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Chemical Bonding JEE Notes

Chemical Bonding JEE Notes provided here will serve as a great study tool for students who are gearing up to appear for competitive exams like JEE Main and JEE Advanced. When students read through the chemical bonding JEE notes, they will gain a better understanding of each topic covered in this unit. Moreover, they will get a detailed overview of all the concepts and important points to remember.

Chemical bonding notes for JEE are based on the latest syllabus, and our subject-matter experts have provided clear and crisp explanations of the concepts. Besides, students can have a quick revision of all the concepts with these notes and ultimately prepare effectively for the exam.

What Is Chemical Bonding?

A chemical bond is defined as an attractive force that holds various constituents, such as atoms and ions, together in different chemical species. Chemical bonding is basically the process of forming bonds between atoms and molecules. 

During the formation of a chemical bond, usually, the electrons that are found in the outermost shell of an atom are active, or they take part in this. One of the important factors in chemical bonding is that an atom tends to attain a stable outer electronic configuration of inert gases. There are three different ways in which atoms can attain stable electronic configuration:

  • By losing electrons
  • By gaining electrons
  • By sharing electrons

Important Chemical Bonding Terms

Octet Rule 

We know that every atom tries to attain the octet configuration (presence of eight electrons) in its valence shell. This is done either by losing, gaining or sharing electrons. The octet rule basically states this fact. A stable arrangement is created when the atom is surrounded by eight electrons.

Students should remember some important categories of exceptions:

  • The odd number of electrons cannot be distributed into pairs. Odd-electron species like O2, NO2, NO, etc., violate the octet rule.
  • Electron-deficient compounds such as AlCl3, BCl3, B2H3, BeCl3, etc., have fewer than eight electrons in their central atom.
  • Some compounds have a central atom which `expands’ its octet. For example, PCl5, SF5, and many transition metal compounds also have 9-18 electrons in their outermost shell.

Lewis Symbols 

The Lewis symbols help us to clearly understand the concept of valence electrons. The valence electrons are usually represented with dots, and they are called Lewis symbols.

Ionic Bond

An ionic bond is the electrostatic force of attraction that holds two oppositely charged ions together. The bond is usually formed as a result of the electrostatic attraction between positive and negative ions. It is also known as an electrovalent bond. The electrovalence is thus equal to the number of unit charge(s) on the ion.

Important conditions for the formation of ionic compounds 

(i) There should be a large electronegativity difference between two combining elements. 

(ii) Electropositive elements must have a low ionization enthalpy. 

(iii) The electronegative element must have a high negative value of electron gain enthalpy.

(iv) The lattice enthalpy of an ionic solid should be high.

Read More: Ionic Bond 

Lattice Energy/Enthalpy

The energy that is required to completely separate one mole of a solid ionic compound into gaseous constituent ions is known as lattice energy. Lattice energy can also be defined as the energy released when gaseous ions form one mole of a solid ionic compound. Lattice energy keeps the cations and anions of the compound in fixed positions in a crystalline solid state. Cations and anions form crystal lattices of ionic crystals in space by the electrostatic force of attraction. However, we cannot experimentally measure lattice energy. Hess’s law of heat summation can be used in estimating the lattice energy.

Sublimation energy is the energy that is used in converting solid metal to a gaseous state.

Ionization energy is used to remove an electron from the outermost shell of metal in a gaseous state. 

Dissociation energy is used in converting the nonmetal molecule to atoms.

Electron affinity is a type of energy that is required during the formation of an anion when an electron is added to a nonmetal in a gaseous state. 

Bond Angle

The angle that exists between the orbitals containing bonding electron pairs around the central atom in a molecule or a complexion is known as a bond angle.

Bond Enthalpy

The amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state is known as bond enthalpy.

Bond Order

Bond order is generally allocated by the number of bonds between two atoms in a molecule. Bond order (BO) is defined as one-half the difference between the number of electrons present in the bonding and the antibonding orbitals.

Born-Haber Cycle

The Born-Haber cycle or process helps us define the different energy terms that are involved during the formation of an ionic compound. It is a thermochemical cycle based on Hess’s law of constant heat summation and largely helps us to describe the formation of ionic compounds from different elements in a more subtle manner. 

Covalent Bond

A covalent bond is usually formed between two electronegative non-metals that are similar in nature. There is sharing of valence electrons to attain the noble gas configuration.

Conditions for the formation of covalent bond 

  • The electronegativity difference between two atoms should be less.
  • Occurs between two non-metals.
  • The combining atom must share at least one electron.
  • Atoms attain the noble gas configuration after bond formation.

Depending upon the number of shared electron pairs, a covalent bond is classified into three types:

  •  Single Covalent Bond
  •  Double Covalent Bond
  •  Triple Covalent Bond

Read More: Covalent Bond

Polar Covalent Bond

When there is unequal sharing of electrons due to the difference in the electronegativity of combining atoms, a polar covalent bond is formed.

Nonpolar Covalent Bond

A nonpolar covalent bond is formed by the equal sharing of electrons between atoms. This bond is formed when the combining atoms have similar electron affinity (diatomic elements). The electronegativity difference between two atoms is usually zero.

Sigma and Pi Bonds 

Sigma (σ) bond is a type of covalent bond that is formed due to the overlapping of atomic orbitals along the internuclear axis.

Pi (π) bond is a type of covalent bond that is formed by sidewise overlapping of atomic orbitals.

Strength of Sigma and Pi Bonds 

The strength of a bond depends upon the extent of overlapping. In the sigma bond, the overlapping of orbitals occurs to a larger extent. In the pi bond, the extent of overlapping occurs to a smaller extent. Therefore, the sigma bond is stronger in comparison to the pi bond.

Hydrogen Bonding

A hydrogen bond is a type of polar covalent bond that exists between oxygen and hydrogen atoms. Here, the hydrogen develops a partial positive charge, and the electrons are pulled towards the more electronegative oxygen atom. This is a weaker form of chemical bonding.

Conditions required for H-bond

(i) An atom with higher electronegativity should be linked to H-atom. 

(ii) The electronegative atom should be smaller in size. 

(iii) The electronegative atom should contain a lone pair.

Lewis Dot Structures

In the simplest manner, we can define Lewis dot structures as diagrams that are used to describe the chemical bonding between atoms in a molecule. It is also called the electron dot structure. Lewis dot structures can be drawn for molecules in which the combining atoms may be similar or different.


Resonance in Chemistry can be generally defined as the phenomenon wherein two or more structures involving an identical position of an atom are written for a particular compound. To put it in simple words, it is a situation in which more than one canonical structure can be written for a species. 

Resonance Hybrid: It is the actual structure of all different possible structures that can be written for the molecule without violating the rules of covalence maxima for the atoms. 


It means the mixing of two atomic orbitals with the same energy level to give a degenerated new type of orbitals bond length. Bond length is the equilibrium distance between the nuclei of two bonded atoms in a molecule.

Salient features of hybridization 

The main features of hybridization are as follows:

  1. The number of hybrid orbitals is equal to the number of atomic orbitals that get hybridized. 
  2. The hybridized orbitals are always equivalent in energy and shape. 
  3. Hybrid orbitals are more effective in forming stable bonds than pure atomic orbitals. 
  4. These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs, and thus a stable arrangement is obtained. Therefore, the type of hybridization indicates the geometry of the molecules. 

Important Conditions for Hybridization

(i) The orbitals present in the valence shell of the atom are hybridized. 

(ii) The orbitals undergoing hybridization should have almost equal energy 

(iii) Promotion of electrons is not an essential condition prior to hybridization. 

(iv) It is the orbital that undergoes hybridization and not the electrons. For example, for orbitals of a nitrogen atom (2.522p„1 2pyl 2 psi) belonging to a valency shell, when hybridized to form four hybrid orbitals, one of which has two electrons (as before) and the other three have one electron each. It is not necessary that only half-filled orbitals participate in hybridization. In some cases, even filled orbitals of valence shells take part in hybridization. 

Determination of hybridization of an atom in a molecule or ion: 

Steric number rule (given by Gillespie): Steric number of an atom = number of atoms bonded with that atom + number of lone pair(s) left on that atom. 

Note: This rule is not applicable to molecules/ions which have odd e-(C102, NO, NO2), free radicals and compounds like B2H6, which involve 3 centres 2e bond (banana bond). 

Valence Bond Theory

Valence bond theory says that electrons in a covalent bond reside in a region that is at the overlap of individual atomic orbitals.

VSEPR Theory

The Lewis concept is unable to explain the shapes of molecules. Sidgwick and Powell provided a useful idea for predicting the shapes and geometries of molecules. The theory was based on the repulsions between electron pairs, known as the valence shell electron pair repulsion (VSEPR) theory.

The main postulates of VSEPR theory are as follows: 

(i) The shape of a molecule depends upon the number of valence shell electron pairs [bonded or nonbonded) around the central atom. 

(ii) Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. 

(iii) These pairs of electrons tend to occupy such positions in space that minimise repulsion and thus maximise the distance between them. 

(iv) The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at a maximum distance from one another. 

(v) A multiple bond is treated as if it is a single electron pair, and the two or three electron pairs of a multiple bond are treated as a single super pair. 

(vi) Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure. 

Molecular Orbital Theory 

The molecular orbital theory was developed by F. Hund and R.S. Mulliken in 1932. According to Molecular Orbital Theory, individual atoms combine to form a molecule. Thus, the electrons of an atom are present in various atomic orbitals and are associated with several nuclei. The salient features are as follows: 

(i) Just like electrons of an atom are present in various atomic orbitals, electrons of the molecule are present in various molecular orbitals. 

(ii) Molecular orbitals are formed by the combination of atomic orbitals of comparable energies and proper symmetry. 

(iii) An electron in an atomic orbital is influenced by one nucleus, while in a molecular orbital, it is influenced by two or more nuclei depending upon the number of atoms in the molecule. Thus, an atomic orbital is monocentric, while a molecular orbital is polycentric. 

(iv) The number of molecular orbitals formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals called bonding molecular orbital and anti-bonding molecular orbital are formed.

(v) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital. 

(vi) Just like the electron probability distribution around a nucleus in an atom is given by an atomic orbital, the electron probability distribution around a group of nuclei in a molecule is given by molecular orbital. 

(vii) The molecular orbitals, like the atomic orbitals, are filled in accordance with the Aufbau principle obeying the Pauli exclusion principle and Hund’s rule of maximum multiplicity. But the filling order of these molecular orbitals is always experimentally decided, there is no rule like the (n 1) rule in the case of atomic orbitals. 

Conditions for the Combination of Atomic Orbitals

  1. The combining atomic orbitals must have the same or nearly the same energy. 
  2. The combining atomic orbitals must have the same symmetry about the molecular axis. 
  3. The combining atomic orbitals must overlap to the maximum extent. 

Contours and energies of bonding and antibonding molecular orbitals formed through combinations of 1s atomic orbitals, 2pz atomic orbitals and 2px atomic orbitals are shown below.

Van der Waals Interactions

Van der Waals interactions are the weakest of all bondings. They are driven by induced electrical interaction between two or more atoms or molecules.

(a) Ionization energy: In the formation of an ionic bond, a metal atom loses electrons to form a cation. This process requires energy equal to the ionization energy. The lesser the value of ionization energy, the greater the tendency of the atom to form a cation. For example, alkali metals form cations quite easily because of the low values of ionization energy.  

(b) Electron affinity: Electron affinity is the energy released when a gaseous atom accepts an electron to form a negative ion. Thus, the value of electron affinity gives the tendency of an atom to form an anion. The greater the value of electron affinity, the more the tendency of an atom to form an anion. For example, halogens having the highest electron affinities within their respective periods form ionic compounds with metals very easily.

(c) Lattice energy: Once the gaseous ions are formed, the ions of opposite charges come close together and pack up three-dimensionally in a definite geometric pattern to form an ionic crystal.

Since the packing of ions of opposite charges takes place as a result of attractive force between them, the process is accompanied by the release of energy, referred to as lattice energy. Lattice energy may be defined as the amount of energy released when one mole of an ionic solid is formed by the close packing of a gaseous ion. 


Hydrolysis means a reaction with water molecules, ultimately leading to the breaking of the O-H bond into H+ and OW ions, while the term hydration means the surrounding of polar molecules or ions by polar molecules of water. In hydrolysis, there is a complex formation with a water molecule or a reaction with the water molecule. Hydrolysis in covalent compounds generally takes place by two mechanisms. 

(a) By coordinate bond formation: Generally in halides of atoms having vacant d-orbitals. 

(b) By H-bond formation: For example, in nitrogen trihalides. 

Coordinate bond (Dative bond):

The bond formed between two atoms in which the contribution of an electron pair is made by one of them while the sharing is done by both.

Formal charge:

The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.

General properties of ionic compounds: 

(a) Physical state: At room temperature, ionic compounds exist either in a solid state or in a solution phase but not in a gaseous state. 

(b) Isomorphism: Simple ionic compounds do not show isomerism, but isomorphism is their important characteristic. Crystals of different ionic compounds having similar crystal structures are known to be isomorphs to each other, and the phenomenon is known as isomorphism. 

(c) Electrical conductivity: 

Ionic solids are almost non-conductors. However, they conduct a very small amount of current due to a crystal defect. All ionic solids are good conductors in the molten state as well as in their aqueous solutions because their ions are free to move. 

(d) The solubility of ionic compounds: Soluble in polar solvents like water which has a high dielectric constant. 

Fajan’s Rule

When an anion and cation approach each other, the valence shell of an anion is pulled towards the cation nucleus, and thus the shape of the anion is deformed. This phenomenon of deformation of an anion by a cation is known as polarisation, and the ability of a cation to polarize a nearby anion is called the polarizing power of a cation. 

Fajan’s pointed out that the greater is the polarization of an anion in a molecule, the more is a covalent character in it. 

The more distortion of the anion, the more polarization, so the covalent character increases. 

Fajan’s rules govern the covalent character in the ionic compounds, and they are as follows:

(i) Size of cation: Size of cation polarisation. 

(ii) Size of anion: Size of anion polarisation.

(iii) Charge on cation: Charge on cation a polarisation. 

(iv) Charge on anion: Charge on anion a polarisation. 

(v) Pseudo inert gas configuration of cation: The cation having a pseudo inert gas configuration has more polarizing power than the cation that has an inert gas configuration. Thus, NaC1, having an inert gas configuration, will be more ionic, whereas CuCl, having a pseudo inert gas configuration, will be more covalent in nature.

Metallic Bond

Most metals crystallize in close-packed structures. The ability of metals to conduct electricity and heat must result from strong electron interactions among 8 to 12 nearest neighbours (which is also called coordination number). Bonding in metals is called metallic bonding. It results from the electrical attractions among positively charged metal ions and mobile, delocalized electrons belonging to the crystal as a whole. 

Two models are considered to explain metallic bonding: 

(A) Electron-sea model 

(B) Band model 

Some special bonding situations: 

(a) Electron-deficient bonding 

There are many compounds in which some electron-deficient bonds are present apart from normal covalent bonds or coordinate bonds. These electron-deficient bonds have fewer electrons than expected, such as three centre-two electron bonds (3c-2e) present in diborane B2H6, Al2(CH3)6, BeH2(s) and bridging metal carbonyls. 

(b) Back Bonding 

Generally, back bonding takes place when out of two bonded atoms, one of the atoms has vacant orbitals (generally, this atom is from the second or third period), and the other bonded atom has some non-bonded electron pair (generally, this atom is from the second period). Back bonding increases the bond strength and decreases the bond length. For example, in BF3, the boron atom completes its octet by accepting two 2p-electrons of fluorine into a 2p empty orbital.


Chemical Bonding


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