Older (pre-1961) historical relative scales based on the atomic mass unit (symbol: a.m.u. Or amu) used either the oxygen-16 relative isotopic mass or else the oxygen relative atomic mass (i.e., atomic weight) for reference. Naturally occurring oxygen is composed of three stable isotopes, 16 O, 17 O, and 18 O, with 16 O being the most abundant (99.762% natural abundance). Depending on the terrestrial source, the standard atomic weight varies within the range of 15.999 03, 15.999 77 (the conventional value is 15.999).
Gauri Karkhanis answered this
In order to calculate relative atomic mass, first of all, is to calculate the 1/12 of carbon-12.. that is 1.661 x 10-27 Kg; then, compare this value with any other atom which needs to be calculated and the obtained ratio is relative atomic mass for that atom. For example, for the oxygen atom, its rest mass is 2.657 x 10-26, divide it by 1.661 x 10-27 (2.657 x 10-26 / 1.661 x 10-27 ) and the answer will approximately be 16 that is the relative atomic mass for oxygen. The contribution of this value is to make calculation much easier.
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History of the atomic mass unit
Stanislao Cannizzaro (1826–1910), the pioneer in this field, adopted the hydrogen atom as a standard of mass and set its atomic weight at 2. Others accepted the idea of using a specific atom as a standard of mass, but preferred a more massive standard in order to reduce experimental error.
As early as 1850, chemists used a unit of atomic weight based on saying the atomic weight of oxygen was 16. Oxygen was chosen because it forms chemical compounds with many other elements, simplifying determination of their atomic weights. Adobe flash player for chrome free download mac. Sixteen was chosen because it was the lowest whole number that could be assigned to oxygen and still have an atomic weight for hydrogen that was not less than 1.
The 0=16 scale was formalized when a committee appointed by the Deutsche Chemische Gesellschaft called for the formation of an international commission on atomic weights in March 1899. A commission of 57 members was formed. Since the commission carried on its business by correspondence, the size proved unwieldy, and the Gesellschaft suggested a smaller committee be elected. A 3-member International Committee of Atomic Weights was duly elected, and in 1903 issued its first report, using the 0=16 scale.5
Taking isotopes into account
The discovery of isotopes complicated the picture. In nature, pure oxygen is composed of a mixture of isotopes: some oxygen atoms are more massive than others.
This was no problem for the chemists’ calculations as long as the relative abundance of the isotopes in their reagents remained constant, though it confirmed that oxygen’s atomic weight was the only one that in principle would be a whole number (hydrogen’s, for example, was 1.000 8). Teleprompter for mac free download.
Physicists, however, dealing with atoms and not reagents, required a unit that distinguished between isotopes. At least as early as 19276 physicists were using an atomic mass unit defined as equal to one-sixteenth of the mass of the oxygen-16 atom (the isotope of oxygen containing a total of 16 protons and neutrons).
In 1919, isotopes of oxygen with mass 17 and 18 were discovered.7 Thus the two amu’s clearly diverged: one based on one-sixteenth of the average mass of the oxygen atoms in the chemist’s laboratory, and the other based on one-sixteenth of the mass of an atom of a particular isotope of oxygen.
Relative Formula Mass Of Oxygen
In 1956, Alfred Nier (at the bar in the Hotel Krasnapolski in Amsterdam) and independently A. Ölander8, both members of the Commission on Atomic Masses of the IUPAP, suggested to Josef Mattauch that the atomic weight scale be based on carbon-12. That would be okay with physicists, since carbon-12 was already used as a standard in mass spectroscopy. The chemists resisted making the amu one-sixteenth the mass of an oxygen-16 atom; it would change their atomic weights by about 275 parts per million. Making the amu one-twelfth the mass of a carbon-12 nucleus, however, would lead to only a 42 parts per million change, which seemed within reason.
Relative Atomic Mass Example
Mattauch set to work enthusiastically proselytizing the physicists, while E. Wichers lobbied the chemists.9 In the years 1959–1961 the chemists and physicists resolved to use the isotope carbon-12 as the standard, setting its atomic mass at 12.
Relative Atomic Mass Of Carbon