The course is also offered to the students of the fields other than those the course is directly associated with.
Fields of study the course is directly associated with
there are 10 fields of study the course is directly associated with, display
The course objective is to introduce students to a variety of chemical principles, in preparation for more detailed chemistry study in later years. Description of concepts and facts seeks the understanding of chemical world on an atomic and molecular level based on qualitative quantum theory. The molecular approach supports the comprehension of macroscopic phenomena and laws firmly depicted by thermodynamics and chemical kinetics.
1. Principles of chemistry, matter, its properties and forms of matter existence, principal chemical laws, chemical formula, chemical materials, purity of compounds, mixtures, physical and chemical characteristics of pure compounds. 2. Atom symbolics, elemental particles, element, nuclide, isotope, isotone, isobare, mass of elements and molecules, mass unit u, mole, molar mass. Atomic nucleus, mass defect, stability of nuclei, radioactivity, radioactivity law, Fajans-Soddy rules, nuclear reactions. 3. Physical differences between micro- and macro world, particle-wave character of microparticles, dualistic view on matter, Heisenberg principle of uncertainty. Bohr theory of hydrogen atom, emission spectrum of H.- atom, X-ray irradiation, Moseley law. Wave equation, wave function, probability of electron occurrence, atomic orbital, quantum numbers, shapes of atomic orbitals, energy states and their degeneration, Aufbau principle of many electron atom, Pauli principle, Hund principle. 4. Periodical law, primary and secondary periodicity. Properties of elements, ionization energy, electron affinity, electronegativity. Formation of ions, ions with 18 and 20 valence electrons, ionic radii, ionic crystals, methods of their study. 5. Covalent and donor-acceptor bonding, wave-mechanic model of chemical bond, overlap of atomic orbitals, overlap integral, types of molecular orbitals (s, p, d), LCAO-MO, molecular diagrams of homo- and heteronuclear biatomic molecules. Polarity, ionic degree, bond order, length and energy of bond. 6. Shape of molecules, hybridization, VSEPR method, delocalization of electron density, resonance, compounds with lack of electrons, weak interactions between molecules (van der Waals forces, H-bonding). 7. Principles of coordination chemistry, central atom, ligand, coordination polyhedra, chelates, chelate effect, polynuclear complexes, clusters, structural isomerism of complexes. Nomenclature of complex compounds. Donor-acceptor properties of ligands, principles of ligand field theory, octahedral, tetrahedral and tetragonal complexes, high- and low spin complexes, Jahn-Teller effect, spectral and magnetic properties of complexes. Complex equilibria, mechanisms of complexing reactions, trans-effect. 8. State equation and simple laws for ideal gas, transport phenomena in gases, Graham law, real gas, critical state, liquefaction of gases, reduced van der Waals state equation, State equation for liquids, surface tension, viscosity. 9. Common properties of solids, crystal lattice, Madelung constant, Born-Haber cycle, lattice energy, symmetry of molecules and ions. Band theory in electronic structure of solids, properties of metals, metallic bond, conductors, semiconductors, insulators. 10. Types and mechanism of chemical reactions, energy changes in course of chemical reactions, fundamental thermodynamical parameters (U,H,G,S) and laws, Hess law, thermodynamical conditions of spontaneous reaction course. Chemical equilibrium, equilibrium constant, influence of temperature and pressure, Le Chatelier's principle. Reaction kinetics, reaction velocity law, reaction molecularity and order. Arrhenius law, activation energy, reaction coordinate, homogeneous and heterogeneous catalysis. 11. Equilibrium in polyphase system. Gibbs phase rule, definition of phase, component and degree of freedom. Solutions, solubility, concentration units, conductivity of solutions, electrolytical dissociation, solvation and association of ions, ionic strength, activity and activity coefficient. Precipitation and solubility product, properties of diluted solutions, Raoult law, ebullioscopy and cryoscopy, distillation, rectification, sublimation, melting. 12. Acid-base theories, solvotheory of acids and bases, superacid media, acidity and basicity of aqueous solutions, pH, hydrolysis of salts, buffers anf their capacity. 13. Fundamentals of electrochemistry, Faraday law, coulometry, electrochemical potential, types of electrodes, standard electrode potentials, standard hydrogen electrode, Nernst and Nernst-Peters equations, galvanic cells. 14. Absorption of elecrtomagnetic irradiation, function of spectrometer. Molecular spectra, IR and Raman spectrometry, electron spectrometry, luminiscence. Magnetic properties og compounds, magnetic moment of atom and nucleus, dia- and paramagnetism, ferro- and antiferromagnetism. X-ray structural analysis, mass spectrometry.
POLÁK, Rudolf and Rudolf ZAHRADNÍK. Obecná chemie : stručný úvod. Vyd. 1. Praha: Academia, 2000. 224 s. ISBN 8020007946. info
ATKINS, P. W. and Loretta JONES. Chemical principles : the quest for insight. 3rd ed. New York: W.H. Freeman and Company, 2005. 1 sv. ISBN 071675701X. info
ZUMDAHL, Steven S. and Susan A. ZUMDAHL. Chemistry. 6th ed. Boston: Houghton Mifflin Company, 2003. xxiv, 1102. ISBN 0618221565. info
HILL, John William. General chemistry. 4th ed. Upper Saddle River, N.J.: Pearson Prentice Hall, 2005. xxvii, 107. ISBN 0131180037. info
HÁLA, Jiří. Pomůcka ke studiu obecné chemie. 1. vyd. Brno: Masarykova univerzita, 1993. 85 s. ISBN 8021002891. info
KLIKORKA, Jiří, Bohumil HÁJEK and Jiří VOTINSKÝ. Obecná a anorganická chemie [Klikorka, 1989] a. 2. nezměn. vyd. Praha: SNTL - Nakladatelství technické literatury, 1989. 592 s. info
The course is composed of 22 lectures. There are two lectures 2 hr each per week.
The course consists of 22 lectures. The final examination is a written test (2 hrs).