CHEMISTRY
Academic Year 2025/2026 - Teacher: VALENTINA SIRACUSAExpected Learning Outcomes
The student will know the main phenomena of basic chemistry and will be able to understand their applications in the engineering field. The student will also learn the ability to understand a scientific text, to rework, even in an original way, what has been studied and to transform and reflect on the knowledge learned.
The student will be able to apply the knowledge acquired to recognize the laws that govern chemical and chemical-physical phenomena, to independently solve complex problems and exercises, even in fields other than chemistry.
The student, using the knowledge of chemical and chemical-physical phenomena acquired during the course, will be able to independently judge the importance of the results obtained, as well as the meaning of the units of measurement and the order of magnitude of the variables that describe the phenomena .
The student who has successfully passed the course will be able to present the knowledge acquired clearly and accurately.
The student who has successfully passed the course will have acquired the ability to study independently and to delve deeper into the concepts studied in full independence.
Course Structure
If the teaching is taught in mixed or remote mode, the necessary variations may be introduced with respect to what was previously declared, in order to respect the planned program and reported in the syllabus.
Required Prerequisites
Attendance of Lessons
Detailed Course Content
1. * Nature of the subject. The matter and the states of aggregation. Homogeneous and heterogeneous systems. Stages and their separations. Elements and chemical compounds. Atoms and molecules. Ponderal laws (Lavoisier, Proust, Dalton). Volumetric laws (Gay-Lussac, Avogadro). Determination of atomic weight (Cannizzaro rule) and molecular weight (gaseous density). Avogadro number. Mole.
2. * Structure of matter. Description of the atom. Protons, neutrons and electrons. Atomic number and mass number. Atomic mass unit. Isotopes. Mass defect. Thomson's experiment. Atomic model of Thomson. Millikan experiment. Rutherford experiment. Atomic model of Rutherford. Electromagnetic radiation. Emission spectrum of the black body. Photoelectric effect. Emission spectrum of the hydrogen atom. Bohr theory. Report by De Broglie. Uncertainty principle of Heisenberg. Wave mechanics. Schrödinger equation. Quantum numbers and energy levels. Orbital. Polyelectronic atoms. Principle of Pauli. Hund rule. Principle of Aufbau. Periodic table. Periodic properties (ionization energy, electronic affinity, atomic radius, electronegativity, metallic character).
3. * Chemical bond. Electron sharing. Covalent bond. Octet rule. Distance and bonding energy. Homeopolar and heteropolar link. Dative tie. Dipoles. Bonds π and s. Hybridization. Bonding angles. VSEPR. Molecular geometry. Resonance. Ionic bond. MO-LCAO theory. Molecular orbitals of diatomic molecules of the second period. Metal tie. Orbitals of Bloch. Weak links. Hydrogen bond.
4. * Chemical compounds and nomenclature. Valence and oxidation number. Oxidation and reduction. Hydrides. Hydrogen acids. Oxides. Peroxides. Hydroxides. Oxyacids. Salts. Chemical equation. Reactions. Redox reactions. Disruption reactions. Combustion reactions. Ponderal relationships. Limiting reagent rule. Calculation examples. Types of formulas (minimum, brute, structure and steric formula). Elementary analysis. Calculation examples.
5. * Thermodynamics. Thermodynamic system. Types of systems. Extensive and intensive variables. Status functions. Work. Heat. Power. Thermal capacity. Work. First principle of thermodynamics. Internal energy and enthalpy. Termochimica. Law of Hess. Second principle of thermodynamics. Heat conversion in work. Entropy. Free energy. Spontaneity of chemical reactions. Third principle of thermodynamics.
6. * States of aggregation of the matter. The gaseous state. Ideal gas and perfect gas. Boyle's law. Law of Gay Lussac. Law of Charles. Avogadro's law. Equations of state of ideal gases. Determination of the molecular weight of a gas. Gaseous diffusion. Partial pressures. Molar heat of gases. Maxwell-Boltzmann speed distribution. Real gases. Van der Waals equation. Liquefaction of gases. Andrews diagram. Numerical exercises. The liquid state. Surface tension. Vapor pressure. Clausius-Clapeyron equation. The solid state. Crystalline and amorphous solids. Isotropy and anisotropy. Primitive cells. Bravais lattices. X-ray diffraction and Bragg equation. Polymorphism. Ionic solids. Covalent solids. Molecular solid. Metallic solids.
7. * State steps and heterogeneous equilibria. State steps: fusion, evaporation, sublimation. Clausius-Clapeyron equation. Variance. Phase rule. State diagrams. One-component systems: water, sulfur, carbon dioxide. Systems with eutectic point.
8. * Status of solution. Types of solution. Solubility of a species. Concentration and way of expressing it. Solute-solvent interaction: ideal and real solutions. Rault law. Relationships between the composition of a mixture of two liquids and that of its vapor. Systems with maximum and minimum azeotrope. Dilute solutions of non-volatile solutes. Colligative properties. Lowering the vapor pressure. Cryoscopic lowering. Ebullioscopic elevation. Osmotic pressure. Numerical exercises.
9. * Chemical equilibria. Law of chemical equilibrium. Le Chatelier's principle. Relationship between free energy and equilibrium constant. Balance constant (Kp and Kc). Relationships between equilibrium constants. Homogeneous and heterogeneous equilibria. Gaseous equilibria. Influence of pressure, temperature and concentration on equilibrium conditions.
10. * Electrolyte solutions. Electrolytic dissociation. Strong and weak electrolytes. Degree of dissociation. Coefficient of Van't Hoff. Conductance. Equivalent conductance. Law of independent ion migration. Acids and bases. Theories of Arrhenius, Bronsted-Lowry and Lewis. Strength of acids and bases. Ionic product of water. Relationship between Ka and Kb. Definition of pH. Calculation of the pH of solution of acids, bases and salts. Buffer solutions. PH indicators. Acid-base titrations. Ampholytic. Solubility equilibria. Solubility product. Ion to common.
11. * Electrochemistry. Oxidation-reduction reactions: electronic ion method. Electrode potentials. Nernst's equation. Standard potential and its measure. Galvanic batteries. Concentration batteries. Electrochemical series of elements. Chemical stacks. Forecasts of redox reactions. Balance constant. Determination of pH, KPS and degree of dissociation. Free reaction energy. Numerical exercises.
12. * Electrolysis. Decomposition voltage. Overvoltage. Faraday laws and numerical exercises. Law of electrochemical equivalents. Electrolysis of molten salts. Water electrolysis. Electrolysis of aqueous solutions. Industrial electrolytic processes. Accumulators. Corrosion. Passivation.
13. * Chemical kinetics. Reaction speed. Kinetic law. Molecularity. Reaction order: reactions of the 1st and 2nd order. Arrhenius equation. Influence of temperature. Activation energy. Catalysts. Kinetic derivation of the equilibrium constant. Chain reactions.
* obligatory skills
- Goal 4: Ensure inclusive and equitable quality education and promote lifelong learning opportunities for all
- Goal 5: Achieve gender equality and empower all women and girls
Textbook Information
1) T.L. Brown, H.E. LeMay, B.E. Bursten, : Fundamentals of Chemistry, fourth Italian edition, EdiSES (or english edition) (new edition)
2) R. Chang – K. Goldsby: “Fundamentals of General Chemistry”, McGrawHill Education - (both new and old editions)
3) P. Atkins – L. Jones: “Fundamentals of General Chemistry”, Zanichelli - (both new and old editions)
Course Planning
| Subjects | Text References | |
|---|---|---|
| 1 | Nature of matter (6 hours) | book 1 (chapters 1, 2); 2, 3 |
| 2 | Structure of the subject and exercises (8 hours) | book 1 (chapters 2, 3, 4, 6, 7); 2, 3 |
| 3 | Chemical bond (9 hours) | book 1 (chapters 8, 9); 2, 3 |
| 4 | Chemical compounds and nomenclature (4 hours) | book 1 (chapters 2, 3), 2, 3 |
| 5 | Thermodynamics and numerical exercises (5 hours) | book 1 (chapters 5, 18); 2, 3 |
| 6 | States of aggregation of matter (10 hours) | book 1 (chapters 10, 11, 12); 2, 3 |
| 7 | State transitions and heterogeneous equilibria (5 hours) | book 1 (chapter 11) ; 2, 3 |
| 8 | Solution state (4 hours) | book 1 (chapter 13); 2, 3 |
| 9 | Chemical equilibria (3 hours) | book 1 (chapter 15); 2, 3 |
| 10 | Electrolyte solutions (5 hours) | book 1 (chapter 16); 2, 3 |
| 11 | book 1 (chapter 16); 2, 3 | book 1 (chapter 19); 2, 3 |
| 12 | Electrolysis (5 hours) | Electrolysis (5 hours) |
| 13 | Chemical kinetics (6 hours) | book 1 (chapter 14); 2, 3 |
Learning Assessment
Learning Assessment Procedures
Two mid-course exams (lasting 90 minutes each) are scheduled, one midway through the course and one at the end of the course. The mid-course exam consists of a written test containing theory questions and numerical exercises with stoichiometric calculations. Each question will be assigned a score, as indicated next to the question. The mid-term tests are considered passed if the score for each is between 9 and 17; otherwise, the student will be required to retake the entire exam on one of the official scheduled test dates. Successful completion of both mid-term tests, with a final score of at least 18/30, entitles the student to be exempted from the final exam. Before the exam begins, the instructor will provide detailed instructions on how to best perform both the mid-term tests and the exam on the official test dates. Reservations for the mid-term tests are mandatory and must be made by the indicated deadline and discussed in class.
The mid-term tests will be administered on the Exam.net platform, using a computer to be brought to class on the scheduled day. If a computer is not available or is accidentally broken, the tests will be taken on paper provided by the instructor.
END-OF-COURSE EXAMS
The end-of-course exam (90 minutes) consists of a written test containing theory questions and numerical exercises using stoichiometric calculations. Each question will be assigned a score indicated next to the question. A pass is considered successful if the score is between 18 and 30 (expressed in thirtieths). Otherwise, the student will be required to retake the entire exam on one of the official scheduled dates. Before the exam, the instructor will provide all the necessary information to ensure the best possible performance. Registration for the exam, for each exam session, is mandatory and must be made exclusively on the student portal (https://studenti.smartedu.unict.it/) by the indicated deadline. Students are free to participate in all scheduled exam sessions during the academic year.
Examples of frequently asked questions and / or exercises
For simplicity's sake, here are some examples of questions:
1. 130 g of Cu are reacted with 530 grams of HNO3 through the following chemical reaction, which must be balanced completely (score = 3):
Cu + HNO3 = Cu(NO3)2 + NO + H2O
Which reagent is in excess?
Calculate the volume in liters of NO, measured at 0°C and 2 atm, that is released from the reaction between Cu and HNO3. Assume ideal gas behavior.
2. The electric charge of an electron is equal to (Score = 1):
☐ -1.602 10-19C ☐ 1.602 10-19C ☐ -1.602 10-24C ☐ -1.602 1023C
3. Isobars are nuclides characterized by the same (Score = 1):
☐ atomic number ☐ type of radioactive decay ☐ mass number ☐ number of neutrons
4. The correct form of Einstein's mass-energy relationship is (Score = 1):
☐ m=Ec2 ☐ E=mc ☐ E2=mc2 ☐ E=mc2
5. The number of protons and electrons in the O- ion is respectively (Score = 1):
☐ 8, 8 ☐ 16, 8 ☐ 16, 16 ☐ 8, 10
6. Balance the following oxidation-reduction reactions (Score = 2)
Zn (s) + HCl (aq) = ZnCl2(aq) + H2 (g)
Al(s) + Cu(NO3)2 (aq) = Cu ( s) + Al(NO3)3 (aq)
7. Balance the following acid-base chemical reactions (Score = 2):
Hg2(NO3)2 + KI Hg2I2 + KNO3
Cl2 + NaOH NaCl + NaClO + H2O