Unit CHEMISTRY
- Course
- Agricultural and environmental sciences
- Study-unit Code
- 80112009
- Curriculum
- In all curricula
- Teacher
- Ferdinando Costantino
- Teachers
-
- Ferdinando Costantino
- Hours
- 90 ore - Ferdinando Costantino
- CFU
- 9
- Course Regulation
- Coorte 2017
- Offered
- 2017/18
- 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
- Introduction to matter and energy.
the structure of 'atom and molecules.
Reactivity of the elements and the periodic table. The chemical nomenclature and chemical reactions.
the chemical bond.
The states of matter (liqiuido, solid, gaseous).
Solutions. The principles of the chemical equilibrium. The equilibrium solubility in aqueous solution and acid-base equilibria.
basic concepts on the structure and reactivity of molecules. Functional groups, preparation, structure and reactivity of akanes, alkenes, alcohols, carboxylic acids, peptides, ammines, carbohydrates, amminoacids and proteins. - Reference texts
- Bettelheim F.A. Brown W.H. et al. Chimica e Propedeutica Biochimica EDISES, II edizione
I. Bertini, C. Luchinat, F. Mani, Stechiometria, Casa Editrice Ambrosiana, Milano, 2009. - Educational objectives
- The course has the following aims: providing to the students an adequate basic knowledge of the following Chemistry contents:
-knowing how to balance chemical equations and to calculate the amount of reagents and products.
-knowing the properties and the reactivity of the elements of the periodic table.
-having a basic understanding of atomic structure and main characteristics of the chemical bond.
-basic stoichiometric calculations about the solutions.
-knowing the properties and the reactivity of the main organic molecules of biological interest. - Prerequisites
- The course of Chemistry is one of the fundamental subjects of the first year of studies. The prerequisites to satisfactorily address the study of Chemistry are:
-a proper knowledge of arithmetic, algebraic rules and logarithmic calculations.
-a proper knowledge of the use of physical and conversion factors.
-suitable analysis capacity to address the problems - Teaching methods
- 78 hours of frontal lectures on topics of basic and theoretical chemistry and numerical exercises on stoichiometry. 12 hours of small laboratory demonstrations to be carried out in the classroom. Students will be provided of didactic materials to integrate the study on basic textbooks, such as: slides of the lectures, numerical exercises, supplementary handouts.
- Learning verification modality
- The examination is divided into two partial tests. The first one is expected around the half of the first semester whereas the second one will be carried out at the end of the semester, tipically at the end of january, half february and end of february.
Only the students which reached a sufficient mark to first partial test can acces to the second one. The final mark will be calcualted as the average of the two tests. The students which do not have reached a suffcient mark at the first partial test can acces to the total one that contains all the arguments of the program.
The total test will be also held during the other dates listed above and in the other calendar dates. The Total examination will also include an oral interview, the date of which will be agreed with the students at the time of the written examination .
At the oral examination can acces only the students which have passed the written test with a score greater than or equal to 18 /30 be accessed after you have .
The written test will contain open-ended questions both theoretical and numerical . For the numerical exercises in the empty space below the text shall be provided the result (with its unit of measurement) and the holding in its essential steps . The only numerical result will not be sufficient to validate the exercise . For theoretical questions will be given an answer brief and comprehensive . If the empty space below the question was not enough, you can continue in the back of one sheet . At the end of a predetermined time must be given to only the first sheet of text so compiled . - 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.
Energy, heat and work. The egocentric Convention. Isolated, closed and open systems. The electromagnetic spectrum. Internal energy. The first law of thermodynamics. Enthalpy and standard enthalpy. Heat and enthalpy. Exothermic and endothermic processes. State functions. Energetics of physical changes of water.
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 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.
Lewis acid-base theory. Definition of Lewis acids and bases. Comparison with Bronsted theory. Typical Lewis acids and bases. Outline of complexation reactions and the formation of coordination compounds according to the acid-base Lewis theory.
Electrochemistry: galvanic cell. Anodic, cathodic and cell reactions . Electromotive force and potential electrodes. Standard hydrogen electrode. The standard reduction potentials and their use. Effect of concentration on electrode potential. Nernst equation. Concentration cells.
Organic Chemistry
Atomic structure. The concept of hybridization of orbitals. The chemical bond. Acidity according to Lewis and Bronsted. Functional group concept. Alkanes. Isomery of
organic molecules. Atomic conformation. Basic principles on the reactivity. Introduction to organic reactivity. Alkenes: structure and reactivity. The reaction of electrophilic addition. Aromatic compounds: aromaticity concept, structure and reactivity of aromatic compounds. The electrophilic aromatic substitution reaction
Elements of organic stereochemistry: stereoisomery, enantiomers and diastereoisomers.
functional groups: amines, alcohols, aldehydes and ketones, carboxylic acids.
Derivatives of carboxylic acids: acid halides, anhydrides, esters and amides. Amino acids, peptides and proteins.