Basic Concept: Chemical Bonding

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1 Basic Concept: Chemical Bonding 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.

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

3 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.

4 Examples of covalent bonding

5 Sigma bond 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.

6 π bond 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.

7 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.

8 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.

9 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 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.

10 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. Three 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)

11 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.

12

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 Linear Trigonal Planar Tetrahedral Trigonal bipyramid Octahedral 90 0

14 Some examples using VSEPR Theory SnCl 2 Lewis model: Shape : bent lp-bp repulsions cause the Cl-Sn-Cl bond angle close to less than (approx 95 0 )

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

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 0

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.

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