2.10 Organic Acids and Organic Bases

Organic Acids and Bases

Organic acids and bases are defined by specific structural features that determine how they behave in reactions. Acids have polarized O–H bonds that can donate a proton, while bases have lone electron pairs (usually on nitrogen or oxygen) that can accept one. Their strength is measured by pKa values, and they form conjugate pairs when reacting. Biological compounds like amino acids showcase both acidic and basic properties, which makes them especially important in living systems.

Structural Features of Organic Acids and Bases

Organic acids contain a polarized O–H bond. Oxygen is more electronegative than hydrogen, so the hydrogen carries a partial positive charge ($\delta+$) and the oxygen carries a partial negative charge ($\delta-$). This polarization makes it easier for the hydrogen to leave as $H^+$.

Common acidic functional groups:

• Carboxylic acids (–COOH): the most common organic acid group
• Phenols (Ar–OH): an –OH group attached directly to an aromatic ring
• Sulfonic acids (–$SO_3H$): strong organic acids found in detergents and dyes

Organic bases contain a lone pair of electrons on nitrogen or oxygen. That lone pair can accept a proton ($H^+$), forming a conjugate acid.

Common basic functional groups:

• Amines: primary ($R\text{-}NH_2$), secondary ($R_2NH$), tertiary ($R_3N$)
• Aromatic amines ($Ar\text{-}NH_2$): the lone pair overlaps with the aromatic ring, making these weaker bases than alkylamines
• Imines ($R_2C\text{=}NR$): the nitrogen lone pair is in an $sp^2$ orbital, also making them weaker bases than typical amines

Charge Stabilization in Conjugate Bases

When an organic acid donates a proton, the conjugate base carries a negative charge. The more stable that negative charge is, the stronger the acid. Two main factors stabilize conjugate bases:

1. Electronegativity of the atom bearing the charge. Oxygen stabilizes a negative charge much better than carbon because oxygen is more electronegative and holds electrons more tightly.
2. Resonance delocalization. If the negative charge can be spread over multiple atoms through resonance, the conjugate base is more stable. More equivalent resonance structures generally means greater stabilization.

Here's how this plays out in key conjugate bases:

• Carboxylate anion ($RCOO^-$): the negative charge is delocalized equally over two oxygen atoms through resonance. This is why carboxylic acids are much more acidic than alcohols.
• Phenoxide anion ($ArO^-$): the negative charge on oxygen delocalizes into the aromatic ring through resonance, though not as effectively as in carboxylates.
• Sulfonate anion ($RSO_3^-$): the negative charge is spread over three oxygen atoms, making sulfonic acids among the strongest organic acids.

Acid-Base Theories and Quantitative Measures

Two theoretical frameworks come up repeatedly in organic chemistry:

• Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors. This is the definition you'll use most often in organic chemistry.
• Lewis theory is broader: acids are electron pair acceptors and bases are electron pair donors. This framework covers reactions where no proton transfer occurs, like $BF_3$ accepting a lone pair from an amine.

pKa is the quantitative measure of acid strength. A lower pKa means a stronger acid. For reference:

• Sulfonic acids: pKa ≈ –1
• Carboxylic acids: pKa ≈ 4–5
• Phenols: pKa ≈ 10
• Alcohols: pKa ≈ 16–18

The Henderson-Hasselbalch equation relates pH to pKa and the ratio of conjugate base to acid:

$pH = pK_a + \log\frac{[A^-]}{[HA]}$

When $pH = pK_a$, the acid and its conjugate base are present in equal concentrations. Buffer solutions exploit this relationship: a mixture of a weak acid and its conjugate base resists pH changes because each component can neutralize added acid or base.

Acid-Base Behavior of Biological Compounds

Nitrogen-containing compounds typically act as bases:

• Amines accept a proton via the lone pair on nitrogen, forming ammonium ions ($R\text{-}NH_3^+$). Alkylamines are moderately strong bases.
• Pyridine and imidazole are aromatic nitrogen heterocycles that also act as bases through a lone pair on nitrogen. Imidazole (found in the amino acid histidine) has a pKa near physiological pH (~7), which makes it especially important in enzyme catalysis.

Oxygen-containing compounds range in acidity:

• Alcohols and phenols are weakly acidic. They can donate $H^+$ to form alkoxide or phenoxide anions, but alcohols require a very strong base to deprotonate.
• Carboxylic acids are moderately strong organic acids and readily donate $H^+$ to form carboxylate anions.

Amino acids are unique because they contain both an acidic carboxyl group (–COOH) and a basic amino group (–$NH_2$). In solution, an internal proton transfer occurs from the carboxyl to the amino group, producing a zwitterion: a molecule with no net charge but with localized positive ($NH_3^+$) and negative ($COO^-$) charges.

The dominant form of an amino acid depends on pH:

1. Low pH (acidic conditions): the amino group is protonated ($NH_3^+$) and the carboxyl group is neutral (–COOH). The molecule carries a net positive charge.
2. At the isoelectric point (pI): the zwitterionic form predominates. The molecule has no net charge.
3. High pH (basic conditions): the carboxyl group is deprotonated ($COO^-$) and the amino group is neutral (–$NH_2$). The molecule carries a net negative charge.