Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams. Scientists arbitrarily define this amount of mass as one atomic mass unit (amu) or one Dalton. Electrons are much smaller in mass than protons, weighing only 9.11 × 10-28 grams, or about 1/1800 of an atomic mass unit. As a result, they do not contribute much to an element's overall atomic mass. This means that, when considering atomic mass, it is customary to ignore the mass of any electrons and calculate the atom's mass based on the number of protons and neutrons alone.
However, since most naturally occurring elements are composed of isotopes, it is important to consider the mass and natural abundance of isotopes when determining the atomic mass of such elements. The atomic mass in such cases is calculated by summing the masses of all the element's isotopes, each multiplied by their natural fractional abundance.
Chemists often use the unit "mole" to determine the number of atoms of a compound that participate in a chemical reaction. A mole of an element is its atomic weight in grams, while a mole of a compound is the sum of the atomic weights of its components, called the molecular weight. An often-used example is calculating a mole of glucose with the chemical formula C6H12O6. The atomic weight of carbon (C) is 12.011 grams, and there are six carbons in glucose for a total atomic weight of 72.066 grams. Doing the same calculations for hydrogen (H) and oxygen (O), the molecular weight of glucose equals 180.156 grams.
This text is partially adapted from Openstax, Biology for AP® courses. Section 2.1 Atoms, Isotopes, Ions, and Molecules: The Building Blocks and Openstax, Anatomy and Physiology 2e, Section 2.4 Inorganic compounds Essential to Human Functioning
