CBSE Class 11 Chemistry Chapter 4 Revision Notes

Chapter 4: Chemical bonding and Molecular structure Revision Notes

Chemical Bond

• Chemical bond is the force that holds distinct atoms in a molecule together.

The Rule of the Octet

• Chemical reactions bring together atoms from various elements to complete their octet or achieve the noble gas configuration.

Valence Electrons

• It is the outermost shell electron that is involved in chemical reactions.

Chemical Combination Modes

• By the transfer of electrons: An electrovalent bond or ionic bond is a chemical bond created by the full transfer of one or more electrons from one atom to another.
• The relationship established by the equal sharing of electrons between one or two atoms is known as a covalent bond. Both parties donate electrons to these connections.
• Dative bond or co-ordinate bond: A dative bond or co-ordinate bond is established when one atom contributes electrons that are shared by both.

Bonds that are ionic or electrovalent

• The full movement of electrons from one atom to another forms an ionic or electrovalent link. It is usually created through the interaction of metals and non-metals.
• We can say that the oppositely charged ions are held together by the electrostatic force of attraction.

Factors Influencing Ionic Bond Formation

• Ionization enthalpy: The amount of energy necessary to remove one electron from the outermost shell of an isolated gaseous atom in order to transform it into a cation is known as the ionisation enthalpy of any element.
• As a result, the lower the ionisation enthalpy, the simpler it will be to generate a cation and the greater the possibility of forming an ionic connection. As a result, alkali metals are more likely to establish an ionic connection.

I.E = 496 kJ/mole in the production of the Na+ ion, for example.

• Magnesium has a specific heat capacity of 743 kJ/mole. As a result, the production of positive ions in sodium is simpler than in magnesium.
• As a result, we may deduce that the lower the ionisation enthalpy, the greater the odds of forming an ionic connection.
• Electron gain enthalpy (Electron affinities): This is the amount of energy produced when a single gaseous atom accepts an electron to create an anion. The easier it is to generate an anion, the higher the negative electron gain enthalpy. As a result, the likelihood of forming an ionic connection rises.
• As an example. Halogens have a strong affinity for electrons. As a result, anion production is rather prevalent in halogens.
• Lattice energy, also known as enthalpy, is the amount of energy necessary to split one mole of an ionic molecule into two oppositely charged ions.
• An ionic compound’s lattice energy is determined by the following factors:

(i) Ion size: The smaller the ions, the higher the lattice energy.

(ii) Ion charge: The larger the ionic charge, the stronger the interionic attraction and 	the higher the lattice energy.

Characteristics of Ionic Compounds

(i) Physical’State: They are crystalline solids with a crystal lattice structure. Ionic compounds, unlike other gaseous molecules such as H2, N2, 02, and Cl2, do not exist as single molecules.

(ii) Melting and boiling points: Ionic compounds have high melting and boiling points due to the high interionic force between them.

(iii) Solubility: They dissolve in polar solvents like water but not in organic solvents like benzene, CCl4, and so forth.

Lewis-Langmuir Concept of Covalent Bond

• The term “covalent bond” refers to a relationship established between two or more atoms through the mutual contribution and sharing of electrons.
• A homoatomic covalent molecule is one in which the joining atoms are the same. Heteroatomic molecules are those that are distinct from one another.

Simple Molecules in Lewis Representation (the Lewis Structures)

The following steps can be used to write the Lewis dot Structure:

(i) Determine the combined atoms’ total number of valence electrons.

(ii) Each anion adds one electron to the equation, whereas each cation subtracts one electron. This provides you the total quantity of electrons you’ll need to distribute.

(iii) By understanding the merging atoms’ chemical symbols.

(iv) After inserting shared pairs of electrons for a single bond, the leftover electrons can be used to account for multiple bonds or lone pairs. It should be emphasised that each atom’s octet should be finished.

Dipole Movement

• Polar molecules are also known as dipole molecules because of their polarity, and they have a dipole moment.
• The product of the magnitude of the positive or negative charge and the distance between the charges is the dipole moment.

Dipole Moment’s Applications

(i) To figure out how polar the molecules are.

(ii) Identifying the molecular shapes.

Molecules having zero dipole moment, for example, will be linear or symmetrical. Molecules with asymmetrical shapes will be bent or angular in form.

(For example, NH3 has a D of 1.47).

(iii) When computing the fraction of polar bonds that are ionic.

Sidgwick and Powell introduced the Valence Shell Electron Pair Repulsion (VSEPR) Theory in 1940, based on the repulsive behaviour of electron pairs in the valence shell of atoms. Nyholm and Gillespie worked on it further (1957).

The following are the main postulates:

I The amount of electron pairs (bound or non-bonded) surrounding the core atoms determines the precise structure of the molecule.

(ii) Because the electron pairs exist surrounding the centre atom and the electron clouds are negatively charged, they have a propensity to reject one other.

(iii) Electron couples strive to align themselves in such a way that the rupulsion between them is minimised.

(iv) The valence shell is modelled as a sphere with electron pairs arranged at their greatest separation.

(v) A multiple bond is considered as a single electron pair, and the electron pairs that make up the bond are treated as single pairs as well.

Bond Theory of Valence

• Heitler and London (1927) established the valence bond hypothesis, which was further expanded by Pauling and others. It is based on the idea of atomic orbitals and atoms’ electrical arrangement.
• Let’s look at how a hydrogen molecule is formed using the valence-bond theory.
• Let’s say two hydrogen atoms A and B have nuclei NA and NB, respectively, and electrons eA and eB.
• New attracting and repulsive forces begin to operate as these two atoms get closer.

(i) Each atom’s nucleus is attracted to its own electron as well as the electron of the other atom, and vice versa.

(ii) Between the electrons of two atoms and the nuclei of two atoms, repulsive forces emerge. Attractive forces pull the two atoms closer together, while repulsive forces pull them apart.

The Concept of Orbital Overlap

• The orbital overlap notion states that when atoms create a covalent connection, orbitals belonging to atoms with opposing spins of electrons overlap.

Orbital Overlap Types

The covalent bonds are classified as sigma () or pi () bonds, depending on the type of overlapping.

(i) Bond (Sigma): The end-to-end (head-on) overlap of bonding orbitals along the internuclear axis forms the sigma bond.

There are three forms of axial overlap between these orbitals:

• s-s overlapping: Two half-filled s-orbitals overlap along the internuclear axis in this example, as seen below:

• s-p overlapping: This sort of overlapping happens when one atom’s half-filled s-orbitals overlap with the half-filled p-orbitals of another atom.

• p-p overlapping: This sort of overlapping occurs between half-filled p-orbitals of two approaching atoms.

(ii) bond (pi): When atomic orbitals overlap in such a way that their axes stay parallel to each other and perpendicular to the internuclear axis, a bond is established. The orbital is generated as a result of lateral or sidewise overlapping.

Sigma and pf Bond Strength

• The axial overlapping of the atomic orbitals forms the sigma bond (bond), whereas the sidewise overlapping of the atomic orbitals forms the -bond. Because axial overlapping is bigger than sidewise overlapping. As a result, the sigma bond is considered to be stronger than the -bond.
• There’s a difference between sigma and n bonds.

Hybridisation is the process of combining orbitals with somewhat varying energies in order to redistribute their energies and generate a new set of orbitals with comparable energy and shapes.

Hybridization’s Key Characteristics:

The hybridisation involves orbitals of almost equal energy.

The number of hybrid orbitals created is equal to the number of atomic orbitals combined; (iii) The kind of hybridisation can reveal the geometry of a covalent molecule.

Hybrid orbitals are more effective than pure atomic orbitals in creating stable bonds. Conditions required for hybridisation: I Valence shell orbitals participate in hybridisation.

The energy of the orbitals engaged in hybridisation should be almost equal.

Prior to hybridisation, electron promotion is not a need.

In some circumstances, filled valence shell orbitals are also involved in hybridization.

Hybridization Types:

• Hybridization of sp: The orbital is known as a sp hybrid orbital, and the kind of hybridisation is known as sp hybridisation, where one s and one p-orbital hybridise to generate two equivalent orbitals.
• Each of the hybrid orbitals generated has 50% s-characers and 50% p-characers. The term “diagonal hybridization” refers to this sort of hybridization.
• sp2 hybridisation: One s and two p-orbitals hybridise to generate three equivalent sp2 hybridised orbitals in this kind of hybridisation.
• All three hybrid orbitals are in the same plane, forming a 120° angle. Example. Sp2 hybridisation occurs in a variety of compounds, including BF3, BH3, BCl3 carbon compounds with a double bond, and others.
• sp3 hybridisation: One s and three p-orbitals in an atom’s valence shell are hybridised to generate four equivalent hybrid orbitals in this type of hybridisation. Each sp3 hybrid orbital has a 25 percent s-character and a 75 percent p-character. The four sp3 orbitals are aimed at the tetrahedron’s four corners.
• The sp3 hybrid orbitals form a 109.5° angle.
• A chemical that undergoes sp3 hybridisation is (CH4). sp3 hybridisation may also be used to explain the structures of NH2 and H20 molecules.

Molecular Orbital Formation: Atomic Orbitals in a Linear Combination (LCAO)

• The linear combination of atomic orbitals helps explain the development of molecular orbitals. Bonding molecular orbitals are generated by adding atomic orbitals, whereas antibonding molecular orbitals are formed by subtracting atomic orbitals.
• The following are the requirements for combining atomic orbitals:

(1) The combining atomic orbitals must have about identical energy.

(2) Around the molecule axis, the combining atomic orbitals must have the same symmetry.

(3) As much as possible, the combining atomic orbitals must overlap.

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