00088 - Chemistry (M-Z)

Learning outcomes

At the end of the course the student acquires basic knowledge of chemistry and is able to apply this knowledge to the solution of numerical problems. Furthermore the student learns that chemistry is a useful and innovative interdisciplinary science.

Course contents

• Fundamental principles of chemistry: the scientific method, properties of matter, measures and units of measures, significant figures. Elements, compounds and mixtures, aggregation states of matter, Lavoisier’s law, Proust’s law, Dalton’s atomic theory. Atoms and atomic mass. The concept of mole, Avogadro’s number.
• Atomic nature ofmatter, elemental particles, mass and charge of the elemental particles, atomic number, mass number, isotopes. Minimum, molecular and structural formula, atomic weight, molecular weight and stoichiometric calculations.
• Recall of atomic theory (Schrodinger, Heisenberg). Quantum numbers. Shape and orientation of orbitals. Concept of probability density for the electron. Multi-electron atoms. Aufbau principle. Hund’s rule and Pauli exclusion principle. Construction of the electron configuration of an atom.
• The Periodic Table. Construction of the periodic table starting with the filling of the quantum levels. Periodic trend of atomic properties. Ionization energy, electron affinity and elettronegativity.
• The chemical bond. Nomenclature and geometry of some inorganic and organic molecules (the list of molecules is present in the teaching materials). The covalent bond according to Lewis theory and description of the molecules with the formal charge. The valence shell electron pair repulsion rule (VSEPR).
• Bond and lone pairs. Covalent and polarized covalent bonds. Valence bond theory. Hybrid orbitals. Construction of sp, sp2 and sp3 hybrids. Examples: molecules of BeH2, BH3, CH4. The molecules of water and ammonia. . Influence of the lone pairs. Multiple bonds. Effect of resonance: the benzene molecule. Molecular orbital theory. Biatomic homonuclear molecules. Bonding and antibonding molecular orbitals. Filling of molecular orbitals. Single, double and triple bonds. Paramagnetism in the O2 molecule. Heteronuclear biatomic molecules. Bond polarization. . Differences in electronegativity. The ionic model. Dipole moments in biatomic and polyatomic molecules. Reference to HOMO, LUMO and delocalized orbitals.
• Ionic bond and lattice energy.
• Metal bonds and band theory. Conductors, insulators and semiconductors. Diodes and photovoltaic cells.
• Bonds and intermolecular interactions (hydrogen bond, Van der Waals , dipole-dipole, ione-dipole interactions…).
• State of matter: solid, liquid, gaseous. Amorphous and crystalline solids. Packing o crystalline solids: metals and ionic solids. Metals with body-centered, face-centered, etc. packings, interstitial sites, defects.
• Ionic solids and determination of their crystalline lattice based on their ionic radii. Reference to molecular solids and crystallography, crystal lattices, diffraction and Bragg’s law.
• Chemical and Physical transformations.
• Chemical reactions. Treatment and balancing of chemical equations. Stoichiometric calculations. Reactions with the limiting reagent. Types of reactions: acid-base, redox, precipitation, exchange. Concentration measure unit. Molarity, molality. Percentage composition by weight, molar fraction.
• Thermochemistry. First law of thermodynamics. Heats of reactions. Enthalpy. Hess law and heats of formation. Standard conditions. Statistical derivation (qualitative) of enthalpy. Second law of thermodynamics. Absolute entropy and Third law of thermodynamics. Gibbs free energy. Calculations of H, S and G. Enthalpy and entropic balance. Enthalpically and entropically favored reactions. Effect of temperature on the calculation of G.
• Chemical equilibrium. Chemical potential and determination of the chemical equilibrium, equilibrium constant and reaction quotient. Le Chatelier principle. Dependence of the equilibrium constant on temperature, effect of the pressure. Effect of the change in concentration. Control of chemical equilibrium. Homogeneous and heterogeneous phase equilibria. Solubility product, sparingly soluble salts and constant of the solubility product (Kps) Reference to phase equilibria. PT state diagram. Vapour pressure. Colligative properties: vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.Electrochemistry. Galvanic cells. Standard reduction potentials. Gibbs free energy and useful work. Nernst equation. Batteries and accumulators. Electrolitic cells, electrolysis of molten salts and electrolysis of solutions. Corrosion.

(*) Please note: the program can undergo changes along the way.

Readings/Bibliography

Authors: Balzani Moggi Prodi Venturi
Title: Chimica: Fondamenti e Prospettive
Editor: Bononia University Press

Authors: S. S. Zumdhal
Title: Chimica (Italian Version, the original english version is entitled "Chemistry")

Editor: Zanichelli

Teaching methods

Lessons and Exercises at the blackboard

Assessment methods

Examination is written. If passed, the student may ask for an additional oral examination.

It is allowed to bring:

• a calculator;
• a periodic table;
• one A4 sheet of paper with notes and formulas.

Students with a DSA certification can contact the teacher (via email, at least 7 days before the exam) with the following office in cc:

https://site.unibo.it/studenti-con-disabilita-e-dsa/it/contatti

for additional time for the examination.

Teaching tools

The teaching material will be available to students on Virtuale.

Office hours

See the website of Damiano Genovese