Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2015

Lecture Notes 29: 8 April

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Molecular Geometry, cont.

VSEPR (Valence Shell Electron Pair Repulsion) Theory

Based on three assumptions (there are more advanced versions, but unnecessary for us):

VSEPR predicts geometry based on these assumptions in a few simple, sequential, steps:

  1. Draw a correct Lewis Structure.
  2. Determine the Steric Number = the number of bonded atoms + the number of lone pairs.
  3. Maximize the angles between electron pairs, placing the lone (unbonded) pairs at the extremes.

For central atoms with eight outer electrons (octets) there are three possible electron pair geometries:

  1. Linear with angles of 180° ( a single pair and a triple bond, or two double bonds).
  2. Trigonal planar with angles of 120° (one double bond and two single pairs).

    image of trigonal planar structure
  3. Tetrahedral with angles of 109.5° (four single pairs). [model]

    drawing of tetrahedron

These three electron pair geometries can lead to five molecular geometries:

  • Trigonal planar (formaldehyde, CH2O)
    • formaldehyde, CH2O
      • valence electrons = 4 + 2x1 + 6 = 12
      • 6y + 2 = 6 x 2 + 2 = 14; so molecule has 2 fewer electrons than required for all single bonds, 1 double bond
      • LS: O can only have two bonds, so C will be central atom, therefore = formaldehyde Lewis Structure
      • Considering C as the central atom, have 3 bonded atoms and no lone-pairs, therefore
      • steric number = 3, so trigonal planar electronic geometry, and 3 atoms so
      • trigonal planar molecular geometry ball and stick model of formaldehyde and a model showing single vs. double bonds: ball and stick model of formaldehyde
  • Tetrahedral (methane, CH4) [model]
      • valence electrons = 4 + 4x1= 8
      • four bonds possible, since only 4 pairs, single bonds because only have H's bound to C.
      • LS: from symmetry C will be central atom, therefore = methane Lewis Structure
      • Considering C as the central atom, have 4 bonded atoms and no lone-pairs, therefore
      • steric number = 4, so tetrahedral electronic geometry, and 4 atoms so
      • tetrahedral molecular geometry, ball and stick model of methane viewd in plane of HCH and rotated for a different view: ball and stick model of methane - offangle view
  • Trigonal pyramidal (ammonia, NH3) [model]
      • valence electrons = 5 + 3x1= 8
      • only 4 pairs, single bonds because only have H's bound to N, 3 bonds, since only 3 H's
      • LS: from symmetry N will be central atom, therefore = ammonia Lewis Structure
      • Considering N as the central atom, have 3 bonded atoms and one lone-pair, therefore
      • steric number = 4, so tetrahedral electronic geometry, but only 3 atoms so
      • trigonal pyramidal molecular geometry ball and stick model of ammonia - side view and rotated to view molecule from below: ball and stick model of ammonia - view from below
  • Bent (water, H2O) [model]
      • valence electrons = 6 + 2x1= 8
      • only 4 pairs, single bonds because only have H's bound to O, 2 bonds, since only 2 H's
      • LS: from symmetry O will be central atom, therefore = water Lewis Structure
      • Considering O as the central atom, have 2 bonded atoms and 2 lone-pairs, therefore
      • steric number = 4, so tetrahedral electronic geometry, but only 2 atoms so
      • bent molecular geometry ball and stick model of water

Exceptions to the "Octet Rule.

There are two categories of exceptions to the octet rule among representative elements:

  1. Small atoms in Periods I & II which cannot accomodate or normally don't accomodate a full set of p-electrons. Thus hydrogen and helium have no ground state p-orbitals and so can only carry a maximum of 2 electrons (a duet). The first three atoms of Period II form a transition between hydrogen (duet) and carbon (octet). Thus lithium, because of its very small size, often forms covalent rather than ionic compounds, with a single bond (e.g. LiH). Similarly beryllium can form compounds with a quartet of electrons around it (e.g. BeH2) and boron can accomodate a sextet (e.g. BH3).
  2. Atoms in the p-block of Periods 3 and higher can have "expanded valence shells" with 10 or 12 electrons in the outermost shell by using some of their "empty" d-orbitals to hold the extra electrons.

To help determine if the octet rule is followed recall Clark's Method (abbreviated) for determining bonding in covalent Lewis Structures:

      • Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge)
        • If Sigma e- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only.
        • If Sigma e- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = Delta/2 (remember a triple bond is 2 multiple bonds!). However, note the exceptions below with small atoms (H, Li, Be, and B).
        • If Sigma e- > 6y + 2 then have an expanded valence shell. Note that if Delta= 2, then pentavalent (10 electrons in the valence shell) , and if Delta= 4, then hexavalent (12 electrons in the valence shell).
      • If you can draw more than one structure, then chose the most symmetrical.
        • If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.

Expanded valence shells

Representative atoms with empty d-shells can also have what are sometimes referred to as expanded valence shells. In these cases the d-orbitals also participate in bonding enabling more bonds to be formed. Two additional electronic geometries are possible:

These two electron pair geometries can lead to six new molecular geometries in addition to another way to make a linear molecule. Last time looked at first two, let's continue today with the remaining geometries:

Trigonal bipyramidal electronic geometry (sn = 5)

  1. Seesaw with angles of 90° & 120° (SF4, one lone pair)

    ball and stick model of sulfur tetrafluoride

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© R A Paselk

Last modified 8 April 2015