VBT & Hybridization Mr. Imran Saleem.pptx

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1. Prepared By : Mr. Imran Saleem Department of Chemistry APSACs Ord. Rd. RWP

2. CONTENT Valence Bond Theory σ-Bond π—Bond Comparison Illustration Hybridization Types of Hybridization sp,sp2,sp3

3. The valence bond theory was proposed by Heitler and London to explain the formation of covalent bond quantitatively using quantum mechanics. Later on, Linus Pauling improved this theory by introducing the concept of hybridization. Valence bond (VB) theory assumes that all bonds are localized bonds formed between two atoms by the donation of an electron from each atom.

4. Valence Bond theory d e s c ri be s covalent b ond f or m a t i o n a s w ell a s t h e el e c t r oni c s tr uc t u r e of molecules. The theory assumes that electrons occupy atomic orbital's that the of individual atoms within a molecule, and electrons of one atom are attracted to the nucleus of another atom.

5. A covalent bond is formed by the overlapping of two half filled valence atomic orbital's of two different atoms. The electrons in the overlappingorbital's get paired and confined between the nuclei of two atoms. The electron density between two bonded atoms increases due to overlapping. This confers stability to the molecule.

6. Greater the extent of overlapping, stronger is the bond formed. The direction of the covalent bond is along the region of overlapping of the atomic orbital's i.e., covalent bond is directional.

7. A S I G M A B O N D ( S Y M B O L : Σ ) I S A CO V A L E N T B O N D FORMED VIA LINEAR OVERLAP OF TWO ORBITAL'S. π-BOND A P i - B O N D ( S Y M B OL : π ) I S A C O V AL E N T BOND FORMED VIA PARALLEL OVERLAP OF TWO ORBITALS. There are two types of covalent bonds based on the pattern of overlapping as follows: Σ-BOND Π bond

8. SIGMA BOND The covalent bond formed due to overlapping of atomic orbital along the inter nucleus axis is called σ-bond. It is a stronger bond and cylindrically symmetrical. Depending on the types of orbital's overlapping, the σbond is divided into following types: (i): σs-s bond, (ii): σp-p bond, (iii): σs-p bond:

9. SIGMA BOND σ s-s bond :

10. SIGMA BOND σ p-p bond:

11. SIGMA BOND σs-p bond:

12. Π- BOND The covalent bond formed by sidewise overlapping of atomic orbital's is called π- bond. In this bond, the electron density is present above and below the inter nuclear axis. It is relatively a weaker bond since the electrons are not strongly attracted by the nuclei of bonding atoms. Note: The 's' orbital's can only form σ-bonds, whereas the p, d & f orbital's can form both σand π-bonds.

13. C OM P ARI S ON σ - b o n d This bond is formed due to the overlap of pure s- s;s-p;p-p (or) hybrid orbitals of two atoms along their internuclear axis. It is a strongest bond because the extent of overlapping of orbitals in sigma bond is greater. Electeron density of a sigma bond is symmetrical about the line joining the two nuclei. π-bond This bond is formed due to lateral or side wise or parallel overlapping of pure ‘p’ orbitals of two atoms. It is weaker than sigma bond because the extent of overlapping of orbitals in pi bond is lesser. Electron density of pi bond is unsymmetrical .

14. C OM P ARISON A sigma bond can present alone. In sigma bond free rotation of atom is possible. A sigma bond possesses high bond energy. A sigma bond is less reactive. A pi bond is always formed in addition to sigma bond. In pi bond free rotation is not possible. A pi bond possesses low bond energy. A pi bond is more reactive.

15. C OM P ARISON A sigma bond has greater bond length. Compound containing sigma bond generally undergo substitution reactions. A sigma bond influence the geometry of molecule. Examples: C H ₄,H ₂ , C l ₂ A pi bond has lesser bond length. Compound containing pi bond usually undergo addition reactions. A pi bond generally has no effect on geometry. Examples: CH₂=CH₂,N≡N,O=O

16. The electronic configuration of hydrogen atom in the ground state is 1s 1 . In the formation of hydrogen molecule, two half filled 1s orbital's of hydrogen atoms overlap along the inter- nuclear axis and thus by forming a σ s-s bond.

17. The electronic configuration of Cl atom in the ground x y z s tat e is [ N e] 3 s 2 3 p 2 3 p 2 3 p 1 . The two half filled 3p z atomic orbital's of two chlorine atoms overlap along the inter-nuclear axis and thus by forming a σ p-p bond.

18. In the ground state, the electronic configuration of hydrogen atom is 1s1 . And the ground state electronic configuration of Cl atom is [Ne]3s2 3px 2 3py 2 3pz 1 . The half filled 1s orbital of hydrogen overlap with the half filled 3pz atomic orbital of chlorine atom along the internuclear axis to form a σs-p bond.

19. The electronic configuration of O in the ground state is [He] 2s2 2px 2 2py 1 2pz 1 . The half filled 2py orbital's of two oxygen atoms overlap along the inter-nuclear axis and form σp-p bond. The remaining half filled 2pz orbital's overlap laterally to form a πp-p bond .

20. Thus a double bond (one σp-p and one πp-p ) is formed between two oxygen atoms.

21. The ground state electronic configuration of N is [He] x y z 2s 2 2p 1 2p 1 2p 1 . Aσ p-p bond is formed between two nitrogen atoms due to overlapping of half filled 2p x atomic orbital's along the inter-nuclear axis.

22. T h e r e m ai n i n g h alf f i l l e d 2 p y a n d 2 p z o r b i t a l ' s f o r m two π p-p bonds due to lateral overlapping. Thus a triple bond (one and two) is formed between two nitrogen atoms.

23. Limitations of VBT: This theory is unable to explain the paramagnetic behavior of O₂ molecule. This theory cannot properly describe the shape and geometry of molecules (e.g. unable to differentiate between square planar and tetrahedral geometry). It does not provide explanation to the formation of co-ordinate bond in which one of the bonded atoms gives both the electrons. This does not account for the formation of odd electron molecule (e.g. NO) or such as H₂+ ion where no pairing of electrons occurs. It does not explain the non existence of noble gas molecules. It does not provide explanation for the presence of fractional bonds in many molecules, e.g. benzene,CO₃¯² ion etc. This theory is also unable to explain bonding in electron deficient molecules.

24. The intermixing of two or more pure atomic orbital's of an atom with almost same energy to give same number of identical and degenerate new type of orbital's is known as hybridization. The new orbital's formed are also known as hybrid orbital's. During hybridization, the atomic orbitals with different characteristics are mixed with each other.

25. sp sp2 sp3

26. Intermixing of one 's' and one 'p' orbital's of almost equal energy to give two identical and degenerate hybrid orbital's is called 'sp' hybridization. These sp-hybrid orbital's are arranged linearly at by making 180 ⁰ of angle. They possess 50% 's' and 50% 'p' character.

27. For example:

28. Intermixing of one 's' and two 'p' orbital's of almost equal energy to give three identical and degenerate hybrid orbital's is known as 'sp2' hybridization. The three sp2 hybrid orbital's are oriented in trigonal planar symmetry at angles of 120 ⁰ to each other. The sp2 hybrid orbital's have 33.3% 's' and 66.6% 'p' character.

29. For example:

30. I n s p 3 h y bridi z a ti o n , o n e ' s ' a n d ' p ' o r b i t al ' s o f a l m o s t e q u a l three e n e r gy i n t e r mix t o g i v e f o u r identical and degenerate hybrid orbital's. These four sp3 hybrid orbital's are oriented in tetrahedral symmetry with 109 ⁰ 28' angle with each other. The sp3 hybrid orbital's have 25% ‘s’ and 75% 'p' character.

31. For example:

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