Skip to main content
Chemistry LibreTexts

4.3.2: Ionic Compounds and Molecular Orbitals

  • Page ID
    243818
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \) \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)\(\newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\) \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\) \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\) \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\) \( \newcommand{\Span}{\mathrm{span}}\) \(\newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\) \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\) \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\) \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\) \( \newcommand{\Span}{\mathrm{span}}\)

    Ionic interactions lie at one extreme on a spectrum of bonding. In the completely covalent bond (homonuclear diatomics) molecular orbitals are formed by equal-energy atomic orbitals, resulting in electron density evenly distributed over the molecule. In the case of polar bonding (heteronuclear diatomics), atomic orbitals of unequal energies contribute unequally to molecular orbitals, resulting in uneven distribution of electron density across the molecule. In the case of polar bonds, the electron density is shifted toward the more electronegative atom since that atom contributes more to the lowest energy bonding molecular orbitals. Molecular orbital diagrams can be drawn for ionic compounds as if they are extremely polar bonds in which electrons are not only shifted toward, but are transferred completely to the more electronegative atom.

    Example: NaCl

    In NaCl, the sodium \(3s\) orbital (-5.2 eV) is significantly higher in energy than the chlorine valence orbitals. The chlorine \(3s\) and \(3p_z\) orbitals have compatible symmetry, yet only the \(3p_z\) orbital (-13.8 eV) is close enough in energy to interact with the Na \(3s\); still the energy difference is large enough to make bonding weak. The Na \(3s\) orbital combines with Cl \(3p_z\) to form the molecular orbitals labeled \(4sigma\) and \(4\sigma^*\) in Figure \(\PageIndex{1}\). The \(4\sigma\) orbital is weakly bonding, but is very close in energy to the Cl \(3p_z\) orbital, and is mostly Cl-like in character. Notice that all \(sigma\) orbitals look very much like either \(s\) or \(p\) orbitals centered on the Cl atom, while the \(sigma^*\) orbital is centered almost entirely on Na. The lack of molecular orbtials that are distributed significantly over both atoms is consistent with a lack of significant covalent bond character in NaCl. The bonding here is characterized by transfer of one electron from Na to Cl and is almost entirely electrostatic. Such bonding in non-directional, unlike true covalent bonding.

    Screen Shot 2020-08-06 at 10.59.07 AM.png
    Figure \(\PageIndex{1}\): The molecular orbital diagram for sodium chloride. Molecular orbital surfaces calculated using Spartan software indicate almost no covalent nature of bonding. (CC-BY-NC-SA, Kathryn Haas)

    Curated or created by Kathryn Haas


    This page titled 4.3.2: Ionic Compounds and Molecular Orbitals is shared under a not declared license and was authored, remixed, and/or curated by Kathryn Haas.

    • Was this article helpful?