General and Inorganic Chemistry
A.Y. 2025/2026
Learning objectives
The course, in line with the educational objectives of the degree course, provides the theoretical and experimental foundations of the chemical sciences and familiarises students with the language of this discipline. It aims to provide students with the basic knowledge in chemistry in order to be able to tackle other chemical and professional courses specific to the degree course (e.g. in biochemistry and molecular biology) with rigour and scientific method.
Expected learning outcomes
The desired learning outcomes at the end of the course will mainly consist of understanding the chemical description of phenomena of both general and more strictly biological interest. These aspects are expected to be complemented by the ability to critically evaluate and understand, also by means of a correct and well-considered bibliographic research, texts of biotechnological interest that require knowledge of both implicit and explicit chemical aspects. The course also aims to offer training, more generally understood though moving from the chemical sphere, in the ability to process quantitative aspects of the scientific method and the ability to communicate these aspects through scientifically correct language and an adequate graphical representation of the data and information. Taken as a whole, these aspects will enable the student to pursue the study of various chemical-biological and quantitative topics independently both in the context of the degree course and in a professional context.
Lesson period: First semester
Assessment methods: Esame
Assessment result: voto verbalizzato in trentesimi
Single course
This course cannot be attended as a single course. Please check our list of single courses to find the ones available for enrolment.
Course syllabus and organization
Single session
Responsible
Lesson period
First semester
Course syllabus
Qualitative and Quantitative Aspects of Chemistry
History and evolution of chemistry from alchemy to the present.
Modern chemistry as a quantitative discipline: units of measurement in chemistry and their dimensionality, precision and accuracy, use of significant figures in measurements, graphical representation of measurements.
Dalton's atomic theory and laws of chemical proportions. Systematization of atomic properties: Mendeleev's periodic table. Atoms, chemical elements, and isotopes: atomic number and atomic weight. Modern interpretation of the periodic table. Overview of nucleogenesis.
Molecules, compounds, and molecular formulas. Molecular mass, molecular weight. Avogadro's number and mole. Molarity.
Mixtures and compounds. Nomenclature of ionic compounds and an introduction to the nomenclature of commonly occurring molecular compounds in inorganic chemistry.
Chemical reactions and equations. Introduction to chemical equilibrium, thermochemistry and reactions in aqueous solution.
Atomic and Molecular Structure of Matter
Structure of the atom. Subatomic particles. Electromagnetic radiation and atomic spectra. Bohr's atom. Quantum mechanical description of the atom and wave functions.
Atomic configuration. Quantum numbers and orbitals. Pauli exclusion principle and Hund's rule. Electronic configuration of elements and the periodic table. Periodic properties: atomic and ionic size, ionization energy, and electron affinity.
Chemical bonding and molecular structure. Electron distribution. Ionic, covalent, and metallic bonding. Lewis symbols and structures. Octet rule. Resonance. Electronegativity. Dipole moment and molecular polarity. Molecular shapes (VSEPR theory). Valence bond theory. Hybrid orbitals. Single and multiple bonds. Some structures of inorganic and organic molecules (amino acids, nitrogenous bases). Molecular orbitals.
Intermolecular interactions: electrostatic forces, van der Waals interactions. Hydrogen bonding.
States of Matter and Chemistry of Aqueous Solutions (I)
General properties of gases. Ideal gas laws. Ideal gas equation of state. Gas mixtures and partial pressures. Kinetic theory of gases. Gas diffusion and Graham's law. Non-ideal gases and van der Waals equation.
Crystal lattices and unit cells in solids. Ionic, covalent, molecular, and metallic solids. Introduction to X-ray diffraction. Lattice energy and thermal decomposition of solids.
General properties of liquids: cohesion, adhesion, surface tension, viscosity. Vapor pressure and enthalpy of evaporation. The boiling point. Critical temperature and pressure.
Phase transitions and phase diagrams.
Solutions and solvation principles. Concentration and its units of measurement. Dilutions. Enthalpy of dissolution. Solubility as a function of pressure and temperature. Raoult's law. Colligative properties of solutions. Colloids.
Control of Chemical Reactions
Energy and its units of measurement. First law of thermodynamics. Internal energy. General concepts of thermochemistry and energy in chemical reactions. Enthalpy and heat of reaction: Hess's law. Specific heat and heat capacity. Calorimetry.
Chemical thermodynamics. Second law of thermodynamics. Gibbs free energy and spontaneity criteria. Chemical potential. Thermodynamic description of phase transitions.
Chemical equilibrium. Law of mass action. Equilibrium constant. Reaction quotient. Le Chatelier's principle.
Chemical kinetics. Reaction rate. Reaction mechanism. Activation energy.
Balancing chemical reactions using weight methods. Reaction yield. Limiting reagent.
Chemistry of Aqueous Solutions (II)
Acid-base theories of Arrhenius and Brönsted-Lowry. Acid-base equilibrium and ion product of water. Concentrations of acidic and basic species in aqueous solution: pH, pOH, and pKw. Strong and weak acids and bases and their relationship to molecular structure. Polyprotic acids.
Lewis acids and bases. "Dative" covalent bonds and coordination compounds. The crystal field theory.
Salt hydrolysis and its influence on pH. Buffer solutions. Acid-base titrations and indicators.
Unsoluble salts, solubility product (Kps) and precipitation. Common ion effect. Effect of pH on solubility.
Electrochemistry
Electrochemical and electrolytic cells. Standard reduction potentials. Electromotive force of a cell. Free energy and electromotive force. Voltaic cells under non-standard conditions: Nernst equation. Electromotive force and equilibrium constant. Concentration cells. Potentiometric determinations of Kps and pH. Indicator and reference electrodes. pH meter.
Electrolysis and its laws. Electrolysis of molten salts and aqueous solutions. Electrolysis of water. Common batteries and accumulators. Metal corrosion.
Nuclear Chemistry
Nuclear reactions and radioactive decay. Stability of atomic nuclei. Rate of radioactive decay. Artificial nuclear reactions. Applications of nuclear chemistry in medicine and analytical fields.
Bioinorganic Chemistry
Essential elements for life and the metallome. Essential and non-essential elements. Role of metals in biological systems. Amino acids and their role as metal ligands in the body.
Entatic state theory in enzymes. Enzyme kinetics. Enzymes stability and function. Competitive and non-competitive inhibitors of the enzyme activity.
Oxygen transport and accumulation in the blood: iron coordination in hemoglobin and myoglobin. Cooperative effect. Influence of pH and Bohr effect.
History and evolution of chemistry from alchemy to the present.
Modern chemistry as a quantitative discipline: units of measurement in chemistry and their dimensionality, precision and accuracy, use of significant figures in measurements, graphical representation of measurements.
Dalton's atomic theory and laws of chemical proportions. Systematization of atomic properties: Mendeleev's periodic table. Atoms, chemical elements, and isotopes: atomic number and atomic weight. Modern interpretation of the periodic table. Overview of nucleogenesis.
Molecules, compounds, and molecular formulas. Molecular mass, molecular weight. Avogadro's number and mole. Molarity.
Mixtures and compounds. Nomenclature of ionic compounds and an introduction to the nomenclature of commonly occurring molecular compounds in inorganic chemistry.
Chemical reactions and equations. Introduction to chemical equilibrium, thermochemistry and reactions in aqueous solution.
Atomic and Molecular Structure of Matter
Structure of the atom. Subatomic particles. Electromagnetic radiation and atomic spectra. Bohr's atom. Quantum mechanical description of the atom and wave functions.
Atomic configuration. Quantum numbers and orbitals. Pauli exclusion principle and Hund's rule. Electronic configuration of elements and the periodic table. Periodic properties: atomic and ionic size, ionization energy, and electron affinity.
Chemical bonding and molecular structure. Electron distribution. Ionic, covalent, and metallic bonding. Lewis symbols and structures. Octet rule. Resonance. Electronegativity. Dipole moment and molecular polarity. Molecular shapes (VSEPR theory). Valence bond theory. Hybrid orbitals. Single and multiple bonds. Some structures of inorganic and organic molecules (amino acids, nitrogenous bases). Molecular orbitals.
Intermolecular interactions: electrostatic forces, van der Waals interactions. Hydrogen bonding.
States of Matter and Chemistry of Aqueous Solutions (I)
General properties of gases. Ideal gas laws. Ideal gas equation of state. Gas mixtures and partial pressures. Kinetic theory of gases. Gas diffusion and Graham's law. Non-ideal gases and van der Waals equation.
Crystal lattices and unit cells in solids. Ionic, covalent, molecular, and metallic solids. Introduction to X-ray diffraction. Lattice energy and thermal decomposition of solids.
General properties of liquids: cohesion, adhesion, surface tension, viscosity. Vapor pressure and enthalpy of evaporation. The boiling point. Critical temperature and pressure.
Phase transitions and phase diagrams.
Solutions and solvation principles. Concentration and its units of measurement. Dilutions. Enthalpy of dissolution. Solubility as a function of pressure and temperature. Raoult's law. Colligative properties of solutions. Colloids.
Control of Chemical Reactions
Energy and its units of measurement. First law of thermodynamics. Internal energy. General concepts of thermochemistry and energy in chemical reactions. Enthalpy and heat of reaction: Hess's law. Specific heat and heat capacity. Calorimetry.
Chemical thermodynamics. Second law of thermodynamics. Gibbs free energy and spontaneity criteria. Chemical potential. Thermodynamic description of phase transitions.
Chemical equilibrium. Law of mass action. Equilibrium constant. Reaction quotient. Le Chatelier's principle.
Chemical kinetics. Reaction rate. Reaction mechanism. Activation energy.
Balancing chemical reactions using weight methods. Reaction yield. Limiting reagent.
Chemistry of Aqueous Solutions (II)
Acid-base theories of Arrhenius and Brönsted-Lowry. Acid-base equilibrium and ion product of water. Concentrations of acidic and basic species in aqueous solution: pH, pOH, and pKw. Strong and weak acids and bases and their relationship to molecular structure. Polyprotic acids.
Lewis acids and bases. "Dative" covalent bonds and coordination compounds. The crystal field theory.
Salt hydrolysis and its influence on pH. Buffer solutions. Acid-base titrations and indicators.
Unsoluble salts, solubility product (Kps) and precipitation. Common ion effect. Effect of pH on solubility.
Electrochemistry
Electrochemical and electrolytic cells. Standard reduction potentials. Electromotive force of a cell. Free energy and electromotive force. Voltaic cells under non-standard conditions: Nernst equation. Electromotive force and equilibrium constant. Concentration cells. Potentiometric determinations of Kps and pH. Indicator and reference electrodes. pH meter.
Electrolysis and its laws. Electrolysis of molten salts and aqueous solutions. Electrolysis of water. Common batteries and accumulators. Metal corrosion.
Nuclear Chemistry
Nuclear reactions and radioactive decay. Stability of atomic nuclei. Rate of radioactive decay. Artificial nuclear reactions. Applications of nuclear chemistry in medicine and analytical fields.
Bioinorganic Chemistry
Essential elements for life and the metallome. Essential and non-essential elements. Role of metals in biological systems. Amino acids and their role as metal ligands in the body.
Entatic state theory in enzymes. Enzyme kinetics. Enzymes stability and function. Competitive and non-competitive inhibitors of the enzyme activity.
Oxygen transport and accumulation in the blood: iron coordination in hemoglobin and myoglobin. Cooperative effect. Influence of pH and Bohr effect.
Prerequisites for admission
The course does not require any prior knowledge, with the exception of mathematical knowledge from upper secondary education.
Teaching methods
The course consists of lectures including sample exercises by the lecturer. Attendance of lectures is compulsory
Teaching Resources
Theory text
N. J. Tro, Chemistry a Molecular Approach; Ed. EdiSES.
J. C. Kotz, Chemistry; Ed.EdiSES.
Textbook on Stoichiometry
G. Marcì, L. Palmisano, F. Ruffo, Stoichiometry; Ed. EdiSES.
P. Michelin Lausarot, G. A. Vaglio, Stoichiometry for General Chemistry; Ed. Piccin.
N. J. Tro, Chemistry a Molecular Approach; Ed. EdiSES.
J. C. Kotz, Chemistry; Ed.EdiSES.
Textbook on Stoichiometry
G. Marcì, L. Palmisano, F. Ruffo, Stoichiometry; Ed. EdiSES.
P. Michelin Lausarot, G. A. Vaglio, Stoichiometry for General Chemistry; Ed. Piccin.
Assessment methods and Criteria
The examination consists of a written test divided into two parts; the first is conducted by means of a multiple-choice test, the second by means of open-ended questions. Passing the first part of the test constitutes a preparatory step for taking the second part. The final mark will be calculated on the basis of the results of the two tests. The final mark is expressed in thirtieths as the sum of the two parts.
CHIM/03 - GENERAL AND INORGANIC CHEMISTRY - University credits: 8
Lessons: 64 hours
Professor:
Mollica Luca
Professor(s)