Talk:Atomic mass/Archive 1

Archive 1

The unit

What is the unit you messure here u or g/mol?Stone 08:39, 24 February 2006 (UTC)

The atomic mass itself is just that, a dimension of mass. Technically you could use any unit of mass you like. The unified atomic mass unit (u), and grams per mole (g/mol) are two equivalent ways of expressing one and the same scale based upon the carbon-12 atom, respectively 12 g of carbon and the number of atoms therein. Femto 12:41, 25 February 2006 (UTC)

History

Nothing in the article to indicate when the concept of atomic mass / atomic weight first developed. -- Jmabel | Talk 06:53, 30 March 2006 (UTC)

Definition

"The atomic mass of a chemical element (also known as the relative atomic mass or average atomic mass or atomic weight) is the average atomic mass of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance." This is a horrendous example of circular logic. To simplify the sentence it says: "The atomic mass, aka the average atomic mass, is the average atomic mass. Even worse is that is goes on to contradict itself. First it says that the atomic mass is the average atomic mass (all isotopes abundance averaged) and then continues to speak about the atomic mass of individual isotopes. I have reworked the beginning of the definition to fix these issues and given external links to the IUPAC definitions as reference. --134.9.228.11 00:35, 4 April 2006 (UTC)

The IUPAC and NIST definitions of "relative atomic mass" do not appear do agree. Could somebody clear this up?--63.145.1.170 15:15, 21 September 2007 (UTC)

Too technical

15-Oct-2007: This article had been tagged "too technical" on 11-Dec-2006. To help simplify the technical presentation, I have added a diagram of a lithium atom, counting the protons, neutrons and electrons. Beyond that diagram, please list specifics, below, about why the article is too technical. I am removing the {technical}-tag at top, which did not simplify this article during the past 11 months. -Wikid77 09:45, 15 October 2007 (UTC)

Name

Woudn't it be better if someone moves all of the data of this article (Atomic mass) to Atomic Mass. Atomic Mass (with a capital M for Mass) would be a better name of an article for an encyclopedia. --Krupted Soul (talk) 10:56, 10 August 2008 (UTC)

Second word of an article title is only capitalized if it is a proper noun per WP:MOSCAPS. Vsmith (talk) 12:38, 10 August 2008 (UTC)

Isotope distribution (extraterrestrial)

Atomic mass is the sum of the total number of protons plus the total number of neutrons. Please disregard the misleading article.

A good example of heterogeneous isotope distribution is Carbon-14 (which is radioactive). Each object in the solar system has a (more or less) steady-state concentration of carbon-14, but since C14 is radioactive, it is generally replenished by solar neutron radiation bombardment and transmutation of atmospheric nitrogen(14). For example, if venus and earth had the same amount of nitrogen, then venus would have more C14 - since the neutron flux at venusian orbital radius is higher. Of course, given similar neutron flux, C14 generation depends on the concentration of atmospheric nitrogen.

Deuterium is known to be more prevalent farther out in the solar system (though it's not entirely known why). Between solar systems, the common isotope ratios are fairly close to within each other, but our data are limited. Presumably isotope ratios between different generation of stars are different, and who knows about different galaxies. Isotopic fractionation in planets is also known, which may contribute to abberations in the isotope ratio.

Ultimately, this means that the average molar mass of a particular compound is dependent on where you are (although the differences are likely to be minute).

isotope ratios in exobiology helium is particularly bad isotopic analysis for understanding origin of terrestrial impact debris

— Preceding unsigned comment added by Foreverstacy (talk • contribs) 05:25, 29 February 2012 (UTC)

See mean atomic mass number which is actually what you're trying to define. What point are you trying to make? Did you read the article on deuterium? Of course atomic weight will depend on where you are in a solar system (or on a planet), for very many reasons. Did you bother to read this article? Or the linked articles which cover things like isotope geochemistry? SBHarris 06:31, 29 February 2012 (UTC)

Many errors

This article has many errors in it...

• incorrect use of units
• incorrect use of technical terms (atomic mass unit, etc.)
• is nonsensical in some places (i.e., "even smaller percentage" w/r/t C-12)

This article looks like lecture notes written by a professor who is bad at explaining things. — Preceding unsigned comment added by 174.136.103.210 (talk) 21:38, 19 January 2014 (UTC)

Your objections are really too general to be able to answer. Incorrect use of units where? As for technical terms, this article does not use "atomic mass unit" (which strictly speaking has been obsolete since 1961). Rather, it uses unified atomic mass unit, and does so correctly. SBHarris 01:19, 10 October 2014 (UTC)

The section on mass number deviation is an overly-broad generalization.

"The amount that the ratio of atomic masses to mass number deviates from 1 is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the fact that nuclear fission in an element heavier than zirconium produces energy, and fission in any element lighter than niobium requires energy. On the other hand, nuclear fusion reactions: fusion of two atoms of an element lighter than scandium produces energy, whereas fusion in elements heavier than calcium requires energy."

This is a generalization that is sort of true most of the time, but shouldn't be stated without qualification. I would suggest it not be stated at all. Its most glaring defect is when it comes to the spike in binding energy for He-4. If this statement were true, we could fuse two He-4 nuclei to produce Be-8 with a release of energy. I don't believe that happens, and it isn't what this graph would predict. In fact, based on another Wikipedia article, I believe the reverse happens: Be-8 decays almost immediately into two He-4's with the release of energy.

A better thing to do here might be to talk about why fission of U-235 produces energy and why fusion of deuterium nuclei produces energy, referring to the graph. I should point out that I am not an expert in nuclear, or any other kind, of physics. I'm just an interested amateur. Otherwise I would be happy to write that text. — Preceding unsigned comment added by Stuart.soloway (talk • contribs) 21:35, 4 April 2014 (UTC)

You're certainly right about the fusion energy bottleneck created by the stability of He-4 so it cannot simply fuse to Be-8. If it weren't for the triple alpha process in stars that gets around that, we wouldn't be here! I'll see if I can generalize the section above so it's not strictly wrong in that regard. SBHarris 01:24, 10 October 2014 (UTC)

Persistent Confusion on Relative Atomic Mass

I have once again corrected the persistent confusion on relative atomic mass. The most important point to remember is that relative atomic mass is practically an antonym of atomic mass. Relative atomic mass is a weighted average and atomic mass is not. I know this is very confusing but please read definitions carefully and make sure you understand the differences.

Correct statement:

Relative atomic mass is a synonym of atomic weight.

Incorrect Statement:

Relative atomic mass and relative isotopic mass are essentially the same.

Why: The relative atomic mass is a weighted average. The relative isotopic mass is not.

Incorrect Statement:

Relative atomic mass and atomic mass are essentially the same.

Why: The relative atomic mass is a weighted average. The atomic mass is not.

Incorrect statement:

The mass defect of all atomic masses above C12 are positive.

Why: They are in fact mostly negative and only become positive at high Z and low Z. The error here is in the confusion between mass defect of atoms and the atomic weights (or more precisely the standard atomic weights) that mostly have masses slightly above the nominal atomic weight. This trend in the decimal places of the atomic weights has to do the relative prevalence of heavier isotopes at higher Z and has little to do with mass defect due to nuclear binding energies.

--Nick Y. (talk) 18:33, 22 September 2009 (UTC)

I find the objections above all correct, and hopefully they have all been fixed in the present version of the article. SBHarris 23:49, 10 October 2014 (UTC)