Lesson 8.4 Resonance
Suggested Reading
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Zumdahl Chapter 8 Section 8.12
Essential Question
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What is delocalized bonding?
Learning Objective
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List and define the three types of bonding.
Differentiate between the localized and delocalized models of bonding.
Create resonance structures for covalent compounds.
Evaluate the stability of resonance structures using the concept of
formal charge.
Introduction
In the previous lesson you learned that chemical bonds are formed when two
nuclei share a pair of electrons between them. These bonds are thought to
result from the overlap of singly occupied orbitals. The bond is strongest when
the two electrons are confined to a region between the two nuclei. This model
of bonding is called valence bond theory and it is used to describe a localized
bond where the bonding electrons are localized in the region between two
atoms. One of the limitations of valence bond theory is that it assumes all
bonds are localized bonds. As we will see, this is not a valid assumption.
Resonance
Consider the ozone molecule, O3. If you try to write the electron dot formula
for this molecule you will see that you can write two formulas.
In formula 1 the bond is on the right, and in formula 2 it is on the left.
Experiment has shown that both formulas are correct. According to theory,
one of the bonding pairs in ozone is spread over the region of all three atoms
rather than associated with a particular O-O bond. This is called delocalized
bonding.
Resonance is a method within the valence bond theory used to describe the
delocalization of electrons within molecules. It is the use of two or more Lewis
structures to represent a particular molecule. The resonance description of a
molecule having delocalized bonding involved writing all possible electron-dot
formulas. These are called resonance formulas or structures. The actual
electron distribution of the molecule is a composite of the resonance
structures. The structures are usually connected by double arrows to reflect
this.
This notation does not mean that the ozone molecule flips back and forth
between forms. There is only one ozone molecule and it is a composite of
these resonance structures. Thus, these structures are models that only
approximate the true nature of the molecule.
Writing electron-dot formulas for various compounds leads to see that
delocalized bonding exists in many molecules. Whenever you can write
several plausible structures, which often differ by the location of only one
electron, you can expect delocalized bonding.
Watch this YouTube Video
https://www.youtube.com/watch?v=MWDL5WCZBzE
Example: Writing Resonance Formulas
Describe the electron structure of the carbonate ion, CO22-, using electron-dot
formulas.
Solution:
There are three possible electron dot formulas. Because you would expect all
of the C-O bonds to be equivalent, you must describe the electron structure in
resonance terms.
Delocalized Bonding in Metals
Metals, which we will discuss in more detail later in this course, are extreme
examples of delocalized bonding. A sodium metal crystal, for example, can be
regarded as an array of Na+ ions surrounded by a "sea" of electron. The
valence, or bonding electrons, are delocalized over the entire metal crystal.
This is often called the electron sea model of metallic bonding. The freedom of
the electrons to move through out the crystal is responsible for many of
metal's properties such as electrical conductivity and malleability.
Formal Charge
Formal charge is used to determine the best Lewis formula when two or more
structures can be written. Even though some structures may look the same,
the formal charge (FC) may not be. Formal charges are hypothetical charges
that are assigned to each atom in a molecule. We then sum the formal
charges for each atom to obtain the overall formal charge of the resonance
structure. We want to choose the resonance structure with the least formal
charges and/or whose charges add up to zero or the charge of the overall
molecule.
The equation for finding formal charge is:
Formal Charge = (number of valence electrons on free atom) - (number of lonepair electrons) - (
number bond pair electrons)
The FC has to equal the molecule's overall charge. For example, CNS- has
an overall charge of -1, so the Lewis structure's formal charge has to equal
-1.
Example: Use Forma Charge to Determine the Best Lewis Formula for
CNSSolution
Step 1. Find the Lewis Structure of the molecule.
Step 2. Write resonance structures: All elements want an octet, and we can
do that in multiple ways by moving the terminal atom's electrons around.
Step 3. Assign formal charges to each atom in each structure.
FC = (number of valence electrons in free atom) - (number of lone-pair
electrons) - (
number bond pair electrons)
Remember to determine the number of valence electrons each atom has
before assigning FC. In this case C
= 4 v.e., N = 5 v.e., and S =
6 v.e.
Step 4. Find the most ideal resonance structure. (Note: It is the one with
the least formal charges and/or whose charges add up to zero or to the
molecule's overall charge.)
As shown above, it is possible to have more than one ideal structure.
HOMEWORK: Book question pg 385 question 73, practice exercise 8.10