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Medical Chemistry

Faculty of Medicine
6 credit points, terminal examination, Gábor Bánhegyi; AOKOVM002_1A

Faculty of Dentistry
6 credit points, final examination, Csaba Sőti; FOKOOVM005_1A

2018/2019 Fall semester

Teaching secretary: Gergely Keszler


Lectures and practical lessons
Two lectures and a laboratory lesson are held every week; detailed schedules can be found under corresponding tabs.

Prerequisites for acknowledging the semester
(1) Participation in the laboratory practicals is compulsory; students are obliged to sign the attendance sheets at the end of each lab lesson. In case of more than three absences from the labs for any reason, the semester cannnot be acknowledged and the student is not going to be allowed to sit for the semifinal exam. Missed practicals can be completed only in the same week with another group; certificate of participation issued by the host teacher needs to be presented by the student to his/her own teacher.
(2) It is compulsory to pass both midterm examinations.

Midterm examinations
Two midterm written examinations will be held in weeks 6 and 13 of the semester, respectively, during lab lessons.
Midterm tests consist of four theoretical questions (10 points each) and four problems (calculations; 10 points each). The material of midterm I covers that of lectures delivered in the first 5 weeks, while midterm II is based on the lecture material of weeks 6-12. Midterm tests will be marked by your own lab teacher.
Grading of midterms:
  0 – 40 points: 1 (fail)
  41 – 50 points: 2 (pass)
  51 – 60 points: 3 (fair)
  61 – 70 points: 4 (good)
  71  or more points: 5 (excellent)

Passing both midterms is a prerequisite for acknowledgement of the semester. Failed midterms might be retaken twice at your own lab teacher. Retakes cannot be performed later than 14th December.

Semifinal examination
Only those students who have fulfilled both acknowledgement criteria, thus obtained an official electronic Neptun signature, are entitled to sit for the semifinal exam.
The semifinal 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
Students having achieved an average of 4.5 or 5.0 of midterm marks will take only 3 topics from groups II, III and IV.

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 curriculum
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.


Location: EOK (Tűzoltó u. 37-47.) Szent-Györgyi Lecture Hall

Duration: 70 min (Mondays 14:40-15:50 & Wednesdays 08:00-09:10)

Week Date TopicLecturer
1.10-14 SeptAtomic structure; periodic table of elementsCsala
Chemical bonds, hybrid statesCsala
2.17-21 SeptElectrolytes; chemical equilibriaCsala
Acid-base theories Csala
3.24-28 SeptpH of strong acids and basesCsala
Weak electrolytes; chemical conductivitySasvári
4.1-5 OctpH of weak acids and bases; titration curvesSasvári
pH of salts and buffersSasvári
5.8-13 OctPhysiological buffersSasvári
Laws of dilute solutions; solubility and conductivityKeszler
Thermodynamics ICsala
6.15-19 OctThermodynamics IICsala
Thermodynamics IIICsala
7.24-26 OctElectrochemistry ICsala
8.29-31 OctElectrochemistry IICsala
Reaction kinetics Hrabák
9.5-10 NovFoundations of organic chemistryRónai
Nomenclature of organic compoundsRónai
10.12-16 NovConstitution, configuration, conformation IRónai
Constitution, configuration, conformation IIRónai
11.19-23 NovSaturated and unsaturated hydrocarbonsMészáros T.
Reactions af alkyl halides and aromatic compoundsMészáros T.
12.26-30 NovHydroxy compoundsKeszler
Oxo compoundsKeszler
13.3-7 DecOrganic acidsSipeki
Reactions of organic acidsSipeki
14.10-14 DecNitrogen-containing organic compoundsSipeki
Sulfur- and phosphorus-containing organic compoundsSipeki


Location: EOK (Tűzoltó u. 37-47.) 1st floor, corridor ‘D’

1.10-14 SeptIntroduction; Safety rules in the laboratory
2.17-21 SeptAcid-base titration I (titration of strong acids)
3.24-28 SeptAcid-base titration II (titration of weak acids)
4.1-5 OctElectrometric titration of acids, plotting titration curves
5.8-13 OctConsultation (general chemistry)
6.15-19 OctMidterm examination I
7.24-26 OctComplexometry*
8.29-31 OctPermanganometry*
9.5-10 NovDetermination of the ionization constant of phenol red by photometry*
10.12-16 NovElectrochemistry*
11.19-23 NovConductometry*
12.26-30 NovPrecipitation titration
13.3-7 DecMidterm examination II
14.10-14 DecPaper and thin layer chromatography

*The order of marked labs varies from group to group. See the detailed schedule (EM or ED).


EM/1Zámbó Thursday8:00-10:40
EM/5Mészáros T.Thursday8:00-10:40
EM/14Keszler Thursday8:00-10:40
ED/3Szeitner Szerda11:00-13:15
ED/5Kállai Szerda11:00-13:15

Formulas & Topics

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 (f = 0.98) 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. A 0.1 N KOH solution was prepared. Upon determining its factor, 5 ml KOH could be neutralized by 20 ml 0.03 N HCl (f = 1.00). Calculate the factor of the KOH solution.
  16. 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%?
  17. 10 ml of an unknown CuSO4 solution react with 8.5 ml 0.02 N EDTA (f=1.00). Calculate the molar concentration of CuSO4.
  18. 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.
  19. 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).
  20. 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.
  21. Calculate the freezing point depression of a solution prepared by mixing 10 ml 0.1 N HCl and 10 ml 0.1 N NaOH.
  22. 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 ℃.
  23. 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℃.
  24. Calculate the osmolarity of a solution obtained by mixing 10 ml 0.1 M sulfuric acid and 20 ml 0.2 N sodium hydroxide.
  25. 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.
  26. The degree of dissociation in a 0.1 M acetic acid solution is 1.3%. At which concentration is α = 90%?
  27. What is the degree of dissociation in a 0.02 N acid solution if Ka = 3 x 10-2?
  28. What is the pH of a 0.01 mM HCl solution?
  29. What is the pH of a mixture of 50 ml 0.45 M sulfuric acid and 50 ml 1 M NaOH?
  30. 20 ml 10 w/w% sulfuric acid (d = 1.08 g/ml) are diluted to 5 liter. Calculate the pH of the diluted solution.
  31. What is the pH of a 0.01 M NaOH solution?
  32. The degree of dissociation of a weak acid in its 0.2 M solution is 0.1 %. What is the pH?
  33. What is the pH and the degree of dissociation in a 1 mM weak acid solution if its Ka = 1.6 x 10-6?
  34. What is the pH of a 0.035 N organic amine solution if its pKa is 9.6?
  35. What is the pH of a 1 w/v % acetic acid solution? Ka = 2 x 10-5
  36. Calculate the pH of an acetate buffer containing 0.1 M acetic acid and 0.05 M sodium acetate. pKa= 4.7
  37. 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
  38. 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
  39. 2 g NaOH are dissolved in 1 liter 0.2 M acetic acid. What is the pH if pKa is 4.7?
  40. Calculate the concentration of acetic acid and acetate ions in a 0.2 M acetate buffer (pH=5.0). pKa= 4.7
  41. Calculate the solubility product of silver bromide if its solubility in water is 88 mM.
  42. How many grams Al(OH)3 can be dissolved in 1.5 liter water if its solubility product is 3.7 x 10-15?
  43. Calculate the solubility of PbI2 in water if its solubility product is 9 x 10-9.
  44. 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.
  45. 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
  46. Calculate the electromotive force of the following voltaic cell (εoMg = – 2.38 V):
    (Pt) H2 / 0.1 N HCl  //  0.001 M MgSO4 / Mg
  47. 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
  48. 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  (-)
  49. 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.
  50. 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.
  51. 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
  52. 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
  53. 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) +
  54. 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.
  55. 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.
  56. 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
  57. 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
  58. 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
  59. 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.
  60. 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.
  61. 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)
  62. 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.
  63. 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.
  64. 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.
  65. 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
  66. 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
  2. Principles of the periodic table of elements
  3. Periodic properties of the elements (atomic radius, ionization energy, electron affinity, electronegativity) and the electronic structure of main-group elements
  4. Formation and classification of ions. The ionic bonding
  5. Definition and classification of the covalent bonding. Transition between ionic and covalent bonding. Polar covalent bonds, dipole moment
  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
  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
  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. Complex ions. Lewis theory and complex formation. Central ions and ligands, coordination number. Geometry and isomerism of complexes. The IUPAC terminology of complexes
  2. Enthalpy of solution of solids and gases. Lattice energy and enthalpy of hydration. Enthalpy of solvation. Effects of temperature and pressure on solubility of solids and gases. Henry’s law. Bunsen (absorption) coefficient
  3. Vapour pressure of solutions. Raoult’s law. Ideal and “real” solutions, vapour pressure depression of solutions of nonvolatile solutes. Vapour pressure depression of dilute solutions of nonvolatile solutes
  4. Gas mixtures. Partial pressure. Composition of air. Ppm as concentration unit. Decompression sickness. Artificial air
  5. Boiling point and freezing point of solutions. Molal freezing point depression and boiling point elevation of aqueous solutions. Colligative properties. Anomalous behaviour of ionic solutions, interionic attractions, the van’t Hoff factor
  6. 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
  7. 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
  8. Enthalpy change of chemical processes (formation and combustion enthalpies). Average bond enthalpy. Energy diagrams of exo- and endothermic processes
  9. Entropy change, spontaneous and reversible processes, the 2nd and 3rd laws of thermodynamics. Absolute and standard entropies
  10. Gibbs’ free enthalpy change, exergonic and endergonic processes. Free enthalpy change under standard and non-standard conditions. Thermodynamic coupling
  11. Spontaneity and velocity of chemical reactions. Reaction rate. Rate equation and rate constant. Collision and transition state theories of the mechanism of chemical reactions
  12. Molecularity and kinetic order of chemical reactions. Single- and multistep reactions. First, pseudo-first, second and zero-order reactions. Half-life of chemical reactions
  13. Factors influencing the velocity of chemical reactions: the Arrhenius equation. Catalysis. Enzymes as biocatalysts
  14. Voltaic cells. Electrode potentials (reduction potentials) and the electromotive force. The Daniell element. Normal and standard electrode potential. Calculation of equilibrium constants from the electromotive force
  15. Dependence of electrode potentials on concentrations: the Nernst equation. Metal and gas electrodes. Concentration cells
  16. Non-polarizable electrodes. Principle of maintaining constant concentration in reference electrodes (calomel and silver/silver chloride electrodes)
  17. Direction of redox reactions. Redox electrodes; biologically important redox systems
  18. Alkali and alkaline earth metals and their medically important compounds
  19. Phosphorus and its compounds: allotropes, oxides, oxiacids, phosphates
  20. Sulfur and its compounds: allotropes, oxides, oxiacids, sulfides, sulfites, sulfates, thiosulfates
  21. Characteristics of halogens and their compounds
  22. 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. Principles of constitution, configuration and conformation isomerism
  3. Types of constitution isomerism: branching (backbone) isomerism, positional isomerism and tautomerism with examples
  4. Configuration in organic chemistry: geometric (cis-trans, Z/E) and optical (stereo) isomerism. Optical activity, chirality, stereogeneic (chiral) centers
  5. Enantiomerism, diastereomerism and epimerism. Racemic mixtures and meso compounds. Prochiral compounds
  6. Relative and absolute configuration, projection rules, the D/L and R/S systems. Stereospecific numbering
  7. Conformation in organic chemistry
  8. Classification of organic compounds according to the main functional groups. Acid-base character of organic compounds
  9. Major reaction types in organic chemistry: radical, electrophilic and nucleophilic substitution
  10. Major reaction types in organic chemistry: electrophilic and nucleophilic addition; elimination
  11. Structure and reactions of alkanes: terminology, conformation isomerism, mechanism of radical substitution
  12. 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.
  13. Synthesis and reactions of halogenated hydrocarbons. SN1 and SN2
  14. Structure and reactions of mono- and polycyclic homoaromatic compounds. Resonance stabilization and the Hückel’s rule
  15. Mechanism of electrophilic substitution of aromatic compounds. Effect of substituents of the aromatic ring on reaction rates and product formation in further substitution reactions
  16. Classification, structure, physical and chemical properties and reactions of organic hydroxyl compounds (alcohols, enols, phenols). Formation of ethers and esters
  17. Structure, terminology, physico-chemical properties and characteristic reactions of oxo compounds (aldehydes, ketones). Typical nucleophilic addition reactions (aldole condensation, formation of hemiacetals/hemiketals, acetals and ketals, Schiff bases)
  18. Structures of the most important mono-, di- and tricarboxylic acids. Ester, lactone, amide and anhydride formation. Decarboxylation of organic acids
  19. Halogen-, hydroxy-, oxo- and amino-derivatives of carboxylic acids
  20. Organic thio-compounds: thioalcohols, disulfides, thioethers, thioesters, sulfinic and sulfonic acids, sulfoxides and sulfones
  21. 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
  22. Structure and importance of heteroaromatic compounds
  23. The amides of carbonic acid. Structure and biological importance of organic phosphates

 V. Labs

  1. The factor of titrating solutions; factorization of HCl
  2. The factor of titrating solutions; factorization of NaOH
  3. Titration of strong acids with NaOH
  4. Titration of acetic acid with NaOH
  5. Titration of gastric fluid
  6. Principles of the electrometric titration of phosphoric acid and plotting the titration curve
  7. Determination of Cl concentration by means of precipitation titration
  8. Permanganometry: principles, factorization of the titrating solution
  9. Permanganometry: determination of Fe2+ concentration
  10. Complexometric titration: determination of unknown Cu2+ concentration
  11. Complexometric titration: determination of Ca2+ and Mg2+ concentration in the same solution
  12. Determination of the ionization constant of acetic acid by conductometry
  13. Spectrophotometric determination of the ionization constant of phenol red
  14. Electrochemistry: measurement of the electromotive force of the Daniell element; studying the effect of electrolyte concentration on the electromotive force
  15. Electrochemistry: experiments with iron redox electrodes as well as with redox systems of biological relevance
  16. Paper and thin layer chromatography

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.


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