A chemical bond is a lasting attraction between atoms, ions, or molecules that enables the formation of chemical compounds. Bonds can result from the electrostatic force of attraction between opposite charges or through the sharing of electrons.
Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals that can form covalent bonds. The type of hybridization (sp, sp2, sp3, etc.) depends on the number of electron pairs around the central atom.
A polar covalent bond is a type of covalent bond where the electrons are unequally shared between two atoms due to a difference in electronegativity. This results in a partial positive charge on one atom and a partial negative charge on the other.
Matter is composed of elements, with atoms typically forming molecules held together by chemical bonds. These bonds arise from attractive forces between atoms. Various theories, including Kössel-Lewis, VSEPR, VBT, and MOT, explain why atoms combine and how molecules form. Ultimately, chemical bonding is nature's way of achieving stability by lowering the energy of systems.
1.0Chemical Bonding
Chemical bonding is the process by which atoms combine to form molecules and compounds. It involves the attraction between atoms, which can be due to the sharing or transfer of electrons. There are two main types of chemical bonds:
Strong bonds- Strong bonds, refer to those with high bond energies, such as covalent bonds and metallic bonds, where electrons are shared or delocalized extensively between atoms. These bonds require significant energy input to break.
Weak bonds- Weak bonds, on the other hand, have lower bond energies and are more easily broken. Examples include hydrogen bonds, van der Waals forces, and dipole-dipole interactions. While weak bonds are important in molecular interactions and structure, they are relatively weaker compared to strong bonds and require less energy to break.
Major types of chemical bonding are here:
Kössel-Lewis Approach:
This approach, also known as the Lewis electron dot structure or simply Lewis structure, is a method used to represent the arrangement of valence electrons within atoms and molecules. In the Kössel-Lewis approach:
Each atom is represented by its chemical symbol.
Valence electrons are depicted as dots surrounding the atom's symbol.
Shared pairs of electrons in covalent bonds are represented by a pair of dots or a line between the symbols of the bonded atoms.
Ionic bonds are illustrated by the transfer of electrons between atoms, resulting in the formation of cations and anions.
Octet Rule:
The octet rule is a fundamental principle in chemistry stating that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons, resembling the electron configuration of noble gases. The octet rule is closely related to the Kössel-Lewis approach because Lewis structures are often used to demonstrate how atoms achieve an octet.
Together, the Kössel-Lewis approach provides a visual representation of electron distribution in molecules and ions, while the octet rule serves as a guiding principle for understanding the stability of these arrangements. The Kössel-Lewis approach helps illustrate how atoms achieve an octet through various bonding patterns, such as Covalent, Ionic bonds and Coordinate bond.
Related Video:
2.0Ionic Bond
An ionic bond is a type of chemical bond that forms between positively and negatively charged ions. It typically occurs between a metal atom (which loses electrons to become positively charged) and a non-metal atom (which gains electrons to become negatively charged). The attraction between these oppositely charged ions results in the formation of an ionic compound.
3.0Covalent Bond
A covalent bond is formed when atoms share electrons to achieve a stable electron configuration, usually that of a noble gas. It commonly occurs between non-metal atoms, resulting in the formation of molecules. Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared. They are directional and can vary in strength based on factors like atom electronegativity and molecular geometry.
Sharing of electrons may occur in three ways:
No. of electrons shared
Electron pair
Bond. between two atoms
Examples
2
1
Single bond(—)
H — H
4
2
Double bond (＝)
O = O
6
3
Triple bond (☰)
N ☰ N
Note- Strengths of Ionic and Covalent Bonds
Chemical bonds are formed either by electron sharing (covalent bonds) or electron transfer (ionic bonds), allowing atoms to achieve a stable electron configuration. Ionic bonds involve the transfer of electrons, while covalent bonds involve electron sharing. Generally, ionic bonds are stronger due to electrostatic forces, while covalent bonds are weaker and held by Van der Waals forces. However, bond strength can vary based on conditions. For instance, diamond, with covalent bonds, is hard, while in biochemical reactions, ionic bonds in water can break easily.
4.0Co-Ordinate Bond (Dative bond)
A coordinate bond, also known as a dative bond, is a type of covalent bond where the shared electron pair comes from one atom. The atom providing the electron pair is called the donor (Lewis base), while the atom accepting the electron pair is called the acceptor (Lewis acid), hence it's often referred to as a donor-acceptor bond.
For example, in BF3, which is an electron-deficient compound, the boron atom acts as the acceptor of the electron pair.
5.0Lewis Symbols and Structures
Lewis structures are diagrams that represent the bonding and electron distribution in molecules and ions. They are constructed using the Kössel-Lewis approach, where each atom is represented by its chemical symbol, and valence electrons are depicted as dots around the atom's symbol.
According to Lewis, only outermost shell electrons partake in chemical bonding, termed valence electrons. Inner shell electrons are shielded and typically do not engage in bonding. Lewis introduced Lewis symbols and structures to represent valence electrons, simplifying the understanding of chemical bonding.
6.0Formal Charges and Resonance
Formal charge is a concept used to determine the distribution of electrons in a molecule or ion within a Lewis structure. It helps to assess the stability and distribution of charge within a molecule.
To calculate the formal charge of an atom in a Lewis structure:
Assign one electron to each bonded atom and one-half of the electrons in each shared pair to the atom.
Subtract the assigned electrons from the total valence electrons of the atom.
The result is the formal charge of the atom.
The formal charge of an atom can be calculated using the formula:
QF=NA−NL.P. −1/2NB.P
Where :
NA= Total number of valence electron in the free atom
NL.P. = Total number of non bonding (lone pair) electrons
NB.P. = Total number of bonding(shared) electrons
Molecule
Structure
Formal Charge
O3
O(1) = 6 –2 – 1/2 (6) = + 1
O(2) = 6 – 4 – 1/2 (4) = 0
O(3) = 6 – 6 – 1/2 (2)
= – 1
CO
O(1) = 6 –2 – 1/2 (6) = + 1
O(2) = 6 – 4 – 1/2 (4) = 0
O(3) = 6 – 6 – 1/2 (2)
= – 1
C = 4 – 2 – 1/2 (6) = –1
O = 6 – 2 – 1/2 (6) = +1
Resonanceeffect
Resonanceis a concept used in chemistry to describe the delocalization of electrons within molecules or ions. It occurs when a Lewis structure can be drawn in multiple ways, called resonance structures, without changing the overall arrangement of atoms. Resonance structures are necessary to accurately describe the electronic structure of certain molecules or ions, especially those containing multiple bonds or charged species.
Here are some examples of molecules and ions with their resonance structures:
the
Benzene is a cyclic aromatic hydrocarbon with six carbon atoms arranged in a ring, each bonded to one hydrogen atom. Its resonance structures involve the delocalization of π (pi) electrons throughout the ring.
Carbonate ion (CO3⁻2)- Carbonate ion consists of a central carbon atom bonded to three oxygen atoms, two of which are single-bonded and one is double-bonded. The resonance structures involve the movement of the double bond between the carbon and oxygen atoms.
7.0Bond Parameters
Bond parameters are key descriptors of chemical bonds in molecules. They include bond length, angle, strength, energy, order, dipole moment, and hybridization. These parameters provide insights into bond characteristics like distance, geometry, stability, polarity, and orbital mixing.
Valence Bond Theory- Valence Bond Theory (VB Theory) provides a deeper understanding of chemical bonding, addressing the limitations of Lewis structures. It was developed by Heitler, London, and later refined by Linus Pauling.
In VB Theory, the formation of chemical bonds is explained by the overlap of atomic orbitals between bonding atoms. This theory relies on concepts such as atomic orbitals, electronic configurations, orbital overlap, hybridization, and superposition principles.
Covalent bonding involves the overlapping of valence shell orbitals, with directional characteristics. Bond strength is proportional to overlap extent, influenced by orbital type and overlapping nature.
Directional orbitals (p, d, f) exhibit stronger bonds, and co-axial overlap is stronger than collateral. Overall, bond strength decreases with increasing principal quantum number (n), and greater overlap occurs with co-axial overlapping.
Type of Overlapping- The covalent bond can be classified into the following types based on the types of overlapping:
(i) Sigma (σ) Bond: Sigma bonds result from the overlap of atomic orbitals along the axis connecting the nuclei of the bonded atoms. This type of bonding allows for free rotation about the bond axis. Sigma bonds are typically formed by the overlap of s-orbitals, as well as by the head-on overlap of p-orbitals and hybrid orbitals.
(ii) Pi (π) Bond: Pi bonds arise from the sideways overlap of parallel p-orbitals that are adjacent to the sigma-bonded atoms. Pi bonds are characterized by the presence of electron density above and below the plane of the nuclei of the bonded atoms. Unlike sigma bonds, pi bonds do not allow for free rotation about the bond axis. Pi bonds are often observed in double and triple bonds alongside sigma bonds.
8.0Hybridization
Hybridization arises when atomic orbitals mix to form new hybrid orbitals that have different shapes and properties from the original atomic orbitals. This phenomenon occurs primarily in covalent bonding situations, where atoms share electrons to form bonds. Hybridization is crucial in explaining the shapes and geometries of molecules, as well as their bond angles.
9.0The Valence Shell Electron Pair Repulsion (VSEPR) Theory
The Lewis concept falls short in explaining molecular shapes, prompting the need for a more comprehensive theory. In 1940, Sidgwick and Powell introduced a theory based on repulsive interactions among electron pairs in the valence shell. This theory was later refined by Nyholm and Gillespie in 1957, providing a simple yet effective method to predict the shapes of covalent molecules.
The VSEPR (Valence Shell Electron Pair Repulsion) theory is a fundamental concept in chemistry that plays a crucial role in predicting the three-dimensional shapes of molecules. This theory is essential for understanding molecular geometry, bond angles, and overall molecular shape, providing insights into chemical properties and reactivity.
Related Video:
10.0Molecular Orbital Theory
The Molecular Orbital (MO) Theory is a fundamental concept in chemistry that describes the behaviour of electrons in molecules using quantum mechanics. In essence, it provides a framework for understanding how electrons are distributed within molecules and how this distribution influences their properties and behaviour.
Molecular Orbitals (MOs) are regions of space within a molecule where electrons are most likely to be found. They are formed by the combination of atomic orbitals (AOs) from the constituent atoms in a molecule.
Formation of Molecular Orbitals:
Molecular orbitals are formed by the combination of atomic orbitals from the constituent atoms in a molecule. This combination can lead to the formation of bonding molecular orbitals (lower energy) and antibonding molecular orbitals (higher energy).
Molecular orbitals (MOs) can be categorized into several types based on their symmetry, orientation, and energy levels. Here are the main types of molecular orbitals:
Molecular Orbitals: Formed by combining atomic orbitals from constituent atoms, resulting in bonding (σ, π) and antibonding (σ*, π*) orbitals. Bonding orbitals have lower energy and contribute to stability, while antibonding orbitals are higher in energy and weaken bonds. Non-bonding orbitals are localized on individual atoms and do not significantly contribute to bonding. Higher-energy orbitals are involved in electronic transitions and spectroscopic phenomena.
For example s-s combination of orbital diagram is given below
11.0Hydrogen Bonding
Hydrogen bonding is a special type of intermolecular force that occurs between a hydrogen atom bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) and another electronegative atom.
