Scheme of Work for Pre U 1 Chemistry (2010) 1st Term: Week 18th Topic Orientation and explaination of packagesand division into classes. Format of STPM 2. Electronic Structure of Atoms 2.1 Concept of electronic energy levels shown by line spectra of atomic hydrogen. 2.2 Atomic orbitals 2.3 Filling of orbitals according to their energy and the pairing of electrons. Objectives Management of the students Explain the STPM syllabus Divide the students into classes. Activity 1. Use LCD and MS Powerpoint to explain. 2. Motivation of the students. 19th 21st Students should be able to 1. explain the formation of the spectrum of atomic hydrogen. 2. calculate the ionisation energy of an atom from the Lyman series converging limit. 3. describe the number and relative energies of the s, p, and d orbitals for the principal quantum numbers 1,2, and 3. 4. describe the shape of the s and p orbitals. 5. predict the electronic configuration of atoms and ions given the proton number (and charge). 6. explain and use the Hund’s rule and Pauli Exclusion Principle in the filling of orbitals.( including the 4s orbitals.) 1. explain the development of the Periodic Table by Newlands, Mendeleev and Moseley. 2. use the Aufbau principlefor electronic configuration of atoms with proton numbers 1 to 30. 3. identify the elements in Gp 1,2,13,14, 15, 16, 17 and 18 and elements in Periods 1 to 4. 4. explain the position of elements in the Periodic Table. 5. explain the position of f-block elements in the Periodic Table. 6.explain the trend and gradation of atomic radii,melting points, boiling points, enthalpy changes, vaporization and electrical conductivities in terms of structure and bonding. 7. explain the factors influencing ionization energies 8. explain the trend in ionisation energy across the second and third periods, and down a group. 9. Predict the electronic configuration and position of unknown elements in the Periodic table from successive values of ionization energies. 10. explain the almost similar physical properties such as density and melting point, in terms of bonding, metallic and ionic radii. 1. Explanation and notes taking. 2. Drawing the line spectrum and spectrometer. 3. Explanation by examples. 4. Written exercises. 22nd – 23rd 3 The Periodic Table 3.1 Development of the Modern Periodic Table 3.2 Building of the Periodic Table 3.3 Elements in Gp 1,2,13,14,16,17, and 18 and Periods 1 to 4. 3.4 Classification of elements into the s,p,d and f blocks. 3.5 Variation in the physical properties with proton number across the second and third periods. 3.6 Variation in physical properties of the first row d-block elements. 1. Set induction to recall the periodic table and pictures of scientist. Student form groups to give presentation on the history of development of Periodic table. Explanation and notes taking. Write the electron configuration of the transition metals. Do the past year questions. 2. 3. 4. 5. 1 octahedral. 2 .. the effects of concentration and temperature on the rate of a reaction. 10. 9.2 Collision theory 5.4 The effect of temperature on rate constants. V-shaped and pyramid.and second-order reactions. Arrhenius’s equation and the Boltzmann distribution.5 The role of catalyst in reactions 5. order of reaction. deduce the order of a reaction by the initial rates method and from concentrationtime graphs. half-life of first-order reaction. 8. Refer to textbooks for exercises. and activation energy..ions. 6. 2. 1.11. 1st Term Holiday (5 June – 20 June 2010) 2 Term: 4. tetrahedral. Reaction Kinetics 5. 5. (d) Co-ordinate covalent bonding (e) VSEPR Model. activation energy. Form teams to surf the internet and do their presentation on MSPowerPoint. 5. and O 7. 3. explain the covalent properties of ionic compounds such as Al2O3. and CN. explain and use the terms rate of reaction.1 Rate of reaction 5. 5. suggest and experimental technique for studying the rate of a given reaction.calculate t for a first-order reaction 7.3 Rate law 5. 4. explain metallic bonding. trigonal planar.2 Covalent bonding (a) Covalent bonding in SO42. explain the Lewis structure of SO42. and O atoms (c) The existence of ionic propertiess in molecules and covalent properties in ionic compounds. Do short notes by own references. Explanation by examples. and catalyst.8 Half-lives of firstorder reactions 1. 4. CO33. rates of reaction. 3. calculate a rate constant from initial rates. explain the variation in successive ionization energies. 5. NO3. 2. 3.and CN. NO3. 8. explain typical properties associated with electrovalent and covalent bonding. van der Waaals’ forces and metallic bonding.. rate determining step. nd 26th – 28th 1. Discuss the answers 4. 2. Draw ‘dot and cross’ diagrams. (a) Van der Waals’ forces (b) Hydrogen bonding and its effect on physical properties 5.ions. explain the concept of overlapping and hybridization of the s and p orbitals for the C.predict and explain the shape of molecules and ions using the principle of electron pairs repulsion.4 Intermolecular forces between molecules. 4. Draw the Boltzman Distribution. 4. first. the formation of conduction and valency bands.1 Electrovalent/ionic bonding 4.6 The order of reaction and rate constants for zero-. eg linear. (b) Hybridisation of the s and p orbitals for the C. predict an initial rate from rate equations and experimental data. Chemical Bonding 24th – 26th 4.. explain hydrogen bonding. N. 5.7 The determination of the orders of reaction and the rate constants. explain qualitatively.. N. Do the worksheets 3. rate equation. CO32. in terms of collision theory.and second-order reactions.3 Metallic bonding 4. AlI3 and LiI. Explanation and notes taking. 9. Written exercises.use integrated forms of rate equations to determine zero-. first. explain the relationship between the rate constants with the activation energy and 1.explain electrovalent and covalent bonding in terms of ‘dot and cross’ diagrams. 5. 2. trigonal bipyramidal. rate constant. 11.5 Buffer solutions 6. pOH . 1.1 The Arrhenius. Explanation and notes taking.1 Chemical equilibria 6. 6. pOH. pKa 3. Show awareness of the importance of an understanding of chemical equilibrium in chemical industry. Kp and Kc .2. 10. pH. 1.2. 3. select suitable indicators for acid-base titrations. 27th 6. 12. 3 . deduce expressions for equilibrium constants in terms of concentration. Written exercises. Calculate the values of Ka. Equilibria 29th – 30th 6.2. explain the significance of the ionic 1. 1. explain qualitatively the different properties of strong and weak electrolytes. explain the distribution of molecular energy with the Boltzmann distribution.2 Ionic Equilibria 31st – 33rd 6. 2.temperature using Arrhenius’s equation.5 Equilibrium constant in terms of partial pressures and concentrations. 2.1 Reversible reaction. Give exercises on the calculations.1. 4. 9.4 Factors affecting chemical equilibriaL Le Chatelier’s principle. 6.2. Ka .1. Kb and pKb.1. explain the effect of catalysts on the rate of a reaction.3 Dissociation constants. Bronsted-Lowry. Kp. 4. Kc and partial pressures. 5. 7. and Lewis theories to explain acids and bases.2 Mass action law and derivation of equilibrium constants. calculate the quantities present at equilibrium from given appropriate data. 2.explain a reversible reaction and dynamic equilibrium in terms of forward and backward reactions. 11. equilibria 6.2. 6.2 the degree of dissociation of weak acids and bases as the basis of strong/weak electrolytes. July Monthly Test 6. 6. Bronstead-Lowry. use Arrhenius. explain how a catalyst lowers the activation energy of a reactions and enzymes as biological catalysts. Explanation and notes taking.3 Homogenous and heterogenous equilibrium constants.6 Heterogenous equilibria of ions. calculate pH from the H3O+ ion concentration for acids(monobasic) and strong and weak bases. explain changes in pH during acid-base titrations in terms of strengths of acids and bases. Discussion of past year questions. 8. identify conjugate acids and bases.1. 6. Ksp and the common ion effect. 6. explain the effect of temperature on equilibrium constants and the equation lnK = -H/RT + C. 5. 3. 10. Explain the graphs drawn. and Lewis theories of acids and bases. Written exercises.1.4 Titration indicators as acids or bases. 6. explain and use the terms pH.pKa . State examples of equilibria which are referred to and studied in general.pKa .2. 3. 2. 6. Redox potential. explain and use the term solubility product. Predict the product liberated during electrolysis based on electrochemistry 14. 3. 13. electroplating of plastics.4 electrode potentials and the formation of the electrochemical series. and its use in calculations. and anodisation 17. 7. Calculate the number of coulomb used. Written exercises. 10. and the charge on the electron 15. explain the methods used to determine standard electrode redox potentials. Explain the concept of Nernst equation.determine the direction of electron flow in a simple cell of given electrode potentials. recycling of aluminium. the mass of material. Electrochemistry. Discussion of past year questions. and/or gas volume liberated during electrolysis. Kw . 14. Ksp and calculate Ksp from given concentration. explain the uses of electrolysis in the manufacturing of aluminium. 7. construct redox equations using halfequations. 12.8 The mechanism of electrolysis as opposed to an electrochemical cell. lower mass and higher voltage. 12. explain the standard hydrogen electrode 4.2 Half-reaction and half-cell.Calculate the quantity of product liberated during electrolysis. 2.9 Faraday’s first and second laws and their use. 16. 7. spontaneous and nonspontaneous electrode reactions. uses and its importance in biological systems. 8.6 Nernst equation and its use. explain the electrochemistry principle in the prevention of rusting and in dental filling. define buffer solutions. 9.11 Manufacturing of chlorine by the electrolysis of brine 1.10 Extraction and manufacturing of aluminium. 7. Explanation and notes taking. treatment of effluent(Ni. Calculate E0cell from the concentration value of solutions using the Nernst equation. the Avogadro constant. calculate the emf of a cell using the E0 values and write the redox potentials.product of water. and Cd). 5.3 Redox reactions and electromotive force (emf) – of cells and the standard hydrogen electrode. predict the possibility of precipitation from given data of concentration of solutions. 16. 7. 4. 15. Mid-term break 4-12 September 2010 Hari Raya Puasa 10 & 11 September 2010 7. 6. 7. 7. predict the stability of aqueous ions from E0value 7.1 Oxidation and reduction reactions in the Daniel cell. 7. 3. explain the advantages of recycling aluminium compared with extracting 1. explain the functions and cell diagram of the Daniel cell 2.5 Cell potentials from the combination of various electrode potentials. 7. predict the feasibility of a reaction from E0cell value. describe the importance of the development of better improved batteries for electric cars in terms of smaller size. 11. 7. calculate the pH of buffer solutions. 4 . 34th – 36th 7. explain the common ion effect including buffer solutions. Cr.7 The use and principles of electrochemistry in the prevention of corrosion. 13. Define and explain the the relationship between the Faraday constant. the energy changes can be exothermic or endothermic. Brief introduction of today experiment. Scheme of work for Pre U 1 Practical Chemistry (2010) 26th Practical 1: Volumetric Analysis – acid base Purpose: To determine the relative molecular weight of a tribasic organic acid. 6. 4. H3A. combustion. 2. Assessment. 6. Explanation by examples. End of the Year Examinations Muet written 13 Nov 2006) 1. neutralization. fusion. 4. 2. use the correct technique of using the burette. 3. 4. Calculate the relative molecular weight of H3A. combustion. ∆H of reactions.05 cm3 in the correct space. 5. July Monthly Test Students should be able to: 1. principally in the form of heat energy. 5. follow the procedure systematically.explain that most chemical reactions are accompanied by energy changes. aluminium 18. solution. 4. 2. explain the manufacturing of chlorine by the electrolysis of brine (12 – 14 & 18 – 20 October. Explanation of procedure. solution. and pipette to measure the volume of a solution. State Hess’ law and its use to find enthalpy changes that cannot be determined directly. Record titration readings to the nearest 0. 2. atomization. calculate the terms enthalpy change from experimental results.4 The Born-Haber cycle for the formation of simple ionic crystals and their aqueous solution.1 Enthalpy changes. 8. use the correct technique of using the burette. eg. 8.05 cm3 in the correct space. Do past year questions. Written exercises. 1. follow the procedure systematically. ionisation energy. 1. Muet Speaking) 39th – 40th 8. 5.either by mercury cathode cell or diaphragm cell. Thermochemistry and chemical Energetics 41st – 42nd 8. 5 . Give exercises on the calculations. 2. 3. Record titration readings to the nearest 0. Practical session.3 Lattice energies for simple ionic crystals. Calculate the concentration of H3A based on the titration readings. Explanation of procedure. 5. Teacher demonstrates how to determine the end-point. 2. Brief introduction of today experiment. define the terms enthalpy change of formation. and pipette to measure the volume of a solution. hydration. 27th 28th Practical 2: Volumetric Analysis – acid base Purpose: To find the exact concentration of a mineral acid X. calculate the heat energy change from experimental measurements using the relationship energy change = mc ∆T 3. 3. Teacher demonstrates how to determine the end-point. 3. Answering questions. Calculate the concentration of 1. 4. Students should be able to: 1. an enthalpy change of formation from enthalpy changes of combustion.2 Hess law 8. neutralization. formation. atomization. 8. 6. Draw the BornHaber Cycle. 3. explain the terms enthalpy change of reaction and standard conditions. A qualitative appreciation of the effects of ionic charge and ionic radius on the magnitude of lattice energy. hydration. Explanation and notes taking.5 The solubility of solids in liquids. follow the procedure systematically. apparatus with the solution. Practical session. Assessment. 2. use the correct technique of using the burette. 4. 31st Experiment 2: Volumetric Analysis Redox Purpose: To determine the ratio of the number of moles of hydroxyammonium ions to the number of moles of iron (III) ions participating in the reaction. Students should be able to: 1. employ the correct method of rinsing the burette and pipette with the solution to be measured before filling up the apparatus with the solution.05 cm3 in the correct space. 4. 2. and pipette to measure the volume of a solution. 6. Practical session. 30th Practical 3: Volumetric Analysis . and to determine the relative atomic mass of the element X. 3. 3. use the correct technique of using the burette. Assessment. Students should be able to: 1. employ the correct method of washing off the solution from the sides of the titration flask with distilled water before reaching the end-point. 3.mineral acid X based on the titration readings. Calculate the concentration of ethanedioate ions and KA1. procedure. 4. employ the correct method of heating. 2. Brief introduction of today experiment. employ the correct method of rinsing 4. employ the correct technique of transferring solution from the pipette into the titration flask. Record titration readings to the nearest 0. and hence identify X. 5. 4. Brief introduction of today experiment. Answering questions. employ the correct technique of transferring solution from the pipette into the titration flask. follow the procedure systematically. 5. employ the correct method of washing off the solution from the sides of the titration flask with distilled water before reaching the end-point. 2. 4. 6. 1. Brief introduction of 1. Answering questions. 3. Assessment. 7. 4. follow the procedure systematically. the burette. employ the correct technique of titration. 7. employ the correct technique of titration. Explanation of procedure. the burette and pipette with the solution 5. use the correct technique of using today experiment. 5. Explanation of procedure. 2. 6. 3. 5. 29th Experiment 1: Volumetric Analysis – acid base Purpose: To determine the exact concentration of a mineral acid.Redox Purpose: To find the concentration of potassium permanganate solution. 6 . calculate the relative atomic mass of the element X. Practical session. Assessment. 3. calculate the number of moles of iron (III) ions required to oxidize 1 mol of hydroxyammonium ions. Teacher demonstrates how to determine the end-point. Practical session. HXO4. Explanation of volume of a solution. 5. 5. and pipette to measure the volume of a solution. Answering to be measured before filling up the questions. and pipette to measure the 2. 1. Students should be able to: 1. Students should be able to: 1. Answering questions. 2. Students should be able to: 1. Use the correct method of holding and shaking the titration flask during the titration process. ethanedioate and 5. calculate the concentration of monobasic acid.. calculate the percentage of sodium ethanedioate in solution KA1. Explanation of procedure. 5. 2. Explanation of procedure. 6. Practical session. Practical session. use the correct technique of using the burette. prepare a solution 4. employ the correct method of ethanedioate used to heating. Answering questions. Checked and verified by. Practical session. Kadir) Pre U Senior Assistant 7 . 3. employ the correct method of hydrated ethanedioic washing off the solution from the sides acid. calculate the concentration of thiosulphate (VI) solution in gdm-3. Experiment 4: Volumetric Analysis Redox Purpose: To standardize sodium thiosulphate (VI) solution using potassium iodide solution. 2. Assessment. 4. and pipette to measure the volume of a solution. HX. Explanation of procedure. 3. 5. 1. 4. 2. HX. 3. Assessment. (Ms. 1. 3. employ the correct technique of containing sodium titration. Brief introduction of today experiment. Ibrahim B. Purpose: To determine 2. Assessment. the mass of sodium 3. 4. and pipette to measure the volume of a solution. Prepared by. employ the correct technique of handling the pipette. use the correct technique of using the burette. 5. 3.33rd Mid-term break 4-12 September 2010 Hari Raya Puasa 10 & 11 September 2010 Experiment 3: Students should be able to: Volumetric Analysis – 1. 4. 34th 1. 2. (En. Answering questions. follow the procedure systematically. 35th Experiment 5: Volumetric Analysis stoichiometry Purpose: To determine the exact concentration of a monobasic acid. Brief introduction of today experiment. 6. Brief introduction of today experiment. ……………………………. of the titration flask with distilled water before reaching the end-point. Lau Chai Wee) (Subject teacher) ……………………………………. use the correct way of adding the indicator. and pipette to measure the volume of a solution. follow the procedure systematically. 5. use the correct technique of using acid base and redox the burette. 4. addition of indicator at the appropriate time. plan the procedures of experiment systematically and use reasonable quantities of substance. 8 .