1. CHEMICAL BONDING

2. 1.Introduction 2.Octet rule 3.Different types of bonding 4.Valency Bond Theory Topics Covered

3. Force of attraction holding group(s) of atoms What is a Chemical Bond? Chemical bonds Na + Cl - Better stability against chemical reagents But why bonds are formed ??

4. Atoms two electrons in the valence shell (1s 2 ) Octet rule noble gas configuration to attain better stability. Na 281 Very reactive Na 2 8 + Ne Cl 287 Very reactive Cl 288 Ar -

5. In SF 6, ‘S’ has twelve electron in its valence shell, leads to minimisation of energy. Other examples are: PCl 5, BF 3 Limitation of octet rule

6. Bonding Ionic Covalent Co-ordinate or dative Metallic Pi bond Sigma bond

7. Formation of ionic bond

8. Covalent bond Formed by mutual sharing of electrons Covalent bonds Double bond Triple bond

9. non-polar covalent bond between two carbon atoms polar covalent bond between carbon and hydrogen atoms. Formation of covalent bond

10. Covalent bonds are called directional while ionic bonds are called non-directional -explain Solution: Illustrative Problem p and d-orbitals generate directional covalent bond. electrostatic force of attraction. Ionic bond overlap of atomic orbitals covalent bond

11. Strength of these sigma bonds is in the order: sigma bond forms due to end-to-end or head-on overlap p-p + s-s + + Orbital Overlap Concept s-p p-p > s-p >s-s

12. This is formed by lateral or sideways overlap which is possible for p or d-orbitals. Sigma bond is stronger than pi bond due to greater extent of overlap. + or Orbital Overlap Concept

13. Difference between sigma and pi bonds Stronger as compared to bond Weaker as compared to bond Formed by head-on overlapping of s-s or s-p or p-p or any hybrid orbital Formed by side ways overlapping of unhybridised p-orbital First bond between any two atoms is always sigma Rest are  bonds In plane of molecule Perpendicular to plane of molecule

14. Valence Bond (VB) Theory, the theory we will explore, describes the placement of electrons into bonding orbitals located around the individual atoms from which they originated. COVALENT BOND FORMATION (VB THEORY) In order for a covalent bond to form between two atoms, overlap must occur between the orbitals containing the valence electrons. The best overlap occurs when two orbitals are allowed to meet “head on” in a straight line. When this occurs, the atomic orbitals merge to form a single bonding orbital and a “single bond” is formed, called a sigma () bond.

16. MAXIMIZING BOND FORMATION In order for “best overlap” to occur, valence electrons need to be re-oriented and electron clouds reshaped to allow optimum contact. To form as many bonds as possible from the available valence electrons, sometimes separation of electron pairs must also occur. We describe the transformation process as “orbital hybridization” and we focus on the central atom in the species...

17. Hybridization of Be in BeCl 2 Atomic Be: 1s 2 2s 2 Hybrid sp orbitals: 1 part s, 1 part p

18. FORMATION OF BeCl 2 : Each Chlorine atom, 1s 2 2s 2 2p 6 3s 2 3p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp orbital of the beryllium to form a sigma bond.

20. “sp 2 ” Hybridization: All 3 Region Species

21. Hybridization of B in BF 3 Atomic B : 1s 2 2s 2 2p 1 Valence e’s Hybrid sp 2 orbitals: 1 part s, 2 parts p

22. FORMATION OF BF 3 : Each fluorine atom, 1s 2 2s 2 2p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp 2 orbital of the boron to form a sigma bond.

24. Hybridization of C in CH 4 Atomic C : 1s 2 2s 2 2p 2 Valence e’s Hybrid sp 3 orbitals: 1 part s, 3 parts p

25. FORMATION OF CH 4 : Each hydrogen atom, 1s 1, has one unshared electron in an s orbital. The half filled s orbital overlaps head-on with a half full hybrid sp 3 orbital of the carbon to form a sigma bond.