Average Atomic Mass

Most elements, as they actually occur in nature, exist not as a single, uniform kind of atom but as a mixture of two or more isotopes — atoms of the same element that carry different individual atomic masses (because they have differing numbers of neutrons in the nucleus, despite having the same number of protons). Since any real, naturally-occurring sample of the element contains all of its isotopes mixed together in fixed natural proportions, the 'atomic mass' actually assigned to that element must be a WEIGHTED AVERAGE across all of its isotopes — weighted according to each isotope's relative abundance, i.e. its percentage occurrence in nature — and this weighted value is what is called the element's average atomic mass. Carbon is a worked example of this in the source text: it occurs naturally as a mixture of three isotopes (carbon-12, carbon-13 and a trace of carbon-14), each with its own atomic mass and relative abundance (see the accompanying isotope table), and combining these using the weighted-average method gives an average atomic mass of 12.011 u for carbon as a whole — which is precisely the value printed for carbon in the periodic table, since periodic-table atomic masses are always these average atomic masses, not the mass of any single isotope. For everyday convenience, this average is ofte …