1. Lecture 22 © slg CHM 151 RESONANCE OCTET VIOLATORS FORMAL CHARGES MOLECULAR SHAPES TOPICS:

2. No OCTET EQUIVALENT RESONANCE THEORY: WHERE TO PLACE THE DOUBLE BOND...

5. In all three cases, O 3, NO 3 -, CO 3 2-, when forming a double bond from a “terminal oxygen” one has a choice of moving e’s from several different O’s to makeup the “central atom’s” octet. Examination of experimental evidence (x ray) shed an interesting light on this topic: When two atoms are bonded together, the distance between their nuclei, their “bond length,” depends on whether the bonds between the two are single, double, or triple.

6. TYPICAL BOND LENGTHS Note that triple bonds are shorter than double and also double shorter than single, as well as being characteristic between any two given atoms. X ray evidence of bond lengths in ozone, nitrate and carbonate ions should therefore prove interesting...

7. 132 pm 121 pm Predicted, “usual” bond lengths: Instead of the predicted bond lengths observed in other compounds, both bonds in x ray showed identical lengths of 127.8 pm, close to an average of 1 1/2 bonds to each O.

8. Linus Pauling proposed the “theory of resonance” to describe this situation: When two or more equivalent Lewis structures can be drawn for a species, differing only in the position of electron pairs, then none are correct: The real structure is a hybrid of all structures drawn.

9. The Lewis structures drawn are called “contributing” or “resonance structures” needed to describe the makeup of the hybrid, which resembles all but is none of the above. A special double headed arrow is drawn between the contributing structures to indicate their hypothetical nature:

10. The hybrid structure, with two equivalent bonds to the central atom, are said to have a bond order of “1.5” or an average of 1 and 1/2 bonds between each O: (Bond order: # bonds/# atoms bonded to central atom) THE HYBRID STRUCTURE OF OZONE

11. Bond Order describes the number of bonds between two atoms in a molecule. Normally, the bond number is 1 (a single bond) or 2 (a double bond) or 3 (a triple bond.) When hybrid structures and resonance situations exist, one must average the number of bonds between all atoms affected, and fractional values arise. In the case of the nitrate and the carbonate ions, the number of bonds to the central atom is averaged out over 3 atoms, and 4 bonds/3 atoms= 1.33 bond order. In both cases, x ray data confirms this theory.

12. The carbonate ion has three equivalent C-O bonds, of a length typical of 1 and 1/3 bond, for a 1.33 bond order.

13. The nitrate ion also has three equivalent N-O bonds, of a length typical of 1 and 1/3 bond, for a 1.33 bond order.

14. OCTET VIOLATORS Another aspect of drawing correct Lewis structures involves the handling of compounds that do not have an octet around the central atom. Three situations exist: 1. More than 4 e - pairs around central atom 2. Less than 4 e - pairs around central atom 3. Molecules with odd number of electrons In all cases we will handle, the irregularity occurs at the central atom; all “terminal atoms” will have normal octet.

15. EXAMPLE: Note: Only the central atom, P, is an “octet violator”

16. Note again: Only the central atom exceeds the octet rule.

17. Case #2: Less than 4 e - pairs around central atom This category specifically applies to the metalloid Boron, but also to metals that form salts that are more covalent in nature than ionic: Beryllium, Aluminum, for example. These elements use their valence e’s to form compounds but do not form an octet in the process and do not accept double bonds to compensate. These “octet deficient” species will react with other atoms however to form polyatomic ions or compounds which relieve the deficiency.

18. EXAMPLE: No Octet, octet rule violator

19. While B will not form a double bond to F to achieve an octet (F’s “don’t do” double bonds), it will accept electron pairs readily from other sources to do so: When one atom donates two electrons for a pair of atoms to share, the bond is called a “coordinate covalent bond” and introduces “charge buildup” in the species formed.

20. To keep track of this kind of charge within a molecule or polyatomic ion, the concept of “FORMAL CHARGE” is introduced. Formal charges look within a molecule or polyatomic ion and determine how the charges are distributed by considering for each atom: the number of valence e’s it started with the number of bonds formed the number of unshared electrons leftover

21. For each atom in species: formal charge = # valence e’s - (#bonds + #unshared e’s) FORMAL CHARGE: The “formal charge” system requires a Lewis Dot Structure and assigns an individual “formal” charge to each atom in the species. Formal charge is an alternate “bookkeeping method” for tracking electron distribution to the “oxidation number” system we met previously.

22. Now let’s return to the compound formed between ammonia and boron trifluoride, and determine formal charges: Formal charges

24. OXIDATION NUMBERS: “Ox #’s” are assigned or calculated based on known fixed positive and negative charges, and can be determined by examination of the formula for the species. Oxidation numbers are useful to identify how charges change in a redox (oxidation-reduction) reaction.

25. To see how both work, let’s look at chloric acid, HClO 3, and see how its charge distribution would be described using both the oxidation number and the formal charge systems.

26. Sum of all charges in compound = 0 Known ox #’s per atom

27. Lewis Structure

28. formal charge = # valence e’s - (#bonds + #unshared e’s) Like ox #’s, the sum of all formal charges in a compound must equal 0.

29. Oxidation numbers Formal Charges

30. Finally, both oxidation numbers and formal charges must add up the same way: For compounds, which are always electrically neutral, the sum of all oxidation numbers or the sum of all formal charges must equal zero. For polyatomic ions, which always have a specific charge, the sum of all oxidation numbers or the sum of all formal charges must equal the charge on the ion.

32. GROUP WORK For each of the species pictured on the next slide, please determine: The oxidation number for each element in the formula; The formal charge for each atom in the Lewis dot structure.

33. 1: Do Ox #’s from formula (same for both!) 2: Do formal charge from Lewis Structure for all atoms: formal charge = # valence e’s -( # bonds+ # unshared e’s)

35. Oxidation Numbers Formal Charge

37. Oxidation Numbers Formal Charge

38. Case #3: Molecules with odd number of electrons This situation occurs primarily with oxides of nitrogen, NO and NO 2 : as ever, give the terminal atoms their octet, and let the central atom, N, handle the lone electron. Species involving unpaired single electrons are generally not stable and react quickly to form octet observing compounds with all e’s paired. Single electron species are called “free radicals.”

40. Like O 2, this “correct Lewis Structure” is inadequate to describe its reactivity, better done by alternate “molecular orbital theory”.

41. Two tasks remain to describe the molecules and polyatomic ions for which we have drawn Lewis Structures: their three dimensional shapes the charge distribution within the species that arises from unequal sharing of bonded pairs which we will call “polarity” Both of these properties require starting with a correct Lewis structure, a skill we are now ready to handle!

42. Molecular and Polyatomic Ion Shapes Once a Lewis structure is drawn, the three - dimensional geometry of the species can easily be determined by utilizing the “valence shell electron pair repulsion theory” called “VSEPR”: “VSEPR” theory is based on the tendency of negatively charged regions to repel each other and align as far apart as possible, resulting in predictable shapes for any covalently bonded species.

43. To utilize “VSEPR”, the number of regions of electron density around the central atom in the species is counted. Count as “one region”: Single Bonds Unshared Pairs Multiple bonds between same two atoms

44. Examples of “four regions”: “three regions”: “two regions”:

45. Basic Shapes predicted by VSEPR: Two regions: Three Regions: Bond Angles Geometry

46. Four Regions: Five Regions: Six Regions: