1. Lecture 21 © slg CHM 151 Lewis Structures: Molecules Ions Oxy Acids Resonance Structures TOPICS:

2. Lewis Structures: Compounds and Polyatomic Ions GUIDELINES Decide on arrangement of atoms. For most species, the element written first in the molecule or ion is the central atom and the remainder of the atoms are grouped around it. Hydrogen is a problem in “oxo acids” where it is written first in the formula. Ignore H, start with the next atom in formula and place the H or H’s on the O or O’s. First step:

4. Second Step Add up all available valence electrons. If species is cation, subtract positive charge from total. If species is anion, add negative charge to total. Divide total by two to determine available number of electron pairs Third Step Place a pair of electrons between each pair of bonded atoms to represent a single bond (use a “dash”!)

5. Step 1 Step 2 Step 3

6. Fourth Step Place leftover electron pairs around “terminal” atoms to achieve their octet (except H). Do central atom last. Fifth Step Examine central atom to determine if a double or triple bond is required to achieve the central atom’s octet. Do so using unshared pairs, IF central atom is: C, N, P, O, S

7. N, octet H, duet Step 4: No Step 5 needed

8. Step 1 Step 2 Step 3

9. Step 4

12. Now let’s try some: GROUP WORK Use 5 steps: Arrange; adds up e’s; draw bonds; assign unshared pairs double bonds if needed to draw Lewis structures for following species: NBr 3 CH 2 Cl 2

13. key

15. Now let’s examine situations requiring the double bond:

16. No octet

18. Either one

20. Be sure to include charge on finished product

21. GROUP WORK Use 5 steps: Arrange; adds up e’s; draw bonds; assign unshared pairs; double bonds if needed to draw Lewis structures for following species: H 3 PO 4 NO 2 1+ ClO 4 1-

22. Key:

23. key

25. Let’s explore the relationship between various “oxo” acids (H, Non metal element, O) and the charge and formula of their anion relative. Recall that acids, by definition, ionize in water to lose one or more H’s as H +. The anion left behind is named according to the name of its “parent” acid. In an acid/base reaction, as we met last unit, acids(H + ) react with bases (OH - ) to form water, leaving behind the anion of the acid and the cation of the base to form a salt.

26. Recall that acids “ionize” in water, or react with a base to form water, in either case leaving behind some “anion”: H- “Anion” + NaOH H 2 O + Na + An - Acid: HCl, HNO 3 H 2 SO 4 etc... Cl -, NO 3 -, SO 4 2- etc... H- “Anion” H + + Anion - H2OH2O

30. We turn next to the structure of the “oxo acids” of Cl, which are identical to those for Br and I. HClO Hypochlorous Acid HClO 2 Chlorous Acid HClO 3 Chloric Acid HClO 4 Perchloric Acid Recall that Cl, Br, I and also F form acids with H, no O’s included: HCl Hydrochloric Acid HF Hydrofluoric Acid

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

36. 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.

37. 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...

38. 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.

39. 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.

40. 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:

41. 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: THE HYBRID STRUCTURE OF OZONE

42. 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.

43. 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.

44. 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.