Chapter 2

1. Organic Chemistry Chapter 2 Polar Covalent Bonds; Acids and Bases

2. Chapter Objectives • Take an in-depth look at polarity of molecules • Use formal charges to designate the distribution of electrons • Represent molecules with resonance structures by ‘pushing’ electrons • Examine the acid-base behavior of molecules • Predict acid-base reactions from pKa values

3. Electronegativity • electronegativity – (EN) a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.

4. Electronegativity Table • ionic bonding - 2.0 < DEN • polar covalent bonding - .5 ≤ DEN ≤ 2.0 • non-polar covalent bonding -DEN < .5

5. Bond Formation Ionic bonding involves the loss of an electron due to a large difference in electronegativity (2.0 < DEN ) Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (DEN < .5) Unequal sharing: polar bond (.5 ≤ DEN ≤ 2.0)

6. Polarity If one side is more electronegative, it tends to have a partial negative charge (δ-) [electron-rich] The other side tends to have a partial positive charge (δ+) [electron-poor] The δ- and δ+ difference along a bond is called a dipole moment δ- δ+

7. Ionic Character

8. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

10. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

11. Electrostatic Potential Maps Red – electron rich (δ-) Blue – electron poor (δ+)

12. Electrostatic Potential Maps You describe it… What molecule do you think it is? Take a guess… Red – electron rich (δ-) Blue – electron poor (δ+)

13. Inductive Effect • inductive effect – the shifting of electrons in a σ (sigma) bond in response to the electronegativity of nearby atoms. 3.5 2.1 Let’s discuss what it means to be an acid. 2.5 2.1 2.1 2.5 3.5 acetic acid (ethanoic acid) (this is a carboxylic acid due to the –OH off the carbonyl group) 2.1

14. Inductive Effect • Why would HCN allow the H+ to be released (proton donor – acid), thus categorizing HCN as an acid, when CH4 is not usually categorized as an acid?

15. Dipole Moment Calculations Section 2.2 • dipole moment (μ – Greek mu) – the magnitude of the charge (Q) at either end of the molecular dipole times the distance (r) between the charges • measured in debyes (D) • μ = Q x r • Just be familiar with magnitude of values and that the D following the value is the unit.

16. Dipole Moments overall dipole moment = 1.70 D 3.5 2.1 2.5 2.5 2.1 2.1 3.5 2.1 acetic acid (ethanoic acid)

17. Dipole Moments overall dipole moment = 1.70 D acetic acid (ethanoic acid)

18. Water and Ammonia

19. Dipole Moment Values

20. Non-Polar Molecules

21. You Try It. • Draw the complete Lewis Structure for the alcohol, methanol (methyl alcohol). Show the general direction of its dipole moment. (μ =1.70)

22. You Try It. • Determine if the following molecules are polar or non polar. Show any bond dipoles. (a) (b) (c)

23. You Try It. • Draw Lewis Structures for each of the following molecules and predict whether each has a dipole moment. If you expect a dipole moment, draw it in the correct direction. (a) C2HF (b) CCl4 (c) CH3CHO

24. Formal Charges Section 2.3 • formal charges – these charges do not imply the presence of actual ionic charges …instead they give insight into the distribution of electrons • calculating the formal charges of each atom in a molecule will help you determine the best, most favorable structure (lowest energy)

25. General Rules of Stability Lewis structures that approximate the actual molecule most closely are those that have: • maximum number of covalent bonds • minimum separation of unlike charges • formal charges of zero are ideal • placement of any negative charges on the most electronegative atom (or any positive charge on the most electropositive atom) • Ex. Oxygen would rather 1- then 1+

26. Formal Charges formal charge is calculated in the following manner: If it violates HONC 1234, then it will have a formal charge on it.

27. DMSO (dimethyl sulfoxide)

28. DMSO (dimethyl sulfoxide)

29. Formal Charges Summary

30. Nitromethane Determine any formal charges on nitromethane, CH3NO2

31. Nitromethane Determine any formal charges on nitromethane, CH3NO2

32. Formal Charges Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there is one less electron CH4 H3O+ NH3BH3

33. Formal Charges Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there is one less electron H2C=N=N O3 [H2CNH2]+ (This has resonance structures.) (1 very likely, 1 less likely, 1 very unlikely)

34. Resonance Structures Some molecules cannot be represented by a single structure. In these cases we draw structures that contribute to the final structure but which differ in the position of the  bond(s) or lone pair(s). Such a structure is delocalized and is represented by resonance forms The resonance forms are connected by a double-headed arrow. Section 2.4

35. Resonance Structures When two or more structures are possible, the molecule will show characteristics of each structure. Experiments show that these two structures are equivalent…both C-O bonds are same length and strength. Since both structures are equally likely, the real structures is most likely a perfect blend of each of these. This is not always the case.

36. Resonance Structures Draw resonance structures for NO3- The “real” structure is a resonance hybrid Each oxygen has a partial negative charge

37. Resonance Structures The “real” structure is said to have its electrons delocalized and is represented by a dotted bond Remember, different resonance structures are not always equivalent. Why would the one on the left be least influential? Carbon wants a complete octet!!! Carbon is KING!

38. Resonance Structures In some cases, one resonance form is more stable than another (one accommodates formal charges better)

39. Rules for Resonance Forms When drawing resonance structures, follow these rules: Individual resonance forms are imaginary, not real Resonance forms differ ONLY in the placement of their pi or non-bonding electrons Different resonance forms of a substance don’t have to be equivalent All resonance forms must be valid Lewis structures and obey normal rules for valency The resonance hybrid is more stable than any individual resonance form Section 2.5

40. When drawing resonance structures, follow these rules: Resonance forms differ ONLY in the placement of their pi or non-bonding electrons Rules for Resonance Forms

41. Rules for Resonance Forms

42. When drawing resonance structures, follow these rules: Different resonance forms of a substance don’t have to be equivalent Rules for Resonance Forms

43. Rules for Resonance Forms

44. When two resonance forms are nonequivalent, the actual structure of the resonance hybrid is closer to the more stable form than to the less stable form. Rules for Resonance Forms

45. When drawing resonance structures, follow these rules: All resonance forms must be valid Lewis structures and obey normal rules for valency Rules for Resonance Forms

46. Rules for Resonance Forms

47. When drawing resonance structures, follow these rules: The resonance hybrid is more stable than any individual resonance form Rules for Resonance Forms

48. In general, the larger the number of resonance forms, the more stable (lower in energy…less reactive) a substance is because electrons are spread out over a larger part of the molecule and are closer to more nuclei. The more the negative charge is spread out, the better! Rules for Resonance Forms

49. General Trends + C • - C + N/O - N/O

50. Remember Curved arrows always represent the movement of electrons, not atoms. Electrons always move towards the more electronegative element or positive charges. Electron pairs can only move to adjacent positions. The Lewis structures that result must be valid and have the same net charges.