FRANCIS ASTON (1877-1945)

1922 – England

‘At the end of the First War, the assistant of JJ THOMSON developed the mass spectrograph for measuring the comparative weights of atoms’

Early Mass Spectrometer

Early Mass Spectrometer

Photo portrait of FRANCIS ASTON ©

FRANCIS ASTON

Whereas Thomson had used a discharge tube to measure the deflection of atomic particles passing through a hole in the anode, Aston refined the instrument by placing photographic plates in the path of the beams emerging through a hole in the cathode. These rays proved to be much harder to deflect from their course, implying they were made of particles thousands of times heavier than electrons, with masses close to those of atoms. These particles were deflected in opposite directions to negative cathode rays, indicating that they carried a positive charge.

Hence hurtling in one direction down a discharge tube were cathode ray electrons occasionally colliding with the atoms of the rarefied gas filling the tube. Drifting in the other direction – much more sluggishly because of their larger mass – were positive gas atoms, or ‘ions’, stripped of an electron or two in the collisions.

Once perfected, this mass spectrograph offered a means of deciding the mass of these atoms to an accuracy of 1 part in 100,000. This was enough to distinguish the existence of different isotopes and to confirm that the ‘rule of thumb’ – that masses of atoms were roughly whole number multiples of the mass of hydrogen – was in actuality accurate.
What it had confirmed was that the fundamental building block had the same mass as the proton, or hydrogen nucleus. When the mass spectrograph was first devised, the proton was the only particle with the mass of a proton, as the neutron was yet to be described by JAMES CHADWICK.

When comparisons of atomic mass were made, the oxygen atom was chosen as the standard with a mass of 16.
Today carbon is used as the atomic mass standard with a weight of 12.

Using this standard it was discovered that although the ratios of atomic masses were indistinguishable from whole numbers, helium being 4, oxygen 16, the atomic mass of hydrogen was anomalously high, being 1.008. The conclusion as to why this should be so had been suggested in the nineteenth century but before Einstein had found little support. After the famous paper of 1905, however, it was not unreasonable to suggest that when hydrogen atoms came together or coalesced to form other elements, mass was lost as energy.

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picture of the Nobel medal - link to nobelprize.org

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MARCUS OLIPHANT (1901-2000)

1934 – UK

‘Hydrogen has three isotopes: hydrogen-1 (ordinary hydrogen: one proton), hydrogen-2 (deuterium: one proton, one neutron) and hydrogen-3 (tritium: one proton, two neutrons)’

They each have one single proton (z = 1), but differ in the number of their neutrons. Hydrogen has no neutron, deuterium has one, and tritium has two neutrons. The isotopes of hydrogen have, respectively, mass numbers of one, two, and three. Their nuclear symbols are therefore 1H, 2H, and 3H. The atoms of these isotopes have one electron to balance the charge of the one proton. Since chemistry depends on the interactions of protons with electrons, the chemical properties of the isotopes are nearly the same.

MARK OLIPHANT

The lightest rare gas, helium, exists in nature in two forms – two isotopes

The usual form is represented as 4He, where the figure 4 stands for the number of nucleons in the atomic nucleus (two protons and two neutrons). In the unusual form, 3He, the atomic nucleus has only one neutron, so it is lighter. In helium that occurs naturally the heavier isotope is more frequent than the lighter one by a factor of about 10 million. That is why it is only in the last 50 years that it has been possible to produce large amounts of 3He, at nuclear power stations, for example. At normal temperatures the gases of the two isotopes differ only in their atomic weights.

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