Lecture #10
  CURMUDGEON GENERAL'S WARNING. These "slides" represent highlights from lecture and are neither complete nor meant to replace lecture. It is advised not to use these as a reliable means to replace missed lecture material. Do so at risk to healthy academic performance in 09-105.
Lecture Outline Electronegativity

Molecular Structure

Lewis Structures

Ionic compounds

Periodic trends of electronegativity of the elements through the first five rows.
A pictorial schematic of valence electrons for the Main Group (s- and p-block) elements is shown here where each "dot" represents an accounting of a valence electron. The scheme was invented by G. N. Lewis and these are referred to as Lewis electron dot pictures of elements.
Electronegativity, a property of atoms that will become important in our discussion of molecules.
As above, but just for the s- and p-block (main group) elements.
Electronegativity trend of the halogens
The most electronegative elements.
The least electronegative elements.
The compounds formed between very electronegative atoms and very electropositive atoms will invariable be ionic in nature. Lithium fluoride is a prime example of this.
The strength of an ionic bond can be understood as due to a hypothetical assembly of the atoms involved. Ionizing Li and putting the released electron onto a separated fluorine atom involves IE(Li) and EA(F). Since IEs are generally much larger than EAs, this dominates what will be a cost in energy. Energy is the released as the potential energy of electrostatic attraction between the Li+ and F- is taken advantage of by bringing the ions close together (to their eventual bond separation distance).
Since even greater energy would be returned if Li2+ and F2- were brought together, you can convince yourself that the cost in making these two ions is exceptionally high. Look at electron configurateions.
In combining atoms to form molecules, Lewis' Octet Rule accounts for how valence electrons are distributed. Note this is not a theory to explain bonding, but merely a book-keeping scheme for tracking valence electrons consistent with what is observed in molecular structure.  
An example of sharing valence electrons to accommodate the Octet Rule.  
In symbolizing valence electrons, a single line can be used to represent a shared pair of electrons; that is, to represent one bonding pair of electrons.  
More than one pair of electrons can be shared to be consistent with the octet rule.  
Here, there are three shared pairs of electrons (six electrons involved) and also two pairs of electrons not involved in bonding at all, and called "lone pairs".  
These are a small selection of bond energies showing how the bond energy increases with increasing bond multiplicity
These values illustrate how bond lengths decrease with increasing bond order.
More than two atoms: When given a molecular formula for a polyatomic molecule, such as ethane, the very first step in constructing the complete Lewis dot structure is to arrive at a framework, a skeleton, of how the nuclei are linked.  
Here are eleven different skeletal structures for a molecule with the formula shown. Not all will prove to be relevant. Some will not work at all as frameworks to be completed. (Note added, one more linear possibility is missing: that with oxygen at each end.)
Here's a slightly more complicated situation in which the formula, C2H4O2, has a number of different skeletal structures possible. Additional information allows you to draw the framework shown in which the linkages have used 14 valence electrons.
Ten more valence electrons must be placed in the structure. They are placed in the atoms whose octets are not yet filled. Ten such electrons are placed here in blue. Other arrangements are possible, but not shown. The carbon in the C-O has only six electrons and violates the octet rule. A pair of electrons from neighboring oxygen can be moved over and shared with that carbon.
This is the correct Lewis structure for the compound identified. (Alternatively, a pair of electrons from the other oxygen would have satisfied the octet rule as well, but another consideration -- formal charge -- would reject that possibility as less favorable. We discuss formal charge later in the lecture.