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Atomic mass, often referred to as atomic weight, is a key concept in understanding elements and their atoms. It's a very precise measurement that reflects the average mass of all the isotopes of an element, considering not only their mass but also their relative abundance in nature. To determine the atomic mass of an element, scientists use a unit called an atomic mass unit (amu).

Let's take a closer look using the element Carbon. With its isotopes Carbon-12 and Carbon-13, the atomic mass is not simply a matter of adding up protons and neutrons. Instead, it's calculated using a weighted average. If you imagine a sample of Carbon with 98.9% Carbon-12 and 1.1% Carbon-13, the atomic mass would result closer to 12.01 amu rather than a whole number. This slight difference accounts for the presence of the heavier Carbon-13 isotope.

The mass number, also known as the nucleon number, is much more straightforward than the atomic mass. It is a total count of the protons and neutrons in an atom's nucleus. This figure is always a whole number and it tells us which isotope of an element we're looking at. Unlike atomic mass, it does not consider the abundance of isotopes or their exact atomic mass values.

Using Carbon as an example again, the mass number for Carbon-12 is 12, because it has 6 protons and 6 neutrons. For the heavier Carbon-13 isotope, the mass number is 13, the sum of its 6 protons and 7 neutrons. This number helps us distinguish between different isotopes of an element.

Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons. This means they have the same atomic number but different mass numbers. Isotopes play a crucial role in defining atomic mass, as they each contribute to the weighted average of an element's atomic mass.

For example, Hydrogen has three isotopes: Protium, with no neutrons; Deuterium, with one neutron; and Tritium, with two neutrons. They each have their own unique mass number and properties. It's the variation in the number of neutrons that leads to a variety of atomic masses among isotopes.

The nucleon number is essentially the same as the mass number—it's the combined total of protons and neutrons in the nucleus of an atom. In our discussions of isotopes and atomic mass, this number is always integral since it uniquely identifies an isotope of an element. Moreover, it's the nucleon number that is linked to the concept of an element's mass number.

By knowing the nucleon number, you can figure out the number of neutrons in an isotope by subtracting the number of protons (the atomic number) from it. For instance, in Oxygen-16, the nucleon number is 16 and the atomic number is 8, meaning it has 8 neutrons.

Atomic mass units (amu) are the standard units used to express the mass of atoms and molecules. One atomic mass unit is defined as one twelfth of the mass of a Carbon-12 atom, making it a universal constant based on an actual physical substance. This small unit helps scientists and chemists calculate and compare the weights of atoms precisely.

When looking at a periodic table, the atomic mass is usually listed below the symbol of each element, illustrating how much one mole of the element weighs in grams, which is essentially the mass expressed in amu.