The
Covalent Bond
Valence Electrons
The electrons in the outermost shell are the valence
electrons
the electrons on an atom that can be gained or
lost in a chemical reaction. Since filled d or f
subshells are seldom disturbed in a chemical reaction, we can
define valence electrons as follows: The electrons on an
atom that are not present in the previous rare gas, ignoring
filled d or f subshells.
Gallium has the following electron
configuration.
Ga: [Ar] 4s2
3d10 4p1
The 4s and 4p electrons can be lost in a
chemical reaction, but not the electrons in the filled 3d
subshell. Gallium therefore has three valence electrons.

The Covalent Bond
Atoms can combine to achieve an octet of valence electrons
by sharing electrons. Two fluorine atoms, for example, can
form a stable F2 molecule in which each atom has
an octet of valence electrons by sharing a pair of electrons.

A pair of oxygen atoms can form an O2 molecule
in which each atom has a total of eight valence electrons by
sharing two pairs of electrons.
The term covalent bond is used to describe the
bonds in compounds that result from the sharing of one or
more pairs of electrons.

How Sharing of
Electrons Bonds Atoms
To understand how sharing a pair of electrons can hold
atoms together, let's look at the simplest covalent bond
the bond that
forms when two isolated hydrogen atoms come together to form
an H2 molecule.
H · + · H
H-H
An isolated hydrogen atom contains one proton and one
electron held together by the force of attraction between
oppositely charged particles. The magnitude of this force is
equal to the product of the charge on the electron (qe)
times the charge on the proton (qp)
divided by the square of the distance between these particles
(r2).

When a pair of isolated hydrogen atoms are brought
together, two new forces of attraction appear because of the
attraction between the electron on one atom and the proton on
the other.

But two forces of repulsion are also created because the
two negatively charged electrons repel each other, as do the
two positively charged protons.

It might seem that the two new repulsive forces would
balance the two new attractive forces. If this happened, the
H2 molecule would be no more stable than a pair of
isolated hydrogen atoms. But there are ways in which the
forces of repulsion can be minimized. As we have seen,
electrons behave as if they were tops spinning on an axis.
Just as there are two ways in which a top can spin, there are
two possible states for the spin of an electron: s = +1/2
and s = -1/2. When electrons are
paired so that they have opposite spins, the force of
repulsion between these electrons is minimized.
The force of repulsion between the protons can be
minimized by placing the pair of electrons between the two
nuclei. The distance between the electron on one atom and the
nucleus of the other is now smaller than the distance between
the two nuclei. As a result, the force of attraction between
each electron and the nucleus of the other atom is larger
than the force of repulsion between the two nuclei, as long
as the nuclei are not brought too close together.
The net result of pairing the electrons and placing them
between the two nuclei is a system that is more stable than a
pair of isolated atoms if the nuclei are close enough
together to share the pair of electrons, but not so close
that repulsion between the nuclei becomes too large. The
hydrogen atoms in an H2 molecule are therefore
held together (or bonded) by the sharing of a pair of
electrons and this bond is the strongest when the distance
between the two nuclei is about 0.074 nm.

Similarities and
Differences Between Ionic and Covalent Compounds
There is a significant difference between the physical
properties of NaCl and Cl2, as shown in the table
below, which results from the difference between the ionic
bonds in NaCl and the covalent bonds in Cl2.
Some Physical Properties of NaCl and
Cl2
|
|
NaCl |
|
Cl2
|
Phase at room
temperature |
|
Solid |
|
Gas |
Density |
|
2.165 g/cm3
|
|
0.003214 g/cm3
|
Melting point |
|
801°C
|
|
-100.98°C |
Boiling point |
|
1413°C |
|
-34.6°C |
Ability of aqueous
solution to conduct electricity |
|
Conducts |
|
Does not conduct
|
Each Na+ ion in NaCl is
surrounded by six Cl- ions, and vice versa, as
shown in the figure below. Removing an ion from this compound
therefore involves breaking at least six bonds. Some of these
bonds would have to be broken to melt NaCl, and they would
all have to be broken to boil this compound. As a result,
ionic compounds such as NaCl tend to have high melting points
and boiling points. Ionic compounds are therefore solids at
room temperature.

Cl2 consists of molecules in which one atom is
tightly bound to another, as shown in the figure above. The
covalent bonds within these molecules are at least as strong
as an ionic bond, but we don't have to break these covalent
bonds to separate one Cl2 molecule from another.
As a result, it is much easier to melt Cl2 to form
a liquid or boil it to form a gas, and Cl2 is a
gas at room temperature.
The difference between ionic and covalent bonds also
explains why aqueous solutions of ionic compounds conduct
electricity, while aqueous solutions of covalent compounds do
not. When a salt dissolves in water, the ions are released
into solution.
|
H2O |
|
NaCl(s) |
 |
Na+(aq) + Cl-(aq) |
These ions can flow through the solution,
producing an electric current that completes the circuit.
When a covalent compound dissolves in water, neutral
molecules are released into the solution, which cannot carry
an electric current.
|
H2O |
|
C12H22O11(s)
|
|
C12H22O11(aq) |
When two chlorine atoms come together to
form a covalent bond, each atom contributes one electron to
form a pair of electrons shared equally by the two atoms, as
shown in the figure below. When a sodium atom combines with a
chlorine atom to form an ionic bond, each atom still
contributes one electron to form a pair of electrons, but
this pair of electrons is not shared by the two atoms. The
electrons spend most of their time on the chlorine atom.

Ionic and covalent bonds differ in the extent to which a
pair of electrons is shared by the atoms that form the bond.
When one of the atoms is much better at drawing electrons
toward itself than the other, the bond is ionic. When
the atoms are approximately equal in their ability to draw
electrons toward themselves, the atoms share the pair of
electrons more or less equally, and the bond is covalent.
As a rule of thumb, metals often react with nonmetals to form
ionic compounds or salts, and nonmetals combine with other
nonmetals to form covalent compounds. This rule of thumb is
useful, but it is also naive, for two reasons.
- The only way to tell whether a compound is ionic or
covalent is to measure the relative ability of the
atoms to draw electrons in a bond toward themselves.
- Any attempt to divide compounds into just two classes
(ionic and covalent) is doomed to failure because the
bonding in many compounds falls between these two
extremes.
The first limitation is the basis of the concept of
electronegativity. The second serves as the basis for the
concept of polarity.

Electronegativity
The relative ability of an atom to draw electrons in a
bond toward itself is called the electronegativity of the
atom. Atoms with large electronegativities (such as F and O)
attract the electrons in a bond better than those that have
small electronegativities (such as Na and Mg). The
electronegativities of the main group elements are given in
the figure below.

When the magnitude of the electronegativities of the main
group elements is added to the periodic table as a third
axis, we get the results shown in the figure below.

There are several clear patterns in the data in the above
two figures.
- Electronegativity increases in a regular fashion from
left to right across a row of the periodic table.
- Electronegativity decreases down a column of the
periodic table.

Using
Electronegativity to Identify Ionic, Covalent, and Polar
Covalent Compounds
When the difference between the electronegativities of the
elements in a compound is relatively large, the compound is
best classified as ionic.
Example: NaCl, LiF, and SrBr2 are good examples
of ionic compounds. In each case, the electronegativity of
the nonmetal is at least two units larger than that of the
metal.
NaCl |
|
|
LiF |
|
|
SrBr2 |
|
|
Cl |
EN = 3.16 |
|
F |
EN = 3.98 |
|
Br |
EN = 2.96 |
|
Na |
EN = 0.93 |
|
Li |
EN = 0.98 |
|
Sr |
EN = 0.95 |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
EN = 2.23 |
|
|
EN = 3.00 |
|
|
EN = 2.01 |
|
We can therefore assume a net transfer of
electrons from the metal to the nonmetal to form positive and
negative ions and write the Lewis structures of these
compounds as shown in in the figure below.

These compounds all have high melting points and boiling
points, as might be expected for ionic compounds.
|
NaCl |
|
LiF |
|
SrBr2 |
|
MP |
801oC |
|
846oC |
|
657oC |
|
BP |
1413oC |
|
1717oC |
|
2146oC |
|
They also dissolve in water to give aqueous
solutions that conduct electricity, as would be expected.
When the electronegativities of the elements in a compound
are about the same, the atoms share electrons, and the
substance is covalent.
Example: Examples of of covalent compounds include methane
(CH4), nitrogen dioxide (NO2), and
sulfur dioxide (SO2).
CH4 |
|
|
NO2 |
|
|
SO2 |
|
|
|
C |
EN = 2.55 |
|
O |
EN = 3.44 |
|
O |
EN = 3.44 |
|
|
H |
EN = 2.20 |
|
N |
EN = 3.04 |
|
S |
EN = 2.58 |
|
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
|
EN = 0.35 |
|
|
EN = 0.40 |
|
|
EN = 0.86 |
|
|
These compounds have relatively low melting
points and boiling points, as might be expected for covalent
compounds, and they are all gases at room temperature.
|
CH4 |
|
NO2 |
|
SO2 |
MP |
-182.5oC |
|
-163.6oC |
|
-75.5oC |
BP |
-161.5oC |
|
-151.8oC |
|
-10oC |
Inevitably, there must be compounds that
fall between these extremes. For these compounds, the
difference between the electronegativities of the elements is
large enough to be significant, but not large enough to
classify the compound as ionic. Consider water, for example.
H2O |
|
|
O |
EN = 3.44 |
|
H |
EN = 2.20 |
|
|
ŻŻŻŻŻŻŻŻŻ |
|
|
EN = 1.24 |
|
Water is neither purely ionic nor purely
covalent. It doesn't contain positive and negative ions, as
indicated by the Lewis structure on the left in the figure
below. But the electrons are not shared equally, as indicated
by the Lewis structure on the right in this figure. Water is
best described as a polar compound. One end, or pole,
of the molecule has a partial positive charge (
+), and the
other end has a partial negative charge (
-).

As a rule, when the difference between
the electronegativities of two elements is less than 1.2, we
assume that the bond between atoms of these elements is covalent.
When the difference is larger than 1.8, the bond is assumed
to be ionic. Compounds for which the electronegativity
difference is between about 1.2 and 1.8 are best described as
polar, or polar covalent.
Covalent: |
|
EN |
< 1.2 |
|
Polar: |
|
1.2 < |
EN |
< 1.8 |
Ionic: |
|
EN |
> 1.8 |
|
Practice Problem 2: Use
electronegativities to decide whether the following
compounds are best described as either covalent,
ionic, or polar.
(a) Sodium cyanide (NaCN)
(b) Tetraphosphorus decasulfide (P4S10)
(c) Carbon monoxide (CO)
(d) Silicon tetrachloride (SiCl4)
Click
here to check your answer to Practice Problem 2
|

Limitations of
the Electronegativity Concept
Electronegativity summarizes the tendency of an element to
gain, lose, or share electrons when it combines with another
element. But there are limits to the success with which it
can be applied. BF3 (
EN =
1.94) and SiF4 (
EN = 2.08), for example, have
electronegativity differences that lead us to expect these
compounds to behave as if they were ionic, but both compounds
are covalent. They are both gases at room temperature, and
their boiling points are -99.9oC and -86oC,
respectively.
The source of this problem is that each element is
assigned only one electronegativity value, which is used for
all of its compounds. But fluorine is less electronegative
when it bonds to semimetals (such as B or Si) or nonmetals
(such as C) than when it bonds to metals (such as Na or Mg).
This problem surfaces once again when we look at elements
that form compounds in more than one oxidation state. TiCl2
and MnO, for example, have many of the properties of ionic
compounds. They are both solids at room temperature, and they
have very high melting points, as expected for ionic
compounds.
|
|
TiCl2 |
|
MnO |
|
|
|
MP = 1035oC |
|
MP = 1785oC |
|
TiCl4 and Mn2O7,
on the other hand, are both liquids at room temperature, with
melting points below 0oC and relatively low
boiling points, as might be expected for covalent compounds.
|
TiCl4 |
|
Mn2O7 |
|
MP = -24.1oC |
|
MP = -20oC |
|
BP = 136.4oC |
|
BP = 25oC |
The principal difference between these
compounds is the oxidation state of the metal. As the
oxidation state of an atom becomes larger, so does its
ability to draw electrons in a bond toward itself. In other
words, titanium atoms in a +4 oxidation state and manganese
atoms in a +7 oxidation state are more electronegative than
titanium and manganese atoms in an oxidation state of +2.
As the oxidation state of the metal becomes larger, the
difference between the electronegativities of the metal and
the nonmetal with which it combines decreases. The bonds in
the compounds these elements form therefore become less ionic
(or more covalent).

The
Difference Between Polar Bonds and Polar Molecules
The difference between the electronegativities of chlorine
(EN = 3.16) and hydrogen (EN = 2.20) is large
enough to assume that the bond in HCl is polar.
Because it contains only this one bond, the
HCl molecule can also be described as polar.
The polarity of a molecule can be determined by measuring
a quantity known as the dipole moment, which depends
on two factors: (1) the magnitude of the separation of charge
and (2) the distance between the negative and positive poles
of the molecule. Dipole moments are reported is units of debye
(d). The dipole moment for HCl is small: µ = 1.08 d.
This can be understood by noting that the separation of
charge in the HCl bond is relatively small (
EN =
0.96) and that the H-Cl bond is relatively short.
C-Cl bonds (
EN = 0.61) are not as polar as H-Cl bonds (
EN =
0.96), but they are significantly longer. As a result, the
dipole moment for CH3Cl is about the same as HCl:
µ = 1.01 d. At first glance, we might expect a
similar dipole moment for carbon tetrachloride (CCl4),
which contains four polar C-Cl bonds. The dipole moment of
CCl4, however, is 0. This can be understood by
considering the structure of CCl4 shown in the
figure below. The individual C-Cl bonds in this molecule are
polar, but the four C-Cl dipoles cancel each other. Carbon
tetrachloride therefore illustrates an important point: Not
all molecules that contain polar bonds have a dipole moment.


