1 Chemistry 12 Specific Curriculum Outcomes (updated September 18, 2016) Unit A: Thermochemistry Overview: In this unit, students study energy as it relates to chemical changes and quantify the energy involved in thermochemical systems, and consider the various aspects of energy use on society. Focusing Questions: How does our society use the energy of chemical changes? What are the impacts of energy use on society and the environment? How do chemists determine how much energy will be produced or absorbed for a given chemical reaction? Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in other courses. The intended level an scope of treatment is defined by the outcomes. enthalpy of formation energy diagrams enthalpy of reaction activation energy H notation catalysts Hess law calorimetry molar enthalpy fuels and energy efficiency endothermic exothermic specific heat kinetic energy potential energy surroundings system General Outcome I Students will determine and interpret energy changes in chemical reactions. 1. given a knowledge of the formulas q = mcδt and q = nδh, calculate the heat gained or lost from a system 2. explain, in a general way, how stored energy in the chemical bonds of hydrocarbons originated from the sun
2 3. define enthalpy and molar enthalpy for chemical reactions 4. compare the molar enthalpies of several combustion reactions involving organic compounds 5. write and balance chemical equations for combustion reactions of alcanes (C n H 2n+2 ), including energy amounts 6. use an interpret H notation to communicate and calculate energy changes in chemical reactions 7. calculate and compare the energy involved in changes of state (physical change) and that in chemical reactions 8. predict the enthalpy change for chemical equations using standard enthalpies of formation 9. explain and use Hess law to calculate energy changes for a net reaction from a series of reactions 10. use calorimetry data to determine the enthalpy changes in chemical reactions 11. classify chemical reactions as endothermic or exothermic General Outcome II Students will explain and communicate energy changes in chemical reactions. 1. define activation energy as the energy barrier that must be overcome for a chemical reaction to occur 2. explain the energy changes that occur during chemical reactions, referring to bonds breaking and forming and changes in potential and kinetic energy 3. analyze and label potential energy diagrams of a chemical reaction, including reactants, products, enthalpy change and activation energy 4. explain that catalysts increase reaction rates by providing alternate pathways for changes, without affecting the net amount of energy involved; e.g., enzymes in living systems
3 Unit B: Solutions, Kinetics, and Equilibrium Solutions Overview: In this unit, students gain insight into the nature of matter through an investigation of change in the context of solutions. Focusing Questions: How is matter as solutions differentiated on the basis of theories, properties and scientific evidence? Why is an understanding of solution chemistry important in our daily lives and in the environment? Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in other courses. The intended level an scope of treatment is defined by the outcomes. homogeneous mixtures solubility electrolyte/nonelectrolyte concentration dilution net ionic equations spectator ions precipitation reaction stoichiometry solute solvent solution saturated unsaturated supersaturated General Outcome I Students will investigate solutions, describing their physical and chemical properties. 1. given a solution, demonstrate an understanding that it is a mixture formed by physically mixing at the particle level and that it does not involve a chemical change 2. explain dissolving as an endothermic or exothermic process with respect to the breaking and forming of bonds 3. explain, with the help of a diagram, the forces of attraction between solute and solvent particles 4. differentiate between electrolytes and nonelectrolytes
4 5. calculate, from empirical data, the concentration of solutions in moles per litre of solution and determine mass or volume from such concentrations 6. demonstrate a knowledge that the [ ] notation always implies concentration in mol/l 7. calculate the concentrations and/or volumes of diluted solutions and the quantities of a solution and water to use when diluting 8. use data and ionization/dissociation equations to calculate the concentration of ions in a solution 9. write balanced ionic and net ionic equations, including identification of spectator ions, for reactions taking place in aqueous solutions 10. use the solubility generalizations to predict the formation of precipitates 11. define solubility and identify related factors; i.e. temperature, pressure and miscibility 12. given solubility-temperature data, interpret solubility vs. temperature (solubility curves) 13. explain a saturated solution in terms of equilibrium; i.e. equal rates of dissolving and crystallization 14. describe the procedures and calculations required for preparing and diluting solutions Kinetics and Chemical Equilibrium Overview: In this unit, the concept that chemical change eventually attains equilibrium is developed. Focusing Questions: What is happening in a system at equilibrium? How do scientists predict shifts in the equilibrium of a system? Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in other courses. The intended level an scope of treatment is defined by the outcomes.
5 rate of reaction chemical equilibrium systems collision theory reversibility of reactions reaction mechanism Le Châtelier s principle equilibrium law expression General Outcome I Students will explain that there is a balance of opposing reactions in chemical equilibrium systems. 1. identify and discuss the properties and situations in which the rate of reaction is a factor 2. identify the factors (temperature, concentration, surface area, and catalysts) that affect rate of the reaction 3. describe collision theory and its connection to factors involved in alerting reaction rates 4. describe the role of the following in reaction rate: nature of reactants, surface area, temperature, and concentration 5. describe the role of a catalyst in a chemical reaction 6. draw and label a potential energy diagram to show the effect of a catalyst on the rate of the reaction 7. define equilibrium and state the criteria that apply to a chemical system at equilibrium; i.e. closed system, constancy of properties, equal rates of forward and reverse reactions 8. identify, write and interpret chemical equations for systems at equilibrium 9. predict, qualitatively, using Le Châtelier s principle, shifts in equilibrium caused by changes in temperature, pressure, volume, concentration or the addition of a catalyst and describe how these changes affect the equilibrium constant 10. explain how a catalyst and the surface area have an effect on the time it takes to reach equilibrium 11. write equilibrium constant (Kc) expressions with the knowledge that solids (s) and pure liquids (l) are not included in the expression
6 12. understand that the constant will vary with temperature 13. define Kc to predict the extent of the reaction and write equilibrium law expressions for given chemical equations, using lowest whole-number coefficients 14. predict whether or not reactant or products are favoured in a reversible reaction, on the basis of the magnitude of the equilibrium constant (Kc) and the reaction quotient (Qc). 15. solve Kc problems involving the initial concentrations, the changes that occur in each substance, and the resulting equilibrium concentrations. (I.C.E.) Unit C: Acids and Bases Overview: In this unit, students gain insight into the nature of matter through an investigation of change in the context of acids and bases. Focusing Questions: How is matter as acids and bases differentiated on the basis of theories, properties and scientific evidence? Why is an understanding of acid-base chemistry important in our daily lives and in the environment? Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in other courses. The intended level an scope of treatment is defined by the outcomes. strong acids and bases weak acids and bases monoprotic/polyprotic acid Arrhenius theory of acids and bases acid-base indicators hydronium ion/ph hydroxide ion/poh neutralization Bronsted-Lowry acids and bases titration curves conjugate pairs of acids and bases amphoteric substances equilibrium constants Kc, Kw, Ka, Kb buffers acid-base equilibrium General Outcome I Students will describe acidic and basic solutions qualitatively and quantitatively.
7 1. recall International Union of Pure and Applied Chemistry (IUPAC) nomenclature of acids and bases 2. recall the empirical definitions of acidic, basic and neutral solutions determined by using indicators, ph and electrical conductivity 3. calculate H 3 O + (aq) and OH - (aq) concentrations and the ph and poh of acidic and basic solutions based on logarithmic expressions; i.e., ph = -log [H 3 O + ] and poh = -log[oh - ] 4. compare magnitude changes in ph and poh with changes in concentration for acids and bases 5. explain how the use of acid-base indicators, ph paper or ph meters can be used to measure H 3 O + (aq) 6. define Arrhenius (modified) acids as substances that produce H 3 O + (aq) in aqueous solutions and recognize that the definition is limited 7. define Arrhenius (modified) bases as substances that produce OH - (aq) in aqueous solutions and recognize that the definition is limited 8. describe Bronsted-Lowry acids as proton donors and bases as proton acceptors 9. write Bronsted-Lowry equations, using indicators, and predict whether reactants or products are favoured for acid-base equilibrium reactions for monoprotic and polyprotic acids and bases 10. identify conjugate pairs and amphoteric substances 11. define neutralization as a reaction between hydronium and hydroxide ions 12. differentiate, qualitatively, between strong and weak acids and between strong and weak bases on the basis of ionization and dissociation; i.e. ph, reaction rate and electrical conductivity 13. identify monoprotic and polyprotic acids and bases and compare their ionization/ dissociation 14. draw and interpret titration curves, using data from titration experiments 15. describe the function and choice of indicators in titrations 16. identify equivalence points for various titration curves and differentiate between
8 the indicator endpoint and the equivalence point General Outcome II Students will determine quantitative relationships in simple acid-base equilibrium systems. 1. recall the concepts of ph and hydronium ion concentration and poh and hydroxide ion concentration, in relation to acids and bases 2. define Kw, Ka, Kb and use these to determine ph, poh, [H 3 O + ] and [OH - ] of acidic and basic solutions 3. calculate equilibrium constants and concentrations for homogeneous systems and Bronsted-Lowry acids and bases (excluding buffers) when concentrations at equilibrium are known initial concentrations and one equilibrium concentration are known (I.C.E.) the equilibrium constant and one equilibrium concentration are known Note: Examples that require the application of the quadratic equation are included. Unit D: Electrochemistry Overview: In this unit, students study electrochemical change and analyze the matter and energy changes within a system. Focusing Questions: What is an electrochemical change? How have scientific knowledge and technological innovation been integrated into the field of electrochemistry? Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in other courses. The intended level an scope of treatment is defined by the outcomes. oxidation spontaneity reduction standard reduction potential oxidizing agent voltaic (electrochemical) cell reducing agent electrolytic cell oxidation-reduction (REDOX) reaction electrolysis
9 oxidation number standard cell potential half-reaction corrosion General Outcome I Students will explain the nature of oxidation-reduction reactions. 1. define oxidation and reduction operationally and theoretically 2. define oxidizing agent, reducing agent, oxidation number and half-reaction 3. differentiate between REDOX reactions and other reactions, using half-reactions and/or oxidation numbers 4. identify electron transfer, oxidizing agents and reducing agents in REDOX reactions that occur in everyday life 5. compare the relative strengths of oxidizing and reducing agents, using empirical data 6. predict the spontaneity of a REDOX reaction, based on standard reduction potentials, and compare their predictions to experimental results 7. write and balance equations for REDOX reactions in acidic and neutral solutions by using half-reaction equations obtained from a standard reduction potential table General Outcome II Students will apply the principles of oxidation-reduction to electrochemical cells. 1. define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte, external circuit, power supply, voltaic (electrochemical) cell and electrolytic cell 2. identify the similarities and differences between the operation of a voltaic (electrochemical) cell and that of an electrolytic cell 3. predict and write the half-reaction equation that occurs at each electrode in an electrochemical cell
10 4. explain that the values of standard reduction potential are all relative to zero (0) volts, as set for the hydrogen electrode at standard conditions 5. calculate the standard cell potential for electrochemical cells 6. predict the spontaneity or nonspontaneity of redox reactions, based on standard cell potential, and the relative positions of half-reaction equations on a standard reduction potential table 7. explain the process of electrolysis and electroplating 8. explain how electrical energy is produced in a hydrogen fuel cell