Chapter 8: Modelling Atoms and Their Electrons

Beginning with experimental evidence, the chapter establishes that electrons exhibit wave-particle duality, meaning they possess characteristics of both particles and waves simultaneously. The photoelectric effect demonstrates the particle nature of light, while electron diffraction experiments reveal the wave-like behavior of moving electrons. Observations of line emission spectra from excited atoms showed that electrons occupy only discrete energy levels rather than a continuous range, a property termed quantization. The Bohr model initially proposed electrons as circular orbits, but this was superseded by the modern quantum mechanical approach. Schrödinger's equation provides the mathematical foundation, treating electrons as three-dimensional standing waves whose distribution is described by three quantum numbers: principal quantum number n, angular momentum quantum number l, and magnetic quantum number m subscript l. These quantum numbers define orbitals, which are regions of probability density rather than fixed paths. Orbitals are classified by shape as s, p, d, or f types. A fourth quantum number, electron spin, represents an intrinsic property generating a small magnetic field. The ground-state electron configuration of atoms follows three organizing principles: the Pauli exclusion principle restricts each orbital to a maximum of two electrons with opposite spins, the Aufbau principle populates orbitals in order of increasing energy, and Hund's rule minimizes electron repulsion by filling degenerate orbitals singly before pairing. This framework explains why elements in the same periodic group share similar valence electron configurations. The chapter then connects atomic structure to periodic trends through the concept of effective nuclear charge, which represents the net positive charge felt by valence electrons after accounting for shielding effects from inner electrons. Atomic radius decreases across a period due to increasing effective nuclear charge and increases down a group due to additional electron shells. Ionization energy, electronegativity, and electron affinity follow similar periodic patterns, generally increasing across periods and up groups. Ionic properties also reflect effective nuclear charge: cations are smaller than neutral atoms while anions are larger due to electron-electron repulsion.