CHEM 1411 Concept Reviews: Covalent Bonding II - Molecular Shapes, VSEPR, & MO Theory

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Bond: A shared pair of electrons between two atoms that holds those atoms together.

Bond Angle:  The angle formed between two adjacent bonds (or the angle formed by the three nuclei that are bonded together by those two adjacent bonds)

Electron Domain:  A region of electron density.  A domain can be either bonding or nonbonding.  A bonding domain can be composed of either a single, double, or triple bond.  (Note that double and triple bonds still only count as ONE electron domain)  A nonbonding domain can be either a lone pair or a free radical (1 unpaired electron).

 

 

VSEPR Theory

The VSEPR (Valence Shell Electron Pair Repulsion) theory states that molecules will arrange themselves in such a way as to maximize the distance between electron domains in each atom.  This allows us to predict electron domain geometry and molecular geometry.  Electron domain geometry is the arrangement of electron domains around the central atom.  Molecular Geometry is the arrangement of ONLY the atoms in a molecule or ion.

 

To predict the shapes of molecules or ions using the VSEPR model, use the following steps:

  1. Draw the Lewis structure of the molecule or ion and count the number of electron domains around the central atom.  (See above definition of Electron Domain)
  2. Predict electron domain geometry by arranging electron domains around the central atom so that the repulsions among them are minimized (i.e. maximum distance between domains).  See table below:
  3. Use the arrangement of bonded atoms to determine the molecular geometry.

 

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Molecular Shape and Polarity

The polarity of a molecule depends on two primary factors, the polarity of its bonds and the arrangements of those bonds (i.e. the molecular shape).  As a reminder from chapter 8, if two bonded atoms have different electronegativities, the bond is polar and contains a dipole moment (a measure of the magnitude and degree of charge separation in a bond or molecule).  If all the bond dipoles (the dipole moment due only to a specific bond) point in geometrically opposite directions, the dipole moments cancel each other out, and there is NO net dipole moment, making the molecule non-polar.  If the bond dipoles are NOT directed in geometrically opposite positions, there will be a net dipole moment, making the molecule polar.  From these principles, the following rules have been developed:

 

Simplified Rules for Determining Molecule Polarity

  1. If a molecule has one lone pair of electrons on the central atom, the molecule is polar. If a molecule has 2 or more lone pairs on the central atom, the molecule will be polar UNLESS those lone pairs are in geometrically opposite positions AND all the atoms bonded to the central atom are the same.
  2. If a molecule has no lone pairs and the atoms bonded to the central atom are all the same, the molecule will be non-polar. If a molecule has no lone pairs and the atoms bonded to the central atom are NOT the same, the molecule will usually be polar.
  3. For any diatomic (2 atom) molecule, the polarity of the bond is the same as the polarity of the molecule.

 

 

 

Bonding and Hybrid Orbitals

According to VSEPR theory, bonding involves the overlap of the valence orbitals of different atoms.  In order to create a more stable molecule, atoms will tend to hybridize its valence orbitals to form orbitals of equal energy according to the number of electron domains around that atom.  This helps to prevent an imbalance in the molecule.  Determining the hybridization of an atom is as easy as following the steps below:

  1. Draw the Lewis structure of the molecule.
  2. Count the number of electron domains around the atom.
  3. Hybridize the number of orbitals equal to the number of electron domains on that atom. See table below:

# of Electron Domains

2

3

4

5

6

Orbitals Hybridized

s, p

s, p, p

s, p, p, p

s, p, p, p, d

s, p, p, p, d, d

Hybridization of Atom

sp

sp2

sp3

sp3d

sp3d2

 

The overlapping of hybridized orbitals (and/or s-orbitals) produces σ bonds (sigma bonds).  The overlapping of p-orbitals produces π bonds (pi bonds).  ALL single bonds are σ bonds.  A double bond contains 1 σ bond and 1 π bond.  A triple bond contains 1 σ bond and 2 π bonds.

 

 

 

Molecular Orbital Theory

The Molecular Orbital Theory is an alternative theory to VSEPR which helps better explain certain experimental observations concerning the effect of bonding on certain properties of molecules, such as magnetism and light absorption.  In this theory, valence orbitals from different atoms hybridize to form molecular hybrid orbitals.  For a good diagram of how these orbitals hybridize, see Figure 10.14 on page 416 of your textbook (Principles of Chemistry: A Molecular Approach, 3rd ed.).  According to this theory, the strength of the bond is determined by the Bond Order, which can be calculated using the following equation:

BondOrder=12(bondingelectrons−antibondingelectrons)