Chemical Sciences: A Manual for CSIR-UGC National Eligibility Test for Lectureship and JRF/Whole number rule
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The Whole Number Rule states that the masses of the elements are whole number multiples of the mass of the hydrogen atom.[1] The rule can be formulated from Prout's hypothesis put forth in 1815.[2] In 1920, Francis W. Aston demonstrated through the use of a mass spectrometer that apparent deviations from the rule are predominantly due to the existence of isotopes;[3] they are secondarily due to binding energy, as mass defect. The modern form of the whole number rule is that the atomic mass of a given isotope is approximately the mass number (number of protons plus neutrons) times an atomic mass unit (approximate mass of a proton, neutron, or hydrogen-1 atom).
References edit
- ↑ Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews. 25 (1): 146–57. doi:10.1002/mas.20061. PMID 16134128.
- ↑ Prout, William (1815). "On the relation between the specific gravities of bodies in their gaseous state and the weights of their atoms". Annals of Philosophy. 6: 321–330. Retrieved 2007-09-08.
{{cite journal}}: Cite has empty unknown parameter:|coauthors=(help) - ↑ Aston, Francis W. (1920). "The constitution of atmospheric neon". Philosophical Magazine. 39 (6): 449–455.
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Harkins WD (1925). "The Separation of Chlorine into Isotopes (Isotopic Elements) and the Whole Number Rule for Atomic Weights". Proc. Natl. Acad. Sci. U.S.A. 11 (10): 624–8. doi:10.1073/pnas.11.10.624. PMC 1086175. PMID 16587053.