1. Bonding II Lewis Structures and Covalent Bond Properties

2. Lewis Structures Introduced by G.N. Lewis (U. C. Berkeley) in 1916 Covalent bonding attributed to shared electron pairs Electron pairs are localized on or between atoms Shared pairs of electrons are “owned” by both atoms C,N,O and F always obey the “Octet Rule” Elements lighter than C have less than 8 e - Elements heavier than F may have more than 8 e - Review method for writing Lewis structures: –“simple” examples:CH 4, H 2 O, CO 2, etc.

3. Resonance Since Lewis model assumes that electrons are localized, a single Lewis structure cannot adequately represent a molecule with delocalized bonding. Resonance forms are drawn by assigning bonding and lone pairs in different ways; the geometric arrangement of atoms does NOT change! (e.g.: SO 2,NO 3 -, etc.) Molecular electronic structure is described as a “hybrid” (mixture) of the individual resonance forms. Equivalent structures contribute equally (same energy). Relative contributions of different structures are determined by assessing their relative energies.

4. Formal Charge Relative merit of different resonance forms may be assessed by assigning formal charge to each atom. Formal charge counting rule: divide bonding pairs of electrons between atoms (i.e., 100% covalent bonding is assumed) Formal charge does NOT reflect the actual charges borne by atoms in a structure! Favored (lowest energy) structure: –Minimizes formal charge separation within the molecule. –Places negative charge (if any) on most electronegative atom. –Places positive charge (if any) on least electronegative atom. Example: cyanate (OCN - ) vs. fulminate (CNO - )

5. Formal Oxidation Number The formal oxidation number represents the number of electrons gained or lost by an element when it forms a compound. Oxidation number counting rule: assign both electrons of each bonding pair to the more electronegative atom (100% ionic) (e.g. NH 3, NCl 3, NO 3 -, etc.) If two atoms of same element are bonded together, divide the bonding electrons (e.g. C 2 Cl 4, N 2 O, etc.) Formal oxidation number is a “book keeping tool”; it DOES NOT represent the “true” charge on any atom!

6. Hypervalence Atoms that necessarily exceed 8 e - in their valence shell to accommodate bonding of peripheral atoms (or lone pairs) are said to be “hypervalent”, or to have “expanded octets” –Examples: PCl 5, PF 6 -, SF 6, BrF 3, etc. –Contrary to popular belief, hypervalence does not necessarily require involvement of d-orbitals in bonding! –Occurrence of hypervalence may have more to do with the relative ease of fitting more atoms around a larger central atom. It may be possible to draw resonance forms with > 8 e - in the valence shell for some molecules (e.g. SO 4 2-, PO 4 3-, (CH 3 ) 2 SO, etc.) but these are NOT properly considered to be “hypervalent” molecules.

7. Molecular Geometry Molecular geometry may be assessed using the Valence Shell Electron Pair Repulsion (VSEPR) model. The description of the molecular geometry is based upon the arrangement of the bonded atoms. The geometry of most common molecules is derived from 5 basic shapes: linear (AB 2 ), trigonal planar (AB 3 ), tetrahedral (AB 4 ), trigonal bipyramidal (AB 5 ), and octahedral (AB 6 ). In this notation, “A” denotes the central atom; “B” groups may be bonded atoms or lone pairs.

8. Molecular Geometries

10. Covalent Bond Properties

11. Bond Length The equilibrium bond length corresponds to the minimum on the molecular potential energy curve The mutual attraction between the valence electrons of each atom and the nucleus of the other pulls the atoms together until the repulsion between the core electrons of the atoms begins to dominate. The sum of the covalent radii of two bonded atoms corresponds to the internuclear separation when the core shells of the two atoms are in contact.

12. Table of Covalent Radii

13. Bond Strength The bond dissociation enthalpy (BDE) is the standard reaction enthalpy for the process: A-B(g)  A(g) + B(g) The mean bond enthalpy is the average BDE taken over a series of A-B bonds in different molecules.

14. Table of Mean Bond Enthalpies

15. Estimation of  H rxn