Hello friends welcome in this website, we shall discuss Chemical bonding in detail in this article. We come across all the types of bonds, regarding its important theories, their limitations and properties of the compounds formed by these bonds. Chemical bonding is considered as basic chemistry. Hence, it is very important to study this chapter so that we can complete the preparation of the whole chemistry very carefully and easily. The preparation of this chapter is also essential because many questions are asked always from this chapter in any competitive exams like engineering, medical, board exams of 12th, 10th, IIT advance and other exams. Let’s we start the notes of chemical bonding.
What is Chemical bonding?
The atom of each element wants to acquire stability like inert gases. As we observe that inert gases atom has eight electrons in there outermost shell and has stability. Therefore, atoms of other elements also involved in chemical bonding for stability.
The force which holds the atoms together within a molecule is called a chemical bond and this process is called chemical bonding.
Important theories about chemical bonding:
(1) Kossel-Lewis approach
(2) VSEPR theory
(3) Valence Bond Theory
(4) Molecular Orbital Theory
Kossel-Lewis Approach To Chemical bonding:
The atoms of different elements combine with each other in order to complete their respective octets or duplets in case of H, Li and Be to attain stable nearest noble gas Configuration. This can occur in two ways.
(1) By complete transfer of one or more electrons from one atom to another.
(2) By sharing of electrons between two atoms of same or different elements.
Types of Chemical bonding:
On the nature of exchange of electrons between two atoms, there are three types of chemical bonding.
(1) Electrovalent bond or Ionic bond
(2) Covalent bond
(3) Coordinate bond.
Lewis symbols- Representing the valence electrons:
G.N. Lewis introduced simple symbols to denote the valence shell electrons in an atom. The outer shell electrons are shown as dots surrounding the symbol of the atom. These symbols are known as Lewis symbols or electron dot symbol. A few examples are given below:—- 
Electrovalent or Ionic bond in Chemical bonding:
A chemical bond that is formed by complete transfer of electrons from one atom to another atom so as to complete their outermost shell by 8 electrons or in some cases by 2 electrons is called electrovalent bond.
When an atom loses electrons, it forms cation while other atom which accepts these electrons form anion. These cation and anion joined together by electrostatic force to form chemical bond and hence this bond is also known as Ionic bond.
Some examples of Ionic bonds:
The number of electrons lost or gained during the formation of electrovalent bond is termed as the electrovalency of the element. For example in formation of NaCl, the electrovalency of Na and Cl is one.
Factors affecting the formation of Ionic bonds:
(1) Ionisation enthalpy: The lower the value of Ionisation enthalpies, greater the chances of ionic bond formation.
(2) Electron gain enthalpy: Higher is the negative value of electron gain enthalpies, easier is the formation of anions and hence greater the chances of Ionic bond formation.
(3) Lattice enthalpy: The higher the value of lattice enthalpy of the resulting ionic compounds, the greater will be the stability of the compound and hence greater will be the ease of its formation.
General characteristics of Ionic compounds in chemical bonding:
(1) Physical state: These compounds usually exist in the solid state.
(2) Crystal structure: In ionic compounds, their ions are arranged in a regular pattern in the three dimensional space to form a lattice.
(3) Melting and boiling point: Ionic compounds possess high melting and boiling point. This is because ions are tightly held together by strong electrostatic forces of attraction.
(4) Solubility: Electrovalent compounds are soluble in solvents like water which are polar in nature and have high dielectric constant.
(5) Electrical conductivity: Ionic compounds are good conductors of electricity in solution or in the molten state movement of free ions.
(6) Ionic reactions: The reactions of Ionic compounds are, in fact, the reaction between the ions produced in solution.
Covalent Bond-Lewis-Langmuir Concept:
The bond formed between the two atoms by mutual sharing of electrons to complete their octet or duplet is called covalent bond and the compounds formed by this bond are called covalent compounds. Some important examples of the formation of covalent compounds.
Bond pairs and lone pairs: The shared pair of electrons present between the atoms are called bond pair of electrons. On the other hand, the valence pair of electrons not involved in bonding are called lone pairs or unshared pair of electrons.
Lewis structure of molecules in Chemical bonding:
Writing of Lewis dot structure of molecules and ions can be done in the following steps.
(1) Calculate the total number of valence electrons of the atoms present.
(2) If the species is an anion, add number of electrons equal to the units of -ve charge and if the species is a cation, subtract number of electrons equal to the units of +ve charge. This gives the total number of electrons to be distributed.
(3) Select the central atom which is generally the least electronegative and draw the skeletal structure by intelligent guess to indicate which atom is linked with other atoms. Also remember that H and F usually occupy terminal position.
(4) Put one shared pair of electrons between every two atoms to represent a single bind between them. Use the remaining pair of electrons either for multiple bonds or to show them as lone pairs, keeping in mind that octet of each atom is completed.
Formal charge on an atom in a molecule/ion:
The formal charge on an atom in a molecule/ion is defined as the difference between the number of valence electrons of that atom in the free state and the number of electrons assigned to that atom in the Lewis structure either as bond pair or lone pair of electrons.
Where FC = V – L + S/2
For example, calculation of formal charge on each O-atom of O3 molecule. The Lewis structure of ozone may be drawn as-
Limitations of the octet rule in Chemical bonding :
(1) Formation of compounds by hydrogen atoms violates the octet rule and follow duplet rule.
(2) Formation of electron deficient compounds like BeCl2, BF3, AlCl3 etc.
(3) Formation of hypervalent compounds like PCl5, SF6, IF7, H2SO4 etc.
(4) Formation of compounds by noble gases like XeF2, XeF4, XeF6, KrF2 etc.
(5) Formation of odd electron molecules like NO, NO2, O2, O2– etc.
VSEPR Theory in Chemical bonding:
VSEPR Theory in Chemical bonding means valence shell electrons pair repulsion theory. This theory was given by Sidgwick and Powell in 1940 and was further improved by Nyholm and Gillespie in 1957. The basic concept of this theory is-
“The electron pairs surrounding the central atom repel one another and move so far apart from one another that there are no further repulsions between them. As a result, the molecule has minimum energy and maximum stability”.
The above concept leads to the following assumptions:
- The shape of a molecule containing only two atoms is always linear.
- For molecules containing 3 or more atoms, one of the atom is considered as central atom to which other atoms are linked.
- If the central atom is linked to similar atoms and is surrounded by bond pair of electrons only, the shape of this molecule is symmetrical and the molecule is said to have a regular geometry.
- If the central atom is linked to different atoms or is surrounded by bond pairs as well as lone pairs of electrons, the shape of molecule is unsymmetrical and the molecule is said to have an irregular geometry.
- The exact shape of the molecule depends upon the total number of electron pairs present around the central atom.
Order of repulsion in valence shell electrons:
Lone pair – Lone pair > Lone pair – Bond pair > Bond pair – Bond pair.
Determination of geometry or shape of molecules or ions by VSEPR theory:
For the determination of geometry or shape, we need to find out the type of molecules in the following way.
(1) Valence shell electron pairs in a molecule or ion = (V.E.+ No. Of monovalent S.A. – C + A)/2
(2) No of bond pairs = No of S.A. to the central atom
(3) No of lone pair of electrons = VSEP – BP
Now fix the type of molecules as ABxLy and follow the chart given below:—- For example of NH4+
VSEP= (5 + 4 – 1)/2 = 4
NO. of bond pairs (BP) = 4
No. of lone pairs (LP) = 4 – 4 = 0
Hence, The type of molecule = AB4L0 and shape will be Tetrahedral.
Valence bond theory in Chemical bonding:
This theory was put forward by Hitler and London in following ways.
(1) In terms of energy consideration
When two atoms of same element or different elements come closer to each other, the new forces come into operation.
⇒The force of repulsion between the nuclei of these combining atoms and between the electrons of these atoms. These forces tend to increase the energy.
⇒The forces of attraction between the nucleus of one atom and electrons of other atom. These forces tend to decrease the energy of the system.
⇒If in a system, these new forces can decrease the energy, then possibility of chemical bonding exists and if these forces lead to increase in energy, the chemical bonding is not possible.
(2) In terms of orbitals overlap concept-
According to this concept, lowering of energy takes place when atomic orbitals of the approaching atoms partially interpenetrate into each other. Hence, A covalent bond is formed by the partial overlap of two half- filled atomic orbitals containing electrons with opposite spin.
Types of covalent bonds: in Chemical bonding
Depending upon the type of overlapping, the covalent bonds are mainly of two types-
(1) Sigma (σ ) bond- When a bond is formed between two atoms by the overlap of their atomic orbitals along the internuclear axis, this bond is called sigma bond.
The sigma bond can take place in the following overlapping:
• s-s overlapping
• s-p overlapping
• pz-pz overlapping along the internuclear axis
(2) Pi (Π ) bond- When a covalent bond is formed by lateral overlapping of p- orbitals in direction at right angles to the internuclear axis is called pi bond.
• A pi bond is seldom formed between atoms unless sigma bond is formed.
• pi bond is formed by px-px and py-py overlapping.
• Sigma bond is stronger than pi bond due to quite large overlapping.
Some Important Bond Characteristics in Chemical bonding:
Bond Length: In Chemical bonding
The equilibrium distance between the centres of the nuclei of the two bonded atoms is called the bond length.
Factors affecting bond length:
(A) Size of the atoms. The bond length increases with increase in the size of atoms. For example, bond length of H-X are in the order.
HI > HBr > HCl > HF.
(B) Multiplicity of bond. The bond length decreases with the Multiplicity of bond. The increasing order of bond length is following.
C ≡ C < C = C < C – C
(C) Type of Hybridisation. As an s-orbital is smaller in the size, greater the s- character, shorter is the hybrid orbitals and hence shorter is the bond length. For example, order of bond length is
sp3 C – H > sp2 C – H > sp C – H
Bond enthalpy (Bond Energy)
The amount of energy required to break one mole of bonds of a particular type so as to separate them into gaseous atoms is called bond enthalpy. Bond enthalpy increases, bond length decreases.
Factors affecting bond enthalpy:
(A) Size of the atoms. Greater the size of the atoms, larger is the bond length and smaller is the bond enthalpy.
(B) Multiplicity of bonds. For the bond between the same two atoms, greater is the multiplicity of bonds, greater is the bond dissociation enthalpy. For example,
H – H < O = O < N≡ N
(C) Number of lone pair of electrons present. Greater the number of lone pair of electrons present on the bonded atoms, greater is the repulsion between the atoms and hence less is the bond dissociation enthalpy. For example, C – C > N – N > O – O > F – F
Bond angles in Chemical bonding
The angle between the lines representing the directions of the bonds is called bond angles.
Factors affecting the bond angles:
(A) Hybridisation. The s- characters in hybridisation increases, the bond angles decreases. For example, sp = 180o, Sp2 = 120o, sp3 = 109.28′
(B) Lone pair of electrons on central atom. The number of lone pair of electrons increases on the central atom, the bond angles decreases. For example, board angle in CH4 = 109.28′, NH3 = 107o, H2O = 104.5o.
(C) Size of the central atom. If molecules or ions have same hybridisation and no. of lone pair of electrons, the size of central atom decreases the bond angles increases. For example, The bond angle of NH3 > PH3 > AsH3 > SbH3
(D) Size of surrounding atom. The size of surrounding atoms increases, the bond angles increases when hybridisation, lone pair of electrons and central atom are same. For example, NF3 < NCl3 < NBr3.
Bond order in Chemical bonding
The number of bonds present between two atoms in any molecules or ions is called bond order. For example, Bond order in H2 (H – H) = 1, in O2 (O = O) = 2
• The isoelectronic species have same bond orders. For example, N2, CO and NO+ have bond orders = 3.
• When a molecule/ion show resonance, the bond order between any two atoms is total no of bonds between any two atoms / no of resonating structure. For example, in Benzene, bond order = 3/2 = 1.5
• For odd electron molecules, as the three electron bond is considered as equivalent to half electron bond. For example, Lewis structure of NO is :N≅ O: , Hence, its bond order = 2.5
Polar and non polar covalent bonds: In Chemical bonding
Non-polar covalent bond. If two similar atoms are bonded together by mutual sharing of electrons, this bond is called polar covalent bond. For example, H2, Cl2, O2 etc.
Polar covalent bond. When two dissimilar atoms having different electronegativities combine together by mutual sharing of electrons, this bond is called polar covalent bond. For example NH3, H2O, HCl etc.
• It may be noted that in case of symmetrical molecules like CO2, CCl4 etc. although there are a number of polar bonds present, yet the molecules on the whole are non- polar. This is because the polar bonds cancel the effect of each other.
Partial Ionic Character of Covalent Bond :
The extent of partial ionic character is determined by the difference in electronegativities of the combining atoms. More is the difference in electronegativity, greater will be the ionic character.
• If the electronegativity difference between two atoms is 1.9, the bond is said to have 50% ionic and 50% covalent character.
• If the electronegativity difference between two atoms is more than 1.9, the partial ionic character of the bond is more than 50% and the bond is taken as ionic.
• If the electronegativity difference between two atoms is less than 1.9, the bond is predominantly covalent.
Dipole moment in Chemical bonding:
The product of magnitude of negative or positive charge (q) and the distance(d) between these two ions is called dipole moment. It is denoted by ‘μ’.
and μ = q × d
Unit of dipole moment- the unit of dipole moment is Debye (D) and
1D = 10-18 esu.cm
again 1D = 3.335 × 10-30 Cm (coulomb metre)
Comparison of dipole moment: in Chemical bonding
• In case of diatomic molecules, the greater is the difference in electronegativity of two covalently bonded atoms, the larger is the dipole moment. For example, H-F > H-Cl > H-Br > H-I.
• In case of polyatomic molecules, the dipole moment is equal to the resultant dipole moment of all the individual bonds called bond moments.
• In case of symmetrical molecules like CO2, BF3, CH4 have zero dipole moment.
In case of unsymmetrical polyatomic molecules, their dipole moment is greater than zero and it increases after increases in lone pair of electrons on central atom or unsymmetriies.
Application of dipole moment: in Chemical bonding
(1) In determining the polarity of bonds. Greater is the magnitude of dipole moment, higher will be the polarity of bond.
(2) In the calculation of percentage ionic character. % ionic character = observed dipole moment/ionic dipole moment × 100
(3) In determining the symmetry or shape of the molecules.
(4) To distinguish between cis and trans isomers.
(5) To distinguish between ortho, meta and para isomers.
Also read : Classification of Elements and Periodicity in properties
Polarisation in Chemical bonding:
when a cation comes closer to an anion, the election cloud of the anion is attracted towards the cation and hence gets distorted. This effect is called polarisation. The power of the cation to polarise the anion is called its polarising power and the tendency of the anion to get polarised is called its polarisability.
Fajan’s Rule in Chemical bonding:
Due to polarisation, the covalent character in the ionic compounds increases. Hence, the properties like melting point, solubility in water, heat of sublimation changes. According to Fajan’s rule, covalent character is favoured by the following factors.
(1) Smaller is the size of cations, greater is the covalent character.
(2) Larger is the size of anions, greater is the covalent character.
(3) Larger the charge on the cation or anion, greater is its polarising power and hence greater is the covalent character
(4) If two cations have the same size and charge, then the one with pseudo noble gas Configuration has greater polarising power than the other with noble gas Configuration. For example CuCl is more covalent than NaCl.
Characteristics of covalent compounds in chemical bonding:
(1) Unlike ionic compounds, the covalent compounds exist in all the three states like solid, liquid and gas.
(2) The crystal structure of covalent compounds differ from that of ionic compounds. They usually consists of molecules rather than ions.
(3) Covalent compounds have low melting and boiling points because their molecules are held rigidly.
(4) Covalent compounds are not soluble in polar solvents like water but soluble in non polar solvents like benzene, ether and other organic solvents.
(5) Covalent compounds are bad conductors of electricity.
(6) Covalent compounds show isomerism because their molecules are directional and have definite geometry.
Hybridisation in Chemical bonding:
The mixing of the atomic orbitals belonging to the same atom but having slightly different energies so that a redistribution of energy takes place between them resulting in the formation of new orbitals of equal energies and identical shapes is called hybridisation and new orbitals thus formed are known as hybrid orbitals.
Determination of hybridisation : in Chemical bonding
First upon all, calculate the number of hybrid orbitals(X) to formed by the central atom as follows:
X = 1/2 [VE + MA – C + A]
Where VE = no of valence electrons of central atom.
MA = no of monovalent atoms/ groups surrounding the central atom.
C = charge on the cations.
A = charge on the anions.
After finding the hybrid orbitals, we decide the hybridisation in the following way: 
Coordinate or Dative bond in Chemical bonding:
Perkin in 1921 suggested a third possible ways by which atoms can combine and form a molecule. According to him, when in the formation of a bond, the electron pair (lone pair) is denoted by one atom but shared by both the atoms to complete their octets; the bond formed is called coordinate bond or dative bond. For example, formation of Sulphuric acid. 
Resonance in Chemical bonding:
In case of certain molecules, a single Lewis structure cannot explain all the properties of the molecule. The molecule is then supposed to have many structures. These structures are called resonating structures or canonical structures and this phenomenon is called resonance. The actual structure in between of all these resonating structures is called resonance hybrid.
For example the structure of ozone can be written as:
Rules for writing of resonating structures.
(1) The contributing structures should have the same position of atoms. They should differ only in the position of electrons.
(2) The contributing structures should have the same number of electrons.
(3) The contributing structures should have nearly same energy.
(4) The contributing structures should have negative charge on the electropositive atom and positive charge on the electronegative atom.
(5) In a contributing structure, like charges should not be present on adjacent atoms while unlike charges should not be widely separated.
Characteristics of Resonance:
(1) The contributing structures do not have any real existence. They are imaginary.
(2) As a result of resonance, the bond lengths in a molecule become equal.
(3) The resonance hybrid has lower energy and hence greater stability than any of the contributing structures.
(4) Greater is the resonance energy, greater is the stability of the molecule.
(5) Greater is the number of canonical structures especially with nearly same energy, greater is the stability of the molecule.
Molecular Orbital Theory: in Chemical bonding.
This theory was developed by Hund and Mulliken regarding Chemical bonding in the following way.
(1) Like valence bond theory, atomic orbitals of different atoms overlap to form new orbitals called molecular orbitals.
(2) Molecular orbitals are the energy state of a molecule in which the electrons are filled just as atomic orbitals in atoms.
(3) Only those atomic orbitals can combine to form molecular orbitals which have comparable energy and proper orientation.
(4) The number of molecular orbitals formed is equal to the number of combining atomic orbitals.
(5) When two atomic orbitals combine, they form two new orbitals called bonding molecular orbitals and anti bonding molecular orbitals.
(6) The bonding molecular orbitals have lower energy and hence greater stability than the corresponding anti bonding molecular orbital.
(7) The bonding molecular orbitals are represented by α, β, γ, etc. While anti bonding molecular orbitals are represented by α*, β*, γ*, etc.
Stability of molecules in molecular orbital theory:
(1) If Nb > Na, the molecule is stable.
(2) If Nb = or < Na, the molecule is unstable.
Where Nb = total number of electrons in bonding molecular orbitals.
Na = total number of electrons in anti bonding molecular orbitals.
Stability of molecules in term of bond order.
Bond order in molecular orbital theory is defined as half of the difference between the number of electrons present in the bonding and anti bonding molecular orbitals, i.e, Bond order (B.O.) = (Nb – Na)/2
The bond order is larger, the stability of the molecule is greater and bond dissociation enthalpy is lower.
Magnetic nature by molecular orbital theory:
When unpaired electrons are present in the molecular orbitals of any molecule, it is paramagnetic and when only paired electrons are present, it is diamagnetic. The greater is the number of unpaired electrons, the paramagnetic nature of this molecule is greater. We can also find magnetic moment with the help of unpaired electrons in the following way.
Magnetic moment (μ) = √n(n+2) BM.
Filling of electrons in molecular orbitals of molecules having electrons < or =14. .
σ1s < σ*1s < σ2s < σ*2s < π2px = π2py < σ2pz < π*2px = π*2py < σ*2pz
Filling of electrons in molecular orbitals of molecules having electrons >14.
σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2px = π2py < π*2px = π*2py < σ*2pz
Hydrogen Bond in Chemical bonding:
In Chemical bonding, when hydrogen atom is linked to a highly electronegative atom like F, O or N, this atom attracts the shared pair of electrons more and so this end of the molecule becomes slightly negative while the other end becomes slightly positive. The negative end of one molecule attract the positive end of the other and as a result, a weak bond is formed between them. This bond is called hydrogen bond.
Conditions for hydrogen bond in Chemical bonding:
(1) The molecule must contain a highly electronegative atom like F, O or N linked to H-atom.
(2) The size of the electronegative should be small.
Types of hydrogen bonding:
In this type of Chemical bonding, there are two types of hydrogen bonding.
(1) Intermolecular hydrogen bond- When hydrogen bond is formed between different molecules of the same or different compounds, it is called intermolecular hydrogen bond.
(2) Intramolecular hydrogen bond-
The hydrogen bonding which takes place within a molecule itself. It is called intramolecular hydrogen bond.
• Intermolecular hydrogen bond is stronger than intramolecular hydrogen bond.