Section 6.1: Covalent Bonds
Objectives:
1. Explain the role and location of electrons in a covalent bond.
2. Describe the change in energy and stability that takes place as a covalent bond forms.
3. Distinguish between nonpoloar and polar covalent bonds based on electronegativity
differences.
4. Compare the physical properties of substances that have different bond types, and relate
bond types to electronegativity differences.
Key Terms:
 covalent bond
 molecular orbital
 bond length
 bond energy



nonpolar covalent bond
polar covalent bond
dipole
Sharing Electrons
When ions form, the electrons are donated between atoms to form stable ions. Oppositely charged ions are then
attracted to each other to form ionic compounds. When covalent compounds form, neutral atoms share
electrons to produce stable compounds. Atoms that are covalently bonded are called molecules. Ionic
compounds are not referred to as molecules.
The simplest example of sharing electrons is found in diatomic molecules, or molecules containing two atoms.
An example is hydrogen gas, H2. The positive nucleus of one hydrogen atom attracts the electron of the other
atom. At the same time, the two atoms’ positive nuclei and electron clouds repel each other. Neither atom will
completely remove the electron from the other atom; instead the two hydrogen atoms share the electrons.
The resulting H2 molecule is more stable than either hydrogen atom by itself. The H2 molecule is stable because
both hydrogen atoms have a shared pair of electrons giving each a full valence shell (filled 1s2 energy level)
similar to the helium atom, a noble gas.
A covalent bond is formed when atoms share one or more pairs of electrons. A molecular orbital is the new
orbital formed where the paired electrons have a high probability of being found. A molecular orbital is a hybrid
combination of the orbitals found in the individual atoms.
Covalent compounds consist of a collection if individual particles that hold atoms that are covalently bonded to
each other. For example, a glass of water is a collection of molecules of the compound H2O. The individual
water molecule consists of two hydrogen and one oxygen atom covalently bonded together. Another example is
the formula C12H22O11 molecule found in sugar. Ionic compounds consist of a collection of individual positive
and negative ions in an orderly array or lattice that is held together by opposite charge.
Energy is release
when atoms form
stable covalent bonds.
The following figure
shows what takes
place as two single
hydrogen atoms bond
to form a diatomic
molecule. In part (3)
of the figure, the two
bonded hydrogen
atoms are at their
lowest potential
energy state. At this
position, the distance
between them is 74
picometers, known as
the bond length.
The energy required to break a bond between tow
atoms is called the bond energy. This table lists the
bond energy of some common bonds. Note how the
bond length decreases as the bond energy increases.
A covalent bond is flexible, and can vibrate back and
forth. As they do, the distance between them
constantly changes. The bond length is the average
distance between the two bonded atoms.
Electronegativity and Covalent Bonding
When the same two atoms are covalently bonded to each other, such as in the H2 molecule, the electron pair is
equally attracted to the nucleus of each atom. In this case, the electron pair is, on average, an equal distance
between the two atoms. The electrons will not be equal distance on average between two covalently bonded
atoms that are not the same. A nonpolar covalent bond is a covalent bond in which the bonding electrons are
equally attracted to both bonded atoms. A polar covalent bond is a covalent bond in which a shared pair of
electrons is held more closely by one of the atoms.
The ability of an atom to attract an electron is measured by its electronegativity. In a covalent bond, the electron
pair will be held closer to the atom whose electronegatity is higher. If the difference in electronegativity
between the two atoms in a bond is great enough, the atom with the higher value may completely remove an
electron from the other atom and an ionic bond will form.
Using the electronegativities, we can predict whether a bond will be non-polar covalent, polar covalent, or ionic.
A non-polar covalent bond is a
covalent bond in which the bonding
electrons are equally attracted to
both bonded atoms.
A polar covalent bond is a covalent
bond in which a shared pair of
electrons sis held more closely by
the atom with the higher
electronegativity.
Polar molecules have positive and negative ends. A molecule in which one end has a partial positive charge and
the other end has a parial negative charge is called a dipole. A partial positive and partial negative charge is
written using the Greek symbol delta (δ+ or δ-). The bonding electrons are held more closely to the atom that has
the partial negative charge.
The greater the difference between the electronegativity values of the two bonded elements, the greater the
polarity of the bond. In addition, greater electronegativity differences tend to be associated with stronger bonds.
Therefore, the magnitude of the polarity is related to bond strength.
Molecule
H—F
H—Cl
H—Br
H—I
Electronegativity Difference
1.8
1.0
0.8
0.5
Bond Energy
570 kJ/mol
432 kJ/mol
366 kJ/mol
298 kJ/mol
The properties of substance depend on the type of bond holding it together.
Ionic
What occupies the
lattice points in the
crystal?
What is the
strongest force
binding them in the
lattice?
Hard or soft?
Brittle or malleable?
High or low melting
point?
Good conductor?
Solubility
Examples
cations and ions
the ionic bond
hard
brittle
high (usually 3001000oC
no (unless melted)
often soluble in
water; usually
insoluble in
nonpolar solvents
NaCl, MgSO4
Covalent
(molecular)
individual
molecules
Macromolecular
(covalent network)
atoms covalently
bonded to another
Metallic
Van der Waals
forces
(intermolecular
attractions)
soft
crumbly
low (usually under
300oC)
no
polar substances
soluble in polar
solvents, nonpolar
in nonpolar solvents
Non-metal
substances
the covalent bond
metal cations (the
valence electrons
are delocalized)
the metallic bond
very hard
very brittle
very high (usually
over 1000oC)
No (insulators)
insoluble
variable
malleable
Variable (-39oC Hg;
3415oC W)
excellent
insoluble
diamond,
pure metals, alloys
gemstones, ceramics
In ionic substances, the overall attraction between all the cations and anions is very strong. Each ion is held in
place in a lattice structure of many oppositely charged neighbors. The forces holding them together are very
strong and hard to break.
In molecular substances the molecules are held together by sharing electrons. The shared electrons are attracted
to the two bonding atoms. They have little attraction to atoms of other nearby molecules except through weak
forces called Van der Waals forces.