Did you know that the same elements can be part of a compound that is either deadly or essential to life depending on how those elements are arranged and the bonding between them? While there are only about 118 known elements, these combine through chemical bonds to form the billions of different substances we encounter in everyday life.
Chemical bonding between atoms results in compounds that can be very different from the parent atoms. This module, the second in a series on chemical reactions, describes how atoms gain, lose, or share electrons to form ionic or covalent bonds. The module lists features of ionic and covalent compounds. Lewis dot structures and dipoles are introduced.
Terms you should knowtoggle-menu
reaction = a chemical change that happens when two or more atoms or molecules combine or break apart to form a new substance
solid = a substance that has definite shape; a state of matter that is not a liquid or a gas
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
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
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!
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.
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 valenceelectron to chlorine (on the right),
a positively charged sodium ion (left) and a negatively charged chlorine ion (right).
Notice that when sodium loses its one valenceelectron 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
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.
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.
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 H2. 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.
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 valenceelectrons 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 H2 and O2 are shown below.
Polar and nonpolar covalent bonding
There are, in fact, two subtypes of covalent bonds. The H2molecule 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 valenceelectrons,
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 its own electrons with each hydrogen, completing the H valence shells.
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.
Because the valenceelectrons in the water molecule spend more time
around the oxygen atom than the hydrogen atoms, the oxygen end of the
molecule develops a partial negative charge (because of the negative
charge on the electrons). For the same reason, the hydrogen end
of the molecule develops a partial positive charge. Ions are not
formed; however, the molecule develops
a partial electrical charge across it called a dipole.
The water dipole is represented by the arrow in the pop-up animation
(above) in which the head of the arrow points toward the electron dense
(negative) end of
the dipole and the cross resides near the electron poor (positive)
end of the molecule.
Anthony Carpi, Ph.D. “Chemical Bonding” Visionlearning Vol. CHE-1 (7), 2003.
The scientist builds slowly and with a gross but solid kind of masonry. If dissatisfied with any of his work, even if it be near the very foundations, he can replace that part without damage to the remainder. -G. N. Lewis 1875-1946