Elements are defined by the number of protons inside their nucleus. The most common and stable form of an element has an equal number of protons and neutrons. However, variations of elements exist in which the number of neutrons outnumbers the number of protons. These atoms are known as isotopes of their particular element. For example, a traditional carbon atom has six protons and six neutrons. A carbon isotope may have six protons and seven or eight neutrons.
Methods of recording isotopes on paper differ from traditional elements. Isotope notation is composed of two parts. The first part outlines the base element using its symbol. The second part is the sum of the isotope's protons and neutrons. For example, a carbon isotope with six protons and seven neutrons would be written as C13.
Some isotope variations are more stable than others. If an isotope has too many or too few neutrons, it may undergo radioactive decay. Radioactive decay causes unstable isotopes to return to more stable forms. For example, C14 is an unstable isotope and, therefore, will slowly decay into a more stable form, C12. Radioactive decay happens at a regular and predictable pace. Carbon, for example, will decay from C14 to C12 in couple thousand years. Decay rates vary depending on the element and range from a couple minutes to several thousands of years.
Atoms are organized on the periodic table according to their atomic mass. The periodic table holds 118 known elements, as of 2011. The atomic mass of an element is the sum of its neutrons and protons. However, atomic masses are never whole numbers. Atomic masses of each element must take its isotope variations into account. Therefore, the atomic mass of elements as listed on the periodic table are the average of an element's traditional atomic mass and that of its isotope variations. Carbon, for example, has an atomic mass of 12.011, an amount taken from the average of its C12, C13 and C14 isotope forms.