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Average Atomic Mass Average Atomic Mass The atomic mass is the mass of a specific isotope, most often expressed in unified atomic mass units. The atomic mass is the total mass of protons, neutrons and electrons in a single atom. The atomic mass is sometimes incorrectly used as a synonym of relative atomic mass, average atomic mass and atomic weight; these differ subtly from the atomic mass. The atomic mass is defined as the mass of an atom, which can only be one isotope at a time and is not an abundance-weighted average as in the case of atomic weight. In the case of many elements that have one dominant isotope the actual numerical similarity/difference between the atomic mass of the most common isotope and the relative atomic mass or standard atomic weights can be very small such that it does not affect most bulk calculations—but such an error can be critical when considering individual atoms. For elements with more than one common isotope the difference even to the most common atomic mass can be half a mass unit or more (e.g. chlorine). The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.

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Standard atomic weight refers to the mean relative atomic mass of an element in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances. These are what are included in a standard periodic table and is what is used in most bulk calculations. An uncertainty in brackets is included which often reflects natural variability in isotopic distribution rather than uncertainty in measurement. For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning. Therefore, for synthetic elements the total nucleon count of the most stable isotope (i.e., the isotope with the longest half-life) is listed in brackets in place of the standard atomic weight. Lithium represents a unique case where the natural abundances of the isotopes have been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources, such as rivers. Relative atomic mass is a synonym for atomic weight and closely related to average atomic mass (but not a synonym for atomic mass), the weighted mean of the atomic masses of all the atoms of a chemical element found in a particular sample, weighted by isotopic abundance. This is frequently used as a synonym for the standard atomic weight and it is correct to do so since the standard atomic weights are relative atomic masses, although it is less specific to do so. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the average or have different certainties (number of significant figures) than the standard atomic weights. Relative isotopic mass is the relative mass of a given isotope (more specifically, any single nuclide), scaled with carbon-12 as exactly 12. No other nuclides other than carbon-12 have exactly whole-number masses in this scale. This is due to two factors: the different mass of neutrons and protons acting to change the total mass in nuclides with proton/neutron ratios

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other than the 1:1 ratio of carbon-12; and an exact whole-number will not be located if there exists a loss/gain of mass to difference in mean binding energy relative to the mean binding energy for carbon-12. Any mass defect due to binding energy is a small fraction (less than 1%) compared to the mass of a nucleon, and even less compared to the average mass per nucleon in carbon-12, which is moderately strongly bound. Since protons and neutrons differ in mass from each by an even smaller fraction (about 0.0014 u), the practice of rounding the atomic mass of any given nuclide or isotope to the nearest whole number, always gives the simple whole number total nucleon count. Neutron count can then be derived by subtracting the atomic number. The mass number of a nuclide is simply the total number of nucleons in the nucleus. It is equal to the number of protons (atomic number) plus the number of neutrons. This number is always a simple whole number. It has units of "nucleons" not atomic mass units. An example is oxygen-16, which has 16 nucleons (8 protons and 8 neutrons).

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