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 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

3. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule Write Lewis dot and dash formulas for hydrogen cyanide, HCN. N = A = S (shared electrons) A-S (lone pair electrons)

4. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule Write Lewis dot and dash formulas for the sulfite ion, SO 3 2-. N = A = S (shared electrons) A-S (lone pairs)

5. © 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?

6. © 2006 Brooks/Cole - Thomson Resonance Write Lewis dot and dash formulas for sulfur trioxide, SO 3. N = A = S A-S

7. © 2006 Brooks/Cole - Thomson Resonance There are three possible structures for SO 3. –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.

8. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule There are some molecules that violate the octet rule. –For these molecules the N - A = S rule does not apply: 1.The covalent compounds of Be. 2.The covalent compounds of the IIIA Group. 3.Species which contain an odd number of electrons. 4.Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents. 5.Compounds of the d- and f-transition metals.

9. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).

10. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule Write dot and dash formulas for BBr 3. –This is an example of exception #2; The covalent compounds of the IIIA Group.

11. © 2006 Brooks/Cole - Thomson Writing Lewis Formulas: Limitations of the Octet Rule Write dot and dash formulas for AsF 5. –This is an example of rule 4; Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents.

12. © 2006 Brooks/Cole - Thomson Covalent Bonding Covalent bonds are formed when atoms share electrons. If the atoms share 2 electrons a single covalent bond is formed. If the atoms share 4 electrons a double covalent bond is formed. If the atoms share 6 electrons a triple covalent bond is formed. –The attraction between the electrons is electrostatic in nature The atoms have a lower potential energy when bound.

13. © 2006 Brooks/Cole - Thomson Formation of Covalent Bonds We can use Lewis dot formulas to show covalent bond formation. 1.H molecule formation representation. 2. HCl molecule formation Some examples of nonpolar covalent bonds. H 2 N 2

14. © 2006 Brooks/Cole - Thomson Polar and Nonpolar Covalent Bonds Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds. –Nonpolar covalent bonds have a symmetrical charge distribution. To be nonpolar the two atoms involved in the bond must be the same element to share equally.

15. © 2006 Brooks/Cole - Thomson Polar and Nonpolar Covalent Bonds Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds –Polar covalent bonds have an asymmetrical charge distribution To be a polar covalent bond the two atoms involved in the bond must have different electronegativities. Some examples of polar covalent bonds. HF 2.1 4.0

16. © 2006 Brooks/Cole - Thomson Polar and Nonpolar Covalent Bonds Compare HF to HI. 2.1 2.5

17. © 2006 Brooks/Cole - Thomson Polar and Nonpolar Covalent Bonds Polar molecules can be attracted by magnetic and electric fields.

18. © 2006 Brooks/Cole - Thomson Dipole Moments Molecules whose centers of positive and negative charge do not coincide, have an asymmetric charge distribution, and are polar. –These molecules have a dipole moment. The dipole moment has the symbol .  is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q. There are some nonpolar molecules that have polar bonds. There are two conditions that must be true for a molecule to be polar. 1.There must be at least one polar bond present or one lone pair of electrons. 2.The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

19. © 2006 Brooks/Cole - Thomson The Continuous Range of Bonding Types Covalent and ionic bonding represent two extremes. 1.In pure covalent bonds electrons are equally shared by the atoms. 2.In pure ionic bonds electrons are completely lost or gained by one of the atoms. Most compounds fall somewhere between these two extremes. All bonds have some ionic and some covalent character. –For example, HI is about 17% ionic The greater the electronegativity differences the more polar the bond.

20. © 2006 Brooks/Cole - Thomson Formal Charges ― The hypothetical charge on an atom in a covalently bonded molecule or ion; bonding electrons are counted as if equally shared. 1.FC = (Group #) – [(number of bonds) + number of unshared e - )] 2.For Lewis dot formulas an atom that has the same number of bonds as its group number has an FC of zero 3.The sum of formal charges is equal to zero for molecules and equal to the charge of the ion for ions. Ammomia – NH 3 Ammonium Ion – NH 4 + Lewis Dot FC =

21. © 2006 Brooks/Cole - Thomson Formal Charges ― The most likely formula for a molecule or ion is usually the one in which the formal charge on each atom is zero or as near zero as possible ―Negative formal charges are more likely to occur on the more electronegative elements ―Lewis dot formulas in which adjacent atoms have formula charges of the same sign are usually not accurate Cl=N-O

22. © 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

23. © 2006 Brooks/Cole - Thomson The same basic approach will be used in every example of molecular structure prediction:

24. © 2006 Brooks/Cole - Thomson

25. 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.

26. © 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.

27. © 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.

28. © 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.

29. © 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.

30. © 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.

31. © 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

32. © 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.

33. © 2006 Brooks/Cole - Thomson Linear Electronic Geometry:AB 2 Species (No Lone Pairs of Electrons on A) 1s2s2p Be   1s sp hybrid 2p    

34. © 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  

35. © 2006 Brooks/Cole - Thomson Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 2s 2p C [He] 

36. © 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

37. © 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 

38. © 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 

39. © 2006 Brooks/Cole - Thomson Compounds Containing Double Bonds Ethene or ethylene, C 2 H 4, is the simplest organic compound containing a double bond. Lewis dot formula N = 2(8) + 4(2) = 24 A = 2(4) + 4(1) = 12 S = 12 Compound must have a double bond to obey octet rule.

40. © 2006 Brooks/Cole - Thomson Compounds Containing Double Bonds VSEPR Theory suggests that the C atoms are at center of trigonal planes. CC H HH H

41. © 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

42. © 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.

43. © 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.

44. © 2006 Brooks/Cole - Thomson Compounds Containing Triple Bonds Ethyne or acetylene, C 2 H 2, is the simplest triple bond containing organic compound. Lewis Dot Formula N = 2(8) + 2(2) = 20 A = 2(4) + 2(1) =10 S = 10 Compound must have a triple bond to obey octet rule.

45. © 2006 Brooks/Cole - Thomson Compounds Containing Triple Bonds Lewis Dot Formula VSEPR Theory suggests regions of high electron density are 180 o apart. HCCH

46. © 2006 Brooks/Cole - Thomson Compounds Containing Triple Bonds Valence Bond Theory (Hybridization) Carbon has 4 electrons. Two of the electrons are in sp hybrids. Two electrons remain in unhybridized p orbitals. 2s2p two sp hybrids 2p C [He] 

47. © 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.

48. © 2006 Brooks/Cole - Thomson Summary of Electronic & Molecular Geometries