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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

2017/2018 Fall semester

Teaching secretary: Gergely Keszler

Intro

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 the 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 with 1, 2, 3, 4 or 5.
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 9th 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 sit for a written exemption exam that takes place in the middle of September. Students are kindly asked to present their official documents (transcripts with exam results and a detailed syllabus on the courses they have completed) to the tutor (Gergely Keszler, EOK building, room 2.132).
Eligible students will be informed on the format of the exemption exam upon registration.

Registration and modification of examination dates
Electronically, via the Semmelweis University Neptun 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 fields 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
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).

Electrochemistry
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.

Thermodynamics
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.

Lectures

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

Duration: 70 min (Wednesdays 9:20-10:30 & Thursdays 16:45-17:55)

Week Date TopicLecturer
1.11-15 Sept.Atomic structure; periodic table of elementsCsala
Chemical bonds, hybrid statesCsala
2.18-22 Sept.Electrolytes; chemical equilibriaCsala
Acid-base theories Csala
3.25-29 Sept.pH of strong acids and basesCsala
Weak electrolytes; chemical conductivitySasvári
4.2-6 Oct.pH of weak acids and bases; titration curvesSasvári
pH of salts and buffersSasvári
5.9-13 Oct.Physiological buffersSasvári
Solubility of salts and gases; laws of dilute solutionsSasvári
6.16-20 Oct.Thermodynamics ICsala
Thermodynamics IICsala
7.24-27 Oct.Thermodynamics IIICsala
Electrochemistry ICsala
8.3 Oct. - 3 Nov.Electrochemistry IICsala
9.6-10 Nov.Reaction kinetics Hrabák
Enzyme kineticsHrabák
10.13-17 Nov.Terminology of organic compoundsRónai
Constitution, configuration, conformationRónai
11.20-24 Nov.Saturated and unsaturated hydrocarbonsRónai
Reactions af alkyl halides and aromatic compoundsMészáros T.
12.27 Nov. - 1 Dec.Hydroxy compoundsKeszler
Oxo compoundsKeszler
13.4-8 Dec.Organic acidsSipeki
Reactions of organic acidsSipeki
14.11-15 Dec.Nitrogen-containing organic compoundsSipeki
Sulfur- and phosphorus-containing organic compoundsSipeki

Labs/Seminars

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

WeekDateLaboratory/Consultation
1.11-15 Sept.Introduction; Safety rules in the laboratory
2.18-22 Sept.Acid-base titration I (titration of strong acids)
3.25-29 Sept.Acid-base titration II (titration of weak acids)
4.02-06 Oct.Electrometric titration of acids, plotting titration curves
5.09-13 Oct.Consultation (general chemistry)
6.16-20 Oct.Midterm examination I
7.24-27 Oct.Complexometry*
8.30 Oct. - 03 Nov.Redox titrations (permanganometry, iodometry)*
9.06-10 Nov.Determination of the ionization constant of phenol red by photometry*
10.13-17 Nov.Electrochemistry*
11.20-24 Nov.Conductometry*
12.27 Nov. - 01 Dec.Determination of the kinetic parameters of urease
13.04 - 08 Dec.Midterm examination II
14.11-15 Dec.Precipitation titration

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

Groups

GroupTeacherDayTime
EM/1StroeThursday8:00 - 10:40
EM/2SzeitnerFriday11:00 - 13:40
EM/3SarnyaiThursday8:00 - 10:40
EM/4BőgelThursday12:50 - 15:30
EM/5CsalaTuesday12:30 - 15:10
EM/6BartaThursday13:50 - 16:30
EM/7HrabákTuesday11:50 - 14:30
EM/8VereczkeiThursday13:40 - 16:20
EM/9MargittaiThursday8:00 - 10:40
EM/10Mészáros T.Wednesday13:40 - 16:20
EM/11SipekiThursday8:00 - 10:40
EM/12ZámbóFriday11:00 - 13:40
EM/13RónaiThursday8:00 - 10:40
EM/14SasváriWednesday10:50 - 13:30
EM/15KeszlerThursday13:50 - 16:30
ED/1KapuyWednesday10:50 - 13:05
ED/2LédecziWednesday10:50 - 13:05
ED/3Simon-Szabó Wednesday10:50 - 13:05
ED/4SpasokukockayaWednesday10:50 - 13:05
ED/5KeszlerTuesday10:00 - 12:15

Formulas & Topics

Essential molecular and structural formulas (.pdf)

I. Problems (calculations)
Problems to be solved in the exam will be selected from the following chapters of the calculation booklet:
   1.1. (except 1.1.6.)
   2.1.2
   3.1.1.
   3.1.2.
   4.1.
   5.1.
   6.1.

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. The Michaelis-Menten model of enzyme kinetics. Initial rate. The Michaelis constant. Maximal velocity. The Lineweaver-Burk plot
  15. Reversible (competitive, non-competitive, uncompetitive) and irreversible inhibition of enzyme activity
  16. Voltaic cells. Electrode potentials (reduction potentials) and the electromotive force. The Daniell cell. Normal and standard electrode potential. Calculation of equilibrium constants from the electromotive force
  17. Dependence of electrode potentials on concentrations: the Nernst equation. Metal and gas electrodes. Concentration cells
  18. Non-polarizable electrodes. Principle of maintaining constant concentration in reference electrodes (calomel and silver/silver chloride electrodes)
  19. Direction of redox reactions. Redox electrodes; biologically important redox systems
  20. Alkali and alkaline earth metals and their medically important compounds
  21. Phosphorus and its compounds: allotropes, oxides, oxiacids, phosphates
  22. Sulfur and its compounds: allotropes, oxides, oxiacids, sulfides, sulfites, sulfates, thiosulfates
  23. Characteristics of halogens and their compounds
  24. 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 rections
  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. Iodometry: principles, factorization of titrating solution
  11. Iodometry: principles, determination of sodium hypochlorite concentration
  12. Complexometric titration: determination of unknown Cu2+ concentration
  13. Complexometric titration: determination of Ca2+ and Mg2+ concentration of the same solution
  14. Determination of the ionization constant of acetic acid by conductometry
  15. Spectrophotometric determination of the ionization constant of phenol red
  16. Electrochemistry: measurement of the electromotive force of the Daniell cell; studying the effect of electrolyte concentration on the electromotive force
  17. Electrochemistry: experiments with iron redox electrodes as well as with redox systems of biological relevance
  18. Determination of the kinetic parameters of urease. Competitive and non-competitive inhibition of urease activity

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, 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, hydrazone, amide, thiol, sulfinic acid, sulfonic acid, sulfoxide, acyl chloride.

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