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Bonding

  

 

Overview of Bonding

  Important Terms:

  • Bonds form because the resulting molecule is more stable than the separated atoms. Atoms achieve maximum stability when their electrons occupy the atomic orbitals of the lowest possible energies. Ions achieve maximum stability when they cluster in ionic solids. Molecules achieve maximum stability when electrons are shared between nuclei.
  • Electron sharing is called covalent bonding
  • The separation distance where the molecule is most stable is known as the bond length
  • The amount of stability at this separation distance is known as the bond energy

    Approaches to Bonding

    There are two ways to think about chemical bonds: localized orbitals and delocalized orbitals. Each description has the following common features:

    1. Each electron is described by a wave function.
    2. No two electrons in a molecule can ever have identical wave functions (Pauli Exclusion Principle).
    3. All electrons have wave functions with the lowest possible energies (Aufbau Principle).
    4. Valence electrons are all that are needed to describe bonding.

  • Valence electrons not involved in chemical bonds are said to be nonbonding electrons and remain in atomic orbitals
  • The extent to which an element attracts bonding electrons is called its electronegativity, chi (c)
  • The greater the difference in electronegativity between two atoms of a molecule, the more polar the bond
  • Electronegativity measures the attraction of a bound electron to an atom to make a polar bond
  • Electron affinity measures the attraction of a free electron to a neutral atom to make an atomic anion
  • Electronegativities increase from the lower left to the upper right of the periodic table
  • Metals generally have low electronegativities and nonmetals have high electronegativities
  • Electronegativity differences (Change in chi) provide a measure of where any particular bond lies on the continuum of bond polarities

    1. When Change in chi (Dc) < 1.6, the compound is classified as polar
    2. When Change in chi (Dc) > 2.0, the compound is classified as ionic
    3. Between 1.6 and 2.0, the compound is ionic if it contains a metal, otherwise it is polar

 

Lewis Structures

  A Lewis structure shows the atoms that are bonded together and the locations withing the molecule of all bonding and nonbonding valence electrons

The Conventions

  1. Only the valence electrons appear in a Lewis structure
  2. A line joining two atoms represents a pair of electrons shared between two atoms
  3. Dots placed next to an atom represent nonbonding electrons

Drawing Lewis Structures

  1. Treat ions separately
  2. Count the valence electrons
  3. Assemble the bonding framework, with two electrons per bond
  4. Place three pairs of nonbonding electrons on each outer atom, except H
  5. Assign the remaining valence electrons to inner atoms
  6. Calculate the formal charge on each atom
  7. Minimize formal charges by shifting electrons to make double bonds and triple bonds

Check out the Drawing Lewis Structures Tutorial for more information.

 

Valence Shell Electron Pair Repulsion (VSEPR) Theory

  The VSEPR theory is based on the following central idea:
    The valence shell electron pairs on an atom repel one another, and these repulsions determine the shapes of polyatomic molecules and ions. From a geometric standpoint, we are most concerned with the repulsions in the valence shell of any inner (or central) atoms.

Around the central atom, we determine electronic geometry by first counting the number of stereoactive sets of electrons in the valence shell according to the following rules:

  1. Each set of electrons that is used to bond another atom to the central atom is counted as one stereoactive set of electrons, whether the bonding is single, double, or triple
  2. Each lone pair of valence electrons on the central atom count as one stereoactive set of electrons

The electronic geometry (or orbital orientation) of the central atom is the resulting arrangement of stereoactive sets of electrons that satisfy the central idea of the VSEPR theory
  Geometry of Stereoactive Sets of Electrons About a Central Atom
  # of Stereoactive Sets Electronic Geometry Angles  
  2 Linear 180°  
  3 Trigonal Planar 120°  
  4 Tetrahedral 109°  
  5 Trigonal Bipyramidal 90°,120°,180°  
  6 Octahedral 90°,180°  

 

Hybridization

 
  • Combining an atom's atomic orbitals to form a special set of directional orbitals is referred to as hybridization. Any hybrid orbital is named from the atomic valence orbitals from which it was constructed.
  • Hybridization is a useful way to describe orbital interactions for inner atoms because inner atoms must form bonds that minimize electron-electron repulsion
  • An atoms coordination number is the number of other atoms to which the atom is bonded in a molecule
  • The steric number (or number of stereoactive sets of electrons) of an inner atom is the sum of its coordination number and the number of its lone pairs

  Steric # Elec. Geom. Hybridization Lone Pairs Mol. Geom. Bond Angles  
  2 Linear sp 0 Linear 180  
  3 Trigonal Planar sp2 0 Trigonal planar 120  
        1 Bent <120  
  4 Tetrahedral sp3 0 Tetrahedral 109  
        1 Trigonal pyramid <109  
        2 Bent <109  
  5 Trigonal Bipyramidal sp3d 0 Trigonal bipyramidal 120, 90  
        1 Seesaw <120,<90  
        2 Tshape <90  
        3 Linear 180  
  6 Octahedral sp3d2 0 Octahedral 90  
        1 Square pyramidal <90  
        2 Square planar 90  

 
  • An atom adopts linear electronic geometry, and its bonding can be represented using sp hybrid orbitals, when its steric number is two. sp hybrid orbitals result from the mixing of an s orbital and one p orbital
  • An atom adopts trigonal planar electronic geometry, and its bonding can be represented using sp2 hybrid orbitals, when its steric number is three. sp2 hybrid orbitals result from the mixing of an s orbital and two of the p orbitals
  • An atom adopts tetrahedral electronic geometry, and its bonding can be represented using sp3 hybrid orbitals, when its steric number is four. sp3 hybrid orbitals result from the mixing of an s orbital and three p orbitals
  • An atom adopts trigonal bipyramidal electronic geometry, and its bonding can be represented using sp3d hybrid orbitals, when its steric number is five. sp3d hybrid orbitals result from the mixing of an s orbital, 3 p orbitals, and one d orbital.

      Unlike the geometries for other steric numbers, the five positions in a trigonal bipyramid are not all equivalent
        Equatorial positions lie on a plane occupying the central atom
        Axial positions lie above or below the equatorial plane
  • An atom adopts octahedral electronic geometry, and its bonding can be represented using sp3d2 hybrid orbitals when its steric number is six. sp3d2 hybrid orbitals result from the mixing of an s orbital, 3 p orbitals, and 2 d orbitals

  • Lone pairs in a molecule cause bond angles to be slightly smaller than predicted for totally symmetrical geometry. A molecule with an asymmetrical distribution of electrons is said to have a dipole moment

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Last Revised : Tuesday, September 9, 1997

Copyright © 1997
Louisiana State University, Department of Chemistry.
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