Slides11c

Lecture #11

Text: Section 13.12 mostly

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

Molecular Structure

Lewis Structures

Formal charge and preferred structures

Role of electronegativity in determining formal charges

Exceptions to the Octet Rule

Incomplete octets

Odd electron numbers

Expanded octets (hypervalency)

A complete sequence of analysis would follow steps you are now at least moderately familiar with. A few topics relating to further details about structure need to be addressed next.

"Formal charge" for our purposes (in this course) are to be considered a required part of a complete Lewis structure.

Calculating formal charge using carbon monoxide as a simple example.

In deciding among various arrangements of valence electrons, formal charges serve to indicate which arrangement is probably closest to the best choice. Such an arrangement will be referred to as a "preferred structure."

An additional consideration for deciding upon a preferred Lewis structure in which there are formal charges.

Using the "azide" ion as another illustration of deciding upon a preferred Lewis structure.

Note that Lewis dot structures are an accounting of where valence electrons are likely to be found. The structures are very useful for giving quick information about molecular structure, but do not represent what a rigorous theory tells us about valence electron distributions.

First exception to be discussed is what had previously been mentioned with respect to hydrogen; the lightest elements often do not complete their octets.

Another exception to the octet rule occurs in molecules that have an odd number of electrons. You can't get an even number (eight) working with an odd number.

There is also the possibility of having more than eight electrons around an atom once the d-levels become accessible. This happens in the third row of the periodic table, most significantly for elements with Z greater than 13. (The term "hypervalency" is sometimes used to describe this effect.)

Benzene is the cyclic hydrocarbon C₆H₆. The two Lewis structures drawn here look to be structural isomers. On the left, the chlorines are adjacent to a double bond. On the right, they are adjacent to a single C-C bond. Yet only one isomer exists. Why? The answer lies in the phenomenon called "resonance."

We could estimate the energy involved in the reaction shown by looking at bonds broken and bonds formed during the electrons' rearrangements. (These energies are called reaction heats and will be the subject of further discussion in the next lecture.)

But the value obtained differs significantly fromwhat is actually measured. Why? ..."Resonance!"