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3: Atoms, Orbitals and Electronic Configurations

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    • 3.1: The Schroedinger equation for the H atom
      Introduces the Full Schroedinger equation for  H-like atoms and the solutions quantum numbers
    • 3.2: Quantum Numbers for Atomic Orbitals
      Schrödinger’s approach uses three quantum numbers (n, l, and ml) to specify any wavefunction. The quantum numbers provide information about the spatial distribution of an electron.
    • 3.3: Atomic Orbitals
      Spatial description of atomic orbitals
    • 3.4: Many-Electron Atoms
      In addition to the three quantum numbers (n, l, ml) dictated by quantum mechanics, a fourth quantum number is required to explain certain properties of atoms. This is the electron spin quantum number (ms), which can have values of +½ or −½ for any electron, corresponding to the two possible orientations of an electron in a magnetic field. This is important for chemistry because the Pauli exclusion principle implies that no orbital can contain more than two electrons (with opposite spin).
    • 3.5: Electron Configurations
      Based on the Pauli principle and a knowledge of orbital energies obtained using hydrogen-like orbitals, it is possible to construct the periodic table by filling up the available orbitals beginning with the lowest-energy orbitals (the aufbau principle), which gives rise to a particular arrangement of electrons for each element (its electron configuration). Hund’s rule says that the lowest-energy arrangement of electrons is the one that places them in degenerate orbitals with parallel spins.
    • 3.6: Electron Configurations and the Periodic Table
      The arrangement of atoms in the periodic table results in blocks corresponding to filling of the ns, np, nd, and nf orbitals to produce the distinctive chemical properties of the elements in the s block, p block, d block, and f block, respectively.

    3: Atoms, Orbitals and Electronic Configurations is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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