Chemistry 103A; Sections 5, 6, 7, 8; Lecture 25, 20 Oct 00

Recall:

Coordinate Covalent Bonds

It is possible and common for some single bonds to be formed with both of the electrons in the electron pair coming from the same atom. Such a bond is called a coordinate covalent bond.

Examples of molecules with coordinate covalent bonds are:

O3, SO2, SO3, HNO3, SO42-

There is a very interesting series of acids containing chlorine.

HCl, HClO, HClO2, HClO3, HClO4.

 In most of the above examples we can see that there are pairs of electrons which are not shared. These are called unshared pairs or sometimes lone pairs.
 
 

Rules for Drawing Lewis Dot Structures

One way to draw Lewis dot structures is to write the atoms (with their Lewis dot valence electrons) in the proper arrangement to form the molecule.

Then use circles to show how the electrons are paired to form bonds.

Form all the single bonds first.

Then form any double or triple bonds.

There is an alternate method to draw Lewis dot structures given beginning on page 383 of our text. The book's method works, but it does not make it quite as obvious how the electrons are shared.

In either case we have to know the arrangement of the atoms in the molecule. Practice and experience help us determine how the atoms are arranged, but there are some general rules that are helpful.

Oxygen is rarely bonded to oxygen. Exceptions are O3 and compounds called peroxides (example H2O2 or HOOH).

C forms many many compounds with C bonded to C (C- C or C=C or Cº C bonds).

C, N, P, S are usually "central" atoms.

F is usually a terminal atom.

Cl, Br, I, etc., can be "terminal" atoms (or central atoms when bonded to O).

H is always a terminal atom because there can be only one bond to H.

 Examples: (I will do the book's examples my way to emphasize the sharing of electrons.)
NH3

ClO-

NO2+

PO43-

Some useful generalizations
For neutral compounds:
C always has four bonds

N always has three bonds

O always has two bonds

F and H always have one bond

For negative molecular ions:
 
N(-) will have two bonds

O(-) will have one bond
 

For positive ions:
N(+) will have four bonds

O(+) will have three bonds


Isoelectronic Series

We often find groups of molecules with the same electronic structure. We say that these species are isoelectronic.

Examples:

NO+, N2, CO, CN-

BH4-, CH4, NH4+

NH3, H3O+

CO2, OCN- , SCN- , N2O, NO2-, COS, CS2


Resonance

Several of the Lewis dot structures we have written could have been written in two or more ways.

For example:

SO2, SO3, O3, HNO3, NO3-, CO32-, etc.
One might ask which structure in each case is the correct structure.

The answer is: none of them and all of them.

That means that 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. 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.