Teaching activities

General Chemistry Lectures

Description Content Requirements Literature

General Chemistry Practice

Description Content Requirements Literature

General Chemistry Practice Supplements

Control Questions Introduction to mathematical expressions

General Chemistry Practice: For Teachers

Advanced physical and polymer chemistry laboratory practical

General Chemistry Lectures Description

The course serves the foundation for further chemistry studies, providing a general introduction to chemistry as a unit of lectures, calculations, and laboratory practices. It helps students with poorer high school chemistry knowledge to catch up, establishes the qualitative and quantitative use of the language of chemistry, empowers students to recognize and interpret simpler chemical phenomena, introduces into inorganic chemistry, physical chemistry, analytical chemistry, and represents an introduction to the description of atoms and molecules.

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General Chemistry Lectures Content

Measurements and units in chemistry. The SI system.

Chemical description of matter, the atomic symbols and chemical formulas.

Physical description of matter, states of matter. Gases, solids and liquids.

Solutions: Composition of solutions. Solubility and saturated solutions. Colloid solutions. Dilute solutions and colligative properties. Ionic solutions.

Chemical equations, stoichiometry: Ionic equations, redox equations.

Thermochemistry: Basic terms in thermochemistry. Heat of reaction and thermochemical equations. Measuring heat of reaction. Hess' law. Enthalpy of formation and the heat of reaction.

Chemical equilibria: General description. Qualitative interpretation of the equilibrium constant. Changing the reaction conditions, LeChatelier's principle. Acid-base equilibria. Acid-base concepts. Self-ionization of water, the pH of a solution. Weak acids and bases. Hydrolysis of salts. Buffers. Multivalent acids and bases. Acid-base titration, indicators. Lewis concept of acids and bases. Complexes and their equilibria. Heterogeneous equilibria.

Electrochemistry: Electrochemical cells, cell reactions. Electrode potential and the standard hydrogen electrode. Nernst equation. Types of electrodes. Electrolysis. Some commercial voltaic cells.

Chemical kinetics: The reaction rate. Dependence of the reaction rate on concentration. Temperature dependence of reaction rates. Catalysis, inhibition. Kinetics of reversible reactions. Reaction mechanism.

Atomic structure: Parts of atoms. Electromagnetic waves. Interaction of atoms and electromagnetic waves. Quantum numbers. Atomic orbitals. Electron configuration and periodicity. Electron structure of atoms. Periodic properties.

Chemical bonding: Ionic bond. Describing covalent bonds. Multiple bonds and delocalized bonding. Coordinate covalent bond, complexes. Metallic bond.

Molecular structure: Intermolecular forces. van der Waals interactions. Hydrogen bonding.

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General Chemistry Lectures Literature:

Árpád Szűcs: General Chemistry (.pdf file)

The lectures, presentations will be uploaded to CooSpace as .pdf files weekly.

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General Chemistry Lectures Requirements

By the end of the semester students should know the basic chemical definitions and relationships. They have to be able to use the language of chemistry, write, read and speak "chemistry". They must be able to understand basic (simple) chemical phenomena, to interpret basic (simple) chemical processes, to see specific areas of chemistry in their context, to obtain synthesized knowledge of the basics of chemistry using both physics and mathematical tools, i.e., to provide quantitative evaluation of chemical events, as well.

Attendance is obligatory/very much advised.

The precondition of taking the exam is the successful completion of General Chemistry Practice.

General Chemistry Lectures: Exam and Topics for the exam

Students have to identify themselves at the exam either with the "Index" or an official ID. Only pen and a normal (not programmable) calculator can be used. Any other aid is illegal!

The exam can be written or oral, which the student must decide when applying for the exam.

In any case it consists of two parts. It starts with calculations (entry). Specifically, this means 3 calculation tasks that could be solved as practiced at the seminars. E.g.,

What is the relative molecular mass of a gas, if 2.46 g of this compound occupies 820 cm³ volume at 35 °C temperature, and at 106.6 kPa pressure? R=8.314 J/(mol K).

Each task is worth 5 points. Getting at least 50 %,( i.e., 7.5 point) means that the student can continue the exam, otherwise the exam is failed.

If a student has chosen oral exam, he/she has to show his/her knowledge in the form of a short presentation by drawing (selecting by chance) two topics from the following list (one from items 1-21 and one from items 22-36). The final mark of the exam is based on the grade of the calculation tasks and the combined grades of the two topics, provided that none of them is failed. If either part is failed, the exam is failed, as well.

The topics of oral exams

1 Gases. Properties of gases, their state transitions. Equation of state for ideal gas, and its versions. The real gases and the van der Waals equation.

2 Solids. Properties of solids, their state transitions. Classification of solids according to their constituents. The crystal systems and their geometric characteristics.

3 Ionic solids. The coordination number. Molecular solids, the allotropes of carbon.

4 Liquids. Properties of liquids, their state transitions. The surface tension and capillarity.

5 Phase diagrams of liquids. Melting point, boiling point, critical point, and triple point.

6 Expression for the composition of solutions and their applications.

7 Solubility of gases in liquids, saturated, oversaturated solutions, immiscible solvents and distribution equilibrium.

8 Characteristics and types of colloids. The structure and application of association colloids.

9 Colligative properties of dilute solutions. Vapor pressure lowering, boiling-point elevation, and freezing-point depression, osmosis.

10 Electrolyte solutions. The specific, the molar, and the ionic molar conductivity.

11 Basic terms in thermochemistry. Energy, work, heat, the internal energy, and enthalpy.

12 Measuring heat. The heat of reactions and Hess's law.

13 General description of chemical equilibria. Various forms of equilibrium constants and their connections.

14 Application of LeChâtelier's principle. The shift in the equilibrium composition by the change in the amount of reactants, in the pressure, and in the temperature.

15 Acid-base concepts. The Arrhenius and the Brønsted-Lowry concept of acids and bases.

16 Weak acids and bases.

17 Hydrolysis of salts.

18 Buffers and multivalent acids and bases.

19 Acid-base titrations. Indicators, titration curves, and how to calculate them.

20 Lewis concepts of acids and bases. Complex compounds and their equilibria.

21 Heterogeneous equilibria. Solubility equilibria.

22 Basic terms in electrochemistry. Electrochemical cells, cell diagrams, cell reactions, half-cell reactions. Cell potential and the electromotive force.

23 The standard hydrogen electrode and the electrode potential.

24 The composition dependence of the electrode potential in various electrode types. Metal electrodes, metal-insoluble salt electrodes, gas electrodes.

25 Comparison of galvanic and electrolytic cells. The batteries.

26 Basic terms in chemical kinetics. The reaction rate, its dependence on the concentration, the order of the reaction. Description of first and second order reactions.

27 Temperature dependence of the reaction rate, catalysis and inhibition.

28 Reaction mechanism. Kinetics of reversible elementary reactions.

29 Interaction between atoms and electromagnetic waves.

30 Quantum mechanics, electron in atoms. The quantum numbers.

31 Atomic orbitals. Their characteristics and description.

32 Electron configuration and periodicity. Periodicity of the chemical properties.

33 Description of ionic, covalent and metallic bondings.

34 Molecular structures. The basics of molecular orbital theory.

35 The basics of valence bond theory. Prediction of molecular structures by the VSEPR model.

36 Intermolecular forces. The van der Waals interactions, and the hydrogen bonding.

If the student has chosen a written exam, he / she will be given a series of 20 questions, which will start with numerical exercises (pop-up). Specifically, this means 3 calculation tasks that could be solved as practiced at the seminars. Each task is worth 5 points. Getting at least 50 %,( i.e., 7.5 point) means that the answer given to the other tasks is taken into account, otherwise the exam is failed.

The rest of the questions are small parts of the topics, including definitions, equations, drawing figures, diagrams, and briefly explaining phenomena.

Illustration of the possible questions of the written exams

1. We put 7.92 g Zn into 200 cm³ hydrochloric acid of 1 mol/dm³ concentration. What volume of hydrogen gas will be evolved at 20 ºC temperature and at 100 kPa pressure? A_r(Zn) = 65.4, R = 8.314 J/(mol K).

2. What volume of chlorine gas (Cl₂) is formed at 101325 Pa and at 273 K, when 20 g of hydrochloric acid (HCl), and excess amount of manganese dioxide (MnO₂) reacts with each other? M_r(HCl) = 36.5, R = 8.314 J/(mol K). Complete the equation: MnO₂ + HCl = MnCl₂ +H₂O + Cl₂

3. What is the rate constant in a first-order reaction if the initial concentration of 0.036 mol/dm³decreases to 0.012 mol/dm³ in 2300 s?

4. What are the derived SI units? Write up five derived physical quantities, give the symbols and the definitions!

5. Describe the structure of graphite! Explain why graphite conducts electric current, and diamond does not?

6. What are the association colloids, and what is the micelle?

7. What is osmosis? What are the isotonic solutions?

8. What are the oxidation numbers of bromine in the following substances? NaBrO₃, HBr, BrO₂, Br₂, BrF₃

9. What is Hess’ law? How would you calculate the reaction enthalpy in the

C(graphite) +1/2 O₂(g) = CO(g) reaction, if the reaction enthalpies of the following reactions are known?

CO(g) + 1/2 O₂(g) = CO₂(g); ΔH_r,1

C(graphite) + O₂(g) = CO₂(g); ΔH_r,2

10. Write up the equilibrium constant for the following reaction, N₂(g) + 3H₂(g) ↔2NH₃(g), with partial pressures and with concentrations. What is the relationship between the two values?

11. What is pH? Dissolving the following salts in water, what would be the pH (equal to, or higher, or lower than that of water), and why? (a) KCl (b) CH₃COONa (c) NH₄NO₃

12. Draw a schematic of the titration curve, when we titrate a weak acid with a strong base! Indicate the pH of the equivalence point! What kind of indicator can be used?

13. Write up the solubility equilibrium of Bi₂S₃, and write up the solubility product! How can we make an oversaturated solution from a saturated Bi₂S₃ solution?

14. What are the cathode and the anode in a lead-acid battery? How can we decide by a density meter, that the battery is completely discharged (flat)?

15. Write up the Arrhenius equation describing the temperature dependence of the rate constant! How can we determine empirically the activation energy?

16. What quantum numbers determine the symmetry of the atomic orbitals, and what are their possible values? Draw a schematic of two kinds of atomic orbitals!

17. What is Pauli’s exclusion principle? Is it possible that ₆C atom has the following orbital diagram?

(↑↑) (↑↑) (↑↑) (↑ )

1s 2s 3s 4s

What is wrong with it?

18. What is a covalent bond? Between which kinds of atoms can it be formed? What is the measure of its strength? Give five examples!

19. According to the VSEPR model, draw the 3D structure and shape of the following molecules! (a) NH₃ (b) XeF₄ (c) ClF₃ (d) HgCl₂ (e) BF₃

20. Write up the empirical and structural formula of water, and hydrogen sulfide! Which one has higher boiling point, and why?

Each completely correct answer is 5 points, each missing or completely wrong answer is 0 points. The actual points given to the questions in proportion to completeness is between the two.

If at least 50% (7.5 points) of the first three questions (calculation tasks) was obtained, the final grade will be calculated on the basis of the total score. Above 50 points the exam is passed (2), above 60 points it is satisfactory (3), above 70 points it is good (4), above 80 points it is excellent (5).

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General Chemistry Lectures Literature

Árpád Szűcs: General Chemistry (.pdf file, password protected)

The lectures, presentations will be uploaded to CooSpace as .pdf files weekly.

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General Chemistry Practice Description

The course provides the basis for further chemistry studies and provides an introduction to chemistry calculations and laboratory practices. It introduces students to the basic laboratory tools, operations, methods of conducting experiments, recording the results of experiments, keeping a laboratory record, and quantifying the observations and results.

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General Chemistry Practice Content

Week	Calculations	Laboratory Practice
1^st week	None	Safety precautions/rules Laboratory rules/Laboratory regulations The necessary safety wares, personal equipments
2^nd week	2 Gases	2.1 Separation from a mixture by dissolution and filtration. Individual 2.2 Comparison of distilled water with tap water. Pair
3^rd week	3 Composition of solutions	3.1 Separation from a mixture by sublimation. Individual 3.2 Overcooling. Pair
4^th week	3 Composition of solutions	4.1 Making a solution. Individual 4.2 Solubility of liquids. Pair 4.3 Solubility of solids. Pair 4.4 Oversaturated solutions. Pair 4.5 Osmosis. Demonstration
5^th week	4 Dilute solutions and colligative properties	5.1 Purification of contaminated alum by recrystallization. Individual 5.2 Determining the molar volume of gases. Pair
6^th week	5 Stoichiometry	6.1 Formation of acids, bases, salts. Pair 6.2 Preparation of a double salt. Individual
7^th week	5 Stoichiometry	7.1 Redox reactions. Pair 7.2 Determining melting point. Individual
8^th week	6 Thermochemistry	8.1 Heat of dissolution. Pair 8.2 Heat capacities. Pair 8.3 Determining boiling point. Individual
9^th week	7 Chemical equilibria	9.1 Studying chemical equilibria. Pair 9.2 Hydrolysis. Pair 9.3 Buffers. Pair
10^th week	7 Chemical equilibria	10.1 Titration of sodium hydroxide. Individual
11^th week	8 Electrochemistry	11.1 Electrochemically explained redox reactions. Pair 11.2 Producing electric current by chemical reaction. Group 11.3 Carrying out chemical reaction by electric current. Group
12^th week	9 Chemical kinetics	12.1 Concentration dependence of the reaction rate. Pair 12.2 Iodine clock reaction. Pair 12.3 Temperature dependence of the reaction rate. Pair 12.4 Catalytic reaction. Pair
13^th week	None	13.2 Spectroscopy, coloring flames. Pair
14^th week	None	Possibility for lab substitution, improvement; Lab closure

(The numbers indicate the chapter numbers in the corresponding books.)

General Chemistry Practice Literature:

Árpád Szűcs: Calculations in General Chemistry (.pdf file)

Árpád Szűcs: Laboratory Practices in General Chemistry (.pdf file)

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General Chemistry Practice Requirements

By the end of the semester students should be able to use basic laboratory equipment. Understand and know how to use basic laboratory methods. Independently perform experiments, independently evaluate experimental results. Design simpler procedures, put together simpler experiments. They need to interpret the simpler chemical events quantitatively using what they have learned in calculation practices.

To check the achievement of expected learning outcomes:

1. Minimum 80% of the practices (i.e., 10) should be completed successfully to acquire the credits for the course. If a student misses or fails more than 2 practices, one of those may be repeated, or resubmitted/improved. If a student misses or fails more than 3 practices (for any reason whatsoever!), no credits are given, and the course should be repeated.

2. Students, who were absent, must show proper excuse (such as an official note from a doctor) when attending the practice the following time. They may have a possibility for a substitution lab which is scheduled for the last week of the semester, but an earlier date may be set up with the instructor. Only one substitution can be done. Missed and not substituted lab, or unexcused absence counts as a failed lab.

3. Students have to arrive on time to the lab well prepared. They have to bring their lab notebook that contains a short summary of the tasks to be carried out and that is ready to enter data, observations, and calculations. This will be checked by the instructor. Unprepared students cannot start the practice and it will count as an unexcused absence.

4. At the very beginning of the practice, a test will be written to check the knowledge of the students about the tasks to be carried out (control questions). Those who passes the test, can start the work in the lab, otherwise (if the knowledge of the student about the practice is insufficient, or the test is not written due to late arrival), the student may not work in the lab, and it counts as a failed lab.

5. Every experiment of the practical course has to be carried out and evaluated. Only those experiments which are fully evaluated, can be considered as successfully completed practices.

6. At the practice, everyone must have personal protective equipment: lab coat, goggles (protective eyeglasses), protective gloves, pipette (filler) ball.

Students who deliberately or severely violate safety rules or use no protective equipment despite warning, will be sent away from the practice and their practice count as failed labs.

7. To the practices, everyone should bring the individual equipment needed for the work: pen, pocket calculator, wiping cloth (for cleaning the table), scissors, self-adhesive label, chemical spoon, matches / lighter (to light, to ignite gas burners). Only those items can be brought into the lab, everything else (bag, coat, telephone, etc.) should be left in the closets in the hallway.

8. The practices last for 4 hours. In the first hour, calculations will be practiced, and the remaining time will be for laboratory practice, for experiments. Leaving the lab for more than 15 minutes requires permission from the instructor. At the end of the lab, the notebook has to be signed by the instructor, thus certifying that the work was finished. If all the tasks are completed ahead of time, the instructor may allow the students to leave earlier.

9. There will be regular tests from calculations, and the lab notebooks will periodically be checked and graded by the instructor.

10. The final grade of the semester will be calculated from the weighted average of three parts: the grades of the control questions (1/6), the grades of the experiments (3/6), and the grades of the calculations (2/6). The final grade is:

5 (excellent), if the average is above 4.50

4 (good), if the average is between 3.76-4.50

3 (satisfactory), if the average is between 2.76-3.75

2 (passed), if the average is between 2.00-2.75

1 (failed), if the average is below 2.00

The course is failed if either the calculation or the practical part, respectively is below 2.00. A student may attempt to improve the final grade once, in the last week of the semester. Only one lab for substitution or improvement is allowed!

11. Students are financially responsible for the laboratory equipment used. Lab-wares obtained at the beginning of the practice, should be given back in the same conditions at the end of the practice (clean and intact), and this should be confirmed by the signature of the technician in the notebook. The instructor or technician should be informed of any broken device/equipment. The way of replacing the device/equipment is determined by the instructor, but the cost is of the student. Missing/lost or broken equipment should be replaced till the end of the semester. It is also a requirement of finishing, completing the course, obtaining the credits.

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