Summary
Francis William Aston’s *Isotopes* (1922) presents the first comprehensive synthesis of his discovery that many chemical elements exist as mixtures of atoms with identical chemical properties but different atomic masses. The central thesis is that isotopes—a term Aston coined—are distinct nuclear species occupying the same place in the periodic table, and that their masses are whole-number multiples of the hydrogen atom’s mass (the whole-number rule). Aston describes his invention of the mass spectrograph, which allowed him to separate and measure isotopes of over 50 elements, and provides tables of isotopic abundances and atomic weights. The book also discusses the implications for atomic theory, including the packing fraction and the stability of nuclei. A reader takes away a foundational understanding of isotopic composition as a universal property of matter, the experimental proof that atomic weights are weighted averages, and the early evidence for nuclear binding energy.
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Key concepts
- Whole-number rule — The observation that isotopic masses are close to integer multiples of the hydrogen atom’s mass, supporting the idea that nuclei are composed of protons.
- Mass spectrograph — Aston’s instrument that uses magnetic and electric fields to deflect ionized atoms, separating isotopes by their mass-to-charge ratio for precise measurement.
- Packing fraction — The fractional difference between an isotope’s actual mass and its nearest whole number, indicating nuclear binding energy and stability.
- Isotopic abundance — The relative proportion of each isotope of an element in a natural sample, which determines the element’s average atomic weight.
- Positive ray analysis — The technique of analyzing beams of positively charged ions to identify isotopes, which Aston refined from J.J. Thomson’s earlier work.
- Non-radioactive isotopes — Aston’s demonstration that stable, non-radioactive elements (e.g., neon, chlorine) also possess isotopes, extending the concept beyond radioactive decay series.