Modern chemistry: the science of the composition and structure of materials and the changes they can undergo; definitions of mass and its conservation, Lavoisier.
Fundamental and derived physical quantities and their units of measurement; SI system. Temperature concept.
Matter: physical and chemical properties; elements, compounds, homogeneous and heterogeneous mixtures; system and environment; states of matter, chemical and physical transformations.
Nomenclature of elements and compounds, introduction to the periodic system; chemical formulas, brute formulas; binary compounds: salts, oxides and hydrides; ternary compounds: oxyacids, ternary ions, anions and cations, hydroxides.
The atom: Proust and Dalton laws and their postulates, laws of multiple proportions; Thomson, Millikan and Rutherford experiments. Particles that make up the atom: electron, proton and neutron and their characteristics, charge, mass etc. The nucleus, atomic number and atomic weights, nuclides and isotopes.
Periodic Table of the Elements: metals, non-metals, semimetals; main and transitional elements; alkali and alkaline earth metals, halogens, chalcogens and noble gases.
Periodic properties of the elements: formation of oxides, hydrides, halides, ionization potentials and electron affinities.
Generic subdivision of substances into molecular and ionic, general concept of formula. Electronegativity. Atomic mass expressed as 1/12 of the carbon atom, molecular mass, mass spectrometer. The concept of mole, Avogadro's number, stoichiometry of a reaction. Main ways of expressing the quantity of a substance in a reaction medium: concept of solute and solvent, molarity, % weight % and volume %.
Empirical or minimal formulas and brute formulas. Chemical reactions or equations: balance. Types of reactions: acid/base, double exchange, redox; definition of oxidation number, rules for assigning the oxidation number, concepts of oxidation and reduction. Redox reactions, balancing rules, half reactions.
Quantum theory of matter; emission spectra, electromagnetic spectrum, wavelength and frequency of an electromagnetic wave. History of light, corpuscular and wave nature, wave particle duality, quantization of Planck energy, photoelectric effect, line spectra, limits of the Rutherford atomic model, Bohr atom, Bohr postulates, Rydberg constant, Balmer sequence , excited electronic states. Quanto-mechanics, energy levels, de Broglie's relation, Heisemberg's uncertainty principle, Schrodinger's equation, wave functions, principal quantum number and quantization of energy levels, atomic orbitals, azimuthal, magnetic and spin quantum numbers. Experience of Stern and Gerlach; electronic configurations and orbital diagrams, Pauli's exclusion principle, Hund's rule, Aufbau's principle, periodicity of the elements, noble gases, alkali metals, halogens, chalcogens, reorganization of the periodic table. Magnetic properties of matter, paramagnetism and diamagnetism; review of the periodic properties, atomic and ionic radius, energy of first, second, and other ionization potentials; electronic affinity.
Chemical bond; importance of electronegativity, valence; ionic bond and covalent bond; loss, acquisition and pooling of valence electrons; Lewis notation, electron doublets, octet rule, Slater's atomic rays, single, double and triple bonds, bond order, limit formulas, formal charge, resonance and resonance energy; VSEPR: method for obtaining molecular geometries.
Bond theories; ionic bond, Coulomb's law, bond distance, isoelectronic ions, coordination numbers and coordination geometry, r+/r- ratio, crystal lattices, Madelung's constant, lattice energy, Born Haber cycle. Covalent bond: molecular orbital and valence bond, molecular wave function, LCAO-MO method; sum and difference of atomic wave functions, bonding and antibonding orbitals, σ and π bonds, application to homonuclear diatomic molecules: H2+, H2, He2+, He2, from Li2 to Ne2, paramagnetism of the O2 molecule. Isoelectronic heteronuclear diatomic species, introduction of hybridization, hybrid orbitals, σ and π bonds, non-bonding orbitals, cis/trans isomers;
Organic chemistry: carbon chemistry, classification of the main organic compounds, hydrocarbons, alkanes, alkenes, alkynes, benzene and its derivatives, notes on their reactivity.
Coordination compounds: definition of ligand, Lewis acid base reactions, some of the most famous polidentate ligands.
The gaseous state, definitions of pressure, volume and temperature, different ways of expressing pressure: bar, mmHg, atm, Pascal; Hg barometer; Celsius, Farenheit and absolute temperature scales; gas density; empirical gas laws, Boyle's, Charles's, Avogadro's law, general gas law, gas constant R and its dimensions, molecular weight of a gaseous substance from gas law, Dalton's law on partial pressures; gaseous concentrations, molar fractions, vapor pressure. Kinetic theory of an ideal gas, postulates of kinetic theory, qualitative interpretation of the gas law, ideal gas law from kinetic theory. Real gases, Van der Waals equation.
Thermodynamics and thermochemistry; energy and energy conservation law, kinetic, potential and internal energy; thermodynamic system, heat capacity, specific heat, heat of reaction, exo and endothermic reactions, enthalpy of reaction, state functions, relationship between enthalpy and internal energy, Hess's law. First law of thermodynamics, isothermal, adiabatic, isochore transformations, etc. Concept of reversibility of a process. Entropy and second principle. Third law of thermodynamics. Spontaneous processes, Gibbs free energy.
State changes: phase diagrams in one-component systems. Phase rule, vapor pressure and boiling point, melting point, latent heats of evaporation and melting, Classius-Clapeyron equation, state diagram of water, CO2, sulfur, triple point, temperature and critical pressure.
Condensed states; solid state: properties, crystalline systems, types of solids: molecular, covalent, ionic, metallic; liquid state: properties, again on the vapor pressure, liquefaction of a gas, Andrews diagram, critical temperature; surface tension, viscosity; Intermolecular interaction forces: Van der Waals, London, dipole-dipole, induced dipole-induced dipole, polarizability of a molecule, hydrogen bond, boiling points of hydrides of the 2nd and 14th, 15th, 16th and 17th group.
Solutions and their behavior; types of solutions: solubility and factors that influence it; Henry's law. Concentration and its units. Ideal solutions. Raoult's law. Totally miscible liquid mixtures: liquid-vapor equilibria. Colligative properties of solutions: cryoscopic lowering, ebullioscopic raising, osmosis.
Chemical equilibrium, chemical potentials, free energy of a chemical reaction and chemical equilibria, reaction quotient, law of mass action, molar free energy of formation; thermodynamic constant of reaction, stoichiometric constants: concentration, pressure and mole fraction; relations between the various stoichiometric equilibrium constants. Homogeneous and heterogeneous equilibria. Use of the equilibrium constant, Le Chatelier's principle; industrial examples: Haber process of ammonia production.
Acids and bases; definition of Arrhenius; acid-base balances, self-ionization of water, ionic product of water, introduction of pH, strong acids and bases. Bronsted-Lowry theory, acid-base conjugate pairs, amphoteric; weak acids and bases, quantitative scale of acidity based on Ka (or Kb), relationship between Ka and Kb in an acid-base pair, application of Le Chatelier's principle to acid-base reactions. Lewis acids and bases. Molecular structure and strength of acids, II period hydrides, VII group hydrides, oxyacids; polyprotic acids.
Neutralization reactions, titration curves strong acid strong base and weak acid - strong base; buffer solutions, strong acids and bases such as buffers, weak acid pairs or bases in the presence of their respective base or conjugate acid; effect of the ion to common, acid base indicators. Notes on the equilibrium of precipitation and complexation, constant of the solubility product of a slightly soluble salt, constant of stability of a complex ion.
Chemical kinetics; rate of disappearance of a reagent and formation of a product, kinetic equation, reaction order, first order reactions, linear logarithmic graph, half-life, second order reactions, straight graph 1 / Conc. against time; stoichiometric coefficients and order of reaction, elementary reactions, mechanism of a reaction, slow step that regulates the global speed; catalysts, activation energy, homogeneous and heterogeneous catalysis, enzymes.
Electrochemistry; Electrical work from redox reactions. Batteries and their electromotive force. Normal potentials and their meaning. Nernst equation; common types of electrodes; concentration batteries. Electrochemical measurement of pH. Electrolysis; electrolysis of molten salts; electrolysis of solutions; stoichiometry of electrolysis.
Basics of nuclear chemistry. Alpha, beta (+ and -), gamma decays. Half-life of a nuclear reaction, radiodating.