Kossel-Lewis approach –<\/strong> Elements other than noble gases, try to attain the completely filled electronic configuration by losing, gaining or sharing one or more electrons from their outer shell.<\/p>\n Octet rule –<\/strong> The atoms transfer or share electrons so that all atoms involved in chemical bonding obtain 8 electrons in their outer shell.<\/p>\n Covalent bond –<\/strong> This type of mutual sharing of one or more pairs of electrons between two combining or more pairs of electrons between two combining atoms results in the formation of a chemical bond called a covalent bond.<\/p>\n Formal charge –<\/strong> Formal charge of an atom in a molecule is the electrical charge difference between the valence electron in an isolated atom and the number of electrons assigned to that atom in the lewis structure. Ionic bond (or) electrovalent bond –<\/strong> The bond formed by the complete transfer of electrons leads to the formation of cation and an anion which are held together by the electrostatic attractive force.<\/p>\n Coordinate covalent bond –<\/strong> In the bond formation, one of the combining atoms donates a pair of electrons and these electrons are shared by both the combining atoms. These type of bonds are called coordinate covalent bond.<\/p>\n Bond length –<\/strong> The distance between the nuclei of the two covalently bonded atoms is called bond length.<\/p>\n Bond order –<\/strong> The number of bonds formed between the two bonded atoms in a molecule is called the bond order.<\/p>\n Bond angle –<\/strong> A fixed angle created by the directional nature of the two covalent bonds which is between the two bonds in a molecule is called bond angle.<\/p>\n Bond enthalpy –<\/strong> It is defined as the minimum amount of energy required to break one mole of a particular bond in the molecule in their gaseous state.<\/p>\n Resonance –<\/strong> The lewis structures which differ only in the position of bonding and lone pair of electrons are called resonance structures and this phenomenon is called resonance.<\/p>\n Dipole moment –<\/strong> \u00b5 = q x 2d, where q = charge, 2d= distance between the two charges.<\/p>\n Unit of dipole moment –<\/strong> Debye unit (D), 1 Debye = 3.336 x 10-30<\/sup> C m (C m = coulomb metre).<\/p>\n Non polar molecules –<\/strong> Molecules that have zero dipole moment are called non polar molecules. For example, H2<\/sub>, O2<\/sub>, F2<\/sub>.<\/p>\n Polar molecules –<\/strong> Molecules that have dipole moment value are called polar molecules. For example, HF, HCl, NO, CO.<\/p>\n Electronegativity difference<\/strong><\/p>\n If XA<\/sub> – XB<\/sub> = 1.7, the bond A – B has 50% ionic character Polarisation –<\/strong> A distortion in the electron cloud of the anion and its electron density drift towards the cation which results in some sharing of the valence electrons between these ions. Thus, a partial covalent character is developed between them. This phenomenon is called polarisation.<\/p>\n Fajan\u2019s rule –<\/strong> The extent of polarisation in an ionic compound is given by Fajan\u2019s rule.<\/p>\n VSEPR theory –<\/strong> Valence Shell Electron Pair Repulsion Theory: It is useful in predicting the shape of the molecules.<\/p>\n The repulsive interaction between the electron pairs –<\/strong> lp – lp > lp – bp > bp – bp lp = long pair, bp = bond pair<\/p>\n Shapes of molecules predicted by VSEPR theory –<\/strong> Sigma bond (\u03c3) –<\/strong> When two atomic orbitals overlap linearly along the axis, the resultant bond is called a bond.<\/p>\n Pi bond (\u03c0) –<\/strong> When two atomic orbitals overlap sideways, the resultant covalent bond is called pi bond (n bond).<\/p>\n Hybridisation –<\/strong> It is a process of mixing of atomic orbitals of the same atom with comparable energy to form equal number of new equivalent orbitals with same energy.<\/p>\n Types of hybridisation with shape and examples<\/strong> Bond order –<\/strong> It is the half difference of number of electrons present in bonding molecular orbitals and the number of electrons present in antibonding molecular orbitals. LCAO –<\/strong> Linear combination of atomic orbitals.<\/p>\n \u03c8 bonding –<\/strong> \u03c8A<\/sub> + \u03c8B<\/sub><\/p>\n \u03c8 antibonding –<\/strong> \u03c8A<\/sub> \u03c8B<\/sub><\/p>\n
\nFormal charge –
\n\\(\\mathrm{N}_{\\mathrm{V}}-\\left[\\mathrm{N}_{l}+\\frac{\\mathrm{N}_{b}}{2}\\right]\\)<\/p>\n
\nIF XA<\/sub> -XB<\/sub> > 1.7, the bond A – B > 50% ionic character
\nIF XA<\/sub> – XB<\/sub> < 1.7, the bond A – B < 50% ionic character<\/p>\n\n
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<\/p>\n
\nBond order = \\(\\frac{\\mathrm{N}_{b}-\\mathrm{N}_{a}}{2}\\)<\/p>\n