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Chapter 1 Introduction and Review

Organic Chemistry , 6 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. Atoms: protons, neutrons, and electrons.

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Chapter 1 Introduction and Review

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  1. Organic Chemistry, 6th 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 • Atoms: protons, neutrons, and electrons. • The number of protons determines the identity of the element. • Some atoms of the same element have a different number of neutrons. These are called isotopes. • Example: 12C, 13C, and 14C Chapter 1

  4. Electronic Structure • Electrons: outside the nucleus, in orbitals. • Electrons have wave properties. • Electron density is the probability of finding the electron in a particular part of an orbital. • Orbitals are grouped into “shells,” at different distances from the nucleus. Chapter 1

  5. First Electron Shell Spherical shape The 1s orbital holds two electrons. Chapter 1

  6. Three p orbitals Second Electron Shell P-orbital (dumbell shaped) S-Orbital (spherical) Second shell (4 orbitals) holds total of 8 electrons Chapter 1

  7. Electronic Configurations • Aufbau principle: Place electrons in lowest energy orbital first. • Hund’s rule: Equal energy orbitals are half-filled, then filled. • 6C: 1s2 2s2 2p2    Chapter 1

  8. Electronic Configurations => Chapter 1

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

  10. Lewis Structures • Bonding electrons represented by a dash line or a pair of electron • Nonbonding electrons or lone pairs are shown Satisfy the octet rule! Chapter 1

  11. Multiple Bonding => Chapter 1

  12. Dipole Moment Amount of electrical charge x bond length. Chapter 1

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

  14. 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 F.C = # valence e – (non-bonding e + 1/2(bonding e) Chapter 1

  15. X => Ionic Structures If the DEN is greater than 2.5, the bond is a ionic bond. Chapter 1

  16. 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. Chapter 1

  17. 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

  18. Major Resonance Form or Major Contributor • 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. Chapter 1

  19. minor contributor, carbon does not have octet majorcontributor => Resonance Hybrid Chapter 1

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

  21. 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

  22. = 1 3.33 3.33 3.33 = 1.98 = 2 = 1 Sample Problem An unknown compound has the following composition: 40.0% C, 6.67% H, and 53.3% O. Find the empirical formula. Empirical Formula: CH2O Chapter 1

  23. 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

  24. conjugate base conjugate acid => acid base BrØnsted-Lowry Acids and Bases • Acid: donates a proton. • Base: accepts a proton. • Conjugate acid-base pairs. Chapter 1

  25. 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

  26. Structural Effects on Acidity • Electronegativity • Size • Resonance stabilization of conjugate base Chapter 1

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

  28. 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

  29. Resonance • Delocalization of electron pairs. • More resonance structures usually mean greater stabilization. Acidity: Chapter 1

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

  31. End of Chapter 1 Chapter 1

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