Re: molecular weight of dry air

[Brian Whatcott <betwys1@SBCGLOBAL.NET>]

   I certainly want to acknowledge to  Michael Edmiston
that I have next to no insight in this area:  still,  without
wishing to seem ungrateful for his corrections to my values
of molecular weights using his values of atomic masses,
and without passing on whether chemists are or are not dumb,
I suppose I could reasonably append a paragraph from
Enc. Chemistry, 3rd Ed., Hampel & Hawley,
 Van Nostrand '66   P117 :

~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~
Atomic weights.

"The atomic weight of an element
is the average mass of the atoms making up
a sample of the element when these masses are
expressed on a proportional scale on which the
average mass of the atoms of a sample of ordinary
oxygen is taken as exactly 16.

On such a scale,
hydrogen has an atomic weight of 1.008. When
using this scale, the whole number or integer nearest
to the mass of a particular atom is called the
mass number of that atom. This scale is sometimes
called the "chemical" system. Another scale called
the "physical" system is also used, especially where
single atoms are to be compared. Some types of
oxygen atoms are heavier than others though all
have the properties of oxygen. The physical scale
selects the mass of the most common form (isotope),
calls it 16, and uses that as a standard.
Under this system the atomic weights come out
about 1.0002 times those based on the chemical
scale.

A "unified"' scale to replace the other two was
adopted by the International Union of Pure and
Applied Physics and the International Union of
Pure and Applied Chemistry. For this scale a
standard value of 12 is used for the most abundant
type of isotope of carbon."
 ~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~

It turns out chemical analyses of atomic weight
based on Avogadro yield only relative molecular weights.
Initially,  Hydrogen was represented as 1, but a scale
 based on Oxygen = 16 came to be preferred   (as offering
many combinations.)   When elements were found to be
a mix of isotopes however,  and the standard, oxygen itself
 was affected,   physicists began to assign the common isotope
examined with the mass spectrometer  the value 16, while
chemists continued with  16 for the natural mix, until the
'quasi-physical' definition of  C = 12 for the most common
 isotope of that material was adopted by IUPAP/IUPAC.

Hence when Michael gives us "correct" values, it would be
well to remember that he has in mind the  values
currently defined to be "correct".

:-)

Brian W


At 07:20 PM 5/12/2004, Michael Edmiston, you wrote:
> Several have indicated a formal statistical procedure
> for determining the molecular masses of things like O2,
> but most did not actually calculate the number.
> One who did, Brian Whatcott, did it incorrectly
> because he used atomic mass numbers
> rather than atomic masses.
>
> The atomic mass of 16O is 15.99492
> The atomic mass of 17O is 16.99913
> The atomic mass of 18O is 17.99916
>
> Using the abundances that others in this thread have used:
>
> 99.757 percent
> 0.038 percent
> 0.205 percent
>
> and multiplying these percentages by the masses
> and summing,
> we get the average atomic mass of natural oxygen as:
>
> 15.99941
>
> and this is the number I find on my periodic table.
>
> If we figure the molecular mass of O2 by
> simply doubling this number, we get:
>
> 31.99881
>
> which is the correct number for the average
> molecular mass of natural O2.
>
> If we do the statistics that John Denker provided,
> and Brian Whatcott calculated, except we use the
> correct atomic masses rather than the mass numbers
> that Brian used, then instead of the 32.01059 that
> Brian got, we actually get:
>
> 31.99881
>
> which is the same number as obtained by simply
> doubling the average atomic mass found on
> the periodic table.
>
> Think about this for a minute.  If the true average
> molecular mass of a molecule could not be found
> by simply adding the average atomic masses found
> in the periodic table, then chemists have been using
> the wrong molecular masses for a long time.
>
> Duh... chemists are not that dumb.  They've been
> doing it right.  The atomic masses on the periodic table
> are the weighted averages (weighted by natural abundance)
> of the atomic masses (not mass numbers) of the naturally
> occurring isotopes of the elements.  It is specifically done
> this way to allow us to find average natural molecular
> masses by simply adding the atomic masses listed
> in the periodic table.
>
> Michael D. Edmiston, Ph.D.
> Professor of Physics and Chemistry
> Bluffton College
> Bluffton, OH 45817
> (419)-358-3270
> edmiston@bluffton.edu


Brian Whatcott    Altus OK    Eureka!