General Chemistry non-Labs (SA)

Overview

1 SECTION A (18 H LECTURES)
1.1 General Chemistry - Elements, Atoms, Ions, Electrons and the Periodic Table (co-taught with CHM1011). 10h lectures
This course aims to introduce the fundamental principles of atoms from the chemists’ viewpoint. Starting from a simple model and using the results of quantum mechanics a more appropriate model of the atom is presented. From this model trends in atomic and ionic properties which enable us to explain differences and similarities and predict the properties of different elements can be deduced. The following topics are covered:
 The Basics: Element, The periodic table, atom, mole.
 Electromagnetic radiation: energy, wavelength and frequency.
 The Atom: The Bohr Atom.
 The Electron: Wave-Particle Duality and The Schrödinger Wave Equation, Probability Density, Radial Distribution Function, Orbitals, Quantum Numbers, s and p Orbitals, Phase, d Orbitals.
 More than One Electron: Filling orbitals, The Aufbau principle, The Pauli Exclusion Principle, Hund’s rules, Penetration, Shielding, Effective Nuclear Charge, Slater’s Rules, Size.
 Trends: Ionization energy, Electron attachment enthalpy (affinity), Electronegativity, Ionic radii, Polarizability and polarizing power, Hydration enthalpies.
 Redox reactions: assigning oxidation numbers, oxidizing and reducing agents, redox potentials.potentials.
 Percentage composition, empirical formula and molecular formula, determining limiting reactant. How to balance simple redox reactions.
1.2 Structure, Bonding and Inorganic Chemistry. 8h lectures.
This course includes some important theories of bonding and an introduction to inorganic chemistry. Theories of bonding are discussed in some detail for discrete molecules. The discussion of bonding in molecular species centres on the valence bond theory. Intermolecular forces between molecules are also discussed.
 Introduction to bonding: Discussion of types of structure and common bonding theories, examples of representative structures.
 Nomenclature: Binary molecular compounds, binary ionic compounds, binary acids, oxoacids.
 Homonuclear Diatomic Molecules: Interatomic distance and covalent radii, Potential energy curves, attractive and repulsive forces, bond energy and enthalpy. Lewis structures, filled shells, the octet rule. Wavefunction, introduction to valence bond theory and molecular orbital theory, Valence bond theory: ionic and covalent contributions, resonance.
 Heteronuclear Diatomic Molecules: Lewis structures, valence bond approach, electronegativity, electric dipole moments, carbon monoxide, isoelectronic molecules.
 Polyatomic Molecules: Metal complexes and covalent polyatomics, coordination number, common geometries, molecules obeying the octet rule, valence bond theory, expanding the octet, hybridization (sp, sp2, sp3), formal charge, single, double and triple carbon-carbon bonds, molecular shapes.
 Valence-shell electron pair repulson theory.
 Macromolecules and intramolecular forces, fullerenes, proteins and hydrogen bonding.
 Intermolecular Forces: Van-der-Waal forces, strength of forces.
 Intermolecular forces and solubility. “like dissolves like”.
 Precipitation and solubility rules.


2 SECTION B (derives from CCE1102) (22 h lectures 8 h seminars);
2.1 Chemical Equilibria (10 hours lectures, 3 hours seminars)
 Introduction to physical chemistry: review states of matter and introduction to ideal gases and ideal solutions, enthalpy and internal energy, Hess cycles, heat capacity and heat transfer.
 Chemical Equilibria: Definitions and calculations involving equilibrium constants Kc and Kp, including examples of homogeneous and heterogeneous equilibria (Ksp). Definitions and calculations involving enthalpy of 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. The Common Ion Effect.
 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. Examples of polyfunctional acids and their behaviour in titrations. Acid-base strength as a function of molecular structure.
2.2 Kinetics (6 hours lectures, 2 hours seminars)
 Key concepts: Elementary reactions, reaction molecularity, molecularity vs. stoichiometry, definition of reaction rates, calculating reaction rates from experimental data, writing differential rate laws, reaction orders, order vs. molecularity, reaction rate constants, initial rates method, integrated rate laws (how and why), collision and transition state theories, Arrhenius rate law and activation energy, reaction kinetics in relation to reaction mechanism. Catalysis. Methods of measuring reaction rates.
 Derivation of rate equations: Derivation of zero, 1st and 2nd rate laws, experimental data analysis and visualisation, Obtaining reaction orders and rate constants by the initial rates method. Analysing data using integrated rate laws and making predictions.
 Collision and transition state theories: The Arrhenius equation. Activation by collision and the collision theory. Measurement of activation energies.
 Classes of reaction: Simple gas phase reactions. Chain and branched chain reactions. Reactions in solution, reactions of solids, catalysed reactions
2.3 Basic Thermodynamics, (6 h lectures; 2 h seminars)
 State function, standard conditions, enthalpy of formation.
 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.
 The second law of thermodynamics. Examples and calculations using standard entropies; entropy changes with volume, temperature, phase transitions and chemical reactions.
 Absolute entropy and the third law of thermodynamics.
 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.

3 PROBLEM CLASSES
 Section A – 2 x 2 h
 Section B – 2 x 2 h

Learning Objectives

On completion of this module a learner should be able to:
 Students should aim to achieve a solid grounding in the fundamental principles of
atomic structure, the principal quantum numbers and s and p orbitals and the periodic table, including the ability to answer problems related to these concepts. They should be able to explain and use the Aufbau principle. They should aim to achieve a good understanding of, and be able to explain, the trends in atomic and ionic properties in the periodic table. They should develop the ability to use these concepts to explain and predict the properties of the different elements.
 Students should be able to draw Lewis structures for simple molecules that obey the octet rule and be able to use hybridisation to describe more complex structures and especially single, double and triple bonds to carbon. They must be able to understand how and why resonance is used to describe a structure.
 Students will become familiar with chemical descriptions of matter. What matter is made up of, how it can be organised into the periodic table and how we can start to understand it from a scientific perspective.
 Explain and use equations to describe chemical systems at equilibrium.
 Understand and apply the basic rules of chemical kinetics.
 Describe and apply the general principles of the first and second laws of thermodynamics.
 Learn how to present scientific concepts in poster form and back up poster with a flash presentation.

Skills

Learners are expected to demonstrate the following on completion of the module:
 Ability to write and predict atomic structure and properties.
 Ability to explain and understand bonding.
 Thermodynamic and kinetic problem solving (including numerical).
In addition,
 Communication – spoken, poster and flash presentation
 Numeracy – basic algebra and calculus.
 Improved independent learning and time management.

Assessment

Course Requirements:
Overall mark must be over 40%

Assessment Profile:
Element type Element weight (%)
1. Class test 50
2. Problem classes 30
3. Presentation project 20