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Chapter 2

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

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Chapter 2

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

  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 – 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. Bond Formation Ionic bonding involves the loss of an electron due to a large difference in electronegativity (DEN>2.0) Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (DEN<.5) Unequal sharing: polar bond (.4<DEN<2.1)

  5. 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+

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

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

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

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

  10. 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)

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

  12. 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

  13. Dipole Moment Values

  14. 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 2.5 2.1 2.1 2.5 3.5 acetic acid (ethanoic acid) 2.1

  15. 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?

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

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

  18. You Try It. • Draw a Lewis Structure of 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

  19. Formal Charges (Section 2.3) • formal charges – these charges don’t 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)

  20. 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+

  21. DMSO (dimethyl sulfoxide)

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

  23. 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

  24. 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 CH3NO2 H2C=N=N O3 [H2CNH2]+ (draw all resonance structures) (1 very likely, 1 less likely, 1 very unlikely)

  25. Resonance Structures formaldehyde

  26. Resonance Structures Some compounds are not adequately represented by a single Lewis structure as we saw in the previous example. When two or more structures are possible, the molecule will show characteristics of each structure.

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

  28. Resonance Structures The “real” structure is said to have its electrons delocalized and is represented by a dotted bond

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

  30. Resonance Structures 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

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

  32. Resonance Structures When drawing resonance structures, follow these rules: Different resonance forms of a substance don’t have to be equivalent

  33. Resonance Structures 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

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

  35. Radicals • radical - (free radical) a neutral substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot (·) leaving it with an odd number of electrons. • Radicals are highly reactive! (octet rule) • Radicals can form from stable molecules and can also react with each other.

  36. Radical Resonance • Resonance forms for radicals will depend upon three-atom groupings that contain a multiple bond next to a p-orbital.

  37. Pentadienyl Radical

  38. Pentadienyl Radical

  39. Pentadienyl Radical

  40. You try it. • Show all the resonance forms for the straight chained C7H9. radical (in line angle). .

  41. Organic Chemistry Chapter 2 Polar Covalent Bonds; Acids and Bases Part II

  42. Define and describe acids and bases based on the Brønsted-Lowry and Lewis definitions Use the curved-arrow formalism to show movement of electrons between Lewis acids and bases Determine conjugate acid-base pairs Predict strength of acids and bases based on size, electronegativity, resonance stabilization, hybridization, and induction Predict reactions using pKa values Section 2.7-2.11 Objectives

  43. Why Study Acids/Bases? At a deeper level, acid/base strength allows us to predict reactivity • Compounds tend to react in such a way that they become more stable (in the long run) • Compounds considered “strong” are called that (technically) because they dissociate completely, but (practically) also because they tend to react quickly. (This occurs because of LOW stability.) • Compounds considered “weak” tend not to react quickly or completely because they are stable (“happy” where they are )

  44. Stability Everything wants to be at the lowest possible energy. (most stable)

  45. Acids & Bases Definitions of acids and bases: Brønsted-Lowry definition Acids donate protons (H+) (Proton donor) Hint for recognizing the acid – look for Hs! Bases accept protons (Proton acceptor) Lewis definition Acids accept electrons (electrophile) Bases donate electrons (nucleophile) Hint for recognizing the base – look for electrons! Either a lone pair or pi bonded electrons

  46. Morphine

  47. Acid Reactions • So, if something loses a hydrogen, it has acted as an acid. It then has the capability to accept a proton. Therefore, what is it called at this point? A BASE! • Acids will donate a proton to become a conjugate base. • Bases will accept a proton to become a conjugate acid.

  48. General Acid-Base Reaction • When writing reactions, we show the movement (called an “attack”) of electrons with an arrow. • Full headed arrow – both electrons • Half headed arrow – one electron

  49. Acid-Base Reactions

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