Writing Lewis Formulas: The Octet Rule

1. Writing Lewis Formulas: The Octet Rule • The octet rule states that representative elements usually attain stable noble gas electron configurations in most of their compounds. • Lewis dot formulas are based on the octet rule. • We need to distinguish between bonding (or shared) electrons and nonbonding (or unshared or lone pairs) of electrons. • N - A = S rule • Simple mathematical relationship to help us write Lewis dot formulas. • N = number of electrons needed to achieve a noble gas configuration. • N usually has a value of 8 for representative elements. • N has a value of 2 for H atoms. • A = number of electrons available in valence shells of the atoms. • A is equal to the periodic group number for each element. • A is equal to 8 for the noble gases. • S = number of electrons shared in bonds. • A-S = number of electrons in unshared, lone, pairs.

2. Writing Lewis Dot Formulas N ever Have a Full Octet Always Have a Full Octet Sometimes Have a Full Octet Sometimes Exceed a Full Octet

3. Writing Lewis Formulas: The Octet Rule • For ions we must adjust the number of electrons available, A. • Add one e- to A for each negative charge. • Subtract one e- from A for each positive charge. • The central atom in a molecule or polyatomic ion is determined by: • The atom that requires the largest number of electrons to complete its octet goes in the center. • For two atoms in the same periodic group, the less electronegative element goes in the center. • Select a reasonable skeleton • The least electronegative is the central atom • Carbon makes 2,3, or 4 bonds • Nitrogen makes 1(rarely), 2,3, or 4 bonds • Oxygen makes 1, 2(usually), or 3 bonds • Oxygen bonds to itself only as O2 or O3, peroxides, or superoxides • Ternary acids (those containing 3 elements) hydrogen bonds to the oxygen, not the central atom, except phosphates • For ions or molecules with more than one central atom the most symmetrical skeleton is used • Calculate N, S, and A

4. Count the number of electrons brought to the party (# of element times group number) • For ions we must adjust the number of electrons available. • Add one e- to A for each negative charge. • Subtract one e- from A for each positive charge. • Select a reasonable skeleton • The least electronegative is the central atom • See prior periodic table for number of electrons involved in bonding • Group I 2 electrons or 1 bond • Group II 4 electrons or up to 2 bonds • Group III Al and B, 6 or 8 electrons up to 3 or 4 bonds • C,N,O,F must have 8 electrons (up to 4 bonds for C, 3 for N, 2 for O, and 1 bond for F). • All others must have at least 8 electrons (up to 4 bonds), but may have more. • The central atom in a molecule or polyatomic ion is determined by: • For ions or molecules with more than one central atom the most symmetrical skeleton is used • The atom that requires the largest number of electrons to complete its octet goes in the center. • For two atoms in the same periodic group, the less electronegative element goes in the center. • Calculate Formal charges, adjust bonds for lowest numbers (zero preferred) and allow for resonance structures

5. Writing Lewis Formulas:The Octet Rule • Write Lewis dot and dash formulas for hydrogen cyanide, HCN.

6. Writing Lewis Formulas:The Octet Rule • Write Lewis dot and dash formulas for the sulfite ion, SO32-.

7. Writing Lewis Formulas:The Octet Rule • What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom?

8. Resonance • Write Lewis dot and dash formulas for sulfur trioxide, SO3.

9. Resonance • There are three possible structures for SO32-. • The double bond can be placed in one of three places. • When two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure. • Double-headed arrows are used to indicate resonance formulas.

10. Writing Lewis Formulas:Limitations of the Octet Rule • Write dot and dash formulas for BBr3.

11. Writing Lewis Formulas:Limitations of the Octet Rule • Write dot and dash formulas for AsF5.

12. Stereochemistry • Stereochemistry is the study of the three dimensional shapes of molecules. • Valence Shell Electron Pair Repulsion Theory • Commonly designated as VSEPR • Principal originator • R. J. Gillespie in the 1950’s • Valence Bond Theory • Involves the use of hybridized atomic orbitals • Principal originator • L. Pauling in the 1930’s & 40’s

13. The same basic approach will be used in every example of molecular structure prediction:

15. Polar Molecules: The Influence of Molecular Geometry • Molecular geometry affects molecular polarity. • Due to the effect of the bond dipoles and how they either cancel or reinforce each other. A B A A B A angular molecule polar linear molecule nonpolar • Polar Molecules must meet two requirements: • One polar bond or one lone pair of electrons on central atom. • Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel.

16. VSEPR Theory • Regions of high electron density around the central atom are arranged as far apart as possible to minimize repulsions. • There are five basic molecular shapes based on the number of regions of high electron density around the central atom. • Lone pairs of electrons (unshared pairs) require more volume than shared pairs. • Consequently, there is an ordering of repulsions of electrons around central atom. • Criteria for the ordering of the repulsions: • Lone pair to lone pair is the strongest repulsion. • Lone pair to bonding pair is intermediate repulsion. • Bonding pair to bonding pair is weakest repulsion. • Mnemonic for repulsion strengths • lp/lp > lp/bp > bp/bp • Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.

17. VSEPR Theory • Frequently, we will describe two geometries for each molecule. • Electronic geometryis determined by the locations of regions of high electron density around the central atom(s). • Molecular geometrydetermined by the arrangement of atoms around the central atom(s). Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used.

18. VSEPR Theory • Two regions of high electron density around the central atom. • Three regions of high electron density around the central atom. • Four regions of high electron density around the central atom.

19. VSEPR Theory • Five regions of high electron density around the central atom. • Six regions of high electron density around the central atom.

20. VSEPR Theory • An example of a molecule that has different electronic and molecular geometries is water - H2O. • Electronic geometry is tetrahedral. • Molecular geometry is bent or angular. • An example of a molecule that has the same electronic and molecular geometries is methane - CH4. • Electronic and molecular geometries are tetrahedral.

21. Valence Bond (VB) Theory

22. Molecular Shapes and Bonding • In the next sections we will use the following terminology: A = central atom B = bonding pairs around central atom U = lone pairs around central atom • For example: AB3U designates that there are 3 bonding pairs and 1 lone pair around the central atom.

23. Linear Electronic Geometry:AB2 Species (No Lone Pairs of Electrons on A) 1s2s2p Be  1ssp hybrid2p    

24. Trigonal Planar Electronic Geometry: AB3 Species (No Lone Pairs of Electrons on A) 1s2s2p B  1ssp2 hybrid    

25. Tetrahedral Electronic Geometry: AB4 Species (No Lone Pairs of Electrons on A) 2s2p C [He]  .

26. Tetrahedral Electronic Geometry: AB4 Species Valence Bond Theory (Hybridization) 2s2p C [He] ­¯­­. four sp3 hybrids Þ ­ ­ ­ ­ . Tetrahedral Electronic Geometry: AB3U Species 2s2p N [He] ­¯­ ­ ­ four sp3 hybrids Þ ­¯­ ­ ­ Tetrahedral Electronic Geometry: AB2U2 Species four sp3 hybrids Þ ­¯­¯ ­ ­ 2s2p O [He] ­¯­¯ ­ ­

27. Tetrahedral Electronic Geometry: ABU3 Species (Three Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) 2s2p F [He] ­¯­¯ ­­ four sp3 hybrids Þ ­¯­¯ ­­

28. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 4s4p4d As [Ar] 3d10 ___ ___ ___ ___ ___ ß five sp3 d hybrids4d  ___ ___ ___ ___

29. Compounds Containing Double Bonds Valence Bond Theory (Hybridization) C atom has four electrons. Three electrons from each C atom are in sp2 hybrids. One electron in each C atom remains in an unhybridized p orbital 2s2pthree sp2 hybrids2p C Þ   • An sp2 hybridized C atom has this shape. • Remember there will be one electron in each of the three lobes. Top view of an sp2 hybrid

30. Compounds Containing Double Bonds • The single 2p orbital is perpendicular to the trigonal planar sp2 lobes. The fourth electron is in the p orbital. Side view of sp2 hybrid with p orbital included.

31. Compounds Containing Double Bonds • Two sp2 hybridized C atoms plus p orbitals in proper orientation to form C=C double bond. • The portion of the double bond formed from the head-on overlap of the sp2 hybrids is designated as a s bond. • The other portion of the double bond, resulting from the side-on overlap of the p orbitals, is designated as a p bond.

32. Compounds Containing Triple Bonds A  bond results from the head-on overlap of two sp hybrid orbitals. The unhybridized p orbitals form two p bonds. Note that a triple bond consists of one  and two p bonds.

33. Summary of Electronic & Molecular Geometries