38837 - General Chemistry

Learning outcomes

At the end of the course, the student knows the theoretical foundations necessary for understanding the structure of matter, in its atomic and molecular constituents. He has a method suitable for the scientific treatment of the phenomena underlying the transformation of matter with models suitable for the rigorous description of the thermodynamic and kinetic aspects.

Course contents

Theoretical teaching unit (30 hours)

Part 1 of 11 - From Matter to the Law of Gases - Establish the characteristic physical quantities of a measurement; Apply the units of measurement of the International System and the related prefixes; Evaluate the precision and accuracy of a measurement; Identify which properties of a sample depend on the size of the sample and which are independent of it; Distinguish between mass and weight; Connect accuracy and precision of a measurement with systematic and accidental errors; Classify matter according to one's physical state; Classify physical states on the basis of the attractive forces that characterize them; Classify physical states from a microscopic point of view; Know the postulates of corpuscular theory; Relate the concentration of a solution to its density; Classify a system as homogeneous or heterogeneous; Identify the most suitable techniques for the separation of mixtures on the basis of the characteristics of the mixture itself; Classify a material as a pure substance or mixture; Understanding what happens when a body is heated; Interpret, according to the kinetic theory, the stops in the thermal analysis curves; Compare different substances on the basis of the temperatures of the changes of state and the values of latent heat; Indicate the experimental evidence underlying Boyle's law; Indicate the experimental evidence underlying Charles's law; Indicate the experimental evidence underlying the Gay-Lussac law; Recognize that the ideal gas is a model; Predict the behavior of a fixed quantity of gas when p, V or T varies; Recognize the behavior of aeriforms as a tool for the determination of molecular formulas and atomic masses.

Part 2 of 11 - Atoms and Molecules - Distinguishing physical transformations from chemical transformations; Distinguish an element from a compound; Indicate the experimental evidence that led Lavoisier to formulate the law of conservation of mass; Indicate the experimental evidence that led Proust to formulate the law of definite proportions; Indicate the experimental evidence that led Dalton to formulate the law of multiple proportions; Knowing how to "read" a formula; Describe the composition of a substance; Connect mass, chemical quantity and number of atoms of a sample; Understanding the relationship between mass percentage composition and atomic composition of a compound; Determine the molar mass of a substance known by the formula; Use the concept of mole to convert the mass / volume of a substance or the number of elementary particles into moles and vice versa; Determine the empirical and molecular formula of a compound; Understanding the difference between real and ideal gases and applying van der Waals law.

Part 3 of 11 - From Atoms to Bonds - Understanding how experimental evidence led to the transition from Thomson's atomic model to Rutherford's; Explain how the composition of the nucleus determines the chemical identity of the atom; Starting from the atomic structure, understand how bonds are formed; Recognize that Bohr's atomic model has as its experimental basis the spectroscopic analysis of the radiation emitted by atoms; Understanding how de Broglie's theory and the uncertainty principle are the basis of a probabilistic conception of matter; Distinguish between wave and corpuscular behavior of electromagnetic radiation; Understanding the meaning of a standing wave and the importance of the wave function; Be aware of the existence of energy levels and sub-levels and their arrangement in order of increasing energy towards the outside; Use the specific symbology and orbital filling rules for writing the electronic configurations of all atoms.

Part 4 of 11 - The Periodic Table - Describe the main properties of metals, semi-metals and non-metals; Identify the position of the various families of elements in the periodic table; Explain the relationship between Z, electronic structure and the position of the elements on the periodic table; Understanding that the law of periodicity has been a tool for both classification and prediction of elements; Discuss the historical development of the concept of periodicity; Explain the trends in the periodic properties of elements in groups and periods.

Part 5 of 11 - The Chemical Bonds - Distinguish and compare the different chemical bonds (ionic, covalent, metallic), Establish on the basis of the external electronic configuration the number and type of bonds that an atom can form, Define the nature of a bond on the basis of the electronegativity difference Describe the observable properties of the materials, on the basis of their microscopic structure. Predict, based on the position in the periodic table, the type of bond that can form between two atoms. Predict, based on the VSEPR theory, the geometry of simple molecules; Identify if a molecule is polar or apolar, after having determined its geometry based on the VSEPR model. Recognize the various types of hybridization.

Part 6 of 11 - Chemical solutions and reactions - Interpret the dissolution processes on the basis of the intermolecular forces that can be established between the solute and solvent particles; Organize data and apply the concept of concentration and colligative properties; Read solubility diagrams (solubility / temperature; solubility / pressure); Know the various ways of expressing the concentrations of solutions; Understanding the colligative properties of solutions; Interpret a chemical equation in terms of quantity of substance; Interpret a chemical equation in terms of quantity of substance; Relate theoretical data and experimental data; Know the various types of chemical reactions; Identify the double exchange reactions in which a precipitate is formed; Recognize a neutralization reaction. Classify the main categories of inorganic compounds into binary / ternary, ionic / molecular; Group the oxides according to their chemical behavior; Group the hydrides according to their chemical behavior; Apply the rules of the IUPAC and traditional nomenclature to name simple compounds and vice versa; Write the formulas of simple compounds; Write the ternary salt formula.

Part 7 of 11 - Chemical equilibria - Predict the evolution of a system, note the Keq values and the exothermic or endothermic character of a reaction; Acquire the conceptual meaning of the Le Châtelier principle; Understanding the historical and conceptual evolution of acid-base theories; Classify compounds as strong, weak, or non-electrolytes; Establish the strength of an acid / base, known the value of Ka / Kb.

Part 8 of 11 - The equilibria in aqueous solution - Identify the pH of a solution; Relate the strength of an acid / base to Ka; Determine the pH for weak acids / bases; Acid-base titration; choose the appropriate relationship to determine the pH; understand the mechanisms of saline hydrolysis; Identify the acid, basic or neutral character of a solution based on the color of the indicator paper; Calculate the pH of strong and weak acid / base solutions; Explain which pairs can form a buffer solution; classify substances according to their solubility using the Kps.

Part 9 of 11 - Thermodynamics - Describe how the potential energy and kinetic energy vary during a transformation; Understanding the meaning of the enthalpy change during a transformation; Relate the spontaneity of a reaction with the variation of enthalpy and entropy; explains how the chemical energy of a system varies during an endothermic / exothermic transformation; Relate the sign of the enthalpy variation with the heat exchanged with the environment; Predict the spontaneity of a reaction, through the variation of free energy of the system; Know the different calorific value of fuels; Distinguish spontaneous transformations; Compare fossil fuels and biofuels; Identify the enthalpy and entropic variation in the spontaneous melting of ice.

Part 10 of 11 - Kinetics - Describe the factors that influence the speed of a reaction; Distinguish between reaction energy and activation energy; Explain how a catalyst works; Explain the reaction kinetics in the light of the collision theory; Arrhenius law.

Part 11 of 11 - Redox reactions - Recognize, in a redox reaction, the agent that oxidizes and that which is reduced; Write the balanced redox equations in both molecular and ionic form; Understanding that spontaneous redox reactions can generate a flow of electrons; Connect the position of a chemical species in the table of standard potentials to its reducing capacity; Establish comparisons between galvanic cells and electrolytic cells; Understanding the importance of redox reactions in the production of electricity

Exercises on problems with numerical calculations (10 hours)

Chemical problems involving numerical calculations will be illustrated and solved in class: balancing of reactions and stoichiometric problems, preparation of solutions, pH calculations, solubility problems. These exercises aim to consolidate the knowledge acquired in the other teaching units and to make the student able to use them in practical application contexts. Laboratory teaching unit (20 hours) The student will be given explanations in the classroom on how it is useful to set up an experiment in the laboratory to highlight particular properties of a substance or sample of interest. Practical exercises will then be carried out in the laboratory with a dual objective: To make the student familiar with the equipment and with the safety regulations of the chemical laboratory; develop in the student the ability to connect the results of an empirical experience to the notions acquired during the theoretical lectures. The aspects on which attention will be focused will be acid base properties, redox properties and solubility problems.

Contents of the laboratory teaching unit (20 hours)

The student will be given explanations in the classroom on how it is useful to set up an experiment in the laboratory to highlight particular properties of a substance or sample of interest. Practical exercises will then be carried out in the laboratory with a dual objective: To make the student familiar with the equipment and with the safety regulations of the chemical laboratory; develop in the student the ability to connect the results of an empirical experience to the notions acquired during the theoretical lectures. The aspects on which attention will be focused will be acid base properties, redox properties and solubility problems.

Readings/Bibliography

The choice of the text on which to prepare the exam is free, provided that the text is of university level.

The texts that I prefer today are those that, contrary to the classical approach, dedicate the first chapter to the description of the atom. Among them, I found

Silvestroni P., Fundamentals of chemistry, CEA

Failla S., Paolesse R. and others, General and Inorganic Chemistry, edi-ermes

Other excellent books are

Burdge J. & Overby J. General Chemistry, EDRA

Atkins P., Jones L., Laverman L. Principles of Chemistry, Zanichelli

Tro, N.J. Chemistry, a molecular approach, EdiSES

Brady, J. E. Holum, R. “Fundamentals of chemistry”, Zanichelli

J. Kotz, et al, “Chemistry”, EdiSES.

R. Chang, “Fundamentals of General Chemistry”, McGraw-Hill

Silberberg M.S, Chemistry, McGraw-Hill

Brown et al., Fundamentals of Chemistry, EDISES.

Palmisano L. et al, Elements of chemistry, EdiSES

Bertani et al, Fundamentals of Chemistry for Technologies, EdiSES

For supplementary exercises: Bertini I et al, Stoichiometry, CEA (Zanichelli)

Teaching methods

The course is constituted by lectures, assisted by the projection of slides and parts of textbooks, animations and films.

There are also classroom exercises with the resolution of numerical problems on the blackboard and practical demonstrations.

Finally, laboratory activities are planned, with experts. The attendance of laboratory activities requires all students of modules 1 and 2 to carry out e-learning [https://www.unibo.it/it/servizi-e-opportunita/salute-e-assistenza/salute- e-safety / safety-and-health-in-study-places-and-internship] and participation in module 3 of specific training on safety and health in study places. Information on dates and methods of attendance of module 3 can be consulted in the specific section of the degree program website.

Assessment methods

The exam consists of a written and an oral test. Each written test is offered on two very close dates, so that the student can choose the one that best suits his or her commitments. Each new test cancels all the previous ones. The result of the last written test remains valid for one year. Only students who have obtained a grade higher than or equal to 18/30 in the written test can access the oral exam.

The end-of-course exam aims to evaluate the achievement of the educational objectives:

• Transfer a chemical problem to a stoichiometric calculation setup. In this setting, both numbers and units of measurement must be considered correctly. Chemistry describes atoms and the interactions between them through conventions, such as strokes that indicate chemical bonds, dots that indicate electrons, subtraction and addition symbols that identify negative and positive charges. These conventions are the equivalent of grammar and spelling for the writer. Errors on these aspects can only be considered with equal severity.

• Describe the composition of matter and the phenomena that lead to its transformations with properties of language and scientific rigor.

Teaching tools

Teaching materials projected in class other than textbook pages will be made available to the student in electronic format on “Virtuale”.

Office hours

See the website of Luca Laghi

See the website of Gianfranco Picone