Module 1

Pharmaceutical Chemistry

Inorganic and Organic Pharmaceutical and Medicinal Chemistry

Kathreen Mae D. Cascabel
BS Pharmacy

General Chemistry

Chemistry: The study of matter,

• Matter: Anything that occupies space and has mass.

Key Concepts of Chemistry

1. Composition: The types of substances and their amounts.
2. Structure: Arrangement of atoms in a substance.
3. Changes that Matter Undergoes: Physical and chemical transformations of matter.
4. Energy Involved in Changes: Interaction and transformation energy in chemical processes.

Definitions

• Mass: Refers to the amount of matter present in a material.
• Weight: $ext{Weight} = ext{Mass} imes ext{Pull of gravity}$

Classification of Matter

• Element: Simplest form of matter, one kind of material or atom.
• Compound: Substances composed of two or more elements united chemically in definite proportions.
• Mixture: Composed of two or more substances that are not chemically combined.

Units of Measurement

• Length: Meter (m)
• Mass: Kilogram (kg)
• Time: Seconds (sec)
• Temperature: Kelvin (K)

Definitions of Plasma and Serum

• Plasma: Colorless fluid part of blood, lymph, or milk in which corpuscles or fat globules are suspended.
• Serum: Amber-colored protein-rich liquid that separates out when blood coagulates; used to provide immunity and as a diagnostic agent.

Classification of Mixtures

A. According to Nature of Particles

• Homogeneous: Uniform composition throughout mixture.
• Heterogeneous: Composition is not uniform.

B. According to the Size of Particles

1. Solution: Uniform mixture (homogeneous) composed of solute and solvent where the solute is soluble.
2. Suspension: Coarse mixture with finely divided solid materials distributed in a liquid, where solid is insoluble.
3. Colloid: Mixture with particles of solute not broken down to molecular size, remaining suspended and evenly dispersed throughout the medium.

Processes of Separating Components of Mixtures

1. Decantation: Based on differences in specific gravity.
2. Distillation: Involves evaporation and condensation.
3. Evaporation: Process of removing liquid to leave solids behind.
4. Magnetic Separation: Uses magnets to separate magnetic materials.
5. Filtration: Uses filters to separate solids from liquids.
6. Sorting: Manual separation based on physical characteristics.
7. Centrifugation: Involves spinning to hasten settlement of precipitates.
8. Fractional Crystallization: Reducing temperature so that the least soluble component crystallizes first.
9. Chromatography: Based on differences in solvent affinity.

Properties of Matter

1. Intrinsic/Intensive Properties: Independent of mass or amount (e.g. density, specific gravity, melting point).
2. Extrinsic/Extensive Properties: Dependent on mass (e.g. weight, volume, pressure, heat content).

Changes that Matter Undergoes

1. Physical Change: Changes in phase without altering chemical composition.
2. Chemical Change: Changes in both intrinsic and extrinsic properties that create new substances; evidenced by several signs:
      - Evolution of gas.
      - Formation of precipitate.
      - Emission of light.
      - Generation of electricity.
      - Production of mechanical energy.
      - Absorption or liberation of heat.

Types of Chemical Reactions

A. Direct Union:
    ext{Fe} + ext{S}
ightarrow ext{FeS} (all sulfides are black).
B. Decomposition/Analysis:
    ext{Na}_2 ext{CO}_3
ightarrow ext{Na}_2 ext{O} + ext{CO}_2
C. Single Replacement:
    ext{Na} + ext{HCl}
ightarrow ext{NaCl} + ext{H}_2
D. Double Displacement:
    ext{NaCl} + ext{AgNO}_3
ightarrow ext{NaNO}_3 + ext{AgCl}

Processes Involved in Chemical Change

1. Oxidation: Involves gaining oxygen.
2. Reduction: Involves losing oxygen or gaining hydrogen.
3. Neutralization: Acid plus base yields salt and water:
       ext{acid} + ext{base}
   ightarrow ext{salt} + ext{water}
4. Hydrolysis: Water plus salt yields acid and base:
       ext{water} + ext{salt}
   ightarrow ext{acid} + ext{base}
5. Saponification: Alkali plus fats produces soap and glycerol:
       ext{alkali} + ext{fats}
   ightarrow ext{soap} + ext{glycerol}
6. Fermentation: Organic substances plus microbes yields alcohol:
       ext{organic substances} + ext{microbes}
   ightarrow ext{alcohol}
7. Nuclear Change: Change in structure, properties, composition of the nucleus, found in transmutation of element:
      - Nuclear Fission: Splitting of a heavy atom.
      - Nuclear Fusion: Union of two light atoms to form a bigger molecule.

Atomic Structure

Similarities and Differences Among Atoms

Isotopes (vary in neutrons), Isotones (vary in protons), Isobars (vary in mass number).

Key Atom Theorists and Their Proposals

• Democritus: Matter is made of small indivisible particles called "atomos".
• John Dalton: Presented the "Billiard Ball Model"; proposed that matter is made of atoms which are indestructible and combine to form compounds.
    - Dalton's Atomic Theory includes:
      - Matter is made up of, composed of indivisible particles (atoms).
      - All atoms of a given element are alike.
      - Atoms combine in simple numerical ratios.
• J.J. Thompson: Developed the "Raisin Bread Model"; discovered electrons.
• Rutherford: Proposed the nuclear model after the Gold Foil Experiment that indicated atoms are mostly empty space.
• Niel Bohr: Developed the planetary model of an atom.
• Erwin Schrodinger: Proposed the Quantum Mechanical Model; referred to as the "Electron Cloud" Model.
• James Chadwick: Discovered the neutron.

Basic Concepts Related to Atomic Structure

• Protons = Electrons = Atomic Number
• Neutrons = Mass Number – Atomic Number
• $ext{Mass Number} = ext{Protons} + ext{Neutrons}$
• Electrons = Protons - Charge
     - Example for ion P = 6, Cation=12:
       - $ext{E} = ext{P} - ext{Charge} = 6 - (+1) = 5$
     - Anion = 12 - 1:
       - $ext{E} = ext{P} - ext{Charge} = 6 - (-1) = 6 - 0 = 7$
  
  Ions - charged atoms:
• Atomic Number: Number of protons
• Mass Number: Protons + neutrons

Allotropism

• Allotropes: Different forms of an element that exist in the same physical state, e.g., carbon exists as diamond and graphite.

Quantum Numbers

• Principal Quantum Number (n): Main energy level (1, 2, …).
• Azimuthal Angular Momentum (l): Subshell shape (s, p, d, f) determined by the value of l which ranges from 0 to n-1.
• Magnetic Quantum Number (m or ml): Orientation of orbitals from -l to +l.
• Spin Quantum Number (s or ms): Direction of electron spin, either -1/2 or +1/2.

Key Chemical Laws Related to Matter

1. Law of Conservation of Mass: Total mass remains constant; mass of reactants equals mass of products.
2. Law of Definite Proportions (Proust’s Law): A chemical compound contains exactly the same proportion of elements by mass.
3. Law of Multiple Proportions: When elements combine they do so in ratios of small whole numbers.

General Principles in Electron Configuration

1. Pauli’s Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers.
      - Each orbital can only hold 2 electrons (e.g., 1s²).
2. Heisenberg’s Uncertainty Principle: It is impossible to determine simultaneously both the position and momentum of an electron.
3. Aufbau Principle: Electrons fill the lowest energy orbitals first.
4. Hund’s Rule: Orbitals of the same energy are filled singly before pairing up.

Periodic Table Development

• Antoine-Laurent Lavoisier: Hydrogen as the first true periodic table.
• Johann Dobereiner: Connected elements into triads.
• Newlands: Developed the Law of Octaves.
• Meyer and Mendeleev: Noted physical/chemical properties as periodic functions depending on atomic weights.
• Henry Mosley: Organized elements by atomic number, the current method used today.

Group Classifications

• Family A (Representative Elements): Includes S block (Alkali Metals, Alkaline Earth) and P block elements (Groups IIIA to VIIIA).
• Family B (Transition Elements): Involves D block transition metals and F block lanthanides (Rare Earth Metals).

Periodic Trends (Excluding Noble Gases)

1. Atomic Radius: Measured as half the distance between the nuclei of two atoms.
2. Ionization Potential: Energy required to remove an electron from an atom.
3. Metallic Property: Tendency to lose electrons.
4. Electron Affinity: Energy change upon adding an electron to a neutral atom.
5. Electronegativity: Measure of ability to attract electrons (increases up a group and across a period).

Kinetic Molecular Theory (KMT)

• Explains phases of matter based on the motion of molecules:
     - SOLID: Defined shape and volume.
     - LIQUID: Defined volume but takes on the shape of the container.
     - GAS: No defined shape or volume.

Solutions

• Definition: Homogeneous mixture (single-phase) consisting of two or more substances.
• Types of Solutions:
     - Saturated: Maximum solute dissolved.
     - Unsaturated: Less than maximum solute.
     - Supersaturated: More solute dissolved than can be held.

Factors Affecting Solubility

1. Nature of Solute and Solvent: Principle of "like dissolves like."
      - Solubility: Maximum amount of solute in grams that dissolves in 100 g of water.
      - Miscibility: Ability of one substance to mix with another.
2. Temperature: Increased temperature generally increases the solubility of solids but decreases that of gases.
3. Pressure (for gases): According to Henry’s Law, solubility of a gas increases with pressure.
4. Particle Size/Surface Area: Smaller particles have larger surface areas, enhancing solubility.
5. Presence of Salts:
      - Salting Out: Salt presence decreases solubility.
      - Salting In: Salt presence increases solubility.

Methods of Expressing Concentration of Solutions

• Percent Solutions:
     - % by mass
     - % by volume
     - % mass/volume
     - ppm (parts per million)
     - proof
• Moles (n): Moles = grams/MW
• Mole Fraction (x): Mole fraction = moles of the substance/total moles of solution
• Molarity (M): Moles of solute per liter of solution.
• Molality (m): Moles of solute per kilogram of solvent.
• Normality (N): Number of equivalents of solute per liter of solution.

Forces of Attraction

Intramolecular Forces (within molecules)

1. Ionic Bond: formed by electron transfer.
2. Covalent Bond: formed by electron sharing.
      - Polar Covalent Bond: unequal sharing of electrons.
      - Nonpolar Covalent Bond: equal sharing of electrons.

Intermolecular Forces (between molecules)

1. Van der Waals Forces:
      - Keesom (Dipole-Dipole): Orientation/Alignment effect.
      - Debye (Dipole-Induced Dipole): Induction.
      - London Dispersion: Extremely close proximity creates temporary dipoles.
2. Ion-Dipole: Attraction between charged ions and polar molecules.
3. Ion-Induced Dipole: Interaction between charged ions and nonpolar molecules.
4. Hydrogen Bond: Occurs between hydrogen and electronegative atoms (F, O, N, Cl, S).

Physical Properties of Systems

1. Additive Property: Depends on the sum (e.g., molecular weight).
2. Constitutive Property: Type and arrangement of components (e.g., optical rotation).
3. Colligative Property: Depend on the number of solute particles in a solution (e.g., Vapor pressure lowering, Boiling point elevation, Freezing point depression, Osmotic pressure).

Key Equations

• Density: $ext{Density} = \frac{ ext{mass}}{ ext{volume}}$
• Specific Gravity: $ext{Specific Gravity} = \frac{ ext{density of sample}}{ ext{density of standard}}$
• Colligative Properties:
     - Vapor Pressure Lowering: $ext{ΔP} = P° imes ext{mole fraction of solute}$
     - Boiling Point Elevation: $ext{ΔT_b} = K_b m$
     - Freezing Point Depression: $ext{ΔT_f} = K_f m$
     - Osmotic Pressure: $ext{π}V = nRT$, $ext{π} = \frac{n}{V}RT$

Gas Law Formulas

1. Boyle's Law: $P_1 V_1 = P_2 V_2$ or $P imes V ext{ is constant}$
2. Charles' Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$
3. Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$
4. Combined Gas Law: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$
5. Ideal Gas Law: $PV = nRT$,
   Where: R = Ideal gas constant (0.08206 L atm/mol K).

Real Gas Deviations

• Van der Waals Equation:
  $ext{(P + \frac{a n^2}{V^2})(V - nb) = nRT}$
  Where a and b are constants.

Acids and Bases

Definitions

• Electrolytes: Substances that conduct electricity (strong electrolytes dissociate completely; weak electrolytes do not).
• Non-Electrolytes: Substances that do not dissociate in solution and cannot conduct electricity.

Acid-Base Theories

• Arrhenius Theory:
     - Acid: Yields H+ or H3O+.
     - Base: Yields OH-.
• Bronsted-Lowry Theory:
     - Acid: Proton donor.
     - Base: Proton acceptor.
• Lewis Theory:
     - Acid: Electron pair acceptor.
     - Base: Electron pair donor.

Properties of Acids and Bases

• Acids: Sour taste, turn litmus from blue to red, react with metals to produce H2, react with carbonates to produce CO2, generally colorless in phenolphthalein.
• Bases: Bitter taste, turn litmus from red to blue, slippery feel, color change in indicator phenolphthalein from colorless to pink.

Neutralization Reaction

• Neutralization: Acid + Base → Salt + Water:
     ext{HA} + ext{BOH}
  ightarrow ext{BA} + ext{H}_2 ext{O}
• Titration: Gradual addition of a standard solution to determine the endpoint where acid and base fully react.

pH Concepts

• pH: Negative logarithm of hydrogen ion concentration:
  $ext{pH} = - ext{log} [ ext{H}^+]$
• For weak acids:
  $ext{pH} = \frac{1}{2} pK_a - ext{log} [ ext{Ca}];$
• For weak bases:
  $ext{pH} = pK_w - pK_b + ext{log} [ ext{base salt}]$

Buffers

Definition

• Buffers: Solutions that resist changes in pH when acids or bases are added; consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.

Capacity of Buffers

• Buffer Capacity: Ability to resist pH change, highest at pH = pKa.
• Henderson-Hasselbalch Equation:
  $ext{pH} = ext{p}K_a + ext{log}\frac{[ ext{Salt}]}{[ ext{Acid}]}$

Thermochemistry

Key Concepts

• Heat (q): Energy transfer due to temperature difference.
• Work (w): Energy transfer occurring as a result of movement against force (compression or expansion).
• Internal Energy (U): Total energy of the particles in a system.
• Enthalpy (H): Energy of a reaction at constant pressure.
• Entropy (S): Measure of disorder in a system, indicative of the number of possible arrangements.

Heat Capacity and Specific Heat

• Heat Capacity (c): Amount of heat required to raise the temperature of a substance by one degree.
• Specific Heat: Heat required to raise 1 gram of a substance by 1°C.

Types of Thermodynamic Processes

1. Endothermic Process: Absorption of heat, +ΔH.
2. Exothermic Process: Release of heat, -ΔH.

Fundamental Laws of Thermodynamics

1. Conservation of Energy: Energy can neither be created nor destroyed, only transformed.
2. Entropy Law: The total entropy of a closed system can only increase over time.
3. Entropy Approach: Entropy approaches a constant at absolute zero temperature.

Alkaline and Alkali Metals

Group IA: Alkali Metals

1. General Properties: Highly reactive, valence 1.
2. Example of Metals:
      - Hydrogen: Lightest and most reactive; isotopes: Protium, Deuterium, Tritium.
      - Lithium: Lightest metal, used in medications (e.g., mood stabilizers).
      - Sodium: Major cation in extracellular fluid; plays a critical role in fluid balance.
      - Potassium: Most abundant intracellular cation, critical in muscle function.
3. Compounds and Actions:
      - Lithium Salts: Used for bipolar disorder.
      - Sodium Compounds: NaCl (common salt), used in electrolytic balance.
      - Potassium Compounds: Various applications including heart health and diuretic functions.

Group IIA: Alkaline Earth Metals

1. General Properties: Less reactive than alkali metals, valence 2.
2. Key Metals:
      - Beryllium: Most toxic, rarely used in medicine.
      - Magnesium: Second most abundant intracellular cation, plays vital roles in enzymatic functions.
      - Calcium: Crucial for bone health, muscle contraction, and neurotransmitter release.

Water and its Medicinal Uses

1. Purified Water: Used in making injections and pharmaceutical products.
2. Water for Injection: Must pass pyrogen tests, used for survival fluid replacement.

Group IIIB and Transition Metals

Features of These Groups

• Coinage Metals: Found in elemental state; includes Copper, Silver, and Gold.
• Key Properties:
     - Copper: Essential for iron utilization; treatment of Wilson's disease.
     - Silver: Used for its germicidal properties (Argyria).
     - Gold: Recognized for its therapeutic uses in arthritis.

Summary of Relevant Compounds

• Copper Compounds: CuSO4 (Blue vitriol, used in medical treatments).
• Silver Compounds: AgNO3 (used as a germicide).
• Gold Compounds: Auranofin (treatment of rheumatoid arthritis).

Chemical Properties of Specific Elements and Their Compounds

• The physical and chemical properties of elements dictate their medicinal applications.
• A detailed analysis of reactions, structural implications, and biological interactions enhances understanding of pharmacological principles in medicinal chemistry.

Common Names of Chemicals

Summary of Organic Chemistry Principles

• Organic Chemistry focuses on carbon-containing compounds, their structural and functional relationships, and reactivity characteristics.
• Proper understanding of hydrocarbons, functional groups, and molecular interactions is essential for drug design and medicinal applications.

Conclusive Remarks

Pharmaceutical Chemistry integrates both Inorganic and Organic Chemistry essential for understanding the chemical basis of drugs, their interactions with biological systems, and the advanced methodologies used in drug development.

This document serves as a comprehensive study guide spanning various principles of chemistry, vital for students pursuing pharmacy and related fields. It encompasses definitions, significant theories, illustrations of chemical properties, examples of chemical reactions, and applications in pharmacology, creating a solid foundation in pharmaceutical chemistry.