General Chemistry

NAME OF THE COURSE General Chemistry



Year of study


Course teacher

Prof Slobodan Brinić
Prof Zoran Grubač

Credits (ECTS)


Associate teachers

Type of instruction (number of hours)






Status of the course


Percentage of application of e-learning

0 %


Course objectives

To familiarize students with the basic chemical laws and principles and to enable students to master the chemical items that follow General Chemistry. To develop students ability to think critically about the experiments performed in the laboratory and about the involvement of of chemistry in everyday life.

Course enrolment requirements and entry competences required for the course


Learning outcomes expected at the level of the course (4 to 10 learning outcomes)

After the the course students will be able to:
1) Understand the nature and properties of the substance, differentiate elementary substances from compounds, distinguish homogeneous from heterogeneous mixtures, assume procedures for separating mixtures into pure substances.
2) Understanding the problem-solving approach to the balance of substances in chemical changes
3) Understand the structure of atoms and existing models of chemical bonds in such way that they can predict certain properties and reactivity of chemical elements and their ionic and covalent compounds
4) Discern the nature of certain chemical reactions.
5) Adopt the concept of pH, and assume direction of the chemical reactions on the basis of knowledge of chemical kinetics and equilibrium.
6) Independently and safely perform simple chemical experiments

Course content broken down in detail by weekly class schedule (syllabus)

1. Introduction - Natural sciences and chemistry. Units of measurement and measurement. Classification of matter. Pure substance. Decomposition of the substance to the pure substance.
2. Properties of a pure substances, physical and chemical properties. Atom and chemical element. The chemical symbols of elements. The laws of chemical combination by weight and volume. The atomic theoryes from the early ideas to John Dalton. Avogadro’s hypothesis.
3. The discovery of the structure of atoms. The discovery of X-rays and radioactivity. Rutherford model of the atom. X-rays and crystal structure. Bragg equation. Isotopes and the structure of the atomic nucleus.
4. The structure of pure substances. The atomic structure of substances. Types of a crystal systems and crystal characteristics. Cubic crystal system. The molecular structure of substances. The nature of the gas. The nature of the fluid. The concept of temperature. The kinetic theory of gases.
5. Gas laws and the equation of state of an ideal gas. Real gases. Relative atomic and molecular weight. Methods for determining relative atomic (Dulong - Petit method, X-ray diffraction, mass spectrograph) and molecular weight (density of the gas, the method of Victor Mayer, Hoffman method). Periodic table of the elements and the periodic law.
6. Electronic structure of atoms - Bohr model of the atom, quantum numbers. Quantum theory of the electronic structure of atoms. Atomic orbitals.
7. Periodic Classification of elements and the periodic table. Periodic changes in physical properties. Atomic radius. Ionization energy. Electron affinity. Electronegativity.
8. Chemical bonding and molecular structure - Electronic valence theory, ionic and covalent compounds. Electronegativity and degree of oxidation. Writing Lewis structures and the octet rule. Formal charges. Exceptions from the octet rule. VSEPR model and geometry of the molecule.
9. Bond characteristics. Valence bond theory and theory of molecular orbitals.
10. Intermolecular forces. Dipole moment, Van der Waals and London forces, hydrogen bond.
11. The structure and properties of the liquid and solid. Physical properties of solutions. Types of solution. Expression of concentration.
12. The liquid in the liquid solution. Solutions of solids in liquids. Solutions of gases in liquids. Effect of temperature on the solubility. Effect of pressure on the solubility of gases. Colligative properties of solutions: nonelectrolyte and electrolyte solution.
13. Chemical reactions - types of chemical reactions, redox reactions, complex reactions (protolytic reactions and precipitation reactions and dissolution), complex reactions.
14. Chemical kinetics, reaction rate, reaction mechanism, the activation energy. Chemical equilibrium - term equilibrium, chemical equilibrium and chemical equilibrium constant. Factors that affect the chemical equilibrium.
15. Equilibrium in homogeneous and heterogeneous systems. Balance in the electrolyte solutions - equilibrium in solutions of acids and bases , the equilibrium of the complex in solution, the equilibrium between the solution and the insoluble crystals, redox balance
1. The oxidation number: definition, rules for determining in ions and molecules. Examples and training.
2. Nomenclature of Inorganic Chemistry. Names of monoatomic cations and monoatomic anions. Names of poliatomic cations and anion. The names of the ligands. Names of complex ions. Names of oxo acid and their salts.
3. Naming of inorganic compounds - training.
4. Balancing chemical equations, balancing redox equations.
5. Writing redox equations - practice.
6. The stoichiometry: Qualitative and quantitative relationships in chemical reactions. Molar method.
7. Stoichiometry: Quantitative relationships. Yield in chemical reactions and processes: the relevant reactant, the reactant in excess of the theoretical amount of reactants, the theoretical amount of product, yield and loss.
8. The stoichiometry: volume and mass in chemical reactions.
9. Electronic configuration of atoms and ions
10. Lewis structural formula
11. Electronic structural formula
12. Chemical equilibrium in homogeneous and heterogeneous systems
13. Chemical equilibrium in electrolyte solutions.
Exercises :
Exercise 1
The basic rules of laboratory work, safety precautions and protection in the lab, basic laboratory equipment. Washing, cleaning and drying of dishes. Basic laboratory operations, chemicals and dealing with them. Decomposition of the substance to the pure substance. Decomposition of heterogeneous and homogeneous mixture
Exercise 2
Decomposition of the mixture to the pure substance, Decomposition of heterogeneous substances, Sedimentation, decanting, centrifuging, filtering, Buchner funnel, distillation and fractional distillation, sublimation of iodine. Extraction of iodine from aqueous solutions
Exercise 3
Physical and chemical changes, the law of conservation of weight, Gay - Lussac’s law of connected volumes. Exercise with models of unit cells. Determining the relative atomic mass of zinc. Determination of the empirical formula of copper chloride.
Exercise 4
Gas Laws: Determination of the molar volume of oxygen, Boyle’s law, Charles Gay - Lussac’s law, the pressure dependence of the temperature in gases.
Exercise 5
Solutions and their properties. Expressing of concentration. Preparation of the solution with given concentration. Solutions of liquids in liquids. Solutions of gases in liquids. Dependence of solubility on the nature (structure) of the substance. Dependence of solubility on temperature. Dissolution of liquids in liquids. Dissolving gases in liquids. Henry’s law. Determination of molar mass by freezing point depression. Illustration of electrolytic dissociation. Illustration of ions traveling to the electrodes. Electrical conductivity of the solution. Redox - reactions of sulfur and oxygen. Redox reaction of dilute nitric acid solution and iron (II) sulfate. Decomposition and formation reactions of complexes. Ligand substitution reaction. Protolytic reactions (acid- base titration).
Exercise 6
Chemical kinetic, effect of concentration of reactants on the rate of chemical reactions. Effect of temperature on the rate of chemical reactions. The catalytic effect on the rate of chemical reactions. Balance in electrolyte solutions. Moving the chemical balance. Determination of the acid dissociation constant, Ka . Determination of pH: Approximately determination of pH using indicators. Determination of pH using pH sensors. Electrolysis - Determination of Faraday’s constant. Electromotive force of galvanic cells - Daniell cell.

Format of instruction:

Student responsibilities

The 80% presence at lectures and seminars, and completed all laboratory exercises.

Screening student work (name the proportion of ECTS credits for eachactivity so that the total number of ECTS credits is equal to the ECTS value of the course):

Class attendance




Practical training


Experimental work








Seminar essay






Oral exam




Written exam






Grading and evaluating student work in class and at the final exam

Prior to joining the laboratory exercises, students’ knowledge of the material concerned exercises will be verified by tests. All exercises must be completed.
Students who obtain a signature from the course General Chemistry can take the exam. The exam consists of a written and oral examination. The student approached the oral exam must first pass a written examination. The written part of the exam lasts two hours. The written part of the exam is evaluated as follows :
Exactly solved more than 55 % - sufficient
Exactly solved more than 70 % - good
Exactly solved more than 80 % - very good
Exactly solved more than 90 % - excellent
After the written exam on the notice board of the Department will be advertised results of the exam and time when students which did not pass the written exam can view tasks and schedule for oral examinations for students which have acquired this right.
A complete examination or part thereof may be installed through three partial tests during the semester. The tests cover material presented in lectures, seminars and exercises. Written tests are evaluated in the following manner:
Exactly solved more than 55 % - released a written exam
Exactly solved by 60 % - freed written and oral - sufficient
Exactly solved by 70 % - freed written and oral - good
Exactly solved by 80 % - freed written and oral - very good
Exactly solved by 90 % - freed written and oral - excellent
It is necessary to pass all tests in order to pass the exam. Students who did not meet any of the tests must take written and oral exam of that part.

Required literature (available in the library and via other media)


Number of copies in the library

Availability via other media

Filipović, I., Lipanović, S., Opća i anorganska kemija I dio, Školska knjiga, Zagreb, 1995


Brinić, Slobodan. „Recenzirana predavanja iz odabranih poglavlja Opće kemije“ Veljača 2012. KTF-Split. 30.1.2014.


Grubač Z.: „Recenzirana predavanja iz odabranih poglavlja Opće kemije“ Veljača 2012. KTF-Split. 30.1.2014.


Sikirica, M., Stehiometrija, Školska knjiga, Zagreb


Vježbe iz Opće kemije (interna skripta), Kemijsko-tehnološki fakultet, Split, 2013.


Optional literature (at the time of submission of study programme proposal)

Darrell D. Ebbing and Steven D. Gammon, General Chemistry, 9th edition, Houghton Mifflin Company, Boston, 2009.
Raymond Chang, Chemistry, 10th edition, McGraw-Hill, New York, 2010.

Quality assurance methods that ensure the acquisition of exit competences

- Information from interviews, observations, and consultation with students during lectures
- Student survey

Other (as the proposer wishes to add)