Medicinal Chemistry: Functional Groups, Chemical Interactions, and Acid-Base Equilibria

Medicinal Chemistry and Functional Groups

Introduction to Medicinal Chemistry

Medicinal chemistry is a branch of chemistry focused on the design, synthesis, and study of the properties of new drugs. It integrates principles from organic chemistry, pharmacology, and biochemistry to optimize the efficacy and safety of pharmaceutical agents.

• Definition: The science of designing and creating new drugs, and understanding their chemical properties.

• Applications: Drug discovery, development, and optimization.

Functional Groups in Organic Molecules

Functional groups are specific groups of atoms within molecules that are responsible for the characteristic chemical reactions of those molecules. They often contain heteroatoms (atoms other than carbon and hydrogen, such as oxygen, nitrogen, or sulfur) and impart unique properties to the compounds.

• Definition: Specially arranged groups of atoms within molecules.

• Heteroatoms: All functional groups contain at least one heteroatom.

• Properties: Functional groups determine reactivity, solubility, and interaction with biological targets.

Example: Functional Groups in Captopril

Captopril is a drug molecule that contains several functional groups, each contributing to its chemical and pharmacological properties.

• Functional Groups Present: Thiol (-SH), carboxylic acid (-COOH), amide (-CONH-), methyl (-CH3).

• Interactions: These groups enable captopril to participate in ionic interactions, hydrogen bonding, and van der Waals forces with biological targets.

• Stability and Solubility: The nature of the functional groups affects the drug's stability at room temperature, solubility in gastrointestinal fluids, and metabolism in the human body.

Additional info: The presence of polar and non-polar functional groups influences both the pharmacokinetics and pharmacodynamics of the drug.

Importance of Functional Groups

Functional groups are the building blocks of organic molecules and are essential in determining the properties and biological activities of drugs and biomolecules.

• Structural Role: Serve as the foundation for molecular architecture.

• Property Determination: Influence solubility, reactivity, and interaction with biological systems.

Chemical Interactions

Types of Chemical Bonds and Interactions

Chemical interactions are fundamental to the structure and function of organic molecules. They vary in strength and are classified as covalent or non-covalent.

• Covalent Bond: Strongest type; electrons are shared between atoms. Typical bond energy is approximately 100 kcal/mol.

• Ionic Bond: Attraction between oppositely charged ions.

• Hydrogen Bond: Attraction between an electropositive hydrogen and an electronegative atom with unshared electron pairs; energy ~1-10 kcal/mol.

• Dipole-Dipole Bonding: Interaction between polar groups with permanent dipoles; energy ~3-7 kcal/mol.

• Van der Waals Forces: Weakest; arise from instantaneous and induced dipoles; energy ~1 kcal/mol.

Order of Bond Strength:

• Covalent bond > Ionic bond > Hydrogen bond > Dipole-dipole bonding > Van der Waals force

Clarification of Terms

• Bond: A stable association between atoms or molecules.

• Bonding: The process of forming bonds.

• Interactions: Includes both covalent and non-covalent associations.

• Forces: Physical phenomena that cause attraction or repulsion between molecules.

• Attraction: The tendency of molecules to come together due to intermolecular forces.

Analogy

• Covalent bonds: Like marriages—strong and enduring.

• Non-covalent bonds: Like dating couples—less permanent, more dynamic.

Drug Interactions with Biological Systems

Drugs interact with themselves, excipients, biological fluids, membranes, proteins, and therapeutic targets through various types of chemical bonds and intermolecular forces.

• Covalent bonds: Often involved in drug-target binding for irreversible inhibitors.

• Intermolecular bonding: Includes hydrogen bonds, ionic interactions, and van der Waals forces, crucial for reversible drug-target interactions.

Covalent Bonds and Dipoles

Nature of Covalent Bonds

Covalent bonds are formed by the sharing of electrons between atoms, providing stability to organic molecules.

• Bond Energy: Approximately 100 kcal/mol.

• Electron Sharing: Can be equal (non-polar) or unequal (polar).

Polar vs. Non-Polar Covalent Bonds

The distribution of electron density in covalent bonds determines whether a bond is polar or non-polar.

• Non-polar covalent bond: Electrons are shared equally; no permanent dipole.

• Polar covalent bond: Electrons are shared unequally due to differences in electronegativity, resulting in a permanent dipole.

• Permanent dipoles: Arise from polar covalent bonds.

• Instantaneous/Induced dipoles: Temporary dipoles due to fluctuations in electron density.

Electronegativity: The tendency of an atom to attract electrons in a bond. For example, fluorine is the most electronegative element.

Atom	Electronegativity
H	2.1
C	2.5
N	3.0
O	3.5
F	4.0

Additional info: The difference in electronegativity between atoms determines the polarity of the bond.

Intermolecular Bonding

Types of Intermolecular Forces

Intermolecular forces are responsible for the physical properties of substances, such as boiling point, solubility, and melting point.

• Dipole-dipole bonding: Occurs between polar groups with permanent dipoles.

• Hydrogen bonding: Involves an electropositive hydrogen and an electronegative atom (e.g., O, N) with unshared electron pairs.

• Van der Waals forces: Occur between non-polar groups due to instantaneous and induced dipoles; contribute to hydrophobic interactions.

• Ionic bond: Occurs between ions of opposite charge.

• Ion-dipole bonding: Occurs between an ion and a polar molecule; important for water solubility of drugs.

Hydrogen Bond Donors and Acceptors

• H-bond donor: A molecule or group that provides the hydrogen atom in a hydrogen bond (e.g., -OH, -NH).

• H-bond acceptor: A molecule or group that provides the electronegative atom with lone pairs (e.g., O, N).

Properties of Water and Solubility

Water as a Solvent

Water is a polar solvent capable of acting as both a hydrogen bond donor and acceptor, making it essential for biological systems and drug solubility.

• Polarity: Water molecules have a permanent dipole.

• Hydrogen bonding: Water can form extensive hydrogen bonds with itself and other molecules.

Solubility of Drugs

• Hydrophilic: Water-loving; readily dissolves in aqueous media.

• Hydrophobic: Water-hating; prefers non-aqueous or lipid environments.

• Lipophilic: Lipid-loving; dissolves in lipid media.

• Lipophobic: Lipid-hating; rarely used term.

Solubility Prediction

• Aqueous solubility: Increases with the number and strength of hydrophilic functional groups.

• Lipid solubility: Increases with the number and strength of hydrophobic functional groups.

Importance: Solubility affects drug formulation, release, absorption, distribution, elimination, and interactions with biological targets.

Acid-Base Equilibria

Water Self-Ionization and Equilibrium Constant

Water can self-ionize to form hydrogen and hydroxide ions, described by the equilibrium constant $K_w$.

• $H_2O \rightleftharpoons H^+ + OH^-$

• $K_w = [H^+][OH^-]$

• At 25°C, $K_w = 10^{-14}$

Brønsted-Lowry Acids and Bases

• Acids: Proton donors ($HA + H_2O \rightleftharpoons H_3O^+ + A^-$; simplified: $HA \rightleftharpoons H^+ + A^-$)

• Bases: Proton acceptors ($B + H_2O \rightleftharpoons BH^+ + OH^-$; simplified: $BH^+ \rightleftharpoons B + H^+$)

Conjugate Acids and Bases

• Conjugate base: The species formed after an acid loses a proton ($A^-$).

• Conjugate acid: The species formed after a base gains a proton ($BH^+$).

Ionization States

• Acids: Un-ionized (uncharged, molecular, free acid) vs. ionized (deprotonated, charged).

• Bases: Un-ionized (uncharged, molecular, free base) vs. ionized (protonated, charged).

Strengths of Acids and Bases

Acid and Base Strengths

• Strong acids: More likely to dissociate and lose $H^+$.

• Weak acids: Less likely to dissociate.

• Strong bases: More likely to attract $H^+$.

• Weak bases: Less likely to attract $H^+$.

• Conjugate strengths: The strength of a conjugate base is inversely related to the strength of its acid, and vice versa.

Quantitative Expression of Strengths

• pH: $pH = -\log[H^+]$

• pKa: $pK_a = -\log K_a$ (acid dissociation constant)

• pKb: $pK_b = -\log K_b$ (base dissociation constant)

• Relationship: $pK_w = pK_a + pK_b = 14$ at 25°C

Acidity and Basicity:

• Strong acids: low pKa

• Strong bases: high pKa (of the conjugate acid)

Common Acidic and Basic Functional Groups

Acidic Functional Group	Typical pKa Range
Sulfonic acids (RSO3H)	-1 to 1
Carboxylic acids (RCOOH)	2 to 6
Phenols (ArOH)	9 to 11
Thiols (RSH)	9 to 11
Imides (RCONHCOR)	9 to 10
Sulfonamides (RNHSO2R)	9 to 10

Basic Functional Group	Typical pKa Range (of conjugate acid)
Guanidines	~13
Aliphatic amines	9 to 11
Aryl amines	1 to 6
N-containing aromatic rings	1 to 6

Additional info: Most weak bases are nitrogen-containing groups, but not all nitrogen-containing groups are basic.