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Introduction to Organic Chemistry: Key Concepts and Principles from Wade's 5th Edition

This overview of Chapter 1 from "Organic Chemistry" by L.G. Wade, Jr. provides a foundational understanding of organic chemistry, highlighting critical definitions, atomic structure, bonding types, and Lewis structures. It covers key topics like ionic and covalent bonds, resonance, electronegativity, and acid-base chemistry, including Bronsted-Lowry definitions. Illustrated concepts such as dipole moments, formal charge calculations, and the significance of molecular formulas assist students in grasping essential organic chemistry principles. Understanding these concepts is vital for any chemistry curriculum.

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Introduction to Organic Chemistry: Key Concepts and Principles from Wade's 5th Edition

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  1. Organic Chemistry, 5th EditionL. G. Wade, Jr. Chapter 1Introduction and Review Jo Blackburn Richland College, Dallas, TX Dallas County Community College District ã 2003,Prentice Hall

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

  3. => Atomic Structure • protons, neutrons, and electrons • isotopes 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. => Bond Formation • Ionic bonding: electrons are transferred. • Covalent bonding: electron pair is shared. Chapter 1

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

  9. Multiple Bonding => Chapter 1

  10. => 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

  11. Electronegativity and Bond Polarity Greater EN means greater polarity => Chapter 1

  12. => 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

  13. X => Ionic Structures Chapter 1

  14. 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

  15. 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

  16. 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

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

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

  19. 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

  20. => 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

  21. 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

  22. 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

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

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

  25. => 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

  26. => 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

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

  28. End of Chapter 1 Chapter 1

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