GROUP 4 Experimental Sciences

Transcription

1 GROUP 4 Experimental Sciences CHEMISTRY The course involves studying fundamental aspects of Chemistry, the knowledge of which allows not only estimate the behaviour of a chemical reaction but also to suggest its direction and predict its possible resultant. For this purpose, the course is designed according to the principle from small to big. It is assumed that it is possible to develop the analytical approach, allowing to better comprehend the chemical properties, develop intuition in Chemistry and make a link between laws in Chemistry and laws in Physics and Biology only when one comprehends the atomic structure, the periodic law, energy laws and many other fields of General Chemistry. Special attention is paid to experiments and the students' participation in them. Through studying any of the group 4 subjects, students should become aware of how scientists work and communicate with each other. While the scientific method may take on a wide variety of forms, it is the emphasis on a practical approach through experimental work that distinguishes the group 4 subjects from other disciplines and characterizes each of the subjects within group 4. It is in this context that all the Diploma Programme experimental science courses should aim to: 1. Provide opportunities for scientific study and creativity within a global context that will stimulate and challenge students 2. Provide a body of knowledge, methods and techniques that characterize science and technology 3. Enable students to apply and use a body of knowledge, methods and techniques that characterize science and technology 4. Develop an ability to analyse, evaluate and synthesize scientific information 5. Engender an awareness of the need for, and the value of, effective collaboration and communication during scientific activities 6. Develop experimental and investigative scientific skills 7. Develop and apply the students information and communication technology skills in the study of science 8. Raise awareness of the moral, ethical, social, economic and environmental implications of using science and technology 9. Develop an appreciation of the possibilities and limitations associated with science and scientists 10. Encourage an understanding of the relationships between scientific disciplines and the overarching nature of the scientific method. It is the intention of all the Diploma Programme experimental science courses that students achieve the following objectives. 1. Demonstrate an understanding of: a. scientific facts and concepts c. scientific terminology d. methods of presenting scientific information. 2. Apply and use: a. scientific facts and concepts c. scientific terminology to communicate effectively d. appropriate methods to present scientific information. 3. Construct, analyse and evaluate: a. hypotheses, research questions and predictions c. scientific explanations.

2 4. Demonstrate the personal skills of cooperation, perseverance and responsibility appropriate for effective scientific investigation and problem solving. 5. Demonstrate the manipulative skills necessary to carry out scientific investigations with precision and safety. COURSE PROGRAMME CHEMISTRY IB DP DP YEAR 1 1. Basics of Chemistry. Qualitative Chemistry. a) Elements and compounds. The mole concept and the Avogadro constant. Relative atomic and molecular mass. Empirical, molecular and structural formulae. b) Experimental methods of formula determination. Percentage VS. formula. The concept of isomers. Chemical equations and stoichiometry. c) The concept of state : solids, liquids, gases. Measuring the amount of a substance. Solutions. Gases and the ideal gas equation. How to determine the molecular mass of a gas experimentally. The Avogadro law. d) Excess and lack of reagents. Acid-base titration. 2. Atomic structure. a) The historical evolution of the atomic theory and the element concept. Thomson and Rutherford s experiments. Nuclones. b) Atomic unit (mass). Atomic number, mass number. Isotopes. Uses of isotopes: carbon dating, isotopes in medicine. Mass spectrometry principles. Deciphering mass spectra. c) Electronic structure. Electromagnetic spectrum. Wavelengths and frequencies, Planck constant. Emission spectrum. Electron transitions. d) Ionization energies: first, second, etc. Sublevels and quantum numbers. Pauli exclusion principle, Hund s rule of maximum multiplicity, the aufbau principle. The shape of s- and p-orbitals. Electron configuration and the Periodic Table. 3. Periodicity. a) History of the periodic law. Groups and periods. The table s organization in accordance with the location of electrons on a certain sublevel. b) The relation between ionization energy and atomic number. Electronegativity. Atomic radius: trends in periods and groups. The Heisenberg principle. The x-ray diffraction. Ionic radius:: trends in periods and groups. Cations and anions. Melting and boiling points: trends in periods and groups. c) Properties of elements belonging to various groups. Alkali metals: reactions with water and halogens. Halogens: their oxidative ability. Comparison of elements properties across a period. 3rd period: how properties of corresponding compounds change across the period. Oxides, chlorides. d) d-elements: the definition of a transition element. Sets of oxidation numbers. Complex ions. Ligands and coordination numbers. Colours of complexes. Catalytic activity. 4. Bonding in chemistry. a) Exothermic and endothermic reactions. Intramolecular and intermolecular forces of attraction. Ionic bond. Electron affinity. Lattice enthalpy. Polyatomic ions. b) Covalent bond. The notion of valence. Lone pairs and vacant orbitals. Double and triple bonds. Lewis structures and various ways to represent them. Coordination bonding (various mechanism for the formation of covalent bonds). Electron density and delocalization. Bond length VS. bond strength. Resonance structures. c) From ionic to covalent: the transition. The concept of bond polarity. VSEPR theory and the shape of molecules. Predicting the shape of molecules with two, three or four negative charge centers. Modeling structures of molecules. d) Intermolecular interactions. Van der Waals bonding and the factors affecting its strength. Dipoledipole interactions, hydrogen bonding as a particular case. How to determine the differences

3 between polar and nonpolar molecules experimentally. Boiling point of a mixture VS. its composition. e) Giant covalent structures AKA atomic lattices (diamond structure as an example). Comparison of diamond s and graphite s structures. Metallic bonding. Using metals in industry. A compound s physical properties dependency on the type of intra- and intermolecular bonding. Basis of valence bond theory and molecular orbital theory. f) Molecular orbitals: bonding and anti-bonding. Sigma- and pi-bonds. Paramagnetism and bond order. g) Hybridization. Electrophilicity. The relation of hybridization type and the shape of a molecule. Shapes related to having five or six negative charge centers in a molecule. Delocalization, delocalization enthalpy. 5. Energetics. a) Exothermic and endothermic reactions. Heat change and enthalpy. Standard reaction enthalpy and standard conditions. Reactions taking place under constant pressure. Calorimetric experiments for homogeneous and heterogeneous reactions. b) Hess s Law. Treating chemical equations as mathematical equations, manipulating reaction cycles. Hess s Law corollaries: enthalpies of formation and combustion. Greenhouse effect. Bond energies and how to employ them to calculate enthalpies of reactions. c) Standard reaction enthalpies and phase changes. Born-Haber Cycle and standard atomization enthalpy. Lattice enthalpy. d) Entropy and free energy. Spontaneous processes. The second law of thermodynamics. 6. Chemical kinetics. a) Reaction rates. Different methods of measuring reaction rates. Kinetic curves and how to elucidate reaction rates graphically. b) Collision theory and the conditions of chemical reactions. Maxwell-Boltzmann distribution, average kinetic energy. c) Factors determining rates of chemical reactions: temperature, concentration, the presence of a catalyst, particle size. Activation energy and it relation to enthalpies of reaction. d) Mathematical equations used to define and calculate chemical reaction rates. Law of mass action and rate constants. Order of reaction. How to find out rate constants using experimental data. Reactions of the zero, first and second order - graphical representation. e) Reaction mechanisms. Activation energy and the Arrhenius equation. The notion of chemical oscillators. 7. Equilibrium. a) Static VS. dynamic equilibrium. Isolated systems. Forward and reverse reactions. The position of equilibrium in time. b) Equilibrium in homogeneous systems. Phases. Equilibrium constants. How the prevalence of products or reagents in the system affects the value of the equilibrium constant. Le Chatelier s principle: temperature, pressure, the concentration of reagents and products, catalyst. c) Equilibrium in chemical industry. The Haber process. Optimal temperature. The contact process for the production of sulfuric acid. d) Liquid-gas equilibrium, steam pressure. Relation between enthalpies of boiling, boiling points and the strength of intermolecular interactions. Shift in the Maxwell-Boltzmann distribution. e) The composition of reaction mixtures at equilibrium: liquid phase, gas phase. Units used for the equilibrium constant. DP YEAR 2 8. Acids and bases. a) Arrhenius, Bronsted-Lowry and Lewis theories. Strong and weak acids. Amphiprotic species, conjugated acids and bases. Lewis acids: complexes (central atoms/ions and ligands). Concentrated and diluted acids and bases. b) Chemical properties of acids and bases: neutralization, reactions with indicators, acids interacting

4 with metals and carbonates, displacement of ammonia from ammonium salts by bases. c) The ph scale and indicators. Acidic and basic media. ph meters. Acid rains. d) The ionic product of water and its dependency on temperature. Calculating ph and poh. Dissociation constants for acids and bases. Hydrolysis of salts VS. acidity of the medium. Buffer solutions. Blood and isotonic solutions. e) Acid-base titration. Various combinations of strong/weak acids and strong/weak bases. 9. Redox reactions. a) The concepts of oxidation and reduction, the concepts of oxidants and reductants. Combustion as an example of oxidation. Half-reactions. Oxidation numbers and the algorithms used to calculate them. Disproportionation and comproportionation. IUPAC and the nomenclature of inorganic compounds. b) Electron method for balancing chemical equations, ion electron method for balancing chemical equations. Typical oxidants and reductants. Reactivity series (metals) and the comparison of halogens reactivity. c) Electrolytic cells and voltaic cells, their structure, half-reactions. Cathodes and anodes. Salt bridges, half-cells. Spontaneous electrochemical reactions, electrolysis of molten salts. d) Significance of electrolysis for industrial uses, yielding elements. Standard electrode potentials and standard electrodes. Finding out if an electrochemical process is spontaneous or not. e) Electrolysis of water solutions. Electrolysis of water and of sodium chloride and copper sulphate solutions. Amalgams. Predicting the products of an electrolytic process. Electroplating. 10. Organic chemistry. a) Properties of carbon atoms that allow them to provide a foundation for organic chemistry. Alkanes, common formula. Homologues and homologous series. Structural isomerism, nomenclature of alkanes and their physical properties, paraffins. b) Incomplete combustion. Homolytic and heterolytic bond fission. Radical substitution and the concept of free radicals. Chain reactions: stages of initiation, propagation and termination. Photoinduced halogenation of alkanes. The ozone layer, freons and their decomposition. c) Naming nomenclature for organic substances with various functional groups. Alcohols, haloalkanes, aldehydes, ketones, carboxylic acids, esters, amines, aromatic compounds. Physical properties of compounds belonging to various classes. Volatile fluids. d) Alkenes, their geometry, addition reactions: hydrogen, halogens, etc. Saturated and unsaturated hydrocarbons. Mineral oil, cracking. Oils and fats, margarine. Examples of various polymers. e) Alcohols: combustion, oxidation to carbonyl and carboxyl compounds. Primary, secondary and tertiary carbons. f) Primary and secondary/tertiary haloalkanes. The notions of electrophiles and nucleophiles. Sn1 and Sn2 mechanisms: monomolecular and bimolecular processes. The concept of inductive effect. Determining mechanisms and reaction pathways for organic reactions. g) Additional classes of compounds: amides, nitriles. Nucleophilic substitution reactions and factors affecting them. Elimination reactions. Condensation reactions (esterification). Amino acids and proteins. Polyamides. h) Stereoisomerization and geometric isomerization. Chiral carbons and enantiomers. Polarimetry, racemic mixtures. 11. Errors and precision. a) Repeatability, reproducibility and the selection of samples. Precision of measurements. Systematic and random errors. b) Rounding and significant digits. Processing graphical data, selecting axes. OPTION A: MODERN ANALYTICAL CHEMISTRY a) Chromatography: adsorption, stationary and mobile phases, partition coefficient. GL, paper (see the concept of eluent ), TL, column chromatography. b) Spectroscopy: emission and absorption. IR and vibrations, wavenumbers, characteristic bands. Mass spectrometry and the analysis of peaks in spectra. NMR and chemical shift, MRT. Atomic absorption spectroscopy and determination of concentrations.

5 c) UV/Vis spectroscopy. The colour wheel and additional colours. Spectrochemical series. UV and organics. High-definition NMR, line multiplicity. HPLC, the structure of a chromatograph and deciphering of chromatograms. OPTION F: FOOD CHEMISTRY а) Vital nutrients and minerals, basic groups of products. The concept of a healthy diet. б) Oils and fats, melting point. Various fatty acids. Expiration date and rancidity. Various mechanism of expiration, hydrolytic and oxidative rancidity. в) Antioxidants, examples. Pigments, colourants, examples. Caramelization, Maillard reaction. г) GMO: pros and cons. Dispersed systems: suspensions and emulsions, foams. Emulsifiers. д) Oxidative rancidity: detailed mechanism. Antioxidants grouped according to mechanisms. Chirality. Detailed description of colourants structures. + Laboratory works STANDARD LEVEL FINAL ASSESSMENT Paper 1: 30 multiple-choice questions on the core Paper 2: Section A: one data-based question and several short-answer questions on the core (all compulsory) Section B: one extended-response question on the core (from a choice of three) Paper 3: Several short-answer questions in each of the two options studied (all compulsory) HIGHER LEVEL Paper 1: 40 multiple-choice questions (±15 common to SL plus about five more on the core and about 20 more on the AHL) Paper 2: Section A: one data-based question and several short-answer questions on the core and the AHL (all compulsory) Section B: two extended-responsequestions on the core and the AHL (from a choice of four) Paper 3: Several short-answer questions and one extended-response question in each of the two options studied (all compulsory) + Reports on laboratory works + Group 4 Project