The student will be able:
1. To write chemical formulas based on a given name; to explain the terms: atom, molecule, formula unit; to differentiate ionic and covalent compounds and their chemical formulas.
2. To differentiate and to explain the physical properties of ionic and covalent (molecular) compounds; to describe basic and simple methods of separating compounds.
3. To explain periodicity of atomic size, ionization energy and electron affinity.
4. To describe chemical bonding models; to list and to explain intermolecular forces and their relationship to compound properties.
5. To explain chemical reactions on a molecular level and to describe the quantitative and qualitative meaning of chemical equations; to predict the products of simple chemical reactions and to write reaction equations.
6. To apply stoichiometry and chemical calculations in solving chemical problems and analyzing results; to explain the terms: physical quantities, units of measure, significant digits, unreliable and reliable figures, reaction yield, limiting reactants.
7. To determine the qualitative and quantitative relationships between matter and energy involved in chemical or physical processes.
8. To explain the meaning of dynamic equilibrium, and the Le Chatelier principle as applied in chemical reactions; to differentiate equilibrium in homogeneous and heterogeneous systems.
9. To explain acids and bases according to the Arrhenius, Bronsted-Lowry and Lewis concepts; to describe ionic equilibria in aqueous systems; to prepare buffers and to explain buffer reactions with strong acids and bases.
10. To list and to describe general laboratory safety procedures; to use and to specify the basic laboratory equipment.
1. Pure substances. Homogeneous and heterogeneous mixtures. Physical and chemical changes. Atomic structure. Mass laws.
2. Chemical periodicity. Electron configurations. Atomic size, ionization energy and electron affinity. Chemistry of metals and nonmetals. Noble gases. Halogens.
Alkali and alkaline earth metals.
3. Chemical bonding models. Ionic and covalent bonding. Polar covalent bonds and electronegativity.
4. Lewis structure formulas. Molecular geometry.
5. Thermochemistry. Thermodynamics. Enthalpy and entropy. Free energy and equilibria. Hess law. Born-Haber cycle.
6. Chemical equilibrium. Equilibrium constant and the le Chatelier principle.
7. Structure and properties of substances. Ionic, molecular and covalent crystals. Types of intermolecular forces. Metallic bonding. Physical properties of liquids. Phase diagrams.
8. Gases and kinetic-molecular theory. Ideal gas law. Gas laws. Real gases and van der Waals model.
9. Physical properties of solutions. Rault law. Osmosis. Electrolyte solutions.
10. Acids and bases. Acidic and basic oxides. Ion hydration. Arrhenius and Bronsted-Lowry concept.
11. Acids and bases. Lewis acid-base definition. Indicators.
12. Acid-base buffer systems. Amphoteric substances. Hydrolysis.
13. Ionic equilibria in aqueous systems. Solubility product.
14. Electrolysis. Faraday laws. Voltaic cells. EMF. Commercial Voltaic cells. Iron corrosion.
15. Chemical reactions - kinetics and rates. Catalysis.
1. Introduction to laboratory work.
2. Introduction to measurements in scientific studies and precision calculations.
3. Measurements and precision calculations - determining density, preparing solutions.
4. Physical and chemical changes.
5. Mixture separation - filtration and recrystallization.
6. Boiling point determination.
7. Mixture separation - distillation.
8. Chemical reactions and stoichiometry - fraction crystallization.
9. Volumetric analysis.
10. Acid-base properties - hydrogen chloride preparation, hydrolysis.
11. Chemical equilibrium - buffers and nitrogen dioxide preparation.
12. Gases - molecular mass determination.
13. Gases - equivalent unit and oxygen preparation.
15. Short oral lecture on a given subject and analysis of results.
1. SI units and precision calculations.
2. Number and amount of substances.
3. Empirical and molecular formulas.
4. Mixture compositions - solids.
5. Mixture compositions - solutions.
6. Equations of chemical reactions.
7. Stoichiometry of chemical reactions and limiting reactants.
8. Chemical equivalence - acid-base and redox reactions.
9. Chemical equivalence - precipitation reactions.
10. Chemical equilibrium - solutions.
11. Chemical equilibrium - heterogeneous systems.
12. Gas laws.
13. Physical properties of solutions.
- I. Filipović, S. Lipanović: Opća i anorganska kemija, Školska knjiga, Zagreb, 1997.
- M. S. Silberberg, Chemistry, 3rd ed., McGraw-Hill, New York, 2005.
(Bilo koji suvremni sveučilišni udžbenik Opće kemije izdan nakon godine 2000. napisan za studente kojima je kemija glavni predmet studija).
- M. Sikirica, Stehiometrija, Školska knjiga, Zagreb, 1989.
- M. Sikirica, B. Korpar-Čolig: Praktikum iz opće i anorganske kemije, Školska knjiga, Zagreb, 2001.
- D. Grdenić, Molekule i kristali, 5. izd., Školska knjiga, Zagreb, 2005.