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Isotopes are atoms of the same element with different numbers of neutrons, leading to different mass numbers. The weighted average of isotope masses determines an element's atomic mass on the periodic table.
Calculate average atomic mass from given isotope abundances and masses.
Thinking all atoms of an element are identical; confusing isotopes with different elements.
From atomic structure basics, you know that an atom consists of a nucleus containing protons and neutrons, surrounded by electrons. The number of protons — the atomic number (Z) — defines which element an atom is. Every carbon atom has exactly 6 protons; every iron atom has exactly 26. But the number of neutrons in the nucleus can vary without changing the element's identity. Atoms of the same element that differ in their neutron count are called isotopes, and this variation is far more common than you might expect — most elements exist naturally as a mixture of two or more isotopes.
Consider carbon. Every carbon atom has 6 protons, but carbon exists primarily as three isotopes: carbon-12 (6 neutrons), carbon-13 (7 neutrons), and carbon-14 (8 neutrons). The mass number (A) — the total count of protons plus neutrons — distinguishes them: ¹²C, ¹³C, and ¹⁴C. Chemically, these isotopes behave almost identically because chemical behavior is determined by electron configuration, which depends on the number of protons (and thus electrons), not neutrons. They form the same bonds, participate in the same reactions, and have the same electronegativity. The difference is physical: they have different masses, and some isotopes (like ¹⁴C) have unstable nuclei that undergo radioactive decay.
The existence of isotopes explains why the atomic masses on the periodic table are not whole numbers. The atomic mass listed for carbon is 12.011 amu, not 12.000, because it is a weighted average of the masses of all naturally occurring isotopes, weighted by their relative abundance. Carbon-12 makes up about 98.9% of natural carbon and carbon-13 about 1.1%, so the average is pulled just slightly above 12. For chlorine, the effect is more dramatic: chlorine-35 (75.8%) and chlorine-37 (24.2%) give a weighted average of approximately 35.45 amu. You calculate this as: average atomic mass = (fraction₁ × mass₁) + (fraction₂ × mass₂) + ..., where the fractions must sum to 1.
Understanding isotopes opens doors to several important areas of chemistry and physics. Isotope ratios are used in radiocarbon dating (measuring the decay of ¹⁴C to determine the age of organic materials), in medical imaging (radioactive isotopes as tracers), and in mass spectrometry (where isotopic signatures help identify unknown compounds). The key takeaway is that the identity of an element is fixed by its protons, but the mass and nuclear stability of a particular atom depend on its neutron count — and this subtle variation has profound practical consequences.