•`color{green}["Chemical bond"]`:The attractive force which holds various constituents (atoms, ions, etc.)together in different chemical species is called a chemical bond.

•`color{green}["Lewis symbols"]`: G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols.

•`color{green}["Electrovalent bond"]`: The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond.

•`color{green}["Octet rule"]`: According to Lewis and kossel, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.

•`color{green}["Covalent bond"]`:The bond formed between the two atoms by mutual sharing of electrons between them so as to complete their octets or duplets in case of elements having only one shell is called covalent bond and the number of electrons contributed by each atom is known as covalency.

•`color{green}["Single Covalent bond"]`:When two atoms share one electron pair they are said to be joined by a single covalent bond.

•`color{green}["Double Covalent bond"]`: If two atoms share two pairs of electrons, the covalent bond between them is called a double bond.

•`color{green}("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.
`color{green}(= [text(total number of valence electrons in the free atom)] - [text{total number of non bonding (lone pair) electrons}] - (1//2) [text{total number of bonding(shared) electrons}])`

•`color{green}("Electrovalent bond ")` :When a bond is formed by complete transference of electrons from one atom to another so as to complete their outermost orbits by acquiring 8 electrons or 2 electrons in case of hydrogen,lithium etc. and hence acquire the stable nearest noble gas configuration,the bond formed is called ionic bond or electrovalent bond.

•`color{green}("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.

•`color{green}("Bond length ")` : Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule.

•`color{green}("Covalent Radii ")`: The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation.

•`color{green}("van der Waal's Radii" )`: The van der Waals radius represents the overall size of the atom which includes its valence shell in a non-bonded situation.

•`color{green}("Bond angle" )`: It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion.

•`color{green}("Bond enthalpy" )`: It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.

•`color{green}("Bond order" )`: In Lewis representation of a molecule or ion, the number of bonds present between two atoms is called bond order.

•`color{green}("Dipole moment" )`:It can be defined as the product of the magnitude of the charge and the distance between the centres of positive and negative charge.

• `color{green}("Sigma (σ) bond")` : This type of covalent bond is formed by the end to end (hand-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap.

•`color{green}("pi (π) bond")` : In the formation of `π` bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis.

•`color{green}("Hybridisation ")` It can be defined as the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape.

•`color{green}("sp Hybridisation ")` :This type of hybridisation involves the mixing of one `s` and one `p` orbital resulting in the formation of two equivalent `sp` hybrid orbitals.

• `color{green}( sp^2 "Hybridisation ")`: In this hybridisation there is involvement of one `s` and two `p`-orbitals in order to form three equivalent `sp^2` hybridised orbitals.

•`color{green}(sp^3 "Hybridisation ")` :This type of hybridisation can be explained by taking the example of `CH_4` molecule in which there is mixing of one `s`-orbital and three `p`-orbitals of the valence shell to form four `sp^3` hybrid orbital of equivalent energies and shape.

•`color{green}("Shape of molecule ")` : A particular arrangement obtained by bonding a number of atoms in definite directions to the central atom of a molecule is called geometry or shape of molecule.

•`color{green}("Bonding molecular orbital ")` : The molecular orbital formed by the additive effect of the atomic orbitals is called bonding molecular orbitals.

•`color{green}("Antibonding molecular orbital ")` : The molecular orbital formed by the subtractive effect of the atomic orbitals is called antibonding molecular orbital.

•`color{green}("Hydrogen bond ")` :Whenever a molecule contains a hydrogen atom linked to a highly electronegative atom(like N,O or F), 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 attracts the positive end of the other and as a result, a weak bond is formed between them. This bond is called hydrogen bond.

•`color{green}("Intramolecular Hydrogen bond ")`: It is formed when hydrogen atom is in between the two highly electronegative (`F`, `O`, `N`) atoms present within the same molecule.

•`color{green}("Intermolecular Hydrogen bond ")`: It is formed between two different molecules of the same or different compounds.

•`color{green}("Resonance ")`: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, each of which can explain most of the properties of the molecule but none can explain all the properties of the molecule. The actual structure is in between of all these contributing structures and is called resonance hybrid and the different individual structures are called resonating structures or canonical forms. The phenomena is called resonance