This theory explains about the cause of chemical reaction/combination on the basis of valence shell electron (i.e electron present in the outermost orbit) & tendency to lose or gain an electron. The outermost shell of an atom is called valence shell &electron present in it is called valence electron. The valence electron of an atom is denoted by Lewis dot symbol.For E.g.
In chemical reaction/combination very atom tries to adjust it’s electronic configuration like nearest noble gas by losing r gaining or sharing of the electron. This rule is called octet rule.
Types of bond
The bond formed by gaining or losing or sharing valence electron with the another reacting atom is called ionic bond or electrovalent bond. The compound having the ionic bond is called the ionic compound. The force of attraction is the electrostatic force.
E.g. NaCl , CaCl2 , AlF3
factors affecting ionic bond
Characteristics of Ionic bond
Formation of Ionic Compound
The bond formed by sharing of the electron between reacting atoms is called the covalent bond. Generally, non-metals form this type of bond. This bond is denoted by (-). The compounds having covalent bond are called covalent compounds.
Characteristics of Covalent bond
Formation of Covalent Bond
Formation of O2 molecule
Coordinate covalent bond
The bond formed of sharing of the electron but shares electron in pairs & thus formed the bond is called Coordinate covalent bond. The bond is formed between electron deficient & electron excess species. This is denoted by an arrow pointing from donor to receptor of the electron.
Formation of coordinate covalent bond
When covalent bond forms then the bonding pair of the electron does not lie exactly midway between the reacting atoms but the pair of electron lies ear toward the atom which has more electronegativity. Since, the electrons shift towards it the atom having more electronegativity bears partial negative charge and the atom having less electronegativity possess partial positive charge. Such molecules are called polar molecules & the bond is called the polar covalent bond.
For e.g. In HCl, chlorine has more electronegativity than hydrogen so shared electron shift towards chlorine and hence chlorine bears partial negative charge whereas hydrogen bears partial positive charge.
It is defined as the product of the magnitude of charge developed in an atom of a polarized molecule & the distance between the combing atom. It is denoted by . Its unit is Debye (D).
If net dipole moment of the molecule is zero then molecule will be non-polar & possess regular geometry & if dipole moment is not equal to zero then molecule will be polar as well as it has the irregular geometry.
Structure of BeF2
In the case of BeF2 molecule , the dipole moment of one beryllium fluorine is cancelled by another beryllium fluorine bond . Hence, the molecule is non-polar and possess linear molecule.
Structure of H2O
The dipole moment of H2O is 1.84 & dipole along H-O is 1.5D. In molecule, there is two H-O bonds not opposing straight to one another which result in distortion of the linear structure of H2O molecule. Two H-O bonds of H2O molecules are inclined to each other & bears an angular V-shape.
Application of dipole moment
Special type of bond
This is defined as the bond formed by the electrostatic force of attracting between polarized hydrogen of one molecule which is already bonded to highly electronegative element.
Types of hydrogen bonding
If hydrogen bond exists between more than two molecules then hydrogen bonding is said to be intermolecular hydrogen bonding.
If hydrogen bond exists within a molecule than it is called intramolecular hydrogen bonding.
It is a special type of bond that exist only in metal. It is the force of attraction between positively charged kernel & sea of mobile electron which is responsible for binding metal ions together. Metallic bonding is responsible for malleability &ductability.
The force of attraction between electrically neutral molecule that collides with each other. It is formed by the temporary attraction between electron rich species and electron deficient species.
The electronic phenomenon in which a molecule shows more than 2 lewis structure due to delocalization of π electron.
poulse, tracy. Introduction to chemistry. u.s.a: flexbook, 2010.
Pathak, Sita Karki. The Text Book of Chemistry. kathmandu: Vidhyarthi Pustak Bhandar, 2012.
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