Biology

Introduction to the course Historical background. The scientific method. Properties of matter. Measurement and units. Significant figures. Fundamental principles of chemistry Elements, compounds, and mixtures. States of matter. Lavoisier’s and Proust’s laws. Dalton’s atomic theory. Atoms and molecules. Basic structure of the atom: atomic number, mass number, isotopes, atomic mass. Compounds: molecular mass, the concept of the mole, Avogadro’s constant. Molar and mass ratios. Percent composition by mass, empirical and molecular formulas. Basic chemical reactions and their balancing. Limiting reagent. Theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic models. Thomson’s model. Waves and the electromagnetic spectrum. Planck’s equation. Photoelectric effect. Bohr’s atom and energy quantization. Uncertainty principle. Wave-particle duality. Introduction to wave mechanics. Schrödinger’s equation. Quantum numbers, wave functions, and atomic orbitals. Aufbau principle, Pauli exclusion principle, and Hund’s rule. Periodic properties The periodic table. Periodic properties of elements: atomic and ionic radius, ionization energy, electron affinity, metallic and magnetic properties. Chemical bonding theory Lewis structures. Octet rule and its exceptions. Formal charge and oxidation state. Bond length, bond energy, and bond order. Polar bonding and electronegativity. Covalent and ionic bonding. Resonance. Valence Bond (VB) theory. Hybrid orbitals and molecular geometry. VSEPR theory. Sigma (σ) and pi (π) bonds. Molecular Orbital (MO) theory. Thermodynamics Heat and work. Internal energy and the first law of thermodynamics. Enthalpy of reaction and formation. Hess’s law and its applications. Reversible and irreversible processes. Spontaneous transformations. Entropy. Second and third laws of thermodynamics. Gibbs free energy. Gaseous state Pressure. Ideal gas laws and the ideal gas equation. Gas mixtures. Dalton’s law. van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Physicochemical properties of liquids: boiling point, enthalpy of vaporization, vapor pressure and its temperature dependence (Clausius-Clapeyron equation). Specific heat. Heat and phase transitions. Heating curve of a pure substance at constant pressure. Phase diagram of pure chemical substances (e.g. H₂O). Structure and properties of water. Types of solids and their characteristics. Born-Haber cycle. Chemical equilibria in the gas phase Reaction quotient and equilibrium constant. Le Chatelier’s principle. Different types of equilibrium constants. Gibbs free energy and the equilibrium constant. Effect of changes in pressure, volume, amount of substance, and temperature (van’t Hoff equation). Heterogeneous equilibria. Solutions Concentration units and conversions. Solubility. Henry’s law. Thermodynamics of solutions. Colligative properties: boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure. Solutions of strong and weak electrolytes. van’t Hoff factor. Equilibria in aqueous solution Acid-base equilibria (Arrhenius, Brønsted-Lowry, and Lewis definitions). pH scale of aqueous solutions. Strong acids and bases. Relationship between structure and chemical behavior. Polyprotic acids. Weak acids and bases. Acid dissociation constants. Salt hydrolysis. Buffers. Titrations. Salt solubility. Solubility product constants (Ksp). Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Oxidation states. Balancing redox reactions using the ionic-electron method. Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of redox reactions. Galvanic and concentration cells. Examples of biologically relevant electrochemical phenomena. Potentiometric pH measurement. Introduction to chemical kinetics Reaction rate. Reaction order. Collision theory. Arrhenius equation. Activation energy. Reaction mechanisms. Catalysis. Enzymatic catalysis. Introduction to systematic inorganic chemistry Preparation and chemical behavior of main group elements.

Students are expected to possess a basic knowledge of mathematics and physics, including: First- and second-degree equations, percentage calculations, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (common and natural) and fundamental concepts of force, potential energy, and kinetic energy.

1) Chimica, VII edizione (EdiSES); Kotz, Treichel, Townsend, Treichel. 2) Chimica, X edizione (Piccin); Whitten, Davis, Peck, Stanley 3) Chimica, V edizione, (McGraw-Hill) Silberberg, Amateis, Licoccia   For stoichiometry:: 1) Manuale di Stechiometria (Piccin); Monti, Stefanelli, Viola. Some lecture slides and other teaching materials will be made available on the Moodle platform.

Attendance at lessons is not mandatory but strongly recommended

Assessment Methods The examination consists of a written test and an oral test, which must both be taken during the same exam session. The written test is designed to assess the student's level of understanding and ability to apply chemical models and methods through theoretical and stoichiometric exercises. It includes 6 to 8 numerical exercises and open-ended questions based on topics covered during the course, in accordance with the official syllabus. The test, which lasts 2 hours, includes one mandatory exercise on nomenclature and structural formulas, which must be passed with a score higher than 1.5 out of 4 in order for the entire test to be considered valid. The oral examination is intended to evaluate the student’s retention of course content, logical reasoning skills, and ability to explain and discuss concepts independently. Students are admitted to the oral exam only if they obtain a minimum score of 18/30 in the written test. Students regularly enrolled in the current academic year who attend the course may voluntarily access an exemption path, consisting of two intermediate tests during the semester, which substitute the final written exam. The indicative dates for these tests are: early December (first test) and mid-January (second test). In order to pass the exemption path, students must obtain a weighted average of at least 18/30 and a minimum score of 15/30 in each test. If the written test is passed through the exemption path, its validity extends to the entire winter exam session. Students with certified Specific Learning Disorders (SLD) are entitled to a 30% reduction in the quantitative content of the written test (or more, the exam procedures are agreed with the responsible offices), while the qualitative level remains unchanged. In these cases as well, the mandatory exercise on nomenclature and structural formulas remains required. The use of notes and textbooks is not permitted during the exam.

N/A

Lectures and in-class numerical exercises, also supported by tutors

Introduction to the course Historical background. The scientific method. Properties of matter. Measurement and units. Significant figures. Fundamental principles of chemistry Elements, compounds, and mixtures. States of matter. Lavoisier’s and Proust’s laws. Dalton’s atomic theory. Atoms and molecules. Basic structure of the atom: atomic number, mass number, isotopes, atomic mass. Compounds: molecular mass, the concept of the mole, Avogadro’s constant. Molar and mass ratios. Percent composition by mass, empirical and molecular formulas. Basic chemical reactions and their balancing. Limiting reagent. Theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic models. Thomson’s model. Waves and the electromagnetic spectrum. Planck’s equation. Photoelectric effect. Bohr’s atom and energy quantization. Uncertainty principle. Wave-particle duality. Introduction to wave mechanics. Schrödinger’s equation. Quantum numbers, wave functions, and atomic orbitals. Aufbau principle, Pauli exclusion principle, and Hund’s rule. Periodic properties The periodic table. Periodic properties of elements: atomic and ionic radius, ionization energy, electron affinity, metallic and magnetic properties. Chemical bonding theory Lewis structures. Octet rule and its exceptions. Formal charge and oxidation state. Bond length, bond energy, and bond order. Polar bonding and electronegativity. Covalent and ionic bonding. Resonance. Valence Bond (VB) theory. Hybrid orbitals and molecular geometry. VSEPR theory. Sigma (σ) and pi (π) bonds. Molecular Orbital (MO) theory. Thermodynamics Heat and work. Internal energy and the first law of thermodynamics. Enthalpy of reaction and formation. Hess’s law and its applications. Reversible and irreversible processes. Spontaneous transformations. Entropy. Second and third laws of thermodynamics. Gibbs free energy. Gaseous state Pressure. Ideal gas laws and the ideal gas equation. Gas mixtures. Dalton’s law. van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Physicochemical properties of liquids: boiling point, enthalpy of vaporization, vapor pressure and its temperature dependence (Clausius-Clapeyron equation). Specific heat. Heat and phase transitions. Heating curve of a pure substance at constant pressure. Phase diagram of pure chemical substances (e.g. H₂O). Structure and properties of water. Types of solids and their characteristics. Born-Haber cycle. Chemical equilibria in the gas phase Reaction quotient and equilibrium constant. Le Chatelier’s principle. Different types of equilibrium constants. Gibbs free energy and the equilibrium constant. Effect of changes in pressure, volume, amount of substance, and temperature (van’t Hoff equation). Heterogeneous equilibria. Solutions Concentration units and conversions. Solubility. Henry’s law. Thermodynamics of solutions. Colligative properties: boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure. Solutions of strong and weak electrolytes. van’t Hoff factor. Equilibria in aqueous solution Acid-base equilibria (Arrhenius, Brønsted-Lowry, and Lewis definitions). pH scale of aqueous solutions. Strong acids and bases. Relationship between structure and chemical behavior. Polyprotic acids. Weak acids and bases. Acid dissociation constants. Salt hydrolysis. Buffers. Titrations. Salt solubility. Solubility product constants (Ksp). Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Oxidation states. Balancing redox reactions using the ionic-electron method. Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of redox reactions. Galvanic and concentration cells. Examples of biologically relevant electrochemical phenomena. Potentiometric pH measurement. Introduction to chemical kinetics Reaction rate. Reaction order. Collision theory. Arrhenius equation. Activation energy. Reaction mechanisms. Catalysis. Enzymatic catalysis. Introduction to systematic inorganic chemistry Preparation and chemical behavior of main group elements.

Students are expected to possess a basic knowledge of mathematics and physics, including: First- and second-degree equations, percentage calculations, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (common and natural) and fundamental concepts of force, potential energy, and kinetic energy.

1) Chimica, VII edizione (EdiSES); Kotz, Treichel, Townsend, Treichel. 2) Chimica, X edizione (Piccin); Whitten, Davis, Peck, Stanley 3) Chimica, V edizione, (McGraw-Hill) Silberberg, Amateis, Licoccia   For stoichiometry:: 1) Manuale di Stechiometria (Piccin); Monti, Stefanelli, Viola. Some lecture slides and other teaching materials will be made available on the Moodle platform.

Attendance at lessons is not mandatory but strongly recommended

Assessment Methods The examination consists of a written test and an oral test, which must both be taken during the same exam session. The written test is designed to assess the student's level of understanding and ability to apply chemical models and methods through theoretical and stoichiometric exercises. It includes 6 to 8 numerical exercises and open-ended questions based on topics covered during the course, in accordance with the official syllabus. The test, which lasts 2 hours, includes one mandatory exercise on nomenclature and structural formulas, which must be passed with a score higher than 1.5 out of 4 in order for the entire test to be considered valid. The oral examination is intended to evaluate the student’s retention of course content, logical reasoning skills, and ability to explain and discuss concepts independently. Students are admitted to the oral exam only if they obtain a minimum score of 18/30 in the written test. Students regularly enrolled in the current academic year who attend the course may voluntarily access an exemption path, consisting of two intermediate tests during the semester, which substitute the final written exam. The indicative dates for these tests are: early December (first test) and mid-January (second test). In order to pass the exemption path, students must obtain a weighted average of at least 18/30 and a minimum score of 15/30 in each test. If the written test is passed through the exemption path, its validity extends to the entire winter exam session. Students with certified Specific Learning Disorders (SLD) are entitled to a 30% reduction in the quantitative content of the written test (or more, the exam procedures are agreed with the responsible offices), while the qualitative level remains unchanged. In these cases as well, the mandatory exercise on nomenclature and structural formulas remains required. The use of notes and textbooks is not permitted during the exam.

N/A

Lectures and in-class numerical exercises, also supported by tutors

Introduction to the course Historical background. The scientific method. Properties of matter. Measurement and units. Significant figures. Fundamental principles of chemistry Elements, compounds, and mixtures. States of matter. Lavoisier’s and Proust’s laws. Dalton’s atomic theory. Atoms and molecules. Basic structure of the atom: atomic number, mass number, isotopes, atomic mass. Compounds: molecular mass, the concept of the mole, Avogadro’s constant. Molar and mass ratios. Percent composition by mass, empirical and molecular formulas. Basic chemical reactions and their balancing. Limiting reagent. Theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic models. Thomson’s model. Waves and the electromagnetic spectrum. Planck’s equation. Photoelectric effect. Bohr’s atom and energy quantization. Uncertainty principle. Wave-particle duality. Introduction to wave mechanics. Schrödinger’s equation. Quantum numbers, wave functions, and atomic orbitals. Aufbau principle, Pauli exclusion principle, and Hund’s rule. Periodic properties The periodic table. Periodic properties of elements: atomic and ionic radius, ionization energy, electron affinity, metallic and magnetic properties. Chemical bonding theory Lewis structures. Octet rule and its exceptions. Formal charge and oxidation state. Bond length, bond energy, and bond order. Polar bonding and electronegativity. Covalent and ionic bonding. Resonance. Valence Bond (VB) theory. Hybrid orbitals and molecular geometry. VSEPR theory. Sigma (σ) and pi (π) bonds. Molecular Orbital (MO) theory. Thermodynamics Heat and work. Internal energy and the first law of thermodynamics. Enthalpy of reaction and formation. Hess’s law and its applications. Reversible and irreversible processes. Spontaneous transformations. Entropy. Second and third laws of thermodynamics. Gibbs free energy. Gaseous state Pressure. Ideal gas laws and the ideal gas equation. Gas mixtures. Dalton’s law. van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Physicochemical properties of liquids: boiling point, enthalpy of vaporization, vapor pressure and its temperature dependence (Clausius-Clapeyron equation). Specific heat. Heat and phase transitions. Heating curve of a pure substance at constant pressure. Phase diagram of pure chemical substances (e.g. H₂O). Structure and properties of water. Types of solids and their characteristics. Born-Haber cycle. Chemical equilibria in the gas phase Reaction quotient and equilibrium constant. Le Chatelier’s principle. Different types of equilibrium constants. Gibbs free energy and the equilibrium constant. Effect of changes in pressure, volume, amount of substance, and temperature (van’t Hoff equation). Heterogeneous equilibria. Solutions Concentration units and conversions. Solubility. Henry’s law. Thermodynamics of solutions. Colligative properties: boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure. Solutions of strong and weak electrolytes. van’t Hoff factor. Equilibria in aqueous solution Acid-base equilibria (Arrhenius, Brønsted-Lowry, and Lewis definitions). pH scale of aqueous solutions. Strong acids and bases. Relationship between structure and chemical behavior. Polyprotic acids. Weak acids and bases. Acid dissociation constants. Salt hydrolysis. Buffers. Titrations. Salt solubility. Solubility product constants (Ksp). Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Oxidation states. Balancing redox reactions using the ionic-electron method. Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of redox reactions. Galvanic and concentration cells. Examples of biologically relevant electrochemical phenomena. Potentiometric pH measurement. Introduction to chemical kinetics Reaction rate. Reaction order. Collision theory. Arrhenius equation. Activation energy. Reaction mechanisms. Catalysis. Enzymatic catalysis. Introduction to systematic inorganic chemistry Preparation and chemical behavior of main group elements.

Students are expected to possess a basic knowledge of mathematics and physics, including: First- and second-degree equations, percentage calculations, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (common and natural) and fundamental concepts of force, potential energy, and kinetic energy.

1) Chimica, VII edizione (EdiSES); Kotz, Treichel, Townsend, Treichel. 2) Chimica, X edizione (Piccin); Whitten, Davis, Peck, Stanley 3) Chimica, V edizione, (McGraw-Hill) Silberberg, Amateis, Licoccia   For stoichiometry:: 1) Manuale di Stechiometria (Piccin); Monti, Stefanelli, Viola. Some lecture slides and other teaching materials will be made available on the Moodle platform.

Attendance at lessons is not mandatory but strongly recommended

Assessment Methods The examination consists of a written test and an oral test, which must both be taken during the same exam session. The written test is designed to assess the student's level of understanding and ability to apply chemical models and methods through theoretical and stoichiometric exercises. It includes 6 to 8 numerical exercises and open-ended questions based on topics covered during the course, in accordance with the official syllabus. The test, which lasts 2 hours, includes one mandatory exercise on nomenclature and structural formulas, which must be passed with a score higher than 1.5 out of 4 in order for the entire test to be considered valid. The oral examination is intended to evaluate the student’s retention of course content, logical reasoning skills, and ability to explain and discuss concepts independently. Students are admitted to the oral exam only if they obtain a minimum score of 18/30 in the written test. Students regularly enrolled in the current academic year who attend the course may voluntarily access an exemption path, consisting of two intermediate tests during the semester, which substitute the final written exam. The indicative dates for these tests are: early December (first test) and mid-January (second test). In order to pass the exemption path, students must obtain a weighted average of at least 18/30 and a minimum score of 15/30 in each test. If the written test is passed through the exemption path, its validity extends to the entire winter exam session. Students with certified Specific Learning Disorders (SLD) are entitled to a 30% reduction in the quantitative content of the written test (or more, the exam procedures are agreed with the responsible offices), while the qualitative level remains unchanged. In these cases as well, the mandatory exercise on nomenclature and structural formulas remains required. The use of notes and textbooks is not permitted during the exam.

N/A

Lectures and in-class numerical exercises, also supported by tutors

Introduction to the course Historical background. The scientific method. Properties of matter. Measurement and units. Significant figures. Fundamental principles of chemistry Elements, compounds, and mixtures. States of matter. Lavoisier’s and Proust’s laws. Dalton’s atomic theory. Atoms and molecules. Basic structure of the atom: atomic number, mass number, isotopes, atomic mass. Compounds: molecular mass, the concept of the mole, Avogadro’s constant. Molar and mass ratios. Percent composition by mass, empirical and molecular formulas. Basic chemical reactions and their balancing. Limiting reagent. Theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic models. Thomson’s model. Waves and the electromagnetic spectrum. Planck’s equation. Photoelectric effect. Bohr’s atom and energy quantization. Uncertainty principle. Wave-particle duality. Introduction to wave mechanics. Schrödinger’s equation. Quantum numbers, wave functions, and atomic orbitals. Aufbau principle, Pauli exclusion principle, and Hund’s rule. Periodic properties The periodic table. Periodic properties of elements: atomic and ionic radius, ionization energy, electron affinity, metallic and magnetic properties. Chemical bonding theory Lewis structures. Octet rule and its exceptions. Formal charge and oxidation state. Bond length, bond energy, and bond order. Polar bonding and electronegativity. Covalent and ionic bonding. Resonance. Valence Bond (VB) theory. Hybrid orbitals and molecular geometry. VSEPR theory. Sigma (σ) and pi (π) bonds. Molecular Orbital (MO) theory. Thermodynamics Heat and work. Internal energy and the first law of thermodynamics. Enthalpy of reaction and formation. Hess’s law and its applications. Reversible and irreversible processes. Spontaneous transformations. Entropy. Second and third laws of thermodynamics. Gibbs free energy. Gaseous state Pressure. Ideal gas laws and the ideal gas equation. Gas mixtures. Dalton’s law. van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Physicochemical properties of liquids: boiling point, enthalpy of vaporization, vapor pressure and its temperature dependence (Clausius-Clapeyron equation). Specific heat. Heat and phase transitions. Heating curve of a pure substance at constant pressure. Phase diagram of pure chemical substances (e.g. H₂O). Structure and properties of water. Types of solids and their characteristics. Born-Haber cycle. Chemical equilibria in the gas phase Reaction quotient and equilibrium constant. Le Chatelier’s principle. Different types of equilibrium constants. Gibbs free energy and the equilibrium constant. Effect of changes in pressure, volume, amount of substance, and temperature (van’t Hoff equation). Heterogeneous equilibria. Solutions Concentration units and conversions. Solubility. Henry’s law. Thermodynamics of solutions. Colligative properties: boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure. Solutions of strong and weak electrolytes. van’t Hoff factor. Equilibria in aqueous solution Acid-base equilibria (Arrhenius, Brønsted-Lowry, and Lewis definitions). pH scale of aqueous solutions. Strong acids and bases. Relationship between structure and chemical behavior. Polyprotic acids. Weak acids and bases. Acid dissociation constants. Salt hydrolysis. Buffers. Titrations. Salt solubility. Solubility product constants (Ksp). Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Oxidation states. Balancing redox reactions using the ionic-electron method. Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of redox reactions. Galvanic and concentration cells. Examples of biologically relevant electrochemical phenomena. Potentiometric pH measurement. Introduction to chemical kinetics Reaction rate. Reaction order. Collision theory. Arrhenius equation. Activation energy. Reaction mechanisms. Catalysis. Enzymatic catalysis. Introduction to systematic inorganic chemistry Preparation and chemical behavior of main group elements.

Students are expected to possess a basic knowledge of mathematics and physics, including: First- and second-degree equations, percentage calculations, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (common and natural) and fundamental concepts of force, potential energy, and kinetic energy.

1) Chimica, VII edizione (EdiSES); Kotz, Treichel, Townsend, Treichel. 2) Chimica, X edizione (Piccin); Whitten, Davis, Peck, Stanley 3) Chimica, V edizione, (McGraw-Hill) Silberberg, Amateis, Licoccia   For stoichiometry:: 1) Manuale di Stechiometria (Piccin); Monti, Stefanelli, Viola. Some lecture slides and other teaching materials will be made available on the Moodle platform.

Attendance at lessons is not mandatory but strongly recommended

Assessment Methods The examination consists of a written test and an oral test, which must both be taken during the same exam session. The written test is designed to assess the student's level of understanding and ability to apply chemical models and methods through theoretical and stoichiometric exercises. It includes 6 to 8 numerical exercises and open-ended questions based on topics covered during the course, in accordance with the official syllabus. The test, which lasts 2 hours, includes one mandatory exercise on nomenclature and structural formulas, which must be passed with a score higher than 1.5 out of 4 in order for the entire test to be considered valid. The oral examination is intended to evaluate the student’s retention of course content, logical reasoning skills, and ability to explain and discuss concepts independently. Students are admitted to the oral exam only if they obtain a minimum score of 18/30 in the written test. Students regularly enrolled in the current academic year who attend the course may voluntarily access an exemption path, consisting of two intermediate tests during the semester, which substitute the final written exam. The indicative dates for these tests are: early December (first test) and mid-January (second test). In order to pass the exemption path, students must obtain a weighted average of at least 18/30 and a minimum score of 15/30 in each test. If the written test is passed through the exemption path, its validity extends to the entire winter exam session. Students with certified Specific Learning Disorders (SLD) are entitled to a 30% reduction in the quantitative content of the written test (or more, the exam procedures are agreed with the responsible offices), while the qualitative level remains unchanged. In these cases as well, the mandatory exercise on nomenclature and structural formulas remains required. The use of notes and textbooks is not permitted during the exam.

N/A

Lectures and in-class numerical exercises, also supported by tutors