8. Draw Lewis structures and determine molecular geometry based on VSEPR Theory

Transcription

1 Chemistry Grade 12 Outcomes 1 Quantum Chemistry and Atomic Structure Unit I 1. Perform calculations on wavelength, frequency and energy. 2. Have an understanding of the electromagnetic spectrum. 3. Relate electron energies to hydrogen atom and electromagnetic spectrum. 4. Relate the structure of the Bohr atomic model to energy and frequency. 5. Write electronic configurations and identify trends. 6. Calculate ionization energy and relate it to periodic trends. 7. Determine bond type and electron arrangement based on electronegativities. 8. Draw Lewis structures and determine molecular geometry based on VSEPR Theory Unit II Kinetics 1. Define rate in terms of some quantity per unit time. 2. Identify how variables (e.g. pressure, temperature, ph, conductivity, etc.) could be used to monitor reaction rates. 3. Design experiments to measure the rate of a chemical reaction. 4. Given data, calculate the average and the instantaneous rates for a chemical reaction using, x/ t. 5. Relate the rate of reaction of a reactant or product to the stoichiometry of a chemical reaction. 6. Identify factors that affect the rate of a reaction, these being: (a) nature of the reactants (b) changes in concentration (c) surface area (d) changes in temperature (e) presence of catalysts 7. Explain how the rate of a reaction can be affected by the "nature of the reactants". 8. Explain using "collision theory" how changes in concentration affect the rate of a reaction. 9. Explain why some reactions take place in a series of steps. 10. Determine the net equation for a reaction, given the reactions in the mechanism. 11. Differentiate between a homogeneous and a heterogeneous reaction and explain how changes in concentration affect the rate of each. 12. Draw and interpret potential energy diagrams for endothermic and exothermic reactions.

2 13. Predict the relative rate of a chemical reaction based on reaction coordinate diagrams with varying activation energies Draw a kinetic energy distribution curve and explain the relationship between temperature and rate. 15. Discuss the effect of temperature changes on the shape of a kinetic energy distribution curve. 16. Given experimental data determine: (i) the order of a reactant in a reaction, (ii) the overall order of the reaction, and/or (iii) the rate law for a reaction 17. Calculate the value of the specific rate constant, given the rate law for a reaction and some experimental data. 18. Construct and interpret rate versus concentration and concentration versus time graphs of zero, first and second order reactions given experimental data. 19. Explain how changes in temperature affect the rate of reaction. 20. Explain how a catalyst affects the rate of the reaction. 21. Explain the effect of a catalyst on the rate of a reaction using reaction coordinate diagram as well as a kinetic energy distribution curve. Unit III Chemical Equilibrium 1. Explain the term "equilibrium" as it applies to physical systems, (e.g. solubility, vapour pressure, dissolving, etc.) 2. Differentiate between a physical system at equilibrium and one that isn't. 3. Sketch graphs of concentration, (amount) versus time for physical equilibrium systems. 4. Compare and contrast physical and chemical equilibria. 5. Define equilibrium constant, Kc, and mass action expression. 6. Distinguish between homogeneous and heterogeneous equilibria. 7. Write an equilibrium law expression, given an equation. 8. Calculate equilibrium constants, Kc, from experimental data and a balanced equilibrium equation. 9. Calculate the equilibrium concentration of a substance in a reaction, given the Kc and experimental data. 10. Calculate a trial Kc, or Q value for an equilibrium equation from given experimental data and predict whether or not the equation is at equilibrium. 11. Relate the magnitude of the equilibrium constant, Kc, to the relative concentrations of reactants and products. 12. Given an equation for a reaction at equilibrium and the H, predict the effect of changes in: (i) concentration, (ii) pressure, (iii) volume and (iv) temperature on the equilibrium state. 13. Apply Le Chatelier's Principle to predict how a system at equilibrium is affected by changes. 14. Interpret shifts in equilibrium using reaction kinetics. 15. Sketch and interpret graphs of rate versus time and concentration versus time for chemical equilibrium's.

3 3 16. Describe the application of the Haber process and its relationship to Le Chatelier's Principle. Unit IV Acid-Base Equilibria 1. Differentiate between (i) electrolytes and non-electrolytes, (ii) strong and weak electrolytes. 2. Distinguish between ionic and molecular compounds and their aqueous solutions. 3. Write balanced equations for the dissolving of pure ionic and molecular compounds in water. 4. Define and compare Arrhenius and BrÖnsted-Lowry acid/base theories and use them to identify acids and bases. 5. Describe the limitations of the Arrhenius acid/base theory in identifying acids and bases. 6. Define and identify conjugate acid-base pairs. 7. Write the formula for a conjugate base or acid given the name or formula of the corresponding acid or base. 8. Write balanced chemical equations for acid-base reactions given the names of the reactants. 9. Write balanced equations for the reactions of acids and bases with water. 10. Write equations for the reaction of ionic compounds with water to produce acids and bases, (hydrolysis). 11. Define the ion product of water, Kw. 12. Relate the concentration of hydronium and hydroxide ion in acidic, basic and neutral solutions. 13. Define ph and measure ph experimentally. 14. Relate the ph scale to the relative acid-base properties of substances and given any one of the values of ph, [H3O + ], or [OH - ], calculate the remaining values. 15. Write equilibrium law expressions for acidic, Ka, and basic, Kb, dissociation reactions. 16. Calculate the acid dissociation constant, Ka, given the ph and the concentration of a weak acid solution. 17. Calculate the ph, given molar concentrations of a strong acid or an ionic hydroxide. 18. Calculate the ph of weak acid and base solutions given their molar concentrations and their equilibrium constants, Ka/Kb's. 19. Distinguish between strong and concentrated, and between weak and dilute for aqueous solutions of acids and bases. 20. Predict whether an aqueous solution of a given ionic compound will be basic, acidic or neutral given the formula. 21. Describe and perform an acid/base titration. 22. Differentiate between end point and equivalence point of a titration. 23. Calculate the amount of acid or base present or the molar mass of an unknown acid, given titration data. 24. Calculate the mass-by-volume percent composition of a solution or solid given titration data.

4 25. Prepare and interpret titration curves for acids and bases. (e.g., strong acid- strong base, weak acid-strong base, and weak base-strong acid). 26. Explain how an indicator works in terms of colour shifts using Le Chatelier's Principle Prepare a standard solution, and standardize it against either a primary standard or a standardized solution. 28. Define an acid/base buffer system and classify a given buffer system as either acidic or basic. 29. Describe how a buffer minimizes a change in ph when either an acid or base is added. 30. Write a balanced equation for the equilibrium reaction in a given buffer solution and describe how the concentration of all components change when H 3 O + or OH - are added. 4 Unit V Solubility Equilibria 1. Describe and write a balanced equation to represent the equilibrium in a saturated aqueous solution of an ionic compound. 2. Write a solubility product expression, given a formula or an equation for the compound. 3. Distinguish between solubility and solubility product constant, Ksp. 4. Calculate the solubility of a compound in water, given its Ksp. 5. Calculate the Ksp of a compound, given the solubility or the concentration of its ions in solution. 6. Calculate the solubility of a compound in a solution containing a known concentration of a common ion, given the Ksp for an ionic for an ionic compound. 7. Given the volumes and concentrations of two solutions that are mixed, calculate the amount of precipitate formed, and the concentration of the ions left in solution. 8. Predict what will happen when pairs of aqueous solutions of ionic compounds are mixed and write balanced ionic equations for any expected reactions. 9. Predict the solubility of a given compound in acid solutions from its formula. 10. Write balanced equations to show the dissolving of compounds in acidic solutions. Unit VI Oxidation - Reduction 1. Relate the role of oxygen to the process of rusting and burning. 2. Define oxidation and reduction as: (i) gain and loss of oxygen, (ii) gain and loss of electrons, and (iii) the increase and decrease in the oxidation numbers. 3. Determine the oxidation numbers for atoms in simple compounds and ions. 4. Given an equation, identify the: (i) substance oxidized, (ii) substance reduced, (iii) oxidizing agent and (iv) reducing agent. 5. Define and use correctly the terms oxidation, reduction, oxidizing agent, and reducing agent.

5 6. Identify the characteristics of a redox reaction in terms of electron transfer and distinguish between redox and non-redox reactions 5 7. Describe single replacement reactions in terms of oxidation and reduction. 8. Given experimental data, develop an activity series and use it to predict spontaneous reactions. 9. Balance neutral, acidic and basic oxidation-reduction reactions using the half-reaction method. 10. Perform stoichiometric calculations based on balanced chemical equations of redox reactions. 11. Describe how redox reactions generate an electric current. 12. Draw and label an electrochemical cell, given the substances for the electrodes. 13. Define anode, cathode, half-cell, and half-cell reaction. 14. Define standard reference electrode and standard electrode potential. 15. Use a table of reduction potentials to (i) write the two half-reactions, (ii) write the over-all reaction and (iii) calculate the cell voltage. 16. Use a table of reduction potentials to predict whether or not two substances will react with each other and write the balanced ionic equation for the reaction that occurs. 17. Diagram an electrolytic cell containing a molten ionic compound and describe its function. 18. Describe the uses of electrolytic cells in the electrolysis of water, aqueous salts, electroplating and in the production and purification of metals. 19. Calculate the amount of product produced by passing a given quantity of electricity through an electrolytic cell using Faraday's laws. 20. Explain how batteries use oxidation and reduction reactions to produce a current in an external circuit. 21. Describe the chemical species and commercial uses of the following batteries: dry cells, lead-acid, mercury oxide, silver oxide and nickel-cadmium. 22. Describe the chemical species, half reactions and commercial uses of a fuel cell. 23. Describe the chemical species and commercial uses of electrolytic cells in the purification of copper and nickel metals. 24. Describe the production of aluminum from bauxite (Hall's process). 25. Distinguish between electrochemical, electrolytic and fuel cells. 26. Explain the chemical process of corrosion and methods of corrosion prevention.