Physical Chemistry 1

Overview

Chemical Equilibria (8 hours lectures, 2 hour seminar):
1.1 Units and dimensional analysis: mass, moles and mole fraction. mass and concentration
calculations
1.2. Chemical Equilibria: Definitions and calculations involving equilibrium constants Kc and Kp, and
the reaction quotient Q, including examples of homogeneous and heterogeneous equilibria
(Ksp). Definitions and calculations involving enthalpy of reaction and solution and lattice energy.
Application of Le Chatelier’s Principle to determine the effect of change in concentration,
pressure, temperature and catalyst on the composition of the reaction mixture and the
equilibrium constant. Solubility rules, precipitation and the Common Ion Effect.
1.3. Acids and Bases: Definitions of conjugate acid and base; strong and weak acids and bases.
Calculation of pH, pKa and pKb. The special case of water Kw and pKw Terminology in
acid/base titrations, calculation of pH at the end point, and indicators. Definition of a buffered
solution and calculation of its pH and on addition of acid or base. Examples of polyfunctional
acids and their behaviour in titrations.

Basic Thermodynamics, (12 hours lectures, 3 hours seminars):
2.1. The zeroth and first laws of thermodynamics. Introduction to zeroth law of thermodynamics, the
first law of thermodynamics, state function, standard conditions, enthalpy of formation.
2.2. The direction of spontaneous change. Spontaneous vs non-spontaneous change. Entropy as
criterion for spontaneous change. Reversible and irreversible processes. Classical and
molecular basis of entropy.
2.3. The second law of thermodynamics. Examples and calculations using standard entropies;
entropy changes with volume, temperature, phase transitions and chemical reactions.
2.4. Absolute entropy and the third law of thermodynamics.
2.5. Chemical equilibrium: Gibbs energy and spontaneity, energy minimum, direction of chemical
change and influence of enthalpy and entropy. Variation of Gibbs energy with temperature,
pressure and concentration. Gibbs energy relationship with equilibrium and the equilibrium
constant. Examples and calculations.
2.6. Gibbs energy and phase equilibria. The thermodynamics of transition. Review of one
component and two component phase diagrams. Liquid-liquid equilibria, phase separation,
critical solution temperature, distillation of partially miscible liquids. Examples of extractions,
separations and molecular interactions.

Phase Equilibria (8 hours lectures, 2 hours seminars):
3.1. States and Properties of Matter: States of matter, solutions, mixtures and colloids. Changes in
state, heating and cooling curves and P-T phase diagrams. Material properties: ideal gas law,
molar volume.
3.2. Relating intermolecular forces and physical properties: ionic, dipole-dipole, London-dispersion
forces, induced dipole, H-bonding (mpt, bpt, vapour pressure, volatility, surface tension).
3.3. One Component Systems: Vapour pressure – temperature relationships: examples using the
Clapeyron and Clausius-Clapeyron Equations, the Antoine Equation.
3.4. Two-component systems (vapour-liquid equilibria of ideal systems): Raoult’s Law,
Dalton’s Law applied to two component systems, mole fraction – Pressure diagrams (ideal and
non-ideal systems relating to intermolecular forces) volatility and relative volatility,
constant pressure diagrams (x,y and T-x,y) and the Lever rule.
3.5. Introduction to non-ideal solutions: An azeotrope being a mixture that vaporizes and
condenses without a change in composition; a eutectic being a mixture that freezes and
melts without change of composition.
3.6. Colligative properties: Relative lowering of vapour pressure, boiling point elevation
and boiling point depression in binary solutions containing non -volatile solutes; osmotic
pressure.

Learning Objectives

On completion of this module a learner should be able to:

 Explain and use equations to describe chemical systems at equilibrium. A1.2, Level B, Point 3.
A2.2 Level B, Point 3 and 4
 Describe the general principles of phase equilibria as applied to single and binary component
systems. A1.2, Level B, Point 3. A2.2 Level B, Point 3 and 4
 Understand and apply the basic rules of chemical kinetics. A1.2, Level B, Point 3
 Describe and apply the general principles of the first and second laws of thermodynamics. A1.2,
Level B, Point 3. A2.2 Level B, Point 3 and 4
 Describe equilibria of electrochemical cells and discuss applications of electrochemical theory.

Skills

Learners are expected to demonstrate the following on completion of the module:
 Basic thermodynamic, kinetic and electrochemistry knowledge.
 Thermodynamic and kinetic problem solving (including numerical)
 Communication – spoken during seminars and written in class tests and exam. A5.2, Level B, Point 2
 Numeracy – basic algebra and calculus.
 Improved independent learning and time management.
 Problem-solving – solving problems in seminars, quizzes and exams and seminars.

Assessment

Assessment Profile
Element type Element weight (%)
1. Examination (2hrs) 3 questions, one from each section 80
2. Quizzes/homeworks 20

Course Requirements:
Examination and overall module mark must both be passed 40% veto.