The mass of a nucleus is determined by the total number of protons and neutrons it contains. Protons and neutrons have nearly the same mass and are collectively called nucleons. The total number of nucleons is called the mass number.
The mass of a proton is about 1.66 x (10 raised to the power -27) grams. Clearly then, the masses of individual atoms are much too small to be expressed in convenient units. Instead, atomic masses are measured by international agreement in "mass units." The mass unit, symbol "u," is defined as exactly 1/12 the mass of the atom C-12, which contains 12 nucleons. The mass of a C-12 is therefore exactly 12u. The masses of all other atoms have decimal values.
Isotopes are of two types: stable and unstable. Unstable isotopes are radioactive, decaying spontaneously into other isotopes. Some elements, such as fluorine and phosphorus, have just one naturally occurring isotope. Many elements, however, have two or more isotopes. For example, the simplest element, hydrogen, has two stable isotopes, 1-H and 2-H. The nucleus of 1-H is a single proton. The nucleus of deuterium, 2-H, is composed of one proton and one neutron. In addition to the naturally-occurring isotopes, many others have been produced artificially by bombarding nuclei with various sub-atomic particles. Artificial isotopes include those that constitute new elements.
Because elements occur naturally as mixtures of their isotopes, the mass of an element is actually its average mass. The average mass depends on the relative abundances of its isotopes. For example, water and other hydrogen-containing compounds are always found to contain about 99.985 percent 1-H and 0.015 percent 2-H. This means the "fractional abundances" of 1-H and 2-H are 0.99985 and 0.00015, respectively. Because the mass of a 1-H atom is 1.00783u and the mass of a 2-H atom is 2.01410u, the "average mass" of a hydrogen atom is the sum of the fractional contributions of its isotopes, which is calculated as follows: 1-H, 0.99985 x 1.00783u = 1.00768u; 2-H, 0.00015 x 2.01410u = 0.00030u; the sum of the contributions of both isotopes is 1.00768u + 0.00030u = 1.00798u.