Alkaline Ionisation and pKa Study Notes

Learning Outcomes and Fundamentals of Acid-Base Behaviour

• After completion of this material, students should be able to:

  • Understand how acid/base behavior fundamentally links to ionisation.

  • Understand the influence ionisation has on drug dissolution and absorption processes.

  • Calculate, predict, and interpret $pKa$ values.

  • Work with the Henderson-Hasselbalch Equation.

  • Sketch and interpret ionisation curves and calculate percentage ($\%$) ionisation values.

• The outcome of acidic or basic behavior is the formation of an ionised entity. Examples provided include:

  • Phenolate ions.

  • Carboxylate ions.

  • Ammonium ions.

• The acid-base behavior of drugs and their associated ionisation influences several critical parameters:

  • Solubility of the compound.

  • Drug analysis methodology.

  • Pharmacokinetics, including absorption, distribution, and excretion.

  • Pharmacodynamics (the drug's effect on the body).

  • Compatibility with other drugs in a formulation.

Impact of Ionisation on Drug Absorption and Dissolution

• Absorption of drugs mainly occurs via passive diffusion through cell membranes along a concentration gradient.

• Lipophilicity vs. Ionisation:

  • The more lipophilic a drug is (i.e., its unionised form), the faster and better its absorption across lipid membranes.

  • However, for drugs administered orally, drug dissolution in the Gastrointestinal Tract (GIT) is often the rate-limiting factor.

• Hydrophilicity vs. Ionisation:

  • The more hydrophilic a drug is (i.e., its ionised form), the faster and better its dissolution in aqueous environments.

• The Pharmaceutical Balance:

  • For effective oral administration, there must be a balance between lipophilic and hydrophilic characteristics.

  • Lipophilic/hydrophilic behavior, and consequently absorption and dissolution, is strongly influenced by the $pH$ of the surrounding environment.

Dissociation Constants $Ka$ and $Kb$

• In contrast to strong acids/bases, weak acids and bases dissociate only partially into their ionised forms in water.

• For weaker acids/bases, an equilibrium exists between conjugated acid-base pairs.

• Equilibrium for Acetic Acid:     $CH_3COOH_{(aq)} + H_2O_{(l)} \rightleftharpoons CH_3COO^{-}_{(aq)} + H_3O^{+}_{(aq)}$

  • In this reaction, Acetic acid acts as the acid and the Acetate ion acts as the conjugated base.

• Equilibrium for Ammonia:     NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH_{4}^{+}_{(aq)} + OH^{-}_{(aq)}

• For weak acids and bases, the dissociation constants $Ka$ and $Kb$ are typically expressed logarithmically.

The $pKa$ Value and Acidity/Basicity Scales

• Definition of $pKa$: The $pKa$ is commonly used to express the strength of acids and bases. The conjugated acid $A^{-}$ acts as a base, and a conjugated base $BH^{+}$ acts as an acid.

• Core Correlation:

  • The stronger the acid, the lower the $pKa$.

  • The stronger the base, the higher the $pKa$.

• Distinction Between $pH$ and $pKa$:

  • $pH = -\log([H_3O^{+}])$. It describes the acidity of a specific solution.

  • $pKa$ is a characteristic constant for a particular chemical compound.

• Typical $pKa$ Values for Functional Groups:

  • Carboxylic acids: approximately $3$ to $5$.

  • $1^{\circ}$ (Primary) aliphatic amines: approximately $7$.

  • Phenols: approximately $9$.

  • $2^{\circ}$ / $3^{\circ}$ (Secondary/Tertiary) aliphatic amines: approximately $9$ to $10$.

  • Aromatic amines: approximately $2$ to $5$.

Structural Influence on Acid and Base Strength

• $pKa$ is strongly influenced by chemical structure.

• General rules to predict the strength of an acid ($HA$):

  • Electronegativity: The stronger the electronegativity of the atom bearing the negative charge in the anion ($A^{-}$), the more stable $A^{-}$ is, making $HA$ a stronger acid.

  • Resonance: The greater the resonance stabilisation of the anion $A^{-}$, the stronger the acid $HA$.

  • Inductive Effects: A negative inductive effect (electron-withdrawing) favors the formation of $A^{-}$, leading to a stronger acid $HA$.

Resonance (Mesomeric) and Inductive Effects

• Resonance (Mesomeric) Effects:

  • Involves the movement of electrons to different atoms, resulting in charge delocalization within a molecule (indicated by curly arrows).

  • Stability Examples:

    • Carboxylate anion: Negative charge is stabilized by delocalization across two oxygen atoms ($pKa = 4.76$).

    • Phenoxide ion: Stabilized by resonance within the aromatic ring ($pKa = 10$).

    • Alkoxide ion: Not resonance stabilized, making the parent alcohol a very weak acid ($pKa = 15.9$).

• Inductive Effects:

  • Refers to the pulling or pushing of electrons within chemical bonds.

  • Positive Inductive Effect ($+I$): Electron-donating effect (e.g., $CH_3 -$ groups).

    • Example: Acetic acid ($pKa = 4.76$) has a $+I$ effect from the methyl group.

  • Negative Inductive Effect ($-I$): Electron-withdrawing effect (e.g., $CF_3 -$ group).

    • Example: Trifluoroacetic acid ($pKa = 0.23$) is significantly more acidic than acetic acid due to the strong $-I$ effect of the fluorine atoms.

Basicity in Amines and Amides

• Aliphatic Amines (Primary vs. Tertiary):

  • The alkyl group ($R$) exerts a positive inductive ($+I$) effect.

  • A $1^{\circ}$ amine has one such effect; a $3^{\circ}$ amine has three.

  • This results in higher electron density on the Nitrogen lone electron pair (EP) in a $3^{\circ}$ amine compared to a $1^{\circ}$ amine.

  • The higher density makes the Nitrogen a more attractive reaction partner for $H^{+}$ ions, resulting in a stronger base.

• Amides:

  • Amides are generally never basic because the Nitrogen lone electron pair is involved in a resonance structure.

  • Because the lone pair is not available to react with $H^{+}$, the molecule has no ability to donate electrons and shows no basic behavior.

  • Rarely, an amide can be weakly acidic. This occurs if the negative charge resulting from the removal of an $H^{+}$ can be delocalized across a large part of the molecule via resonance. Such wide delocalization encourages the acidic reaction.

The Henderson-Hasselbalch Equation and Ionisation

• The equilibrium between an acid and its conjugated base is influenced by the $pH$ of the environment (e.g., stomach, intestine, blood, or urine).

• The relationship is expressed by the Henderson-Hasselbalch equation:

  • For Acids: $pH = pKa + \log\left(\frac{[Ionised]}{[Unionised]}\right)$

  • For Bases: $pH = pKa + \log\left(\frac{[Unionised]}{[Ionised]}\right)$

• Key Observation: If $pH = pKa$, then equal amounts ($50\% / 50\%$) of the acid and its conjugated base are present.

• If the $pKa$ of a drug is known, the position of the equilibrium (the ratio of ionised to unionised drug) can be calculated for any specific $pH$.

Percentage Ionisation and Ionisation Curves

• Because only one species (either $HA/A^{-}$ for acids or $B/BH^{+}$ for bases) is ionised, the equilibrium position is vital for predicting solubility and absorption.

• Calculation Formulae:

  • For Acids: $\%\text{ ionisation} = \frac{[A^{-}]}{[A^{-}] + [HA]} \times 100$

  • For Bases: $\%\text{ ionisation} = \frac{[BH^{+}]}{[BH^{+}] + [B]} \times 100$

• Ionisation Curve Characteristics:

  • Acidic Drug ($HA \rightarrow A^{-} + H^{+}$):

    • If pH > pKa, the drug is mainly ionised.

    • If pH < pKa, the drug is mainly unionised.

  • Basic Drug ($B + H^{+} \rightarrow BH^{+}$):

    • If pH > pKa, the drug is mainly unionised.

    • If pH < pKa, the drug is mainly ionised.

    • A stronger base (larger $pKa$) will be ionised over a wider range of the $pH$ scale.

  • Amphoteric Drug:

    • Features both acidic and basic groups (e.g., $B + H^{+} \rightarrow BH^{+}$ and $HA \rightarrow A^{-} + H^{+}$).

    • These drugs are highly ionised over the entire $pH$ range due to the presence of both functional groups.

Further Reading

• Material can be found in Chapter 1 of 'Essentials of Pharmaceutical Chemistry' by Cairns (Fourth edition).