Università degli Studi di Perugia

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Unit GENERAL AND INORGANIC CHEMISTRY

Course
Pharmacy
Study-unit Code
50170105
Location
PERUGIA
Curriculum
In all curricula
Teacher
Riccardo Vivani
Teachers
  • Riccardo Vivani - Didattica Ufficiale
Hours
  • 80 ore - Didattica Ufficiale - Riccardo Vivani
CFU
10
Course Regulation
Coorte 2018
Offered
2018/19
Learning activities
Base
Area
Discipline chimiche
Sector
CHIM/03
Type of study-unit
Obbligatorio (Required)
Type of learning activities
Attività formativa monodisciplinare
Language of instruction
Italian
Contents
Periodic table of elements, Electronic structure and reactivity of elements. Compounds and stoichiometric relationships. Bases of chemical reactivity. Principia of chemical equilibrium
Reference texts
M. Schiavello, L. Palmisano, Fondamenti di Chimica, EdiSES, Third (or Fourth) edition, Napoli 2010 (2013).
Additional didactic material (slides, exercises and examples, and other material) availableon-line at the web site http://www.unistudium.unipg.it.
Educational objectives
Teaching is the first rigorous approach to general and inorganic chemistry. The main objective of the course is to provide students with the basic concepts of general chemistry as a description of nature, an appropriate scientific language and the ability to study in a critical and rational way.
The main knowledge gained will be:
- Atomic theory and electronic structure of atoms.
- Chemical bond and molecular geometries.
- Intermolecular forces.
- Chemical reactions.
- Chemical equilibrium in the gas phase and in aqueous solution.
The main skills (ability to apply knowledge acquired):
- Identify and be able to write formulas of inorganic compounds;
- Represent molecules or molecular ions inorganic highlighting the orientation of the atoms and the links between them;
- Predict the polarity and the physical state of molecules;
- Predict the reactivity of inorganic compounds both in redox reactions and that of non-redox;
- Write and describe the qualitative and quantitative aspects (stoichiometric) of a chemical reaction in relation to chemical equilibrium homogeneous and heterogeneous
Prerequisites
In order to understand and achieve the intended learning is necessary that the student possesses skills in mathematics and physics. In particular, the student should know and be able to use some basic mathematical tools (equivalence, equation of first and second degree, logarithmic, exponential, inequalities, derivatives, integrals) and concepts of basic physics (units, force, energy).
Teaching methods
The course is organized as follows:
- Lectures on all the topics of the course. The lessons will be conducted with the help of the blackboard and through slide shows.
- Numerical exercises in the classroom for the solution of numerical exercises guided with the aid of the blackboard.
The teaching material (slides, exercises during numerical exercises, texts of previous written tests, other material) are made available to students on the e-studium platform after registration.
Other information
At the exam the students are advised to bring student card and proof of identity.
- Essentials allowed for the written exam: periodic table, traditional calculator (scientific), pen.
- Not Allowed: mini computers or PDAs, cell phones, notebooks, books, notes or other information.
- The test must be written in non-erasable pen.
- The student who had not passed a test, can access any of the subsequent tests.
Learning verification modality
The exam consists of a written test.
The duration of the written test is three hours. It consists of 10 to 15 open-ended questions both theoretical and numerical, in which case it must contain the procedure followed to obtain the result. The theoretical questions require a clear answer concise and comprehensive, according to the knowledge transmitted during the lectures.
The exam is designed to ascertain the knowledge, understanding, language acquisition own discipline and the ability to display and summary about the theoretical aspects, and ensure the ability to apply the acquired skills to solve numerical problems related to practical cases.
Extended program
Introduction to matter and energy. Homogeneous and etherogeneous systems. Solutions, simple and compound systems. Constitution of the atom, atomic number, mass number, nuclea, isotopes, elemnts. Atomic mass. Scale of atomic masses. Isotopic abundancy. Mass defect. Average atomic mass. Avogadro constant and the mol. Molar scale of atomic masses. Minimal and molecular chemical formula. Combination ratios. Weigth percent composition. Outline on elemental analysis. Chemical reactions. Principle of atom conservation. Complete and with limiting reactant reactions. Solutions. Concentration. Measure unit:% by weight,% by volume, mole fraction, molarity, molality. Density and dilution procedures.
Energy, heat and work. The egocentric Convention. Isolated, closed and open systems. Internal energy. The first law of thermodynamics. Enthalpy and standard formation enthalpy. Heat and enthalpy. Exothermic and endothermic processes. State functions. Energetics of physical changes of water. Reactions of formation standard of compounds. Standard Enthalpy of formation. Thermodynamic cycles. Hess Law.
Ondulatory phenomena. Interference of waves. The electromagnetic spectrum.
Fundamentals on the atomic theory: The atomic spectra. The hydrogen atom according to Bohr theory. Fundamentals of Quantum Mechanics. Wave nature of the electron. The wavefunction. The quantum numbers and spin. Orbital and energy levels. The principle of Aufbau, the Hund rule, the Pauli exclusion principle. Electronic structure of the elements. Electronic configurations. The periodic table. Periodic properties. Effect of shielding and effective nuclear charge. Atomic and ionic radius and volume. Ionization energy. Electron affinity. Valence, electronegativity and oxidation number. Simple methods to determine the number of oxidation.
Basic reactivity of elements. Metallic, semimetallic and metallic character. Hydrides and oxides. Basic oxides and hydroxides. Acid oxides (anhydrides), and oxoacids (oxoanions). Nomenclature. Formation of salts. Classification of chemical reactions. Redox and non-redox reactions. Formation , decomposition, combustion, displacement, and exchang ereactions. Balancing of redox reactions with the ionic-electronic method in acid and basic aqueous solutions. Disproportion.
Molecular structure and chemical bonds: Notes on ionic bonding. Description of the covalent bond with the valence bond method. Sigma and pi bonds. Octet rule. Simple, double and triple bond. Dative bond. Electrondeficient molecules. Expansion of valence sphere and the octet rule violation. V.S.E.P.R. method and molecular geometry. Hybridization. Structural formula of common molecules and molecular ions. Resonance. Polarity of bonds, polarity of molecules.
Determination of molecular geometry of organic molecules. Hybridization of atomic orbitals. sp, sp2 and sp3 hybrids. Determination of oxidation numbers with the bond to bond method. Relationship with average oxidation numbers. Pi bonds and free or prevented rotation around a bond axis. Cis and trans geometrical isomers. Molecular resonance. Resonance limit formulas. Case of benzene, carbonate and sulfate. Delocalization of electric charge and pi bonds. Resonance energy.
Polarity of bonds. Dipole moment. Homopolar and polar covalent bond. Molecular dipole moment as a vector sum of dipole moments of individual bonds. Study of the polarity of molecules. Polarity and molecular symmetry.
Intermolecular interactions: ion-ion, ion-dipole (dissociation and solvation of ionic solids), permanent dipole-permanent dipole, permanent dipole-induced dipole, instantaneous dipole-induced dipole (van der Waals forces and London dispersion forces). Concept of polarizability. Examples of typical systems. Hydrogen bond. Examples of intermolecular / intramolecular hydrogen bonds. The hydrogen bond in water and its chemical and physical effects. The hydrogen bond in proteins and in nucleic acids. Energies of intermolecular interactions.
The solid state: General types of solids, classified by the nature of chemical bonding: Solid metal (basic properties, model of the sea of electrons), ionic (basic properties, dissociation, solvation), covalent (basic properties, dimensionality in covalent solids depending on the direction of propagation in space of covalent bonds), molecular (basic properties, intermolecular interactions).
Gases: the nature and definition of pressure. Units of measurement. Atmospheric pressure. The model of perfect gases. Energy of ideal gas. Maxwell-Boltzmann distribution law of molecular energies. Temperature dependence of the distribution curves of molecular energies according to the Maxwell-Boltzmann Law. Ideal Gas Laws. State equation of perfect gases. Mixtures of perfect gases. Partial pressure. Dalton's law. Mole fraction. Partial volume.
Liquids. Vapor pressure of a liquid. Liquid-vapor equilibrium. Definition of equilibrium state. Dependence of vapor pressure on temperature. Clausius-Clapeyron. Solid-vapor equilibrium. Phase equilibria for systems with a component. Phase diagrams. Triple point. Normal melting and boiling temperatures. The phase diagram of water and carbon dioxide. Concept of variance. Perturbations of equilibria. Le Chatelier's principle of mobile equilibrium. Applications to phase equilibria.
Solutions. Concentration. Measure unit:% by weight,% by volume, mole fraction, molarity, molality. Ideal solutions. Definition of solution. Mixing Enthalpy. Deviations from the ideal behaviour. Dissociation of solutes. Types of solutes: strong electrolytes, weak electrolytes, non-electrolytes. Degree of dissociation. Van t'Hoff Binomial.
Colligative properties: vapor pressure of solutions Raoult's law , both volatile two components and the case of two components, one of which is non-volatile. Lowering the vapor pressure of the solvent. Cryoscopic lowering of melting temperature and ebullioscopic increase of boiling temperature of the solvent. Osmotic pressure. Operational definition. Semipermeable membranes. Isotonicity. Reverse osmosis.
Chemical Kinetics: The rate of reaction. Elementary reactions. Reaction mechanism. Dependence of the rate of reaction by the concentration. Differential and integrated rate laws of I and II order. Half-life. Dependence of the reaction rate on temperature. Arrhenius equation. Activated complex and transition state. Catalysts. Enzymes. Michaelis-Menten equation.
Chemical equilibrium. Characteristics of chemical equilibrium. Equilibrium constant and its properties. Prediction of reactivity based on the principle of Le Chatelier mobile equilibrium. Effects of perturbation of equilibrium: variation of concentration, pressure, volume and temperature. Equilibrium constant and reaction quotient. Reversible reactions and spontaneous reactions. Entropy. Definition. Second law of thermodynamics. Entropy change for the system and the environment. Criterion of spontaneity and reversibility based on the entropy change. Concept of entropy and order-disorder. Microscopic interpretation of entropy. Boltzmann equation. Concept of microstate. Qualitative assessment of the change in entropy for some chemical reactions. Third law of thermodynamics and absolute scale of entropies of substances. Free energy G. Definition. Criterion of spontaneity and reversibility based on free energy change at constant pressure and temperature. Standard free energies of formation and standard enthalpies of formation. Thermodynamic tables and their use. Dependence of G on pressure and concentration of the components of a chemical reaction. Relationship between standard free energy change and equilibrium constant.
Dependence of equilibrium constant on temperature. Van t'Hoff Equation.
Equilibrium solubility in aqueous solution. Concept of solubility. Solubility product. Calculation of concentrations of ionic balance. Effect of salt stoichiometry. Effect of an ion in common.
Acid-base equilibria. Definition of acid and base according to Bronsted-Lowry. Acid-base reactions. Ampholytes. Autoprotolysis reactions. Water autoprotolysis Equilibrium. Kw. Strong and weak acids and bases. Relative strength of acids and bases. Dissociation constants Ka and Kb and Kw and their relationship. Polyprotic acids. pH and the pH scale. Calculations of equilibrium concentration of typical aqueous acid-base systems: strong and weak acid solutions and bases. Neutralization reaction. Determination of equilibrium constant of the reaction of neutralization. Buffer solutions and mechanism of the buffering. Calculations for the determination of the equilibrium concentration in the buffers. Henderson-Hasselbach equation.
Electrochemistry. Galvanic cells. Anodic reactions, and cathode of the cell. Electromotive force and potential electrode. The standard hydrogen electrode. The standard potentials and their use. Effect of concentration on the electrode potential. Nernst equation. Concentration cells. Electrodes of I, II and III species.
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