The document discusses chemical bonding theories, primarily focusing on Lewis's theory, which explains how electrons are transferred or shared to form ionic and covalent bonds. It includes details about Lewis structures, coordinate covalent bonds, polar covalent bonds, and electronegativity, emphasizing how these concepts impact the stability and characteristics of molecules. Additionally, it notes the relationship between electronegativity and the metallic/nonmetallic character of elements in the periodic table.

1. Chemical Bonding Theories
Dr. K. Shahzad Baig
Memorial University of Newfoundland
(MUN), Canada
Petrucci, et al. 2011. General Chemistry: Principles and Modern Applications. Pearson Canada Inc., Toronto, Ontario.
Tro, N.J. 2010. Principles of Chemistry. : a molecular approach. Pearson Education, Inc

2. Lewis s theory
1. Electrons, especially those of the outermost (valence) electronic shell, play a
fundamental role in chemical bonding.
2. In some cases, electrons are transferred from one atom to another. Positive and
negative ions are formed and attract each other through electrostatic forces called
ionic bonds.
3. In other cases, one or more pairs of electrons are shared between atoms. A bond
formed by the sharing of electrons between atoms is called a covalent bond.
Some fundamental ideas associated with Lewis s theory follow:
Electrons are transferred or shared in such a way that each atom acquires an especially
stable electron configuration. Usually this is a noble gas configuration, one with eight
outer-shell electrons, or an octet.

3. Lewis structure and Lewis symbols
A Lewis structure is a combination of Lewis symbols that represents either the
transfer or the sharing of electrons in a chemical bond.
A Lewis Symbol consists of the:
• element symbol
• the number of electrons in the outer energy level as
represented by a Bohr Diagram
• electrons are placed outside the symbol as "dots"

6. Lewis Structures for Ionic Compounds
For an ionic compound of a main-group element, the Lewis symbol of
(1) the metal ion has no dots if all the valence electrons are lost,
(2) the ionic charges of both cations and anions are shown
•When sodium loses its only valence electron to become an ion,
the Lewis structure shows it with no dots (electrons).
The Na and Cl are near each other but the two dots from
the Cl should not be interpreted as a covalent bond.

7. Lewis structures for the compounds:
(a) BaO; (b) MgCl2 ; (c) aluminum oxide

8. Covalent Bonding: An Introduction

9. Coordinate covalent bond
A covalent bond in which a single atom contributes both of the electrons to a shared
pair is called a coordinate covalent bond.

10. The H atom of HCl leaves its electron with the Cl atom and,
as H+, attaches itself to the NH3 molecule through the lone-pair electrons on the N atom.

11. The ions NH3 and Cl are formed and are formed.

12. Multiple Covalent Bonds
Some molecules are not able to satisfy the octet rule by making only single covalent bonds
between the atoms.

13. Each N atom appears to have only six outer-shell electrons,
not the expected eight
The situation can be corrected by bringing the four unpaired electrons into the region
between the N atoms and using them for additional bond pairs.
The bond between the N atoms in is a triple covalent bond Double and Triple covalent
bonds are known as Multiple covalent bonds

14. Polar Covalent Bonds and
Electrostatic Potential Maps
In such a bond, electrons are displaced toward the more nonmetallic element. The
unequal sharing of the electrons leads to a partial negative charge on the more
nonmetallic element, 𝛿−and a corresponding partial positive charge on the more
metallic element, 𝛿+
A covalent bond in which electrons are not shared equally between two atoms is called a
polar covalent bond.
The polar bond in HCl by a Lewis structure is represented

15. An electrostatic potential map gives information about the
distribution of electron charge in a molecule
Red, is used for regions of the
most negative electrostatic
potential
Blue is used for regions of
the most positive
electrostatic potential
Cl2 has a uniform
distribution of electron
charge density as depicted
by the uniform color
distribution in the
electrostatic potential map.

16. Electronegativity
Electronegativity (EN) of an atom is its ability to attract electrons of other atoms to
which it is bonded.
Electronegativity is related to ionization energy (I) and electron affinity (EA).
Consider the reaction between two hypothetical elements, A and B, which could give the
products A+B- or A-B+
𝐴 + 𝐵 → 𝐴+ 𝐵− ∆𝐸1 = 𝐼𝐴 − 𝐸𝐴 𝐵
𝐴 + 𝐵 → 𝐴− 𝐵+ ∆𝐸2 = 𝐼 𝐵 − 𝐸𝐴 𝐴
It is, expect that ∆𝐸1 = ∆𝐸2

17. Assuming that the resultant bond is nonpolar
𝐼𝐴 + 𝐸𝐴 𝐵 = 𝐼 𝐵 + 𝐸𝐴 𝐴
Equation (10.12) tells us that a nonpolar bond will result when the difference between the
ionization energy and the electron affinity is the same for both atoms involved in the bond.
the electronegativity of the atom
𝐸𝑁𝐴 ∝ 𝐼𝐴 − 𝐸𝐴 𝐴
𝐼𝐴 − 𝐸𝐴 𝐴 = 𝐼 𝐵 − 𝐸𝐴 𝐵 after collecting terms for each atom(10.12)

18. Pauling’s EN values range from about 0.7 to 4.0.
In general,
the lower its EN, the more metallic the element is, and
the higher the EN, the more nonmetallic it is.
The EN decreases from top to bottom in a group and
increases from left to right in a period of the periodic table.
as the ionization energy (I) increases across the period
If ∆EN is large, the bond is ionic.
For ∆EN intermediate, the bond is polar covalent.
Large ∆EN are found between the more metallic and the more nonmetallic elements.
Assess the metallic/nonmetallic characters of the bonded elements from the periodic table

19. Writing Lewis Structures
we have already encountered.
• All the valence electrons of the atoms in a Lewis structure must appear in the
structure.
• Usually, all the electrons in a Lewis structure are paired.
• Usually, each atom acquires an outer-shell octet of electrons. Hydrogen, however, is
limited to two outer-shell electrons.
• Sometimes, multiple covalent bonds (double or triple bonds) are needed. Multiple
covalent bonds are formed most readily by C, N, O, P, and S atoms.