Chemical Bonding

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Chemical Bonding
by: Anthony Carpi, Ph.D.
U. of North Carolina
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While there are only 118 or so elements listed on the periodic table,
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. 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!
Cornell U.
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Biography
Gilbert N. Lewis
Classics
The Atom and the
Molecule
G. N. Lewis - 1916 paper
from the American
Chemical Society
+
sodium metal
chlorine gas
table salt
Research
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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 contained 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
database allows
searching of the scholarly
literature.
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 18751946
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 non-metals,
In naming simple ionic compounds, the metal is always first,
the non-metal second (ie. 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 at right. 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 becomes blurred in ionic crystals as the solid exists
as one, continuous system. Forces between molecules are
comparable to the forces within the molecule, and ionic compounds
tend to form crystal solids with high melting points as a result.
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
Sodium Chloride Crystal 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 non-metals bond together. Because both of the non-metals 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 others 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 H2 and O2 are
shown below.
H2
H :H
H -H
or
O2
Polar and Non-Polar Covalent Bonding
There are, in fact, two sub-types of covalent bonds. The H2
molecule is a good example of the first type of covalent bond, the
non-polar 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 non-polar covalent bond is
formed. Whenever two atoms of the same element bond together, a
non-polar 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 6 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 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 HH 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.
The Dipole
Because the valence electrons 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.
Related Modules
• Chemical Reactions
• Chemical Equations
Resources
Chemical Bonding Quiz
Visionlearning - an interactive practice exercise.
Quizzes are only available for registered
members of Visionlearning. If you are already a
member log in to access this feature.
Chemical Bonding Quiz
Further Exploration
Ionic, Covalent & Polar Bonding
P. Young, U. Illinois - Lesson containing a brief
summary of chemical bonding.
Lewis Structures
B. Wojciechowski, P. Cerpovicz, Georgia Southern
U. - Concise lesson on writing and using Lewis
structures.
R. Schneider, Quia.com - An interactive quiz on
chemical bonding and compounds.
Visionlearning Glossary
An alphabetical glossary of relevant scientific
terms.
Anthony Carpi, Ph.D. "Chemical Bonding," Visionlearning Vol. CHE-1 (7), 2003.
http://www.visionlearning.com/library/module_viewer.php?mid=55
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