Topic list


This list is aimed to help you get prepared for the semifinal exam and serves as a collection of potential questions for midterm examinations.

General chemistry


  1. The electronic structure of the atom. Quantum numbers, orbital filling and Hund's rule. Electronic structure of elements (examples!). Half filled and filled subshells. (Ebbing pp. 276-281, 290-307)
  2. Properties of atoms and the ionic bond. Ionization energy, electronaffinity. The ionic bond. Lattice energies. Nomenclature of ionic compounds. (Ebbing pp. 330-341)
  3. The covalent bond. Covalent bonding. Transition between ionic and covalent bonding. Electronegativity. Molecular geometry of covalent compounds. (Ebbing pp. 341-350, 374-398)
  4. Intramolecular forces of attraction. Interactions betwween polar and nonpolar molecules. The hydrogen bond (e.g. H2O). Requirements for strong hydrogen bonding. (Ebbing pp. 437-444)
  5. Solutions. Solute, solvent. Solubility. The solution process. Solubility of iodine in organic solvents and in water. Lugol solution. Solution of ionic crystals (NaCl) and crystals of polar substances (glucose) in water. Hydrated ions. (Ebbing pp. 481-497).
  6. Enthalpy of solution of solids and gases. Lattice energy and enthalpy of hydration. Enthalpy of solvation. Role of the change of entropy in the solution process. Effects of temperature and pressure on solubility of solids and gases. Henry's law. Bunsen (absorption) coefficient. Calculation of molar concentration of dissolved gases. (Ebbing pp. 485-491)
  7. Vapor pressure of solutions. Raoult's law. Ideal and "real" solutions, vapor pressure depression of solutions of nonvolatile solutes. Mole fraction and molality. Vapor pressure depression of dilute solutions of nonvolatile solutes. (Ebbing pp. 493-494, 497-500)
  8. Solutions of gas in gas. Partial pressure. Composition of air. ppm as concentration unit. Decompression sickness. Arteficial air. (Ebbing pp. 184-185, 192, 199-200, 281, 466-468, 480-481)
  9. Boiling point and freezing point of solutions. Molal freezing point depression and boiling point elevation of aqueous solutions. Colligative properties. Anomalous behavior of ionic solutions,interionic attractions, van't Hoff factor. Formula mass of ionic compounds. Determination of concentration or molar mass by freezing point depression measurements. (Ebbing pp. 500-504)
  10. The phenomenon of osmosis. Osmotic pressure, dependence on temperature, solute concentration and ionic dissociation. Isotonic, hypertonic and hypotonic solutions. Determination of molecular mass or concentration by measuring osmotic pressure. Biological and medical significance of osmosis. (Ebbing pp. 505-509)
  11. Exergonic and endergonic reactions. Definition of the equilibrium constant (Keq). The low of mass action. The range of Keq and pKeq for spontaneous and non-spontaneous reactions. Predicting the direction of a reaction. Le Chatelier principle. Effects of changes in concentration, pressure and temperature for the equilibrium (examples). (Ebbing pp. 621-649).
  12. The ionization of water. Water product, definition of pH and pOH. The pH scale. Calculation of pH for strong electrolytes. Examples of strong acids and bases.
  13. Acid-base theories. The Arrhenius concept. The Bronsted-Lowry concept (examples). The Lewis concept (examples). Acidic strength and the molecular structure. Hydrides, oxiacids (examples).
  14. pH of weak acids and bases. Degree of dissociation and the ionization constant. Definition of pKa and pKb. Conjugated acids and bases. Examples.
  15. Common ion effect: pH dependence of acid/base equilibriums. Effect of strong acid on the ionization of a weak acid. Acid-base indicators used for titration of strong and weak acids. The Henderson-Hasselbalch equation. Effect of pH on the ionization of weak electrolytes (examples: acetic acid, ammonia, amino acids).
  16. Polyprotic acids. pH dependence of their ionization. Examples: carbonic acid and phosphoric acid.
  17. pH of salts. Anion-hydrolysis and cation-hydrolysis (examples). pH of acidic salts (examples: NaHSO4, NaHCO3, NaH2PO4 and NaHPO4).
  18. Buffers. Principle of maintaining a constant pH (examples). Buffer capacity. Comparison of acid and base capacity. The phosphate buffers.
  19. Titration curves. Comparison of strong and weak electrolytes. Titration curves of polyprotic acids. pH of the equivalence points. Buffer ranges.
  20. Buffers of physiological importance. Buffer effect of the phosphate group. The carbonic acid / hydrogen carbonate buffer. The pH-bicarbonate diagram. Effect of stabilization of carbon dioxide and bicarbonate concentration on the buffer capacity. Total acidity of the urine. Respiratory acidosis.
  21. Solubility of salts. The solubility product. Saturated solutions, solubility. Conditions for precipitation. Examples of well soluble and mainly water insoluble compounds.
  22. Complex ions. Complex and double salts. Nature of central ion and ligands. Coordination number. Geometry of complexes. IUPAC nomenclature of complexes. (Ebbing pp. 980-990 + lecture)
  23. Unidentate, bidentate, ambidentate and multidentate ligands in complexes. Chelate complexes: EDTA, EGTA, heme, vitamin B12, calmodulin, proteins, EF hand. Elimination of heavy metal ions from the body. (Ebbing pp. 980-990 + lecture)
  24. Enthalpy. Enthalpy change of chemical reactions (reaction heat). Standard enthalpy of formation. Law of Hess. Thermochemical equations. (Ebbing pp. 213-240)
  25. Combustion heat, heat of neutralization and dissociation. Determination of standard enthalpy offormation by calorimetry and from combustion heat data. Calculation of change of enthalpy from change in internal energy. Vectorial summation of enthalpy changes during combustion. Energy of fats, carbohydrates and proteins. (Ebbing pp. 228-230, 240-241, 744-747)
  26. Enthalpy change of physical processes. Calculation of standard enthalpy change of reactions from molar combustion heat data, and from the standard enthalpy of formation data. Bond dissociation energies of di- and polyatomic molecules. Bond energy. Calculation of enthalpy of formation from bond energy and dissociation energies. (Ebbing pp. 360-363, 747-748 + lecture)
  27. Second law of thermodynamics. Spontaneity of chemical reactions. Relationship between entropy and heat transfer. Unit of entropy. Difference between entropy and energy. (Ebbing pp. 749-752)
  28. Third law of thermodynamics: absolute entropies. Standard entropy values. Determination of standard entropy change of a chemical reaction. Prediction of the direction of entropy change of a chemical reaction. (Ebbing pp. 754-757)
  29. Gibbs free energy (free enthalpy) (G) and its relation to the spontaneity of chemical reactions. Effect of signs of H and S on reaction spontaneity. Standard free energy of formation ( Gof). (Ebbing pp. 757-765, 768-769)
  30. Gibbs free energy and equilibrium. Relationship between equilibrium constant and temperature. Gibbs free energy of concentration gradient. (Ebbing pp. 765-767, 769-770 + lecture)
  31. Spontaneity and speed of chemical reactions. Reaction rate. Rate equation, rate law. Rate constant and its unit, initial rate. (Ebbing pp. 570-578, 581)
  32. Molecularity and order of chemical reactions. Single and multistep reactions. First, pseudo-first, second, third, zero and fractional orders. Rate limiting step. Half-life of chemical reactions. (Ebbing pp. 578-579, 582-590, 597-606)
  33. Reaction rate and temperature. Activation energy. Collision and transition state theories of the mechanism of chemical reactions. Potential energy diagrams. Catalysis. (Ebbing pp. 590-594, 595-596, 606-609)
  34. Specific and equivalent conductance. Determination of the degree of dissociation and the ionization constant by conductometry. (lecture)
  35. Voltaic cells. The Daniell cell. Normal/standard electrode potential. Reduction potential. Electromotive force. (Ebbing 779-795)
  36. Effect of the concentration on the electrode potential. Concentration cells made from cation- and anion electrodes. The hydrogen electrode. Measurement of pH, the glass electrode.
  37. Non-polarizable electrodes. Principle of maintaining constant concentration in reference electrodes. Examples: the calomel electrode and the silver electrode. (lecture)
  38. Gibbs free energy change and the redox reaction. The Nernst equation. Direction of redox reactions, reversibility/ irreversibility. Biologically important redox systems. (Ebbing pp. 795-808)




  1. Alkali metals and their compounds.
  2. The alkaline earth metals and their compounds.
  3. Boron and aluminium family metals.
  4. Allotropes of carbon. CO.
  5. Carbon dioxide. Carbonic acid and its salts. Cyanides.
  6. Silicon and derivatives. Tin and lead and their compounds.
  7. Properties of nitrogen. The nitrogen cycle. Ammonia, hydrazine and hydroxylamine.
  8. Oxides of nitrogen. Oxiacids containing nitrogen. Nitrites and nitrates.
  9. Phosphorus and its compounds: allotropes, oxides, oxiacids, phosphates.
  10. Arsenic, antimony, bismuth and their compounds.
  11. Oxygen and its compounds: allotropes, oxides, peroxides, superoxides.
  12. Properties of water.
  13. Sulfur and its compounds: allotropes, oxides, oxiacids, sulfides, sulfites, sulfates, thiosulfates.
  14. Characteristics of halogens. Fluorine and its compounds.
  15. Chlorine and its compounds.
  16. Bromine, iodine, their compounds.
  17. Hydrogen. Noble gases.
  18. Transition elements. Manganese, iron, cobalt and their compounds.
  19. Copper, zinc, mercury and their compounds. Precious metals.
  20. Air pollution.




  1. The concept of constitution, configuration and conformation in organic chemistry. Constitutional isomerism and stereoisomerism. Chain and position isomerism.
  2. The principal bond types in organic compounds: their characteristic properties and features. Electron polarisation. The tetrahedrical and planar structure of organic compounds: sp3, sp2 hybrid orbitals in carbon containing compounds.
  3. Conformational conditions in open chain (aliphatic) and non aromatic (alicyclic) organic compounds: eclipsed and open forms, periplanar and clinal forms, syn and anti forms.
  4. Cis-trans isomerism in aliphatic and alicyclic compounds.
  5. Optical isomerism. The concept, cause and measurement of optical activity: chiral and achiral compounds. Enantiomers and diastereomers.
  6. Representation of chiral compounds: projective formula. Relative and absolute configuration: the D-L and R-S system.
  7. Delocalisation of bonds (resonance) in organic compounds: structural isomerism and tautomerism.
  8. The principal reaction types (substitution, addition and elimination). Reaction mechanisms (radical, electrophilic and nucleophilic) in organic chemistry.
  9. The structure, properties and reactions of alkenes and cyclic alkenes.
  10. The structure, properties, reactions and several representatives of alkenes and cyclis alkenes.
  11. The aromatic character. The oxidation and reduction of aromatic compounds.
  12. The principal representatives of homo- and heteroaromatic monocyclic and polycyclic compounds.
  13. Substitution type reactions of aromatic compounds. Direction rules in the case of repeated substitution.
  14. The halogen derivatives of hydrocarbons: alkyl- and aryl-halogenides: their synthesis, their role in the preparation of O-, S-, and N-containing organic compounds, some important representatives.
  15. Organic compounds containing -OH groups: alcohols, enols, phenols: their synthesis, grouping principles.
  16. The principal characteristics and reactions of alcohols. The most important representatives of alcohols. Ethers as alcohol derivatives.
  17. The principal characteristics and reactions of phenols: their most important representatives. The -OH derivatives of purine and pyrimidine.
  18. The synthesis and principal characteristics of oxo-compounds (aldehydes, ketones, quinones). The oxidation and reduction of oxo-compounds.
  19. Addition and condensation type reactions of aldehydes and ketones (addition of simple inorganic molecules; dimerisation, polymerisation, aldole formation, acetal formation, formation of ketimines, oximes, hydrazones, Schiff's bases)
  20. The electronic structure of the carboxylate anion, the most important mono-, di- and tricarboxylic acids.
  21. Condensation type reactions of organic acids: ester- and anhydrid-formation, lactones.
  22. Halogen-, hydroxy--, oxo- and amino derivatives of carboxylic acids.
  23. Decarboxylation. The decarboxylated products of amino-, hydroxy- and keto acids.
  24. Organic thio-compounds. Thioalcohols, thioethers, sulfonic acids.
  25. Amino- and imino-derivatives of hydrocarbons: their formation, classification and base character.
  26. The principal reactions of organic amines: acylation, reaction with HNO2, deamination, transamination.
  27. The most important representatives of organic amines in living organisms. The amine derivatives of carbonic acid: carbamoyl-P, urea, guanidine, creatine, barbiturate.