Pharmaceutical Analytical Chemistry - Lecture 1

Pharmaceutical Analytical Chemistry - First Level, Lecture 1

• Lecture focuses on a review of fundamental concepts vital to analytical chemistry, specifically concerning solutions.
• The lecture covers physical properties of solutions, including the solution process itself, different concentration units, and relevant concepts like electrolytes and nonelectrolytes.

The Solution Process

• Intermolecular forces significantly influence the solution process.
• Appropriate solvents can be predicted based on the nature of the substance to be dissolved.
• Solution processes can be exothermic (releasing heat) or endothermic (absorbing heat), depending on the interactions involved. The enthalpy change (ΔH_soln) for a solution process is the sum of the energies involved in breaking solute-solute interactions, solvent-solvent interactions, and forming solvent-solute interactions.

Types of Solutions

• Solutions are homogenous mixtures of two or more substances.
• The solvent is the component present in greater quantity and retains its original state.
• The solute is the component present in lesser quantity, and its state may or may not change in solution.
• Different types of solutions exist, including gas in liquid, liquid in liquid, solid in liquid, and gas in solid, each with examples (such as air, seltzer water, alloys).

Electrolytes and Nonelectrolytes

• Electrolytes are substances that produce electrically conducting solutions when dissolved in water.
• Nonelectrolytes do not yield electrically conductive solutions on dissolution in water.
• Strong electrolytes completely ionize in solution, whereas weak electrolytes only partially ionize.

Polarity and Solubility

• "Like dissolves like" is a fundamental concept; polar substances dissolve other polar substances, and nonpolar substances dissolve other nonpolar substances.
• Water and oil are immiscible (do not mix) due to their different polarities.
• Substances like sodium chloride and water are soluble (able to dissolve) because their polarities are compatible.

The Solution Process (Further Detail)

• Solutions formation depends on interactions between solvent-solvent, solvent-solute, and solute-solute molecules.
• The mixing of solvent and solute molecules may be exothermic or endothermic, as per Hess's law, which states that enthalpy change for a reaction is the sum of the enthalpy changes of the steps in the reaction.

Concentration Units

• Molarity (M) is moles of solute per liter of solution.
• Percent by mass is the mass of solute divided by the total mass of solution, multiplied by 100.
• Mole fraction (X_A) is the moles of component A divided by the total moles of all components.
• Molality (m) is moles of solute per kilogram of solvent.

Colligative Properties

• Colligative properties depend on the number of solute particles, not their identity.
• Examples include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.

Chemical Equilibrium

• Chemical equilibrium is a state where the forward and reverse reaction rates are equal, and the net change in concentrations of reactants and products is zero.
• The equilibrium constant (K) quantifies the relative concentrations of reactants and products at equilibrium.

Rules for Equilibrium Constants

• Equilibrium expressions are written by considering the stoichiometry coefficients of the respective chemical species. The equilibrium expression includes concentration of the products and the reactants raised to the power of the respective stoichiometric coefficient.
• Concentrations of pure solids, liquids, and solvents are excluded from the expression as they do not change significantly.
• The equilibrium constant (K) remains constant at a given temperature.
• The value of K will change for the reverse reaction.

Le Chatelier's Principle

• If a change is made to a reaction at equilibrium, the reaction will shift to counteract that change to re-establish equilibrium.
• Changes causing a reaction shift can include changes in reactant or product concentration, pressure or volume, and temperature.
• The addition of a catalyst will increase the rate of reaching equilibrium, but it will not affect the position of the equilibrium.

Equilibrium and Thermodynamics

• Thermodynamics helps interpret equilibrium, relating it to enthalpy and entropy changes.
• Enthalpy (ΔH): the heat absorbed or released during a reaction.
• Entropy (ΔS): the degree of disorder in a system.
• Free energy (ΔG): combines enthalpy and entropy to determine spontaneity (whether a process will occur naturally).

Solubility Product

• The solubility product (K_sp) is an equilibrium constant that describes the solubility of a solid in an equilibrium with its ions.
• Rules govern the calculation of equilibrium constant values.
• The reaction quotient (Q) helps ascertain whether a precipitate will form.

Problems Concerning Solubility Product

• Different types of problems involving solubility product calculations (type I, type II, type III, Type IV) are described including situations involving common ions.
• Detailed examples illustrate how to calculate K_sp from known ion concentrations, and how to calculate ion concentrations from K_sp under different conditions.

Homework Problems

• Several problems related to solubility product, concentration calculations and equilibrium concepts, are included to help practice problem solving. These problems test understanding and application of the various concepts learned in the lecture.

This lecture provides a comprehensive review of essential concepts in analytical chemistry, focusing on solutions and their physical properties. Key topics include the solution process, concentration units, and the role of electrolytes and nonelectrolytes. Understanding the thermodynamics behind solution processes is also highlighted.