Unit 3: Chemical Bonding

Section 1: An Overview of Bonding

* Introduction
* Bonding In Perspective
* Covalent Bonding
* Ionic Bonding
* Polar Bonds
* Identifying Bonds

Introduction

What holds an atom together? What are the forces that make a liquid remain a
liquid, a solid remain a solid, and a gas a gas? In this unit we will
explore the forces that keep atoms together and the properties of these
bonds.

Bonding In Perspective

The way atoms are bonded can help determine the molecular formula of a
compound from the bonds between those elements. However, a chemical analysis
of a compound is the only sure way of knowing the formula. Bonding does help
explain why certain compounds form, and thereby reinforces the law of
multiple proportions (e.g. H2O, H2O2, and CO, CO2).

The shape of molecules can contribute to the type of bonds they form. If a
molecule has the wrong shape, if may not bond with other molecule.
Similarly, if a key has the wrong shape, it cannot fit into a lock.
Therefore, a good knowledge of how a substance bonds can help explain
macroscopic observations such as the digestion of food, oil in water, and
food production in green plants.

When a bond is formed, or a bond is broken, bond energy is involved. A
strong bond may require more energy to break than a weak bond. This can lead
to strong molecules as in a solid, and to weak molecules as in a gas.
Knowing the nature of the bond, and the bond energy involved, can help you
better characterize the nature of a molecule.

Covalent Bonding

What happens when two atoms approach each other? We know that each atom has
protons with positive charges and electrons with negative charges. As two
atoms approach each other, a proton of one atom attracts an electron of the
other, and vice versa, until the atoms are joined. The attractive forces of
the atoms play an important role on each other (e.g. a postive charge
attracts a negative charge, a negative charge attracts a positive charge).
The result is a chemical bond in which two atoms join together by the above
processes.

As the two atoms draw near each other, their attractive forces pull them
together. In some molecules, the repulsion of the atom's electrons began to
repel each other at a certain distance. At this point, the proton of one
atom has drawn the electron of the other atom as close as possible before
the repulsion forces act. Two atoms of similar properties can form a bond in
such a way that they are actually sharing electrons. This type of bond is
called a covalent bond. The two atoms would count each electron in the pair
as one of their own. For example, in a H2 molecule, each atom would claim to
have two electrons at a given point in time, even though each atom had only
one electron before the bond formed. In essence, a pair of electrons is
being shared between the two atoms involved in the covalent bond.

Ionic Bonding

In the last paragraph, we talked about covalent bonds being formed between
two atoms will similar properties. In most cases, atoms in a molecule may
not have similar properties. One atom may tend to have more of an attractive
pull than other. As in a tug of war, one side would end up with the piece of
rope. In a ionic bond, that piece of rope would be electrons. When two atoms
with unequal properties form, one atom would tend to grab all the electrons
in a molecule, leaving the other with none. The difference in the attractive
properties of the atoms is called electronegativity. An atom with the
greater electronegativity would tend to grab all of the electrons in an
ionic compound. For example, in NaCl, the chlorine (Cl) has a greater
electronegativity than sodium (Na), so chlorine would take all the valence
electrons (outer electons) of sodium in the compound. In the fact, in a
compound with an electronegativity difference of 1.7 or more, the compound
is most likely ionic.

NaCl also forms a crystal due to its ionic nature. In an NaCl molecule, each
sodium atom loses one electron to become a 1+ ion, and each chlorine atom
gains one electron to become a -1 ion. A crystal forms between sodium ions
and chloride ions because each sodium ion draws six chloride ions and each
chloride ion draws six sodium ions. The crystal continues to form releasing
energy as long as equal numbers of sodium and chlorine atoms are present.
The three dimesional bonding of negative and postive ions is called a
crystal lattice structure.

Polar Bonding

In a polar molecule, the electrons are shared unequally. One atom has a
greater electronegativity than the other, resulting in one end of the
molecule having more of the electron pair. The atom with the higher
electronegativity pulls on the electrons more, and thus it has a partial
negative charge. This also means that the other atom, that has the lower
electronegativity, has a slight positive charge. This distrubition of charge
among the poles (ends) of the molecule make up a dipole molecule. For
example in a HCL molecule, hydrogen has a 2.1 electronegativity and chlorine
has a 3.0 electronegativity. The difference in electronegativity is 0.9.
Chlorine would take up a partial negative charge (between 0 and 1-) and
hydrogen would have a partial postive charge (between 0 and 1+).

Water (H2O), also behaves in the same manner. That is why dipole and polar
properties are associated with H2O. Other molecules that exhibit dipole
properties include SO2, and NH3. However, the only sure way to find out if a
compound is polar, is to carry out a chemical analysis of the compound.

Identifying Bonds

In most cases, you can use a table of electronegativity values to determine
the type of bonds that will occur between atoms. A simple of rule of thumb
goes as follows:

1. If the electronegativity difference is between 0 and 0.2, it is
probably covalent
2. If the electronegativity difference is between .2 and 1.7, it is
probably polar
3. If the electronegativity difference is greater than 1.7, it is probably
ionic