JEE Chemistry Chemical Bonding and Molecular Structure Notes Set 06

Covalent Bond and Octet Rule

A covalent bond forms when two atoms share electrons; sharing 2, 4 or 6 electrons gives a single, double or triple bond. Atoms tend to reach eight electrons in the outer shell (two for hydrogen), like the nearest noble gas. Noble gases themselves have the stable \( ns^2np^6 \) arrangement.
Drawing a Lewis structure: (1) add the valence electrons of all atoms; (2) add one electron for each negative charge and remove one for each positive charge; (3) choose the central atom, usually the least electronegative or the one able to form most bonds (hydrogen and fluorine are never central); (4) join atoms by single bonds; (5) place the remaining electrons as lone pairs so that each atom gets an octet, using multiple bonds if needed.

Formal Charge

Formal charge = valence electrons of the free atom - non-bonding electrons - half the bonding electrons. In \( \ce{O3} \) the central O is +1, one terminal O is 0 and the other -1. Among alternative Lewis structures the one with the smallest formal charges, and with negative charge on the more electronegative atom, is preferred. Adjacent +1 and -1 charges represent a coordinate bond.

Exceptions to the Octet Rule

1. Incomplete octet: electron-deficient central atoms such as \( \ce{BeH2} \), \( \ce{BF3} \), \( \ce{BCl3} \) and \( \ce{AlCl3} \).
2. Odd-electron molecules: NO, \( \ce{NO2} \), \( \ce{ClO2} \).
3. Expanded octet: elements of period 3 and beyond can hold more than eight electrons, as in \( \ce{PCl5} \), \( \ce{SF6} \), \( \ce{IF7} \) and \( \ce{H2SO4} \).
The octet rule also fails to explain noble gas compounds such as \( \ce{XeF2} \), gives no information about shapes and says nothing about the energy of molecules.

Resonance and Bond Order

When one Lewis structure cannot describe a species, two or more structures that differ only in the position of electrons (same atomic positions, similar energy) are written. The real molecule is a resonance hybrid, more stable than any one structure; the stabilisation is the resonance energy and it grows with the number of covalent bonds and the number of equivalent structures.
Average bond order = total number of bonds between the two atoms in all structures divided by the number of structures. Examples: \( \ce{O3} \) 1.5; \( \ce{CO3^2-} \) 4/3; \( \ce{NO3-} \) 4/3; \( \ce{SO4^2-} \) 1.5; \( \ce{PO4^3-} \) 1.25; benzene C-C 1.5. The higher the bond order, the shorter and stronger the bond.

Valence Bond Theory: Sigma and Pi Bonds

Heitler and London treated the hydrogen molecule: as two atoms approach, attraction between each nucleus and the other electron outweighs repulsion until energy is a minimum at the bond length (74 pm for H2, bond energy about 436 kJ/mol).
A sigma bond results from head-on overlap along the internuclear axis (s-s, s-p, p-p along the axis); it is strong and allows free rotation. A pi bond results from sideways overlap of p (or d) orbitals; it is weaker, and has a nodal plane through the bond. A single bond is one sigma bond, a double bond is one sigma and one pi, and a triple bond is one sigma and two pi. With z as the bond axis, s-\( p_z \), \( p_z \)-\( p_z \) give sigma overlap, \( p_x \)-\( p_x \) and \( p_y \)-\( p_y \) give pi overlap, and s with \( p_x \) gives no net overlap.

Counting Lone Pairs and Bond Pairs

Central-atom lone pairs = (valence electrons - electrons used in bonds)/2. Examples: \( \ce{XeF2} \) 3, \( \ce{XeF4} \) 2, \( \ce{XeF6} \) 1, \( \ce{ClF3} \) 2, \( \ce{SF4} \) 1, \( \ce{BrF5} \) 1, \( \ce{I3-} \) 3 on the central iodine. The lone pairs decide the shape: see-saw \( \ce{SF4} \), T-shape \( \ce{ClF3} \), linear \( \ce{XeF2} \), square planar \( \ce{XeF4} \), square pyramid \( \ce{BrF5} \).