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AMU
Name: Myron G.
Status: educator
Age: 60s
Location: N/A
Country: N/A
Date: 2000
Question:
1 amu is defined as 1/12 the mass of carbon 12. Carbon
12 is assigned an amu of 12.00000. Yet, 1 amu is reported as greater
than 1.0000!
Values given are in the area of 1.008.
Equally confusing, the amu of both the proton and the neutron is reorted
as greater than 1.000 (reported values for both are essentially the
same). There are 12 of them (in total) in Carbon 12. So, Carbon 12 would
have an amu greater than 12.0000!
Can you explain this apparent contradition to me?
Replies:
One amu. is defined as 1/12 the mass of the mass 12 isotope of carbon,
exactly.
The number that you refer to as 1 amu. = 1.008 is for the mass of the
element hydrogen, H. But this is a weighted average of the naturally
occurring isotope deuterium. The same thing happens for most elements. Only
a few, like fluorine, have only a single naturally occurring isotope. Things
are further complicated by the fact that the distribution of naturally
occurring isotopes is not constant, but can vary depending upon the origin
of the sample. In addition, the rates of chemical reactions does
discriminate slightly from one isotope to another of the same element, so
even chemical reaction can cause a change in the isotopic abundance.
The fact that the rest mass of a proton or a neutron is not exactly 1 amu.
arises from the fact that these particles are themselves made up of more
"elementary" particles. So, for example, the mass of a neutron is not
exactly equal to the mass of a proton + the mass of an electron. These mass
"discrepancies" occur because of the equivalence of energy and mass --
E=m*c^2 -- for quantum mechanical reasons.
You can obtain the "best" values of these fundamental constants by doing a
WEB search on the subject: "recommended values fundamental physical
constants".
There is no contradiction, but only some small, but real differences due to
some subtle quantum mechanical effects.
Vince Calder
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Update: February 2012
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