RBSE Class 11 Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure

These comprehensive RBSE Class 11 Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure will give a brief overview of all the concepts.

RBSE Class 11 Chemistry Chapter 4 Notes Chemical Bonding and Molecular Structure

→ Chemical Bond: The attractive force which holds various atoms or ions together in different chemical species is called chemical bond.

→ Ionic or Electrovalent Bond: The bond formed as . a result of the electrostatic attraction between the positive ions and negative ions is termed as ionic bond.

→ Lattice Energy: This is the amount of energy required to separate one mole of a solid ionic compound completely into gaseous constituents.

→ Solvation Energy: The amount of energy released during the process of dissociation of ionic compounds into its constituent in presence of solvent.

→ Covalent Bond: It is the bond formed between two atoms by sharing the electrons or electron pair. Each combining atom contributes atleast one electron to shared pair.

→ Homoatomic Molecule: When two same type of atoms combine to form molecule, it is called homoatomic molecule, e.g. ,H₂, O₂, Cl₂, N₂ etc.

→ Heteroatomic Molecule: When different types of atoms combine, then they form heteroatomic molecule, e.g., HCl, HF, NH₃ etc.

→ Non polar Covalent Bond: When shared pair of electron situated at centre of both atoms, then both atoms are neutral. This type of bond is known as non polar covalent bond. For example: H₂, N₂, Cl₂, O₂ etc.

 

→ Polar Covalent Bond: When shared pair of electron shifted towards more electronegative atom and it attains partial negative charge (δ^-) while other atom becomes partially positively charged (δ⁺). Such type of covalent bond is known as polar covalent bond, e.g., HF, HCl, NH₃ etc.

→ Coordinate or Dative or Semipolar Bond: In this type of bond, the shared electron pair comes from one atom only, other atom only shares it. Both the atoms have equal sharing on the electron and thus both atoms acquire nearest noble gas configuration.

→ Bond length: Bond length is defined as the equilibrium distance between the nuclei of'two bonded atoms in a molecule.

→ Bond Angle: It is defined as the angle between the orbitals containing bonding electron pair around the central atom in a molecule/complex ion. Bond angle is expressed in degree.

→ Bond Enthalpy: It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol-1.

→ Bond Order: In the Lewis description of covalent bond, the bond order is given by the number of bonds between the two atoms in a molecule.

→ Resonance Energy: Resonance energy is equal to the difference between actual bond energy and energy of most stable resonating structure. Resonance Energy = Actual bond energy - Energy of most stable resonating structure.

→ Resonance: Whenever a single Lewis structure can not describe the properties of a molecule accurately, a number of structures with similar energy, position of nuclei, bonding and non bonding pairs of electrons are taken as the canonical structure of the hybrid which describe the properties of molecule accurately.

→ Regular Geometry: Molecule in which the central atom has no lone pair, then its geometry will be called as regular geometry.

→ Irregular Geometry: In a molecule, if there is lone pair along with bonded electron pair at its central atom. The geometry of molecule will be irregular or distorted.

→ Dipole Moment: It is defined as the product of the magnitude of the charge and the distance between the centre of positive and negative charge.

→ Sigma Bond: This type of covalent bond is formed by the end to end (head on) overlapping of bonding orbital along the internuclear axis.

→ pi-bond: It is formed when atomic orbitals overlap in such a way that their axis remain parallel to each other and perpendicular to the internuclear axis. π-bonds are formed by sidewise overlapping.

→ Hybridization: Hybridization is the phenomenon of intermixing of orbitals of slightly different energies so as to redistribute their energies to give new set of orbitals of equivalent energy and shape. The atomic orbitals taking part in hybridisation are called hybrid orbital.

→ Molecular Orbital: When two atomic orbitals intermix or overlap then they lose their identity and form new orbital which is called molecular orbital. Molecular orbital is the space around nucleus of atom where probability of finding of electron is maximum.

→ Bonding Molecular Orbital: The bonding molecular orbital has lower energy than combining atomic orbitals.

→ Anti-bonding Molecular Orbital: The energy of antibonding molecular orbital is higher than combining atomic orbitals.

→ Bond Order: It is defined as one half of the difference between the number of electrons present in bonding and antibonding molecular orbitals.
Bond order = \(\frac{1}{2}\)(N_b - N_a).

→ Hydrogen Bond: Hydrogen bond is attraction force which binds one H-atom of a molecule with more electronegative atom (F,0,N). It is denoted by dotted line ( ).

→ Intermolecular Hydrogen Bond: This type of hydrogen bond is formed between two different molecules of the same or different compounds.

→ Intramolecular Hydrogen Bond: It is formed when hydrogen atom is in between the two highly electronegative (F, 0, N) atoms present within the same molecule.

→ Some Important Facts

• Atoms take part in chemical combination to complete their octet or to attain electronic configuration of noble gases.
• Metals and non-metals combine with each other to form ionic bonds. For example, alkali metals combine with halogens to form ionic bonds.
• Non-metals combine with each other and form covalent bond.
• The electrons of valence shell are represented by dots in Lewis structure.
• Sigma bonds are formed by axial overlapping of orbital while pi-bonds are formed by lateral overlapping.
• Sigma bonds are stronger bond as compared to pi-bonds.
• pi bond is always formed along with sigma bond.
• In single bond, there is only sigma bond, in double bond there is one sigma bond and one pi bond whereas in triple bond, there is one sigma bond and two pi-bonds.
• VSEPR theory tells about geometry and shape of co’valent compounds.
• Hybridisation can be expressed by following formula:
• \(\frac{1}{2}\)[Number of electron in valence shell of central atom] - [Number of valence electrons around central atom]
• Formal charge = [Total number of valence electrons in the free atom] - [Total number of non-bonding (lone pair) electrons] - \(\frac{1}{2}\) [Total number of bonding (shared) electrons]
• Any ionic or covalent bond is not 100% ionic or covalent in nature. Ionic bond has some covalent character and covalent bond has some ionic character. This character can be determined by Fajan’s rule.
• The electronegativity of elements can be expressed by both Pauling and Mulliken’s scale.
• Valence bond theory was first proposed by Heitler and London and it was developed by Pauling and Slater.
• Molecular orbital theory was given by Hund and Mulliken.
• Percentage of ionic character = \(\frac{\text { Experimental value of dipole moment }}{\text { Theoretical value of dipole moment }}\) × 100

→ Hybridisation gives information about geometry of molecules:

Hybridization	Geometry
sp	Linear
sp2	Trigonal planar
sp3	Tetrahedral
sp3d	Trigonal bipyramidal
sp3d2	Octahedral
sp3d3	Pentagonal bipyramidal
dsp2	Square planar

→ Hydrogen bonds are of two types:

• Intermolecular hydrogen bond.
• Intramolecular hydrogen bond.

→ Bond order = \(\frac{N_b-N_a}{2}\)
N_b = Number of electrons in bonding molecular orbital.
N_a = Number of electrons in antibonding molecular orbital.

→ Bond order ∝ Stability of molecule or ion

→ Bond order ∝ \(\frac{1}{\text { Bond length }}\)

→ Bond order oc Bond dissociation energy.

→ Bond order may be whole number or fractional. Molecule which have positive bond order are stable.

→ If there are unpaired electrons in molecular orbitals then molecule will be paramagnetic in nature while absence of unpaired electron shows the diamagnetic nature of molecule/ion.