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Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers but the same atomic number. Atomic mass is the weighted average of all naturally occurring isotopes' masses. Isotopes have different physical properties but similar chemical properties because chemistry depends on electron configuration.
From your study of atomic structure, you know that an atom's identity is defined by its number of protons — the atomic number (Z). Every carbon atom has 6 protons; every oxygen atom has 8. But the nucleus also contains neutrons, and here is the key: the number of neutrons can vary. Isotopes are atoms of the same element that differ in their neutron count. Carbon-12 has 6 protons and 6 neutrons (mass number 12), while carbon-13 has 6 protons and 7 neutrons (mass number 13). Both are carbon — same atomic number, same electron configuration, same chemical behavior. But they have different masses, which means different physical properties like density and rate of diffusion.
The notation is straightforward. The mass number (A) is the total count of protons plus neutrons. You write isotopes as the element symbol with the mass number as a superscript (¹²C, ¹³C, ¹⁴C) or in hyphenated form (carbon-12, carbon-13, carbon-14). The number of neutrons is simply A − Z. Since the periodic table lists elements by atomic number, and isotopes share the same atomic number, all isotopes of an element occupy the same box on the periodic table.
Now look at the atomic mass listed on the periodic table — for carbon, it reads 12.011 amu, not 12.000. That is because atomic mass is the weighted average of all naturally occurring isotopes, accounting for each isotope's mass and its natural abundance. Carbon is 98.9% carbon-12 (mass 12.000 amu) and 1.1% carbon-13 (mass 13.003 amu), with a trace of radioactive carbon-14. The weighted average calculation is: (0.989 × 12.000) + (0.011 × 13.003) = 12.011 amu. This is why no element has an atomic mass that is a whole number — the average always reflects the mixture of isotopes found in nature.
Understanding isotopes matters beyond just reading the periodic table. In mass spectrometry, which you will encounter later, isotopes produce distinct peaks that reveal molecular composition. Radioactive isotopes like carbon-14 are used in radiometric dating because they decay at known rates. In medicine, radioactive iodine-131 targets the thyroid gland for imaging and treatment. And the concept of weighted averages of isotopic masses is essential for converting between mass and moles — the molar mass you use in stoichiometry comes directly from these averaged atomic masses.