Inchemistry, thewhole number rule states that the masses of the isotopes arewhole number multiples of the mass of thehydrogen atom.[1] The rule is a modified version ofProut's hypothesis proposed in 1815, to the effect thatatomic weights are multiples of the weight of the hydrogen atom.[2] It is also known as theAston whole number rule[3] afterFrancis W. Aston who was awarded the Nobel Prize in Chemistry in 1922 "for his discovery, by means of hismass spectrograph, of isotopes, in a large number of non-radioactive elements, and for his enunciation of the whole-number rule".[4]
Thelaw of definite proportions was formulated byJoseph Proust around 1800[5] and states that all samples of a chemical compound will have the same elemental composition by mass. Theatomic theory ofJohn Dalton expanded this concept and explained matter as consisting of discreteatoms with one kind of atom for each element combined in fixed proportions to form compounds.[6]
In 1815,William Prout reported on his observation that theatomic masses of the elements were whole multiples of the atomic mass ofhydrogen.[7][8] He then hypothesized that the hydrogen atom was the fundamental object and that the other elements were a combination of different numbers of hydrogen atoms.[9]
In 1920, Francis W. Aston demonstrated through the use of amass spectrometer that apparent deviations from Prout's hypothesis are predominantly due to the existence ofisotopes.[10] For example, Aston discovered that neon has two isotopes with masses very close to 20 and 22 as per the whole number rule, and proposed that the non-integer value 20.2 for the atomic weight of neon is due to the fact that natural neon is a mixture of about 90% neon-20 and 10% neon-22). A secondary cause of deviations is thebinding energy ormass defect of the individual isotopes.
During the 1920s, it was thought that the atomic nucleus was made of protons and electrons, which would account for the disparity between theatomic number of an atom and itsatomic mass.[11][12] In 1932,James Chadwick discovered an uncharged particle of approximately the mass as the proton, which he called theneutron.[13] The fact that the atomic nucleus is composed of protons and neutrons was rapidly accepted and Chadwick was awarded theNobel Prize in Physics in 1935 for his discovery.[14]
The modern form of the whole number rule is that theatomic mass of a given elementalisotope is approximately themass number (number of protons plus neutrons) times adalton (approximate mass of a proton, neutron, or hydrogen-1 atom). This rule predicts theatomic mass ofnuclides and isotopes with an error of at most 1%, with most of the error explained by the mass deficit caused bynuclear binding energy.
