Organic acids and bases
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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

Organic Acids and Bases

Br nsted 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+

    • “proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton

K a the acidity constant
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

P k a s of some common acids
pKa’s of Some Common Acids

P k a the acid strength scale
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

2 10 organic acids and organic bases
2.10 Organic Acids and Organic Bases

  • Organic Acids:

  • characterized by the presence of positively polarized hydrogen atom

Organic acids
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)

Organic bases
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

2 11 acids and bases the lewis definition
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

Lewis acids and the curved arrow formalism
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

Illustration of curved arrows in following lewis acid base reactions
Illustration of Curved Arrows in Following Lewis Acid-Base Reactions

Lewis bases
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

Dissociation of carboxylic acids
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

Acidity constant and p k a
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)

Substituent effects on acidity
Substituent Effects on Acidity

  • Electronegative substituents promote formation of the carboxylate ion

Inductive effects on acidity
Inductive Effects on Acidity

  • Fluoroacetic, chloroacetic, bromoacetic, and iodoacetic acids are stronger acids than acetic acid

  • Multiple electronegative substituents have synergistic effects on acidity

20 3 biological acids and the henderson hasselbalch equation
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

Aromatic substituent effects
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