Unit 3: Chemical Bonding

Section 2: Molecular Structures

* Introduction
* Electron Dot Structures
* The Octet Rule
* Exceptions to the Octet Rule
* Bond Energy

Introduction

Have you tried picturing the chemical bonds between water moleucles? This
can prove to be difficult. Therefore, scientists have devised ways to
represent chemical bonds on paper. In this section, we will examine some
common structural models that scientists use to describe chemical bonds.
These models can help to better visualize chemical bonds.

Electron Dot Structures

An American chemist, G.N. Lewis, developed a system used to describe
chemical bonds in 1916. This system is called electron dot structures or
Lewis structures. In this system, the symbol of the element represents the
nucleus of the atom. Dots are then placed around the the symbol. The dots
represents the valence electrons (the outermost electrons in the atom that
are used for chemical bonding) of an element. However, only two dots may be
placed on each of the four sides of the symbol (top, left, bottom, right).
This represents the electron pairs. A maximum of eight dots can fit one
symbol. Each element in a family have the same number of dots around it. For
example, sodium (Na) has one electron around it, potassium (K) has one
valence electron (one dot), and the same holds for the remaining elements in
the family.

The Octet Rule

When forming a compound, a metal would like to lose electrons, and a
nonmetal would like to gain electrons. Why is this? Looking at the hydrogen
electron dot structure, we see that hydrogen (H) has one dot (which side the
dot is placed has no significance). When hydrogen combines with fluorine (F)
it forms hydrogen fluoride. Fluorine (F) has seven dots in its electron dot
structure. Since the compound is an polar molecule, we also know that H will
tend to lose its electron, and F will pull on it stronger. According to the
octet rule, nonmetals tend have eight electrons in a chemical bond. This is
due to the fact that nonmetals try to have eight valence electrons like
noble gases (eight dots) which are the most stable elements. To represent HF
in terms of Lewis dot structures, we would write H with a dot on the right
side, and F with a dot on the left side, two dots on the top side, two dots
on the right side, and two dots on the bottom. The common link between the
right side dot on the H atom and the left side dot of the F atom signifies
the chemical bond. The octet rule tells us that this compound is possible.
However, the octet rule does not accurately predict every stable
configuration of all molecules and compounds.

Exceptions to the Octet Rule

The Lewis dot structure can serve as the underlying concept in understanding
the chemical bonds in molecules. However, the octet rule, which governs the
writing of lewis dot structure, has expections.

Not every nonmetal, nor metal, can form compounds in which each element
satisifies the octet rule. For example in BF3, boron has only three valence
electrons which can bind to three fluorine atoms. As a result of covalent
bonding, boron shares three electron pairs with fluorine. Furthermore, boron
ends up with six valence electrons and not with eight electrons, as stated
by the octet rule. Boron is an example as to why the octet rule does not
apply to all metals. As for the nonmetals, exceptions also exist. In SF6,
sulfur forms six bonds with fluorine, resulting in 12 shared electrons (not
eight). Experimental data shows that BF3 and SF6 form stable molecules. Thus
the octet rule should not be the only rule in determining the nature of
bonds and shapes of molecules and compounds.

Bond Energy

Bonds formed between elements have energy. The energy required to break
these bonds is called bond energy. For strong bonds, more energy is
required, and in weak bonds, less energy is required. Similarly, the energy
released when a strong bond is broken is more than the energy released by a
weak bond. To measure the bond energy between atoms and molecules,
scientists rely on greater quantities than single atoms or molecules.
Instead, they use 6.02 x 1023 molecules or atoms. This number is very
important in chemistry, and it represents one mole of a substance. It would
prove much easier to measure 6.02 x 1023 atoms or molecules than one single
small atom or molecule!

To break a bond, energy must be put into the system. Thus, an endothermic
reaction is required to measure the energy of a chemical bond. Experimental
data has shown that it takes 436 kJ (in chemistry, the unit for energy most
commonly used is the joule, abbreviated "J") to break one mole of H2
molecules into H atoms. Thus the bond energy for H2 is 436 kJ/mol. This is
the energy of one mole of H2 molecules. Bonds must also be broken to form
molecules and compounds. If a reaction involves a molecule breaking up to
form a new compound, the bonds that are broken release energy. However,
energy is also required to form the bonds in the new compound.

In a molecule where a triple bond, double bond, or single bond is formed,
the triple bond would be the strongest. As in the bonds between to two
carbon atoms, a triple bond is the strongest of the possible single, double,
and triple bonds that could be formed. A triple bond between two carbon
atoms also satisify the octet rule. Experimental data has also shown that
the atoms are closer together in a triple bond than in a single or double
bond. In experiments with a compound containing a single, double, or triple
carbon bond (i.e. C2H6, C2H4, and C2H2), the compound containing a double
carbon bond (C2H4) was more reactive than the compound containing a single
carbon bond (C2H6), and the triple carbon bond compound (C2H2) was more
reactive than the double carbon bond. The compound with the triple bond
released the most energy (highest bond energy) when reacting with a
different compound because bonds were broken when the reaction was taking
place; the stored energy of the triple carbon bond was more than a single or
double carbon bond.