L10. Isotopes and Atomic Mass

Atoms and the Periodic Table

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L10. Isotopes and Atomic Mass

If every atom of an element has the same number of protons, why do some atoms of that element weigh more than others?

A curious classroom sample

You pour a few drops of saltwater into a dish and test its properties. Chemically it behaves like sodium chloride — the same element types and same bonding. But if you could weigh single atoms, some chlorine atoms would be slightly heavier than others. That difference shows up in precise balances and in the tiny mass numbers scientists report on the periodic table. This lesson connects the particle model you already learned (protons, neutrons, electrons) to a real observation: atoms of one element can have different masses. We will use simple particle drawings and one clear calculation to explain why, and then look at one everyday use of isotopes.

What an isotope is — a particle-level view

An isotope is one of several versions of the same element that differ only in the number of neutrons in their nuclei. The number of protons (the atomic number) stays the same, so isotopes of an element have the same chemical identity but different mass. Think of it in terms of the particles you already know:

• Protons: same count for all isotopes of an element (defines the element). Neutrons: vary between isotopes (change the mass). Electrons: usually match protons for a neutral atom (same chemical behavior).

How atomic mass on the periodic table is calculated

Worked example (chlorine): Chlorine has two common isotopes: 35Cl with mass ≈ 34.9689 u and about 75.8% abundance, and 37Cl with mass ≈ 36.9659 u and about 24.2% abundance. Calculation: convert percents to decimals: 0.758 and 0.242. Contribution from 35Cl: 34.9689 u × 0.758 = 26.50 u (rounded). Contribution from 37Cl: 36.9659 u × 0.242 = 8.95 u (rounded). Add contributions: 26.50 u + 8.95 u = 35.45 u. The unit is atomic mass unit (u), so the average atomic mass = 35.45 u. That value appears on the periodic table because it reflects how many atoms of each isotope you expect in a typical sample of chlorine on Earth.

• Step 1 — List isotopes with their mass numbers (rounded masses in atomic mass units, u) and natural abundances (percent). Step 2 — Convert each percent abundance to a decimal fraction (percent ÷ 100). Step 3 — Multiply each isotope’s mass (u) by its decimal abundance to get the isotope’s contribution (u). Step 4 — Add all contributions to get the average atomic mass (u).

Everyday uses and safety

Isotopes matter in medicine, geology, and the environment. A few examples at a level you can picture: - Radiocarbon dating uses the small amount of carbon-14 (an unstable isotope) in once-living things to estimate age. Scientists count radioactive decays to infer time since death. - In medicine, doctors sometimes use isotopes as tracers. A safe amount of a radioactive isotope can be added to a molecule that travels in the body; detectors trace where the molecule goes without changing chemistry. - Environmental scientists measure ratios of stable isotopes (like different oxygen isotopes in water) to learn about climate history or water sources. All of these applications depend on knowing which isotopes exist, how common they are, and how their masses differ. That knowledge comes directly from the particle model (protons and neutrons) plus careful measurements of abundance and mass in units of u.

Big idea and where this leads

Isotopes are versions of the same element that differ in neutron number; their different masses combine into the average atomic mass shown on the periodic table. You can predict and calculate that average by multiplying each isotope’s mass (in u) by its abundance and summing the results.

Understanding isotopes prepares you to read the periodic table more deeply (next in the course) and to follow real-world uses like dating fossils and tracking molecules in medicine. The particle-level view you learned earlier (protons, neutrons, electrons) is the key model that makes isotopes predictable and measurable.