Chemical Sciences

Introduction to chemistry The scientific method. International System of Units (SI), fundamental and derived physical quantities. Symbols, dimensional analysis, tables and plots of physical quantities. Accuracy and errors. Precision and significant figures of an experimental quantity. Elements, atoms, ions. Atomic mass. Atomic mass units (u.m.a.). Compounds, molecules. Molecular mass. Substances. Homogeneous and heterogeneous mixtures of substances. Chemical nomenclature. Chemical reactions. Chemical equations. Periodic table of the elements. Physical and chemical properties of substances. Stoichiometric calculations The concept of mole. Elementary analysis of a substance. Minimal formula and molecular formula (Avogadro's hypothesis and Cannizzaro’s law). Calculation of percentage composition of a compound. Chemical reactions, chemical equations and their balance. Reagents in excess and in defect. Combustion reactions of hydrocarbons. Calculation of the quantitative relations between different substances. Yield of a reaction. Notes on the atomic structure and fundamental theories Dalton's atomic theory. Nuclear model of the atom (Rutherford). Protons. Electrons. Neutrons. Isotopes. Energy of first ionization. Periodic properties of the elements. Line spectra of atoms. Planck's equation, photoelectric effect, energy quantization. Wave-particle dualism. De Broglie equation. Bohr's theory for the hydrogen atom. Behavior of electrons in atoms. Electronic affinity. Concept of electronegativity. Quantum theory and atomic structure. Heisenberg's uncertainty principle (outline). Schrödinger equation (outline). Quantum numbers. Internal electrons and valence electrons. Pauli exclusion principle. Hund's rule. Electronic configurations and periodicity. Atomic orbitals s, p, d, f. Chemical bonding and molecular geometry Ionic bond. Ionic charges and chemical formulas. Transition metal ions. Introduction to the concept of chemical bond through the Lewis formulas. Covalent bond. Octet rule. Expansion of the octet. Covalent coordination bond. Formal charge and real charge of an atom in a molecule. Concept of resonance structures. Bond lengths. Energy of chemical bonds. Concept of electronegativity. Dipole moments. Gillespie model based on the repulsion between electron pairs in valence orbitals (VSEPR) for the prediction of molecular geometries. Concept of electronegativity applied to an atom in a molecule. Binding order. Valence bond theory to determine the molecular structure: hybridization of atomic orbitals (examples of hybrid orbitals: sp, sp2, sp3, sp3d, sp3d2). Molecular orbitals (outline). Chemical reactions Combination, decomposition, single substitution and double substitution reactions. Acids and bases. Acid-base reactions. Traditional nomenclature and according to the rules of the International Union of Pure and Applied Chemistry (IUPAC) of the main inorganic compounds. Concept of oxidation number. Redox reactions (redox). Balancing of redox reactions with the ionic-electronic method in acidic and alkaline environments. Redox reactions and chemical analysis Stoichiometric calculations in solution Solubility rules in water for inorganic compounds. Saturated solution concept. Electrolytes and non-electrolytes. Strong and weak electrolytes. Dissociation reactions, concept of degree of dissociation. Density of solvent and solutions. Concentration of a solute: molarity, percentage by weight. Notes on the conductivity of aqueous solutions of electrolytes. Concentrated solutions: relationship between chemical activity and concentration. States of aggregation of matter The gaseous state Properties of gases (pressure and temperature). Derivation of the equation of state for gases with ideal behavior obtained by combining Boyle's and Charles's laws, Gay-Lussac's law of combining volumes and Avogadro's hypothesis. Molar volume of gases. Gas density. Dalton's law of partial pressures. Kinetic theory of ideal gases. Temperature and kinetic energy. Maxwell's molecular velocity distribution. Van der Waals equation to describe the behavior of real gases. The liquid state and solutions Processes of melting and evaporation. Molar enthalpies of fusion, evaporation, sublimation. Molecular dipoles. Electronic polarizability. Intermolecular bonds: dipole-dipole interactions and London forces (van der Waals forces); hydrogen bond, ion-dipole interactions. Vapor pressure of a liquid. Viscosity, surface tension, capillarity, polarity (hints). State diagrams (water, carbon dioxide). State changes: Clausius-Clapeyron equation. Concentration of solutions: mole fraction and molality. Definition of the van’t Hoff binomial. Concept of ideal solutions: Raoult's law. Colligative properties of the solutions: lowering of the vapor pressure, raising of the boiling point, lowering of the solidification point, osmotic pressure. Fractional distillation. Solubility of a gas in a liquid (Henry's law). The solid state Crystalline solids and amorphous solids. Classification of solids in relation to the nature of the chemical bonds that determine their structure and properties: molecular, ionic, covalent, metallic solids. Metallic bond. Introduction to chemical thermodynamics First law of thermodynamics. Heat and temperature. Specific heat and molar heat of a substance at constant pressure or volume. Enthalpy concept. Thermochemistry. Hess's law. Second and third law of thermodynamics. Entropy, free energy. Spontaneous transformations and useful work. Calculation of the equilibrium constant and temperature effect: van’t Hoff equation. Chemical equilibrium First introduction of the concept of chemical equilibrium. Dynamic balance. Concept of the degree of progress of a reaction, reactions limited by chemical equilibrium. Expression of the equilibrium constant. Homogeneous and heterogeneous chemical equilibria. Reactions in the gas phase limited by the chemical equilibrium. Concept of molar fraction and partial pressure of a gas in a gaseous mixture. Le Chatelier's principle. Reaction quotient. Thermal dissociation of gases. Degree of dissociation. Dissociation factor. Chemical equilibrium in aqueous solutions Ionic product of water. Acids and bases in aqueous solutions and the concept of pH. Arrhenius and Broensted definitions and relation with the general Lewis definition. Saline hydrolysis. Polyprotic acids. Buffer solutions. Measurement of pH by means of indicators. Neutralization reactions. Acid-base titrations (strong acid-strong base, weak acid-strong base, weak base-strong acid). Not very soluble salts. Solubility product. Effect of the common ion. Formation of complex ions. Influence of acidity on the solubility of poorly soluble hydroxides and salts of weak acids. Fractional precipitation. Amphoteric hydroxides. Introduction to chemical kinetics Degree of progress of a reaction as a function of time: definition of the speed of a chemical reaction. Kinetic equations. Order of a reaction. Molecular energy distribution: Maxwell-Boltzmann law. Activation energy. Arrhenius equation. Impact theory (collisions). Transition state theory (outline). Homogeneous and heterogeneous catalysts. Reaction mechanisms (outline). Electrochemistry Pila di Volta. Chemical and concentration batteries. Cathodic and anodic semi-elements. Nernst equation. Electrode potential and definition of standard potential. Electromotive force (emf) and potential difference of a cell. Emf measurements to determine the concentration of chemical species, thermodynamic quantities and equilibrium constants. PH measurement using the pH meter. Fuel cells. Alkaline battery, lead accumulator (outline). Corrosion (notes). Possible chemical reactions with the use of electricity: electrolysis and Faraday's law. Electrolysis: hydrogen and oxygen from water. Preparation of sodium and chlorine by electrolysis of molten NaCl. Chlorine-soda process and sodium hypochlorite synthesis. Hall process for aluminum production. Properties of the elements Essential chemical characteristics of the most common elements of the periodic system of blocks s, p, d, and f. Classroom and laboratory activities Chemical problems related to all the topics will be discussed in class for a better understanding of the theoretical topics. Three laboratory exercises will be carried out: a) copper reactions; b) effect of temperature on chemical equilibrium, application of the van’t Hoff equation; c) construction of some piles and application of the Nernst equation.

Preliminary knowledges recommended in order to proficiently follow the course consist of basic mathematical skills which should be already acquired from high school teaching programs. More in particular, students should know the concept of function of one variable, being able to write and solve first and second degree algebraic equations. They should know how to use a scientific calculator to do simple calculations, including elevating to power, roots and logarithms. Preliminary basic chemical knowledge is useful but it is not necessary to follow the course.

CHIMICA GENERALE, D. A. McQuarrie, P. A. Rock, E. B. Gallogly, Seconda edizione italiana condotta sulla quarta edizione americana (con sito web), Zanichelli editore S.p.A., Bologna, 2012.

Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. On the contrary, all tutorials in the laboratory are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.

The exam consists of a written and an oral test. The written test requires to solve 5 chemical problems within an allocated time of 3 hours, and is passed with a grade ≥18/30. The topics of the test are: 1) Stoichiometric calculations in a redox reaction; 2) Calculations of the empirical and molecular formula of a substance starting from chemical analysis and colligative properties of a solution containing the substance; 3) Obtaining the equilibrium composition from data of gas phase or heterogeneous equilibria; 4) Obtaining chemical data from a chemical equilibrium in aqueous solution, such as pH of solutions of acid, bases, salts, buffers, slightly soluble salts; 5) Extracting relevant data from an electrochemical cell by application of Nernst's equation. The oral exam follows the written test by a few days, and consists of a discussion and correction of the errors in written test, plus questions about a few relevant topics of the program. This exam worths 12 ‘universitary credits’ and enables students to apply to the Inorganic chemistry 1 examination, a course given in the second semester of 1st year course.

General Chemistry 1) John C. Kotz, Paul M. Trichel, John R. Townsend, Chimica, Edises, VII edizione, 2021 2) R. H. Petrucci, F. G. Herring, J. D. Madura, C. Bissonette, Chimica generale, Piccin, 11° edizione, 2018 3) Ivano Bertini, Claudio Luchinat, Fabrizio Mani, Enrico Ravera, Chimica: struttura, proprietà e trasformazioni della materia, Casa Editrice Ambrosiana, 2022 4) Peter Atkins, Loretta Jones, Leroy Laverman, Principi di Chimica, Zanichelli, 2018 5) Paolo Silvestroni, Fondamenti di chimica, undicesima edizione, Casa Editrice Ambrosiana, 2020 Stechiometry 1) I. Bertini, C. Luchinat, F. Mani, E. Ravera Stechiometria - Un avvio allo studio della chimica, Casa Editrice Ambrosiana, sesta edizione, 2020 2) F. Cacace, M. Schiavello, Stechiometria, Bulzoni Editore, 1992 3) A. Caselli, S. Rizzato, F. Tessore, Stechiometria, Edises, VI edizione, 2021 4) M. Giomini, E. Balestrieri e M. Giustini, Fondamenti di Stechiometria, Edises, II Edizione, 2009 5) A. Paterno Parsi, A. Parsi, T. Pintauer, L. Gelmini, R. M. Hilts, Esercizi svolti. Chimica generale. Principi ed applicazioni moderne (R. H. Petrucci, F. G. Herring, J. D. Madura, Bissonette), Piccin,10°edizione 2015

This 12 CFU (Universitary Learning Credits) class consists of lectures of two hours each, four times per week from the end of September to the end of January, followed by a series of laboratory experiments with a duration of about 3 hours each. The course follows an integration of experimental and theoretical several approaches. In the first part of the program the lessons in class are dedicated to illustrate the fundamentals subject of chemistry (molecules, atoms and chemical reactivity). After two weeks there will be lessons dedicated to solve chemical problems relative to reactions and chemical calculations (the classification of the chemical reactions, the concept of amount of substance and its unit, the mole). The final part of the program consists to do several experiences in the laboratory, to make it clear to the students the experimental nature of chemistry. Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. On the contrary, all tutorials in the laboratory are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.

Recalls fundamental notions of classical physics: mass, heat and energy; concept of waves and their properties. Limits of classical physics in the description of some fundamental phenomena such as: black body radiation, photoelectric effect, quantization of electronic energies in the emission spectra of the hydrogen atom (Bohr model). Introduction to Atomic Theory. Heisenberg's uncertainty principle, Schrödinger's wave equation and the structure of the hydrogen atom. Quantum numbers. Atomic orbitals and their energy levels. The Aufbau principle. Polyelectronic atoms and their electronic configuration. Relationship between electronic configurations of elements and their properties. Atomic radii, ionic radii, ionization energy and electronic affinity. The Periodic Table. The chemical bond: ionic, covalent. The concept of electronegativity and polarity of bonds. Gillespie model based on repulsion between electron pairs in valence orbitals (VSEPR) for predicting molecular geometries. Valence bond theory. Hybridization and resonance. Structures of simple molecules; homonuclear and heteronuclear diatomic molecules. Description of the structure of simple polyatomic molecules of fundamental importance (structures of the most common acids and bases). Concept of oxidation number and formal charge. Length, angle and bond strength. Introduction to the theory of molecular orbitals, application to the description of simple homonuclear and heteronuclear diatomic molecules. Metallic bond (outlines). Intermolecular forces: ion-dipole, dipole-dipole, interactions between induced dipoles (Van der Waals interactions and London dispersion forces). Hydrogen bonding: nature and effect on the structure of some condensed phases. Examples of the structure of some condensed phases (ionic, covalent, molecular solids). Gases, equation of state of ideal gases and applications. Notes on the kinetic theory of gases. Real gases, van der Waals equation. The Principles of Thermodynamics and applications. The chemical equilibrium. Relationship between free energy and equilibrium constants (Kp, Kc). Studies of chemical equilibria in the homogeneous gas phase and in the heterogeneous phase. Homogeneous equilibria in aqueous solution: Acid-base theories and applications. Definition of pH. Autoprotolysis of water. Strength of acids and bases. Relationship between structure and acid or basic strength. Study of the acid-base behavior of some salts. Buffer solutions. Homogeneous equilibria in aqueous solution: Slightly soluble salts and solubility equilibria. Enthalpies of solution and hydration of ions, their relationship with the solubility of ionic compounds. Lewis acid-base theory (outline). Electrochemistry and oxidation-reduction reactions. Electrode potentials and electromotive force of an electrochemical cell. Nernst's law and its thermodynamic meaning. Standard potentials. Some examples of batteries and applications. Electrolysis; Faraday's laws. Physical equilibrium: concept of vapor pressure and Clapeyron's law. State diagrams (H2O, CO2). Raoult's law. Ideal and non-ideal solutions. Colligative properties. Notes on chemical kinetics; reaction order and kinetic laws. Effect of temperature on the rate of a reaction, Arrhenius equation and Activation Energy. Determination of the mechanism of a reaction by kinetic study. The role of catalysts in chemical reactions. Examples and applications to some processes of fundamental importance. Notes on inorganic chemistry: General chemical-physical and reactivity properties of the main group elements and noble gases.

Basic knowledge required The scientific method. International system of units of measurement (SI), fundamental and derived physical quantities. Symbols, dimensional analysis, tables and graphs of physical quantities. Accuracy and percentage error. Accuracy and significant figures. Elements, atoms, ions. Atomic mass; Atomic mass unit (a.u.); concept and definition of mole Compounds, molecules. Molecular mass. Substances. Homogeneous and heterogeneous mixtures of substances. Nuclear model of the atom (Rutherford). Protons. Electrons. Neutrons. Isotopes. Chemical nomenclature. Chemical reactions. Chemical equations. Physical properties and chemical properties of substances. Periodic system of elements. The concept of mole, atomic and molecular mass. Stoichiometric calculations. Calculation of the percentage composition of a compound. Elementary analysis of a substance. Minimal formula and molecular formula. Combustion reactions of hydrocarbons and determination of their molecular formula. Chemical reactions, chemical equations and their balancing. Oxidation-reduction reactions. Stoichiometric reactions, limiting reagent. Calculation of quantitative relationships between substances. Reaction yields.

Testi proposti Gli argomenti svolti a lezione sono presenti in TUTTI i testi di Chimica Generale a livello universitario. Nella biblioteca del Dipartimento di Chimica sono presenti diversi testi di Chimica Generale disponibili per il prestito e la consultazione. R. H. Petrucci, F. G. Herring, J. D. Madura, C. Bissonette. Chimica Generale, principi ed applicazioni moderne, Ed. Piccin. D.A McQuarrie, P.A. Rock, E.P. Gallogly, Chimica Generale, Ed. Zanichelli P. Silvestroni, “Fondamenti di Chimica”, XI edizione. Casa Editrice Ambrosiana. P.W. Atkins, L. Jones; Principi di Chimica, III Edizione italiana, Ed. Zanichelli. S. Borocci, M Crucianelli, M.L. Di Vona, C. Fraschetti, S. Lamponi, G. Leone, A. Magnani, D. Monti, L. Rossi, Le Basi della Chimica, Ed. A.L.E. F. Cacace, M. Schiavello, “Stechiometria”, Ed. Bulzoni. P.M. Lausarot, G.A. Vaglio, Stechiometria per la Chimica Generale, Ed. Piccin.

Attendance at lessons is free. Attendance of educational-practical exercises (laboratory activities) is mandatory. The initial lesson dedicated to safety in the chemical laboratory is mandatory. The dates will be promptly communicated during the course.

Procedure for carrying out the exam To take the exam you must book on the INFOSTUD Sapienza University of Rome website: https://stud.infostud.uniroma1.it/Sest/Log/ The exam consists of a written test that involves the solution of 5 chemical problems, having 3 hours available. The topics of the test are generally the following: 1) redox reactions and stoichiometric calculations; 2) calculation of the minimal and molecular formula of a substance using data from chemical analysis and colligative properties of solutions; 3) resolution of chemical equilibria in the gaseous or heterogeneous phase; 4) resolution of chemical balances in aqueous solution: pH of solutions of acids, bases, salts, buffer solutions, slightly soluble salts; 5) calculation of species concentrations or equilibrium constants by means of electrochemical batteries and application of the Nernst equation. The written test is considered passed if a grade of ≥ 18/30 is obtained, and allows access to the final oral test. This consists of a discussion of the written test and an interview on some of the main topics of the program. Passing the exam entails the acquisition of 12 university credits and the possibility of taking the Inorganic Chemistry 1 exam, taken in the following semester.

n.a.

The teaching is carried out in "frontal mode" and consists of lessons on the topics reported in the program, accompanied by stoichiometry exercises and chemical calculations aimed at elucidating and deepening the topics covered. Further lessons/tutorials will be conducted by the teacher and by tutors aimed at carrying out guiding exercises to prepare for the written test. These can possibly be delivered in mixed mode (frontal and remote).

Introduction to chemistry. Elements, atoms, ions. Atomic mass. Substances and compounds, molecules. Molecular mass. Chemical nomenclature. Chemical reactions. Chemical equations. Periodic system of the elements. Physical properties and chemical properties of substances. Stoichiometric calculations. The concept of the mole. Elemental analysis of a substance. Empirical formula and molecular formula. Calculation of the percentage composition of a compound. Chemical reactions, chemical equations and their balancing. Stoichiometry in reactions. Reagents in deficiency and in excess, limiting reagent. Reaction yields. Chemical reactions. Combination reactions, decomposition reactions, single displacement and double displacement reactions. Acids and bases. Acid-base reactions. Concept of oxidation number. Oxidation-reduction (redox) reactions. Balancing of redox reactions with the ion-electron method in acidic and alkaline environment. Redox reactions and chemical analysis. Outlines of atomic structure and fundamental theories. Dalton's atomic theory. Nuclear model of the atom (Rutherford). Protons. Electrons. Neutrons. Isotopes. Periodic properties of the elements. Quantum theory and atomic structure. Heisenberg's uncertainty principle (outlines). Schrödinger's equation (outlines). Quantum numbers. Core electrons and valence electrons. Pauli exclusion principle. Hund's rule. Electronic configurations and periodicity. Substances and chemical bonding. Nomenclature of substances, stoichiometry and composition. Ionic bond. Ionic charges and chemical formulas. Transition metal ions. Bonding in molecules: valence bond model. Covalent contribution, ionic contribution, resonance contribution, electronegativity, hybridization, sigma and pi bonds, resonance hybrids. Lewis model and structural formulas. Outlines of the molecular orbital model, outlines of bonding in metals. Stoichiometric calculations in solution. Solubility rules in water for inorganic compounds. Concept of saturated solution. Electrolytes and non-electrolytes. Strong and weak electrolytes. Dissociation reactions, concept of degree of dissociation. Density of the solvent and of solutions. Concentration of a solute: molarity, weight percentage. Outlines on the conductivity of aqueous electrolyte solutions. Concentrated solutions: relationship between chemical activity and concentration. The gaseous state. Properties of gases. Equation of state for ideal gases. Molar volume of gases. Density of gases. Dalton's law of partial pressures. Kinetic theory of ideal gases. Temperature and kinetic energy. Van der Waals equation to describe real gases. The condensed states. Processes of fusion, evaporation and sublimation. Intermolecular bonds: dipole-dipole interactions and London forces (van der Waals forces); hydrogen bond, ion-dipole interactions. Vapor pressure of a liquid. Viscosity, surface tension, description of crystalline solids. Phase diagrams (water, carbon dioxide). Phase transitions: Clausius-Clapeyron equation. Introduction to chemical thermodynamics. First law of thermodynamics. Heat and temperature. Thermochemistry and enthalpy. Hess's law. Second and third law of thermodynamics. Entropy, Free energy. Spontaneous transformations and useful work. Solutions. Concentration of solutions. Concept of ideal solutions: Raoult's law. Colligative properties of solutions. Fractional distillation. Solubility of a gas in a liquid (Henry's law). Chemical equilibrium in the gas phase. Dynamic equilibrium. Degree of reaction. Law of mass action. Homogeneous and heterogeneous gas equilibria. Le Chatelier's principle. Reaction quotient. Thermal dissociation of gases. Degree of dissociation. Effect of temperature on the equilibrium constant: van't Hoff equation. Chemical equilibrium in solution. Ionic product of water. Acids and bases in aqueous solutions and concept of pH. Salt hydrolysis. Polyprotic acids. Buffer solutions. Neutralization reactions. Acid-base titrations. Sparingly soluble salts. Solubility product. Common ion effect. Formation of complex ions. Electrochemistry. Cathodic and anodic half-cells. Chemical and concentration cells. Electrode potential and definition of standard potential. Nernst equation. Electrolysis and Faraday's law. Introduction to chemical kinetics. Definition of the rate of a chemical reaction. Rate equations. Order of a reaction. Molecular energy distribution: Maxwell-Boltzmann law. Activation energy. Arrhenius equation. Collision theory. Transition state theory (outlines). Homogeneous and heterogeneous catalysts. Reaction mechanisms (outlines). Laboratory activities. The course includes three laboratory exercises: a) reactions of copper; b) effect of temperature on chemical equilibrium, application of the van't Hoff equation; c) construction of some electrochemical cells and application of the Nernst equation.

The preliminary knowledge recommended in order to proficiently follow the course consists of basic mathematical skills which should be already acquired from high school teaching programs. In particular, students should know the concept of function of one variable, being able to write and solve first and second degree algebraic equations. They should know how to use a scientific calculator to do simple calculations, including elevating to power, roots and logarithms. Preliminary basic chemical and physical knowledge is useful to follow the course.

the following books are suggested: 1) Petrucci et al. - CGeneral Chemistry - ed. Piccin 2) Atkins P. Jones L. L. Laverman - Chemistry principles - ed. Zanichelli 3) Tro N.J. - Chemistry, a molecular approach - ed. EdiSES 4) Silberberg et al. - Chemistry- ed. Mc Graw Hill 5) Robinson et al. - General chemistry (VIII edizione) - ed. Pearson Students are advised to take advantage of the service offered by the department library in addition to having a textbook. Students who already own a book other than those listed above can ask the teacher to evaluate it.

N/C

Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. On the contrary, all tutorials in the laboratory are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.

The exam consists of a written and an oral test. The written test requires to solve 5 stoicochemical problems within an allocated time of 3 hours, and is passed with a grade ≥18/30. This exam worths 12 ‘universitary credits’ and enables students to apply to the Inorganic chemistry 1 examination, a course given in the second semester of 1st year course.

N/C

This 12 CFU (Universitary Learning Credits) class consists of lectures of two hours each, four times per week from the end of September to the end of January, followed by a series of laboratory experiments with a duration of about 3 hours each. The course follows an integration of experimental and theoretical several approaches. In the first part of the program the lessons in class are dedicated to illustrate the fundamentals subject of chemistry (molecules, atoms and chemical reactivity). After two weeks there will be lessons dedicated to solve chemical problems relative to reactions and chemical calculations (the classification of the chemical reactions, the concept of amount of substance and its unit, the mole). The final part of the program consists to do several experiences in the laboratory, to make it clear to the students the experimental nature of chemistry.

Introduction to chemistry. The scientific method. International System of Units (SI), fundamental and derived physical quantities. Symbols, dimensional analysis, tables and plots of physical quantities. Accuracy and errors. Precision and significant figures of an experimental quantity. Elements, atoms, ions. Atomic mass. Atomic mass units (u.m.a.). Compounds, molecules. Molecular mass. Substances. Homogeneous and heterogeneous mixtures of substances. Chemical nomenclature. Chemical reactions. Chemical equations. Periodic table of the elements. Physical and chemical properties of substances. Stoichiometric calculations. The concept of mole. Elementary analysis of a substance. Minimal formula and molecular formula (Avogadro's hypothesis and Cannizzaro’s law). Calculation of percentage composition of a compound. Chemical reactions, chemical equations and their balance. Reagents in excess and in defect. Combustion reactions of hydrocarbons. Calculation of the quantitative relations between different substances. Yield of a reaction. Atomic Structure. Dalton's atomic theory. Nuclear model of the atom (Rutherford). Protons. Electrons. Neutrons. Isotopes. Energy of first ionisation. Periodic properties of the elements. Line spectra of atoms. Wave-particle duality. Bohr model for the hydrogen atom. Quantization. Behavior of electrons in atoms. Electron affinity. Concept of electronegativity. Introduction to quantum theory and atomic structure. Quantum numbers. Internal electrons and valence electrons. Pauli exclusion principle. Hund's rule. Electronic configurations and periodicity. Atomic orbitals. Chemical bonding and molecular geometry. Introduction to the concept of chemical bonding by means of the formulas of Lewis. Gillespie model based on the repulsion between electron pairs in valence orbitals (VSEPR). Bond order. Energy of chemical bonds. Concept of electronegativity applied to an atom in a molecule. Concept of oxidation number. Formal charge and actual charge of an atom in a molecule. Covalent bonding, ionic, metallic. Coordination bond, hydrogen bonding. Notes on the method of valence bond structure for determining the hybridization of atomic orbitals. Concept of resonance structures. Notes on molecular orbital theory. Metal bond. Chemical reactivity – generalities. Acid-base reactions and oxidation-reduction (redox) reactions according to Lewis. Traditional nomenclature and the rules of the International Union of Pure and Applied Chemistry (IUPAC) applied to inorganic compounds. Concept of degree of advancement of a reaction. Stoichiometric reactions, limiting reactant, chemical equilibrium limited reactions. Chemical equilibrium: homogeneous and heterogeneous chemical reactions. Reactions which need use of electricity: Faraday electrolysis. States of matter and introduction to chemical thermodynamics. The liquid state and solutions. Molecular dipoles. Electronic polarizability. Van der Waals forces and intermolecular bonds. Water and aqueous solutions: the importance of hydrogen bond and of ion-dipole and dipole-dipole interactions. Rules of solubility in water for the inorganic compounds. Concept of saturated solution. Electrolytes and non-electrolytes. Strong and weak electrolytes. Reactions of dissociation, the concept of degree of dissociation and the definition of the binomium of van't Hoff. Density of the solvent and solutions. Concentration of a solute: fraction (percentage) in the mass, mass per unit volume, fraction (percentage) by volume, molarity, molality. Conductance of aqueous solutions of electrolytes. Molar conductance of the solutions and molar conductance of single ions at infinite dilution. Concentrated solutions: relationship between chemical activity and concentration. The gaseous state. Derivation of the equation of state of ideal gas behavior obtained by combining the laws of Boyle, Gay-Lussac and Avogadro's hypothesis. Kinetic theory of ideal gases. Temperature and kinetic energy. Distribution of molecular velocities of Maxwell. Van der Waals equation to describe the behavior of real gases. Reactions in the gas phase chemical equilibrium limited. Concept of mole fraction and partial pressure of a gas in a gaseous mixture.The solid state. Classification of solids in relation to the nature of the chemical bonds that determine the structure and properties: molecular solids, ionic, covalent, metallic. First law of thermodynamics. Heat and temperature. Specific heat and molar heat of a substance under pressure or constant volume. Concept of enthalpy. Thermochemistry. Hess's law. Second and third law of thermodynamics. Entropy, free energy. Spontaneous transformations and useful work. Calculation of the equilibrium constant and the effect of temperature: van't Hoff equation. State diagrams (water, carbon dioxide). Concept of ideal solutions: Raoult's law. Change of state: Clausius-Clapeyron law. Colligative properties of solutions: vapor pressure, raising of the boiling point, lowering of the freezing point, osmotic pressure. Chemical equilibrium in aqueous solutions. Ionic product of water. Acids and bases in aqueous solutions and the concept of pH. Definitions of acidity and basicity given by Arrhenius and Broensted, and relationship with the definition of Lewis. pH measurement by means of indicators. Hydrolysis and buffer solutions. Partially soluble salts. Physical and chemical parameters that influence solubility. Solubility of gases in water. Henry's Law. Distribution of a substance between two immiscible solvents. Introduction to chemical kinetics. Time evolution of a chemical reaction: definition of speed of a chemical reaction. Kinetic equations and reaction order. Zero-order, first-order, second-order reactions. Arrhenius equation. Collision theory. Notes on transition state theory and reaction mechanism. Homogeneous and heterogeneous catalysts. Examples. Electrochemistry. Cathodic and anodic half cells. Nernst equation. Electrode potential and definition of standard potential. Electromotive force (emf) and the potential difference of a pile. Voltaic pile. Electrolysis: hydrogen and oxygen from water by means of Volta's pile. Measurements of emf to determine the concentration of chemical species, and of thermodynamic equilibrium constants. Measurement of pH using pH-meter. Alkaline battery. Lead accumulator. Overvoltage, corrosion. Properties of the elements. The essential chemical characteristic of the most common elements of the periodic system of s, p and d blocks will be discussed. Tutorials. Chemical problems related to all the arguments will be presented in tutorials. Laboratory exercises are selected from the following: a) preparation of a complex salt; b) oxidation-reduction reactions; c) effect of the temperature on chemical equilibirum, application of the equation of van't Hoff; d) construction of some piles and application of the Nernst equation.

Preliminary knowledges recommended in order to proficiently follow the course consist of basic mathematical skills which should be already acquired from high school teaching programs. More in particular, students should know the concept of function of one variable, being able to write and solve first and second degree algebraic equations. They should know how to use a scientific calculator to do simple calculations, including elevating to power, roots and logarithms. Preliminary basic chemical knowledge is useful but it is not necessary to follow the course.

1) Silberberg et al. Chimica (McGraw-Hill) 2) Petrucci et al. Chimica generale (Piccin) 3) Tro. Chimica – un approccio molecolare (EdiSES) or any universitary Chemistry textbook. Books for stoichiometric calculations: 1) D. Monti, M. Stefanelli, E. Viola, Manuale di Stechiometria (Piccin) 2) M. Giomini, E. Balestrieri e M. Giustini, Fondamenti di Stechiometria (Edises)

This 12 CFU (Universitary Learning Credits) class consists of lectures of two hours each, five times per week from the end of September to the end of January, followed by a series of laboratory experiments with a duration of about 3 hours each. The course follows an integration of experimental and theoretical several approaches. In the first part of the program the lessons in class are dedicated to illustrate the fundamentals subject of chemistry (molecules, atoms and chemical reactivity). After two weeks there will be lessons dedicated to solve chemical problems relative to reactions and chemical calculations (the classification of the chemical reactions, the concept of amount of substance and its unit, the mole). The final part of the program consists to do several experiences in the laboratory, to make it clear to the students the experimental nature of chemistry. Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. All tutorials in the laboratory (in-person or online) are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.

This class consists of lectures of one or two hours each, 9 hours per week from the end of September to mid-January, followed by a series of laboratory experiments with a duration of about 3 hours each. Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. All tutorials in the laboratory (in-person or online) are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.

The exam consists of a written and an oral test. The written test requires to solve 5 chemical problems within an allocated time of 3 hours, and is passed with a grade ≥18/30. The topics of the test are: 1) Stoichiometric calculations in a redox reaction; 2) Calculations of the empirical and molecular formula of a substance starting from chemical analysis and colligative properties of a solution containing the substance; 3) Obtaining the equilibrium composition from data of gas phase or heterogeneous equilibria; 4) Obtaining chemical data from a chemical equilibrium in aqueous solution, such as pH of solutions of acid, bases, salts, buffers, slightly soluble salts; 5) Extracting relevant data from an electrochemical cell by application of Nernst's equation. The oral exam follows the written test by a few days, and consists of a discussion and correction of the errors in written test, plus questions about a few relevant topics of the program. This exam worths 12 ‘universitary credits’ and enables students to apply to the Inorganic chemistry 1 examination, a course given in the second semester of 1st year course.

This 12 CFU (Universitary Learning Credits) class consists of lectures of one or two hours each, 9 hours per week from the end of September to mid-January, followed by a series of laboratory experiments with a duration of about 3 hours each. The course follows an integration of experimental and theoretical several approaches. In the first part of the program the lessons in class are dedicated to illustrate the fundamentals subject of chemistry (molecules, atoms and chemical reactivity). After two weeks there will be lessons dedicated to solve chemical problems relative to reactions and chemical calculations (the classification of the chemical reactions, the concept of amount of substance and its unit, the mole). The final part of the program consists to do several experiences in the laboratory, to make it clear to the students the experimental nature of chemistry. Students are not obliged to follow the lessons and tutorials on chemical calculation, but it is strongly suggested that they attend them. All tutorials in the laboratory (in-person or online) are to be attended, including the preliminary lesson about safety rules in a chemical laboratory.