UNIT 2
CHEMICAL BONDING
INTRODUCTION
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Almost everything a person sees or touches in daily life, like the air we breathe, the food we
eat, the clothes we wear, are the result of chemical bonds.
It is observed that the atoms of all the elements, except those for noble gases, tend to remain
in combined state with the atoms of same or other element. They do not exist as single atoms
under ordinary conditions. Such atomic aggregates occur as molecules or as giant network.
The attractive force which holds atoms, ions, and molecules together is called a chemical
bond.
Octet Rule
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Chemical bond is the attractive force that binds atoms together in a molecule, or a compound.
Atoms combine to form chemical bonds to be stable (to minimize energy).
Atoms join together in order to achieve the nearest noble gas configuration.
The process of achieving a stable electron configuration of eight electrons in the outer most
shell is called octet rule.
 Atoms with 8 electrons (except He) in the outer most shell are stable.
 Atoms can achieve the noble gas configuration by gaining, losing, or sharing electrons.
 The type of electrons that participate in bond formation is valence electrons.
TYPES OF CHEMICAL BONDING
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The forces of attraction that hold atoms together are called chemical bonds.
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There are three main types of chemical bonds covalent, ionic and metallic bonds.
1. IONIC BONDS ((Electrovalent Bond)
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When an atom either loses or gains electrons, it becomes an ion.
Any atom or molecule with a net charge, either positive or negative, is known as an ion.
Net electric charge = number of protons – number of electrons
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Metals tend to lose these electrons and become positively charged ions (Cations).
Non-metals tend to gain electrons to form negatively charged ion(anions).
Ionic bond is the electrostatic force of attraction between oppositely charged ion (cations
and anions).
The bond is produced when electrons are transferred from the outermost shell of a metal
atom to the outermost shell of a nonmetal atom.
Ionic compounds are usually formed when metal cations bond with non-metal anions. The
only common exception is ammonium ion which is not a metal, but it forms ionic
compounds.
Example

Sodium metal(Na) reacts with chlorine gas (Cl2), to form sodium chloride (NaCl)
2Na(s) + Cl2(g) → 2NaCl(s)

A sodium atom, Na, by losing an electron, forms a sodium ion, Na+, which has the same
electron configuration as the noble gas neon.
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
A chlorine atom, Cl, by gaining an electron, forms a chloride ion, Clˉ , which has the same
electron configuration as the noble gas argon.
Cl


+
eˉ
→ Clˉ
When particles have opposite electric charges, a force of attraction exists between them;
known as electrostatic force of attraction.
In sodium chloride, the sodium ions and the chloride ions are held together by electrostatic
force of attraction, thus forming an ionic bond or electrovalent bond.
Exercise 2.1 Page - 70
Lewis Electron-Dot Symbols
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Lewis formulas (Electron dot notation) consist of the symbol of the element plus dots equal
to the number of valence electrons in the atom or ion.
Chemical reactions between two combining atoms are represented by means of electron dot
diagram or atomic diagram.
Example

Lewis symbols showing the transfer of electrons during the formation of ionic compounds.
Exercise 2.2 Page - 71
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Formation of Ionic Bonding
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The formation of ionic compounds is the result of

low ionization energies and

high affinities for electrons,

It is always an exothermic process
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The compound is formed because it is more stable (lower in energy) than its elements.
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Much of the stability of ionic compounds result from the packing of the oppositely charged
positive and negative ions together.
A measure of just how much stabilization results from this packing is given by the lattice
energy (U)
Lattice Energy (U)
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Lattice energy is the amount of energy released when gaseous ions combine to form an ionic
solid.
It can also be defined as the energy required to separate one mole of a solid ionic compound
into its gaseous ions.
Lattice energy is a measure of the strength of the forces between the ions in an ionic solid;

higher lattice energy indicates stronger ionic bonds and greater stability of the ionic
compound.
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It is an important indication of the strength of ionic interactions
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It is a major factor influencing melting points, hardness, solubility of ionic compounds.
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The lattice energies of many compounds have been determined using Hess's law of heat
summation
states that an overall reaction's enthalpy change is the sum of the enthalpy changes for the
individual reactions that make it up:
∆H total = ∆H1 + ∆H2 + ∆H3 +...
Born-Haber Cycle
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The Born–Haber cycle is an approach to analyze reaction energies.
The cycle is concerned with the formation of an ionic compound from the reaction of a metal
(often a Group I or Group II element) with a halogen.
Born–Haber cycles are used primarily as a means of calculating lattice energies (or more
precisely enthalpies) which cannot otherwise be measured directly.
The lattice enthalpy is the enthalpy change involved in the formation of an ionic compound
from gaseous ions
A Born–Haber cycle applies Hess's law to calculate the lattice enthalpy by comparing the
standard enthalpy change of formation of the ionic compound (from the elements) to the
enthalpy required to make gaseous ions from the elements.
The enthalpy of formation of lithium fluoride (NaF) from its elements lithium and fluorine in
their stable forms is modeled in five steps in the diagram:
1. Enthalpy change of atomization enthalpy of sodium
2. Ionization enthalpy of sodium
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3. Atomization enthalpy of fluorine
4. Electron affinity of fluorine
5. Lattice enthalpy
The Born-Haber cycle for NaF(s).
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Consider the Born-Haber cycle for the formation of sodium fluoride.
There are two routes to follow: either the direct combination reaction (Route A) or the multistep cycle (Route B)
From Hess's law, it is known that both routes involve the same overall enthalpy change. The
overall enthalpy change for the process (∆fH°) is calculated as the sum of the enthalpy
changes ∆H°step1, through ∆H°step4.
Therefore, ∆H°step5, the lattice energy (UNaF), can be calculated.
∆fH° of NaF(s) (Route A) = sum of ∆H° for steps in cycle (Route B)
To preview Route B
Step 1: Metallic sodium is vaporized to a gas of sodium atom: (atomization)
Na(s) → Na(g) ∆H°step1 = + 108 kJ
Step 2: Converting fluorine molecule to fluorine atoms involves breaking the
covalent bond in F2, so it requires energy:
One mole of F atoms are needed to form one mole of NaF, so start with ½ mol F2:
1/2 F2(g) → F(g) ∆H°step2 = 1/2 bond energy (BE) of F2
=1/2 (154 KJ) = 77.0 KJ
Step 3 : Removing the 3s electron from Na to form Na+ requires energy
Na(g) → Na+(g) + e– ∆H°step3 = IE1 = + 495 kJ/mol
Step 4: Adding an electron to F to form F– releases energy:
F(g) + e– → F– (g) ∆H°step4 = EA = –328 kJ
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Step 5: Forming the crystalline ionic solid from the gaseous ions is the step whose
enthalpy changes (the lattice energy) is unknown:
Na+(g) + F– (g) → NaF(s) ∆H°step5 = U NaF (lattice energy) = ?
 It is known that the enthalpy change of the formation reaction (Route A),=∆H°f = –574 kJ
Therefore, we calculate the lattice energy by using Hess's Law:
Solving for UNaF gives
UNaF = ∆fH° – (∆H° step1 +∆H° step2 ∆H° step3 +∆H° step4 )
= –574 kJ/mol – [108 kJ /mol + 77 kJ/ mol + 495 kJ/mol + (–328 kJ/mol)]
= –926 kJ/mol
Exercise 2.3 Page - 74
Factors Affecting Formation of Ionic Bonding
Ionization energy (IE):
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low ionization energy of metallic elements favours the formation of an ionic bond.
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That is why alkali and alkaline earth metals form ionic compounds.
Electron affinity (EA):
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The elements having higher electron affinity favour the formation of an ionic bond.
Halogens have high electron affinities, and therefore halogens generally form ionic
compounds when they react with metals.
Lattice energy: (U )
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Larger lattice energy would favour the formation of an ionic bond.
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Lattice energy thus is a measure of coulombic attractive force between the combining ions.
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The lattice energy of an ionic compound depends directly on the product of the ionic charges
(q1 × q2), and inversely on the distance (r) between them.
where q1 and q2 are the charges on +ve and –ve ions respectively, r is the distance between
the charges q1 and q2.
Thus, small ions having higher ionic charge shall have larger lattice energy.
If the total energy released is more than that which is absorbed, then the formation of ionic
compound is favored.
Exceptions to Octet Rule in Ionic Compounds
i. Less than Octet (Central Atom is Deficient of Electrons)
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Ions of some elements which are near to helium in the periodic table do not obey the octet
rule.
The tendency of these atoms (H, Li, Be and B) is to attain an arrangement of two electrons
like the noble gas He (duplet configuration), which is also a stable configuration.
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Hydride ion (H+ ), lithium ion (Li+), beryllium ion (Be2+) and boron ion (B3+) are
isoelectronic with He.
Therefore, compounds like LiH, BeCl2 and BF3 are stable in spite of short of electrons
around the central, number of electrons around Li, Be and B is 2, 4 and 6 respectively.
ii. More than Octet (18-Electron Rule):
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The ions of some transition elements and post-transition elements do not usually obey the
octet rule.
For transition metals, the 18-electron rule replaces the octet rule, due to the involvement of d
orbitals of these atoms.
The atoms of these elements would have to lose a large number of electrons to achieve the
noble-gas configurations.
It is important to note that when these atoms form positive ions, electrons are always lost
first from the shells with the highest value of the principal quantum number (n).
Electron Configurations of Iron:
26Fe : 1s
2
2s2 2p6 3s2 3p6 4s2 3d6
26Fe
2+
: 1s2 2s2 2p6 3s2 3p6 3d6
26Fe
3+
: 1s2 2s2 2p6 3s2 3p6 3d5
Electron configurations of zinc:
30Zn
2+

: 1s2 2s2 2p6 3s2 3p6 3d10
Generally, these properties are exhibited by ions of elements from
i.
Group IB and Group IIB (transition elements) and
ii. Group IIIA and Group IVA (heavier post-transition elements)
Exercise 2.4 Page - 78
Properties of Ionic Compounds
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Ionic compounds are crystalline solids at room temperature.
Crystalline ionic solids are usually brittle and non-conductors of electricity, although
molten crystals may be good conductors.
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They usually have high melting and boiling points.
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Ionic compounds are nonvolatile.
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Ionic compounds are usually soluble in inorganic solvents (water is the most common
solvent for ionic compounds) but insoluble in organic solvents like benzene, ethanol and
carbon tetrachloride.
Ionic compounds are very resistant to heat but many will be easily broken by water
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COVALENT BONDING & MOLECULAR GEOMETRY
Formation of Covalent Bonding
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A covalent bond is a bond that is formed by the mutual sharing of one or more pairs of
valence electrons between two combining atoms.
It is an electrostatic force of attraction that exists between positively charged nuclei of atoms
and the shared pairs of electrons in a molecule.
Covalent bonding most commonly occurs between nonmetal atoms which has high
ionization energy or high electronegativity.
Chemical compounds, formed as a result of mutual sharing of electrons thereby establishing
covalent bonds, are called covalent or molecular compounds.
Generally substances that contain covalent bonds are called molecules.
The Lewis formula of Covalent Bonds
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The representation of covalent bonding through Lewis symbols and shared pairs of electrons
is called a Lewis structure.
Lewis electron-dot symbol is a notation in which the symbol of an element represents the
nucleus and inner electrons, and surrounding dots represent the valence electrons.
The Lewis formula;
Show all the valence electrons of the combining atoms.
Illustrate the bonding and non bonding (lone pair) electrons by dots.
Clearly show the satisfaction of octet rule.
Pairs of electrons that is not used for bonding are called lone-pair electrons.
Pairs that are used for bonding are called bonding-pair electrons.
Types of Covalent Bond
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Single bond ; is formed when two atoms are held together by one electron pair.

It is denoted by one dash ( __ )
Eg. H__H , H__Cl
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Double bond is formed when two atoms share two pairs of electrons.
 It is denoted by two dash (= )
Eg. CO2 , O2

Triple bond is formed when two atoms share three pairs of electrons.
 It is denoted by three dash (≡)
Eg. N2, HCN
:N≡N:
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1.
2.
3.
4.
5.
1.
Steps To Write Thr Lewis Formula Of Covalent Bonds
Determine the total number of valence (outer shell) electrons.
 For cations, subtract one electron for each positive charge.
 For anions, add one electron for each negative charge.
Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom.
(Generally, the least electronegative element should be placed in the center.) Connect each
atom to the central atom with a single bond (one electron pair).
Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen),
completing an octet around each atom.
Place all remaining electrons on the central atom.
Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in
order to obtain octets wherever possible
EXAMPLE
Write the Lewis structure of the following and indicate the lone pair and bond pairs of
electrons.
a. CH4
f. SO2
2
b. CO3¯
g. C2H4
c. CCl4
h. CS2
1
d. NO3¯
i . C2H2
e. HCN
Exercise 2.5 Page - 90
Coordinate-Covalent Bonding
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A covalent bond which is formed between two atoms where both the shared pair of electrons
are provided by only one atom is called a coordinate covalent (or dative )bond.
The atom that shares the electron pair is the acceptor atom.
The donor atom must contain at least one lone pair of electrons.
The acceptor atom contains at least one vacant orbital (empty space in its valence shell).
For example,
1.
The ammonium ion, NH4+ , is formed by a coordinate covalent bond in which the two nonbonding electrons on NH3 are bonded with a hydrogen ion, H+ , which has no electron to
contribute.
2.
Sometimes, two or more stable molecules combine to form a molecular complex.
E.g. Combination of NH3 and BF3
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Resonance Structures
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A resonance form is another way of drawing a Lewis dot structure for a given compound.
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Equivalent Lewis structures are called resonance forms.
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They are used when there is more than one way to place double bonds and lone pairs on
atoms.
Resonance structures have the same relative placement of atoms but different locations of
bonding and lone electron pairs.
If two or more possible Lewis structures can be written and the true structure is a composite
or hybrid of them.
The different structures used to represent the molecule or ion are called resonance forms or
resonance contributors or resonance structures.and a two headed resonance arrow (↔)
between them.
Example
 The three resonance structures for carbonate anion (CO3¯2) are drawn below
 The three resonance structures for nitrate anion (NO3¯1):

The resonance structures and resonance hybrid of ozone (O3) are drawn below
Exercise 2.6 Page - 93
Exceptions to the Octet rule in Covalent Bonding
i. Electron-deficient Molecules
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Electron-deficient molecules have a central atom that has fewer electrons than needed for a
noble gas configuration.
The molecules containing central atoms from Group IIA and IIIA. BeCl2, BF3 and AlCl3,
whose Lewis formula are shown below, are typical examples:
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ii.
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More than octet (central atom has excess of electrons):
Molecules whose central atoms have more than eight electrons (Expanded octet).
This group consists of molecules containing central atoms from periods 3, 4, 5, and 6. PF5,
SF6 and XeF4 are typical examples of this type;
iii. Molecules containing an odd number of electrons:

Even if stable molecules of this kind are rare, they do exist.
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Some examples are ClO2 , NO and NO2 having 19, 11 and 17 valence electrons respectively.

The best way to suggest the Lewis-like structure for these molecules is:
E.g. ClO2
Exercise 2.7 Page - 94
Polar and Non-Polar Covalent Molecules
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A covalent bond in which the electrons are shared equally between the two atoms is called a
non-polar covalent bond.
Non-polar is formed when two similar atoms or atoms which differ slightly in their
electronegativity.
Eg. H2, O2, F2, CH4, etc.
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A chemical bond in which shared electrons spend more time in the vicinity of one atom than
the other is called a polar covalent bond, or simply a polar bond
Polar covalent bond is formed between atoms of different electronegativity , the electrons
are not equally shared.
Eg. HCl, HF, H2O, NH3 etc.
The minimum electronegativity value difference between the two atoms in a polar covalent
bond is generally considered to be around 0.4 on the Pauling scale..
When the difference in electronegativity is between approximately 0.4 and 1.7, the bond is
classified as polar covalent.
If the electronegativity difference is less than 0.4, the bond is usually considered non-polar
covalent, while a difference greater than 1.7 typically indicates an ionic bond.
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Bond Polarity and Dipole Moment
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The bond polarity can also be shown by drawing an arrow so that the head points toward the
negative end (F) of the bond and the crossed tail indicates the positive end (H).
The dipole moment (µ) is defined as the product of the magnitude of the charge (δ) at either
end of the dipole multiplied by the distance (d) that separates the charge.
µ= δ× d

The SI-unit of dipole moment is coulomb-metre (C.m).

Dipole moments are often expressed in the non-SI unit debye (D),
where 1D = 3.33564 × 10–30 C.m.
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A very polar molecule is one with a large dipole moment, while a non-polar molecule will
have a zero-dipole moment.
Bond Polarity is a measure of inequality in the sharing of bonding electrons.
The Bond Polarity is a vector, pointing from the atom with less electro negativity to the atom
with larger one.
If the vector sum of Bond Polarities of a molecule is not zero as in HCCl3, but in CCl4 the
vector sum of Bond Polarities of a molecule is zero.
Exercise 2.8 nPage - 96
Properties of Covalent Compounds
Covalent compounds share the following properties:
1.
2.
3.
4.
5.
Covalent compounds are generally liquids or gases at ordinary temperature.However, some
covalent molecules like iodine are solids at room temperature.
Eg. CO2, CH4,NH3 etc are gases.
Ethanol, H2O etc are liquid
Covalent compounds have relatively lower melting points and boiling points, as compared to
ionic compounds.
Covalent compounds are generally insoluble in water. Most covalent compounds are soluble
in non-polar solvents. However, some polar covalent compounds can dissolve in water.
Most covalent compounds do not conduct electricity in solid or liquid form because they do
not have free ions or electrons. However, some covalent compounds can conduct electricity
when dissolved in water if they ionize (e.g., acids).
Covalent compounds are volatile.
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