History
The atomic weight scale has traditionally been a relative scale, that is without an explicit unit, with the first atomic weight basis suggested by John Dalton in 1803 as 1H. Despite the initial mass of 1H being used as the natural unit for atomic weight, it was suggested by Wilhelm Ostwald that atomic weights would be best expressed in terms in units of 1/16 weight of oxygen. This evaluation was made prior to the discovery of the existence of elemental isotopes, which occurred in 1912.
The discovery of isotopic oxygen in 1929 led to a divergence in atomic weight representation, with isotopically weighted oxygen (i.e., naturally occurring oxygen atomic weight) given a value of exactly 16 atomic mass units (amu) in chemistry, while pure 16O (oxygen-16) was given the mass value of exactly 16 amu in physics.
The divergence of these values could result in errors in computations, and was unwieldy. The chemistry amu, based on the atomic weight of natural oxygen (including the heavy naturally-occurring isotopes 17O and 18O), was about 1.000282 more massive than the physics amu, based on pure isotopic 16O.
For these and other reasons, the reference standard for both physics and chemistry was changed to carbon-12 in 1961. The choice of carbon-12 was made to minimise further divergence with prior literature. The new and current unit was referred to as the "unified atomic mass unit" u. and given a new symbol "u," which replaced the now deprecated "amu" that had been connected to the old oxygen-based system.
However, modern sources often still use "amu" in place of "u" (as a synonym) and define "amu" in terms of carbon-12. In general, "amu" probably does not refer to the old oxygen standard amu, unless the material is older than the 1960s.
The "unified atomic mass unit" u was defined as:
Read more about this topic: Atomic Mass Unit
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