Notes on Chemical Bonding for JEE Main

By Prashant Kumar|Updated : September 14th, 2018

The attractive force which holds various constituents such as atoms, ions etc., together in different chemical species is called a chemical bond. This is an important topic in order to score more in the chemistry section of the JEE Main exam.

In the formation of a chemical bond, only the electrons of the outermost shell of an atom are involved. In the process, each atom attains a stable outer electronic configuration of inert gases.

Ionic or Electrovalent Bond :

The formation of an ionic compound would primarily depend upon :

* The ease of formation of the positive and negative ions from the respective neutral atoms.

* The arrangement of the positive and negative ions in the solid, that is the lattice of the crystalline compound.

Conditions for the Formation of Ionic Compounds :

(i) Electronegativity difference between two combining elements must be larger.

(ii) Ionization enthalpy (M(g) → M+(g) + e) of electropositive element must be low.

(iii) The negative value of electron gain enthalpy (X(g) + e → X (g)) of the electronegative element should be high.

(iv) Lattice enthalpy (M+(g) + X (g) → MX (s)) of an ionic solid must be high.

Lattice Enthalpy :

The lattice enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

The process involves both the attractive forces between ions of opposite charges and the repulsive forces between ions of like charge. The solid crystal being three dimensional; it is not possible to calculate lattice enthalpy directly from the interaction of forces of attraction and repulsion only. Factors associated with the crystal geometry have to be included.

Factors affecting lattice energy of an ionic compound :

(i) Lattice energy byjusexamprep where (r+ + r) = Inter-ionic Distance.


(ii) Lattice energy byjusexamprep

byjusexamprep charge on cation in terms electronic charge.

byjusexamprep charge on anion in terms electronic charge.  


Determination of lattice energy :

Born-Haber Cycle :

It inter relate’s the various energy terms involved during the formation of an ionic compound.

It is a thermochemical cycle based on the Hess’s law of constant heat summation.

Hess’s Law is the net enthalpy change of a chemical reaction or of any process always remain the same whether the reaction takes place in on step or many steps.

Hydration :

All the simple salts dissolve in water, producing ions, and consequently the solution conduct electricity. Since Li+ is very small, it is heavily hydrated. This makes the radius of hydrated Li+ ion large and hence it moves only slowly. In contrast, Cs+ is the least hydrated because of its bigger size and thus the radius of the Cs+ ion is smaller than the radius of hydrated Li+, and hence hydrated Cs+ moves faster, and conducts electricity more readily.

Hydrolysis :

Hydrolysis means reaction with water molecules ultimately leading to breaking of O-H bond into H+ and OH- ions. While the terms 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 reaction with the water molecule.

Hydrolysis in covalent compounds takes place generally 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.

General properties of ionic compounds :

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

(b) Isomorphism: Simple ionic compound does 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 little 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 high dielectric constant

The covalent character in ionic compounds (Fajan’s rule) :

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

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

More distortion of anion, more will be polarisation then covalent character increases.

Fajan’s gives some rules which govern the covalent character in the ionic compounds, which are as follows :

(i) Size of cation: Size of cation a1 / polarisation.

(ii) Size of anion: Size of anion a 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: Cation having pseudo inert gas configuration has more polarizing power than the cation that has inert gas configuration. Thus NaCl having inert gas configuration will be more ionic whereas CuCl having pseudo inert gas configuration will be more covalent in nature.

Covalent Bond :

It forms by sharing of valence electrons between atoms to form molecules e.g., formation of Cl2 molecule :


Chlorine atoms attain the outer shell octet of the nearest noble gas (i.e., argon). The dots represent electrons. Such structures are referred to as Lewis dot structures.


The Lewis dot structures can be written for other molecules also, in which the combining atoms may be identical or different. The important condition being that :

(i) Each bond is formed as a result of the sharing of an electron pair between the atoms.

(ii) Each combining atom contributes at least one electron to the shared pair.

(iii) The combining atoms attain the outer-shell noble gas configurations as a result of the sharing of electrons.

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


Other examples: H2SO4, HNO3, H3O+, N2O [Cu(NH3)4]2+

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. It is expressed as :


Let us consider the ozone molecule (O3). The Lewis structure of O3, may be drawn as :


The atoms have been numbered as 1.2 and 3. The formal charge on :

The central O atom marked byjusexamprep The terminal O atom marked byjusexamprep


The terminal O atom marked byjusexamprep

Hence, we represent O3 along with the formal charges as follows :


Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. Generally, the lowest energy structure is the one with the smallest formal charges on the atoms.


Limitations of the Octet Rule :

The octet rule, though useful, is not universal. It is quite useful for understanding. The structures of most of the organic compounds and it apply mainly to the second-period elements of the periodic table. There are three types of exceptions to the octet rule.


1. The incomplete octet of the central atom

In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH2 and BCl3.


Some other such compounds are AlCl3 and BF3.


2. Odd-electron molecules

In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide. No2, the octet rule is not satisfied for all the atoms.



3. The expanded octet

Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3rd orbitals also available for bonding. In a number of compounds of these elements, there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obviously, the octet rule does not apply in such cases.

Some of the examples of such compounds are: PF5, SF6, H2SO4 and a number of coordination compounds.

byjusexamprep byjusexamprep byjusexamprep


4. Other drawbacks of the octet theory

(i) It is clear that the octet rule is based upon the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF2, KrF2, XeOF2, etc.

(ii) This theory does not account for the shape of molecules.

(iii) It does not explain the relative stability of the molecules being totally silent about the energy of the molecules.


Valence bond theory (VBT) :

Since the energy gets released when the bond is formed between two hydrogen atoms, the hydrogen molecule is more stable than that of isolated hydrogen atoms. The energy so released is called as bond enthalpy, which is corresponding to a minimum in the curve depicted in the figure. Conversely. 435.8 kJ of energy is required to dissociate one mole of H2 molecule.



Figure: The potential energy curve for the formation of H2 molecule as a function of the internuclear distance of the H atoms. The minimum in the curve corresponds to the most stable state of H2.

Orbital Overlap Concept :

In the formation of the hydrogen molecule, there is a minimum energy state when two hydrogen atoms are so near that their atomic orbitals undergo partial interpenetration. This partial merging of atomic orbitals is called overlapping of atomic orbitals which results in the pairing of electrons. The extent of overlap decides the strength of a covalent bond.

In general, greater the overlap the stronger is the bond formed between two atoms. Therefore, according to orbital overlap concept, the formation of a covalent bond between two atoms results by the pairing of electrons present, in the valence shell having opposite spins.

Directional Properties of Bonds :

The valence bond theory explains the formation and directional properties of bonds in polyatomic molecules like CH4, NH3 and H2O, etc. in terms of overlap and hybridisation of atomic orbitals.

Types of Overlapping and Nature of Covalent Bonds :

The covalent bond may be classified into two types depending upon the types of overlapping :

(i) sigma (s) bond, and (ii) pi (p) bond

(i) sigma (s) bond : This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.

· s-s overlapping : In this case, there is overlap of two half filled s-orbitals along the internuclear axis as shown below :


· s-p overlapping : This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.


· p-p overlapping : This type of overlap takes place between half filled p-orbitals of the two approaching atoms.


(ii) pi (p) bond : In the formation of p bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to sidewise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.


Strength of Sigma and pi Bonds :

Basically the strength of a bond depends upon the extent of overlapping – In case of sigma bond, the overlapping of orbitals takes place to a larger extent. Hence, it is stronger as compared to the pi bond where the extent of overlapping occurs to a smaller extent. Further, it is important to note that pi bond between two atoms is formed in addition to a sigma bond. It is always present in the molecules containing multiple bonds (double or triple bonds)

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 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.

The repulsive interaction of electron pairs decreases in the order :

lone pair (lp) – lone pair (lp) > lone pair (lp) – bond pair (bp) > bond pair (bp) – bond pair (bp)


Hybridisation :

Salient features of hybridisation : The main features of hybridisation are as under :

1. The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.

2. The hybridised orbitals are always equivalent in energy and shape.

3. The hybrid orbitals are more effective in forming stable bonds than the 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 hybridisation indicates the geometry of the molecules.

Important conditions for hybridisation :

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

(ii) The orbitals undergoing hybridisation should have almost equal energy.

(iii) Promotion of electron is not an essential condition prior to hybridisation.

(iv) It is the orbital that undergo hybridization and not the electrons. For example, for orbitals of nitrogen atom byjusexamprep belonging to valency shell when hybridize to form four hybrid orbitals, one of which has two electrons (as before) and other three have one electron each. It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals of valence shell take part in hybridisation.

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

Steric number rule (given by Gillespie) :

Steric No. of an atom = number of atom 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(ClO2, NO, NO2), free radicals and compounds like B2H6 which involve 3 centre 2e bond (banana bond).

Table – 3

Steric Number

Types of Hybridisation







Trigonal planar






Trigonal bipyramidal






Pentagonal bipyramidal

Hybridization Involving d-orbital : 

Type of ‘d’ orbital involved

sp3 d


sp3 d2

dx2 – y2 & dz2

sp3 d3

dx2 – y2 , dz2 & dxy



Molecular Orbital Theory (MOT) :

The molecular orbital theory was developed by F. Hund and R.S. Mulliken in 1932. The salient features are :

(i) Just as electrons of any 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 the 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 as 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 the Hund’s Rule of Maximum Multiplicity. But the filling order of these molecular orbitals is always experimentally decided, there is no rule like (n + l) rule in case 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.


The energy levels of molecular orbitals have been determined experimentally from spectroscopic data for homonuclear diatomic molecules of second row elements of the periodic table. The increasing order of energies of various molecular orbitals for O2 and F2 is given below :


The increasing order of energies of various molecular orbitals for Be2, B2, C2, N2 etc., is :


The important characteristic feature of this order is that the energy of byjusexamprep molecular orbital is higher than that of byjusexamprep molecular orbitals.



Bond order (b.o.) is defined as one half the difference between the number of electrons present in the bonding and the antibonding orbitals i.e., Bond order (b.o.) = 1/2 (Nb – Na)

A positive bond order (i.e., Nb > Na) means a stable molecule while a negative (i.e., Nb < Na) or zero (i.e., Nb = Na) bond order means an unstable molecule.



Integral bond order values of 1,2 or 3 correspond to single, double or triple bonds respectively.



The bond order between two atoms in a molecule may be taken as an approximate measure of the bond length. The bond length decreases as bond order increases.



If all the molecular orbitals in a molecule are doubly occupied, the substance is diamagnetic (repelled by magnetic field) e.g., N2 molecule. However if one or more molecular orbitals are single occupied it is paramagnetic (attracted by magnetic field), e.g., O2 molecule.


Dipole moment :

It is defined as the product of magnitude of the partial charge byjusexamprep developed on any of the covalently bonded atoms atoms and the distance between two atoms.

Dipole moment (m) = Magnitude of charge (q) × distance of separation (d).

Dipole moment is usually expressed in Debye units (D). The conversion factors are

· 1 D = 3.33564 × 10–30 cm, where C is coulomb and m is meter.

· 1 Debye = 1 × 10–18 e.s.u. cm.

Further dipole moment is a vector quantity and is depicted by a small arrow with tail on the positive centre and head pointing towards the negative centre. For example the dipole moment of HF may be represented as


The shift in electron density is represented by crossed arrow byjusexamprep above the Lewis structure to indicate the direction of the shift.

In case of polyatomic molecules the dipole moment not only depend upon the individual dipole moments of bonds known as bond dipoles but also on the spatial arrangement of various bonds in the molecule. In such case, the dipole moment of a molecule is the vector sum of the dipole moment of various bonds. i.e. a molecule will have a dipole moment if the summation of all of the individual moment vector is non-zero.

byjusexamprep byjusexamprep where R is resultant dipole moment.


Resonance :

Definition : Resonance may be defined as the phenomenon in which two or more structures involving in identical position of atom, can be written for a particular compound.

For example, the ozone, O3 molecule can be equally represented by the structures I and II shown below :


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.



Hydrogen Bond :

Nitrogen, oxygen and fluorine are the highly electronegative elements. When they are tied to a hydrogen atom to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. This partially positively charged hydrogen atom forms a bond with the other more electronegative atom. This bond is called as hydrogen bond and is weaker than covalent bond. For example, in HF molecule, the hydrogen bond exists between hydrogen atom of one molecule and fluorine atom of another molecule as given below :



Conditions required for H-bond:

(i) Molecule should have more electronegative atom (F,O,N) linked to H-atom.

(ii) Size of electronegative atom should be smaller.

(iii) A lone pair should be present on electronegative atom.


Order of H-bond strength




(A) Intramolecular H-Bonding :

This type of H-bonding occurs when polar H and electronegative atom are present in the same molecule i.e., it is formed when hydrogen atom is present in between the two highly electronegative (F,O,N) atoms within the same molecule.


It has lower boiling point (i.e. more volatile) than its para-derivative (where association of molecules takes place using intermolecular H-bonding) because it exists as discrete molecules.


Necessary conditions for the formation of intramolecular hydrogen-bonding :

(a) the ring formed as a result of hydrogen bonding should be planar.

(b) a 5- or 6- membered ring should be formed.

(c) interacting atoms should be placed in such a way that there is minimum strain during the ring closure.


(B) Intermolecular H-Bonding :

Exists between the negative and positive ends of different molecules of the same or different substances i.e, it is formed between two different molecules of the same or different compounds.


(a) In water molecules

Due to polar nature of H2O there is association of water molecules giving a liquid state of abnormally high boiling point.


(b) The hydrogen bonds in HF link the F atom of one molecule with the H-atom of another molecule, thus forming a zig-zag chain (HF)n in both the solid and also in the liquid.


Some hydrogen bonding also occurs in the gas, which consists of a mixture of cyclic (HF)6 polymers, dimeric (HF)2, and monomeric HF.

Very strong hydrogen bonding occurs in the alkali metal hydrogen fluorides of formula M[HF2]; in KHF2, for example, an X-ray diffraction study together with a neutrons diffraction study shows that there is a liner symmetrical anion having an over all, F – H – F distance of 2.26 Å, which may be compared with the H-F bond length of 0.92Å in hydrogen fluoride monomer.


Intermolecular forces (Vander Waal’s Forces):

International attractions hold two or more molecules together. These are weakest chemical forces and can be of following types.

(a) lon-dipole attraction :

(b) Dipole-dipole attraction :

(c) lon-induced dipole attraction :

(d) Dipole-induced dipole attraction :

(e) Instantaneous dipole – Instantaneous induced dipole attraction : (Dispersion force or London forces)

· Strength of vander wall force ∝ molecular mass.

· Van der Wall’s force ∝ boiling point.


Metallic bond :

Most metals crystallise in close-packed structures. The ability of metals to conduct electricity and heat must result from strong electrons 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, delocalised 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 less number of electrons than the 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 :

Back bonding generally takes place when out of two bonded atoms one of the atom has vacant orbitals (generally this atom is from second or third period) and the other bonded atom is having 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 2p empty orbital.


The extent of back bonding is much larger if the orbitals involved in the back bonding are of same size, for example the extent of back bonding in boron trihalides is as follows :

BF3 > BCl3 > BBr3

More from us:

JEE Main Syllabus with weightage

JEE Main Question Paper 2019 with Solutions

Download Gradeup, the best IIT JEE Preparation App
Attempt subject wise questions in our new practice section

All the best!!

Team Gradeup


write a comment
Load Previous Comments
Nihal Tiwari

Nihal TiwariSep 15, 2018
watch this video for chemical bonding
Tushar Atrish
Sir thoda advanced level ka provide krao

AnjaniJul 8, 2021

Very useful

Follow us for latest updates