

                      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


Separation of isotopes
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!




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Isotopes and relative atomic masses
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.



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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?




Uses of radioisotopes
Radioactive isotopes can be very useful. They are used in:

1. Medicine for both treatment and diagnosis
2. Archaeological and geological dating using carbon 14 or uranium
3. Fluid flow measurement - water, blood, mud, sewage etc.
4. Thickness testing of materials such as polythene
5. Radiographs of metal castings
6. Sterilisation of food and insects
7. Tracers in fertilisers used in agriculture
8. Smoke alarms in houses
9. Tracing phosphate fertilisers using phosphorus 32
10. Checking the silver content of coins
11. Atomic lights using krypton 85
12. Testing for leaks in pipes




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