The Covalent Bond
Valence Electrons
The electrons in the outermost shell are the valence electronsthe
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
Phase at room
temperature
Density
Melting point
Boiling point
Cl2
Solid
Gas
g/cm3
2.165
801°C
0.003214 g/cm3
1413°C
-34.6°C
-100.98°C
Ability of aqueous
solution to conduct
Conducts
Does not conduct
electricity
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.
NaC
l
Cl
Na
Li
F
F
Li
SrBr
2
Br
Sr
EN = 3.16
EN = 3.98
EN = 2.96
EN = 0.93
EN = 0.98
EN = 0.95
¯¯¯¯¯¯¯¯¯
¯¯¯¯¯¯¯¯¯
¯¯¯¯¯¯¯¯¯
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
846oC
657oC
MP 801oC
1717oC
2146oC
BP 1413oC
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
C
H
SO
2
O
S
EN = 2.55
O
EN = 3.44
EN = 3.44
EN = 2.20
N
EN = 3.04
EN = 2.58
¯¯¯¯¯¯¯¯¯
¯¯¯¯¯¯¯¯¯
¯¯¯¯¯¯¯¯¯
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
-163.6oC
-75.5oC
MP -182.5oC
o
o
-151.8 C
-10oC
BP -161.5 C
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
H
EN = 3.44
EN = 2.20
¯¯¯¯¯¯¯¯¯
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
Ionic:
EN
>
1.8
<
1.8
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
o
BP = 136.4 C
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.
+
-
C
l
Because it contains only this one bond, the HCl molecule can also be
described as polar.
H
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.