Lecture 6: Molecular Shape

Read section 1.19, 1.20 from your textbook. By Sunday, January 28 at 10 PM, you should take a quiz (Quiz 3) on this material.

VSEPR

Single crystal X-ray diffraction or neutron diffraction of molecules in the solid state and electron diffraction of molecules in the gas phase can show us the bond distances and angles. We use Lewis structures along with Valence Shell Electron Pair Repulsion Theory to predict the structures of molecules. The idea behind this is that electrons in filled orbitals will repel each other because they have the same charge (just as magnets with the same polarity repel).
  1. All paires of electrons, both bonding pairs and lone pairs, are important in determining the shape of a molecule.

  2. Bonding pairs are smaller than lone pairs because there are 2 positively charged nuclei pulling them in.

  3. Single bonds are smaller than double bonds and double bonds are smaller than triple bonds.

  4. If a central atom (A) is surrounded by different atoms (B and C) in the molecule ABxCy, the relative sizes of B and C can affect the structure of the molecule.
The first step is to construct the best Lewis structure of the molecule. Let's look at a few examples: CH4, NH3, BH3



The electron pairs on the central atom will be arranged in such a way as to maximize their distance to the others. Two pairs will always be 180 degrees apart, in a linear arrangement. Three pairs will be 120 degrees apart in a trigonal arrangement. Four pairs will be arranged in a tetrahedron, 109 degrees apart. When there are 5 pairs of electrons, there are two possible arrangements: trigonal bipyramidal (90 and 120 degree angles) and square pyramidal (90 degree angles). Trigonal bipyramidal is the lowest energy, but the square pyramidal structure is pretty close and is also important. When there are 6 pairs of electrons, they occupy the vertices of an octahedron (90 degree angles).



Methane and ammonia both have 4 electron pairs, arranged in a tetrahedron. Only three of those pairs are bonded to another atom in ammonia. Borane has 3 electron pairs and must be trigonal.



Coordination Geometry

Both bonding and non-bonding electron pairs determine the structure but we name the geometry of molecules according to the arrangement of atoms.

Electron Pairs 0 lone pairs 1 lone pair 2 lone pairs 3 lone pairs
2 e- pairs linear

linear
3 e- pairs trigonal

bent

linear
4 e- pairs tetrahedral

trigonal pyramidal

bent

linear

5 e- pairs trigonal bipyramidal

disphenoidal

T-shaped

linear

6 e- pairs octahedral

square pyramidal

square planar

T-shaped



The true bond angles will usually be distorted from the idealized angles in the pictures above because all bonds and non-bonding electron pairs don't have the same "size".

lone pairs > triple bond > double bond > single bond


Also, atoms that are bonded to a central atom make a difference. The I atoms are much larger than the H atoms in CH2I2 and the H-H angle is smaller than the ideal 109 deg while the I-I angle is larger.

Isomers

Isomers are molecules that have the same formula but different structures. There are a number of different categories of isomers:
  1. Ionization isomers
    These are complex salts where the elements are the same but the identity of the cation/anion pair varies. For example: [Co(NH3)5Br][SO4] and [Co(NH3)5SO4][Br]

  2. Hydration isomers
    These are similar to ionization isomers but differ by the number of water molecules bound to the cation.

  3. Linkage isomers
    These differ in the order of elements. Acetonitrile and methyl isonitrile are linkage isomers.

  4. Geometric isomers
    These differ in the arrangement of groups about a central atom.

  5. Optical isomers
    Optical isomers differ in stereochemistry. Enantiomers are examples of stereoisomers.