# Thread: atomic mass units - how so?

1. Revising my (very basic) physics at present.

I am led to understand that the atomic mass unit (or 'standard atomic weight as defined in the periodic table of an element is some form of average of the distribution of the various isotopes of the element on the planet - ie. carbon is 12.01 meaning most of the planets carbon is carbon 12 and some is carbon 14
I realise that in special cases the physicist would use the atomic mass of a specific isotope, but why the 'standard' atomic weight as defined in the periodic table? It just puzzles me why such an exact mathematical science as physics should use such a ostensively bizarre concept.
Surely the atomic mass, as defined in this way, is merely an arbitrary number, as it must be the case that on other planets/stars or in the universe as a whole the distribution of the isotopes would vary and hence the so called atomic mass would calculate differently in that context?
Secondly if you are in a real world situation then the atomic mass of (e.g) the carbon in your test tube would be specific and not a function of the average on the planet?

I cant see why such a measure would be functionally useful?

2.

3. Originally Posted by ZeroZero
Revising my (very basic) physics at present.

I am led to understand that the atomic mass unit (or 'standard atomic weight as defined in the periodic table of an element is some form of average of the distribution of the various isotopes of the element on the planet - ie. carbon is 12.01 meaning most of the planets carbon is carbon 12 and some is carbon 14
I realise that in special cases the physicist would use the atomic mass of a specific isotope, but why the 'standard' atomic weight as defined in the periodic table? It just puzzles me why such an exact mathematical science as physics should use such a ostensively bizarre concept.
Surely the atomic mass, as defined in this way, is merely an arbitrary number, as it must be the case that on other planets/stars or in the universe as a whole the distribution of the isotopes would vary and hence the so called atomic mass would calculate differently in that context?
Secondly if you are in a real world situation then the atomic mass of (e.g) the carbon in your test tube would be specific and not a function of the average on the planet?

I cant see why such a measure would be functionally useful?
The atomic mass in the periodic temple is of more utility in chemistry than in physics. Chemical experiments usually involve quite a few atoms, and the usual abundance of isotopes is meaningful in that context.

4. But surely their would be variation of the consituant isotopes in the test tube (i.e. sample) which would be to a significant extent dissimilar to the random distribution average across the earths crust, making the determining of the average standard weight useless to a calculation? why not simply go with the atomic weight of thev isotope you believe the sample to be?

5. Originally Posted by ZeroZero
But surely their would be variation of the consituant isotopes in the test tube (i.e. sample) which would be to a significant extent dissimilar to the random distribution average across the earths crust, making the determining of the average standard weight useless to a calculation? why not simply go with the atomic weight of thev isotope you believe the sample to be?
The relative abundance of isotopes is pretty much the same in any sample unless there is some process to change it -- as with Carbon 14 dating that relies on cosmic ray interaction with nitrogen to produce the neutrons that maintain a small amount of carbon 14 in the atmosphere, while entrapped carbon loses carbon 14 due to natural radioactive decay..

How would you even begin to guess what the distribution of isotopes in your test tube would be ? Unless you have processed the material in a very sophisticated and expensive way, you will not have anything like a single isotope unless that is the natural state. The uranium enrichment segment of the Manhattan project, for instance, was a huge undertaking.

6. So you are saying that these standardised units are pretty much accurate across all of the elements in the periodic table? I have no empirical evidence but this seems,
theoritically to be a pretty odd state of affairs?
surely there must be some, or even many, elements that do vary from these standards - across samples, in significant ways commonly?

Your second point seems to support my (highly amatuer) thesis

7. This is what Wikipedia says about it.
Atomic weights, unlike atomic masses (the masses of individual atoms), are not physical constants and vary from sample to sample. Nevertheless, they are sufficiently constant in "normal" samples to be of fundamental importance in chemistry.
The definition deliberately specifies "An atomic weight…", as an element will have different atomic weights depending on the source. For example, boron from Turkey has a lower atomic weight than boron from California, because of its different isotopic composition.[6][7] Nevertheless, given the cost and difficulty of isotope analysis, it is usual to use the tabulated values of standard atomic weights which are ubiquitous in chemical laboratories.
http://en.wikipedia.org/wiki/Atomic_weight

8. It's less surprising if you know a little statistics and take a look at the exponent on Avogadro's number. In even a small sample, there are a huge number of atoms. The standard deviation would be tiny compared to what would matter to almost any common experiment.

9. In the context of mass spectrometry, an atomic mass unit is defined as one twelfth the mass of the carbon 12 isotope atom. It isn't a physical unit in the normal sense but simply a means of enumerating the peaks that appear in a mass spectrum. A somewhat puzzling aspect of this definition is that uncharged atoms never appear in mass spectra - only ions do. A peak with a mass of twelve amu (or any other integer) is never seen in a high resolution spectrum.

(I'm old enough to remember when it was defined as one sixteenth of Oxygen 16 !)

10. Originally Posted by Old Fool
(I'm old enough to remember when it was defined as one sixteenth of Oxygen 16 !)
Joseph Priestley and I went to different schools together.

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