Atomic theory, periodic table organization, electron configuration, periodic trends, and isotopes/radioactivity. Covers Dalton through quantum mechanical models, electron shell filling (Aufbau, Hund, Pauli), periodic law and block structure, trend prediction (electronegativity, ionization energy, atomic radius, electron affinity), isotope notation, nuclear stability, and radioactive decay modes. Use when teaching, problem-solving, or reasoning about atomic-level chemistry.
All chemistry rests on atoms. Understanding how atoms are built — their subatomic particles, electron arrangements, and position in the periodic table — is the foundation for predicting chemical behavior. This skill covers atomic models from Dalton to quantum mechanics, electron configuration rules, periodic table organization, periodic trends, and isotopes with radioactivity.
Agent affinity: mendeleev (periodic/inorganic chemistry, primary), curie-m (nuclear/radiochemistry, for isotope and decay topics)
Concept IDs: chem-atomic-structure, chem-periodic-table-organization, chem-periodic-trends, chem-isotopes-radioactivity
| # | Model | Year | Key idea | Limitation |
|---|---|---|---|---|
| 1 | Dalton | 1803 | Indivisible solid spheres | No internal structure |
| 2 | Thomson | 1897 | "Plum pudding" — electrons in positive matrix | No nucleus |
| 3 | Rutherford | 1911 |
| Dense positive nucleus, electrons orbit |
| Classical orbits radiate energy, collapse |
| 4 | Bohr | 1913 | Quantized circular orbits, energy levels | Only works for hydrogen |
| 5 | Quantum mechanical | 1926 | Probability orbitals (Schrodinger equation) | Full model — currently accepted |
Each model was displaced by experimental evidence the previous model could not explain. Rutherford's gold foil experiment demolished Thomson's model. Bohr's model explained hydrogen line spectra but failed for multi-electron atoms. The quantum mechanical model, based on the Schrodinger equation, treats electrons as probability clouds (orbitals) rather than particles on fixed paths.
| Particle | Symbol | Charge | Mass (amu) | Location |
|---|---|---|---|---|
| Proton | p+ | +1 | 1.0073 | Nucleus |
| Neutron | n0 | 0 | 1.0087 | Nucleus |
| Electron | e- | -1 | 0.00055 | Electron cloud |
Atomic number (Z): Number of protons. Defines the element. Carbon is always Z = 6.
Mass number (A): Protons + neutrons. Written as superscript-left: 12-C or carbon-12.
Charge: Protons minus electrons. Neutral atom: protons = electrons.
Each electron in an atom is described by four quantum numbers:
| Quantum number | Symbol | Values | Describes |
|---|---|---|---|
| Principal | n | 1, 2, 3, ... | Energy level (shell) |
| Angular momentum | l | 0 to n-1 | Orbital shape (s, p, d, f) |
| Magnetic | m_l | -l to +l | Orbital orientation |
| Spin | m_s | +1/2 or -1/2 | Electron spin direction |
Orbital shapes: s = spherical (l=0), p = dumbbell (l=1), d = cloverleaf (l=2), f = complex multi-lobed (l=3).
Capacity per subshell: 2(2l + 1). So s holds 2, p holds 6, d holds 10, f holds 14.
Three rules govern how electrons fill orbitals:
1. Aufbau Principle. Electrons fill the lowest-energy orbitals first. The filling order follows increasing (n + l), with lower n breaking ties:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
2. Pauli Exclusion Principle. No two electrons in the same atom can have all four quantum numbers identical. Each orbital holds at most 2 electrons with opposite spins.
3. Hund's Rule. Within a subshell of equal-energy orbitals, electrons occupy empty orbitals first (one per orbital, same spin) before pairing.
Problem. Write the full electron configuration and orbital diagram for iron.
Solution. Fill orbitals in Aufbau order, distributing 26 electrons:
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6
Check: 2 + 2 + 6 + 2 + 6 + 2 + 6 = 26. Correct.
Noble gas shorthand: [Ar] 4s^2 3d^6, where [Ar] = 1s^2 2s^2 2p^6 3s^2 3p^6 (18 electrons).
Orbital diagram for 3d subshell (Hund's rule):