Subatomic Particles and Atomic Mass vs. Atomic Weight
High-Yield Summary
- Atoms are built from protons (+1, nucleus), neutrons (0, nucleus), and electrons (−1, orbiting) — valence electrons drive bonding.
- Atomic number (Z) = protons (defines the element); mass number (A) = protons + neutrons; A − Z = neutrons.
- Isotopes share Z but differ in A. Atomic weight = the abundance-weighted average mass (in amu) of all naturally occurring isotopes — not the mass of any single atom.
- 1 amu = 1/12 the mass of carbon-12 ≈ 1.66×10⁻²⁷ kg. A mole = 6.022×10²³ entities (Avogadro's number); molar mass (g/mol) links moles to grams.
- Ordinary chemical reactions change only electron count/arrangement — proton/neutron counts change only in nuclear reactions.
- An ion forms when proton count ≠ electron count: cation (positive, fewer electrons) or anion (negative, more electrons).
The Three Subatomic Particles
| Particle | Charge, Mass, Location |
|---|---|
| Proton | +1, ≈1 amu, nucleus |
| Neutron | 0, slightly more than 1 amu, nucleus |
| Electron | −1, ≈1/1,836 of a proton's mass, orbiting the nucleus |
Key Terms
- Atomic number (Z)
- Number of protons; determines the element's chemical identity.
- Mass number (A)
- Protons + neutrons (nucleons). Not the same as atomic weight.
- Isotopes
- Atoms of the same element (same Z) with different numbers of neutrons (different A).
- Atomic weight
- Abundance-weighted average mass, in amu, of an element's naturally occurring isotopes.
- Atomic mass unit (amu)
- 1/12 the mass of a carbon-12 atom; ≈1.66×10⁻²⁷ kg.
- Mole / Avogadro's number
- 1 mole = 6.022×10²³ entities; molar mass (g/mol) converts moles to grams.
- Cation / Anion
- Cation: more protons than electrons (+). Anion: more electrons than protons (−).
Atomic Weight from Isotope Data
Atomic weight = Σ (isotope mass × natural abundance)
- isotope mass = Exact mass of that isotope, in amu
- natural abundance = Fraction of that isotope found in nature (%, as a decimal)
- Example: Isotope A (10 amu, 90% abundant) + Isotope B (11 amu, 10% abundant) → (10 × 0.90) + (11 × 0.10) = 10.1 amu.
- This weighted average is what appears on the periodic table — no single atom necessarily weighs exactly that amount.
Common MCAT Trap
- Mass number (A, a whole-number count of nucleons in ONE atom) is not atomic weight (a decimal, abundance-weighted average across ALL isotopes) — don't conflate them.
- In ordinary chemical reactions, only electrons move — protons/neutrons change only in nuclear reactions. A charge change (ion formation) never means the element changed.
- Electron mass is negligible (~1/1,836 of a proton's), so electrostatic attraction — not gravity — dominates atomic-scale interactions.
Quick Recall
How do you find the number of neutrons in an atom?
Why isn't atomic weight the mass of any single atom?
What changes during an ordinary chemical reaction — protons, neutrons, or electrons?
An atom has more electrons than protons — cation or anion?