Chemical Bonding B.Sc. 1st Semester

Chemical Bonding B.Sc. 1st Semester

Chemical Bonding B.Sc. 1st Semester

Chemical Bonding

(i) lonic bond: General characteristics, types of ions size effects, radius ratio rule and its limitations. Packing of ions in crystals. Bom-Lande equation with derivation and importance of Kapustinskii expression for lattice energy. Madelung constant, Born Haber cycle and its application, Solvation energy.

(ii) Covalent bond: Lewis structure, Valence Bond Theory (Heitler-London approach). Energetics of hybridization, equivalent and non-equivalent hybrid orbitals. Bent's rule, Resonance and resonance energy, Molecular orbital theory. Molecular orbital diagrams of diatomic and simple polyatomic molecules N2, 02,C2, B2, F2 & CO, NO and their ions; hydrogen chloride, berrylium fluoride, carbon dioxide, (idea of s-p mixing and orbital interaction to be given). Formal charge, Valence shell electron pair repulsion theory (VSEPR), shapes of simple molecules and ions containing lone pairs and bond pairs of electrons, multiple bonding (sigma and pi bond approach) and bond lengths.

Covalent character in ionic compounds, polarizing power and polaizability. Fajan's rule and consequences of polarization. Ionic character in covalent compounds: Bond moment, dipole moment and electronegativity difference.

(iii) Metallic bond: Qualitative idea of valence bond and band theories. Semiconductors and insulators, defects in solids.

(iv) Weak Chemical Forces: van der Waals forces, ion-dipole forces, dipole-dipole interactions, induced dipole interactions, Instantaneous dipole-induced dipole interactions. Repulsive forces, Hydrogen Bonding (theories of hydrogen bonding. valence bond treatment) Effects of chemical force, melting and boiling points, solubility energetics of dissolution process.

Hydrogen Bonding

Hydrogen bond is a type of dipole-dipole interaction between very high electronegative atom (i.e.N, O and F) and a hydrogen atom bonded to another electronegative atom. This bond always involves a hydrogen atom. Hydrogen bonds can occur between molecules or within a single molecule.
Hydrogen bond is stronger than van der Waals forces, but weaker than covalent bonds or ionic bonds. It is about 5% the strength of the normal covalent bond formed between O-H.

Types of Hydrogen Bonding

There are two types of Hydrogen bonds. They are-
1. Intermolecular Hydrogen Bonding
2. Intramolecular Hydrogen Bonding

1. Intermolecular Hydrogen Bonding:

When hydrogen bonding takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding. For example – hydrogen bonding in water, alcohol, ammonia, p-nitrophenol etc.
Intermolecular Hydrogen Bonding

2. Intramolecular Hydrogen Bonding:

The hydrogen bonding which takes place within a molecule itself is called intramolecular hydrogen bonding.
It takes place in compounds containing two groups such that one group contains hydrogen atom linked to an electronegative atom and the other group contains a highly electronegative atom linked to a lesser electronegative atom of the other group. The bond is formed between the hydrogen atoms of one group with the more electronegative atom of the other group.
For example – ortho-nitrophenol, Salicylic acid, Salicyldehyde etc.
Intramolecular Hydrogen Bonding

Limiting Radius Ratio:

The limiting radius ratio is the minimum allowable value for the ratio of cationic radii to anionic radii (ρ=r+/r-) for the structure to be stable. Here, r+ is the radius of the cation and r- is the radius of the surrounding anions.
Limiting radius ratio

Formal Charge

Formal charge is the difference between the number of valence electrons in an isolated atom and number of electrons assigned to that atoms in Lewis structure.
Formal charge = [Total number of valence electrons in the free atom ) - (Total number of nonbonding electrons) -1/2(Total number of shared electrons i.e. bonding electrons)]
Formal Charge in Ozone

Formal charge of the oxygen atom 1 = 6 - 4 - 2 = 0
So the formal charge on oxygen atom 1 is zero.
Formal charge on oxygen atom 2 = 6 - 2 - 3 = +1
So the formal charge on oxygen atom 2 is +1.
Formal charge on oxygen atom 3 = 6 - 6 - 1 = -1
So the formal charge on oxygen atom 3 is -1
Now the total formal charge of the ozone = 0 + 1 - 1 = 0
Therefore the formal charge of ozone is '0'.

Fajan's Rule

Kazimierz Fajans in 1923, gave some important points to predict whether a chemical bond is expected to be predominantly ionic or covalent.They are given below-
1. Size of Cations
2. Size of Anions
3. Charge on ions (Cations and Anions)
4. 18 electron configuration

1. Size of Cations
Smaller the size of cation, greater the covalent character.
Example: LiCl, NaCl, KCl, RbCl, CsCl
LiCl is most covalent among the given IA chlorides as the size of Li+ is smaller than that of other IA cations.
The order of covalent character is-
LiCl > NaCl > KCl > RbCl > CsCl
2. Size of Anions
Larger the size of anions, greater the covalent character.
Example: LiF, LiCl, LiBr, LiI
LiI is most covalent among the given IA halides as the sixe of iodide ion is larges.
So, the order of covalent character is-
LiI > LiBr > LiCl > LiF

3. Charge on ions (Cations and Anions)
Higher the charge on ions, greater the covalent character.
Example: NaCl, MgCl2, AlCl3
AlCl3 is most covalent among the given molecules as the chage on Al is +3 highest (charge on Na is +1 and on Mg is +2).
So, the order of covalent character is-
AlCl3 > MgCl2 > NaCl
4. 18 electron configuration
Molecules which follows 18 electron configuration is more covalent in nature than those molecules which follows 8 electron configuration.
Example: NaCl and CuCl
CuCl is more covalent as it follows 18 electron configuration while NaCl is more ionic because it follows 8 electron configuration.

According to Fajan's rules, the covalent nature of ionic compounds is favoured by

a. Large cation and small anion
b. Large cation and large anion
c. Small cation and large anion
d. Small cation and small anion

Maximum covalent character is associated with the compound

a. NaI
b. MgI2
c. AlCl3
d. AlI3

Amongst LiCl, RbCl, BeCl2 and MgCl2 the compounds with the greatest and the least ionic character, respectively, are

a. LiCl and RbCl
b. RbCl and BeCl2
c. RbCl and MgCl2
d. MgCl2 and BeCl2

Born Haber Cycle

Energy change associated with the formation of an ionic compound in a regular crystal lattice in a systemic manner constitute a cycle called Born Haber Cycle named after Max Born and Fritz Haber, who used this method for calculating lattice energies of crystals.
The lattice energy of sodium chloride, for example is the change in enthalpy, ∆H, when Na+ and Cl- ions in the gas phase come together to form one mole of NaCl crystal.
Sodium chloride can be obtained in the following steps-

Valence Bond Theory (VBT)

This theory was proposed by Heitler and London to explain the formation of covalent bond quantitatively using quantum mechanics. Later on, Linus Pauling improved this theory by introducing the concept of hybridization.
The main postulates of this theory are as follows-
1. It deals with the electronic configuration of the elements.
2. The valency of an element is the number of unpaired electrons present in valence shell of ita atom.
3. The paired electrons of the valence shell does not take part in the bond formation.
4. A covalent bond is formed by the overlapping of two half filled valence atomic orbitals of two different atoms.

5. The electrons in the overlapping orbitals get paired and confined between the nuclei of two atoms.
6. The electron density between two bonded atoms increases due to overlapping. This confers stability to the molecule.
7. Greater the extent of overlapping, stronger is the bond formed.
8. The direction of the covalent bond is along the region of overlapping of the atomic orbitals i.e., covalent bond is directional.
9. There are two types of covalent bonds (sigma and pi) formed based on the pattern of overlapping.

Limitations of Valence Bond Theory

Followings are the limitations of this theory-

1. It does not explain the paramagnatism of oxygen and some other molecules.
2. It does not explain the coordinate bond formation.
3. It fails to explain the bond formation in B2H6, H2+, He2+ etc.
4. It totally fails to explain color, kinetics, thermodynamics and structural properties of compounds.

Valence Sheel Electron Pair Repulsion Theory (VSEPR Theroy)

This theory was proposed by Sidwick and Powell in 1940 and later on modified by Gillespie and Nyholm in 1957.
This theory predict the shape of simple molecules and ions on the basis of repulsion of electron pairs present in the valence shell of their central atom.

Some Important Postulates of this Theory are given below
1. Electrons involved in the bond formation is called bonding electrons or bond pair (B.P.) and the rest electrons are called lone pairs (L.P.).
2. Electron pairs in valence shell repel one-another as electron clouds are negatively charged.
3. These electron pairs occupy the space at maximum distance for minimum repulsion.
4. The most stable geometrical arrangement of 2,3,4,5,6 electron pairs is linear, triangular, tetrahedral, triangular bipyramidal and octahedral respectively.
5. The central atom in a molecule is surrounded by only B.P. the molecule has regular or symmetrical geometry but in case of B.P. and L.P. the molecule does not have regular geometry.
6. A lone pair occupies more space than a bond pair because lone pair attached with only one atom. Hence the order of repulsion is-
L.P - L.P. > L.P - B.P. > B.P - B.P.

Greater the repulsion, smaller the bond angle.
7. Multiple bonds does not affect the gross geometry of the molecules rather the geometry is exclusively decided by B.P. and L.P.
8. A lone pair and double bond repulsion is much greater than a lone electron and double bond repulsion.
9. A lone pair and a single bond repulsion is larger than a lone pair and double bond repulsion.

Geometries of molecules from VSEPR Theory

Geometries of molecules from VSEPR Theory

Molecular Orbital Theory (MOT)

This theory was proposed by F.Hund and R.S.Mulliken in 1932 and the basic features of this theory are given below-

1. Electrons of the molecule are present in various molecular orbital just like the electrons of an atom are present in atomic orbitals.
2. Molecular orbitals are formed by mixing of atomic orbitals of comparable energies and proper symmetry.
3. An electron in an atomic orbital is influenced by only one nucleus and thus it is monocentric while in a molecular orbital it is influenced by two or more nuclei of the molecule and thus it is polycentric. Continue...

Molecular Orbital Diagram

Molecular orbital diagram is a qualitative descriptive tool explaining chemical bonding in molecules in terms of molecular orbital theory in general and the linear combination of atomic orbitals (LCAO) method in particular.
Moecular Orbital Diagram
Molecular orbital diagram is a molecular orbital energy levels diagram, shown as short horizontal lines in the center, flanked by constituent atomic orbital energy levels through dotted lines for comparison, with the energy levels increasing from the bottom to the top. Degenerate energy levels are commonly shown side by side. Appropriate atomic orbital and molecular orbital levels are filled with electrons by the Pauli Exclusion Principle, symbolized by small vertical arrows whose directions indicate the electron spins. Continue...

Dipole - Dipole Forces

Dipole-dipole forces act between molecules possessing the permanent dipole (i.e. polar molecules). The polar molecules interact with neighbouring molecules. This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. The attractive force decreases with the increase of distance between the dipoles. The interaction energy is also inversely proportional to distance between polar molecules. Dipole-dipole interaction energy between stationary polar molecules (as in solids) is proportional to 1/r3 and that between rotating polar molecules is proportional to 1/r6, where r is the distance between polar molecules. Besides dipole-dipole interaction, polar molecules can interact by London forces also. Thus cumulative effect is that the total of intermolecular forces in polar molecules increase.
dipole-dipole interaction