Fulfillment of the experiment

 Pour 5-6 drops of 0,5N potassium dichromate solution into the tube and by drops add the concentrated alkali solution (NaOH, KOH). Look at the color change. When the solution becomes yellow, add into this one, by drops, concentrated sulphuric acid solution and contemplate the appearance of orange color. In what medium are dichromate ions Cr₂O₇^2- stable, are the chromate ions CrO₄^2- stable?

  Make up a conclusion about the direction of equilibrium shift, using the Le-Shatelie principle.

 

Examination questions

1. The base values of thermodynamics. Internal energy. Work and heat – are the two forms of energy transmission.

2. Classification of thermodynamic systems and processes. Isobaric and isochoric processes. Exothermic and endothermic reactions.

3. The first law of thermodynamics. Enthalpy. Standard heat of formation and combustion of the substances. The Hess^’s law and its consequences.

4. The second law of thermodynamics. Entropy. Gibbs energy.

5.  What tendency does the a) enthalpy factor b) entropy factor express? What function of the system stay gives the quantitative characteristic of simultaneous influence of those both factors? What equation expresses this one?

6. Prediction of spontaneously proceeding processes direction. The role of enthalpy and entropy factors.

7.   The rate of chemical reaction. Average rate for any interval, true rate.

8. Kinetic classification of the reactions. Molecularity and order of the reactions. Half-reaction time period.

9. Relationship of reaction rate and reactants concentrations. Action masses law. Kinetic equations of reactions.

10. Dependence of reaction rate on the time of its proceeding, on the temperature. Temperature coefficient.

11. Activation energy. Arrenius equation. Energy curve of reaction. The role of spatial factor. The theory of transitional stay.

12. Mechanism of chemical reaction. Kinetics of simple and complex reactions. Chain reactions. Photochemical reactions. Photosinthesis.

13. Catalysis. Homogeneous and heterogeneous catalysis. Energy curve of catalytic reaction.

14. Chemical equilibrium. Reversible and irreversible reactions. Kinetic conditions of equilibrium. Equilibrium constant. Equilibrium concentrations.

15. Prediction of the direction of chemical equilibrium shift. Le-Shatelie principal. Stationary state of the organism and its subsystems.

16. Solutions theory. Water and solutions in living systems. Water as universal solvent. Autoprotolysis of water. Autoprotolysis constant. Ion product of water and its physical meaning. Hydrogen ion and hydroxide ion exponents.

17. Thermodynamics of dissolution. Solubility of solid and liquid in liquids. Definition of ideal solution.

18. Solubility of gases in liquids. Henry^,s and Dalton^,s laws. Influence of electrolytes onto the solubility of gases. Sechenov^,s law.

19. Colligative properties of nonelectrolyte solutions. Saturated vapor pressure of solvent above deluted solution of nonelectrolyte. Rauolt^,s law.

20. Consequenses from Rauolt^,s law. Freezing point depression and boiling point elevation of diluted solutions of nonelectrolytes compared with pure solvents.

21. Osmosis, osmotic pressures and their biological role. Vant-Goff^,s law. Determination of osmotic pressure according to cruoscopic dates.

22. Weak electrolytes solutions. Degree and constant of electrolytic dissociation. Ostvald^,s law of dilution. Electrolytes in human organism.

23. Acid – base equilibriums. Equilibriums in weak acid solutions. Equation, used for calculation of [H⁺] and pH in the solutions of weak acids.

24. Equilibriums in weak bases solutions. Equation for calculation of [H⁺] and pH in weak base solutions.

25. Strong electrolytes solutions. Activity, activity coefficient. Ionic strength of the solution.

26. Colligative properties of electrolytes solutions. Isotonic coefficient and its relationship with dissociation degree of weak electrolytes.

27. Osmotic properties of electrolyte solutions. Hypotonic, hypertensive and isotonic solutions. Hemolysis and plasmolysis.

28. Protolytic theory of acids and bases. Dependance of the strength of acids and bases on solvent nature. Protolytic equilibriums and processes.

29. Buffer solutions and their classification. Calculations of [H⁺] and pH for various types of buffers.

30. Mechanism of buffer action, using acetic and ammonium buffers. Buffer capacity.

31. Acid-base equilibriun and buffer systems in human organism.

32. Hydrolysis. Classification. Acidity of the solution as a result of hydrolysis.

33. Hydrolysis degree and constant. Derivation of the expressions for hydrolysis constant of some different types of hydrolysis

34. Relationship of weak acids and bases dissociation constants and hydrolysis constants. Meaning.

35. Heterogeneous equilibriums. Solubility product. The conditions for dissolution and precipitation. Completeness and sequence of the precipitation.

36. Base principles of volumetric analysis. Classification of the methods of analysis. Calculations.

37. Requirements to reactions, used in volumetric analysis. Working solutions, their preparation (with prepared titre and established titre)

38. Neutralization method. Working solutions. Primary substanses, requirements for them. Application in clinical and sanitary-hygienic researchers.

39. Acid-base indicators, mechanism of their activity. Change color interval and titration index. Choice of indicator.

40. Ox-red reactions, their direction. Equivalents of oxidants and reductants. Electron balance method and ion-electron balance or half-reaction method of choosing of the coefficients.

41. Oxidimetry. Permanganatometry, its peculiarities. Equivalent of KMnO₄ in various mediums. Application in clinical sanitary-hygienic researchers.

42. Iodimetry, its peculiarities. Primary substances. Indicator of this method.

43. Coordination theory of Verner, base principles. General and secondary valency, coordination, spatial configuration. Modern view to bonds and structure of complex compounds.

44. Classification, isomery, nomenclature of complexes, their significance in biological processes.

45. Equilibrium in their solutions. Primary and secondary dissociation, unstable and stable constants.

46. Chelatometry and its application in medicine.

47. Quantum-mechanical model of atom structure. Description of energy state of electron with help of the quantum numbers system.

48. The order of electrons distribution in atoms (Pauli^,s principle and Gund^,s rule)

49. Ordinary and exiting states of atoms and ions. Electron configuration of valence electrons, s-, p-, d-, f- elements

50. Atom and ion radius. Ionisation energy and electron affinity. Relative electronegativity.

51. Nature of chemical bond. Length, energy and valence angle of the bond. Geometry of the bond and molecule.

52. Types of the bonds: covalent – formation mechanism (exchanging and donor-acceptor). Properties of covalent bond.

53. Ionic bond – as completely polarized covalent bond, its properties.

54. Metallic bond as particular kind of chemical bond.

55. Hydrogen bond (inter- and inmolecular)

56. Intermolecular actions (dispersion, induction, orientation)

57. Spreading of chemical elements in earth crust. Macro- and microelements. Biogeneous elements. Biosphere, cycle of biogeneous elements.

58. S – elements: electron structure of atoms and cations, biological role of Na,   K,Ca, Mg. Chemical affinity and biological antagonism of these elements (Na-K, Mg-Ca)

59. Chemistry of d-elements: electron structure of atoms and cations, most biologically important elements: Cr, Cu, Mg, Fe, Zn, Co, Mo.

60. Chemistry of p-elements: electron structure of atoms and cations, biogeneous p-elements: C, N, P, O, S, H, halogens.

 

                                                                                                            Appendix

                                                                                                         Table 1.

Concentration units

 

The name	Dimension	Calculating formula
Mass fraction, ω(X)	_	ω = m(X) / m(solution)
Volume fraction, φ(X)	_	φ(X) = V(X) / V(solution)
Mole fraction, χ(X)	_	χ(X) = ν(Xi) / ∑ν(Xi)
Molarity, C(X)	mol/l	C(X)= ν(X) / V(solution)= =m(X) / M(X)^.V(s)
Molar concentration of equivalent, normality,  C(1/z X)	mol/l	C(1/zX)= ν(1/z X) /V(solution)= =m(X) / M(1/z X)^.V(s)
Molality, b(X) or Cm(X)	mol/kg	b(X)= ν(X) ^.1000 / m(solvent)= =m(X)^.1000 / M(X)^.m(solvent)
Titre, t(X)	g/ml	t(X)=m(X) / V(solution); t(X)=ρ(s-n)^.ω(X)

 

 

                                                                                                                             

 

Appendix

                                                                                                                  Table 2