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Chemical Bonding Theories Explained
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"
4.
5.
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.
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.
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.
Editor's Notes
we will describe the interactions between atoms called chemical bonds. Most of the discussion centers on the Lewis theory, which provides one of the simplest methods of representing chemical bonding.
We almost never write Lewis structures for ionic compounds, except when we want to emphasize the ratio in which the ions combine. The structures of ionic compounds are much more complicated than is suggested by the Lewis structure. See, for example, the structure of NaCl
we have used a single pair of electrons between two atoms to describe a single covalent bond.
Often, however, more than one pair of electrons must be shared if an atom is to attain an octet (noble gas electron configuration).
N has 7 electrons = 1s2, 2s2, 2p3
So, the outer shell has 5 electrons
The triple covalent bond in is a very strong bond that is difficult to break in a chemical reaction. The unusual strength of this bond makes quite inert.
As a result, coexists with in the atmosphere and forms oxides of nitrogen only in trace amounts at high temperatures.
The lack of reactivity of with is an essential condition for life on Earth. The inertness of also makes it difficult to synthesize nitrogen compounds.
Most chemical bonds fall between the two extremes of 100% ionic and 100% covalent.
that has a uniform distribution of electron charge density as depicted by the uniform color distribution in the electrostatic potential map. This is typical for a nonpolar covalent bond and occurs in all diatomic molecules containing identical atoms. The sodium chloride molecule, conversely, exhibits a highly nonuniform distribution of electron charge density. The sodium atom is almost exclusively in the blue extreme of positive charge and the chlorine in the red extreme of negative charge. This electrostatic potential map is typical of an ionic bond, yet it is clear from the map that the transfer of electron density from the sodium atom to the chlorine atom is not complete. That is, the NaCl bond is not completely ionic. Experiments show that the bond is only about 80% ionic. The molecule HCl also has an unsymmetrical distribution of electron charge density, as indicated by the gradation of color in the electrostatic potential map. Note, however, that in this case the chlorine atom is not completely in the extreme dark red corresponding to a large negative charge. Instead, it is in the orange-red region indicating a partial negative charge. Correspondingly, the hydrogen atom has a partial positive charge, as indicated by the pale blue. The electrostatic potential map clearly depicts the polar nature of the bond in HCl.
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used.
One widely used electronegativity scale, with values given in Figure 10-6, is that devised by Linus Pauling