The most fundamental relationship in nuclear physics is Albert Einstein's relationship between mass and energy:

E=mc2

This relationship tells us that mass can be converted to energy and vice-versa. While to Einstein this was probably quite obvious, to the rest of us it is likely not so.

Before jumping into some simple example problems based on this concept, let's review some basic definitions including some units and prefixes to units that we frequently use:

A **mole** is defined as the amount of
substance that contains as many elementary entities (e.g., atoms,
molecules, ions, electrons) as there are atoms in 12 g of the
isotope carbon-12 (12C). Thus, by definition, one mole of pure 12C
has a mass of exactly 12 g. Therefore, an atomic mass unit (amu) is
typically in units of grams/mole: Carbon-12 has a mass of 12 amu or
one mole of pure Uranium-235 has a mass of 235.043923 grams.

**Avogadro's number** is a
universal constant commonly used in nuclear physics for the number
of atoms (or molecules) in one 1 mole of substance. The value is
6.022×1023. This number is
extensively used in calculating number densities and energy
conversions from mass.

If a mass is given in amu and it must be converted to grams, then dividing by Avogadro's number will give grams/atom.

If you are still confused about Avogadro's number, take a look at this video.

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