Chemical Bonding
by Anthony Carpi, Ph.D.
Though the periodic table has only 118 or so elements, there are obviously more
substances in nature than 118 pure elements. This is because atoms can react with one
another to form new substances called compounds (see our Chemical Reactions module).
Formed when two or more atoms chemically bond together, the resulting compound is
unique both chemically and physically from its parent atoms.
Let's look at an example. The element sodium is a silver-colored metal that reacts so
violently with water that flames are produced when sodium gets wet. The element
chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in
World War I. When chemically bonded together, these two dangerous substances form
the compound sodium chloride, a compound so safe that we eat it every day - common
table salt!
enlarge image
In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are
formed between atoms because electrons from the atoms interact with each other. Lewis
had observed that many elements are most stable when they contain eight electrons in
their valence shell. He suggested that atoms with fewer than eight valence electrons bond
together to share electrons and complete their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that
electrons occupy cube-shaped orbitals), his work established the basis of what is known
today about chemical bonding. We now know that there are two main types of chemical
bonding; ionic bonding and covalent bonding.
Ionic bonding
In ionic bonding, electrons are completely transferred from one atom to another. In the
process of either losing or gaining negatively charged electrons, the reacting atoms form
ions. The oppositely charged ions are attracted to each other by electrostatic forces, which
are the basis of the ionic bond.
For example, during the reaction of sodium with chlorine:
sodium (on the left) loses
its one valence electron to
chlorine (on the right),
resulting in
a positively charged sodium
ion (left) and a negatively
charged chlorine ion (right).
The reaction of sodium with chlorine
Concept simulation - Reenacts the reaction of sodium with chlorine.
(Flash required)
Notice that when sodium loses its one valence electron it gets smaller in size, while
chlorine grows larger when it gains an additional valence electron. This is typical of the
relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms
while negative ions tend to be larger than their parent. After the reaction takes place, the
charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic
bond. Ionic compounds share many features in common:

Ionic bonds form between metals and nonmetals.

In naming simple ionic compounds, the metal is always first, the nonmetal second
(e.g., sodium chloride).

Ionic compounds dissolve easily in water and other polar solvents.

In solution, ionic compounds easily conduct electricity.

Ionic compounds tend to form crystalline solids with high melting temperatures.
This last feature, the fact that ionic compounds are solids, results from the intermolecular
forces (forces between molecules) in ionic solids. If we consider a solid crystal of sodium
chloride, the solid is made up of many positively charged sodium ions (pictured below as
small gray spheres) and an equal number of negatively charged chlorine ions (green
spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are
arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is
attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to
sodium attraction. The concept of a single molecule does not apply to ionic crystals
because the solid exists as one continuous system. Ionic solids form crystals with high
melting points because of the strong forces between neighboring ions.
Sodium Chloride Crystal
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
Cl-1
Na+1
NaCl Crystal Schematic
Covalent bonding
The second major type of atomic bonding occurs when atoms share electrons. As opposed
to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs
when two (or more) elements share electrons. Covalent bonding occurs because the atoms
in the compound have a similar tendency for electrons (generally to gain electrons). This
most commonly occurs when two nonmetals bond together. Because both of the
nonmetals will want to gain electrons, the elements involved will share electrons in an
effort to fill their valence shells. A good example of a covalent bond is that which occurs
between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their
first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom
will "want" to pick up a second electron. In an effort to pick up a second electron,
hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H 2.
Because the hydrogen compound is a combination of equally matched atoms, the atoms
will share each other's single electron, forming one covalent bond. In this way, both atoms
share the stability of a full valence shell.
Covalent bonding between hydrogen atoms
Concept simulation - Recreates covalent bonding between hydrogen atoms.
(Flash required)
Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are
shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules
are not strongly attracted to one another. As a result, covalent molecules move about
freely and tend to exist as liquids or gases at room temperature.
Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent
bond is formed. Some atoms can share multiple pairs of electrons, forming multiple
covalent bonds. For example, oxygen (which has six valence electrons) needs two
electrons to complete its valence shell. When two oxygen atoms form the compound O2,
they share two pairs of electrons, forming two covalent bonds.
Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence
electrons of an atom. The structures are written as the element symbol surrounded by
dots that represent the valence electrons. The Lewis structures for the elements in the first
two periods of the periodic table are shown below.
Lewis Dot Structures
Lewis structures can also be used to show bonding between atoms. The bonding electrons
are placed between the atoms and can be represented by a pair of dots or a dash (each
dash represents one pair of electrons, or one bond). Lewis structures for H 2 and O2 are
shown below.
H2
:
-
H H
H H
or
O2
Polar and nonpolar covalent bonding
There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of
the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule
have an equal attraction (or affinity) for electrons, the bonding electrons are equally
shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of
the same element bond together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between two atoms. Polar
covalent bonding occurs because one atom has a stronger affinity for electrons than the
other (yet not enough to pull the electrons away completely and form an ion). In a polar
covalent bond, the bonding electrons will spend a greater amount of time around the atom
that has the stronger affinity for electrons. A good example of a polar covalent bond is the
hydrogen-oxygen bond in the water molecule.
Water molecules contain two hydrogen atoms (pictured in
red) bonded to one oxygen atom (blue). Oxygen, with six
valence electrons, needs two additional electrons to complete
its valence shell. Each hydrogen contains one electron. Thus
oxygen shares the electrons from two hydrogen atoms to
complete its own valence shell, and in return shares two of
H2O: a water molecule
its own electrons with each hydrogen, completing the H valence shells.
Polar covalent bonding simulated in water
The primary difference between the H-O bond in water and the H-H bond is the degree of
electron sharing. The large oxygen atom has a stronger affinity for electrons than the
small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it
preoccupies their time, and this leads to unequal sharing and the formation of a polar
covalent bond.