1. © 2006 Brooks/Cole - Thomson 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. © 2006 Brooks/Cole - Thomson 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. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule 1.For ions we must adjust the number of electrons available, A. a.Add one e - to A for each negative charge. b.Subtract one e - from A for each positive charge. 2.The central atom in a molecule or polyatomic ion is determined by: a.The atom that requires the largest number of electrons to complete its octet goes in the center. b.For two atoms in the same periodic group, the less electronegative element goes in the center. 3.Select a reasonable skeleton a.The least electronegative is the central atom b.Carbon makes 2,3, or 4 bonds c.Nitrogen makes 1(rarely), 2,3, or 4 bonds d.Oxygen makes 1, 2(usually), or 3 bonds e.Oxygen bonds to itself only as O 2 or O 3, peroxides, or superoxides f.Ternary acids (those containing 3 elements) hydrogen bonds to the oxygen, not the central atom, except phosphates g.For ions or molecules with more than one central atom the most symmetrical skeleton is used 4.Calculate N, S, and A

4. © 2006 Brooks/Cole - Thomson 1.Count the number of electrons brought to the party (# of element times group number) 2.For ions we must adjust the number of electrons available. a.Add one e - to A for each negative charge. b.Subtract one e - from A for each positive charge. 3.Select a reasonable skeleton a.The least electronegative is the central atom b.See prior periodic table for number of electrons involved in bonding a.Group I 2 electrons or 1 bond b.Group II 4 electrons or up to 2 bonds c.Group III Al and B, 6 or 8 electrons up to 3 or 4 bonds d.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). e.All others must have at least 8 electrons (up to 4 bonds), but may have more. 4.The central atom in a molecule or polyatomic ion is determined by: a.For ions or molecules with more than one central atom the most symmetrical skeleton is used b.The atom that requires the largest number of electrons to complete its octet goes in the center. c.For two atoms in the same periodic group, the less electronegative element goes in the center. 5.Calculate Formal charges, adjust bonds for lowest numbers (zero preferred) and allow for resonance structures

5. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule Write Lewis dot and dash formulas for hydrogen cyanide, HCN.

6. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule Write Lewis dot and dash formulas for the sulfite ion, SO 3 2-.

7. © 2006 Brooks/Cole - Thomson 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. © 2006 Brooks/Cole - Thomson Resonance Write Lewis dot and dash formulas for sulfur trioxide, SO 3.

9. © 2006 Brooks/Cole - Thomson Resonance There are three possible structures for SO 3 2-. –The double bond can be placed in one of three places. oWhen 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. oDouble-headed arrows are used to indicate resonance formulas.

10. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule Write dot and dash formulas for BBr 3.

11. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule Write dot and dash formulas for AsF 5.

12. © 2006 Brooks/Cole - Thomson 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. © 2006 Brooks/Cole - Thomson The same basic approach will be used in every example of molecular structure prediction:

14. © 2006 Brooks/Cole - Thomson

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 linear molecule nonpolar A B A angular molecule polar Polar Molecules must meet two requirements: 1.One polar bond or one lone pair of electrons on central atom. 2.Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel.

16. © 2006 Brooks/Cole - Thomson 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: 1Lone pair to lone pair is the strongest repulsion. 2Lone pair to bonding pair is intermediate repulsion. 3Bonding 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 109.5 o.

17. © 2006 Brooks/Cole - Thomson VSEPR Theory Frequently, we will describe two geometries for each molecule. 1.Electronic geometry 1.Electronic geometry is determined by the locations of regions of high electron density around the central atom(s). 2.Molecular geometry 2.Molecular geometry determined 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. © 2006 Brooks/Cole - Thomson 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. © 2006 Brooks/Cole - Thomson VSEPR Theory Five regions of high electron density around the central atom. Six regions of high electron density around the central atom.

20. © 2006 Brooks/Cole - Thomson VSEPR Theory An example of a molecule that has different electronic and molecular geometries is water - H 2 O. 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 - CH 4. Electronic and molecular geometries are tetrahedral.

21. © 2006 Brooks/Cole - Thomson Valence Bond (VB) Theory Regions of High Electron Density Electronic GeometryHybridization 2Linearsp 3Trigonal planarsp 2 4Tetrahedralsp 3 5Trigonal bipyramidalsp 3 d 6Octahedralsp 3 d 2

22. © 2006 Brooks/Cole - Thomson 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: AB 3 U designates that there are 3 bonding pairs and 1 lone pair around the central atom.

23. © 2006 Brooks/Cole - Thomson Linear Electronic Geometry:AB 2 Species (No Lone Pairs of Electrons on A) 1s2s2p Be   1s sp hybrid 2p    

24. © 2006 Brooks/Cole - Thomson Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) 1s 2s 2p B  1s sp 2 hybrid  

25. © 2006 Brooks/Cole - Thomson Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 2s 2p C [He] 

26. © 2006 Brooks/Cole - Thomson Tetrahedral Electronic Geometry: AB 4 Species Tetrahedral Electronic Geometry: AB 3 U Species Valence Bond Theory (Hybridization) 2s 2p N [He]  four sp 3 hybrids  2s 2p C [He]  four sp 3 hybrids  2s 2p O [He]  four sp 3 hybrids  Tetrahedral Electronic Geometry: AB 2 U 2 Species

27. © 2006 Brooks/Cole - Thomson Tetrahedral Electronic Geometry: ABU 3 Species (Three Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) 2s 2p F [He]  four sp 3 hybrids 

28. © 2006 Brooks/Cole - Thomson Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U2, and AB 2 U 3 4s 4p4d As [Ar] 3d 10   five sp 3 d hybrids 4d 

29. © 2006 Brooks/Cole - Thomson Compounds Containing Double Bonds Valence Bond Theory (Hybridization) C atom has four electrons. Three electrons from each C atom are in sp 2 hybrids. One electron in each C atom remains in an unhybridized p orbital 2s 2p three sp 2 hybrids 2p C   An sp 2 hybridized C atom has this shape. Remember there will be one electron in each of the three lobes. Top view of an sp 2 hybrid

30. © 2006 Brooks/Cole - Thomson Compounds Containing Double Bonds The single 2p orbital is perpendicular to the trigonal planar sp 2 lobes. The fourth electron is in the p orbital. Side view of sp 2 hybrid with p orbital included.

31. © 2006 Brooks/Cole - Thomson Compounds Containing Double Bonds Two sp 2 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 sp 2 hybrids is designated as a  bond. The other portion of the double bond, resulting from the side-on overlap of the p orbitals, is designated as a  bond.

32. © 2006 Brooks/Cole - Thomson Compounds Containing Triple Bonds A  bond results from the head-on overlap of two sp hybrid orbitals. Note that a triple bond consists of one  and two p bonds. The unhybridized p orbitals form two  bonds.

33. © 2006 Brooks/Cole - Thomson Summary of Electronic & Molecular Geometries