Learning objectives
The student must demonstrate knowledge of the basic concepts of general chemistry and know how to apply them in the resolution of exercises (stoichiometry, balance of reactions, calculations of thermodynamic quantities) as well as in the description of the phenomena studied in the teaching module.
It is required to be able to:
1. Be able to use the scientific language of the topics covered in the chemistry course. (knowledge and understanding)
2. demonstrate an adequate knowledge of the basic laws of general chemistry and to know their application in real cases ( applying knowledge and understanding)
3. express in a concise and precise form the basic concepts of general chemistry in the written test. (Communication skills )
4. to create links between the various chapters dealt with and apply the acquired knowledge to the solution of problems of stoichiometric calculation and to demonstrate the understanding of the basics of chemistry by illustrating some of the basic laws and by the formulation of some examples (making judgments)
5. to integrate the didactic material provided with textbooks is required in order to formulate a synthesis in order to construct a base of preparatory knowledge to face the comprehension of successive courses of chemistry. ( learning skills)
Prerequisites
No prerequisites are required
Course unit content
Classification of the subject. Elements substances and mixtures. Mass conservation laws
Structure of the atom (neutron proton and electron) atomic number, atomic mass. Elements and Isotopes. Orbitals and rules of electron distribution. Periodic table and correlation with the electronic configuration. Periodic properties.
Chemical bond. Ionic and covalent bond. Oxidation number. Valence bond and octet rule. Lewis structures. Resonance formulas and formal charge. Orbitalic hybridization .Polarity in molecular bonds. Nomenclature of cations, anions and inorganic compounds.
Chemical reaction. Formula weight, molecular weight, size and use in the calculation of mass reactions. Stoichiometry. Balancing different types of chemical reactions. Ox-redox reactions.
Introduction to thermodynamics. First and second principles of thermodynamics. Reaction heat. Exothermic and endothermic reactions. Gibbs free energy, spontaneity of a process.
Chemical equilibrium. K equilibrium constant, Le Chậtelier's law. Factors that influence chemical balance.
States of the matter. Gaseous state: perfect gas state equation. Liquid state: vapor pressure, surface tension, boiling point, intermolecular forces (van der Waals forces, hydrogen bond). Notes on the solid state. Phase changes and phase diagrams, supercritical state.
Solutions. Solvent and solute, solubility. Colloidal solutions, emulsions. Methods to indicate the concentration (molarity, molality, normality, percentages). Henry's law. Colligative properties (cryoscopic lowering, ebullioscopic elevation, osmotic pressure).
Acids and Bases. Definition according to Brønsted. Conjugated acid and base. Ionic product of water, pH, calculation of Ka and Kb and their use in determining the strength of an acid or base. Strong acids and bases, weak acids and bases. Examples of hydrolysis of some salts, buffer solutions, pH of a buffer.
Overview of kinetics. Reaction speed, effect of concentration and temperature, Arrhenius's law, activation energy, catalysts (homogeneous, heterogeneous, enzymatic).
Full programme
Introduction to the scientific method. Classification of matter. Elements substances and mixtures. Mass conservation laws (Lavoisier, Proust). Measurement systems for matter and energy.
Structure of the atom (proton, neutron and electron) atomic number, atomic mass. Elements and Isotopes. Orbitals and electron distribution rules. Periodic table and correlation with the electronic configuration. Periodic properties (ionization energy, atomic radius, affinity). Non-metal metals.
Chemical bond. Ionic and covalent bond. Oxidation number. Covalent bond and octet rule and valence evening expansion. Lewis structures. Formulas of resonance and formal charge. Molecular geometries and theory of minimum repulsion (VSEPR). Simple and multiple bonds. Σ and π bonds. Hybridization, hybrid orbitals and molecular geometries. Polar covalent bond, electronegativity, polarity of molecules.
Nomenclature of cations, anions and inorganic compounds
Chemical reaction. Formula weight, molecular weight, mole and use in the calculation of mass reactions. Stoichiometry. Balancing of different types of chemical reactions. Ox-redox reactions. Weight relationships: stoichiometric calculations, limiting reagent, non-quantitative reactions and yields.
Introduction to thermodynamics. First and second law of thermodynamics. Heat of reaction. Exothermic and endothermic reactions. Gibbs free energy.
Chemical equilibrium. Equilibrium constant K. Definition of Kp and K and their dependence. Le Chậtelier's law. Factors that affect the chemical balance. Heterogeneous equilibria.
States of matter. Gaseous state: the laws of Boyle, Gay-Lussac and Charles, perfect gas equation of state. Liquid state: vapor pressure, surface tension, boiling point, intermolecular forces. Perhaps by van der Waals, Hydrogen bond.
State diagrams. Equilibrium between phases. State changes (freezing point and boiling point). One component state diagrams (water and CO2). Heating curves
Solutions. Solvent and solute, solubility. Concentration of solutions (molarity, molality,% by weight and volume, mole fraction). Calculations on dilutions. Henry's Law. Colligative properties (cryoscopic lowering, ebullioscopic raising, osmotic pressure).
Acids and Bases. Definition according to Brønsted. Conjugated acid and base. Ionic product of water, pH, calculation of Ka and Kb and their use in determining the strength of an acid or base. Strong acids and bases, weak acids and bases. Examples of hydrolysis of some salts, buffer solutions, pH of a buffer solution. Calculation of pH for strong acids and bases. Calculation for acids, weak bases and for reactions in solution.
Introduction to kinetics. Rate of reaction, effect of concentration and temperature, catalysts (homogeneous, heterogeneous, enzymatic). Arrhenius law and activation energy, activated complex theory.
Bibliography
Any textbook of General Chemistry for undergraduate students
Teaching methods
The course takes place over 40 hours of lectures and 10 hours of exercises. The lessons will be carried out by explaining the concepts summarized in slides that will be made available to students. During the lessons practical examples will be given regarding the application of fundamental concepts of General Chemistry. The course includes practice lessons on the execution of stoichiometry exercises and on other topics of the course program.
The lessons will be organized in person with the possibility of enjoying the lessons also remotely in synchronous mode (via Teams) and possibly also in asynchronous mode (uploaded on the Elly page of the course). The slides will be available online on the Elly portal (https://elly.saf.unipr.it/2020/) in pdf format for students.
Assessment methods and criteria
The access the exam is obtained by registering through the ESSE3 platform. The assessment procedure will be a written test that comprises ten questions including both the request to illustrate theoretical concepts and the stoichiometry exercises. Each question will be evaluated for a maximum of 3 points. For partial answers or partially completed exercises will be assigned scores from 1 to 2 points. The exam will last for 2 hours. The outcome will be communicated and recorded on the ESSE3 platform only after acceptance of the vote achieved. Students can view the exam by appointment with the teacher.
If it is impossible to carry out the written exam in person due to Covid-19 protocol imposed by the University, the exam will be carried out remotely through an oral interview through Teams.
Other information
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2030 agenda goals for sustainable development
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