In general chemistry you learned that having eight electrons (an octet) in the valence shell of the noble gases (group 8A) gives those elements special stability. We also have seen that the reactions of most elements are driven by their desire to attain a noble gas electronic configuration. The alkali metals in group 1A have a single electron in their valence shell and usually give up this electron in reactions to achieve a noble gas electronic configuration. Unlike the alkali metals which tend to give up their valence electron to achieve a noble gas configuration, the halogens in group 7A (with 7 valence electrons) tend to gain an electron in reactions to attain a full octet of electrons. Elements on the right side of the periodic table (nonmetals) react by gaining electrons and elements on the left side of the periodic table (metals) react by losing electrons. The type of bonding that occurs when a nonmetal accepts an electron from a metal is referred to as ionic bonding and the bond is called an ionic bond. This is the type of bonding that occurs in ionic solids like sodium chloride, NaCl.
In the previous section we saw that elements on the left and right sides of the periodic table form ionic bonds by gaining or losing electrons. It may at first glance seem that elements in the middle of the periodic table would form bonds in a similar manner, but that is not the case. Carbon and other elements in the middle of the periodic table form bonds by sharing electrons in covalent bonds. Carbon forms bonds by sharing electrons because it would be very difficult to either gain or lose the four electrons necessary to achieve a noble gas electronic configuration. (Covalent and ionic bonding are extremes in the bonding picture and most bonds have some characteristics of both ionic bonds and covalent bonds.) These shared electron bonds are easily drawn using electron-dot structures or Lewis structures. A Lewis structure for carbon tetrachloride, CCl4, is shown below.
Valence electrons not used for bonding are referred to as lone pair electrons or nonbonding electrons. In Lewis structures all valence electrons including lone pairs are always shown. Notice that each chlorine atom in carbon tetrachloride shown has 3 lone pairs of electrons (nonbonding electrons) associated with it. We will see later that these lone pairs of electrons play a role in determining the chemistry of organo-halogen compounds. We explain the reactions and bonding that occurs in organic molecules by monitoring the valence electrons associated with the atoms. In addition to Lewis structures, molecules can also be drawn using "Kekule" structures (also known as line-bond structures) in which the two-electron covalent bond is shown as a line joining the two atoms. Although nonbonding electrons are usually not shown in "Kekule" structures it is very important to keep track of them since they are the electrons that are involved in organic reactions. A "Kekule" structure of carbon tetrachloride is shown below. Notice that the lone pairs of electrons on the four chlorine atoms are not shown but are understood to be there.
Since the majority of the bonds we will encounter
in organic chemistry are covalent bonds, we need to look at the possible
ways of explaining how covalent bonds are formed. There are two models
which explain covalent bonding, Valence Bond Theory and Molecular Orbital
Theory. In Valence Bond Theory, bonding is viewed as occurring by the overlap
of two atomic orbitals, one from each atom. Each atom's orbital contains
a single electron and a bond is formed by the electrons, now paired in
overlapping orbitals, holding the two nuclei together. Every covalent bond
has a characteristic bond strength and bond length. The bond length is
defined as the distance between the two nuclei. Bond strength is
defined as the amount of energy needed to homolytically break a bond (each
of the nuclei takes one of the two electrons from the bond). A normal
carbon-chlorine bond has a bond length of 1.78 Angstroms (1.78 X 10-10
meters) and a bond strength of (339 kJ/mol or 81 kcal/mol).
In Molecular Orbital Theory, bonding is explained in terms of the mathematical combination of atomic orbitals to form molecular orbitals. The newly formed orbitals are called molecular orbitals because they belong to the entire molecule. The combination of two atomic orbitals leads to two molecular orbitals, a bonding molecular orbital and an antibonding molecular orbital.
1. How many hydrogens would you expect carbon to bond with in methane, CH??
a. 2 b.
3 c. 4
2. Which of the following is the correct Lewis structure
for boron tribromide, BBr3?
3. Which of the labeled bonds in butyllithium (shown below) would
you expect to be the most ionic?
a. 1 b.
2 c. 3 d.
4. Which of the following is a correct line-bond structure for
5. How many nonbonding pairs of electrons are missing in
the "Kekule" structure of dimethylformamide, DMF, shown below?
a. 1 b. 2 c. 3 d. 4