Medical Chemistry - requirements


Medical Chemistry

2013/14 Fall Semester

Information on the course and examination

Name of the Department:

Department of Medical Biochemistry, Faculty of Medicine, Semmelweis University


Name of the course: Medical Chemistry

code: AOKOBI001_1A

credits: 6

Director of the course:

Prof. Veronika Ádám M.D., Ph. D., member of Hungarian Academy of Sciences

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 basics of general, organic and inorganic 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. 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. I. Halogenated hydrocarbons, their physicochemical properties.

II. Organic compounds containing hydroxyl groups. Classification. Alcohols,   physical properties, chemical reactions. Enols and phenols, their chemical reactions. Synthesis of ethers, their reactions.

III. 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. IV. 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.

V. Organic compounds containing sulfur: thiols, thiophenols and thioethers, their synthesis and physicochemical properties.

VI. 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.


III. 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.

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.




Requirements for acknowledging the semester


Participation in the laboratory practicals is obligatory; students should sign the attendance sheets at the end of the practicals. In case of more than three absences from the practicals for any reason, the semester will not be acknowledged and the student is not allowed to sit for the semifinal exam. Missed practicals can be completed only in the same week at another group; certificate from the host teacher should be presented by the student to the assigned teacher.


Midterm examinations


Two midterm written examinations will be held on the 6th and 13th week of the semester, respectively, during regular laboratory practicals. Passing the midterms is not a prerequisite to acknowledgement of the semester.

Midterm tests consist of five theoretical questions and five problems (calculations). The material of midterm I covers the subject of lectures given in the first 5 weeks, while midterm II is based on the lecture material of weeks 6-12. Midterm tests will be evaluated by lab teachers and marked as 0, 2, 3, 4 or 5. These ’bonus points’ are added to the scores achieved at the semifinal exam (see below). Retakes are possible; please consult your laboratory teacher on the conditions.


Semifinal exams


Only students who successfully completed the semester, thus obtained an official electronic Neptun signature, are entitled to sit for the semifinal exam.


The semifinal is a written exam that consists of two theoretical parts and a practical exam.


First theoretical part (50 min): drawing 10 structures (both inorganic and organic, 1 point each) and solving five chemical calculations (2 points each). The list of structures to be memorized can be found on the last page of this document. Please note that any inorganic base or salt might be asked that can be formed by combining any cations and anions provided there. Moreover, any normal or branched-chain alkane, alkene or alkyne (up to eight carbon atoms) can be asked such as 2,3-dimethyl-pentane, 3-methyl-2-hexene etc.

Second theoretical part (80 min): 40 multiple choice questions (1 point each).

Practical exam (15 min): an essay question on a laboratory experiment performed during the semester (evaluation: 0, unacceptable; 1 point, minor mistakes; 2 points, clear, detailed and correct). Exact quantities (mass, volume of reagents, incubation times etc.) are not expected here.


The total score is 20 + 40 + 2 = 62.


The exam is unsuccessful with

-          10 or less points in part 1, OR

-          20 or less points in part 2, OR

-          0 point from the practical exam.


Students who pass both part 1 AND part 2 but fail the practical essay have to retake only the practical essay when they repeat the semifinal exam. Those who want a better grade are entitled to rewrite also the first 2 parts, however, with the risk of performing worse.

Students who pass the practical exam but fail either part 1 or part 2 (or both parts) have to retake both theoretical parts but not the practical exam.


In case of successful exams, i. e. when both theoretical units and the practical exam are successfully completed (at least 11, 21 and 1 points are obtained in blocks 1, 2 and the practical essay, respectively), bonus points from the midterms (10 maximum) are added to the total points collected during the exam. Therefore, successful semifinals will be evaluated as follows:


33-39 points = grade 2 (pass)

40-49 points = grade 3 (satisfactory)

50-59 points = grade 4 (good)

60-72 points = grade 5 (excellent).


It is possible to write the practical essay in week 14, in the first 15 minutes of the last laboratory practical of the semester. Students successfully completing this test (getting 1 or 2 points) are exempted from writing the practical exam at the semifinal exam.

It is to note that this is an extra opportunity for passing the practical exam prior to the beginning of the exam period and in case of failure the semifinal exam should proceed as outlined above.




Those students who have passed BOTH midterm examinations with a grade of 3 or better are entitled to participate in the competition. Students intending to participate should register in advance by sending an e-mail to the tutor ( The competition is organized in week 14 (the exact date and venue will be announced later). It is based on the whole material of the semester and has the same structure as the written exam. The winners will be exempted from the semifinal.


Exemption from the semifinal exam

Students who learned general, inorganic and organic chemistry at an academic institution prior to the commencement of their studies at Semmelweis University might sit for an exemption exam that takes place at the beginning of October. Students are kindly asked to present their official documents (transcripts with exam results and a detailed description of the curriculum they completed) to the tutor (Gergely Keszler, EOK building, room 2.132; deadline: 27th September, 2013).

The exemption exam will correspond to parts 1 and 2 of the semifinal (structures, multiple choice tests and calculations).


Registration and modification of examination dates:


electronically, via the Semmelweis University Neptun System.

Retakes are not possible within 3 days following the exam. Only unsuccessful exams can be repeated in the last exam week („extension period”).

All our examination rules comply with the official examination regulations of the Semmelweis University.


Recommended textbooks, manuscripts, handouts:


General chemistry: Ebbing: General Chemistry,  latest edition

Organic chemistry: Hrabák-Csermely-Bauer: Principles of Organic Chemistry (manuscript);

Sasvári: Bioorganic compounds (manuscript)

Inorganic chemistry: Tóth: Concise inorganic chemistry for Medical Students (manuscript)

Laboratory: Hrabák: Selected Collection of Chemical Calculations (manuscript); Medical Chemistry and Biochemistry Laboratory Manual (manuscript)


Manuscripts can be purchased in the bookshops of Semmelweis Publisher (on the ground floor of the NET and EOK buildings).






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.


The 10 structures asked at the semifinal exam will be selected from the following list

Inorganic acids and other 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 formed by the following cations and anions:

Cations: ammonium, sodium, potassium, magnesium, calcium, ferrous, ferric, cuprous, cupric, zinc, silver, aluminium, mercurous, mercury, 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, naphtalene, phenantrene, pyrrole, thiophene, furane, thiazole, oxazole, imidazole, pyrazole, pyridine, pyrane, pyrazine, pyrimidine, purine, indole, pteridine, acridine

Small 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, b-hydroxybutyric acid, pyruvic acid, acetoacetic acid, citric acid, cis-aconitic acid, isocitric acid, a-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, oxime, hydrazone, hydroxamic acid, amide, thiol, sulfinic acid, sulfonic acid, sulfoxide, acyl chloride.



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