GENERAL AND INORGANIC CHEMISTRY

Course objectives

The students at the end of the course will be able: A) Knowledge and understanding To know the most common chemical elements, their properties and the way they behave in simple chemical reactions, being able to solve exercises on stoichiometric calculations and equilibrium reactions in solution; To know the principal classes of substances (acids, bases and salts) and their behaviour in aqueous solutions by applying fundamental thermodynamic properties; To understand qualitative and quantitative aspects of chemical transformations, as described in lessons and tutorials; To use reference basic chemistry manuals and handbooks to understand more advanced courses. B) Applying knowledge and understanding To apply the basic chemical knowledge to correlate the macroscopic properties and the elemental structure of matter at the level of chemical elements and simple molecules. C) Making judgements To critical think through the historical survey of the main discoveries in chemistry To learn by questioning. D) Communication skills To communicate the acquired theoretical and the experimental knowledge. E) Learning skills To learn the specific terminology; To make the logical connections between the topics covered; To identify the most relevant topics.

Channel 1
ANDREA CICCIOLI Lecturers' profile
ANDREA CICCIOLI Lecturers' profile
Channel 2
GIUSEPPE POMARICO Lecturers' profile

Program - Frequency - Exams

Course program
Introduction to the course. Historical notes. Scientific method, properties of matter, measurement and units of measurement, significant figures. Fundamental principles of chemistry Elements, compounds and mixtures, states of aggregation of matter, Lavoisier's law, Proust's law, Dalton's atomic theory. Atoms, molecules, concept of mole and Avogadro's number. Molar and ponderal ratios. Percentage composition by weight, minimum formula and molecular formula. Simple chemical reactions and their balancing, limiting reagent, theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic structure, Thomson model, waves and electromagnetic spectrum, atomic spectra, Planck equation, photoelectric effect, Bohr atom and quantization of energy. Uncertainty principle, wave-particle duality, notes on wave mechanics, Schrodinger equation. Quantum numbers, wave functions and atomic orbitals. Construction (Aufbau) of the electronic structure of atoms, Pauli exclusion principle and Hund's rule. Periodic Properties Periodic table, periodic properties of the elements (atomic radius, ionic radius, ionization energy, electron affinity, metallic properties and magnetic properties. Chemical bond theory Lewis theory, octet rule and exceptions, formal charge and oxidation state, order, length and bond energy. Polar bond and electronegativity, covalent and ionic bond. Resonance. Valence bond theory (VB), hybrid orbitals and shape of molecules, VSEPR theory. Covalent bonds sigma (σ) and pi (π). Molecular orbital theory (MO). Thermodynamics Heat and work. Internal energy and 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 law of thermodynamics. Gibbs free energy. Gaseous state Pressure, ideal gas laws and equation of state of ideal gases, gas mixtures, Dalton's law, van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Chemical-physical properties of liquids (boiling point, evaporation enthalpy, vapor pressure and its dependence on temperature: Clausius-Clapeyron equation). Specific heat. Heat and phase transitions; heating diagram of a pure species at constant pressure; state diagram of pure chemical species (H2O); structure and properties of water. Chemical equilibria in the gaseous phase Reaction quotient and equilibrium constant, Le Chatelier's principle; different types of equilibrium constant. Gibbs free energy and equilibrium constant, effect of the variation of P, V, n and T (van't Hoff equation). Heterogeneous equilibria. Solutions Units of measurement of concentration, conversion between the different units of measurement of concentration, solubility. Henry's law, thermodynamics of solutions, colligative properties (bullioscopic rise, freezing point lowering, lowering of vapor pressure, osmotic pressure). Solutions of strong and weak electrolytes, van't Hoff coefficient. Solution equilibria Acid-base equilibria (Arrhenius, Brønsted-Lowry and Lewis definitions). pH scale of aqueous solutions, strong acids and bases, correlations between structure and chemical behavior, polyprotic acids. Weak acids and bases, acidity constants, salt hydrolysis, buffers, titrations. Solubility of salts. Solubility products. Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of a redox reaction. Chemical and concentration cells. Examples of biological phenomena with electrochemical implications. Potentiometric measurement of pH. Balancing redox reactions. Notes on chemical kinetics. Reaction rate, reaction order, collision theory, Arrhenius law; activation energy; reaction mechanism; catalysis; enzymatic catalysis. Notes on systematic inorganic chemistry Preparation and chemical behavior of the elements of the main groups.
Prerequisites
Basic knowledge of mathematics and physics: first and second degree equations, calculation of percentages, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (decimal, natural), concept of force, potential and kinetic energy.
Books
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 
Frequency
Attendance at lessons is not mandatory but strongly recommended
Exam mode
The exam consists of a separate written test and oral test. The written test, aimed at assessing the level of learning and understanding in the application of chemical models and methods in theoretical and stoichiometric exercises, consists of 6-8 numerical exercises and open questions on topics covered during the lessons, as per the course program. The written test (duration 2 hours), will contain a mandatory exercise for the purpose of passing the test itself (score higher than 1.5/4), regarding nomenclature and structural formulas. The oral test is aimed at assessing the degree of memorization of the contents and the logical skills acquired by the student as well as his ability to expose and explain the concepts independently. Those who pass the written test with a score of at least 18/30 are admitted to take the oral test. The oral test must be taken in the same session in which the written test was passed, under penalty of forfeiture of validity of the latter. Attending students can voluntarily access the exemption path (2 intermediate tests during the semester in place of the final written test; indicative dates: early December (1st exemption), mid-January (2nd exemption). The exemption test is passed with a score of at least 18/30 and the average of the two exemption tests constitutes the score of the written test. In the case of a written test passed through exemptions, its validity is extended to the entire winter session. For students with DSA, the exam procedures are agreed with the responsible offices.
Bibliography
N/A
Lesson mode
Lectures and numerical exercises in class.
GIUSEPPE POMARICO Lecturers' profile

Program - Frequency - Exams

Course program
Introduction to the course. Historical notes. Scientific method, properties of matter, measurement and units of measurement, significant figures. Fundamental principles of chemistry Elements, compounds and mixtures, states of aggregation of matter, Lavoisier's law, Proust's law, Dalton's atomic theory. Atoms, molecules, concept of mole and Avogadro's number. Molar and ponderal ratios. Percentage composition by weight, minimum formula and molecular formula. Simple chemical reactions and their balancing, limiting reagent, theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic structure, Thomson model, waves and electromagnetic spectrum, atomic spectra, Planck equation, photoelectric effect, Bohr atom and quantization of energy. Uncertainty principle, wave-particle duality, notes on wave mechanics, Schrodinger equation. Quantum numbers, wave functions and atomic orbitals. Construction (Aufbau) of the electronic structure of atoms, Pauli exclusion principle and Hund's rule. Periodic Properties Periodic table, periodic properties of the elements (atomic radius, ionic radius, ionization energy, electron affinity, metallic properties and magnetic properties. Chemical bond theory Lewis theory, octet rule and exceptions, formal charge and oxidation state, order, length and bond energy. Polar bond and electronegativity, covalent and ionic bond. Resonance. Valence bond theory (VB), hybrid orbitals and shape of molecules, VSEPR theory. Covalent bonds sigma (σ) and pi (π). Molecular orbital theory (MO). Thermodynamics Heat and work. Internal energy and 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 law of thermodynamics. Gibbs free energy. Gaseous state Pressure, ideal gas laws and equation of state of ideal gases, gas mixtures, Dalton's law, van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Chemical-physical properties of liquids (boiling point, evaporation enthalpy, vapor pressure and its dependence on temperature: Clausius-Clapeyron equation). Specific heat. Heat and phase transitions; heating diagram of a pure species at constant pressure; state diagram of pure chemical species (H2O); structure and properties of water. Chemical equilibria in the gaseous phase Reaction quotient and equilibrium constant, Le Chatelier's principle; different types of equilibrium constant. Gibbs free energy and equilibrium constant, effect of the variation of P, V, n and T (van't Hoff equation). Heterogeneous equilibria. Solutions Units of measurement of concentration, conversion between the different units of measurement of concentration, solubility. Henry's law, thermodynamics of solutions, colligative properties (bullioscopic rise, freezing point lowering, lowering of vapor pressure, osmotic pressure). Solutions of strong and weak electrolytes, van't Hoff coefficient. Solution equilibria Acid-base equilibria (Arrhenius, Brønsted-Lowry and Lewis definitions). pH scale of aqueous solutions, strong acids and bases, correlations between structure and chemical behavior, polyprotic acids. Weak acids and bases, acidity constants, salt hydrolysis, buffers, titrations. Solubility of salts. Solubility products. Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of a redox reaction. Chemical and concentration cells. Examples of biological phenomena with electrochemical implications. Potentiometric measurement of pH. Balancing redox reactions. Notes on chemical kinetics. Reaction rate, reaction order, collision theory, Arrhenius law; activation energy; reaction mechanism; catalysis; enzymatic catalysis. Notes on systematic inorganic chemistry Preparation and chemical behavior of the elements of the main groups.
Prerequisites
Basic knowledge of mathematics and physics: first and second degree equations, calculation of percentages, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (decimal, natural), concept of force, potential and kinetic energy.
Books
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 
Frequency
Attendance at lessons is not mandatory but strongly recommended
Exam mode
The exam consists of a separate written test and oral test. The written test, aimed at assessing the level of learning and understanding in the application of chemical models and methods in theoretical and stoichiometric exercises, consists of 6-8 numerical exercises and open questions on topics covered during the lessons, as per the course program. The written test (duration 2 hours), will contain a mandatory exercise for the purpose of passing the test itself (score higher than 1.5/4), regarding nomenclature and structural formulas. The oral test is aimed at assessing the degree of memorization of the contents and the logical skills acquired by the student as well as his ability to expose and explain the concepts independently. Those who pass the written test with a score of at least 18/30 are admitted to take the oral test. The oral test must be taken in the same session in which the written test was passed, under penalty of forfeiture of validity of the latter. Attending students can voluntarily access the exemption path (2 intermediate tests during the semester in place of the final written test; indicative dates: early December (1st exemption), mid-January (2nd exemption). The exemption test is passed with a score of at least 18/30 and the average of the two exemption tests constitutes the score of the written test. In the case of a written test passed through exemptions, its validity is extended to the entire winter session. For students with DSA, the exam procedures are agreed with the responsible offices.
Bibliography
N/A
Lesson mode
Lectures and numerical exercises in class.
Channel 3
MARIA CHIARA DI GREGORIO Lecturers' profile
MARIA CHIARA DI GREGORIO Lecturers' profile
Channel 4
GIUSEPPE POMARICO Lecturers' profile

Program - Frequency - Exams

Course program
Introduction to the course. Historical notes. Scientific method, properties of matter, measurement and units of measurement, significant figures. Fundamental principles of chemistry Elements, compounds and mixtures, states of aggregation of matter, Lavoisier's law, Proust's law, Dalton's atomic theory. Atoms, molecules, concept of mole and Avogadro's number. Molar and ponderal ratios. Percentage composition by weight, minimum formula and molecular formula. Simple chemical reactions and their balancing, limiting reagent, theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic structure, Thomson model, waves and electromagnetic spectrum, atomic spectra, Planck equation, photoelectric effect, Bohr atom and quantization of energy. Uncertainty principle, wave-particle duality, notes on wave mechanics, Schrodinger equation. Quantum numbers, wave functions and atomic orbitals. Construction (Aufbau) of the electronic structure of atoms, Pauli exclusion principle and Hund's rule. Periodic Properties Periodic table, periodic properties of the elements (atomic radius, ionic radius, ionization energy, electron affinity, metallic properties and magnetic properties. Chemical bond theory Lewis theory, octet rule and exceptions, formal charge and oxidation state, order, length and bond energy. Polar bond and electronegativity, covalent and ionic bond. Resonance. Valence bond theory (VB), hybrid orbitals and shape of molecules, VSEPR theory. Covalent bonds sigma (σ) and pi (π). Molecular orbital theory (MO). Thermodynamics Heat and work. Internal energy and 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 law of thermodynamics. Gibbs free energy. Gaseous state Pressure, ideal gas laws and equation of state of ideal gases, gas mixtures, Dalton's law, van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Chemical-physical properties of liquids (boiling point, evaporation enthalpy, vapor pressure and its dependence on temperature: Clausius-Clapeyron equation). Specific heat. Heat and phase transitions; heating diagram of a pure species at constant pressure; state diagram of pure chemical species (H2O); structure and properties of water. Chemical equilibria in the gaseous phase Reaction quotient and equilibrium constant, Le Chatelier's principle; different types of equilibrium constant. Gibbs free energy and equilibrium constant, effect of the variation of P, V, n and T (van't Hoff equation). Heterogeneous equilibria. Solutions Units of measurement of concentration, conversion between the different units of measurement of concentration, solubility. Henry's law, thermodynamics of solutions, colligative properties (bullioscopic rise, freezing point lowering, lowering of vapor pressure, osmotic pressure). Solutions of strong and weak electrolytes, van't Hoff coefficient. Solution equilibria Acid-base equilibria (Arrhenius, Brønsted-Lowry and Lewis definitions). pH scale of aqueous solutions, strong acids and bases, correlations between structure and chemical behavior, polyprotic acids. Weak acids and bases, acidity constants, salt hydrolysis, buffers, titrations. Solubility of salts. Solubility products. Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of a redox reaction. Chemical and concentration cells. Examples of biological phenomena with electrochemical implications. Potentiometric measurement of pH. Balancing redox reactions. Notes on chemical kinetics. Reaction rate, reaction order, collision theory, Arrhenius law; activation energy; reaction mechanism; catalysis; enzymatic catalysis. Notes on systematic inorganic chemistry Preparation and chemical behavior of the elements of the main groups.
Prerequisites
Basic knowledge of mathematics and physics: first and second degree equations, calculation of percentages, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (decimal, natural), concept of force, potential and kinetic energy.
Books
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 
Frequency
Attendance at lessons is not mandatory but strongly recommended
Exam mode
The exam consists of a separate written test and oral test. The written test, aimed at assessing the level of learning and understanding in the application of chemical models and methods in theoretical and stoichiometric exercises, consists of 6-8 numerical exercises and open questions on topics covered during the lessons, as per the course program. The written test (duration 2 hours), will contain a mandatory exercise for the purpose of passing the test itself (score higher than 1.5/4), regarding nomenclature and structural formulas. The oral test is aimed at assessing the degree of memorization of the contents and the logical skills acquired by the student as well as his ability to expose and explain the concepts independently. Those who pass the written test with a score of at least 18/30 are admitted to take the oral test. The oral test must be taken in the same session in which the written test was passed, under penalty of forfeiture of validity of the latter. Attending students can voluntarily access the exemption path (2 intermediate tests during the semester in place of the final written test; indicative dates: early December (1st exemption), mid-January (2nd exemption). The exemption test is passed with a score of at least 18/30 and the average of the two exemption tests constitutes the score of the written test. In the case of a written test passed through exemptions, its validity is extended to the entire winter session. For students with DSA, the exam procedures are agreed with the responsible offices.
Bibliography
N/A
Lesson mode
Lectures and numerical exercises in class.
GIUSEPPE POMARICO Lecturers' profile

Program - Frequency - Exams

Course program
Introduction to the course. Historical notes. Scientific method, properties of matter, measurement and units of measurement, significant figures. Fundamental principles of chemistry Elements, compounds and mixtures, states of aggregation of matter, Lavoisier's law, Proust's law, Dalton's atomic theory. Atoms, molecules, concept of mole and Avogadro's number. Molar and ponderal ratios. Percentage composition by weight, minimum formula and molecular formula. Simple chemical reactions and their balancing, limiting reagent, theoretical and percentage yield. Traditional and IUPAC nomenclature of the main inorganic compounds. Atomic structure Atomic structure, Thomson model, waves and electromagnetic spectrum, atomic spectra, Planck equation, photoelectric effect, Bohr atom and quantization of energy. Uncertainty principle, wave-particle duality, notes on wave mechanics, Schrodinger equation. Quantum numbers, wave functions and atomic orbitals. Construction (Aufbau) of the electronic structure of atoms, Pauli exclusion principle and Hund's rule. Periodic Properties Periodic table, periodic properties of the elements (atomic radius, ionic radius, ionization energy, electron affinity, metallic properties and magnetic properties. Chemical bond theory Lewis theory, octet rule and exceptions, formal charge and oxidation state, order, length and bond energy. Polar bond and electronegativity, covalent and ionic bond. Resonance. Valence bond theory (VB), hybrid orbitals and shape of molecules, VSEPR theory. Covalent bonds sigma (σ) and pi (π). Molecular orbital theory (MO). Thermodynamics Heat and work. Internal energy and 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 law of thermodynamics. Gibbs free energy. Gaseous state Pressure, ideal gas laws and equation of state of ideal gases, gas mixtures, Dalton's law, van der Waals equation for real gases. Solids and liquids Intermolecular forces and hydrogen bonding. Chemical-physical properties of liquids (boiling point, evaporation enthalpy, vapor pressure and its dependence on temperature: Clausius-Clapeyron equation). Specific heat. Heat and phase transitions; heating diagram of a pure species at constant pressure; state diagram of pure chemical species (H2O); structure and properties of water. Chemical equilibria in the gaseous phase Reaction quotient and equilibrium constant, Le Chatelier's principle; different types of equilibrium constant. Gibbs free energy and equilibrium constant, effect of the variation of P, V, n and T (van't Hoff equation). Heterogeneous equilibria. Solutions Units of measurement of concentration, conversion between the different units of measurement of concentration, solubility. Henry's law, thermodynamics of solutions, colligative properties (bullioscopic rise, freezing point lowering, lowering of vapor pressure, osmotic pressure). Solutions of strong and weak electrolytes, van't Hoff coefficient. Solution equilibria Acid-base equilibria (Arrhenius, Brønsted-Lowry and Lewis definitions). pH scale of aqueous solutions, strong acids and bases, correlations between structure and chemical behavior, polyprotic acids. Weak acids and bases, acidity constants, salt hydrolysis, buffers, titrations. Solubility of salts. Solubility products. Common ion effect, pH effect, complexation effect. Fractional precipitation. Electrochemistry Conversion of chemical energy into electrical energy. Standard reduction potentials and their use. Nernst equation and spontaneity of a redox reaction. Chemical and concentration cells. Examples of biological phenomena with electrochemical implications. Potentiometric measurement of pH. Balancing redox reactions. Notes on chemical kinetics. Reaction rate, reaction order, collision theory, Arrhenius law; activation energy; reaction mechanism; catalysis; enzymatic catalysis. Notes on systematic inorganic chemistry Preparation and chemical behavior of the elements of the main groups.
Prerequisites
Basic knowledge of mathematics and physics: first and second degree equations, calculation of percentages, arithmetic and weighted averages, scientific (exponential) notation of numbers, logarithms (decimal, natural), concept of force, potential and kinetic energy.
Books
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 
Frequency
Attendance at lessons is not mandatory but strongly recommended
Exam mode
The exam consists of a separate written test and oral test. The written test, aimed at assessing the level of learning and understanding in the application of chemical models and methods in theoretical and stoichiometric exercises, consists of 6-8 numerical exercises and open questions on topics covered during the lessons, as per the course program. The written test (duration 2 hours), will contain a mandatory exercise for the purpose of passing the test itself (score higher than 1.5/4), regarding nomenclature and structural formulas. The oral test is aimed at assessing the degree of memorization of the contents and the logical skills acquired by the student as well as his ability to expose and explain the concepts independently. Those who pass the written test with a score of at least 18/30 are admitted to take the oral test. The oral test must be taken in the same session in which the written test was passed, under penalty of forfeiture of validity of the latter. Attending students can voluntarily access the exemption path (2 intermediate tests during the semester in place of the final written test; indicative dates: early December (1st exemption), mid-January (2nd exemption). The exemption test is passed with a score of at least 18/30 and the average of the two exemption tests constitutes the score of the written test. In the case of a written test passed through exemptions, its validity is extended to the entire winter session. For students with DSA, the exam procedures are agreed with the responsible offices.
Bibliography
N/A
Lesson mode
Lectures and numerical exercises in class.
  • Lesson code1016546
  • Academic year2025/2026
  • Coursecorso|33586
  • CurriculumBiotecnologico
  • Year1st year
  • Semester1st semester
  • SSDCHIM/03
  • CFU6
  • Subject areaDiscipline chimiche