Chemical classification and the atomic nature of matter. Atomic structure and quantum theory. Periodic properties.
The chemical bond. Ionic, metallic and covalent bond.
Chemical reactions. The gases.
Entropy, free energy, spontaneous transformations.
Acids and bases. pH calculation.
Solubility and solubility product.
Redox equilibria. Electrochemical cells. Electrolysis. Chemical kinetics. Stoichiometric calculations. Basic laboratory experiences.
One of the followng text-books (any edition):
Atkinson-Jones "Fondamenti di Chimica generale"
M.S. Silberberg "Chimica"
One of the following text-books for stoichiometry :
Breschi - Massagli "Stechiometria"
Ferri "Calcoli Stechiometrici"
The course aims to provide the basic knowledge of general and inorganic chemistry necessary to deal with successive courses of the chemical and chemical-pharmaceutical area.
Basic competences in mathematics such as, operations, logarithms, exponential notation, the first and second order equations, system of linear equations.
Lectures, exercises in stoichiometry, laboratory practicals.
Teaching materials (slides and excercises) available on the e-learning platform. Meeting and discussion with students are offered by teachers with a large availability during the year, after oral or e-mail requests.
Type of Assessment
During the course, 3 written tests are organized on the subject of lectures carried out up to that moment. Tentatively these are hold in late October, late November and before Christmas or soon after. Each test is composed by stoichiometry exercises and open questions directed to the assessment of General Chemistry concepts. In each test, the student has available 30 points. If, at the end of the 3 in-progress tests the student has accumulated a score equal or higher than 54, he/she is directly admitted to the oral examination at the time of the session. This must be done before the end of the September session, on penalty of forfeiture of bonus obtained. If the student has achieved a lower score, it must take a written test before being admitted to the oral (score 16/30). The written test is of the same type of in-progress tests performed during the course, only it applies to the entire program of the course and not to parts of it, like those hold during the stream of the course. Every year there are three exam sessions with 6 exams in institutional periods (January / February; June / July; September) and one or two extra sessions, usually just before Christmas and immediately after Easter. Within each session the student can attempt the written part twice. A student who had passed the written test but had failed the oral twice, must again pass the written test to be admitted to the oral part.
Chemical classification and the atomic nature of matter. Properties of matter. The states of matter. Homogeneous and heterogeneous systems. Physical quantities and metric systems. The fundamental laws of chemical combinations. The mole concept and Avogadro's number. Atomic weight. The molecules. Molecular weight. Chemical nomenclature: Oxides, hydroxides, acids, salts. Brute formulas and structural formulas.
Atomic structure and quantum theory. Quantization of energy. Planck's constant. The photoelectric effect. Hydrogen spectral series. Corpuscular and wave nature of light. The photon. Corpuscular and wave nature of the electron. The uncertainty principle of Heisenberg. The Schrödinger equation. Wave functions and their interpretation. The eigenfunctions and eigenvalues ??of the hydrogen atom. Symmetry of the hydrogen atom eigenfunctions: s, p, d, f ... First, secondary azimuth spin quantum numbers.
Periodic properties of the elements. The multi-electron atoms and the Aufbau principle. The periodic table. Electronic structure of atoms. Ionization energies. Electron affinity. Electronegativity. The chemical bond. Interactions between two atoms of hydrogen and helium. The covalent bond. Overlap between orbitals and formation of bonds. Symmetry of the overlap between orbitals. Sigma and pi- bonds. Homo- and heteronuclear diatomic molecules. Polyatomic molecules. The model of repulsion of valence shell electronic pairs. Meaning of chemical formulas. Resonance. Hybrid orbitals. Structure-property relationships. Polarity of bonds.
Chemical bonding and aggregation states. The ionic bond. The metals. Coordination compounds. Molecular solids and liquids: Van der Waals Forces. Dipole-dipole interactions. The hydrogen bond. Madelung constant. Types of ionic lattice.
The chemical reactions. Conservation of mass. Stoichiometric calculations. Balances of the reactions. Yield reactions.
Matter in its gaseous state. The ideal gas law. Kinetic theory of gases. Temperature and average kinetic energy. Determination of molecular weights in the gaseous state. Dalton's Law. Molecular weight of gaseous or easily vaporizable substances.
Chemical thermodynamics. The three principles of thermodynamics. Energies at play in chemical reactions. Heat of reaction and Hess law. Enthalpy of reaction. Standard enthalpy of formation. Entropy, free energy and spontaneous transformations. Heat, energy and molecular motions. Entropy and disorder. Criteria for evaluating the entropy changes in a transformation. Standard free energy of formation.
Equilibria involving liquids and solids. Melting, evaporation and sublimation. Vapor pressure. Phase diagrams of water. Solutions and Raoult's law. Colligative properties. Lowering of vapor pressure. Raising and lowering of the boiling temperature of the melting temperature, osmotic pressure. Molecular weight of non-volatile substances. The chemical equilibrium. Spontaneous reactions. Balance, equilibrium constant. Standard free energy. Equilibrium and rate of reaction. Equilibria involving gases, liquids and solids. Homogeneous and heterogeneous equilibria. Equilibrium constant as a function of the partial pressures and the molar fractions. Factors affecting the balance. The principle of Le Chatelier.
Solution equilibria. Content of solutions. Equilibria in aqueous solutions. Ionization of water and pH scale. Strong acids and weak acids. Strong bases and weak bases. Bronsted theory. Hydroxides and acids. Criteria to predict strength of hydracids, hydroxides and oxoacids. Neutralization and acid-base titrations. Buffer solutions. Indicators. PH measurement with indicator method. Hydrolysis. Dissociation of polybasic acids. Acids and bases according to Lewis. pH calculation. Equilibria involving poorly soluble salts. Solubility and solubility product. Simultaneous equilibria: effect of acidity on solubility.
Redox equilibria. Electrochemical cells. The Daniell battery. Electrode potentials. Standard hydrogen electrode. Using the standard potentials to predict the spontaneous sense of a redox reaction. Nernst equation. Frost diagrams. Electromotive force of a cell, standard free energy and equilibrium constant. Electrochemistry. Electrolysis.
Elements of kinetics in chemicstry. Chemical reaction rates and mechanisms. Order of the reaction rate. Reactions of the first order. Stoichiometry and expression of reaction rate. activation energy and Arrenhius diagram. Catalysis and catalysts.
Inorganic Chemistry. General properties of the groups of the periodic table. Industrial processes: preparation of sodium carbonate, sodium hydroxide, chlorine, of hypochlorite, ammonia, nitric acid, sulfuric acid.