General Inorganic Chemistry and Stoichiometry
A.Y. 2024/2025
Learning objectives
The course provides students with the fundamentals of General and Inorganic Chemistry that form the essential background for understanding the subsequent chemical teachings of the CdS Pharmacy. Students are also introduced to the main computational and reasoning methodologies for solving stoichiometry problems.
Expected learning outcomes
Knowledge of theoretical principles and understanding of the basic structure of atoms and molecules, the three states of matter, inorganic nomenclature, periodic properties of elements, chemical reactions and chemical equilibria in the gas phase and in aqueous solution, including their thermochemical basis. Understanding and solving major stoichiometry problems.
Lesson period: First semester
Assessment methods: Esame
Assessment result: voto verbalizzato in trentesimi
Single course
This course can be attended as a single course.
Course syllabus and organization
Linea AL
Responsible
Lesson period
First semester
Course syllabus
Matter: states of aggregation; physical and chemical properties; intensive and extensive properties. Compounds and mixtures. Chemical formulas. Unit of measure.
Dealing with numbers: exact numbers, approximate numbers. Accuracy and precision. Errors and uncertainties in measurements. Significant figures.
Atoms and elements: The atomic theory. Structure of the atom. Atoms and isotopes. Atomic mass. Periodic table. Molar mass.
Molecules and compounds: Ions and molecules. Nomenclature
Molecular mass. Avogadro's number and the mole concept. Percentage composition of compounds. Determination of the empirical formula and the molecular formula.
Chemical reactions: Chemical equations. Balancing chemical equations. Limiting reagent. Reaction yield. Examples of chemical reactions.
Solutions: Concentration. Stoichiometry of solutions. Reactions in aqueous solution. Precipitation reactions. Acid-base reaction. Reactions with evolution of gas. Oxidation-reduction reaction (Oxidation states and standard reduction potentials).
The gas. The pressure. Gas laws (Boyle, Charles, Gay-Lussac, Avogadro). Ideal gas equation and applications. Gas mixtures, Dalton's law. Molecular kinetic theory. Outline on real gases (van der Waals equation).
Outline on thermochemistry: Energy in chemical reactions. Enthalpy of reactions. Specific heat and heat capacity.
Quantum mechanical model of the atom: Electromagnetic radiation. Atomic emission spectroscopy. Bohr's model. Wave-particle duality. Heisenberg principle. Atomic orbitals. Electronic spin. Pauli exclusion principle. Orbitals and energy levels in multi-electron atoms. Electronic configuration and Aufbau principle.
Periodic properties of elements. The periodic table. Metals, non-metals and metalloids. Groups and periods. Periodicity. Atomic rays. Ionization energy. Electronic affinity.
The chemical bond: Ionic and covalent bond. Lewis structures and formal charge. The resonance. Exceptions to the octet rule. Electronegativity. Reactivity and polarity of bonds. Bond energies and lengths.
VSEPR theory. Molecular geometries. Polarity of molecules. Outlines of valence bond theory. Hybridization. Outlines on the theory of molecular orbitals. The metallic bond.
Intermolecular forces. State transitions. Phase diagrams. The water. Structure of water. Outlines on solids: ionic, molecular and atomic solids.
Solutions: solutions and solubility. Energetics of solutions. Colligative properties.
Outlines on chemical kinetics. Reaction rates and kinetic laws. Integrated velocity laws. Temperature effect (Arrhenius equation) and collision model. Catalysts.
Notes on thermodynamics: Entropy and Gibbs free energy
Chemical equilibrium. The concept of equilibrium. Homogeneous and heterogeneous equilibria. The equilibrium constant. Calculation of equilibrium concentrations. Le Châtelier's principle. Factors influencing chemical equilibrium. Equilibrium and Gibbs free energy.
Acids and bases. Acids and bases concept. Lewis acids. Strength of acids. pH scale. Determination of the pH of a solution. Acid-base properties of ions and salts. Acid-base titration. Indicators. Buffer solutions.
Solubility equilibria. Effect of the common ion.
Thermodynamics: Functions of State. The First Law of Thermodynamics. Enthalpy. Entropy. The Second Law of Thermodynamics. Standard Free Energy Change. Free Energy Change and Equilibrium.
Electrochemistry: Electrochemical cells, standard reduction potentials
Dealing with numbers: exact numbers, approximate numbers. Accuracy and precision. Errors and uncertainties in measurements. Significant figures.
Atoms and elements: The atomic theory. Structure of the atom. Atoms and isotopes. Atomic mass. Periodic table. Molar mass.
Molecules and compounds: Ions and molecules. Nomenclature
Molecular mass. Avogadro's number and the mole concept. Percentage composition of compounds. Determination of the empirical formula and the molecular formula.
Chemical reactions: Chemical equations. Balancing chemical equations. Limiting reagent. Reaction yield. Examples of chemical reactions.
Solutions: Concentration. Stoichiometry of solutions. Reactions in aqueous solution. Precipitation reactions. Acid-base reaction. Reactions with evolution of gas. Oxidation-reduction reaction (Oxidation states and standard reduction potentials).
The gas. The pressure. Gas laws (Boyle, Charles, Gay-Lussac, Avogadro). Ideal gas equation and applications. Gas mixtures, Dalton's law. Molecular kinetic theory. Outline on real gases (van der Waals equation).
Outline on thermochemistry: Energy in chemical reactions. Enthalpy of reactions. Specific heat and heat capacity.
Quantum mechanical model of the atom: Electromagnetic radiation. Atomic emission spectroscopy. Bohr's model. Wave-particle duality. Heisenberg principle. Atomic orbitals. Electronic spin. Pauli exclusion principle. Orbitals and energy levels in multi-electron atoms. Electronic configuration and Aufbau principle.
Periodic properties of elements. The periodic table. Metals, non-metals and metalloids. Groups and periods. Periodicity. Atomic rays. Ionization energy. Electronic affinity.
The chemical bond: Ionic and covalent bond. Lewis structures and formal charge. The resonance. Exceptions to the octet rule. Electronegativity. Reactivity and polarity of bonds. Bond energies and lengths.
VSEPR theory. Molecular geometries. Polarity of molecules. Outlines of valence bond theory. Hybridization. Outlines on the theory of molecular orbitals. The metallic bond.
Intermolecular forces. State transitions. Phase diagrams. The water. Structure of water. Outlines on solids: ionic, molecular and atomic solids.
Solutions: solutions and solubility. Energetics of solutions. Colligative properties.
Outlines on chemical kinetics. Reaction rates and kinetic laws. Integrated velocity laws. Temperature effect (Arrhenius equation) and collision model. Catalysts.
Notes on thermodynamics: Entropy and Gibbs free energy
Chemical equilibrium. The concept of equilibrium. Homogeneous and heterogeneous equilibria. The equilibrium constant. Calculation of equilibrium concentrations. Le Châtelier's principle. Factors influencing chemical equilibrium. Equilibrium and Gibbs free energy.
Acids and bases. Acids and bases concept. Lewis acids. Strength of acids. pH scale. Determination of the pH of a solution. Acid-base properties of ions and salts. Acid-base titration. Indicators. Buffer solutions.
Solubility equilibria. Effect of the common ion.
Thermodynamics: Functions of State. The First Law of Thermodynamics. Enthalpy. Entropy. The Second Law of Thermodynamics. Standard Free Energy Change. Free Energy Change and Equilibrium.
Electrochemistry: Electrochemical cells, standard reduction potentials
Prerequisites for admission
Basic math knowledge (I e II degree equations, logarithms, Integrals).
SI measurement units and their conversions.
SI measurement units and their conversions.
Teaching methods
Traditional Lectures: Classroom lectures with the aid of slides concerning both the theoretical part and the resolution of stoichiometry exercises. Classroom practices on solving stoichiometry exercises.
Smart interactive learning: Use of the "exam manager" online platform to assign stoichiometry exercises and theoretical questions. Classroom simulations of exam tests. The teaching materials (slides and exam tests) are made available after classroom lessons on the teacher's myAriel website.
Smart interactive learning: Use of the "exam manager" online platform to assign stoichiometry exercises and theoretical questions. Classroom simulations of exam tests. The teaching materials (slides and exam tests) are made available after classroom lessons on the teacher's myAriel website.
Teaching Resources
Copies of the lecture slides, including the homework problems, will be made available on the teacher's myAriel website. Any general chemistry textbook for university courses.
Recommended textbooks (any of the following):
- N. J. Tro, Chimica Un approccio molecolare, EdiSES
- F. Demartin, Fondamenti di Chimica Generale, EdiSES
- R. Chang, Fondamenti di Chimica Generale, McGraw-Hill
- Petrucci; Herring; Madura; Bissonnette CHIMICA GENERALE PRINCIPI ED APPLICAZIONI MODERNE, Editore: Piccin
- J. Kotz John, P. Treichel jr., G.C. Weaver, CHIMICA, Editore: Edises
Exercise book:
- A. Caselli, S. Rizzato, F. Tessore "Stechiometria dal testo di M. Freni e A. Sacco" Editore: Eises
- P. D'Arrigo, A. Famulari , C. Gambarotti , M. Scotti «Chimica. Esercizi e casi pratici» EdiSES
- A. Del Zotto «Esercizi di chimica generale» EdiSES
Recommended textbooks (any of the following):
- N. J. Tro, Chimica Un approccio molecolare, EdiSES
- F. Demartin, Fondamenti di Chimica Generale, EdiSES
- R. Chang, Fondamenti di Chimica Generale, McGraw-Hill
- Petrucci; Herring; Madura; Bissonnette CHIMICA GENERALE PRINCIPI ED APPLICAZIONI MODERNE, Editore: Piccin
- J. Kotz John, P. Treichel jr., G.C. Weaver, CHIMICA, Editore: Edises
Exercise book:
- A. Caselli, S. Rizzato, F. Tessore "Stechiometria dal testo di M. Freni e A. Sacco" Editore: Eises
- P. D'Arrigo, A. Famulari , C. Gambarotti , M. Scotti «Chimica. Esercizi e casi pratici» EdiSES
- A. Del Zotto «Esercizi di chimica generale» EdiSES
Assessment methods and Criteria
The exam is written (3 hours) and divided into three parts:
A) Basic knowledge on chemical science [nomenclature, identification and classification of chemical species]: 11 questions (word completion test) [11 points]
B) Description of fundamental properties of atoms, molecules and chemical compounds: 3 questions (open-ended questions) [11 points]
C) Problem-solving ability and application skills: 3 exercises (stoichiometry, pH / solubility / redox ...) [11 points]
A minimum score of 6 points in part A and 18 points as sum of parts A / B and C are compulsory to pass the test.
A) Basic knowledge on chemical science [nomenclature, identification and classification of chemical species]: 11 questions (word completion test) [11 points]
B) Description of fundamental properties of atoms, molecules and chemical compounds: 3 questions (open-ended questions) [11 points]
C) Problem-solving ability and application skills: 3 exercises (stoichiometry, pH / solubility / redox ...) [11 points]
A minimum score of 6 points in part A and 18 points as sum of parts A / B and C are compulsory to pass the test.
CHIM/03 - GENERAL AND INORGANIC CHEMISTRY - University credits: 8
Practicals: 32 hours
Lessons: 48 hours
Lessons: 48 hours
Professor:
Di Carlo Gabriele
Linea MZ
Responsible
Lesson period
First semester
Course syllabus
Frontal teaching (6 CFU, 48 hours frontal theorical lessons + 2 CFU, 32 hours frontal excercises)
Matter and its fundamental laws: states of aggregation; physical and chemical properties; intensive and extensive properties. Main units of measurement. Atoms, isotopes, molecules, ions, elements. Structure of the atom (introduction). Atomic and molecular mass. Compounds and mixtures. The mole and Avogadro's Number. Chemical formulae. Determination of the empirical formula and the molecular formula. Oxidation number and rules for its assignment. Nomenclature of inorganic compounds. Percentage weight composition of chemical compounds. Periodic table and periodic properties of elements (introduction). Chemical reactions and their balancing 'by trial and error'. Oxidation-reduction (redox) reactions and their balancing according to the 'semi-reaction method'. Definition of stoichiometry. Limiting reagent. Reaction yield. Chemical solutions. Concentration of solutions. Reactions in aqueous solutions (precipitation, acid-base and redox). Gaseous systems. Steam, ideal gas and real gas. Ideal gas laws (Boyle, Charles and Gay-Lussac) and the Ideal Gas Equation. Gas mixtures, Dalton's law of partial pressures. Molecular kinetic theory (outline). Van der Waals equation of real gases. Brief history of Atomic Theories: electromagnetic radiation; emission and absorption spectroscopy; Bohr's atomic model. Wave-particle duality. Heisenberg Principle. Atomic orbitals. Electronic spin. Pauli exclusion principle. Orbitals and energy levels in multi-electron atoms. Electronic configuration and the Aufbau principle. Periodic table and periodic properties of the elements (deepening): metals, non-metals and metalloids; groups and periods; atomic and ionic radii; ionisation energy and electronic affinity; electronegativity. The chemical bond. Ionic, metallic and covalent bonding. Bond polarity. Valence bond theory. Lewis structures and formal charge. Resonance. Exceptions to the octet rule. Hybridisation. VSEPR theory and molecular geometry. Theory of molecular orbitals (outline). Intermolecular forces. Phase diagrams (overview). Colligative properties (outline). Chemical equilibrium in liquid, gaseous and solid phases: homogeneous and heterogeneous equilibria; the equilibrium constant; Le Châtelier principle; factors influencing chemical equilibrium. Acids and bases according to the definitions of Arrhenius, Broensted/Lowry and Lewis. Strength of acids. pH scale. Determination of the pH of an aqueous solution: acid-base; salt hydrolysis; buffer solutions. Acid-base titration. Indicators. Solubility equilibria. Common ion effect. Brief introduction to electrochemistry: voltaic cells (batteries) and electrolytic cells, redox potentials under standard conditions, electromotive force, Nernst equation. Brief introduction to thermochemistry and kinetics: enthalpy, entropy, free energy and speed of reactions, spontaneity of a reaction and its prediction, dependence of the speed of a reaction on temperature (Arrhenius equation).
Matter and its fundamental laws: states of aggregation; physical and chemical properties; intensive and extensive properties. Main units of measurement. Atoms, isotopes, molecules, ions, elements. Structure of the atom (introduction). Atomic and molecular mass. Compounds and mixtures. The mole and Avogadro's Number. Chemical formulae. Determination of the empirical formula and the molecular formula. Oxidation number and rules for its assignment. Nomenclature of inorganic compounds. Percentage weight composition of chemical compounds. Periodic table and periodic properties of elements (introduction). Chemical reactions and their balancing 'by trial and error'. Oxidation-reduction (redox) reactions and their balancing according to the 'semi-reaction method'. Definition of stoichiometry. Limiting reagent. Reaction yield. Chemical solutions. Concentration of solutions. Reactions in aqueous solutions (precipitation, acid-base and redox). Gaseous systems. Steam, ideal gas and real gas. Ideal gas laws (Boyle, Charles and Gay-Lussac) and the Ideal Gas Equation. Gas mixtures, Dalton's law of partial pressures. Molecular kinetic theory (outline). Van der Waals equation of real gases. Brief history of Atomic Theories: electromagnetic radiation; emission and absorption spectroscopy; Bohr's atomic model. Wave-particle duality. Heisenberg Principle. Atomic orbitals. Electronic spin. Pauli exclusion principle. Orbitals and energy levels in multi-electron atoms. Electronic configuration and the Aufbau principle. Periodic table and periodic properties of the elements (deepening): metals, non-metals and metalloids; groups and periods; atomic and ionic radii; ionisation energy and electronic affinity; electronegativity. The chemical bond. Ionic, metallic and covalent bonding. Bond polarity. Valence bond theory. Lewis structures and formal charge. Resonance. Exceptions to the octet rule. Hybridisation. VSEPR theory and molecular geometry. Theory of molecular orbitals (outline). Intermolecular forces. Phase diagrams (overview). Colligative properties (outline). Chemical equilibrium in liquid, gaseous and solid phases: homogeneous and heterogeneous equilibria; the equilibrium constant; Le Châtelier principle; factors influencing chemical equilibrium. Acids and bases according to the definitions of Arrhenius, Broensted/Lowry and Lewis. Strength of acids. pH scale. Determination of the pH of an aqueous solution: acid-base; salt hydrolysis; buffer solutions. Acid-base titration. Indicators. Solubility equilibria. Common ion effect. Brief introduction to electrochemistry: voltaic cells (batteries) and electrolytic cells, redox potentials under standard conditions, electromotive force, Nernst equation. Brief introduction to thermochemistry and kinetics: enthalpy, entropy, free energy and speed of reactions, spontaneity of a reaction and its prediction, dependence of the speed of a reaction on temperature (Arrhenius equation).
Prerequisites for admission
The teaching of General, Inorganic Chemistry and Stoichiometry is the first Chemistry course faced by the student at the beginning of the degree program; therefore, it does not require prerequisites at university level. However, basic knowledge of Mathematics, Physics and Chemistry acquired in the secondary school is necessary for a more rapid and effective understanding of the topics.
Teaching methods
Lectures and frontal exercises (6 CFU, 48 hours frontal theoretical lessons + 2 CFU, 32 hours frontal excercises) take place in the classroom with projection of 'slides' in power point format and use of some innovative teaching tools (Wooclap, digital students' engagement platform, and Labster virtual lab simulation platform). Students have access to the professor's Ariel website, where various teaching materials, such as slides of lectures including frontal exercises, are regularly posted.
Teaching Resources
Lecture slides and exercises available on the professor's Ariel website.
Any university textbook in General Chemistry that addresses the program topics.
The following is a suggestion:
- Barone, Castellano, Della Pina, Demartin, Grasso, Maccato, Pepi, Piccinelli, Ruffo, Zampella. Fundamentals of General Chemistry; Ed. EdiSES
Any university textbook in General Chemistry that addresses the program topics.
The following is a suggestion:
- Barone, Castellano, Della Pina, Demartin, Grasso, Maccato, Pepi, Piccinelli, Ruffo, Zampella. Fundamentals of General Chemistry; Ed. EdiSES
Assessment methods and Criteria
The examination is a written test consisting of 5 stoichiometric exercises and 2 open theoretical questions, for a maximum total score of 30/30 (to which "cum laude" may be added, if applicable). The written test lasts 2 hours and 30 minutes. A minimum score of 18/30 is required to pass the exam. There are no oral tests. During the course, the student may have access to 2 in-process tests, i.e., 2 partial examination tests (the first one at the end of November and the second one at the end of the course). If the average mark of the two tests reaches at least 18/30, the exam is considered passed. These tests, therefore, allow the examination to be divided into two parts and reflect the same type of exercises as in the ordinary and extraordinary sessions. In addition to the 2 in-process tests, 7 examinations are set for each academic year in the ordinary and extraordinary sessions.
CHIM/03 - GENERAL AND INORGANIC CHEMISTRY - University credits: 8
Practicals: 32 hours
Lessons: 48 hours
Lessons: 48 hours
Professor:
Della Pina Cristina
Professor(s)
Reception:
prior appointment by e-mail
Room R102, Dip. Chimica, corpo A piano terra