Chemical bonding

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1. Essential Reading: 1. P. W. Atkins, Elements of Physical Chemistry , 4th Ed., Oxford University Press, 2007. 2. F. A. Carey, R. M. Guuliano , Organic Chemistry , Mcgraw -Hill, 6th edition, 2006. 3. J.D. Lee, Concise Inorganic Chemistry, 5th edition, Blackwell Publishing, 2008. 4. Fundamentals of molecular spectroscopy, C. N. Banwell , Tata McGraw-Hill Education, 1994.   Supplementary Reading: 1. J. Singh, L.D.S. Yadav , Advanced Organic Chemistry , PragatiPrakashan , 2009. 2. J. E. Huheey , E. A. Keiter and R. L. Keiter , Inorganic Chemistry, Principles of structure and reactivity, Harper Collins, 1993. 3. Clayden , Greeves , Warren and Wothers , Organic Chemistry, Oxford, 2001. 4. B. R. Puri , L. R. Sharma, M. S. Pathania , Principles of physical Chemistry, ShobanLalNagin Chand & Co., 2001. Recommended Text Books

2. Ionic bond: Type of chemical bond  that involves the  electrostatic  attraction between oppositely charged  ions. These ions represent atoms that have lost one or more  electrons   ( cations ) and atoms that have gained one or more electrons (anions). Here Sodium  molecule is donating its 1 valence electron to the Chlorine molecule.  This creates a Sodium  cation  and a Chlorine anion.  Notice that the net charge of the compound is 0. Basic Concept: Chemical Bonding

3. Some examples of ionic bonds and ionic compounds: NaBr - sodium bromide NaF - sodium fluoride KI - potassium iodide KCl - potassium chloride CaCl 2 - calcium chloride KBr - potassium bromide Ionic bonding in sodium chloride Formation of ionic bond in lithium fluoride

4. Covalent bond : A chemical bond  that involves the sharing of  electron pairs between atoms. The stable balance of attractive and repulsive forces between atoms when they share electrons is known as covalent bonding. Here Phosphorous molecule is sharing its 3 unpaired electrons with 3 Chlorine atoms. In the end product, all four of these molecules have 8 valence electrons and satisfy the octet rule.

5. Examples of covalent bonding

6. In chemistry, sigma bonds ( σ bonds ) are the strongest type of covalent chemical bond. They are formed by head-on overlapping between atomic orbitals. Sigma bonding is most clearly defined for diatomic molecules. S igma bond

7. Pi bonds (π bonds) are covalent chemical bonds where two lobes of one involved atomic orbital overlap two lobes of the other involved atomic orbital. Each of these atomic orbitals is zero at a shared nodal plane, passing through the two bonded nuclei. π bond

8. Coordinate bond : A  dipolar bond, more commonly known as a dative covalent bond  or coordinate bond is a kind of 2-center, 2-electron covalent bond in which the two electrons derive from the same atom.

9. Metallic Bond : Metallic bonding constitutes the electrostatic attractive forces between the delocalized electrons, called conduction electrons, gathered in an electron cloud and the positively charged metal ions .

10. Valence Shell Electron Pair Repulsion (VSEPR) Theory Valence shell electron pair repulsion (VSEPR) theory is a model in chemistry, which is used for predicting the shapes of individual molecules. The theory was suggested by Sidgwick and Powell in 1940 and was developed by Gillespie and Nyholm in 1957. It is also called the Gillespie- Nyholm Theory after the two main developers. VSEPR theory is based on the idea that the geometry of a molecule or polyatomic ion is determined primarily by repulsion among the pairs of electrons associated with a central atom. The pairs of electrons may be bonding or nonbonding (also called lone pairs ). Only valence electrons of the central atom influence the molecular shape in a meaningful way.

11. VSEPR theory may be summarized as: The shape of the molecule is determined by repulsions between all of the electron pairs present in the valence shell. A lone pair of electrons takes up more space around the central atom than a bond pair. T hree types of repulsion take place between the electrons of a molecule: The lone pair-lone pair repulsion ( lp-lp ) The lone pair-bonding pair repulsion ( lp-bp ) The bonding pair-bonding pair repulsion . ( bp-bp )

12. The best spatial arrangement of the bonding pairs of electrons in the valence orbitals is one in which the repulsions are minimized . lp-lp > lp-bp > bp-bp The magnitude of the repulsions between bonding pairs of electrons depends on the electronegativity difference between central atom and other atoms . Double bonds cause more repulsion than single bonds, and triple bonds cause more repulsion than a double bond.

13. Predicted molecular shapes from Sidgwick - Powell Theory: No. of electron pairs in outer shell Arrangement of electron pairs Electron-pair geometry Bond angles 2 3 4 5 6 Linear Trigonal Planar Tetrahedral Trigonal bipyramid Octahedral 180 120 109.5 90 120 90

14. Some examples using VSEPR Theory SnCl 2 Lewis model: Shape : bent l p-bp repulsions cause the Cl-Sn-Cl bond angle close to less than 120 ( approx 95 )

15. NH 3 Lewis model: Shape : Trigonal Pyramid lp-bp repulsions cause the H-N-H angles to close to less than 109.5 o ( 107.3 o ).

16. H 2 O Lewis model: Shape : Bent lp-bp repulsions cause the H-O-H angle to be lesser than 109.5 (104.5 )

17. ClF 3 Lewis model: Shape : T shape Lone pairs occupy equatorial positions of trigonal bipyramid lp-bp repulsions cause F-C-F angle to be lesser than 90

18. Limitations of VSEPR Theory It fails to predict the shapes of isoelectronic species [ CH 4 and NH 4 + ] and transition metal compounds. The model does not take relative size of substituents. Atomic orbitals overlap cannot be explained by VSEPR theory. The theory makes no predictions about the lengths of the bonds, which is another aspect of the shape of a molecule.

19. Bent’s Rule In a molecule, smaller bond angles are formed between electronegative ligands since the central atom, to which the ligands are attached, tends to direct bonding hybrid orbitals of greater p character towards its more electronegative substituents. Structure of water illustrating how the bond angle deviates from the tetrahedral angle of 109.5°.

20. The carbon atoms are directing sp 3 , sp 2 , and sp orbitals towards the hydrogen substituents. This simple system demonstrates that hybridised atomic orbitals with higher p character will have a smaller angle between them.

21. Limitations of Valence Bond Theory: It involves a number of assumptions . (ii) It does not give quantitative interpretation of magnetic data. (iii) It does not explain the color exhibited by coordination compounds . (iv) It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds . ( v) It does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes . ( vi) It does not distinguish between weak and strong ligands.

22. Molecular Orbital Theory Molecular orbitals result from the combination of atomic orbitals. Since orbitals are wave functions, they can combine either constructively (forming a bonding molecular orbital), or destructively (forming an antibonding molecular orbital ). Consider the H 2 molecule, for example. One of the molecular orbitals in this molecule is constructed by adding the mathematical functions for the two 1 s atomic orbitals that come together to form this molecule. Another orbital is formed by subtracting one of these functions from the other

23. Bonding Molecular Orbital Theory The bonding orbital results in increased electron density between the two nuclei, and is of lower energy than the two separate atomic orbitals.

24. Antibonding Molecular Orbital Theory The antibonding orbital results in a node between the two nuclei, and is of greater energy than the two separate atomic orbitals.

25. Overlap of s & p Orbitals + - + + + + - - - - + + + - - -

26. Sigma bonding orbitals From s orbitals on separate atoms + + s orbital s orbital + + + + Sigma bonding molecular orbital

27. Molecular Orbitals of the Second Energy Level If we arbitrarily define the Z axis of the coordinate system for the O 2 molecule as the axis along which the bond forms, the 2 p z orbitals on the adjacent atoms will meet head-on to form a 2 p bonding and a 2 p * antibonding molecular orbital

28. Sigma bonding orbitals From p orbitals on separate atoms p orbital p orbital Sigma bonding molecular orbital      

29. The 2 p x orbitals on one atom interact with the 2 p x orbitals on the other to form molecular orbitals that have a different shape. These molecular orbitals are called pi ( π ) orbitals because they look like p orbitals when viewed along the bond.

30. Pi bonding orbitals P orbitals on separate atoms         Pi bonding molecular orbital

31. s-p mixing Molecular Orbital Diagram

32. Molecular Orbital Diagram of H 2

33. Molecular Orbital Diagram of N₂

34. Molecular Orbital Diagram of O 2