Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110

General Chemistry

Summer 2006

Lecture Notes::Lec 13_15 June

© R. Paselk 2006


VSEPR Theory and Molecular Geometry, cont.

Polarity in Covalent Molecules

Polarity: So now we can predict bonding and shape in representative group molecules (and thus most biomolecules), how about electron density and thus charge distribution? Need two bits of information:


Molecule Geometry Structure Electronegativities Bond Dipoles Molecular Dipole
Carbon monoxide  linear   
ENC= 2.5,
ENO= 3.5
Carbon dioxide  linear  
ENC= 2.5,
ENO= 3.5
   None: two dipoles are of equal magnitude, but opposite in direction and cancel.
 Water  bent


ENH= 2.1,
ENO= 3.5
 Ammonia  trigonal pyramidal  


ENH= 2.1,
ENN= 3.0
Ammonium ion tetrahedral   


ENH= 2.1,
ENN= 3.0
 None: four dipoles are symmetrically arranged to cancel each other out and give a spherically charged but non-polar ion.

A Quantum View of Bonding

With the successes and failures of classical bonding models in mind, let's explore how we might view covalent bonding from a more modern, quantum, point of view.

To do this we will need to look briefly again at atomic orbitals and ask what they can tell us about how atoms might share electrons.

But before we even do that, I want to look at some simple atoms and molecules calculated at the highest level of theory, and thus the best approximation we have of what real atoms and molecules look and behave like. In order to do these calculations, we are assuming our atoms or molecules are in a vacuum, and essentially alone in the Universe. First let's look at ionic bonding using the example of sodium chloride (one of the best "pure" ionic compounds). The images. animations etc. are available in the initial section on ionic bonds in the Bonding Supplement.

We have looked at a quantum model for ionic bond formation, now I want to continue our discussion with a model for covalent bond formation using two well studied diatomic molecules: Cl2 and H2. The animations and images are available in the Bonding Supplement.

In viewing these models we should keep in mind that:

With these thoughts in mind, lets look further at bonding and bond formation.

For both Cl2 and H2 you will note that we have a cylindrical distribution of the electrons in the single bond around the axis between the nuclei. Obviously in both cases the shapes of the orbitals have changed.

In order to understand this change, let's go back and review the shapes and electron distribution of atomic orbitals. The animations and images from this discussion are available at the Atomic Orbital Supplement.

For our discussion of bonding we need to look at s, p, and d orbitals. Higher orbitals are not involved in any of the substances we are interested in in this course.

Electronic Energy Levels Review:

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Last modified 16 June 2006