Kossel – Lewis Approach to Chemical Bond
* In the nature, the system with more energy is less stable. In other words, the lowest energy state corresponds to the most stable state of the system.
* So every system tends to attain stability by losing energy.
* Atoms attain stability by the means of bond formation. The process of bond formation is associated with the lowering of energy of the system.
* Kossel and Lewis developed electronic theory of valence or theory of chemical bonding to explain the chemical bond formation between the two atoms.
* According to the electronic theory of valence, every atom tries to attain octet in its valence shell by losing or gaining or by sharing of electrons. This is known as the “Octet Rule”.
* The electrostatic forces of attraction that holds the two oppositely charged ions together is known as “electrovalent bond”.
Ionic Bond
* The formation of ionic bond depends upon the following three key factors: first, the ease of formation of a cation, second the ease of formation of an anion, and finally, the arrangement of the positive and negative ions in the ionic solid.
* Lower the ionisation enthalpy and higher the negative value of electron gain enthalpy, greater is the ease of formation of a cation and anion, that is, greater is the ease of formation of an ionic bond.
* The magnitude of energy released during the formation of crystal lattice called lattice enthalpy provides a measure for determination of the stability of the ionic compound.
Covalent Bond-I
* The bond formed by mutual sharing of electrons between the combining atoms of same or different elements is called a covalent bond.
The following conditions are to be satisfied for the formation of a covalent bond in a molecule:
• Each of the bonded atoms must contribute equal number of electrons to the shared pair.
• A bond is said to be formed when an electron pair is shared between the bonded atoms.
• All the bonded atoms must attain the nearest inert gas configuration as a result of sharing of electrons.
• In case a covalent molecule is formed between different atoms, then the atom with the least electronegativity is placed at the centre and the other atoms will be placed at the terminal positions for a carbon dioxide molecule.
* Formal charge is equal to the total number of valence electrons in the free atom minus total number of non bonding electrons minus half of the total number of bonding electrons.
* Octet rule is not applicable to all the molecules.
* Octet rule is not applicable for the molecules containing odd electrons. It is also not applicable for the molecules in which the central atom has less than eight electrons, that is, molecules that have contracted octets.
* Octet rule is not applicable for the molecules such as phosphorus penta chloride and sulphur hexa fluoride etc.
Covalent Bond –II
* In the Lewis structure, electrons are drawn as dots and the bonding of a pair of electrons is represented by a line.
* The Lewis structure for an atom is the chemical symbol for the atom, surrounded by as many dots as the number of valence electrons.
* The difference between the number of electrons present in that atom in its isolated state and the number of electrons assigned to the atom in the Lewis structure is called the formal charge of the atom.
* In covalent bonding, atoms share electrons so that the number of valence electrons in each atom becomes eight.
* Formation of covalent bonds in gaseous elements involves:
•Formation of cation
•Formation of anion
•Combination of opposite charges
Bond Parameters
* The factors that affect chemical bonds are known as bond parameters.
* The bond parameters are bond length, bond angle, bond enthalpy and bond order.
* Bond length is the distance between the nuclei of the atoms that are bonded together.
* Half the distance between two similar atoms in separate molecules in a solid is called the Van der Waals radius.
* The number of chemical bonds between a pair of atoms in a molecule is called the bond order of the molecule.
* The angle between the orbitals present around the central atom in a molecule is called the bond angle.
* Bond enthalpy is the amount of energy required to break one chemical bond between two gaseous atoms in a molecule.
Resonance Structures
* Resonance structures are two or more Lewis structures for the same molecule that have similar arrangements of atoms and the same number of electrons, and differ only in the arrangements of the electrons.
* The actual structure of a molecule is a hybrid of all the resonance structures possible. It is called a resonance hybrid. This resonance hybrid is more stable than any of the contributing resonating structures.
* The energy of the resonance hybrid is less than any single canonical structure. Thus, resonance stabilises the molecule.
Polarity of Bonds
* A covalent molecule with two oppositely charged poles is called a dipole.
* The crossed arrow above a Lewis structure indicates the direction of electron shift.
* Dipole moment is defined as: “The product of the magnitude of the charge and the distance between the centres of the positive and negative charges”.
* Dipole moment is a vector quantity, because it has both magnitude and direction.
* The tendency of an anion to get polarized by a cation is called “polarisability”.
* If the extent of polarisation is small, then the bond is more ionic in nature, but if the extent of polarisation is more, then the bond is more covalent in nature.
Vserp’s Theory
* According to VSERP theory, a molecule will adopt a geometry such that the repulsions between the valence electrons around the central atom would be minimum.
* The geometry of the molecule gets distorted due to the presence of lone pair of electrons around the central atom in a molecule.
* On the central atom of a molecule, the repulsion between two lone pairs of electrons is the greatest.
* The repulsion between a bond pair and another bond pair is the weakest, and that between a lone pair and a bond pair is between these two. Thus, lp – lp > lp – bp > bp – bp.
* The molecules having only bond pair of electrons possess regular geometry where as the molecules having one or more lone pair of electrons in addition to bond pairs of electrons possess distorted geometry.
Valence Bond Theory
* According to Valence Bond Theory:
• A covalent bond is formed when pure, valence atomic orbital of one atom overlaps with anotherpure, valence atomic orbital of another atom.
• Each of the overlapping orbitals contains the unpaired electron of opposite spin.
• The electron pair is shared by both the atoms.
• A strong bond is formed when the orbitals of the two atoms overlap to the maximum extent.
* Based on the overlapping of orbitals, two types of covalent bonds are formed. These are known as sigma and pi bonds.
* Sigma bonds are formed by the end-to-end overlap of atomic orbitals along the inter-nuclear axis known as a head-on or axial overlap.
* A pi bond is formed when atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the inter-nuclear axis.
* Sigma bonds are stronger than pi bonds.
* Valence Bond theory fails to determine the actual shapes of the polyatomic molecules, such as methane, ammonia and water which are Tetrahedral, Pyramidal and Bent respectively.
Hybridisation
* Hybridisation is defined as the intermixing of atomic orbitals of slightly differing energies to form a new set of orbitals having equivalent energies.
* The number of hybrid orbitals formed is equal to the number of atomic orbitals combining to form hybrid orbitals.
* Hybrid orbitals form more stable bonds than pure atomic orbitals.
* Hybrid orbitals are formed by the mixing of s, p and d atomic orbitals.
* Hybrid orbitals give a characteristic geometric shape to polyatomic molecules.
• The shape of an sp hybrid orbital is linear.
• The shape of an sp2 hybrid orbital is trigonal planar.
• The shape of an sp3 hybrid orbital is tetrahedral.
• The shape of an sp3 d hybrid orbital is trional bipyramidal.
• The shape of an sp3 d2 hybrid orbital is octahedral.
Molecular Orbital Theory: Features, Lcao
* Atomic orbitals of equivalent energies combine to form molecular orbitals.
* Two atomic orbitals combine to form two molecular orbitals, known as bonding molecular orbital, denoted by ψb and antibonding molecular orbital denoted by ψa.
* Linear Combination of Atomic Orbitals or LCAO is an approximate method that has been adopted to explain systems containing more than one electron.
* The linear combination of atomic orbitals helps describe the formation of molecular orbitals, mathematically.
* When a bonding molecular orbital forms, there is constructive interference because the two electron waves of the bonding atoms reinforce each other.
* The electron waves cancel each other due the destructive interference in an antibonding molecular orbital.
* The total energy of the two resultant molecular orbitals equals that of the original atomic orbitals.
* A linear combination of atomic orbitals forms molecular orbitals, only if they satisfy some conditions.
Molecular Orbital Theory: Energy Level Diagram
* Molecular orbitals are obtained by combining the atomic orbitals of the atoms in a molecule.
* The molecular orbital formed due to constructive interference is called a bonding molecular orbital as the electrons belonging to it are placed between the two nuclei.
* Sigma orbitals stabilise a molecule.
* The molecular orbital formed due to destructive interference is called an anti-bonding or sigma star molecular orbital as the electrons belonging to it are placed away from the region between the two nuclei, making the molecule less stable.
* As in atoms, electrons in a molecule too first occupy lower energy levels.
* For molecules formed by atoms containing more than two electrons, we will only consider the molecular orbitals formed by their valence shell and not the ones formed by their core orbitals to understand their stability.
* The bond number obtained by using the molecular orbital model is different from the bond number obtained using the Lewis structure because in the Lewis structure, all the electrons are paired and here, two of the electrons are unpaired.
* When electrons in an atom or molecule are paired, they form a diamagnetic atom or molecule and when the electrons in an atom or molecule are not paired, they form a paramagnetic atom or molecule.
Molecular Orbital Theory: Bonding In Homonuclear Diatomic Molecules
* The electronic configuration of a hydrogen molecule is σ 1s2 and the bond order is 1.
* A hydrogen molecule is diamagnetic, as no unpaired electrons are present in it.
* The electronic configuration of a helium molecule is σ 1s2, σ*1s2.
* The bond order of helium is 0, which makes it unstable, and, therefore, nonexistent.
* The electronic configuration of a lithium molecule is σ 1s2, σ*1s2, σ 2s2.
* The bond order of lithium is 1, it is a stable molecule and is diamagnetic in nature.
* The electronic configuration of a carbon molecule is σ 1s2, σ*1s2, σ 2s2 , σ*2s2 , π2 p_x^2 = π2p_y^2
* The bond order of carbon is 2 and it is diamagnetic in nature.
* The bond order of oxygen is 2 and it is paramagnetic in nature.
Hydrogen Bonding
* A bond that forms between a partially positive charged hydrogen atom and a more electronegative atom is known as a hydrogen bond.
* Hydrogen bonds are weaker than covalent bonds.
* The physical state of a compound determines the strength of a hydrogen bond.
* Hydrogen bonding is very strong in the solid state but minimal in the gaseous state.
* Hydrogen bonds are either intermolecular or intramolecular.