Valence Bond Theory (VBT).pptx-365377368765

pptx

1. Prepared By : 1-Dler Ismail Rasul 2-Saz Zirak Hamad 3-Nawa Masud Rasul 4-Nihad Khalil Khalid

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

3. Valence bond theory came into account when G.N Lewis proposed the geometrical structures of compounds in 1916. He explained how the valence electrons are arranged among the atoms in a molecule. Using Lewis theory, people started drawing Lewis structures that allowed them to predict many properties of molecules such as molecular shape and polarity. Later, in 1927 Hitler and London attempted to explain the bonding properties by applying quantum mechanics. They took the hydrogen atoms as an example to explain chemical bond formation using Schrodinger's wave equation. History

4. 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.

5. 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.

6. Valence bond theory is used to explain the formation of covalent bonds between two atoms They explained the overlap of atomic orbitals This theory explained the structure of molecules based on the hybridisation phenomenon It gave the idea that maximum intersection gives rise to the development of strong possible bonds Applications of Valence

7. 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 .

8. 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.

9. 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

10. SIGMA BOND σ s-s bond :

11. SIGMA BOND σ p-p bond:

12. SIGMA BOND σs-p bond:

13. 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.

14. 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 .

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

16. 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.

17. 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.

18. 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.

19. sp sp2 sp3

20. 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.

21. For example:

22. 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.

23. For example:

24. 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.

25. For example: