John E. McMurry Paul D. Adams University of Arkansas Organic 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 Brønsted Acids and Bases

 The concentration of water as a solvent does not change significantly when it is protonated  The molecular weight of H 2 O is 18 and one liter weighs 1000 grams, so the concentration is ~ 55.6 M at 25°  The acidity constant, K a for HA is K eq times 55.6 M (leaving [water] out of the expression)  K a ranges from for the strongest acids to very small values ( ) for the weakest K a – the Acidity Constant

pK a ’s of Some Common Acids

 pK a = –log K a  The free energy in an equilibrium is related to –log of K eq (  G = –RT log K eq )  A smaller value of pK a indicates a stronger acid and is proportional to the energy difference between products and reactants  The pK a of water is pK a – the Acid Strength Scale

 Organic Acids: - characterized by the presence of positively polarized hydrogen atom 2.10 Organic Acids and Organic Bases

 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) Organic Acids

 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 Organic Bases

 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 pK a 2.11 Acids and Bases: The Lewis Definition

 The Lewis definition of acidity includes metal cations, such as Mg 2+  They accept a pair of electrons when they form a bond to a base  Group 3A elements, such as BF 3 and AlCl 3, are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases  Transition-metal compounds, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4, 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 Lewis Acids and the Curved Arrow Formalism

Illustration of Curved Arrows in Following Lewis Acid-Base Reactions

 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 Lewis Bases

 Carboxylic acids are proton donors toward weak and strong bases, producing metal carboxylate salts, RCO 2  + M  Carboxylic acids with more than six carbons are only slightly soluble in water, but their conjugate base salts are water-soluble Dissociation of Carboxylic Acids

 Carboxylic acids transfer a proton to water to give H 3 O + and carboxylate anions, RCO 2 , but H 3 O + is a much stronger acid  The acidity constant, K a,, is about for a typical carboxylic acid (pK a ~ 5) Acidity Constant and pK a

 Electronegative substituents promote formation of the carboxylate ion Substituent Effects on Acidity

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

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

20.4 Substituent Effects on Acidity

 An electron-withdrawing group (-NO 2 ) increases acidity by stabilizing the carboxylate anion, and an electron-donating (activating) group (OCH 3 ) decreases acidity by destabilizing the carboxylate anion  We can use relative pK a ’s as a calibration for effects on relative free energies of reactions with the same substituents Aromatic Substituent Effects