GENERAL CHEMISTRY AND PHYSICAL CHEMISTRY
Learning outcomes of the course unit
the course objectives include giving the student those instrutments necessary for understanding the basic features and behaviour of matter and teaching the student the basic principles of physics and chemistry and the solving of stechiometric problems
to establish and to develop the principles for the explanation and the interpretation of chemical reactions, by means of models, peculiarity of the Physical Chemistry;
to provide knowledge of thermodynamics and kinetics, especially with regard to the study and the interpretation of biochemical and biological processes;
to show that thermodynamics is the foundation for understanding the main machinery related to food preservation;
to give an outline of colloidal chemistry, foundation of structural and functional features of foods.
Course contents summary
1) definition of system and environment, isolated, closed and open systems; 2) definition of state functions; 3) energy: kinetic and potential energy; heat and work; enthalpy; free energy and enthropy. 4) the atom: protons, neutrons, electrons; eletronic structure: quantum theory; 5) the periodic table; 6) chemical bonds: ionic bond, covalent bond (Lewis’s theory, Valence bond theory and molecular orbital theory): molecular form: hydrogen bond: van der waals and London’s dispersion interactions; 7) chemical reactions: stechiometric caclulations; 8) the gas state: the ideal gas and real gases; 9) the liquid state: evaporation and vapor pressure; boiling temperature and its variation with pressure; 10) solid state: sublimation and vapore pressure: melting temperaO2; 12) solution state: solutes and solvents: enthralpic and enthropic balances during the solution process: solubility, the ideal solution and real solutions: Rault’s law and Henry’s law: state diagramms of acqueous solutions with non volatile solutes; fractionated distillation; 13) chemical equilibrium: homogeneous and hetereogeneous equilibria, definition of activity: Kc, Kp, reaction quotient Q; principle of mobile equilibrium; 14) acids and bases according to Arrhenius, Bronstead and Lewis; forces of acids and bases; solution equilibrium: Ka, Kb, pH; hydrolysis; buffer solutions, acid/base titration; indicators; 15) insoluble and solubility equilibria: Kps; 16) thermodynamics: first principle: dU and dH: calorimetry; bond, formation and combustion enthalpies, Hess’s law; second principle: definition of enthropy and its physical meaning; relationship bewteen G,H and S; definition of ÄGr and Q: extent of reactions: “ÄH driven ” “ÄS driven” reactions; 17) electrochemistry: the electrochemical cells; the Nernst equation; relationship between ÄGr and the potential difference between the half cells: concentration cells; the pH-meter; electrolysis; 18) chemical kinetics: kinetic equations; reaction order and its determination; zero, I° and II° reaction orders; half-life; elementary reactions and reaction meccanisms; collision theory; 20) inorganic chemistry: hydrogen, nitrogen oxygen sulphur, phosphoros, halogens and principle transtion elements.
1. Equilibrium thermodynamics applied to chemical and biological systems with a statistical thermodynamics outline. Variables and state functions. The laws of thermodynamics. The temperature and pressure dependence of thermodynamic quantities. Thermochemistry. Calorimetry. Outline of statistical Thermodynamics. Exercises.
2. Changes of state: physical transformations of pure substances. Phase diagrams. Clapeyron and Clausius-Clapeyron equations. Vapour-liquid phase transition and critical phenomena. The principle of corresponding states. Gibbs phase rule
3. Changes of state: physical transformations of simple mixtures. Open systems and partial molar quantities. Ideal and real solutions. Raoult and Henry laws. Fugacity and activity. The water activity in foods. Regular solutions. Ideal mixing and excess functions. Phase equilibria in binary systems. Fractional distillation. Azeotropes, eutectic, partially miscible liquids, binary mixtures compounds forming. Solvent chemical potential. Colligative properties. Osmotic pressure. Molecular weight measurements. Membrane equilibria. Solutions of macromolecules. Dialysis equilibrium. Donnan equilibrium. Transport phenomena: passive and active transport.
4. Equilibria of chemical reactions. Thermodynamics of chemical equilibrium. Gibbs free energy and equilibrium constant. Activity and ionic strength. Standard state. Distribution diagrams. Binding curves. Cooperativity.
5. Electrochemistry. Electrochemical cells. Electrodes. Nernst equation. Standard reduction potentials. The potentiometer.
P.W. Atkins, L. Jones: Chimica Generale, Zanichelli (1998);
S. S. Zumdahl: Chimica, Zanichelli (1993)
P. Michelin Lausarot, G. A. Vaglio: Stechiometria per la Chimica Generale, Piccin (2005)
Gianluigi Ingletto: Esercizi di Chimica, UNI.NOVA (2006)
A. Schiraldi, , Elementi di Chimica Fisica, Cisalpino, Milano (1993)
- A. Immirzi, Chimica Fisica (Termodinamica), CUES, Salerno (2002)
- Laidler, Meiser, Chimica Fisica, Editoriale Grasso, Bologna, 1999