C3420 Physical Chemistry

Faculty of Science
Spring 2023
Extent and Intensity
3/0/0. 3 credit(s) (plus extra credits for completion). Type of Completion: zk (examination).
Teacher(s)
doc. Mgr. Jana Pavlů, Ph.D. (lecturer)
doc. RNDr. Pavel Brož, Ph.D. (lecturer)
Guaranteed by
doc. Mgr. Jana Pavlů, Ph.D.
Department of Chemistry – Chemistry Section – Faculty of Science
Supplier department: Department of Chemistry – Chemistry Section – Faculty of Science
Timetable
Wed 12:00–13:50 C12/311, Wed 16:00–16:50 C12/311
Prerequisites
C1020 General Chemistry && (! C4660 Physical Chemistry I ) && (! C4020 Physical Chemistry II ) && !NOWANY( C4660 Physical Chemistry I , C4020 Physical Chemistry II )
Basic university level knowledge of general chemistry and mathematics (contained in courses: C1020, M0001).
Course Enrolment Limitations
The course is also offered to the students of the fields other than those the course is directly associated with.
fields of study / plans the course is directly associated with
Course objectives
A. Understanding the basic concepts of physical chemistry with respect to their usual time allocation in secondary school subjects;

B. Gradual building of abstraction (which the sequence of topics is subordinated to);

C. Providing a background in physical chemistry usable in simultaneous and subsequent university courses
Learning outcomes
Student will be able to:
- to orient in various areas of physical chemistry such as thermodynamics, kinetics and structure and to describe the relations between the findings from these fields of study;
- accurately explain the terms and principles of physical chemistry;
- describe the relationships between physical quantities by mathematical formulas and graphs;
- derive some of the mathematical formulas used in physical chemistry;
- explain the possibilities of utilisation of the knowledge of physical chemistry in practice;
- apply the acquired knowledge in solving practical problems and questions in chemistry;
- propose a method of teaching of topics from physical chemistry at secondary schools;
Syllabus
  • 1. States of matter. Gas, liquid, solid. 1.1 Properties of gases. Concept of pressure, pressure measurement, concept of temperature, the zero law of thermodynamics. 1.1.1 Ideal gas. Boyle's law, Charles's law, Gay-Lussac's law, and Avogadro's principle. Function of one variable. Perfect gas equation, concept of function of two variables. 1.1.2 Real gas. Intermolecular interactions, virial equation of state, van der Waals equation of real gas.

    2. Thermodynamics I 2.1 Basic concepts. Heat, work and energy. The concept of internal energy and an example of its calculation for monoatomic gas. Concept of state and non-state variables. 2.2 The first law of thermodynamics. Expansion work, concept of increase (differential) of distance, volume and work, concept of a definite integral. The concept of reversible changes. 2.3 Enthalpy and its change, relation to change of internal energy. 2.4 Heat transfer: calorimetry, heat capacity and molar heat capacity. The change of enthalpy and internal energy with temperature. Enthalpy and internal energy of transformation. 2.5 Thermochemistry. Standard enthalpy change, Hess's law, standard enthalpy of formation. Kirchhoff's equation.

    3. Thermodynamics II 3.1 The second law of thermodynamics. Spontaneous direction of change. The concept of energy dispersion, entropy as a reversible change of heat at a given temperature, the second law of thermodynamics. Boltzmann formula for entropy. Calculation of entropy change for isothermal expansion of ideal gas, Carnot cycle, Clausius inequality. Efficiency of heat engine. 3.2 The third law of thermodynamics. 3.3 Focus on the system. Helmholtz (A) and Gibbs (G) energy. Increase of G, A at constant temperature and criteria of spontaneity at constant T, V and constant T, p. Maximum work and maximum non-expansion work. Calculation of standard Gibbs energy of reaction, thermodynamic cycles. 3.4 Combining the first and second laws of thermodynamics. Reversible change of internal energy in closed system. Fundamental equation of chemical thermodynamics. Change of Gibbs energy with temperature and pressure. 3.5 Chemical potential of the pure substance.

    4. Phase equilibria and mixtures 4.1 Phase diagrams of pure substances. 4.1.1 Phase, phase transition, phase stability criteria. Dependence of chemical potential on temperature and pressure, change of melting temperature with pressure. 4.1.2 Phase diagram of pure substance, Gibbs phase rule, coexistence curve. Types of phase transitions. 4.2 Simple mixtures. Description of composition in thermodynamics. Partial molar volumes of components of mixture (e.g. water and ethanol). Partial molar Gibbs energy - chemical potential. 4.2.1 Mixtures of gases. Gibbs energy of mixing of ideal gases, composition of vapour from Dalton's law. 4.2.2 Mixtures of liquids. Ideal solution, saturation vapor pressure of mixture (benzene-toluene example), Raoult's law, Henry's law. Gibbs energy of mixing of liquids, excess functions. Colligative properties of solutions: Ebulioscopic and cryoscopic effect, osmotic pressure. 4.3 Phase diagrams of two-component systems. Binary ideal solution: interpretation of phase diagrams, lever rule, distillation. 4.4 Real solutions. Activity, molar fraction and activity coefficient of component in solution. Electrolytes. Mean activity coefficient of electrolyte. Debye-Hückel limiting Law.

    5. Chemical equilibrium 5.1 Spontaneous chemical reactions. Minimum Gibbs energy, extent of reaction. Course of function of one variable, concept of slope, slope of increasing and decreasing function and of function at extrema. Reaction Gibbs energy. Description of equilibrium of transformation of ideal gas and general reaction. Reaction coefficient x reaction quotient. Dependence of reaction Gibbs energy on the reaction quotient, thermodynamic equilibrium constant and its expression and calculation. Estimating the degree of dissociation at equilibrium. Molecular origin of equilibrium constant. 5.2 The response of equilibria to changing conditions. Response of equilibria to change of pressure. Le Chatelier's principle. Response of exothermic and endothermic reactions to temperature. Van 't Hoff's reaction isotherm and isobar. Measurement of reaction enthalpy. K ~ value at different temperatures.

    6 Electrochemistry. 6.1 Equilibrium electrochemistry. Basic concepts. Half-reactions, electrodes and their types, cathode and anode. Notation for reactions. Cells. Types and notation of cells. Electromotive force and its relationship to reaction Gibbs energy (Nernst equation). Utilisation of Nernst equation. Cells in equilibrium: standard potential, calculation of equilibrium constant from standard cell potential. Application of standard potentials: (a) electrochemical series, energy conversion in biological cells - respiratory chain and oxidative phosphorylation; (b) calculation of equilibrium constants; (c) electrochemical determination of Gibbs and Helmholtz energy and entropy. Types of electrodes and potentiometry. 6.2 Dynamic electrochemistry. Processes at electrodes. Butler-Volmer equation. Polarisation. Electrolysis. Faraday's laws. Galvanic cells. Accumulators. Fuel cells. Corrosion.

    7. Molecules in motion. 7.1 Molecular motion in gases. Pressure and molecular speed, Maxwell distribution of speeds, collision frequency, mean free path, collision flux. Effusion. Transport properties of perfect gas: flux, particle flux in the concentration gradient direction, Fick's first law of diffusion, diffusion coefficient. 7.2 Molecular motion in liquids. 7.3 Electrolyte solutions. Conductivity of electrolyte solutions, molar conductivity - its calculation and relation to conductivity, Kohlraush's law, mobility of ions, Debye-Hückel-Onsager theory, ion channels. 7.4 Diffusion in solutions. Thermodynamic view: thermodynamic force of concentration gradient. Fick's first and second law of diffusion, diffusion with convection, statistical view.

    8. Rates of chemical reactions I. 8.1 Empirical chemical kinetics. Basic concepts and principles: Definition of rate of reaction. Rate of reactant consumption and product formation. Rate equation and rate constant. Molecularity and reaction order. Half-life of reaction. 8.2 Types of chemical reactions. Isolated and simultaneous reactions. 8.3 Kinetics of basic chemical reactions. 8.4 Methods of monitoring the progress of reaction. 8.5 Chemical reactors. 8.6 Temperature dependence of reaction rates: Arrhenius equation. Reaction coordinate, collision theory and activated complex theory.

    9. Rates of chemical reactions II. 9.1 Kinetics of complex reactions. Chain reactions. Photochemical reactions. Catalytic reactions and inhibition. Adsorption. Enzyme-catalysed reactions.

    10. Structure I 1.10 Particles, elementary particles and their properties, bosons and fermions, Fermi-Dirac and Bose-Einstein statistics. Physical fields, types of interactions. Energy of particles. 10.2 Nucleus. Binding energy of nucleus, potential barrier. Magnetic properties of nuclei, quantum numbers of nuclear spins, energy levels of nuclear spin in the magnetic field, Larmor frequency. 10.3 Nuclear magnetic resonance, transitions between energy levels and intensities of lines, chemical shift: delta-scale, shielding, spin-spin interaction: 1:1 doublet, 1:2:1 triplet and 1:3:3:1 quartet. NMR applications in medicine: Magnetic Resonance Imaging.

    11. Structure II 11.1 Introduction to quantum theory. Failure of classical mechanics. Electromagnetic radiation, frequency and energy of radiation, spectrum of radiation. Wave-particle duality, photoelectric effect, electron diffraction. Wave function and its properties, Born interpretation, probability density, normalisation, Schrödinger equation, Hamiltonian operator, eigenfunctions and eigenvalues. 11.2 Electronic structure of atoms. 11.2.1 Hydrogen atom. Atomic orbitals and quantum numbers, energy levels of each AO, wave function: radial and angular part. 11.2.2 Ions of hydrogen type. 11.2.3 Atoms with many electrons. Many-electron and one-electron wave function, energy of orbitals, Pauli and building-up principles, Hund's rule, exchange interaction and exchange energy. Shielding. 11.3 Properties of elements: radii and energies of AO, ionization energy. 11.4 Atomic spectra.

    12. Structure III 12.1 Electronic structure of molecules. Born-Oppenheimer approximation, potential energy curve and potential energy hypersurface, bond dissociation energy. 12.1.1 Valence-bond theory, overlap of AO, hybrid orbitals. 12.1.2 Molecular orbital (MO) theory, LCAO-MO method: bonding and antibonding MO from the point of view of energy and electron density distribution, overlap integral, MO and inversion symmetry. Interaction diagrams of molecules: H2+, H2, He2; HHe+; A2; AB. Occupation of MO. 12.2 Molecular spectra. 12.2.1. Electronic spectra of molecules. Colour, frequency and energy of visible light. Intensities of lines and the Lambert-Beer law. 12.2.2 Vibrational structure. Vibrations of diatomic molecule in harmonic and anharmonic approximation, selection rules. Franck-Condon principle. Types of vibration, normal modes of vibration and their number. 12.2.3 Rotational structure of vibrational spectra. Polyatomic molecules, typical vibrational frequencies of functional groups. 12.2.4 Fate of electronically excited states: fluorescence and phosphorescence.

    13. Further methods of structural analysis 13.1 X-ray diffraction. 13.2 Microscopy.
Literature
    recommended literature
  • ATKINS, P. W. and Julio DE PAULA. Fyzikální chemie. Vyd. 1. Praha: Vysoká škola chemicko-technologická v Praze, 2013, xxvi, 915. ISBN 9788070808306. info
  • JEAN, Yves and François VOLATRON. An introduction to molecular orbitals. Edited by Jeremy K. Burdett. New York: Oxford University Press, 1993, xiv, 337. ISBN 0195069188. info
Teaching methods
lectures, class discussion

In case the COVID-19 measure does not allow contact teaching, the teaching method will be adjusted as follows: teaching will be conducted online in the MS Teams program or through recorded lectures (commented electronic presentations). If interested, the lectures will be supplemented by online consultations.
Assessment methods
The examination with a range corresponding to the syllabus of the subject can be realized in one of two forms: 1) in-class oral or 2) remote oral via MS Teams.

ad 1) Written and oral examination in the extent of the course syllabus. Written part (not electronic) is formed by a test of duration of approx. 15+60 minutes. The open and closed (multiple-choice) tasks are included from both computational and theoretical field. Writing tools only are available. After passing the written part of the exam, the oral part follows, which focuses on the discussion of the test topics.

ad 2) Distance form using online test and subsequent oral examination in MS Teams. The test will be implemented through a ROPOT in the IS MU; duration of approx. 60 minutes. This will be followed by an oral examination in the MS Teams internet application according to the published schedule. Due to the distance form, the examination will not include written preparation and the examinee will directly answer the given questions related to the studied topics.
Language of instruction
Czech
Further Comments
Study Materials
The course is taught annually.
Listed among pre-requisites of other courses
The course is also listed under the following terms Spring 2016, Spring 2017, spring 2018, Spring 2019, Spring 2020, Spring 2021, Spring 2022, Spring 2024, Spring 2025.
  • Enrolment Statistics (Spring 2023, recent)
  • Permalink: https://is.muni.cz/course/sci/spring2023/C3420