Brønsted-Lowry and Lewis Acids and Bases

Brønsted-Lowry Acids and Bases

• A Brønsted-Lowry acid is defined as a proton (H+) donor.
• A Brønsted-Lowry base is defined as a proton (H+) acceptor.
• All Brønsted-Lowry acids contain a proton.
• The net charge on Brønsted-Lowry acids can be zero, positive (+), or negative (-).
• All Brønsted-Lowry bases contain a lone pair of electrons or a π bond.
• The net charge on Brønsted-Lowry bases can be zero or negative (-).

Lewis Acids and Bases

• A Lewis acid is a substance capable of accepting a pair of electrons, with examples including H+, AlCl₃, and ZnCl₂.
• A Lewis base donates a pair of electrons, like OH- and NH₃.

Molecules Acting as Acids or Bases

• Some molecules, can act as either acids or bases based on the reaction conditions.
• Morphine is an example of such a molecule; it contains both hydrogen atoms and lone pairs.
• The OH bonds in morphine make it an acid.
• The lone pairs and π bonds in morphine make it a base.

Reactions of Brønsted-Lowry Acids and Bases

• A Brønsted-Lowry acid-base reaction involves the transfer of a proton from an acid to a base.
• In this reaction, one bond is broken and another is formed.
• The electron pair of the base forms a new bond to the proton of the acid.
• The acid loses a proton, leaving the electron pair in the H-A bond on A.
• Electron movement in reactions is illustrated using curved arrows; two curved arrows are needed as two electron pairs are involved.
• Loss of a proton from an acid forms its conjugate base.
• Gain of a proton by a base forms its conjugate acid.
• A double reaction arrow indicates the reaction can proceed in both forward and reverse directions, indicating equilibrium.

Lewis Acids and Bases Reactions

• The Lewis concept of acid-base reactions involves the transfer of electrons.
• Many organic reactions can be considered reactions of Lewis acids or bases.

Basicity of Amines

• Amines are considered organic bases due to the presence of a nitrogen atom with a lone pair of unshared valence electrons.
• Methyl amine (CH₃NH₂) is an example.

Carboxylic Acids

• Carboxylic acids are organic acids, with ethanoic acid (CH₃COOH) as an example.
• They contain the -COOH functional group.
• In aqueous solutions, carboxylic acids ionize, with H₂O acting as a base.

Acidity and Basicity Constant

• Acidity strength reflects the tendency of an acid to donate a proton.
• The more readily a compound donates a proton, the stronger the acid.
• Acidity is measured by an equilibrium constant.
• When a Brønsted-Lowry acid (H-A) dissolves in water, an acid-base reaction occurs, and an equilibrium constant can be written for the reaction.
• Acidity constant is given by Ka = [H3O+][A-]/[H-A].
• It's more convenient to use pKa values than Ka values, where pKa= -log Ka.
• Strong acids have large Ka values and low pKa values.
• Weak acids have small Ka values and high pKa values.
• An increase in Ka (stronger acid) correlates with a decrease in pKa.

Factors Affecting Acid Strength

• Key factors affecting acidity: structure of the compound, electronegativity of the atom bonded to hydrogen, hybridization, inductive effect, and resonance stability.
• Compounds that stabilize a conjugate base A⁻ make the starting acid H-A more acidic.
• If HA is more stable than A⁻, then HA is a weak acid.
• Methanol (CH₃O-H) has a pKa of 16, while methane (CH₃-H) has a pKa of 43.
• Methanol is more acidic than methane due to presence of oxygen, which is more electronegative than carbon.

Hybridization and its Effect on Acidity

• Increased s-character of the hybrid orbital of a carbon atom increases its electronegativity.
• This leads to increased polarity and acid strength of the C-H bond.
• The strongest acid yields the weakest base, and the weakest acid yields the strongest base.

Inductive Effect on Acidity

• Ethanoic acid (CH₃COOH) has a pKa of 4.75, while chloroethanoic acid (ClCH₂COOH) has a pKa of 2.81.
• CICH₂COOH is more acidic than CH₃COOH.
• The difference in pKa arises from the inductive effect of the chlorine atom on the molecule.

Resonance Stability

• The greater the resonance stability of the conjugate base of a compound, the more acidic the compound.
• Phenol is more acidic than alcohol due to the resonance stability of the phenoxide ion compared to the alkoxide ion.

Basicity

• In water, a weak base (NH₃) dissociates and the equilibrium constant is given by K = [NH₄+][OH-]/[NH₃].
• pKb = - log Kb.
• Key factors affecting base strength: structure of the compound, hybridization, and resonance.

Amines

• Amines are basic compounds due to the nitrogen atom's lone pair of electrons.
• In aqueous solution, an amine is a weak base and accepts a proton from water in a reversible acid-base reaction.
• If the free amine is stabilized relative to the cation, the amine is less basic.
• If the cation is stabilized relative to the free amine, the amine is a stronger base.
• An electron-releasing group on the nitrogen increases basicity by increasing the availability of electrons on nitrogen.

Hybridization and Basicity

• An sp² orbital has more s-character than an sp³ orbital, making an sp² nitrogen more electronegative.
• A molecule with an sp² nitrogen is less basic because its unshared electrons are more tightly held.
• The lone pair electrons are not available for sharing with a proton.

Resonance and Basicity

• Resonance stabilization also affects the base strength of an amine.
• In an amide, the acyl group's electron-withdrawing effect makes the electron pair on the nitrogen less available.
• Resonance stabilization makes amides weaker bases; they are only weakly basic in aqueous solutions.