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Organic Acids and Bases

Organic Acids and Bases. Brønsted Acids and Bases. “ Brønsted-Lowry ” is usually shortened to “ Brønsted ” A Brønsted acid is a substance that donates a hydrogen cation (H + ) A Brønsted base is a substance that accepts the H +

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Organic Acids and Bases

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  1. Organic Acids and Bases

  2. Brønsted Acids and Bases • “Brønsted-Lowry” is usually shortened to “Brønsted” • A Brønsted acid is a substance that donates a hydrogen cation (H+) • A Brønsted base is a substance that accepts the H+ • “proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton

  3. Ka – the Acidity Constant • The concentration of water as a solvent does not change significantly when it is protonated • The molecular weight of H2O is 18 and one liter weighs 1000 grams, so the concentration is ~ 55.6 M at 25° • The acidity constant,Ka for HA is Keq times 55.6 M (leaving [water] out of the expression) • Ka ranges from 1015 for the strongest acids to very small values (10-60) for the weakest

  4. pKa’s of Some Common Acids

  5. pKa – the Acid Strength Scale • pKa = –log Ka • The free energy in an equilibrium is related to –log of Keq (DG = –RT log Keq) • A smaller value of pKa indicates a stronger acid and is proportional to the energy difference between products and reactants • The pKa of water is 15.74

  6. 2.10 Organic Acids and Organic Bases • Organic Acids: • characterized by the presence of positively polarized hydrogen atom

  7. Organic Acids • Those that lose a proton from O–H, such as methanol and acetic acid • Those that lose a proton from C–H, usually from a carbon atom next to a C=O double bond (O=C–C–H)

  8. Organic Bases • Have an atom with a lone pair of electrons that can bond to H+ • Nitrogen-containing compounds derived from ammonia are the most common organic bases • Oxygen-containing compounds can react as bases when with a strong acid or as acids with strong bases

  9. 2.11 Acids and Bases: The Lewis Definition • Lewis acids are electron pair acceptors and Lewis bases are electron pair donors • Brønsted acids are not Lewis acids because they cannot accept an electron pair directly (only a proton would be a Lewis acid) • The Lewis definition leads to a general description of many reaction patterns but there is no scale of strengths as in the Brønsted definition of pKa

  10. Lewis Acids and the Curved Arrow Formalism • The Lewis definition of acidity includes metal cations, such as Mg2+ • They accept a pair of electrons when they form a bond to a base • Group 3A elements, such as BF3 and AlCl3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases • Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids • Organic compounds that undergo addition reactions with Lewis bases (discussed later) are called electrophiles and therefore Lewis Acids • The combination of a Lewis acid and a Lewis base can shown with a curved arrow from base to acid

  11. Illustration of Curved Arrows in Following Lewis Acid-Base Reactions

  12. Lewis Bases • Lewis bases can accept protons as well as Lewis acids, therefore the definition encompasses that for Brønsted bases • Most oxygen- and nitrogen-containing organic compounds are Lewis bases because they have lone pairs of electrons • Some compounds can act as both acids and bases, depending on the reaction

  13. Dissociation of Carboxylic Acids • Carboxylic acids are proton donors toward weak and strong bases, producing metal carboxylate salts, RCO2+M • Carboxylic acids with more than six carbons are only slightly soluble in water, but their conjugate base salts are water-soluble

  14. Acidity Constant and pKa • Carboxylic acids transfer a proton to water to give H3O+ and carboxylate anions, RCO2, but H3O+ is a much stronger acid • The acidity constant, Ka,, is about 10-5 for a typical carboxylic acid (pKa ~ 5)

  15. Substituent Effects on Acidity • Electronegative substituents promote formation of the carboxylate ion

  16. Inductive Effects on Acidity • Fluoroacetic, chloroacetic, bromoacetic, and iodoacetic acids are stronger acids than acetic acid • Multiple electronegative substituents have synergistic effects on acidity

  17. 20.3 Biological Acids and the Henderson-Hasselbalch Equation • If pKa of given acid and the pH of the medium are known, % of dissociated and undissociated forms can be calculated using the Henderson-Hasselbalch eqn Henderson-Hasselbalch equation

  18. 20.4 Substituent Effects on Acidity

  19. Aromatic Substituent Effects • An electron-withdrawing group (-NO2) increases acidity by stabilizing the carboxylate anion, and an electron-donating (activating) group (OCH3) decreases acidity by destabilizing the carboxylate anion • We can use relative pKa’s as a calibration for effects on relative free energies of reactions with the same substituents

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