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Medical Chemistry

Faculty of Medicine
6 credit points, end term examination, Csaba Sőti; AOKMBT829_1A

Faculty of Dentistry
4 credit points, end term examination, Csaba Sőti; FOKOMBT304_1A

Fall semesters

Teaching secretary: Gergely Keszler

For details, schedule, Zoom links, downloads, etc., please go to the Moodle pages of the subject.


Lectures and practical lessons
Two lectures are held every week. Practical lab lessons and seminars are held in alternating weeks.

Prerequisites for acknowledging the semester
Participation in the laboratory practicals and seminars is compulsory; students are obliged to sign the attendance sheets at the end of each lesson. In case of more than (i) two missed labs and a missed seminar, or (ii) a missed lab and three absences from seminars, or (iii) five missed seminars for any reason, the semester cannot be acknowledged and the student is not going to be allowed to sit for the end term exam. Missed practicals and/or seminars can be made up only in the same week with another group; certificate of participation issued by the host teacher has to be presented by the student to his/her own teacher.

End term examination
Only those students whose semester is acknowledged by an official electronic Neptun signature are entitled to sit for the end term exam.
The end term examination is an oral exam conducted by a two-member examination committee.
Students take one topic from each of the following groups of topics:
  I. Problems (calculations)
  II. General and inorganic chemistry 1
  III. General and inorganic chemistry 2
  IV. Organic chemistry
  V. Labs

Exemption from attending the course
Students who learned general, inorganic and organic chemistry at university levels prior to the commencement of their studies at Semmelweis University might be exempted from attending the Medical Chemistry course. Students are kindly asked to present their official documents (academical transcripts and a detailed syllabus on the courses they have completed) to the tutor (Gergely Keszler, EOK building, room 2.132).

Registration and modification of examination dates
Electronically, via the Semmelweis University Neptune System.
Retakes are not possible within 3 days following the exam.
All our examination rules comply with the official examination regulations of the Semmelweis University.

Recommended textbooks
General chemistry:
Ebbing-Gammon: General Chemistry, latest edition
Organic chemistry:
Hrabák-Csermely-Bauer: Principles of Organic Chemistry (2nd edition, 2007, editor: A. Hrabák)
Sasvári: Bioorganic compounds
Inorganic chemistry:
Tóth: Concise inorganic chemistry for medical students
Lab lessons:
Hrabák: Laboratory Manual – Medical Chemistry, Biochemistry and Molecular Biology (fourth edition, 2015)
Hrabák: Selected Collection of Chemical Calculations and Biochemical Exercises (latest edition)

Description of the subject
The principal aim of the course is to prepare students for the understanding of Biochemistry and Molecular Biology. This requires a firm knowledge of the foundations of general, organic and inorganic chemistry.
The Medical Chemistry course encompasses the following chapters of Chemistry:

I. General Chemistry
Structure of atoms, ions and molecules. Chemical bonds
Relation of atomic radius, ionization energy, electron affinity and electronegativity to the periodic table. Ionic bond, ion radius, ions. Covalent bonding, s and p bonds, hybrid orbitals, hybridization of carbon. Electron pair repulsion, geometry of molecules, bond angle. Molecular orbital theory.

Polar covalent bonds. Molecules composed of more than two atoms. Coordinative bond. Structure and geometry of ions. Metallic bonding. Interactions between molecules: electrostatic interactions, van der Waals and hydrogen bonds. Structure of water, its properties. Physical states. Types of crystals, characteristic crystal lattices.

Solutions, laws of aqueous solutions, their biological and medical aspects
Solute, solvent, solution. The solution process. Solubility of ions in water, dissociation. Enthalpy of hydration. Concentration, % and molar concentration, normality, molality, molar fraction. Saturated solutions. Solubility, partition, solubility product. Demonstration on calculation problems. Laws of dilute solutions. Vapor pressure, freezing point, boiling point of pure solvents. Vapor pressure of solutions, Raoult’s law. Freezing point depression and boiling point elevation of aqueous solutions. Osmotic pressure, dependence on temperature, solute concentration and ionic dissociation. Biological and medical importance of osmosis.

Electrolytes, degree of dissociation and the ionization constant, their correlation. Conductance of electrolytes, specific and equivalent conductance of strong and weak electrolytes. Acid-base theories. The Arrhenius theory. Classification of acids and bases, their anhydrides. The Bronsted-Lowry concept. The Lewis concept (e.g. coordination compounds). Acidic strenght and the molecular structure. The ionization of water. Water product, definition of pH and pOH. The pH scale. Calculation of pH for strong electrolytes. The effect of strong acids and bases on the ionization of weak acids and bases, respectively. The effect of strong acids and bases on the salts of weak acids and bases. Buffers, calculation of pH of buffers. Buffers of polyprotic acids. Buffers of physiological importance. The carbonic acid/hydrogencarbonate buffer.
Buffer capacity. Acid-base indicators. Titration curves of strong and weak electrolytes. The selection of indicator for titrations. The amphoteric character. Basic and acidic salts. Double salts, complexes. Geometry of complexes, chelates. Reaction of salts with water (hydrolysis).

Redox processes. Oxidation number, its definition. redox equations. The electrode potential, its explanation. Normal and standard potentials. Galvanic cells, Nernst equation. Concentration cells, the principle of electrometric pH measurement. Non-polarizable electrodes, their utilization in practice. Biological redox potential, redox electrodes. The application of redoxi potential for biological processes, the principle of mitochondrial energy production. Electrolysis.

Chemical thermodynamics. Internal energy and enthalpy, reaction heat, standard enthalpy. Hess’ law. Combustion heat, atomic and molecular enthalpy of formation. Bonding energy. The I. and II. laws of thermodynamics, entropy, free energy and free enthalpy. Relation between electromotive force and free enthalpy change. Exergonic and endergonic processes. The equilibrium constant. The direction of the processes and its relation to free energy change.

Chemical kinetics
Reaction kinetics, rate of reaction, order and molecularity. Half-time of reactions. The van’t Hoff rule. Activated complex, transition state, activation energy. The Arrhenius equation. Catalysis, catalysts. Reversible processes, the law of mass action, equilibrium constant and its relation to free energy change. Consecutive reactions, the importance of rate-limiting steps in metabolic processes.

II. Inorganic chemistry
Properties of non-metals
Group of halogens, their biological significance. Oxygen group, oxygen, free radicals containing oxygen, air, air pollution, ozone. Sulfur, its compounds. The nitrogen group. Nitrogen, its important inorganic compounds. Nitrogen cycle. Phosphorus and its compounds. Carbon group, carbon and its important inorganic compounds. The air polluting effect of carbon dioxide. Hydrogen and noble gases. Inorganic compounds of medical importance.

Properties of metals
Alkali metals and their compounds. Alkali earth metals and their compounds, the biological significance of calcium and magnesium. Earth metals. Heavy metals and their biological importance. Precious metals. Medically important metals and metal-containing compounds.

III. Organic chemistry
General properties of organic compounds
Introduction, definition of organic compounds, their composition. Homologous series, constitution, constitution isomerism. Classification according to carbon skeletons and functional groups. Characterization of bondings in organic compounds, bonding energy, distance of atoms, dipole moment. Apolar and polar character, inductive and inductomeric, mezomeric and electromeric effects. The vectorial character of dipole moment. Optical isomerism: structural principles of rotation. Chirality, chiral carbon atoms, configuration, enantiomers. Principle of relative and absolute configuration. Projected formulas. Compounds with more than one chiral center: diastereomerism, mezo-forms. Separation of optical isomers.

Classification of hydrocarbons based on their carbon backbone
Alkanes, cycloalkanes, their homologous series. Steric forms, conformations, conformational isomerism. Physicochemical properties of paraffines. Steric structure of cycloalkanes. Alkenes, their homologous series. Constitutional and configurational isomerism. Chemical properties of alkenes, possible mechanisms of addition reactions. Hydrocarbones containing more double bonds, delocalization of p-electrons in compounds containing conjugated double bonds. Acetylene: physicochemical properties. Aromatic hydrocarbons: homologous series, isomerism. The explanation of the aromatic character by the electronic structure. Chemical behavior of benzene and its homologues. Substitution, oxidation, reduction, direction rules in repeated substitutions. General characterization of heteroaromatic compounds, important heteroaromatic compounds.

Functional groups. Classification and chemical characterization of compounds containing various functional groups
Classification of organic compounds according to their functional groups.

  1. Halogenated hydrocarbons, their physicochemical properties.
  2. Organic compounds containing hydroxyl groups. Classification. Alcohols, physical properties, chemical reactions. Enols and phenols, their chemical reactions. Synthesis of ethers, their reactions.
  3. Oxo compounds: classification, nomenclature, physical properties. Chemical reactions of aldehydes and ketones, nucleophilic addition reactions. Condensation reactions of oxo-compounds, oxidation reduction, substitution on the carbon chain.
  4. Carboxylic acids and their derivatives. Classification, nomenclature, their synthesis, physical properties. The explanation of the acidic character of carboxylic group, the effects of substituents on the acidic character. Chemical reactions of monoprotic carboxylic acids, formation of esters, haloids, amides and anhydrides. Substitution of the carbon chain: synthesis of halogenated, hydroxy-, keto- and amino acids. Acidic character of dicarboxylix acids, important reactions. Chemical reactions of hydroxy- and ketoacids. Important representatives of dicarboxylic, hydroxy- and ketoacids.
  5. Organic compounds containing sulfur: thiols, thiophenols and thioethers, their synthesis and physicochemical properties.
  6. Organic compounds containing nitrogen: classification, physicochemical properties of nitro compounds. Amines, classification, synthesis, basicity. Important chemical reactions of amines (e.g. Schiff base formations). Amides of carbonic acids.

Formulas & Topics - YEAR 2019/20

Essential molecular and structural formulas (.pdf)

I. Problems (calculations) (download them in .pdf)

  1. How many grams of pure, solid NaOH are necessary to prepare 700 ml 16 w/w% solution if the density of the solution is 1.17 g/ml?
  2. How many grams of NaOH are dissolved in 1 liter 30 w/w% solution? The density of the solution is 1.39 g/ml.
  3. How many grams NaCl should be dissolved in 500 g water to prepare a 20 w/w% NaCl solution?
  4. 200 g NaCl have been dissolved in 1 kg water. Calculate the concentration of the solution both in w/w% and w/v% if its density is 1.15 g/ml.
  5. What is the molarity and the molality of the solution prepared by dissolving 15 g NaOH in 400 g water (density of the solution is 1.0 g/ml)?
  6. The molar mass of NaCl is 58.5 g/mol. What is the molality of the solution prepared by dissolving 2 g NaCl in 100 ml water?
  7. What is the molarity of a 28 w/w% KOH solution (d = 1.27 g/ml)?
  8. A 2 molal glycine solution was prepared by dissolving 15 g glycine in 100 ml water. What is the molar mass of glycine?
  9. Calculate the normality of a 30 w/w% KOH (molar mass: 56 g/mol) solution if its density is 1.27 g/ml.
  10. What is the mole fraction of HCl in its 36 w/w% aqueous solution? The molar mass of HCl is 36.5 g/mol and that of water is 18 g/mol.
  11. How many gram sulfuric acid (molar mass: 98 g/mol) are there in 5 liter 0.2 N H2SO4 solution?
  12. Calculate the [H+] concentration of a solution prepared by mixing 40 ml 0.2 N sulfuric acid and 40 ml 0.8 M NaOH. Is this mixture acidic or basic?
  13. 16.6 ml 0.01 N silver nitrate solution were consumed upon titration of 10 ml NaCl solution. What is the concentration of NaCl in w/v%? The molar mass of NaCl is 58.5.
  14. Provide the concentration of a sulfuric acid solution in w/v% if 10 ml were neutralized by 17.5 ml, 0.1 N NaOH.
  15. 10 ml oxalic acid solution [molar mass: 90 g/mol] reacts with 16.6 ml of 0.1 N KMnO4. What is its concentration in w/v%?
  16. 10 ml of an unknown CuSO4 solution react with 8.5 ml 0.02 N EDTA. Calculate the molar concentration of CuSO4.
  17. 8.5 ml 0.1 N HCl have been used to neutralize 80 mg KHCO3 (molar mass: 100 g/mol). Calculate the factor of the HCl solution.
  18. Calculate the freezing point depression of the 1 w/v % aqueous solution of urea (molar mass: 60 g/mol) and NaCl (molar mass: 58.5 g/mol).
  19. Calculate the freezing point depression of an aqueous solution prepared by mixing 50 ml 0.2 M KCl and 50 ml 0.04 M Na2SO4.
  20. Calculate the freezing point depression of a solution prepared by mixing 10 ml 0.1 N HCl and 10 ml 0.1 N NaOH.
  21. Compare the osmotic pressure of a 0.1 w/v% NaCl (molar mass: 58.5 g/mol) and a 0.1 w/v% glucose (molar mass: 180 g/mol) solutions at 0 ℃.
  22. Calculate the osmotic concentration and osmotic pressure of a solution prepared by mixing 100 ml 0.2 M K2SO4 and 100 ml 0.1 M NaCl at 0℃.
  23. Calculate the osmolarity of a solution obtained by mixing 10 ml 0.1 M sulfuric acid and 20 ml 0.2 N sodium hydroxide.
  24. The [H+] concentration of a 0.01 M monoprotic organic acid solution is 10-4. Calculate the degree of dissociation and the dissociation constant of this acid.
  25. The degree of dissociation in a 0.1 M acetic acid solution is 1.3%. At which concentration is α = 90%?
  26. What is the degree of dissociation in a 0.02 N acid solution if Ka = 3 x 10-2?
  27. What is the pH of a 0.01 mM HCl solution?
  28. What is the pH of a mixture of 50 ml 0.45 M sulfuric acid and 50 ml 1 M NaOH?
  29. What is the pH of a 0.01 M NaOH solution?
  30. The degree of dissociation of a weak acid in its 0.2 M solution is 0.1 %. What is the pH?
  31. What is the pH and the degree of dissociation in a 1 mM weak acid solution if its Ka = 1.6 x 10-6?
  32. What is the pH of a 0.035 N organic amine solution if its pKa is 9.6?
  33. What is the pH of a 1 w/v % acetic acid solution? Ka = 2 x 10-5
  34. Calculate the pH of an acetate buffer containing 0.1 M acetic acid and 0.05 M sodium acetate. pKa= 4.7
  35. A buffer is composed of 0.25 M ammonia and 0.5 M NH4Cl. 20 ml 0.2 M HCl are added to 100 ml buffer. Calculate the pH change. pKb = 4.7
  36. How would you prepare 1 liter 50 mM buffer (pH=7.4) using 1 M KH2PO4 and 1 M K2HPO4 stock solutions? pKa = 7.2
  37. 2 g NaOH are dissolved in 1 liter 0.2 M acetic acid. What is the pH if pKa is 4.7?
  38. Calculate the concentration of acetic acid and acetate ions in a 0.2 M acetate buffer (pH=5.0). pKa= 4.7
  39. Calculate the solubility product of silver bromide if its solubility in water is 88 mM.
  40. How many grams Al(OH)3 can be dissolved in 1.5 liter water if its solubility product is 3.7 x 10-15?
  41. Calculate the solubility of PbI2 in water if its solubility product is 9 x 10-9.
  42. How many ml 1 mM NaCl should be added to 10 ml 1 mM AgNO3 to initiate precipitation? The solubility product of AgCl is 1.6 x10-10.
  43. What is the electromotive force of a voltaic cell whose electrodes contain 1 N HCl and 0.02 M MgSO4, respectively? εoMg = – 2.38 V
  44. Calculate the electromotive force of the following voltaic cell (εoMg = – 2.38 V):
    (Pt) H2 / 0.1 N HCl  //  0.001 M MgSO4 / Mg
  45. Calculate the electromotive force for the reaction of a galvanic cell consisting of the following electrodes:
    Cu / 1.6 w/v% CuSO4  //  0.2 M MgSO4 / Mg
    εoMg = – 2.38 V; εoCu = + 0.34 V; molar mass of CuSO4 is 160 g/mol
  46. The electromotive force of the voltaic cell below is 1.65 V. Calculate the Zn2+ concentration. εoAg = + 0.80 V; εoZn = -0.76 V
    (+)  Ag / 1 M AgNO3  //  ZnSO4 / Zn  (-)
  47. Calculate the ZnSO4 concentration in the voltaic cell comprising a Zn2+/Zn and a standard H-electrode if the electromotive force is 0.85 V. The standard potential of Zn is – 0.76 V.
  48. An iron plate is immersed into a CuSO4 solution. In a certain time, the weight of the plate increased by  2 g. How many grams copper was reduced on the iron plate? Atomic masses are 64 for Cu and 56 for Fe.
  49. What is the electromotive force of a hydrogen concentration cell consisting of 0.1 N HCl and 0.01 N acetic acid electrolytes? pKa = 4.7
  50. What is the electromotive force of the following concentration cell:
    – (Pt) / H2 / buffer (pH = 6.5)  //  0.1 M valeric acid / H2 (Pt) +
    The ionization constant of valeric acid is 1.6 x 10-5
  51. Calculate the molarity of acetic acid in the concentration cell below if the electromotive force is 0.15 V. pKa = 4.7
    – (Pt) /H2 /  acetic acid  //  0.001 N HCl  / H2 (Pt) +
  52. The electromotive force of a concentration cell is 0.1 V. The electrodes contain 0.1 N HCl and 0.02 N HCOOH (pKa = 3.7). Calculate the pH and the degree of ionization of the formic acid.
  53. A concentration cell is composed of chlorine electrodes containing 1 N and 0.0001 N HCl. Which electrode is the negative pole? What is the electromotive force?  Set up equations for half-cell reactions. The standard reduction potential of chlorine is + 1.36 V.
  54. A concentration cell contains acetate buffer (1:1) and 0.5 M acetic acid at its hydrogen electrodes, respectively. Calculate the electromotive force. pKa = 4.7
  55. What is the electrode potential of a redox electrode which contains 80 % Fe3+ and 20 % Fe2+ ions in 0.1 M sulfate salt solution? The standard reduction potential is + 0.77 V
  56. Calculate the electromotive force and the free enthalpy change for the reaction of the following cell:
    (Pt) / H2 / C4H4O5 : C4H6O5 (1:1)  //  0.1 M HCl / H2 / (Pt)
    The standard reduction potential of the oxaloacetate/malate electrode is – 0.17 V
  57. Calculate the electromotive force of a redox cell consisting of copper and iron redox electrodes. [Cu2+] = 100 [Cu+]; [Fe2+] = 100 [Fe3+]. Standard reduction potentials are + 0.15 V for the copper and + 0.77 V for the iron redox electrode, respectively. Calculate the free enthalpy change, too.
  58. Calculate the enthalpy of formation of ethanol if its combustion heat is – 1364.4 kJ/mol, the formation enthalpies of CO2 and H2O are – 393.3 kJ/mol and – 285.5 kJ/mol, respectively.
  59. How much is the heat of formation of urea if the enthalpy of formation of ammonia is – 45.8 kJ/mol and the reaction heat is + 119.1 kJ?
    H4N2CO(s) + H2O(l)  = CO2(g) + 2 NH3(g)
  60. How much is the enthalpy of formation of benzene (C6H6) if its combustion heat is – 41.8 kJ/g? The formation enthalpies of CO2 and H2O are – 393.3 kJ/mol and – 285.5 kJ/mol, respectively.
  61. Calculate the enthalpy of formation of acetylene (ethyne) if its combustion heat is – 1300 kJ/mol. The formation enthalpies of CO2 and H2O are – 393.3 kJ/mol and – 285.5 kJ/mol, respectively.
  62. Calculate the standard entropy change of the burning of sulfur if standard entropies are 167.2 J/Kmol for S, 204.8 J/Kmol for oxygen and 248.3 J/ Kmol for SO2. Calculate the Gibbs free enthalpy change, too, if the enthalpy change during the reaction is – 393 kJ/mol.
  63. The equilibrium constant of the reaction below is 120. How much is the standard Gibbs free enthalpy change in this reaction? Is this reaction sponteneous or not?
    2 NH3 + CO2 + H2O  =  (NH4)2CO3
  64. In which direction is the following reaction spontaneous?
    CO + H2O(g)  =  CO2 + H2
    Standard free enthalpies are – 137.1 kJ/mol for CO, – 394 kJ/mol for CO2, – 228 kJ/mol for water and 0 kJ/mol for hydrogen.

II. General and inorganic chemistry 1.

  1. The nuclear and electronic structure of atoms. Quantum numbers and atomic orbitals (only for EM students)
  2. Principles of the periodic table of elements (only for EM students)
  3. Periodic properties of the elements (atomic radius, ionization energy, electron affinity, electronegativity) and the electronic structure of main-group elements (only for EM students)
  4. Formation and classification of ions. The ionic bonding (only for EM students)
  5. Definition and classification of the covalent bonding. Transition between ionic and covalent bonding. Polar covalent bonds, dipole moment (only for EM students)
  6. Intermolecular forces: dipole-dipole forces, London (dispersion) forces, Van der Waals forces, hydrogen bonds
  7. Chemical equilibria. The equilibrium constant. The law of mass action. The LeChatelier principle
  8. The mole concept. Calculation of various concentrations (percentage concentrations, molarity, molality, normality, mole fraction)
  9. Acid-base theories: the Arrhenius, Brønsted-Lowry and Lewis concepts (only for EM students)
  10. Self ionization of water, the pH and pOH of solutions. The pH scale. Calculation of pH for strong acids and bases
  11. pH of weak acids and bases. Degree of dissociation (a) and dissociation constants (Ka and Kb). Definition of pKa and pKb. Acid-base indicators
  12. Specific and equivalent conductance of strong and weak electrolytes (only for EM students)
  13. Titration curves of strong electrolytes. Relative strength of acids and bases. Acidic strength and the molecular structure of hydrogen halides and oxoacids
  14. Titration curves of monoprotic and polyprotic (phosphoric and carbonic acid) weak acids
  15. Principle of maintaining a constant pH (examples). Buffer range and buffer capacity. Comparison of acid and base capacity
  16. pH of buffers: the Henderson-Hasselbalch equation
  17. Buffers of physiological importance: the phosphate and carbonic acid/hydrogen carbonate buffers
  18. pH of salt solutions. Anion hydrolysis (example: acetate) and cation hydrolysis (example: ammonium ion). pH of acidic salts ( NaHSO4, NaHCO3, NaH2PO4 and NaHPO4)
  19. Solubility of salts. The solubility product. Common ion effect
  20. Hydrogen and its inorganic compounds. Noble gases. Air and air pollution
  21. Properties of water
  22. Oxygen and its compounds: allotropes, oxides, peroxides, superoxides
  23. Properties of nitrogen. The nitrogen cycle. Ammonia, hydrazine and hydroxylamine. Oxides of nitrogen. Nitrogen-containing oxiacids. Nitrites and nitrates
  24. Allotropes of carbon. Properties of carbon monoxide and -dioxide, carbonic acid and cyanides

 III. General and inorganic chemistry 2.

  1. Vapour pressure of solutions. Raoult’s law. Ideal and “real” solutions, vapour pressure depression of dilute solutions of nonvolatile solutes.
  2. Gas mixtures. Partial pressure. Composition of air. Ppm as concentration unit. Decompression sickness. Artificial air (only for EM students)
  3. Boiling point and freezing point of solutions. Molal freezing point depression and boiling point elevation of aqueous solutions. Colligative properties.
  4. The phenomenon of osmosis. Osmotic pressure, its dependence on temperature, solute concentration and ionic dissociation. Isotonic, hypertonic and hypotonic solutions. Biological and medical significance of osmosis
  5. The system and its surroundings. Internal energy, mechanical work and reaction heat, the first law of thermodynamics. Enthalpy and the law of Hess. Standard enthalpy change
  6. Enthalpy change of chemical processes (formation and combustion enthalpies). Average bond enthalpy. Energy diagrams of exo- and endothermic processes
  7. Entropy change, spontaneous and reversible processes, the 2nd and 3rd laws of thermodynamics. Absolute and standard entropies
  8. Gibbs’ free enthalpy change, exergonic and endergonic processes. Free enthalpy change under standard and non-standard conditions. Thermodynamic coupling
  9. Spontaneity and velocity of chemical reactions. Reaction rate. Rate equation and rate constant. Collision and transition state theories of the mechanism of chemical reactions (only for EM students)
  10. Molecularity and kinetic order of chemical reactions. Single- and multistep reactions. First, pseudo-first, second and zero-order reactions. Half-life of chemical reactions (only for EM students)
  11. Factors influencing the velocity of chemical reactions: the Arrhenius equation. Catalysis. Enzymes as biocatalysts (only for EM students)
  12. Voltaic cells. Standard electrode potentials and the electromotive force. The Daniell element. Calculation of equilibrium constants from the electromotive force
  13. Dependence of electrode potentials on concentrations: the Nernst equation. Metal and gas electrodes. Concentration cells
  14. Non-polarizable electrodes. Principle of maintaining constant concentration in reference electrodes (calomel and silver/silver chloride electrodes) (only for EM students)
  15. Direction of redox reactions. Redox electrodes; biologically important redox systems
  16. Alkali and alkaline earth metals and their medically important compounds
  17. Phosphorus and its compounds: allotropes, oxides, oxiacids, phosphates
  18. Sulfur and its compounds: allotropes, oxides, oxiacids, sulfides, sulfites, sulfates, thiosulfates
  19. Characteristics of halogens and their compounds
  20. Properties of transition elements, heavy metals and their medically important compounds

IV. Organic chemistry

  1. The central role of carbon in organic chemistry. The hybrid states of carbon, resonance and delocalization in organic compounds
  2. Types of constitution isomerism: branching (backbone) isomerism, positional isomerism and tautomerism with examples
  3. Configuration in organic chemistry: geometric (cis-trans, Z/E) and optical (stereo) isomerism. Optical activity, chirality, stereogeneic (chiral) centers
  4. Relative and absolute configuration, projection rules, the D/L and R/S systems. Enantiomerism, diastereomerism and epimerism. Racemic mixtures and meso compounds. Prochiral compounds
  5. Conformation in organic chemistry
  6. Classification of organic compounds according to the main functional groups. Acid-base character of organic compounds
  7. Major reaction types in organic chemistry: radical, electrophilic and nucleophilic substitution
  8. Major reaction types in organic chemistry: electrophilic and nucleophilic addition
  9. Structure and reactions of alkanes and cycloalkanes: terminology, conformation isomerism, mechanism of radical substitution (only for EM students)
  10. Structure and reactions of alkenes and alkynes: electrophilic addition reactions, (hydro)halogenation, Markownikow’s rule. Dienes, conjugation and resonance. Electrophilic addition of 1,3-butadiene. (only for EM students)
  11. Synthesis and reactions of halogenated hydrocarbons. SN1 and SN2 (only for EM students)
  12. Structure and reactions of mono- and polycyclic homoaromatic compounds. Resonance stabilization and the Hückel’s rule
  13. Structure, classification and physico-chemical properties and of organic hydroxyl (alcohols, enols, phenols) and oxo compounds (aldehydes, ketones).  Formation of ethers and esters
  14. Reactions of hydroxy and oxo compounds: formation of ethers and esters, nucleophilic addition reactions (aldole condensation, formation of hemiacetals/hemiketals, acetals and ketals, Schiff bases)
  15. Structures of the most important mono-, di- and tricarboxylic acids. Ester, lactone, amide and anhydride formation. Decarboxylation of organic acids
  16. Halogen-, hydroxy-, oxo- and amino-derivatives of carboxylic acids
  17. Organic thio-compounds: thioalcohols, disulfides, thioethers, thioesters, sulfinic and sulfonic acids, sulfoxides and sulfones
  18. Nitrogen containing organic compounds: amines and imines, nitro and nitroso derivatives. Terminology and basic character of amines. Principal reactions of organic amines: acylation, deamination, Schiff base formation
  19. Structure and importance of heteroaromatic compounds
  20. The amides of carbonic acid. Structure and biological importance of organic phosphates

 V. Labs

  1. Titration of strong acids with NaOH
  2. Titration of acetic acid with NaOH
  3. Titration of gastric fluid
  4. Principles of the electrometric titration of phosphoric acid and plotting the titration curve
  5. Determination of Cl concentration by means of precipitation titration
  6. Permanganometry: determination of Fe2+ and oxalic acid concentrations
  7. Complexometric titration: determination of unknown Cu2+ concentration
  8. Determination of the ionization constant of acetic acid by conductometry
  9. Spectrophotometric determination of the ionization constant of phenol red
  10. Electrochemistry: measurement of the electromotive force of the Daniell element; the effect of electrolyte concentration on the electromotive force
  11. Electrochemistry: experiments with iron redox electrodes as well as with redox systems of biological relevance

List of compulsory structures
Inorganic acids and other inorganic compounds: sulfuric acid, sulfurous acid, nitric acid, nitrous acid, hydrochloric acid, hydrobromic acid, hypochlorous acid, chlorous acid, chloric acid, perchloric acid, hypobromous acid, bromous acid, bromic acid, perbromic acid, hydrogen cyanide, metaphosphoric acid, orthophosphoric acid, boric acid, carbonic acid, water, ammonia, hydrazine, hydroxylamine, hydrogen peroxide, superoxide anion, pyrophosphate anion, hydrogen sulfide, carbon monoxide, carbon dioxide, nitrous oxide, nitric oxide, sulfur dioxide, sulfur trioxide, hydroxyapatite, fluoroapatite, ferrous ammonium sulfate

Any inorganic salts and bases consisting of  the following cations and anions:
Cations: ammonium, sodium, potassium, magnesium, calcium, ferrous, ferric, cuprous, cupric, zinc, silver, aluminium, mercurous, mercuric, manganese
Anions: hydroxide, oxide, fluoride, chloride, bromide, sulfide, sulfate, sulfite, hydrogen sulfate, thiosulfate, nitrate, nitrite, hypochlorite, chlorite, chlorate, perchlorate, hypobromite, bromite, bromate, perbromate, cyanide, phosphate, monohydrogen phosphate, dihydrogen phosphate, carbonate, hydrogen carbonate (bicarbonate), permanganate, chromate, ferricyanide

Hydrocarbons: alkanes, alkenes and alkynes (up to carbon number 8, both normal- and branched-chain isomers); 1,3-butadiene, 2-methyl-1,3-butadiene (isoprene)

Aromatic rings: benzene, naphthalene, anthracene, phenanthrene, pyrrole, thiophene, furane, thiazole, oxazole, imidazole, pyrazole, pyridine, pyrane, pyrazine, pyrimidine, purine, indole, pteridine, acridine

Simple organic compounds: methanol, ethanol, propanol, isopropanol, n-butanol, ethylene glycol, glycerol, inositol, phenol, diethylether, formaldehyde, acetaldehyde, acetone, mercaptoethanol, aniline, urea, guanidine

Organic acids: formic acid, acetic acid, propionic acid, butyric acid, valeric acid, caproic acid, oxalic acid, malonic acid, succinic acid, glutaric acid, maleic acid, fumaric acid, lactic acid, β-hydroxybutyric acid, pyruvic acid, acetoacetic acid, citric acid, cis-aconitic acid, isocitric acid, α-ketoglutaric acid, malic acid, oxaloacetic acid

Types of bondings and derivatives: ether, phenolether, thioether, ester, lactone, thioester, anhydride (including mixed and phosphoric acid anhydrides), hemiacetale, hemiketale (cyclic forms included), Schiff-base, hydrazone, amide, thiol, sulfinic acid, sulfonic acid, sulfoxide, acyl chloride.