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Chapter 1 Introduction and Review

Organic Chemistry , 5 th Edition L. G. Wade, Jr. Chapter 1 Introduction and Review. =>. Definitions. Old: “derived from living organisms” New: “chemistry of carbon compounds” From inorganic to organic, Wöhler, 1828. Atomic Structure. Protons (+ charges),

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Chapter 1 Introduction and Review

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  1. Organic Chemistry, 5th EditionL. G. Wade, Jr. Chapter 1Introduction and Review

  2. => Definitions • Old: “derived from living organisms” • New: “chemistry of carbon compounds” • From inorganic to organic, Wöhler, 1828 Chapter 1

  3. Atomic Structure • Protons (+ charges), • Neutrons (no electrical charge), and electrons (- charges) • Number of protons = number of electrons • Isotopes: Atoms of a given element that differ in the number of neutrons Chapter 1

  4. 2s orbital (spherical) => 2p orbital Atomic Orbitals Chapter 1

  5. => Electronic Configurations • Aufbau principle: Place electrons in lowest energy orbital first. • Hund’s rule: Equal energy orbitals are half-filled, then filled.    Chapter 1

  6. Table 1-1 => Chapter 1

  7. Chapter 1

  8. => Bond Formation • Ionic bonding: electrons are transferred. • Covalent bonding: electron pair is shared. Chapter 1

  9. Chapter 1

  10. Lewis Structures • Bonding electrons • Nonbonding electrons or lone pairs Satisfy the octet rule!=> Chapter 1

  11. Chapter 1

  12. Multiple Bonding => Chapter 1

  13. => Dipole Moment • Amount of electrical charge x bond length. • Charge separation shown by electrostatic potential map (EPM). • Red indicates a partially negative region and blue indicates a partially positive region. Chapter 1

  14. Chapter 1

  15. Electronegativity and Bond Polarity Greater EN means greater bond polarity => Pauling Scale Chapter 1

  16. Chapter 1

  17. => Calculating Formal Charge • For each atom in a valid Lewis structure: • Count the number of valence electrons • Subtract all its nonbonding electrons • Subtract half of its bonding electrons Chapter 1

  18. X => Ionic Structures Chapter 1

  19. Chapter 1

  20. Chapter 1

  21. Chapter 1

  22. Resonance • Only electrons can be moved (usually lone pairs or pi electrons). • Nuclei positions and bond angles remain the same. • The number of unpaired electrons remains the same. • Resonance causes a delocalization of electrical charge. Example=> Chapter 1

  23. Resonance Example • The real structure is a resonance hybrid. • All the bond lengths are the same. • Each oxygen has a -1/3 electrical charge. => Chapter 1

  24. Chapter 1

  25. Major Resonance Form • has as many octets as possible. • has as many bonds as possible. • has the negative charge on the most electronegative atom. • has as little charge separation as possible. Example=> Chapter 1

  26. major minor, carbon does not have octet. => Major Contributor? Chapter 1

  27. Chapter 1

  28. Full structural formula (no lone pairs shown) Line-angle formula Condensed structural formula Molecular formula Empirical formula CH3COOH C2H4O2 CH2O => Chemical Formulas Chapter 1

  29. Chapter 1

  30. Chapter 1

  31. Calculating Empirical Formulas • Given % composition for each element, assume 100 grams. • Convert the grams of each element to moles. • Divide by the smallest moles to get ratio. • Molecular formula may be a multiple of the empirical formula. => Chapter 1

  32. => Arrhenius Acids and Bases • Acids dissociate in water to give H3O+ ions. • Bases dissociate in water to give OH- ions. • Kw = [H3O+ ][OH- ] = 1.0 x 10-14 at 24°C • pH = -log [H3O+ ] • Strong acids and bases are 100% dissociated. Chapter 1

  33. conjugate acid conjugate base acid base => BrØnsted-Lowry Acids and Bases • Acids can donate a proton. • Bases can accept a proton. • Conjugate acid-base pairs. Chapter 1

  34. pKa 4.74 pKb 3.36 pKb 9.26 pKa 10.64 => Acid and Base Strength • Acid dissociation constant, Ka • Base dissociation constant, Kb • For conjugate pairs, (Ka)(Kb) = Kw • Spontaneous acid-base reactions proceed from stronger to weaker. Chapter 1

  35. Determining Relative Acidity • Electronegativity • Size • Resonance stabilization of conjugate base => Chapter 1

  36. => Electronegativity As the bond to H becomes more polarized, H becomes more positive and the bond is easier to break. Chapter 1

  37. => Size • As size increases, the H is more loosely held and the bond is easier to break. • A larger size also stabilizes the anion. Chapter 1

  38. => Resonance • Delocalization of the negative charge on the conjugate base will stabilize the anion, so the substance is a stronger acid. • More resonance structures usually mean greater stabilization. Chapter 1

  39. Chapter 1

  40. nucleophile electrophile => Lewis Acids and Bases • Acids accept electron pairs = electrophile • Bases donate electron pairs = nucleophile Chapter 1

  41. Chapter 1

  42. End of Chapter 1 Chapter 1

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