4: Quantum Mechanics Describes Atomic Structure

CHEM 111 and CHEM 113 (General Chemistry I and II) · LibreTexts

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    • 4.1: Overview and Objectives
      This page discusses the shift from classical physics to quantum mechanics, crucial for comprehending the behavior of particles at the nanoscale. It outlines the limitations of classical physics for small objects and introduces fundamental quantum concepts like wave-particle duality. The chapter also sets objectives for skills such as converting physical properties, comparing atomic models, and predicting quantum behaviors, encouraging a deeper appreciation of small-scale physics complexities.
    • 4.2: Electromagnetic Energy
      Light and other forms of electromagnetic radiation move through a vacuum with a constant speed, c. This radiation shows wavelike behavior, which can be characterized by a frequency, ν, and a wavelength, λ, such that c = λν. Light is an example of a travelling wave. Other important wave phenomena include standing waves, periodic oscillations, and vibrations. Standing waves exhibit quantization, since their wavelengths are limited to discrete integer multiples of some characteristic lengths.
    • 4.3: Coulomb's Law and the Electrostatic Potential
    • 4.4: The Bohr Model
      Bohr incorporated Planck’s and Einstein’s quantization ideas into a model of the hydrogen atom that resolved the paradox of atom stability and discrete spectra. The Bohr model of the hydrogen atom explains the connection between the quantization of photons and the quantized emission from atoms. Bohr described the hydrogen atom in terms of an electron moving in a circular orbit about a nucleus. He postulated that the electron was restricted to certain orbits characterized by discrete energies.
    • 4.5: Development of Quantum Theory
      Macroscopic objects act as particles. Microscopic objects (such as electrons) have properties of both a particle and a wave. but their exact trajectories cannot be determined. The quantum mechanical model of atoms describes the 3D  position of the electron in a probabilistic manner according to a mathematical function called a wavefunction, often denoted as ψ. The squared magnitude of the wavefunction describes the distribution of the probability of finding the electron in a particular region in
    • 4.6: Electronic Structure of Atoms (Electron Configurations)
      The relative energy of the subshells determine the order in which atomic orbitals are filled. Electron configurations and orbital diagrams can be determined by applying the Pauli exclusion principle (no two electrons can have the same set of four quantum numbers) and Hund’s rule (whenever possible, electrons retain unpaired spins in degenerate orbitals). Electrons in the outermost orbitals, called valence electrons, are responsible for most of the chemical behavior of elements.
    • 4.7: Periodic Variations in Element Properties
      Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller.

    This page titled 4: Quantum Mechanics Describes Atomic Structure was last modified on Thu, 13 Aug 2026 01:22:16 GMT and is shared under a not declared license and was authored, remixed, and/or curated by Kathy Davis, Terrie Salupo-Bryant, Mark Bryant and Jeff Osborne.