|CH301H - Principles
of Chemistry I: Honors
Fall 2011, Unique 51040
Lecture Summary, 6 September 2011
|Covalent Bonding Continued: We
wrapped up our discussion of the classical model of covalent bonding by
relating bond length to bond energy, and found a number of trends for
homonuclear and heteronuclear diatomic molecules. In general bond
energy decreases with increasing bond length, but there are a couple
exceptions to this trend that we can't yet explain. We also saw
that the bond length between two atoms is generally the same, even if
those bonds are in very different molecules. This is a remarkable
property that is very useful for predicting the structure and behavior
of molecules we don't know based on the molecules that we do.
We also spent some time talking about polar covalent bonds, in which the 2 bonding electrons are not shared completely equally between the two atoms. This occurs when a bond forms between two atoms of different electronegativies, but not so different that the atoms are completely ionized. A polar covalent bond leads to a distribution of partial charges throughout a molecule, which in turns leads to a dipole moment. We can use the definition of dipole moment to determine the "extent" of ionic character of a bond: 0% means the bond is completely covalent, while 100% means the bond is completely covalent. Most bonds we have looked at fall somewhere in between those two extremes.
Lewis Dot Structures: Today we also did a brief review of the rules for drawing Lewis dot structures. Lewis dot structures are a formalism for determining atomic connectivity in a molecule, which in turn will help us figure out molecular geometry. To draw an accurate Lewis dot structure, make a series of covalent bonds, each containing 2 electrons, in such a way that all atoms have 8 electrons in their filled valence shell (except for H, which will have 2 electrons in its valence shell). Here are my general rules for Lewis dot structures:
1. Hydrogen and halides can only form one bond and are always terminal atoms on a molecule.
2. Write out each atom with its own valence electrons and make an initial guess about the structure. In general, atoms with the fewest valence electrons will be the central atom. Chemists very often write the molecular formula with the central atom listed first, although this is not always true.
3. Start making bonds, either between single electrons on two different atoms, or with both electrons from a single atom. Remeber to either add or remove electrons as needed to achieve the appropriate molecular charge.
4. Assign formal charges. Add up the formal charges to make sure it equals the known molecular charge.
5. If you have multiple reasonable structures, in general, the "correct" structure is the one with the fewest nonzero formal charges, and where formal charges are the lowest (+/-1).
6. Then draw resonance structures.
These rules are useful, but will not guide you through any possible senario. There are a number of exceptions to these rules, the most interesting being expanded octets.