CHEMISTRY 1

Acquisition of important concepts of general chemistry, inorganic chemistry, stoichiometry and notes on organic matter. The course is organized so as to provide a good knowledge of basic chemistry (inorganic and stoichiometry) and is divided into a theoretical part and a part consisting of exercises aimed at solving chemistry problems (knowledge and understanding). The aim of the course is to acquire reasoning skills to deal with the study of chemical phenomena with analytical and numerical methods (applying knowledge and understanding). At the end of the course students should have their own judgment: ability to propose their own numerical exercises (making judgments) on specific topics.

In particular:

This course is specifically theoretical and contains many numerical exercises.

The specific training objectives of this course are:

Understanding the atomic structure;

understanding the mechanisms of chemical bond formation;

to know the chemical interactions in solids, liquids and understanding the gas state equations;

to know the main thermodynamic and kinetic quantities involved in chemical reactions;

To assess the conditions of chemical equilibrium;

to quantitatively determine the gaseous equilibria, in aqueous solutions and in electrochemical systems;

to discuss all proposed activities with scientific method and appropriate language.

To acquire the ability to set and correctly perform exercises on the various types of chemical reactions

Furthermore, in reference to the so-called Dublin Descriptors,

this course helps to acquire the following transversal skills:

Knowledge and understanding:

Capacity of inductive and deductive reasoning.

Ability to outline a chemical reaction in qualitative and quantitative terms.

Ability to set a problem using appropriate relationships between chemical and physical quantities and solve it with analytical methods.

Ability to apply knowledge:

Ability to apply the acquired knowledge for the description of chemical phenomena using the scientific method with rigor.

Capacity for quantitative calculation of reagents and products in chemical reactions.

Quantitative calculations on homogeneous solutions, of their colligative properties, of pH, and of electro-chemical phenomena.

Autonomy of judgment:

Ability to critical reasoning.

Ability to identify the most appropriate solutions to solve chemical problems.

Ability to identify the predictions of a theory or a model.

Ability to assess the accuracy required for stoichiometric calculations.

Communication skills:

Ability to describe a scientific topic in oral form, with properties of language and terminological rigor, illustrating the reasons and results.

The course includes 5 CFU (35 hours) of lectures and 1 CFU (15 hours) of classroom exercises.

Theoretical lessons will be alternate with numerical exercises in the classroom. Students will be actively involved in numerical exercises on the board. If possible, the teacher will experimentally perform simple and harmless chemical reactions in the classroom such as the preparation of cupric oxide, cuprous oxide, acid-base titrations with pH indicators, a redox titration (KMnO₄ - H₂C₂O₄) and the preparation of a chemical pile.

(Indispensable) the quantities, dimensions and units of measurement;

(Important) The basic concepts of chemistry and chemical symbols;

(Useful) the chemical nomenclature.

Starting course on the nomenclature.

1 - STRUCTURE OF THE ATOM (4 hours)
Subatomic particles: Electron, proton, neutron - atomic number, mass number - isotopes - atomic mass unit - Bohr / Rutherford atomic model.   Wave Mechanical atom model. - Atomic orbitals - quantum numbers - Principle of exclusion of Pauli - principle of maximum multiplicity - the principle of aufbau.

2 - PERIODIC SYSTEM OF ELEMENTS (2 hours)
Periodical classification and electronic configuration of the elements - Periodic properties: atomic and ionic radii, ionization energy, electron affinity, and electronegativity.

3 - STECHIOMETRY (15 hours)

The concept of mass and fundamental laws that govern the course of chemical reactions. The chemical equation. Classification of the main chemical reactions. Quantitative relationships in chemical reactions. Ionic compounds and molecular compounds; electrolytic dissociation; oxidation numbers, concepts of oxidant and reducing agent. Redox reactions and their balance. Determination of the formula of any compound. Numerical applications.

4 - CHEMICAL BOND (6 hours)

Concepts on ionic, covalent and dative bonds. Lewis formalism and valence bond theory. Structural formulas of the most common compounds: N2, O2, CO, NO, NO +, HCl, CO2, NO2, O3, NO2-, HCN, CH2O, CO3--, NO3-, CH4, C2H6, C2H4, C2H2, NH3, N2H4, NH4 +, NH2-, H2O, H2O2, H3O +, OH-, HF, CNO-, F2O; expanded valence shells: XeF2, PCl5, SO2, SO3, H2SO4, SO4--, SiF5-, SiF6--, PF5, PF6-, SF4, SF6, ClF3, BrF5, XeF4; common oxyacids and their anions; common isoelectronic molecules and ions; resonance concept. Electronegativity of atoms and polarity of bonds. Dipole moments of HF, HCl, HBr, HI, H2, NH3, NF3, BF3, H2O, H2S, SO2, CO2, CH4, CH3Cl, CH2Cl2, CHCl3, CCl4. Molecular geometries: state of carbon promotion, hybridization theory, sp3, sp2, sp, methane, ethane, ethylene, acetylene, benzene hybrid orbitals.

Molecular geometry and V.S.E.P.R theory: H2O, NH3, CH4, BF3, BeCl2, PCl5, TeCl4, ClF3, I3-, SF4, SF6, IF5, ICl4-, PF5, XeF2, BrF5, XeF4. Variation of potential energy during the formation of H2. Lattice energy and ionic bond. LiF molecule. Born-Haber cycle. Physical properties of ionic compounds. Chemical bond and theory of molecular orbitals, sigma and pi orbitals; diatomic molecules H2 +, H2, He2 +, Li2, B2, C2, N2, O2, F2, N2 +, N2-, O2 +, O2-. Order, length, force constant and binding energy. Dative bond. Metallic bond and hints of band theory. Heteronuclear diatomic molecules: HF, CN, CO, NO, O3, Benzene. Physical properties of covalent compounds. Metallic bond and hints of band theory.

5 - INTERMOLECULAR FORCES (2 hours)

Intermolecular forces: ion - ion, ion - dipole, dipole - dipole, Van der Waals and London forces. Hydrogen bond and boiling points of H2O, H2S, H2Se. Surface tension.

6 - ELEMENTS OF THERMODYNAMICS (2 hours)

7 - GAS LAWS (2 hours)
General characteristics of the gaseous state. Ideal gas. Ideal gas laws. State law of the ideal gas. Law of partial pressures and  volumes. Gas diffusion. Real gases. Numerical applications.

8 - CONDENSED STATES AND STATE CHANGES (2 hours)

Characteristics of the solid state as a function of the chemical bond - Characteristics of the liquid state, vapor pressure and boiling temperature of pure liquids. Phase changes. Phase diagrams of water and carbon dioxide - Principle of mobile equilibrium.

9 - WATER SOLUTIONS (4 hours)

Solvent, solute, gaseous, liquid and solid solutions, interactions between solute and solvent, hydrogen bond, unit of measurement of concentration: % by mass and volume, equivalent weight and normality, molarity, molality, molar fractions. Colligative properties of electrolytes in solution: Vapor pressure of ideal and real solutions, Raoult's law, distillation, azeotropes. Solutions of non-volatile solutes. Degree and factor of dissociation. Relative lowering of solvent vapor pressure, ebullioscopy and cryoscopy. Phase diagrams of granita and brine. Osmosis and osmotic pressure. Electrolyte solutions and degree of dissociation. Solubility and Henry's law. Numerical applications

10 - CHEMICAL EQUILIBRIUM (4 hours)

Elements of chemical thermodynamics. Equilibrium in homogeneous chemical systems - Mass action law and equilibrium constants: Kp, Kc, Kx. Factors affecting equilibrium: temperature and pressure. Ionic equilibria in aqueous solution. Theories of acids and bases. Arrhenius, Bronsted and Lewis acids and bases. Ampholytes. Complex ions in aqueous solution, polyprotic acids and bases. Acid-base reactions and stoichiometry of solutions. Strength of acids and bases. Strong and weak acids and bases. Self-dissociation of water and definitions of pH and pOH and pKw. HCl-NaOH titration. pH of strong and weak acids and bases: definitions of pKa, pKb. Hydrolysis of salts and pH calculations. Acid, basic buffer solutions. PH indicators. Heterogeneous equilibria - Solubility product. Numerical applications.

11 - ELECTROCHEMISTRY (3 hours)

Metal electrodes, measurements of d.d.p., galvanic cells, Nerst equation. Second species electrodes (Ag / AgCl / KCl). Third type electrodes: Pt / Fe2 + - Fe3 +; Pt / Mn2 + and MnO4-, Pt / Quinhydrone. Normal standard hydrogen electrode. Saturated calomel electrode. Standard potential series of semi-elements, predictions of red-ox reactions.

Reactions of disproportion. Stacks of concentration. Electrochemical determination of pH. Electrolysis and Faraday's laws. Overvoltage, electrolysis of water, of NaCl solutions, of ZnSO4; lead-acid accumulator. Numerical applications.

12 - CHEMICAL KINETICS (2 hours)

Reaction rate, reaction order, half-life, reaction molecularity, collision theory, Arrhenius equation and activation energy, factors affecting the reaction rate, catalysts.

13 - INORGANIC CHEMISTRY (2 hours)

Metals and non-metals: general information on chemical and physical properties. General characteristics of each group of the periodic system. Alkaline and alkaline earth metals. Main oxidation states and compounds of Hydrogen, Oxygen, Carbon, Nitrogen, Phosphorus, Sulfur and Chlorine. Transition elements: generalities. Coordination compounds. Binders. Coordination number and geometry. Nomenclature. Notes on the theories of the chemical bond in coordination compounds.

Contribution of teaching to the objectives of the 2030 Agenda for Sustainable Development

GOAL 3: HEALTH AND WELL-BEING - Ensure health and well-being for all and for all ages

Target: 3.9 By 2030, substantially reduce the number of deaths and illnesses from hazardous chemicals and from pollution and contamination of air, water and soil.

Method: lesson

GOAL 7: CLEAN AND ACCESSIBLE ENERGY - Ensure everyone has access to affordable, reliable, sustainable and modern energy systems

Target: 7.2 By 2030, significantly increase the share of renewable energy in the global energy mix

Method: lesson

Texts 1-4 are equivalent and the student is free to choose others not listed.
Texts of exercises 5-7 are equivalent and the student is free to choose others not listed.

1. KOTZ, TREICHEL, TOWNSEND  Chimica V edizione - EdiSES
2. SILVESTRONI, Fondamenti di Chimica 11° Ed. 2020. Ambrosiana. Distribuzione esclusiva Zanichelli
3. PETRUCCI, HERRING, MADURA, BISSONNETTE,  Chimica Generale - Piccin
4. EBBING, Chimica Generale - Editoriale Grasso
5. NOBILE, MASTRORILLI, Vol.1 e 2, Esercizi di Chimica -  Ambrosiana
6. GIOMINI, BALESTRIERI, GIUSTINI, Fondamenti di Stechiometria – EdiSES
7. P.MICHELIN LAUSAROT, G.A. VAGLIO, Fondamenti di Stechiometria - Piccin

The exam tends to ascertain the level of overall knowledge acquired by the candidate, his/her ability to critically address the topics studied and to relate the various parts of the program. The exam consists of a written test and an oral exam. The written exam consists in solving 3 simple stoichiometry problems related to nomenclature, chemical reactions, gas transformations and the study of chemical phenomena in aqueous media and electrochemistry. This test tends to verify the possession of the basic notions of the discipline. The time available for this test will be 75 minutes. The written test will be considered passed if the student has solved exactly at least two out of three exercises. During the written exam, students will be allowed to use books, notes, tables, calculators and specific teaching material but not mobile phones.

There are no ongoing tests.

The oral exam will consist of questions relating to the various parts of the program to ascertain the level of overall knowledge acquired by the candidate.

Criteria for the attribution of the final grade:

The mastery shown in the qualitative and quantitative arguments, the critical vision of the topics addressed during the course and the ability to implement the various parts of the program will contribute in equal measure to the formulation of the final grade. At the end, the average of the marks obtained in the two tests, written and oral, will be calculated, which will constitute the final mark.

Information for students with disabilities and/or DSA:

To guarantee equal opportunities and in compliance with the laws in force, the students concerned can, upon request, opt for a personal interview during which to implement any compensatory and/or dispensatory measures, based on the educational objectives and specific needs.

Date of Examination

As a rule, 8 exam sessions are set for each Academic Year; consult the Exam Calendar for the Bachelor's Degree Course in Physics: http://www.dfa.unict.it/corsi/L-30/esami.

Exam dates can be either at both written and oral exams.

The questions below are not an exhaustive list but are just a few examples Formulas of compounds. Atomic model of Bohr / Rutherford. Atomic orbitals. Quantum numbers. Periodic classification of the elements. Electronic configuration of the elements. Periodic properties. Ionic bond. Covalent bond. V.S.E.P.R. Theory Hybrid orbitals and molecular geometry. Molecular orbital theory. Ideal gases and laws of ideal gases. State diagrams. Concentration units. Colligative properties. Osmotic pressure. Law of mass action. Factors affecting the chemical equilibrium. pH. Acid-base theories. Ampholytes. Hydrolysis. Buffer solutions. Galvanic cells. Nernst equation. Standard potential series. Electrolysis. Faraday's laws.