Atoms: Subatomic Particles & Electron Configuration
Protons, neutrons, electrons — and the rules that govern where electrons live. Isotopes, the Bohr model, modern orbitals, and how to write the configuration of any element you'll see on the CBE.
Every atom has three kinds of particles
- Protons — positive charge (+1), located in the nucleus, mass ≈ 1 amu. The proton count defines the element.
- Neutrons — no charge, in the nucleus, mass ≈ 1 amu (slightly heavier than proton). Hold the nucleus together.
- Electrons — negative charge (−1), occupy regions outside the nucleus, mass ≈ 1/1836 amu (negligible).
Key quantities:
- Atomic number (Z) = number of protons. Defines the element.
- Mass number (A) = protons + neutrons. The "weight" tag.
- Number of neutrons = A − Z.
- Charge of the atom = protons − electrons. Neutral atoms have equal counts; ions don't.
Isotopes — same element, different mass
Atoms of the same element with DIFFERENT neutron counts (and therefore different mass numbers) are called isotopes. Carbon-12, carbon-13, and carbon-14 all have 6 protons but 6, 7, and 8 neutrons respectively. They are still all carbon.
The atomic mass shown on the periodic table (e.g., chlorine ≈ 35.45) is a weighted average of an element's natural isotope mix. To calculate:
avg mass = Σ (isotope mass × fractional abundance)
Example for chlorine: (34.97 × 0.7577) + (36.97 × 0.2423) = 35.45 amu ✓.
Atomic theory — short history
- Dalton (1808): Atoms are indivisible, identical for an element, combine in fixed ratios.
- Thomson (1897): Discovered the electron via cathode rays — atoms have internal structure.
- Rutherford (1911): Gold-foil experiment — most of the atom is empty space; a small dense positive nucleus exists.
- Bohr (1913): Electrons occupy fixed quantized energy levels; jumps between levels emit/absorb photons.
- Schrödinger (1926): Quantum mechanical model — electrons occupy 3D probability regions (orbitals), not fixed orbits.
- Chadwick (1932): Discovered the neutron.
Electron configuration — the address system
Electrons fill orbitals according to three rules:
- Aufbau principle — fill the LOWEST energy level first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d.
- Pauli exclusion — no two electrons in the same atom can have the same set of four quantum numbers. In practice: each orbital holds at most 2 electrons, with opposite spins.
- Hund's rule — for orbitals of equal energy (the three p orbitals, for example), fill each with ONE electron (parallel spins) before doubling up.
Sublevel capacities:
- s: 2 electrons (1 orbital)
- p: 6 electrons (3 orbitals)
- d: 10 electrons (5 orbitals)
- f: 14 electrons (7 orbitals)
Example — sulfur (Z = 16):
1s² 2s² 2p⁶ 3s² 3p⁴
Count: 2 + 2 + 6 + 2 + 4 = 16 ✓. Sulfur has 6 valence electrons (3s² + 3p⁴), explaining why it forms 2 bonds in compounds like H₂S and typically gains 2 electrons to become S²⁻.
Ions form by gaining or losing electrons
Metals tend to LOSE electrons (form cations: Na⁺, Mg²⁺, Al³⁺). Nonmetals tend to GAIN electrons (form anions: F⁻, O²⁻, N³⁻). The driving force is usually achieving a noble-gas electron configuration — full valence shell.
Note that ion formation does NOT change the proton or neutron count — only the electron count. Cations are SMALLER than the neutral atom; anions are LARGER.