General and Inorganic chemistry 2

Course structure includes frontal, collaborative and/or cooperative lessons.  Group works and individual research projects will be also encouraged. 

In particular, the course includes 42 hours of lectures and 36 hours of classroom exercises, during which exercises that are useful for solving the written exam will be conducted. Students will also be taken to the laboratory once to perform an experiment that aids in the understanding of some fundamental course concepts. An interim test is scheduled to be conducted before the Christmas break.

Should the circumstances require online or blended teaching, appropriate modifications to what is hereby stated may be introduced, in order to achieve the main objectives of the course.

Basic mathematics, basic physics, basic chemistry.  To attend the chemistry course in-depth preliminary knowledge is not needed.
However, it will be necessary to follow the lessons regarding the various topics consistently, avoiding being left behind in the study. Chemistry is a discipline that is built gradually and it will not be possible to understand it if it is not done a regular and constant study during the course. 

It will be very difficult to prepare it by yourself because many topics need time to be well received and matured. Students are strongly advised to take the zero courses in mathematics, physics, and chemistry that will be given before and/or during the courses.

1. WHAT CHEMISTRY IS: scientific method, matter, states of aggregation, substances, and compounds.

2 - STOICHIOMETRY - The concept of mole - stoichiometry law- Determination of the formula of a compound - The chemical equation and its equilibrium- Identification of redox reactions - Balancing of redox reactions - Stoichiometry: Quantitative relationships in chemical reactions - limiting reagent, numerical applications

3 - STRUCTURE OF ATOM - Subatomic particles: Electron, proton, neutron - atomic number, mass number - isotopes - atomic mass unit - Atomic model of Bohr / Rutherford - Quantum Mechanical Description of the atom - Atomic orbitals - Quantum Numbers - the Pauli exclusion principle - Principle of maximum multiplicity

4 - PERIODIC SYSTEM OF ELEMENTS - Periodic classification and electronic configuration of the elements- Periodic properties: atomic and ionic radii, ionization energy, electron affinity, electronegativity.

5 - CHEMICAL BOND - Ionic bond - covalent bond - valence bond theory - Electronegativity of atoms and polarity of bonds - Oxidation number - Dative bond - VSEPR Theory: molecular geometry hybrid -Orbitals - Resonance - Chemical bonding and structural formulas of the most common inorganic compounds.

6 - Intermolecular forces - Van der Waals and London Forces - hydrogen bond.

7 - GASEOUS STATE - General characteristics of gaseous state - ideal or perfect Gas - ideal gas laws – Gas Equation -  Gas Diffusion - Real Gas.  Numerical applications.

8 - CONDENSED PHASES AND STATE CHANGES - Outline of solid state properties as a function of the chemical bond - the liquid state Features - State changes - Vapor Pressure - State diagram of the water and carbon dioxide- Le Chatelier’s principle -.

9 - WATER SOLUTIONS - Types of solutions - Concentration units - Solubility (with particular reference to the solubility of ionic compounds) - Henry's Law - colligative properties of solutions: Lowering of vapor pressure and Raoult's Law - Cryoscopy and ebullioscopy - Osmosis and osmotic pressure - electrolyte solutions. Colligative properties of electrolytes - Dissociation Grade, Numerical applications

10 - KINETICS - Factors affecting the rate of reaction - kinetic equation and reaction order - Graphical Treatment of 1st order reaction - elementary reactions: the reaction rate-limiting step - Activation energy - Catalysts

11 - CHEMICAL BALANCE - The balance in homogeneous systems – chemical equilibrium Law - Factors affecting the balance - Ionic equilibria in aqueous solution - Dissociation of water and pH -Theory of acids and bases: Acids and Bases according to Arrhenius, Bronsted, and Lewis - ampholytes - pH of salt solutions (hydrolysis) - Buffer solutions - Calculation of pH in the solution of acids, bases, salts, and buffers - pH indicators. Numerical applications.

12 - Electrochemistry - Galvanic cells - Nernst equation - Series of potential standards and its importance - - Electrolysis – Faraday law.

13 - ELEMENTS OF THERMODYNAMICS - Enthalpy - Hess's Law (to be treated prior to chemical bonding) - Entropy (to be treated before the aqueous solutions) - Free Energy - Temperature Role in the spontaneity of chemical reactions (to be treated before electrochemistry)

14 - Inorganic Chemistry - metals and non-metals: general information on the chemical and physical properties, natural and biological relevance. General characteristics of each group of the periodic Table. Main oxidation states and compounds of hydrogen, alkali metals, and alkaline earth metals, Carbon, Nitrogen, Phosphorus, Oxygen, Sulfur, and Chlorine. Transition elements: general information. Coordination compounds of biological relevance.

The underlined parts, and everything else that is necessary to pass the test for the programmed number (see the website of Biological Sciences),  are discussed in depth at the basic courses to be held in the afternoons of the first two weeks of the course. 

• Chemistry & Chemical Reactivity by Kotz, John C., Treichel, Paul M., Townsend, John, Treichel, David. A.
• The material provided on Studium

The exam consists of a written test and, in case of a favorable outcome, a subsequent oral interview. Both focus exclusively on the topics covered in class. All types of exercises for the exam will be covered during the course. The written exam consists of exercises and open-ended questions. IT IS NOT POSSIBLE TO TAKE THE WRITTEN TEST IN A DIFFERENT SESSION FROM THE ONE IN WHICH YOU WANT TO TAKE THE ORAL TEST. A midterm exam will be given.

The student must approach the written exam with a complete preparation of the entire subject matter. The minimum passing grade for the written test is 18/30. Evaluation criteria for the written test include the correctness of numerical results, the explanation of the procedures used to obtain them, the internal coherence between logically interdependent results, and the precision in using the correct units of measurement associated with the physical quantities used. Evaluation criteria for the oral exam are quantitative rigor in demonstrations, the depth of understanding of the topics, and the ability to establish connections between different aspects of a chemical phenomenon. The acquisition of knowledge from laboratory experiments is assessed based on correctness, completeness, conciseness, and the quality of expression in the preparation of reports. Knowledge of CHEMICAL NOMENCLATURE is an essential requirement without which it is impossible to pass the exam.

Textbooks, formula sheets, course notes, or periodic tables are not allowed in the written test, nor is the use of mobile phones, even in calculator mode. Only the use of a calculator is permitted. The student must present a valid identification document.

Nomenclature

Lewis structures of model molecules: energy diagrams, molecular geometry according to VSEPR, and central atom hybridization

Stoichiometric calculations

Colligative properties

pH calculation in saline and buffer solutions: acid-base reactivity

Balancing redox reactions

Chemical Equilibrium

Solutions and colligative properties

The periodic table

Electronic configuration of elements

First and Second Law of Thermodynamics

Entropy

Gibbs energy

Description of the chemistry of main group elements

Chemical battery

Chemical kinetics

Mixtures of acids and bases