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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AMEDEO AVOGADRO (1776-1856)

1811 – Italy

‘Equal volumes of all gases at the same temperature and pressure contain the same number of molecules’

In 1811, when Avogadro proposed his HYPOTHESIS, very little was known about atoms and molecules. Avogadro claimed that the same volume of any gas under identical conditions would always contain the same number of fundamental particles, or molecules. A litre of hydrogen would contain exactly the same number of molecules as a litre of oxygen or a litre of carbon dioxide.

Drawing of AVOGADRO ©

In 1814 ANDRE AMPERE was credited with discovering that if a gas consisted of a single element, its atoms could clump in pairs. The molecules of oxygen consisted of pairs of oxygen atoms, and the molecules of chlorine, pairs of chlorine atoms.
Diatomic gases possess a total of six degrees of simple freedom per molecule that are related to atomic motion.

This provides a way of comparing the weights of different molecules. It was only necessary to weigh equal volumes of different gases and compare them. This would be exactly the same as comparing the weights of the individual molecules of each gas.

Avogadro realised that GAY-LUSSAC‘s law provided a way of proving that an atom and a molecule are not the same. He suggested that the particles (molecules) of which nitrogen gas is composed consist of two atoms, thus the molecule of nitrogen is N2. When one volume (one molecule) of nitrogen combines with three volumes (three molecules) of hydrogen, two volumes (two molecules) of ammonia, NH3, are produced.

N2 + 3H2 ↔ 2NH3

However, the idea of a molecule consisting of two or more atoms bound together was not understood at that time.

Avogadro’s law was forgotten until 1860 when the Italian chemist STANISLAO CANNIZZARO (1826-1910) explained the necessity of distinguishing between atoms and molecules.

Avogadro’s constant
From Avogadro’s law it can be deduced that the same number of molecules of all gases at the same temperature and pressure should have the same volume. This number has been determined experimentally: it’s value is 6.022 1367(36) × 1023AVOGADRO’S NUMBERAvogadro's_number_in_e_notation

That at the same temperature and pressure, equal volumes of all gases have the same number of molecules allows a simple calculation for the combining ratios of all gases – by measuring their percentages by volume in any compound. This in turn facilitates simple calculation of the relative atomic masses of the elements of which it is composed.

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