A Guided Approach to Learning Chemistry
5.1
INTRODUCTION
5.2.
IONIC BONDING
5.3
COVALENT BONDING
5.3.1 Introduction
5.3.2 Lewis theory of covalent bonding
5.3.3 Molecular orbital theory of covalent bonding
5.3.4 Polar covalent bonding, polar molecules
5.4
METALLIC BONDING
5.5
INTERMOLECULAR BONDING
5.5.1 Dipole-dipole bonds
5.5.2 Dipole-induced dipole bonds
5.5.3 Mutually induced dipole-dipole bonds
5.5.4 Hydrogen bonds
5.6
INTERMEDIATE BONDING
Before studying a chapter, go through the list of contents to get an overview. After studying a chapter, go
through the list of contents once again:
(a) to identify and summarize the main concepts and principles;
(b) to obtain a coordinated view of the entire chapter, and the logical interrelationships between the various
parts;
(c) to identify the importance and relevance of the chapter and how it relates to the previous chapters.
Check whether you have done all these satisfactorily by reviewing the text.
143
Chapter 5: Chemical Bonding
5.1
INTRODUCTION
An introduction to the properties of atoms, ions and molecules was given in chapters 3 and 4.
In this chapter we discuss one of these properties in greater detail: the bonding together of
atoms, ions and molecules. This topic is known as chemical bonding.
Bonding is always associated with the rearrangement of some of the electrons present in
the outermost energy level (outermost shell) of the combining atoms (visualize the planetary
model of the atom, section 4.2). Since only the electrons in the outermost shell are involved
in bonding, this shell is often known as the valence shell. The electrons in the valence shell
are called valence electrons.
The electron configurations in the noble gas atoms are found to be particularly stable.
Bonding also often leads to the bonded atoms attaining the electron configurations found in
noble gas atoms by the loss, gain or sharing of electrons.
Bonding can take place between atoms, ions and molecules. According to a simple
model, bonding between atoms of non-metals involves the mutual sharing of electrons; this is
known as covalent bonding. Bonding between atoms of metals is known as metallic bonding,
that between ions is called ionic bonding, and that between molecules is intermolecular
bonding. Models for all these types of bonds are discussed in this chapter. First we consider
ionic bonding (section 5.2), covalent bonding (section 5.3) and metallic bonding (section
5.4). Then we discuss, in section 5.5, the various types of bonds between molecules. Finally,
in section 5.6, we emphasize that bonding is often intermediate between the “pure” types of
bonds.
5.2 IONIC BONDING
To form this type of bond, positive and negative ions must first be formed from the reacting
atoms. The oppositely charged ions then attract one another and bond together. Bonding
between oppositely charged ions is known as ionic bonding.
We illustrate the formation of this type of bond with a typical example: the reaction
between sodium and chlorine atoms. A sodium atom has the electron configuration (table 4.4)
Na:
1s2
2s2
2p6
3s1
The outermost electron, the 3s electron, is more loosely bound to the atomic nucleus than the
other electrons in the atom. A chlorine atom has the electron configuration
Cl:
1s2
2s2
2p6
3s2
3p5
and it readily accepts an electron to form a Cl ion.
When sodium reacts with chlorine, an important step is the transfer of the 3s electron
from a sodium atom to a chlorine atom (figure 5.1). The loss of an electron from an Na atom
(Na → Na+ + e- ) will result in the formation of an Na+ ion whose electron configuration is
Na+:
1s2
2s2
2p6
(a noble gas electron configuration)
while the gain of an electron by a Cl atom (Cl + e- → Cl-) will result in a Cl- ion which has
the electron configuration
Cl-:
144
1s2
2s2
2p6
3s2
3p6
(a noble gas electron configuration)
A Guided Approach to Learning Chemistry
Na+ and Cl- ions, because they have opposite charges, will attract one another and get
bonded. The bond is an ionic bond.
Figure 5.1: Illustration, using the Rutherford-Bohr model, of the formation of
ions from atoms by electron transfer
(a) The electron in the outermost energy level of an Na atom is transferred to a Cl atom.
Na+and Cl- ions are then formed.
(b) The two electrons in the outermost energy level of an Mg atom are transferred to an O
atom to give Mg2+ and O2- ions.
(c) Two electrons from an Mg atom are transferred to two Cl atoms. An Mg2+ ion and two
Cl- ions will then be obtained.
In each case, oppositely charged ions will bond together by ionic bonding, threedimensionally (figure 5.3).
145
Chapter 5: Chemical Bonding
For ionic bonding, positive and negative ions must first be formed. The more readily
positive and negative ions can form; the more likely it is that ionic bonding will occur.
Positive ions are formed readily from the atoms of many metals (for example group 1
and 2 elements), and negative ions are formed readily from the atoms of many non-metals
(for example group 17 and 16 elements). Hence the bonding in most of the compounds
formed by the reaction of a metal with a non-metal is ionic (for example salts, metal oxides).
As a general rule, all salts are built up of ions, and the bonding in them is ionic.
We now consider energies associated with bonds. The bond energy is the energy that
must be supplied to break completely the bond holding the two particles (figure 5.2).
The relative energies of two particles when they are bonded together, and when they are
separated from each other, is shown. The energy of the bonded particles is always lower than
that of the separated particles.
The bond energy is the difference between the energies of the bonded particles and of the
separated particles.
The bond energy between two particles will evidently depend on the force of attraction
between them. The force (F) of attraction between two oppositely charged ions is given by
Coulomb's law (equation 3.5):
F k
q1q2
r2
(3.5)
The equation shows that F, and hence the bond energy between two oppositely charged ions,
would depend only on:
 q1 and q2, the magnitudes of the two charges;
 r, the distance between the centres of the two ions.
Forces between ions do not depend on the chemical nature or identity of the ions.
146
A Guided Approach to Learning Chemistry
EXERCISE 5.1
Which is more stable: an Na atom or an Na+ ion? Give reasons for your answer.
EXERCISE 5.2
The chemical reactivity of the elements of group 1, 2 and the first two elements of group 13
(boron and aluminium) can often be correlated with their atoms attaining the electron
configuration found in the neighbouring noble gas atoms by the loss of electrons. On this
basis, interpret the valency of calcium (Z = 20) and aluminium (Z = 13). How many electrons
must be lost from a calcium atom for it to have the electron configuration of argon, and from
an aluminium atom to attain the electron configuration of neon? In each case, what will be
the charges on the ions formed?
EXERCISE 5.3
Write the electron configuration of the atoms of the element with atomic number 38. Use
Table 4.4. What will be the formula of the most stable ion formed from this element? With
how many chlorine atoms will one atom of this element react?
EXERCISE 5.4
In terms of electron configurations, interpret the reaction between magnesium (Z = 12) and
oxygen (Z = 8) to form magnesium oxide.
EXERCISE 5.5
Consider the compounds formed by the reaction between the following elements:
(a) nitrogen and hydrogen;
(b) hydrogen and oxygen;
(c) calcium and bromine;
(d) magnesium and oxygen;
(e) aluminium and oxygen;
(f) carbon and chlorine;
(g) sulphur and oxygen;
(h) sodium and oxygen.
In which of the compounds formed by the above reactions would you expect bonding to be
ionic/covalent? In each case, write down the formula of the compound formed.
EXERCISE 5.6
In which compound would you expect the bonds to be stronger? Explain (use Coulomb's
law).
(a) NaCl or LiCl
(b) CaO or MgO
Thus far we have considered bonding between two individual ions. Such bonding, which will
lead to the formation of ion pairs (for example Na+Cl-), can, however, take place only in the
147
Chapter 5: Chemical Bonding
vapour phase. Bonding in solids is more complicated since it occurs three-dimensionally. We
shall consider this now.
Bonding and structure in ionic solids
In an ionic solid (for example all salts, some acids, some bases, and some metal compounds),
the particles present are positive and negative ions. Since oppositely charged ions attract one
another, and ions of the same type of charge repel one another, each ion in a crystal will be
bonded with, and will therefore be surrounded on all sides by, oppositely charged ions. The
entire crystal is a regular network of positive and negative ions, and the crystal is said to have
a network structure.
How the constituent particles (molecules, atoms, ions) are arranged in a solid can be
determined by some experimental techniques (for example X-ray diffraction, section 6.8).
These techniques reveal, for example, that in a crystal of sodium chloride each sodium ion is
surrounded by six chloride ions, and that each chloride ion is surrounded by six sodium ions
(figure 5.3).
The number of nearest neighbours around an ion is known as its coordination number.
Since each Na+ ion has 6 Cl- ions as its nearest neighbours, the coordination number of an
Na+ ion in sodium chloride is 6. Similarly, the coordination number of a Cl- ion in sodium
chloride is also 6.
Larger spheres represent Cl- ions, smaller spheres Na+ ions.
(a) The three-dimensional arrangement of Na+ and Cl- ions in a crystal of sodium chloride.
(b) An Na+ ion surrounded by 6 Cl- ions. Think three-dimensionally. The Cl- ions labelled 1
and 2 are along the x-axis, 3 and 4 are along the y-axis and 5 and 6 are along the z-axis.
(c) A Cl- ion surrounded by 6 Na+ ions. The Na+ ions labelled a, b, c and d are in the xy
plane, ions a, b, e and f are in the xz plane and ions c, d, e and f are in the yz plane.
148
A Guided Approach to Learning Chemistry
EXERCISE 5.7
The coordination number of Ba2+ ions in solid BaCl2 is 6. What would be the coordination
number of Cl- ions in this solid? Give reasons for your answer.
EXERCISE 5.8
With the help of Coulomb's law explain why:
(a) sodium chloride has a higher melting point than potassium chloride;
(b) calcium oxide has a higher melting point than barium oxide.
5.3 COVALENT BONDING
5.3.1 Introduction
The bonding together of some atoms can be interpreted in terms of a model that involves the
sharing of electrons between the bonded atoms. This type of bond is known as a covalent
bond.
This type of bond occurs when atoms of non-metals (that is, elements of group 14, 15, 16
and 17), combine with one another. The bond is known as a:



single bond if two electrons are shared (for example H:H, also represented H-H);
double bond if four electrons are shared (for example 0::0, also represented 0=0);
triple bond if six electrons are shared (for example N:::N, also represented N =N)).
Covalent bonding generally leads to the formation of molecules, both small molecules
(for example H2O, CO2, C2HsOH, C6H12O6) and large molecules (macromolecules found in
starches, celluloses, plastics, fibres, rubbers, proteins, etc). In just a few instances, this type of
bonding gives network structures (for example diamond, C; silicon carbide, SiC).
Covalent bonding differs from ionic bonding in that the number of covalent bonds
formed by an atom, and also the directions in which these bonds are formed, depends on the
chemical nature of the atom (figure 5.4). The number of covalent bonds formed, for example,
by a hydrogen atom is always 1, by an oxygen atom is 2, by a boron atom is 3 and by a
carbon atom is 4. This is in contrast to ionic bonding where the number of bonds formed does
not depend on the chemical nature of the ion; it depends mainly on the sizes of the ions
(figure 5.4). The directions in which bonds are formed also depend, in covalent bonding, on
the chemical nature of the atom. For example, if two single bonds are formed by an oxygen
atom they would be at an angle of 1080, if three single bonds are formed by a boron atom
they would be at angles of 1200, and if four single bonds are formed by a carbon atom they
would be separated by angles of 1080.
In contrast to ionic bonding, which is easily understood in terms of coulombic attraction
between oppositely charged ions, covalent bonding is not easy to understand. Many theories
(models) have been put forward to describe and explain covalent bonding. We shall consider
only two of these: the Lewis theory and the molecular orbital theory.
149
Chapter 5: Chemical Bonding
(a) Covalent bonding: The type of atom (its electron configuration) determines (i) the
number of covalent bonds formed, (ii) the directions in which the bonds are formed and
hence bond angles.
The figure shows the structures of three molecules formed from an oxygen atom (H2O,
F2O and Cl2O) and from a boron atom (BH3, BF3 and BCl3). An oxygen atom (O), when it
forms single bonds with two other atoms, forms bonds with bond angle of about 1090, and a
boron atom (B) forms three bonds with bond angles of 1200. The bond angle and molecular
shape do not depend very much on the other atoms with which the O or B atoms bond.
(b) Ionic bonding: The number of bonds formed by an ion, and also the directions of these
bonds, is independent of the type of ion. They depend mainly on the sizes of the bonding ions.
Two two-dimensional representations of the bonds formed by spherical positive ions of two
different sizes are shown. In (i), the positive and negative ions have the same size. Then
twelve negative ions can be packed physically around the central positive ion; only six of
these can be seen in the two-dimensional drawing shown. In figure (ii), the negative ion is
much larger than the positive ion. Then only six negative ions can be packed physically
around a positive ion; only four of these can be seen in a two-dimensional drawing.
150
A Guided Approach to Learning Chemistry
5.3.2 Lewis theory of covalent bonding
The Lewis theory (1916) interprets covalent bonding in terms of the combining atoms
sharing one or more pairs of electrons. The electrons shared are only those in the valence
shell (outermost shell). In many molecules, the sharing of electrons leads to each of the
combining atoms attaining the electron configuration found in a noble gas atom.
Table 5.1: Bonding according to the Lewis theory
151
Chapter 5: Chemical Bonding
We illustrate the Lewis theory with some examples (table 5.1). The second column in the
table indicates the electrons in the valence shells of the combining atoms, each electron being
represented either by a dot (∙) or a cross (×). The third column shows how some of the
valence electrons are shared in the combined atoms.
Consider the first example in table 5.1. Seven electrons are present in the valence shell of
a chlorine atom (electron configuration, ls2 2s2 2p6 3s2 3p5). In a chlorine molecule (see
column 3) the two electrons in the shaded region, one originating from each atom, are shared
by the two atoms. Each atom in the molecule may then be considered to have eight electrons
in the valence shell (count the shared pair for each of the atoms), and thus has the electron
configuration of neon. The shared pair of electrons is believed in Lewis' theory to bind the
atoms together. Since two electrons are shared, the bond is a single bond.
The other examples in table 5.1 can be interpreted in a similar manner. The electron
structures for H2O, NH3 and CC14 indicate that there are two single bonds (O-H bonds) in
H2O, three single bonds (N-H bonds) in NH3, and four single bonds (C-Cl bonds) in CC14. In
all these molecules, each atom has attained the electron configuration of a noble gas by the
sharing of electrons.
Consider now the electron structure of O2 (see table). Four electrons are shared between
two oxygen atoms: there is thus a double bond (O=O). Four electrons have to be shared if
each oxygen atom is to be surrounded by eight electrons. The CO2 molecule also has double
bonds. To have 8 electrons in the valence shell of each atom, it is necessary for this molecule
to have the electron structure O = C = O.
In the nitrogen molecule, six electrons have to be shared if each atom is to have 8
electrons in the valence shell. There will thus be a triple bond (N≡N).
In the examples considered thus far, the shared electrons are provided by both atoms. It
is possible, however, for one atom to provide both the shared electrons. A covalent bond in
which both electrons are provided by one atom is known as a coordinate bond and is
represented by an arrowed line (→). An example of a molecule that has a coordinate bond is
sulphur trioxide, SO3. Its electron structure (see table) shows the presence of a coordinate
bond and two double bonds. Two other examples where coordinate bonds are present are
NH4+ (H3N→H+) and NH3BF3 (H3N→BF3); this can be seen from the following equations:
152
A Guided Approach to Learning Chemistry
The rule we have used for writing electron structures (that each atom in a molecule
attains the electron configuration of a noble gas atom) is, however, not always applicable. A
simple example is BCl3, a stable molecule, where the boron (B) atom has only six electrons in
its valence shell (table 5.1).
The representation of the valence shell electrons of a molecule, as shown in column 3 of
table 5.1, is known as the Lewis structure, or the electron-dot formula, of the molecule.
EXERCISE 5.9
The chemical reactivity of the group 14, 15, 16 and 17 elements can be interpreted in terms of
their atoms attaining the noble gas electron configuration, either by the gain or by the sharing
of electrons. On this basis interpret the reaction between:
(a) oxygen (Z = 8) and potassium (Z = 19) to form K2O;
(b) chlorine (Z = 17) and calcium (Z = 20) to form CaCl2;
(c) chlorine (Z = 17) and carbon (Z = 6) to form CCl4;
(d) oxygen (Z = 8) and carbon (Z = 6) to form CO2;
(e) nitrogen (Z = 7) and magnesium (Z = 12) to form Mg3N2.
In which of the above products is the bonding (i) covalent, (ii) ionic?
EXERCISE 5.10
Write the electron-dot formulas (Lewis structures) of
necessary data from a periodic table.
(a) H2S
(b) H2O
(c) BH3
(d)
+
(f) C2H2
(g) NH3
(h) NH4
(i)
(k) HNO3
(l) SO2
(m) CO2
(n)
the following species. Obtain any
CH4
H2SO3
ClO2-
(e) C2H4
(j) H2SO4
(o) CO32-
5.3.3 Molecular orbital theory of covalent bonding
The molecular orbital theory is presently the most widely applicable theory for the
explanation of covalent bonding. It is one of the theories based on quantum mechanics
(section 4.2).
According to quantum mechanics, the formation and stability of a covalent bond is
mainly due to the simultaneous attraction of the bonding electrons by two (and sometimes
more than two) atomic nuclei. Just as the simultaneous pulling of the two ends of a bone by
two dogs will "bind" the dogs together, so the simultaneous attraction of an electron (or
electrons) by two nuclei will bind the nuclei together.
The molecular orbital theory, like all theories, is based on certain hypotheses
(assumptions). An important hypothesis of the molecular orbital theory is that electrons in
molecules, as in atoms, move about in orbitals. Since the orbitals involve molecules, they are
called molecular orbitals. Because molecules are formed by the combination of atoms, it is
generally assumed that molecular orbitals may be obtained by the combination of the atomic
153
Chapter 5: Chemical Bonding
orbitals present in the combining atoms. This is known as the Linear Combination of Atomic
Orbitals (abbreviated LCAO) method for the construction of molecular orbitals.
We now illustrate this method. We shall consider only homonuclear diatomic molecules,
which are molecules built up of two atoms of the same element (for example H2,
N2, O2). The molecular orbitals are then formed by the combination of two similar atomic
orbitals. We have stated already (table 4.3) that the orbitals present in atoms, in the order of
increasing energy, are
1s, 2s, 2px, 2py, 2pz, 3s, 3px, 3py, 3pz, ....
How these atomic orbitals are combined together to give molecular orbitals, and also the
shapes of these orbitals, is shown schematically in figure 5.5.
Quantum mechanical reasoning suggests that the combination of two atomic orbitals will
give two molecular orbitals, having different energies and different shapes. One molecular
orbital has an energy that is lower than that of the combining atomic orbitals. It is thus stable
and is known as a bonding orbital. The second molecular orbital has an energy higher than
that of the combining atomic orbitals. This is an unstable orbital and is known as an
antibonding orbital. An asterisk (*) is used to specify antibonding orbitals.
Consider first the combination of two 1s atomic orbitals (figure 5.5, last row). The
shapes of the two molecular orbitals formed are shown in the second column. The molecular
orbital shown lower down is the bonding orbital. It is formed by the fusion together of the
electron clouds (waves) of the two atomic orbitals. This molecular orbital can be seen to be
symmetrical about the internuclear axis (the line joining the two nuclei); such an orbital is
known as a sigma (σ) orbital. Since this molecular orbital is formed from 1s atomic orbitals,
it is given the symbol 1s σ. The second molecular orbital is the antibonding orbital. This
orbital is also symmetrical about the internuclear axis, and is thus a σ orbital. It is given the
symbol ls σ*; the asterisk (*) specifies that it is an antibonding orbital.
The two molecular orbitals obtained by the combination of two 2s orbitals are
represented by 2s σ and 2s σ*. The former is bonding and the latter is antibonding.
The two molecular orbitals obtained by the linear combination of two 2px orbitals are
also symmetrical about the internuclear axis (figure 5.5), and are thus σ orbitals. They are
designated 2px σ and 2px σ*, the former orbital being bonding and the latter antibonding.
The two molecular orbitals obtained by the combination of two 2py atomic orbitals do
not lie along the internuclear axis. They are known as π orbitals and they are given the
symbols 2py π and 2py π*. Similarly, the two molecular orbitals formed from two 2pz
atomic orbitals are π-orbitals; 2pz π and 2pz π * orbitals.
The energies of these molecular orbitals, which may be obtained from spectroscopic
data, are found to be in the following order for the first few orbitals. The 1s σ orbital has
the lowest energy.
1s σ < 1s σ* < 2s σ < 2s σ* < 2py π = 2pz π < 2px σ < 2py π* = 2pz π* < 2px σ*
154
A Guided Approach to Learning Chemistry
Atomic
orbital
Atomic
orbital
Shape of molecular orbitals
Symbol for
molecular
orbitals
2py π*
2py
2py
2py π
2px σ*
2px
2px
2px σ
2s σ*
2s
2s
2s σ
1s σ*
1s
1s
1s σ
155
Chapter 5: Chemical Bonding
Electron structures of molecules according to the molecular orbital theory
The molecular orbital electron structure of a molecule is a representation of the electrons
present in the various molecular orbitals.
The method for writing this electron structure is similar to that used for writing electron
structures (electron configurations) of atoms (section 4.4.7). The information required is the
energies of the various molecular orbitals. The electrons present in the molecule are then
assigned to these orbitals, in the order of increasing energy, starting with the orbital that has
the least energy (1sσ). During this assignment, one should use the Pauli exclusion principle
(not more than two electrons can be present in any orbital) and Hund's rule (in orbitals having
the same energy, electrons occupy different orbitals rather than the same orbital).
The electron structures of some molecules, and molecular ions, written on this basis are
given in table 5.2.
Table 5.2: Electron structures of some simple molecules and molecular
ions
Molecule Electron structure
H2+
1sσ1
H2
1sσ2
He2
1sσ2
1sσ* 2
Li2
1sσ2
1sσ* 2
2sσ2
Be2
1sσ2
1sσ* 2
2sσ2
2sσ* 2
N2
1sσ2
1sσ* 2
2sσ2
2sσ* 2
2pyπ2
2pzπ2
2pxσ2
O2
1sσ2
1sσ* 2
2sσ2
2sσ* 2
2pyπ2
2pzπ2
2pxσ2
2pyπ* 1
2pzπ* 1
F2
1sσ2
1sσ* 2
2sσ2
2sσ* 2
2pyπ2
2pzπ2
2pxσ2
2pyπ* 2
2pzπ* 2
Consider first the hydrogen molecule; this has two electrons. The molecular orbital that
has least energy is 1sσ, and the two electrons will be in this orbital. The electron structure is
thus written as 1sσ2. According to the molecular orbital theory, electrons present in any
bonding orbital bind the atoms together, and thus two hydrogen atoms will be bonded
together to give a stable hydrogen molecule.
Consider next the possibility of forming a stable He2 molecule. Its electron structure,
2
1sσ 1sσ* 2, indicates two electrons in a bonding orbital (1sσ), and two electrons in an
antibonding orbital (1sσ*). According to the molecular orbital theory, a stable bond will be
formed between two atoms only if the total number of electrons present in all the bonding
orbitals exceeds the total number of electrons present in all the antibonding orbitals. In He2,
the bonding effect of the two electrons present in the 1sσ orbital is cancelled by the
antibonding effect of the two electrons in the 1sσ* orbital. Thus the He2 molecule will not be
expected to be stable, according to the molecular orbital theory.
As another example, consider the electron structure of the N2 molecule (see table 5.2).
The bonding effect of the electrons in the 1sσ and 2sσ orbitals will be cancelled by the
156
A Guided Approach to Learning Chemistry
antibonding effect of the electrons in the 1sσ* and 2sσ* orbitals. So we need to consider only
the six electrons present in the 2pyπ, 2pzπ and 2pxσ orbitals, all bonding orbitals. The triple
bond (N≡N) in the nitrogen molecule thus consists, according to the molecular orbital theory,
of a σ bond (2pxσ ) and two π bonds (2pyπ and 2pzπ).
The bonding in the oxygen molecule is more complicated (table 5.2). The excess of
bonding electrons over antibonding electrons is four, and so we can say that four electrons are
involved in bonding. Two of them are present in a σ orbital (2pxσ) and the other two electrons
are in the 2pπ orbitals: one in 2pyπ and the other in 2pzπ. (Note: the bonding effect of one
electron in the 2pyπ orbital is cancelled by the antibonding effect of the electron in the 2pyπ*
orbital; similarly for the 2pzπ orbitals.)
EXERCISE 5.11
How many bonding electrons, and how many antibonding electrons, would be present in the
species Li2+? Would you expect this species to be stable?
EXERCISE 5.12
In terms of the molecular orbital theory, which of the following species, in the gaseous state,
would you expect to be stable?
(a) Li2
(b) Li2+
(c) Be2
(d) C2
(e) F2
(f) Ne2
Give reasons for your answers.
5.3.4 Polar covalent bonding, polar molecules
A covalent bond is formed by the sharing of electrons between atoms. A pure (100%)
covalent bond, known as a non-polar covalent bond, can be formed only if the two atoms are
from the same element (that is, in homonuclear molecules such as H-H, Cl-Cl, O=O, N =N).
Only then will the bonding electrons be shared equally between the two atoms; only then will
the electrons move about symmetrically around both nuclei.
If the two atoms are from different elements, the bond will not be purely (100%)
covalent. It will be partly ionic and the bond is known as a polar covalent bond. A polar
covalent bond may be considered to be partly ionic and partly covalent. It is intermediate
between a pure ionic and a pure covalent bond.
Consider, as an example, an HCl molecule (figure 5.6). A chlorine atom is found to
attract electrons more strongly than does a hydrogen atom (chlorine is thus said to be more
electronegative than hydrogen). The two bonding electrons will be found, more often, closer
to the chlorine nucleus than to the hydrogen nucleus. The chlorine atom will thus have a
greater share of these two electrons. Since electrons are negatively charged, it follows that the
chlorine atom will be partially negative (-) and the hydrogen atom will be partially positive
(+). The molecule will thus have a positive centre and a negative centre. It is said to be a
157
Chapter 5: Chemical Bonding
polar molecule and the bond between the atoms is a polar covalent bond. Since there are two
poles (+ and -), the molecule is also said to have a dipole.
Some molecules are more polar than others. This depends on the difference in the
electronegativities (electron attracting power) of the two atoms. The greater the difference in
electronegativities, the more polar will be the bond. For example, HCl is more polar than
HBr; this is because Cl is more electronegative than Br. The polarity of a diatomic molecule
is usually expressed in terms of a quantity known as the dipole moment (μ), which is defined
by the equation
μ = qr
(5.1)
where q is the net charge at each atomic nucleus, and r is the distance separating the two
nuclei.
Figure 5.6: Schematic representations of an HCl molecule
(a) Lewis representation: The two bonding electrons are closer to the Cl nucleus than to the
H nucleus.
(b) Molecular orbital representation: The orbital is not symmetrical; the electron cloud is
concentrated more around the Cl nucleus.
Diatomic molecules have only one bond and therefore the dipole moment of a molecule
is the same as the dipole moment of the bond. Polyatomic molecules, in contrast, have two or
more bonds and a molecule's dipole moment is the net effect of the dipole moments of the
various bonds. It is thus possible for a polyatomic molecule to be non-polar, even when it
contains polar bonds. To illustrate this consider, as an example, the CO2 molecule
(-)
O = (+)C = O(-)
Oxygen atoms attract the bonded electrons (four electrons in each double bond) more
towards themselves. The oxygen atoms will thus be negatively charged and the carbon atom
positively charged. However, because the molecule is linear, the dipoles associated with the
two bonds will act in opposite directions and will cancel each other. The molecule as a whole
will therefore be non-polar.
The polarity of molecules is important because it influences many physical and chemical
properties of substances. For example, if the constituent molecules are polar, the opposite
poles in different molecules will attract one another (figure 5.7). The molecules will then be
158
A Guided Approach to Learning Chemistry
held together more strongly. More energy will then be needed to separate the molecules;
substances having polar molecules (for example H-O-H) will therefore have relatively high
melting points, boiling points, and heats of vaporization.
EXERCISE 5.13
Many methods are available for determining dipole moments, and dipole moment data
provide an important method for deducing molecular shapes. This is based on the principle
that the dipole moment of a molecule is the net effect (vector sum) of the dipole moments of
all the bonds present in that molecule. By making use of this principle, deduce whether:
(a) a water molecule (H-O-H) is linear or bent, given that it has a dipole moment;
(b) BeCl2 molecule is linear or bent, given that its dipole moment is zero;
(c) NH3 molecule is planar or non-planar, given that it has a dipole moment.
EXERCISE 5.14
In which of the following molecules would you expect polar covalent bonding between
atoms?
(a) H-O-H
(b) O=O
(c) H-H
(d) N≡N
(e) H-F
(f) F-Be-F
(g) NH3
(h) CCl4
EXERCISE 5.15
Arrange the following molecules in the order of decreasing dipole moment. Make use of the
fact that the electronegativity of an atom decreases with increasing atomic number in a group
in the periodic table.
HCl,
HI,
HBr,
HF
5.4 METALLIC BONDING
We have already described a simple model for the internal structure of metals (figure 2.8).
There is a regular arrangement of positive ions (for example Cu2+ ions in copper metal, Ag+
ions in silver, Zn2+ ions in zinc, Mg2+ ions in magnesium) and these are surrounded by a
common “pool” or “sea” of electrons that move about throughout the metal.
In metallic bonding, positive ions are bonded together. This is not easy to understand,
since positive ions would be expected to repel one another. An understanding of metallic
bonding is possible only by quantum mechanical reasoning which suggests that bonding is
the result of the simultaneous attraction of the electrons by the various positive ions.
5.5 INTERMOLECULAR BONDING
Thus far we have considered the types of bonds (ionic, covalent, metallic) between atoms and
ions.
159
Chapter 5: Chemical Bonding
We now consider bonding between molecules - that is, intermolecular bonding.
Intermolecular bonds have relatively low energies (generally less than 40 kJ mol-1). This is in
contrast to ionic, covalent and metallic bonds which have much higher energies (generally
more than 200 kJ mol-1).
Intermolecular bonds are of four types:
 dipole-dipole bonds;
 dipole-induced dipole bonds;
 mutually induced dipole-dipole bonds;
 hydrogen bonds.
5.5.1 Dipole-dipole bonds
These exist between polar molecules - that is, molecules that have dipole moments (for
example H2O, HCl). The positive end of the dipole in one molecule attracts and thus bonds
with the negative end of the dipole of another molecule (figure 5.7 (a)). An important type of
a dipole-dipole bond is the hydrogen bond; this will be discussed at the end of this section.
5.5.2 Dipole-induced dipole bonds
Bonding between a polar and a non-polar molecule is due to this type of bond. A dipole is
first induced in the non-polar molecule, and then attraction and bonding takes place. An
oxygen molecule, for example, is non-polar because the electrons move symmetrically
around the two oxygen nuclei. But in the presence of a polar molecule a dipole will be
induced. Suppose, for example, that an H2O molecule approaches an oxygen molecule with
the oxygen atom pointing towards it (figure 5.7 (b)). Since the oxygen atom in H2O is
negatively charged, it will repel the electron cloud (negative charge) in the oxygen molecule
thereby distorting it. That part of the oxygen molecule close to H20 will then become
positively charged. Attraction between the negative charge of the original dipole (in H2O) and
the positive charge induced in O2 will bond the H2O and O2. The bond formed between a
polar and a non-polar molecule is always a dipole-induced dipole bond.
5.5.3 Mutually induced dipole-dipole bonds
Even non-polar molecules such as O2, N2, He and CH4 attract one another, although this
attraction is relatively weak. This is clear from the fact that all these molecules stick together,
if the temperature is low enough, to give liquids.
The bonding between non-polar molecules may be understood in terms of mutually
induced dipole-dipole forces. Here a dipole is first induced in each of the bonding molecules.
This is possible because electron clouds in atoms and molecules are not rigid structures; they
are easily deformed by forces. Consider, for simplicity, a spherical molecule. The electrons
move around the central nucleus symmetrically, and the electron cloud around the nucleus is
spherically symmetrical (figure 5.7, (c), (i)). The molecule will then be non-polar. In the
close presence of another molecule, the electron clouds of both the molecules can be
160
A Guided Approach to Learning Chemistry
deformed, thereby inducing dipoles in them (figure (c), (ii)). The opposite poles in the
induced dipoles will then attract and thereby bond the molecules together.
Figure 5.7: Schematic representation of some types of bonds between
molecules (the dotted lines represent the bonds between the molecules
(a) Dipole-dipole bonds: Two polar molecules will attract one another, the negatively
charged atom (e.g. O in H2O) of one molecule attracting the positively charged atom (e.g. H
in H2O) of another molecule.
(b) Dipole-induced dipole bonds: Here a polar molecule (e.g. H2O) first induces a dipole in a
non-polar molecule. The opposite poles then attract and this leads to bonding.
(c) Mutually induced dipole-dipole bonds: Here bonding is due to mutually induced dipoledipole forces.
(i) A non-polar spherical molecule. The electron cloud is symmetrical and the molecule
does not have a dipole.
(ii) Shows the distortion of the electron clouds in two adjoining molecules.
Dipoles are induced in each molecule. Bonding then takes place between the positive end of
one dipole and the negative end of another dipole.
5.5.4 Hydrogen bonds
This is a dipole-dipole bond in which a hydrogen atom is involved. A special name is given
because this type of bond is extremely important in influencing the internal structures and
properties of many substances such as proteins and enzymes. Hydrogen bonds are formed
only in compounds where a hydrogen atom is bonded to the strongly electronegative atoms
fluorine, oxygen or nitrogen. In these compounds the bond is strongly polar.
To illustrate and explain hydrogen bonding, consider a sample of hydrogen fluoride. It
consists of HF molecules, the H-F bond being strongly polar (H(+)-F(-)). The F atom, being
negatively charged, will attract the H atom of a neighbouring H-F molecule to form
... H(+)−F(-)... H(+)−F(-)... H(+)−F(-)... H(+)−F(-)
161
Chapter 5: Chemical Bonding
The dotted lines represent hydrogen bonds; they bond HF molecules. Another compound that
shows hydrogen bonding is water. Each H2O molecule has two O-H bonds which are strongly
polar (H(+)-O(-)-H(+)). The opposite poles (+ and-) in neighbouring water molecules will
attract one another. There is experimental evidence that each water molecule links, for very
short periods of time, with four other water molecules as indicated schematically in figure
5.8. The dotted lines represent hydrogen bonds. The shape and electron structure of an H2O
molecule (see table 6.1) is such that the four hydrogen bonds are arranged tetrahedrally.
As already mentioned, hydrogen bonding is important because it influences the
properties of many substances. The relatively high melting point, boiling point and heat of
vaporization of water (compared with the analogous substance H2S) is due to hydrogen bonds
which help to hold the molecules strongly together. Hydrogen bonds also link molecules in
many living systems. For example, they link protein molecules. A protein molecule has -C=O
and -N-H groups and there is hydrogen bonding between these groups (-C=O... H-N-). This
bond occurs within the same molecule (intramolecular hydrogen bond, and not
intermolecular hydrogen bond as in H2O), and is responsible for maintaining the coiled shape
(figure 5.9) of protein molecules.
Figure 5.8: Schematic representation of hydrogen bonding in water
The dotted lines represent hydrogen bonds. Each water molecule is bonded, by hydrogen
bonds, to four other water molecules. Two of these hydrogen bonds involve the oxygen atom
(bonding is to a hydrogen atom in two neighbouring water molecules), while the other two
involve each of the hydrogen atoms (bonding is to the oxygen atom in two neighbouring
water molecules).
The four hydrogen bonds that link any H2O molecule with other H2O molecules are
arranged tetrahedrally. The tetrahedral arrangement is to be expected from the electron
structure of a water molecule (see table 6.1). The tetrahedral structure is not permanent: it
forms and breaks up very rapidly all the time.
162
A Guided Approach to Learning Chemistry
Figure 5.9: Intramolecular hydrogen bonds in a small part of a protein
molecule
A protein molecule is a macromolecule and it has a coiled structure. The dotted lines between
the –N-H groups and O-C- groups represent hydrogen bonds. These bonds help to maintain
the characteristic coiled shape of protein molecules.
EXERCISE 5.16
In which of the following molecules would you expect hydrogen bonding? State the criterion
you used for making the decision.
CH4, CH3OH, NO2, H2O, CH3COOH, CH2F2, CH3NH2
5.6 INTERMEDIATE BONDING
All types of bonds are essentially similar in that they ultimately depend on the attraction
between unlike charges and repulsion between like charges. Because of this similarity in the
various types of bonds, it is not surprising that there is a continuous variation from one bond
type to another.
Consider two atoms A and B. If they have the same electronegativity (the same ability to
attract electrons; this happens when A and B are the same type of atom), the bond between
them will be non-polar (that is, 100% covalent bond). If the electronegativity of B exceeds
that of A, the bonding electrons will, on the average, be closer to atom B and the bond will be
polar. As the difference in the electronegativities of B and A increases, the polarity of the
bond increases. If the electronegativity difference is very large, the electron pair will not be
shared but will be associated exclusively with atom B. This corresponds to a negative ion B
and a positive ion A, or to a 100% ionic bond. No substance, however, has 100% ionic
bonding.
163
Chapter 5: Chemical Bonding
In many crystals too, the bonding is intermediate in character. For example, in a silver
chloride crystal the bonding between the silver and chloride ions is not entirely ionic; it is
partly covalent.
In some metals, the bonding is also intermediate in character. For example, in tin (Sn)
and germanium (Ge) there are four directed bonds, a property characteristic of covalent
bonds. Some mobile electrons, characteristic of metallic bonding, are also present, since these
metals conduct electricity to some extent. The bonding in these metals is thus intermediate
between covalent and metallic.
164