C1400 General Chemistry

Faculty of Science
Autumn 2000
Extent and Intensity
4/0/0. 7 credit(s). Type of Completion: zk (examination).
Teacher(s)
prof. RNDr. Jiří Hála, CSc. (lecturer)
Guaranteed by
prof. RNDr. Jiří Hála, CSc.
Chemistry Section – Faculty of Science
Prerequisites (in Czech)
Z
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
The course consists of lectures covering basic principles and concepts of chemistry. It is aimed at students in the teacher- training program, and covers the following topics: subatomic particles and atomic nucleus; electrons in atoms; electron configurations of atoms and periodicity in chemistry; energy in chemistry; covalent bond and molecules; metals and metallic bond; ions and ionic bond; weak bonding interactions; chemical reactions and equilibrium; solutions and solubility; acid-base reactions; oxido-reduction reactions; donor- acceptor bond and complex compounds; gases and liquids; solid state of matter; heterogeneous equilibria.
Syllabus
  • 1. Subatomic particles and atomic nucleus Fundamental particles: leptons, quarks, their properties. Hadrons, mesons, baryons. Laws of conservation of baryon and lepton numbers. Atom: proton, neutron and nucleon numbers. Nuclide, isotope, element. Mass unit, relative nuclide and atomic mass. Determination of atomic masses and isotopic abundances of elements - mass spectrometer, mass spectrum. Atomic nucleus: mass deffect, binding energy, potential curve. Principles and types of radioactivity, existence regions of stable and radioactive nuclides, occurence of radioactive nuclides in the Nature, radioactive elements. 2. Electrons in atoms: Quantum mechanical model: wave function, wave equation, its radial and azimutal parts, probability of finding an electron. Principal, orbital and magnetic quantum numbers. Atomic orbitals, shapes and number of orbitals, boundary surface, nodal plane. Electron spin, spin pairing, spin and spin magnetic quantum numbers. 3. Many-Electron Atoms Effective atomic number, shielding of outer electrons. Energy of atomic orbitals. Electron configuration of atoms, build-up principle, Pauli's principle, Hund's rule, half- and fully occupied shells, energy of 4s and 3d orbitals in atoms of transition elements. Emission spectra, selection rules. Periodic system of elements, types of elements according to their electron configurations, valence shell. Periodicity of properties of elements: atomic radii, ionization energy, electron affinity, electronegativity. X-rays: origin, continuous and line spectrum, Moseley's law, importance of x-rays in chemistry. Magnetic properties of substances. 4. Covalent Bond and Covalent Compounds Principle of covalent bonding and of the theory of molecular orbitals: overlap of atomic orbitals, rules for the construction of molecular orbitals, types of molecular orbitals (s, p, d) and their shapes. The pp - dp bond. Energy of molecular orbitals of diatomic molecules, bond order, its relation to bond length and energy. Ionization energy of molecules, molecular ions, radicals, excitation of molecules. Shapes of polyatomic molecules - the VSEPR method. Polar covalent bond, dipole moment, ionic character of covalent bond. Inductive and mesomeric effects. Hybrid atomic orbitals, origin and properties, hybridization in molecules with multiple bonds. Delocalized p-molecular orbitals, resonance. Optical properties of covalent substances: rotational, vibration-rotational and electronic spectra, chromophores, fluorescence and phosphorescence, charge transfer spectra. 5. Principles of Thermochemistry and Thermodynamics Forms of energy in chemistry; intensive, extensive and state properties. Internal energy, the first law of thermodynamics, work of expansion. Heats of reaction at constant volume and constant pressure, enthalpy, enthalpy of formation, reaction enthalpy. Hess's law and its use, thermochemical cycles of covalent compounds. Entropy as a measure of disorder. Gibbs free energy, spontaneity of chemical reactions. Ellingham diagrams. Chemical equilibrium: thermodynamic condition, dynamic character of chemical equilibrium, equilibrium constant, its relation to Gibbs free energy. Affecting chemical equilibrium: LeChatelier's principle. 6. Ions and Ionic Bond Types of ions, electron configurations of monoatomic ions, surface charge density. Electrostatic theory of ionic bond, lattice energy. Thermodynamics of ionic compounds - the Born - Haber cycle. Polarizability of ions, highly charged ions. 7. Metallic Bond and Metals Band theory of electronic structure of solids, band theory for metals, metallic bond. Electrical and thermal conduction of metals. Intermetallic and interstitial compounds, alloys. 8. The Solid State Crystal lattice, unit cell. Crystal structures of metals and simple ionic compounds (NaCl, CsCl, ZnS), electrical and optical properties of ionic compounds. Crystal structures of covalent compounds - molecular and covalent crystals, electrical conduction of graphite and non-conduction of diamond. Semiconductors. Polarizability of covalent molecules. Allotropy, polymorphy, isomorphy. Clathrates. Liquid crystals. 9. Weak Bonding Interactions van der Waals's interactions - London, dipole - dipole and dipole - induced-dipole interactions. Hydrogen bond: conditions for the formation, inter- and intramolecular hydrogen bond. Ion-dipole interaction: hydration of ions, hydration and solution enthalpy of ionic compounds. 10. Gaseous and Liquid States Kinetic theory of gases, gas pressure, molecular speeds, Maxwell-Boltzman distribution of speeds, mean kinetic energy of molecules in gases. Gas laws, ideal gas law, determination of relative molecular mass and density of gases. Dalton's law of partial pressures. Van der Waals' equation of real gases, Joule-Thomson effect, liquefaction of gases. The liquid state: evaporation, vapor pressure, Clausius-Clapeyron equation, evaporation of mixtures of two liquids, Raoult's law, vapor pressure lowering, behavior of non-ideal mixtures. Miscibility of liquids. 11. Solutions and Solubility Classification of solutions and solvents, concentration units. Osmotic pressure of a solution, osmotic coefficient. Electrolytes and electrolytic dissociation, weak and strong electrolytes, dissociation constant, electrolytic conductance, activity of an electrolyte, activity coefficient, ionic strength. Factors affecting solubility. Sparingly soluble ionic compounds, solubility product. Solubility of gases in liquids, Henry's law. Effect of temperature on solubility. Vapor pressure over solutions. 12. Acid-Base Reactions and Equilibria Brönsted's theory of acids and bases, strength of acids and bases, dissociation constants, polyprotic acids, dissociation degree of a weak acid. Autoionization of water, autoionization constant, the pH scale, pH of solutions of weak acids. Neutralization. Amphiprotic substances. Acid-base indicators. Buffer solutions. Hydrolysis. Acid-base equilibria in non-aqueous solutions, solvotheory of acids and bases. 13. Fundamentals of Electrochemistry Standard potentials of metals, first order electrode, hydrogen electrode. Galvanic cells, the electrochemical series. Oxidation and reduction, standard oxido-reduction potentials, oxidation and reduction agents, oxidoreduction reactions and standard potentials, displacement oxidoreduction reactions. Disproportionation reactions. Gibbs free energy and equilibrium constant of electrochemical reactions. Electrolysis, electrode reactions, electrolysis of molten salts and aqueous solutions, Faraday's laws. 14. Heterogeneous Equilibria Gibbs phase rule, phase diagram of a pure substance, diagrams of water and sulfur, sublimation, metastable states. Phase diagrams of two-component systems: isobaric and isothermic diagram of a mixture of two miscible liquids, fractional distillation, azeotropic mixtures. Phase diagram of the salt - water system, salt hydrates, cooling mixtures. Deliquescence and efflorescence. Phase diagrams of condensed systems. Adsorption and chemisorption. 15. Coordinate Bond and Coordination Compounds Donors, acceptors, central ion, ligand, complex compound. Coordinate bonding. Bonding in complexes of d0 and d10 cations. Ligand field theory - bonding in complexes of d1 - d9 cations. Electron configurations, magnetic and optical properties of complexes with octahedral symmetry, spectrochemical series of ligands. p-donor and p-acceptor complexes, low oxidation states of metals. Complexes in solutions, stability constant, influence of complexes on solubility and oxidoreduction potentials. Chelates, chelate effect. Polymer complex compounds. Lewis and Pearson theories of acids and bases. 16. Rate of Chemical Reactions Molecular-kinetic theory of reaction rate, activated complex, activation energy, reaction profile. Effect of temperature on reaction rate, Arrhenius equation. Rate constant. Reaction order, molecularity of chemical reaction, reaction mechanism. Reversible and irreversible reactions, radical and polymerization reactions. Catalysis: role of catalysts, profile of catalyzed reaction, homogeneous, acid-base, oxidoreduction and heterogeneous catalysis.
Literature
  • Hála, Jíří. Pomůcka ke studiu obecné chemie: MU Brno, 1999.
  • 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
Language of instruction
Czech
Further Comments
The course can also be completed outside the examination period.
The course is taught annually.
The course is taught: every week.
The course is also listed under the following terms Autumn 1999, Autumn 2001.
  • Enrolment Statistics (Autumn 2000, recent)
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