Organic Chemistry for Dummies 2nd Edition by Arthur Winter, PhD

Chapter 4: Covering the Bases (And the Acids)

A Defining Moment: Acid-Base Definitions

Arrhenius acids and bases: A little watery

• Arrhenius acids: molecules that dissociate in water to make the hydronium ion, H3O+.

  • Strong Arrhenius acids: completely dissociate in water to make hydronium ions.

    • Example: Nitric acid (HNO3) is a strong acid because it completely dissociates in water to make hydronium ions.

  • Weak Arrhenius acids: acids that only partially dissociate in water.

    • Example: Acetic acid (CH3COOH) only partially dissociates in water and is a weak acid.

• Arrhenius bases: molecules that dissociate to make hydroxide ions, OH–.

  • Strong Arrhenius bases: bases that dissociate completely to generate hydroxide ions.

    • Example: Potassium hydroxide (KOH) is a strong base because it completely dissociates in water to make hydroxide ions.

  • Weak Arrhenius bases: bases that only partially dissociate to generate hydroxide ions.

    • Example: Beryllium hydroxide (Be[OH]2) is a weak base because it only partially dissociates in water.

Pulling for protons: Brønsted-Lowry acids and bases

• Brønsted-Lowry acid: molecule that donates a proton (H+) to a base.

• Brønsted-Lowry base: molecule that accepts a proton from an acid.

  • Note: An H+ ion is called a proton because the hydrogen atom has no neutrons or electrons — just a single proton at the nucleus.

• Conjugate base: what the deprotonated acid becomes

• Conjugate acid: what the protonated base becomes

Electron lovers and haters: Lewis acids and bases

• Lewis acid: molecule that accepts a pair of electrons to make a covalent bond.

  • Also called electrophiles: molecules that can accept electrons (Lewis acids) to make a new bond and are electron lovers.

• Lewis base: molecule that donates electrons to make a covalent bond.

  • Also called nucleophiles: molecules that can donate electrons (Lewis bases) to form bonds and are nucleus lovers.

    • Example:

      • Borane (BH3) is an example of a Lewis acid. Borane doesn’t have a full octet of valence electrons Because it doesn’t have a full octet of electrons, it can accept a lone pair from a molecule like methylamine (CH3NH2), which has a lone pair of electrons, and use this electron pair to fill its octet.

        • Because BH3 accepts a pair of electrons to make a covalent bond, it’s said to be a Lewis acid.

        • Because methylamine donates electrons to make a bond, it’s said to be a Lewis base.

  • Note: Any Brønsted acid will also be a Lewis acid, and any Brønsted base will also be a Lewis base.

Comparing Acidities of Organic Molecules

• The strength of an acid is directly proportional to the stability of the acid’s conjugate base.

  • An acid that has a more stable conjugate base will be more acidic than an acid that has a less stable conjugate base.

    • Strong acids have stable conjugate bases.

• Acidic molecules have structural features that allow the anion in the conjugate base to delocalize the charge over a larger space.

  • Delocalization of the negative charge (so one atom doesn’t have to bear the full negative charge) makes the molecule more stable.

Comparing atoms

• Which atom does the negative charge of the acid’s conjugate base stay?

  • It prefers to rest on electronegative (electron-loving) elements.

    • Therefore, a negative charge is more stable on oxygen than it is on nitrogen.

    • “For that reason, alcohols (R—OH) are more acidic than amines (R—NH2), which in turn are more acidic than alkanes (R—CH3).

  • Atomic size also stabilizes the negative charge. Charges prefer to be on larger atoms than on smaller atoms.

    • This preference results from large atoms allowing the negative charge to delocalize over a much larger region of space, instead of being concentrated in a small region (as it would on a small atom).

  • Atom size trumps electronegativity considerations.

    • Example: Fluorine is a more electronegative atom than iodine, HI is more acidic than HF.

      • The much larger iodine atom allows the negative charge to delocalize over a larger space than does the much smaller fluorine atom, and thus makes hydrogen iodide more acidic.

Seeing atom hybridization

• The orbital on which the lone-pair anion rests also affects the acidity.

  • Lone-pair anions prefer to reside in orbitals that have more s character than p character, because s orbitals are closer to the atom’s nucleus than p orbitals, and the electrons are stabilized by being closer to the nucleus.

  • Orbitals that are sp hybridized have 50 percent s character.

  • Orbitals that are sp2 hybridized have 33 percent s character.

  • Orbitals that are sp3 hybridized have 25 percent s character.

    • Therefore, anions prefer to be in orbitals that are sp hybridized over those that are sp2 hybridized orbitals. They prefer to be in orbitals that are sp2 hybridized over those that are sp3 hybridized.

Seeing electronegativity effects

• Electron-withdrawing groups on an acid also stabilize the conjugate base anion by allowing some of the charge on the anion to delocalize to other parts of the molecule.

  • Example: Trifluoroethanol is more acidic than ethanol.

    • The highly electronegative fluorine atoms on trifluoroethanol pull electron density away from the anion, taking away some of the negative charge from the oxygen, thereby stabilizing the molecule.

Seeing resonance effects

• Acids with conjugate bases that allow the negative charge to be delocalized through resonance are stronger acids than acids whose conjugate bases don’t have resonance structures.

  • Example: acetic acid is much more acidic than ethanol because the conjugate base anion of acetic acid can delocalize the negative charge through resonance.

Defining pKa: A Qualitative Scale of Acidity

• pKa value of an acid: a quantitative measurement of a molecule’s acidity.

  • pKa is derived from the equilibrium constant for the acid’s dissociation reaction, Ka, and uses a logarithmic scale to allow the pKa values to span wide ranges.

    • pKa = –log Ka

      • Lower the pKa value of an acid, the stronger the acid.

      • Higher the pKa value, the weaker the acid.

      • Very strong acids have pKa values of less than zero.

      • Weak acids generally have pKa values of between 0 and 9.

Problem Solving: Predicting the Direction of Acid-Base Reactions at Equilibrium

• Weak acids and bases are lower in energy than strong acids and bases,

  • Equilibria favor the reaction side with the lowest-energy species; acid-base reactions will go to the side with the weakest acids and bases.

• Example: Predict the direction of the acid-base reaction between hydrogen cyanide (HCN) and acetate (C2H3O2–)

  • Hydrogen cyanide (pKa = 9) has a higher pKa value than acetic acid (pKa = 5), the equilibrium will lie to the left, in the direction of the weaker acid and base.

• If you know the pKa values of the two acids on both sides of the equation, equilibrium will favor the side with the acid that has the highest pKa.