Engineering chemistry textbook chapter 1 chemical bonding

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Chapter 1: Chemical Bonding
“Chemical bonding between atoms reveals the undisclosed nature of matter around us and in the
universe”.
1.0 Introduction
In this chapter you will learn about the most basic concept of chemistry, chemical bonding.
Chemical bonding is termed as an attraction force between two or more atoms/ions to form a
molecule. These bonds are formed by the electromagnetic force of attraction between them.
Chemical bonding is important because:
i) Chemical bonds help in understanding the nature and formation of chemical compounds like
water, salt, etc.
ii) Chemical bonds help us determine properties of various chemical compounds.
For example − If a chemical compound is formed by an ionic bond, then it has a high melting
point.
iii) Knowledge of chemical bonds help in discovering new chemical compounds, like
pharmaceutical drugs.
Chemical Bonding occurs when two atoms having electron deficiency, share, combine or remove
electrons to complete their valence shells. The nature of bonding depends on electronegativity of
the participating atoms. Bonding may be covalent (sharing of electrons) or ionic (removing of
electrons) in nature.
All things in universe are composed of matter, which is made out of the bonding between
different atoms. But then, the question arises, “how do atoms combine to form chemical bonds”.
This, and many more such questions have been answered in the chapter.
1.0.1 History
There are many scientists who have given various theories on chemical bonding. Some of the
more popular theories being
• Kossel-Lewis theory
• Valence Shell Electron Repulsion Theory (VESPER)
• Valence Bond Theory (VB)
• Molecular Bond Theory (MB)
You will learn all these theories in detail later in this chapter. Let us first discuss about bond
energy, which is an important concept in chemical bonding.
1.0.2 Bond Energy
The amount of energy required to break or form one mole of a molecule into its individual atoms
is termed as Bond Energy. To break molecules in a chemical reaction, bond energy is required.
Let us take the example of water:
2 2 22H O 2H + O→

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Here, a water molecule breaks to form two separate molecules, 22H and 2O . The amount of
energy required to break 22H O into 22H and 2O is the bond energy.
In order to break a strong bond, we would need a proportionately higher amount of energy.
1.0.3 FAJAN’S Rule
• Developed by Kazimierz Fajans in 1923.
• This helps in deciding the covalent and ionic nature of chemical bonds, that depends on
the cation charge and relative sizes of the cation and anion.
• Fajan believed that partial ionic character is present in covalent bonds and a partial
covalent character is present in ionic bonds.
• According to Fajan:
o Ionic bond shows more covalent character when cation has high positive charge
with small size while the anion has large size.
o Covalent bond shows more ionic character when cation has low positive charge
with large size and small anion size.
Example 1: NaCl (sodium chloride) is ionic with a low positive charge (+1) and large cation of
size ~1 Å and relatively small anion of size 2 Å. Aluminium iodide ( 3AlI ) is covalent with a
high positive charge (+3) and a large anion.
Fajan’s Rule can be written as:
Ionic Character Covalent Character
Low positive charge High positive charge
Large cation Small cation
Small anion Large anion
Polarization
Polarization deals with sharing of electrons in covalent bond. When two atoms of different
electro negativity bond together, the atom with high electro negativity pulls the shared pair of
electron closer to itself, hence becoming a polar bond.
Example 2: In 3HCF or trifluoromethane, the Carbon electrons are closer to Fluorine because of
higher electronegativity of Carbon. The structure of 3HCF is as shown:
1.0.4 Polar and Non-Polar Molecules
Polar Molecules
A polar molecule is a molecule that has positive and negative ends, or polarity. This polarity is
due to:

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• Unequal electronegativity of atoms.
• Uneven sharing of electron.
When atoms having different electronegativity share electrons, the most electronegative atom
among them pulls the shared pair of electron towards itself; and as such, the molecule becomes
polar.
Example 3: Water is a polar molecule because oxygen shares more electrons than hydrogen due
to its higher electro negativity. The polarity in 2H O molecule is as shown:
Non-Polar Molecules
A non-polar molecule is a molecule that does not have positive or negative ends. This is due to:
• Equal electronegativity of atoms.
• Even sharing of electrons.
When atoms having same electronegativity share electrons, the shared pair occupies symmetrical
position so none of the atoms have charge.
Example 4: The non-polarity in Carbon dioxide or 2CO is as shown:
1.0.5 Orbital Overlapping
The concept of orbital overlapping was given by Heitler and London in 1972 which was further
developed by Pauling.