C1B- T2 – Bonding & Molecular Structure,

Comparative Study of P-Block Elements

1. Chemical Bonding and Molecular Structure:

Ionic Bonding: General characteristics of ionic bonding. Energy considerations in ionic
bonding, lattice energy and solvation energy and their importance in the context of stability
and solubility of ionic compounds. Statement of Born-Landé equation for calculation of
lattice energy, Born-Haber cycle and its applications, polarizing power and polarizability.
Fajan’s rules, ionic character in covalent compounds, bond moment, dipole moment and
percentage ionic character.

Dr. Subhradeep Mistry

Asst. Prof. Ramananda college
 Types of bonds
Atoms may attain a stable electronic configuration in three different ways: by losing
electrons, by gaining electrons, or by sharing electrons. Elements may be divided into:

I. Electropositive elements, whose atoms give up one or more electrons fairly readily.
II. Electronegative elements. which will accept electrons.
III. Elements which have little tendency to lose or gain electrons.

Three different types of bond may be formed, depending on the electropositive or

electronegative character of the atoms involved.

Electropositive element + Electronegative element = Ionic bond

Ionic bonding involves the complete transfer of one or more electrons from one
atom to another.
 Ionic bonds
Ionic bonds are formed when electropositive elements react with electronegative elements.
We will take an example of Ionic bonds in Sodium chloride (NaCl).

2 Na(s)+Cl2 (g) → 2NaCl(s)

Electronic configuration of Na: 1s2 2s2 2p6 3s1

Na (sodium atom) → Na+ (sodium ion) + electron
Electronic configuration of Na+: 1s2 2s2 2p6 (Noble gas Neon configuration)

Electronic configuration of Cl: ls2 2s2 2p6 3s2 3p5

Cl (chlorine atom ) + electron - Cl- (chloride ion)
Electronic configuration of Cl-: ls2 2s2 2p6 3s2 3p6 (Noble gas Argon configuration)

Electrostatic attraction between the positive and negative ions holds the ions together in a
crystal lattice. The process is energetically favourable as both sorts of atoms attain the
stable noble gas configuration, and sodium chloride Na+C1- is formed readily.
Electron configuration of a Ca ion is : 1s2 2s2 2p6 3s2 3p6 4s2 = [Ar] 4s2
Electron configuration of a Ca2+ ion is : 1s2 2s2 2p6 3s2 3p6
Electronic configuration of Cl-: ls2 2s2 2p6 3s2 3p6
 Salient features of Ionic bonds
Melting point:
Ionic compounds are typically solids and usually have high melting and boiling points. In contrast
covalent compounds are typically gases, liquids or low melting solids. These differences occur because
of differences in bonding and structure.

Ionic compounds, are made up of positive and negative ions arranged in a regular way in a lattice.
The attraction between ions is electrostatic, and is non-directional, extending equally for all
directions. Melting the compound involves breaking the lattice. This requires considerable energy,
and so the melting point and boiling point are usually high, and the compounds are very hard.

Conductivity:
Ionic solids conduct electricity poorly, but when melted or dissolved are good electrical conductors.
Conduction is achieved by the ions migrating towards the electrodes due to the electric potential.

If an electric current is passed through a solution of sodium chloride, Na+ ions are attracted to the -ve
electrode (cathode), where they gain an electron and form Na atoms. The Cl- ions are attracted to the
+ve electrode (anode), where they lose an electron and become Cl2 atoms. This process is called
electrolysis.
Solubility
If they dissolve at all, ionic compounds are usually soluble in polar solvents. These are solvents of
high, dielectric constant such as water, or the mineral acids.

Speed of reactions
Ionic compounds usually react very rapidly, whilst covalent compounds usually react slowly. For
ionic reactions to occur, the reacting species are ions, and as these already exist, they have only to collide
with the other type of ion. For example, when testing a solution for chloride ions (by adding silver
nitrate solution), precipitation of AgCI is very rapid. Reactions of covalent compounds usually involve
breaking a bond and then substituting or adding another group.

❑ It is important to realize that bonds are not necessarily 100% covalent or 100% ionic, and that
bonds of intermediate character exist. If a molecule is made up of two identical atoms, both atoms
have the same electronegativity (EN), and so have an equal tendency to gain electrons. This is
sometimes called a non-polar covalent bond.

❑ Mostly molecules are formed with different types of atoms, and the EN of the two atoms differs.
For example in HF, Fluorine is the most EN atom, and it attracts electrons more strongly when
covalently bonded. Thus, the F atom has a very small -ve charge δ- and the H has a small +ve
charge δ+. These bonds are sometimes called polar covalent bonds.
 Lattice energy
The lattice energy (U) of a crystal is the energy evolved when one gram molecule of the
crystal is formed from gaseous ions:

Na+(g) + Cl-(g) -> NaCl(crystal) U = - 782 kJ mo1-1

Lattice energies cannot be measured directly. but experimental values are obtained from
thermodynamic data using the Born--Haber cycle.

Theoretical values for lattice energy may be calculated. The ions are treated as point charges,
and the electrostatic (coulombic) energy E between two ions of opposite charge is calculated:

E = - Z+ Z- e2 / r

Where, z+ and z- are the charges on the positive and negative ions
e is the charge on an electron
r is the inter-ionic distance
 For more than two ions, the electrostatic energy depends on the number of ions. and also on A
their arrangement in space. For one mole. the attractive energy is

E = - N0 A Z+ Z- e2 /r [Where, N0 is the Avogadro constant - 6.023 X 1023 mol-1 and A is the

Madelung constant, which depends on the geometry of the crystal.]
❑ The equation for the attractive forces between the ions gives a negative value for energy. that is
energy is given out when a crystal is formed.

❑ The inter-ionic distance r occurs in the denominator of the equation. Thus the smaller the value
of r, the greater the amount of energy evolved when the crystal lattice is formed, and hence
the more stable the crystal will be.

❑ Mathematically. the equation suggests that an infinite amount of energy should be evolved if
the distance r is 0. When the inter-ionic distance becomes small enough, they begin to repel
each other. This repulsion originates from the mutual repulsion of the electron clouds.

❑ The repulsive force is given by B/rn, where B and n both are constant. B depends on structure,
and n is called the Born exponent. For one gram molecule the total repulsive force is (N0B)/rn.
 Born-Lande equation
The total lattice energy is the sum of the attractive and the repulsive forces.

U = - N0AZ+Z-e2/r (Attractive force) + N0B/rn (repulsive force) … (1)

The equilibrium distance between ions is determined by the balance between the attractive and
repulsion terms. At equilibrium. dU/dr = 0, and the equilibrium distance r = r0

𝐝𝑼
= N0AZ+Z-e2/r02 – nN0B/r0n+1 = 0 … (2)
𝐝𝒓

By rearranging the above equation we get, B = AZ+Z-e2r0n-1 / n … (3)

From equation 1 and 3 we get, U = - N0AZ+Z-e2 / r0 (1- 1/n)

This equation is called the Born-Lande equation. It allows the lattice energy to be calculated from a
knowledge of the geometry of the crystal, and hence the Madelung constant, the charges z+ and z-,
and the inter-ionic distance. When using SI units, the equation takes the form

U = - N0AZ+Z-e2 / 4πε0r0 (1- 1/n) [ε0 is the permittivity of free space = 8.854 x 10-12 Fm-1]
 Born-Haber Cycle
❑ This cycle devised by Born and Haber in 1919 relates the lattice energy of a crystal to other
thermochemical data.

❑ The elements in their standard state are first converted to gaseous atoms, and then to ions,
and finally packed into the crystal lattice.

❑ The enthalpies of sublimation and dissociation and the ionization energy are positive
since energy is supplied to the system. The electron affinity and lattice energy ate negative
since energy is evolved in these processes.

❑ According to Hess's law, the overall energy change in a process depends only on the
energy of the initial and final states and not on the route taken. Thus the enthalpy of
formation ΔHf is equal to the sum of the terms going the other way round the cycle.
The Born-Haber Cycle can be reduced to a
single equation:

Heat of formation= Heat of atomization+

Dissociation energy + (sum of Ionization
energies)+ (sum of Electron affinities) +
Lattice energy

Or - ΔHf = ΔHs + I + ½ ΔHd - E – U

For NaCl,
-381.2 = +108.4 + 495.4 + 120.9 – (-348.6) -
U
Hence, U = - 757.3 kjmol-1

From, this we can both extract the values of

E or U. If U value is known, we can get E
and vice-versa. Born-Haber Cycle for the formation of NaCl
❑ It is useful to know the lattice energy, as a guide to the solubility of the crystal.

❑ When a solid dissolves, the crystal lattice must be broken up (which requires that energy is
put in). The ions so formed are solvated (with the evolution of energy).

❑ When the lattice energy is high a large amount of energy is required to break the lattice.

❑ lt is unlikely that the enthalpy of solvation will be big enough (and evolve sufficient energy
to offset this), so the substance will probably be insoluble.

❑ Lattice energies may also provide some information about the ionic/covalent nature of the
bonding. If the lattice energy is calculated theoretically assuming ionic bonding then the
value can be compared with the experimental value for the lattice energy obtained from the
experimentally measured quantities in the Born-Haber cycle. Close agreement indicates
that the assumption that bonding is ionic is in fact true, whilst poor agreement may
indicate that the bonding is not ionic.
 Polarizing power and polarizability
Ionic and covalent bonding are two extreme types of bonding, and almost always the bonds
formed are intermediate in type, and this is explained in terms of polarizing (that is
deforming) the shape of the ions.

The type of bond between A+ and B- depends on the effect one ion has on the other. The
positive ion attracts the electrons on the negative ion and at the same time it repels the
nucleus, thus distorting or polarizing the negative ion.

The negative ion will also polarize the positive ion, but since anions are usually large, and
cations small, the effect of a large ion on a small one will be much less pronounced.

If the degree of polarization is quite small., then the bond remains largely ionic; If the
degree of polarization is large; electrons are drawn from the negative ion towards the positive
ion, resulting in a high concentration of electrons between the two nuclei, and a large degree
of covalent character results.
 Fajan’s Rule
Fajan put forward four rules which summarize the factors favouring polarization and
hence covalency.

1. A small positive ion favours covalency. In small ions the positive charge is concentrated
over a small area. This makes the ion highly polarizing, and very good at distorting the
negative ion.

2. A large negative ion favours covalency. Large ions are highly polarizable, that is easily
distorted by the positive ion, because the. outermost electrons are shielded from the
charge on the nucleus by filled shells of electrons.

3. Large charges on either ion, or on both ions, favour covalency. This is because a high
charge increases the amount of polarization.

4. Polarization, and hence covalency, is favoured if the positive ion does not have a noble
gas configuration.
 Dipole Moments
When non-identical atoms are joined in a covalent bond, the electron pair will be
attracted more strongly to the atom that has the higher electronegativity. As a
consequence, the electrons will not be shared equally. Such bonds are said to be polar and to
possess partial ionic character, and they may confer a polar nature on the molecule as a
whole.

The larger the difference in electronegativity, the larger the dipole moment. The distance
between the charge separation is also a deciding factor into the size of the dipole moment.
The dipole moment is a measure of the polarity of the molecule.

One of the most common examples is the water molecule,

made up of one oxygen atom and two hydrogen atoms.
The differences in electronegativity and lone electrons
give oxygen a partial negative charge and each
hydrogen a partial positive charge.
When two electrical charges, of opposite sign and equal magnitude, are separated by a distance,
an electric dipole is established. The size of a dipole is measured by its dipole moment (μ).
Dipole moment is measured in Debye units, which is equal to the distance between the charges
multiplied by the charge (1 Debye = 3.34 X 10-30 C m). The dipole moment of a molecule can
be calculated by Equation
μ dipole moment vector
qi is the magnitude of the charge, and ri is the vector representing
the position of ith charge
The above equation can be simplified for a simple separated two-charge system like
diatomic molecules or when considering a bond dipole within a molecule

μdiatomic = Q x r

This bond dipole is interpreted as the dipole from a charge separation over a distance r between
the partial charges Q+ and Q- (or the more commonly used terms δ+ δ-); the orientation of the
dipole is along the axis of the bond.
Consider a simple system of a single electron and proton separated by a fix distance. When
proton and electron close together, the dipole moment (degree of polarity) decreases.
However, as proton and electron get farther apart, the dipole moment increases. In this
case, the dipole moment calculated as

μ = Q x r = (1.60 × 10-19 C)(1.00 × 10-10 m) = 1.60 × 10−29 C ⋅ m

The Debye characterizes size of dipole moment. When a proton & electron 100 pm apart,
the dipole moment is 4.80 D.

μ = (1.60 ×10-29 C ⋅ m) (1D / 3.34 ×10−30 C ⋅ m) = 4.80 D

 Bond moments
The polyatomic molecules contain number of bonds between the atoms. Each bond has a
definite dipole moment and hence makes a definite contribution to the overall dipole
moment of the molecule. This dipole moment of a chemical bond is called bond moment.
It does not depend on the nature of molecule.
The dipole moment of a molecule is really vectorial sum of the individual bond moments
present in it. No direct method is known to obtain the magnitude of a bond moment.

The dipole moment acts in the direction of the vector quantity!!!

Net Dipole Moment and %
Ionic character calculation
VVI