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So we say that 1 Amu (u) = 1/12 of the mass of a Carbon 12 atom.

Im not sure about this, but I believe that first we measured the mass of 1 carbon 12 atom (This includes the mass defects that occur during the formation of nucleus) and then defined that 1 atomic mass unit has a value of 1/12 of the carbon 12 atom mass... Therefore one Carbon 12 atom has exactly 12 atomic mass units.

Until here everything is fine.

Now my question is... To define all the other elements isotopes atomic mass did we first measured each isotope atom mass and then said " If 1/12 of the Carbon 12 atom is equal to X grams and this X grams are equal to 1 atomic unit the then my isotope atom which has Y grams must have Y atomic mass units " did we work it out this way or did we use calculation to define other isotopes mass based on carbon 12 atomic mass ?

Basically what I want to know is, do we need to have the value of an atoms mass in grams or kilograms before we know how many Amu´s it has, or can we calculate it ?

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    $\begingroup$ Yes we do need to measure each isotope of each element. I guess you've heard of mass spectrometry. $\endgroup$
    – Ivan Neretin
    Commented Jan 19, 2018 at 4:52
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    $\begingroup$ One other odd point. We know the mass of the individual isotopes much more accurately than we know the average atomic mass of the element. The problem is that the isotopic ratios have variation. $\endgroup$
    – MaxW
    Commented Jan 19, 2018 at 6:00
  • $\begingroup$ @IvanNeretin I wont say that I completely understand it but yes I´ve heard about it. Thank you for making it clear ;) $\endgroup$
    – Goncalo Fonseca
    Commented Jan 19, 2018 at 12:51
  • $\begingroup$ Not entirely sure what you're getting at here. As far as I can tell the process of "calculate isotope mass from ratio with C12 amu" and "calculate isotope mass by first converting to grams" are mathematically equivalent. What do you expect to differ between the two? $\endgroup$
    – chipbuster
    Commented Jan 19, 2018 at 18:47

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Historically, the relative masses of the isotopes were measured relative to O=16.

There is a good historical explanation in A New Mass-Spectrograph and the Whole Number Rule (1927).

At this time (1927), though it was understood that many elements had more than one isotope, only one isotope of oxygen was known.

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