Unit GENERAL CHEMISTRY AND INORGANIC CHEMISTRY

GENERAL CHEMISTRY

Code	GP003081
CFU	6
Teacher	Morena Nocchetti
Teachers	Morena Nocchetti
Hours	48 ore - Morena Nocchetti
Learning activities	Base
Area	Discipline chimiche
Sector	CHEM-03/A
Type of study-unit	Obbligatorio (Required)
Language of instruction	Italian
Contents	Introduction to the general chemistry. Atomic theory and electronic structure of the atoms. Molecular geometries. Ionic and covalent bond theories. Intermolecular forces. Chemical reactions. State of matter: solid, liquid, gas. Equilibrium. Acids, Bases, and Salts-Ionic Equilibria. Electrochemistry. Kinetic elements.
Reference texts	M. Schiavello, L. Palmisano, Fondamenti di Chimica, EdiSES. lecture notes
Educational objectives	This teaching is the first rigorous approach to the general and inorganic chemistry. The main objective of the course is provide to the students the basic concepts of general chemistry as a description of nature, an appropriate scientific language and the ability to study in a critical and reasoned way. The main knowledge gained will be: - Atomic theory and electronic structure of atoms. - Chemical bond and molecular geometries. - Intermolecular forces. - Chemical reactions. - Chemical equilibrium in the gas phase and in aqueous solution. The main skills (ability to apply the knowledge acquired): - Identify and be able to write formulas of inorganic compounds; - Represent inorganic molecules or molecular ions highlighting the orientation of the atoms and the bonds between them; - Predict the polarity and the physical state of molecules; - Predict the reactivity of inorganic compounds both in redox and in non-redox reactions; - Write and describe the qualitative and quantitative aspects (stoichiometric) of a chemical reaction in relation to chemical homogeneous and heterogeneous equilibrium.
Prerequisites	In order to understand and achieve the expected learning targets the student should possess skills of mathematics and physics. In particular, the student should know and be able to use some basic mathematical tools (equivalence, linear and quadratic equations, logarithm, exponential function, inequations, derivatives, integrals) and notions of fundamental physics (unit of measurement, force, energy).
Teaching methods	The course is organized as follows: - Lectures on all the topics of the course. The lessons will be conducted with the help of the blackboard and by the projection of slides. - Numerical exercitations in classroom for the guided solution of numerical exercises with the aid of the blackboard. The teaching material (slides, exercises proposed during numerical exercitations, texts of previous written exams) are made available to students on the platform unistudium after registration.
Learning verification modality	The evaluation of the actual acquisition by students of the learning outcomes will be done through a written exam and an oral exam. The written exam will be administered at the end of the lessons, according to the academic calendar, and will be aimed at ascertaining the student's ability to use the acquired skills to solve numerical problems related to practical cases. The exam lasts three hours and will consist of 13 numeric problems in the form of closed-response testing. The final oral exam will take if the mark of the written exam is greater than or equal to 18/30. The oral exam lasts about 45 minutes and consists of questions on theoretical aspects related to the issues addressed in teaching and reported in the detailed program of the course. The purpose of the oral exam is to assess the knowledge, the understanding and the discipline language acquisition. Moreover, the ability of the student to explain the theoretical aspects and to apply the skills acquired in more complex systems, correlated to the program of teaching, is verified. The exam final judgment will take account of the marks obtained in the written and the oral exam.
Extended program	Generalities and elements of stoichiometry Intensive and extensive properties of matter. S.I measurement system, unit conversions. Forms of energy. Classification of matter: pure substances, elements, compounds, homogeneous and heterogeneous mixtures. Constitution of the atom, atomic number, mass number, nuclides, isotopes, elements. Atomic masses: absolute, relative and molar mass of atomic masses. Avogadro's constant and mole concept. Chemical formula (minimum and molecular), molecular weight (formula). Stoichiometry of mixtures. Chemical reactions, chemical equation, balancing of a chemical reaction, limiting reagent. Numerical exercises. Fundamentals of atomic theory Electromagnetic radiation and electromagnetic spectrum. The hydrogen atom according to N. Bohr: postulated, introduction of the main quantum number, electronic transitions, interpretation of the hydrogen spectra. Eisenberg's Indefinite Principle. Dual-wave particle and De Broglie's relationship. Quantum mechanics: Shrödinger equation, wave functions, and quantum numbers. Physical meaning of wave function, radial nodes and angular nodes of wave function, radial probability density curves, atomic orbitals, spin. Multi-electron atoms: approximate resolution of the Shrödinger equation by using Zeff. Electronic configurations, Aufbau principle, Pauli exclusion principle. Periodic table. Periodic properties of elements. Ionization energy, electron affinity, atomic rays and ionic rays, Ion radius calculation by the Pauling method. Periodicity in the chemical properties of hydrides and oxides. Numerical exercises. Chemical bonds and molecular structure Bonding energy, bond length, bonding angles. Ionic bond. Information on ionic lattices, ionic model and lattice energy, and calculation of lattice energy using thermodynamic cycles (Born-Haber). Covalent bond: electron pair theory, octet rule, homo and hetero nuclear bonds. Dative covalent bond. Exceptions to the octet rule. Valence bond theory. Sigma and pi greek bonds. Polyatomic Molecules: Method V.S.E.P.R. and molecular geometry. Hybridization. Formal charge. Electronegativity, electronegativity scales according to Mulliken and Pauling. Polarity of the bonds, percentage of ionic character of a polar covalent bond, polarity of the molecules. Formulas of the structure of the most common molecules and of the most common molecular ions. Resonance. Chemical nomenclature Oxidation number. Nomenclature of the most common compounds. Basic Oxides, Hydroxides, Acid Oxides, Acids, Salts. Classification of chemical reactions. Redox reactions: balance of the chemical reactions with the ionic-electronic method. Disproportionation reactions. Intermolecular bonds Ion-dipole interaction, dipole-dipole. Dipole-induced dipole, induced dipole-induced dipole, hydrogen bond and its consequences. The gaseous state Ideal model of perfect gas, perfect gas laws (ideal), Avogadro's law, perfect gas state equation (ideal). Gaseous mixtures: partial pressures and volumes. Dalton's Law. % By mass, by volume, the average molecular weight of a gaseous mixture. Distribution of molecular velocities (energies) according to Maxwell and Boltzmann. Numerical exercises. Solid state and liquid state Classification of solids: ionic solids, covalent solids, molecular solids, metallic solids. Properties of liquids. Viscosity and surface tension. The vapor pressure. The equilibrium and its characteristics. Liquid-liquid equilibrium, solid-liquid, solid-steam. Clausius-Clapeyron equation. Normal boiling and melting temperature. State diagrams (H2O, CO2). Concept of variance. Relative humidity. Solutions Definitions, solubility and temperature effect on solubility. Concentration (% in mass, molarity, molar fraction, molarity, normality). Convertion of the units. The equivalence principle. Classification of solutes, electrolytes and electrolytic dissociation. Colligative properties of solutions: lowering the vapor pressure. Raoult law. Cryoscopy, ebullioscopy, osmotic pressure. Numerical exercises. Chemical equilibrium Equilibrium characteristics, homogeneous and heterogeneous equilibria. Kp and Kc. Use of the equilibrium constant, Le Châtelier's principle. Effects of temperature, pressure and concentration on chemical equilibria. Gaseous dissociation: degree of dissociation and binomial of Van't Hoff. Ionic equilibria in aqueous solution: factors that influence salt solubility, insoluble salts, solubility product. Selective Precipitation. Numerical exercises.

INORGANIC CHEMISTRY

Code	GP003082
CFU	6
Teacher	Monica Pica
Teachers	Monica Pica Monica Pica
Hours	48 ore - Monica Pica 10 ore - Monica Pica
Learning activities	Base
Area	Discipline chimiche
Sector	CHEM-03/A
Type of study-unit	Obbligatorio (Required)
Language of instruction	Italian
Contents	Reactivity and strength of acids and bases. Electrochemistry. Thermodynamics and chemical kinetics. The chemical bond described by MO-LCAO method. Fundamentals of solid state and crystallography. The chemistry of coordination compounds.
Reference texts	M. Schiavello, L. Palmisano, Fondamenti di Chimica, EdiSES, Terza edizione, Napoli 2010. M. Schiavello, L. Palmisano, Fondamenti di Chimica, EdiSES, Terza edizione, Napoli 2010.
Educational objectives	Learning objectives The main educational objective of the course is to provide students with the basic concepts of inorganic chemistry, which are preparatory to several other disciplines, and to develop an appropriate scientific language as well as the ability to study in a critical and reasoned manner. The main knowledge acquired will concern: the reactivity of acid–base and electrochemical systems; chemical thermodynamics and kinetics; description of the chemical bond using the MO–LCAO method; principles of solid-state chemistry; coordination compounds. Main skills (ability to apply the acquired knowledge): ability to predict the reactivity of inorganic compounds in acid–base and electrochemical systems; ability to predict the spontaneity of chemical and physical processes; ability to interpret kinetic data; ability to predict the optical and magnetic properties of diatomic molecules using the MO–LCAO method and of complex ions using crystal field theory.
Prerequisites	Prerequisites: Knowledge of the basic concepts and contents of general chemistry (stoichiometry, atomic theory, chemical bonding, chemical equilibrium); Basic knowledge of mathematics (powers and logarithms, derivatives and integrals, first- and second-degree equations and their graphical representation); Basic knowledge of classical physics (work, heat, energy, force, pressure).
Teaching methods	Teaching methods The course is organized as follows: Lectures covering all the topics of the course. Lectures are delivered using the blackboard and supported by slide presentations. Numerical exercise sessions, consisting of guided problem-solving activities carried out with the support of the blackboard. Teaching materials (slides, exercises proposed during the problem-solving sessions, and texts of previous written examinations) are made available to students on the Unistudium platform.
Other information	Attendance is mandatory and integrated with the General Chemistry module. Supplementary teaching activities: numerical exercise sessions consisting of guided problem-solving activities carried out on the blackboard (2 hours per week).
Learning verification modality	Assessment methods A written examination (3 hours) consisting of numerical exercises, followed by an oral examination, both aimed at assessing the knowledge and skills acquired. The written examination is considered passed with a grade of 18/30 or higher.
Extended program	Equilibria in aqueous solution: acids and bases. Acid–base theories: Arrhenius; Brønsted–Lowry theory (conjugate acid–base pairs, strong and weak acids and bases); classification of solvents; leveling effect of the solvent; Ostwald’s dilution law. Autoprotolysis of water and KwK_wKw¿; pH, pKa, pKb; structure–acidity relationships; polyprotic acids. Acid–base reactions; salt solutions and their acid–base properties. Calculation of the pH of simple systems (solutions of strong acids and bases, weak acids and bases, polyprotic acids). Buffer solutions, Henderson–Hasselbalch equation, universal buffers, buffering properties of amino acids. Lewis acid–base theory: Lewis acid–base adducts, nucleophilic and electrophilic species. Acid–base properties of aquo-complexes of polyvalent metals, acid–base behavior of boric acid, amphoterism. Electrochemistry. Galvanic and electrolytic cells; conductors of the first and second kind; galvanic half-cells; the salt bridge; the Daniell cell. Notation of galvanic cells; classification of galvanic half-cells; electrode potential; galvanic cells under open- and closed-circuit conditions; electromotive force (EMF) of a galvanic cell. Standard electromotive force of a galvanic cell; standard electrode potentials; the standard hydrogen electrode; the scale of standard electrode potentials and its use in predicting the spontaneity of redox reactions under standard conditions. Dependence of the electromotive force on temperature and concentration: the Nernst equation and its applications. Equilibrium constant of a redox reaction. Electrolysis: brief introduction. Fundamentals of chemical thermodynamics. Definition of open, closed, and isolated systems; definition of thermodynamic state functions. Heat and mechanical work; Joule’s experiment. Specific heat and heat capacity at constant pressure and constant volume; heat capacities of ideal gases. First law of thermodynamics; internal energy; enthalpy; calorimetry; standard molar enthalpy of formation; Hess’s law and thermochemistry. Entropy and the second law of thermodynamics. Dependence of ¿S on temperature; entropy changes in phase transitions. Criteria of spontaneity and reversibility of chemical and physical processes based on ¿S. Third law of thermodynamics; standard absolute entropy. Gibbs free energy and standard Gibbs free energy; spontaneity of chemical reactions. Partial molar free energy. Chemical potential. Thermodynamic derivation of the equilibrium constant. Fundamentals of chemical kinetics. Elementary reactions and reaction mechanisms. Average and instantaneous reaction rates. Reaction rate and factors affecting it. Concentration and reaction rate. Differential rate laws; reaction order; specific rate constant. Integration of rate laws for zero-, first-, and second-order reactions. Half-life. Archaeological dating. Temperature and reaction rate; Arrhenius equation. Determination of activation energy. Reaction mechanisms. Molecularity and elementary reactions. Collision theory and activated complex theory. Reaction coordinate. Hammond postulate. Introduction to photochemistry. Introduction to homogeneous and heterogeneous catalysis. Chemical bonding and molecular orbital theory. Covalent bonding described using the MO–LCAO method. Bonding, antibonding, and nonbonding molecular orbitals. s and p symmetry of molecular orbitals. Electronic configurations of selected homo- and heteronuclear diatomic molecules. Bond order. Magnetic properties. HOMO and LUMO concepts. Hückel method. Coordination compounds and complex ions. Nomenclature. Central atom and mono- and polydentate ligands. Description of the metal–ligand bond according to the valence bond (VB) model. Crystal field theory for octahedral, square-planar, and tetrahedral complexes. The spectrochemical series. Optical and magnetic properties. Ligand field theory: brief introduction. Solid state. Introduction to crystallography: unit cell, the seven crystal systems and the 14 Bravais lattices. Bragg’s law of diffraction.
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