Chemistry 103A; Sections 5, 6, 7, 8; Lecture 26, 23 Oct 00

Recall:

Resonance

We have seen that some of the Lewis dot structures we have written can be written in two or more ways.

For example:

SO2, SO3, O3, HNO3, NO3-, CO32-, etc.


None of the structures is correct by itself, but that the correct structure is a "mixture" of all of them.

This process of "mixing" structures is called resonance.

(The concept of resonance as applied to molecular structure was first proposed by Linus Pauling. For some reason resonance was determined to be in conflict with Marxist/Leninist principles by the Russian authorities so that Pauling's books were banned from the Soviet Union for many years.) The mixture is sometimes said to be a "linear combination" of all the possible structures.

The official name for this linear combination or mixtures is: resonance hybrid.

The individual structures which combine to make the resonance hybrid are referred to as "contributing structures."

Resonance confers extra stability on a molecule. This is because electrons don't like to be confined so the more room you give an electron to move around the more stability you gain.

The rules are:

You can't move nuclei to make a resonance form.

You can only move electrons.

All of the contributing forms must be valid Lewis dot structures.

Probably the most important molecule that is resonance stabilized is benzene, C6H6.

Exceptions to the Octet Rule

1. We have already seen that H, Li, and Be do not obey the octet rule because they can gain stability with the helium, 1s2 configuration.

Boron forms some compounds in which there are only six electrons around the B atom.

For example, BF3, and BCl3. BH3 does not exist as a stable compound, but B2H6 does exist. However, it does not have the C2H6 structure. Rather B2H6 has a very interesting bridged structure.

BF3 forms a compound with NH3 with a coordinate covalent bond using the line pair electrons on ammonia.

 2. We have already stated that as you go down a column in the periodic table reactivity becomes more varied. One manifestation of this increased variety of reactivity is Si, P and S (and elements below them in the periodic table) can have expanded octets. Some compounds of Si, P and S can have ten or twelve electrons in their "expanded octet."

Examples:

SiF51-and SiF62-

PF5 and PF61-

SF4 and SF6

ClF3 and BrF5

We have mentioned that the noble gases are very unreactive and the compounds of the noble gases have only been prepared relatively recently. Some of the first compounds of the noble gases to be prepared were XeF2 and XeF4.

 An elementary explanation for the ability of the heavier elements to expand their octets is that there are empty s and d orbitals lying just above the usual filled octet. These orbitals are close enough in energy to the filled octet orbitals that they can be involved in bonding.
 

3. There are some molecules which have an odd number of electrons available from the valence electrons of their constituent atoms. For these molecules it is impossible for every atom to have a complete octet.

The most common examples are: NO and NO2.

NO2 will "dimerize" at low temperatures to give N2O4 which does follow the octet rule.

NO does not dimerize. It is stable alone. NO is an example of a free radical. Free radicals have an unpaired electron. Most free radicals are unstable and will react to pair up all the electrons. NO is stable as a free radical.

NO has recently been shown to be important in brain function as a neurotransmitter. Science (the Journal of the American Association for the Advancement of Science) named NO the "Molecule of the Year" for 1999 in honor of its newly discovered brain function activity.
  Some Properties of Bonds

We will discuss three properties of chemical bonds in this course, bond order, bond length, and bond energy.

Bond Order

Bond order is a measure of how many pairs of electrons are shared between two atoms to bond those two atoms together. Recall that a chemical bond is a sharing of electron pairs between two atoms.

For molecules without resonance the bond order is easy.

A single bond has a bond order of 1 because there is one pair of electrons being shared between the two atoms.

A double bond has a bond order of 2 because the two atoms share two pairs of electrons.

Likewise, in a triple bond the bond order is 3 because three pairs of electrons are shared between the two atoms bound together.

Examples:

 
H2O

CO2

C2H2

 Things get interesting in molecules with resonance. To determine the bond order when there is resonance we must recognize that some of the electron pairs are split between more than one pair of atoms. For example, O3 is a resonance hybrid molecule.

Each O-O pair has a single bond and then there is a second shared pair of electrons that resonates between the two pairs of atoms.

That is, the second half of the bond in the double bond is split between the two O-O pairs.

So the bond order of bonds between each pair is 1.5. The 1 comes from the single bond and the 0.5 comes from the pair involved in the resonance.

Examples:
 

SO2,

SO3.

 In more complicated cases we can calculate the bond order by a formula.

If we let natom pairs = the number of atom pairs of a specific type involved in the reonance, and nelectron pairs = the number of pairs of electrons involved in the resonance then the bond order is nelectron pairs/natom pairs.

Examples:

NO3-

C6H6.


Bond Length

The bond length is defined as the distance between the nuclei of the two atoms involved in the bond.

There are three effects that govern bond lengths:

Bonds with higher bond orders tend to be shorter.

So, for example, a CºC bond is shorter than a C=C bond which is shorter than a C-C bond.   Likewise, a carbon-carbon bond in benzene (C6H6) is shorter than an unhybridized carbon-carbon bond, but longer than an unhybridized C=C bond. Bonds lengths tend to increase as you go down a column in the periodic table. That is, bonds between larger atoms tend to be longer. An S-H bond would be expected to be longer than an O-H bond because S is a larger atom than O.   Likewise, the H-I bond is longer than the H-Br bond which is longer than the H-Cl bond which in turn is longer than the H-F bond. Bond lengths tend to decrease as you go to the right in a row in the periodic table. C-H bonds are longer than N-H bonds which are longer than O-H bonds which are longer than F-H bonds. (There are two tables of average bond lengths on page 401 of your text to help you check these rules. I will not ask you to compare bond lengths between the two tables, but only within a table.)

Bond Energy

The bond energy is the energy required to break a chemical bond. That is, it is the energy required to split the molecule into two parts at the bond in question.

Breaking chemical bonds is an endothermic process. So DH for breaking bonds is positive. (Making a bond is the reverse of breaking a bond. Making bonds is an exothermic process, with a negative DH.) Bond energies are usually given for a mole of bonds.

We will represent the bond energy (per mole) by the symbol, D (for bond Dissociation energy). Since bond dissociation energies are endothermic, all Ds are positive.

We can use bond energies to make estimates of DHs of chemical reactions.

DHreaction = SDbonds broken - SDbonds formed.
Example:
N2(g) + 3 H2(g) ® 2 NH3(g).
The bond energies are (from the table on p 403 of our text),
DNºN = 945 kJ/mol

DH-H = 436 kJ/mol

DN-H = 391 kJ/mol.