macmillan chemistry pathways 2001 update · web viewanalyse data to explain the conditions under...

School: 1

MACMILLAN CHEMISTRY PATHWAYS 1 PRELIMINARY COURSE & 2 HSC COURSETEACHING PROGRAMS FOR YEARS 11 AND 12

INTRODUCTION

In the 18 months since the implementation of the New HSC Chemistry syllabus, some anomalies and problems remain. In their attempts to teach in the spirit of the document, teachers have found that there is insufficient time to cover all the dot points even when column 2 and column 3 are integrated. The syllabus requirement to teach chemistry in context can lead to vacuums of knowledge, which are required to be filled for further progress by students.

Teachers need to find ways of covering this extensive content and achieve the outcomes. The major stumbling block is in the Preliminary Course, where four modules have to be taught in three terms in order to begin the HSC course in term 4 of Year 11. In addition, the syllabus has not allowed time for consolidation and reflection to ensure that students master the concepts before proceeding to a new idea. There is also little time to integrate the mandatory perspectives which are required in a contextual program.

To assist teachers and students using the texts Macmillan Chemistry Pathways 1 and Macmillan Chemistry Pathways 2 by Geoffrey Thickett, the following programs have been developed. These programs have been trialed in several schools in term 1, 2001 and found to be most successful in achieving the majority of content outcomes as

well as balancing time constraints. Some dot point pruning has been required to achieve a program that can be taught in the time available.

These programs make reference to page numbers in Chemistry Pathways 1 & 2, and the content statements have been organised in such a way as to ensure the development of fundamental knowledge which will assist students in comprehending the contextual components of the syllabus. The modules ‘The Chemical Earth’ and ‘Metals’ have been integrated as there is considerable overlap of content. The modules ‘Water’ and ‘Energy’ have been similarly integrated, although to a lesser extent.

The Chemistry Pathways 1 & 2 series has delivered to teachers a large amount of additional information, to provide extension activities and enrichment for students, as well as materials to address the first- and second-hand experiences in column 3 of the syllabus. Reference to these data processing activities is made in the sample programs.

The teaching program is provided as a Microsoft Word 6 file to allow teachers to make their own modifications to it, and permission is granted to do so and to copy it into teachers’ own documents. The school name can be inserted in the left side of the header.

CONTENTSWe set files up to print on A4 paper, but the exact position of page breaks depends on your system and printer setup.

Check page breaks in Page Layout View or Print Preview before printing and adjust if necessary.

Preliminary Course Module 1: The Elements of the Earth 2Preliminary Course Module 2: Water and Energy 7HSC Course Module 1: The Identification and Production of Materials 11HSC Course Module 2: The Acidic Environment 15HSC Course Module 3: Chemical Monitoring and Management 19HSC Course Module: Option – Shipwrecks and Salvage 23

Teaching Programs for Macmillan Chemistry Pathways 1 & 2 by Geoffrey Thickett

Outcomes or other information quoted from syllabuses should be checked against the full syllabus document published by the NSW Board of Studies.The information in this document is not endorsed by the Board of Studies.

Teaching Programs for Macmillan Chemistry Pathways 1 & 2 by Geoffrey Thickett

School: 3

SENIOR PROGRAM PRELIMINARY COURSE 2 UNIT CHEMISTRY

MODULE 1: 8.2/8.3 THE ELEMENTS OF THE EARTH (14 WEEKS)Note: This module combines content and outcomes from ‘The Chemical Earth’ and ‘Metals’.

Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come

Week Content Compulsory ExperiencesMacmillan Chemistry

Pathways 1P2P3P11P12P13P14P15P16

1 Physical Separation of Mixtures1. Identify that mixtures exist in the lithosphere, hydrosphere, atmosphere and biosphere. Identify and describe how mixtures can be separated into their component pure substances by physical separation techniques such as filtration, evaporation and distillation.2. Recall that elements and compounds are pure substances that have fixed physical properties. Mixtures do not have fixed physical properties.3. Recall that the abundance of elements varies and identify the five most abundant elements in the universe, lithosphere, hydrosphere, atmosphere and biosphere.

P: Perform a gravimetric analysis of a mixture of sand, salt and water to give % composition.A: Report on an industrial separation process such as the fractional distillation of liquid air. (p15–16)P: Investigate physical properties of pure and impure substances – freezing point curve of a pure and impure sample of a molten organic compound (eg lauric acid).A: Plot data on the abundance of elements in nature (eg earth’s crust; sea water; cells).(p43–46, 48–49)

1. p5–6p13–16

2. p39–43

3. p43–46

P6P11P12P13P14P16

2–3 Atomic Structure and Bonding1. Recall that elements are composed of unique atoms and that atoms are composed of electrons, protons and neutrons.2. Describe qualitatively that electrons are arranged in shells (energy levels) around the nucleus. 3. Describe the formation of ions when atoms gain or lose electrons. Recall that ionisation energy is the energy required to remove an electron from a gaseous atom.4. Apply Lewis electron-dot diagrams to visualise ion formation and electron sharing in molecules.5. Recall that cations and anions combine in fixed ratios to form ionic compounds. Identify that the attraction between

A: Report on the historical development of ideas about atomic structure. (p155–160)A: Write chemical formulae using valency rules and name the compounds. (p56–57, 81–83) A: Plot data on ionisation energy versus atomic number to reveal trends; relate these trends to the structure of the atom. (p151–152)

1. p51–56

2. p57–60

3. p54–55p147–148

4/5. p60–62

School: Preliminary Course Module 1: The Elements of the Earth 4

Out-come


Pathways 1oppositely charged ions is called the ionic bond.6. Describe the formation of covalent molecular substances when non-metal atoms link to other non-metal atoms by a sharing of electron pairs called a covalent bond.7. Valency rules help us to construct chemical formulae for ionic and molecular compounds. Valency can be determined from a knowledge of electron configurations. 8. Apply systematic naming of inorganic compounds.

6. p63–65

7. p56

8. p62, p65

P10P11P12P13P14P16

4–5 Compounds and Chemical Extraction1. Explain that compounds can be separated into their component elements by chemical separation techniques (roasting, smelting, photolysis and electrolysis). Explain that energy is required for these processes. 2. Describe using historical contexts that metals such as copper were extracted from their ores using thermal decomposition (roasting) reactions or reactions with carbon (smelting). These techniques have been known for thousands of years.3. Identify that iron requires more energy to extract it from its ores than copper. Blast furnaces are required for the extraction of iron.4. Identify that aluminium requires considerable electrical energy to extract it from bauxite using the technique of electrolysis.5. Construct balanced chemical equations to describe the extraction of metals from their minerals.

P: thermal decomposition of MgCO3 with identification of CO2 (limewater); gravimetric determination of % loss in weight. (p92–93)P: Demonstration: observe (i) darkening of silver salts on exposure to light, (ii) electrolysis of water to show decomposition of a compound.P: Dissolve copper carbonate in sulfuric acid and recover the copper from solution by electrolysis. (p22–24, 130–132)A: Report the differences between the boiling of water and the electrolysis of water in terms of energy changes and particles. (p87)

1. p 75–80p84–87

2. p128–130 p187–190

3. p 191–194

4. p199–201

5. p139–140


Out-come


Pathways 1P11P12P13P14P16

6–7 Classifying the Elements1. Classify elements as metals, semi-metals and non-metals according to their physical properties. These properties determine their use.2. Outline the contributions of chemists such as Dobereiner and Mendeleev to the historical development of ideas about the classification of elements in the periodic table.3. Recognise patterns and trends to arrange elements in periods and groups according to increasing atomic number in the modern periodic table.4. Explain that the physical and chemical properties of the elements and their components show trends across periods and down groups.

A: Use second-hand data to plot physical data for various elements to show they can be classified into metals, semi-metals and non-metals and as solids, liquids and gases. (p183–185)P: Investigate the properties of a range of elements to show periodicity.A: Report on periodic trends (conductivity; hardness; atomic radius; m.p.; b.p.; valency; electronegativity). (p180–182)

1. p 36–43

2. p161–164

3. p169–173

4. p 180–185

P2 P3P6P13P14

8–9 Metals and their Chemical Properties1. Outline the uses that metals have and explain that these uses are related to their specific properties. Describe the commercial price of metals and relate this to their cost of production and abundance. Discuss metals as resources and why metals need to be recycled.2. Describe the use of common alloys . Explain that alloying leads to changes in properties. Explain that alloys find increasing use in modern society.3. Justify the criteria used to place metals in an order of activity of metals. These criteria include the rate at which they react with water, oxygen and dilute replacement acids. Balanced equations (including ionic equations) can be used to describe these reactions.4. Explain that the ease of thermal decomposition of metallic compounds is related to the activity series.5. Identify that some metals react with both acids and

P: Demonstration: observe the production of an alloy such as solder.A: Use second-hand data to investigate and compare the differences in properties of metals and alloys. (p124–127)A: Report on the chronology of the ages of metals and developments in metallurgy through to the modern era. (p128–130)P: Investigate the activity series of metals using reactions of metals and acids.A: Write balanced and ionic equations for reactions involving metals and non-metals; metals and oxygen; metals and acid; metals and alkalis. (p150)

A: Use half-equations to describe the formation of

1. p121–127p195–196

2. p124–127

3. p142–147

4. p144

5. p146–147


Out-come


Pathways 1alkalis to form ionic compounds in solution. 6. Recall that metals react with non-metals to form ionic compounds.

metal ions when metals dissolve in acids. (p150–151) 6. p150–151

P11P12P13P14P16

10–11 Relating Properties to Structure1. Identify that elements and compounds differ in their physical and chemical properties. Identify that the uses of materials is determined by these properties.2. Identify that crystalline solids can be classified as continuous lattices or molecular lattices. 3. Describe metallic, ionic and covalent network compounds as examples of continuous lattices. Their properties can be related to these lattice structures.4. Describe how covalent molecules form molecular lattices. These lattices have distinctly different properties from continuous lattices.

P/A: Compare properties of elements with properties of their compounds (eg Mg, O and MgO; Al, O and Al2O3). (p96–97)A: Data processing: analyse data on the bonding and properties of different crystal types.(p104–106) P/A: Construct models and draw diagrams of different lattices and relate bonding and structure to properties.

1. p95–104

2. p101

3. p101–102

4. p103–104

P2P10P14

12–14 Chemical Analysis1. Recall that mass is conserved in a chemical change. Describe how gravimetric analysis is a useful technique used by chemists to analyse the chemical composition of materials.2. Distinguish between empirical and molecular formulae. Show that balanced symbolic equations can be expressed in terms of empirical, molecular and structural formulae.3. Define the amount of a chemical substance by the ‘mole’ in terms of Avogadro’s number.4. Relate the mass (m) of a substance to the number of moles (n) and the molar weight (M) by the equation:n = m/M.

P: Determine mass ratios and the empirical formula of a simple compound using chemical analysis (eg. Mg/O for the combustion of magnesium). (p137)A: Solve problems using the equation: n = m/M (p217–222)A: Solve problems related to the volumes of gases produced by metal/acid reactions and relate this to the mole theory. (p223–224).

1. p76, p131–133

2. p210–214

3. p214–223

4. p223–225


Out-come


Pathways 15. Explain that chemists in their analysis of chemical composition use formula weight and percentage composition calculations.6. Explain that calculations involving the molar volumes of gases are useful in reactions involving gases.

5. p227 p231–232

6. p225–226

School: 8

SENIOR PROGRAM PRELIMINARY COURSE 2 UNIT CHEMISTRY

MODULE 2: 8.4/8.5 WATER AND ENERGY (14 WEEKS)Note: This module combines content and outcomes from ‘Water’ and ‘Energy’.

Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come


Pathways 1P2P6P7P12P13

1–2 Water and Intermolecular Forces1. Describe the attractive forces that exist between molecules. Recall that these attractive forces can be classified as dispersion forces, dipole-dipole attraction and hydrogen bonding.2. Recall that water is a unique molecule in terms of its structure and physical properties. Explain that the high electronegativity of oxygen causes water molecules to have an asymmetric charge distribution.3. Identify that water is a polar molecule. Compare the structure and polarity of water to other molecules such as ammonia, hydrogen sulfide and methane.4. Describe hydrogen bonding as an important intermolecular force. Explain that many of the physical properties of water can be related to the strong hydrogen bonding between its molecules.

P/A: Investigate and calculate the density of ice and liquid water. Explain the lower density of ice using a model of ice’s structure. (p 255, 271)P/A: Investigate the effect of impurities such as salt on the melting point/freezing point of water. A: Graph melting and boiling point data for water and other molecules of similar molar weight to show the effect of hydrogen bonding. (p268)P/A: Investigate/demonstrate and report on some properties of water including surface tension, viscosity and cohesion/adhesion. (p272–275, 280–281)

1. p258–259

2. p248p262–265

3. p264–265

4. p267–268 p270

P2P4P6P7P10P11P13P14

3–4 Solutions and Solubility1. Recall that water is an excellent solvent for many ionic and some molecular compounds. Describe the ion/dipole or dipole-dipole interactions involved. Covalent network solids (eg SiO2) and polymeric molecules (eg starch/cellulose) do not dissolve in water.2. Identify that some molecules (such as HCl and H2SO4) ionise when they dissolve in water. Other molecules (eg

P: Test the solubility of various solutes in water P: Prepare a solution to a known concentration and systematically dilute the solution using appropriate apparatus.A: Perform calculations involving concentration of solutions and dilutions. (p292)

1. p277–280

2. p287–290

School: Preliminary Course Module 2: Water And Energy 9

Out-come


Pathways 1glucose) remain unionised on dissolution. 3. Explain that solutes can dissolve in water to form dilute, concentrated or saturated solutions. Identify the dynamic and reversible nature of ion movement in a saturated solution.4. State that the concentration of a solution can be expressed in a variety of ways including mass/volume, volume/volume, moles per litre and parts per million.5. Describe the molarity (c) of a solution using the equation:

c = n/V6. Describe how a solution can be diluted to a specified concentration.

A: Report on the monitoring of heavy metal concentrations in waterways and in industrial waste discharges. (p291; 197–198)

3. p293–299

4. p283–286

5. p285–286

6. p285

P2P6P10P14

5–6 Ions and Precipitation1. Recall that strong electrolytes contain a high concentration of ions in solution. 2. Identify that mixing electrolyte solutions sometimes produces precipitates due to the insolubility of certain cation/anion combinations.3. Use solubility data tables to predict precipitation reactions. Compare the solubility of a variety of salts through precipitation reactions. 4. Write ionic equations and perform mass/volume calculations for precipitation reactions.

P: Compare the solubility of salts using a range of precipitation reactions.A: Perform calculations involving the concentrations and mass relationships in precipitation reactions. (p307–308)A: Write balanced equations using appropriate phase descriptors [(s); (l); (g); (aq)].(p297)

1. p288–290

2. p293–297

3. p296 p303–305

4. p297

P2P4P6P7P8P12

7–8 Chemical Energy and Reaction Rates1. Explain that energy changes are involved in chemical reactions. Some reactions are exothermic (release heat) and others are endothermic (absorb heat). Define the term ‘enthalpy’ and ‘enthalpy change’.2. Relate enthalpy changes to the bond-breaking and bond-forming processes during a reaction.

P: Investigate a range of exothermic and endothermic reactions. (p314–316)P: Investigate the effect of variables (concentration, temperature, surface area, catalysts) on the rate of a reaction. (p342–344)A: Report on the use of catalysts in industry

1. p309–311

2. p313–314


Out-come


Pathways 1P13 3. Describe the energy required to initiate a reaction as the

activation energy. Energy profile diagrams can be drawn for exothermic and endothermic reactions. 4. Relate the rate of a chemical change to the activation energy for that reaction. Increasing the temperature of a system leads to an increase in kinetic energy of the reactants and an increase in reaction rate. 5. Identify that reaction rates can be altered by changing the concentration or the surface area of solid reactants.6. Explain that catalysts can alter the rate of a reaction by changing the activation energy.

(p341–343) 3. p314 p379–380

4. p336–337

5. p337–339

6. p339–340

P2P4P7P8P9P10P14P16

9–11 Calorimetry and Combustion1. Define the term ‘specific heat’ (C) and compare the specific heat of water to a variety of other substances.2. Explain the technique of calorimetry for determining the energy changes in a reaction.3. Explain and use the calorimetry equations:

Dh = –mCDT; DH = plus or minus h/n4. Identify combustion as an exothermic chemical reaction. Describe combustion reactions as slow, spontaneous and explosive.

5. Identify that carbon compounds can undergo combustion to produce a variety of products. Combustion of carbon compounds can be described as complete or incomplete.Incomplete combustion leads to atmospheric pollution.6. Explain the relationship between ignition temperature and activation energy for a combustion reaction.7. Identify the need for a wick in some combustion reactions involving hydrocarbons.

A: Report on water’s ability to absorb heat and how this is important to aquatic systems on earth; Discuss problems of thermal pollution. (p247–9, 316–317)P: Use calorimetry techniques to measure the heat of dissolution of ammonium chloride and sodium hydroxide in water. (p318–319)P: Use calorimetry techniques to measure the heat of combustion of kerosene. Discuss the errors involved due to heat losses.P: Demonstration: observe the combustion of a candle and compare with the combustion of kerosene and natural gas in a Bunsen burner. (p385)A: Report on conditions in which explosive reactions occur and the need for safety in work environments where fine particles mix with air. (p389)

1. p311–314

2. p314–316

3. p314–316

4. p389–390

5. p383–385

p389–390

6. p381–383

7. p385


Out-come


Pathways 1

P3P4P7P9P10

12–14 Carbon Compounds and Fuels1. Identify the position of carbon in the periodic table and describe its electron configuration.2. Describe the structure and properties of the allotropes of carbon.3. Identify that carbon can form single, soluble and triple bonds. Carbon has a unique ability to bond with itself and form a huge variety of compounds.4. Define the terms ‘homologous series’ and ‘functional group’. Systematically name alkanes, alkenes and alkynes and compare and contrast their properties.5. Relate the melting point, boiling point and volatility of hydrocarbons to their non-polar nature and dispersion forces between their molecules.6. Describe the geological origin of fuels such as coal, oil and natural gas. Explain how the chemical energy stored in these fuels is the result of energy transformations involving photosynthesis in prehistoric environments. 7. Describe the process of fractional distillation and the properties of the fractions produced at an oil refinery.

P: Model the structures of carbon allotropes and relate their properties to their structure. (p346)P: Model simple hydrocarbon molecules and use IUPAC rules to name these molecules. (p350)P: Use quick-fit apparatus to model distillation using a mixture of ethanol and water. (p363–4)A: Report on the sources of hydrocarbon fuels in Australia. Identify the range of compounds that can be obtained from fossil fuels. (p368–376)Identify safety issues associated with the storage and use of hydrocarbons. (p360, 381–3)

1. p346

2. p347–348

3. p348–355

4. p357–362

5. p357–360

6. p368–372

7. p355–357

School: 12

SENIOR PROGRAM HSC COURSE 2 UNIT CHEMISTRY

HSC MODULE 1: 9.2 THE IDENTIFICATION AND PRODUCTION OF MATERIALS (9 WEEKS)Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come


Pathways 2H6 H7H9H13 H14

1 Hydrocarbons and Carbon Nomenclature1. Identify that carbon can form single, double and triple covalent bonds.2. Use IUPAC nomenclature to name alkanes, alkenes and alkynes and their halogenated derivatives.3. Draw structural formulae for hydrocarbons and their simple halogenated derivatives.4. Identify the non-polar nature of hydrocarbons and the weak dispersion forces between their molecules that help us to explain the physical properties of alkanes and alkenes.

A: Construct and name some models of hydrocarbons.

1. p6–7(+ Book 1 p348–349)2. Book 1 p350–352Book 2 p113/4. p6–7

H4H5H6H7H8H9H10H11H12H13H14H15H16

2–5 Ethene, Polymers and Biofuels1. Identify how ethene is prepared industrially by the process of catalytic cracking of petroleum fractions.2. Identify that catalytic cracking involves surface reactions on an inorganic catalyst such as a zeolite (aluminium silicate).3. Relate the reactivity of ethene to the ease of addition reactions at the double bond. 4. Identify ethene as an example of a monomer that can undergo addition polymerisation. Identify the steps in the commercial manufacture of polyethene. 5. Identify the structures of monomers such as vinyl chloride, acrylonitrile and styrene and name them systematically. Name and draw structures for the polymers formed from these monomers and indicate uses of these

P: Addition of bromine water across double bonds. (p12)A: Write balanced structural equations for a range of reactions related to this topic.P/A: Use model kits and Internet programs to simulate the formation of polymers from monomers and to visualise the coiling of polymer strands.P: Microbial decay of grass clippings (cellulose) produces heat and other by-products.A: Demonstration of the identification of cellulose.A: Analyse secondary sources to determine the progress in the development of a biopolymer such

1. p7–11

2. p10

3. p10–13

4. p14–16

5/6. p16–19

School: HSC Course Module 1: The Identification and Production Of Materials 13

Out-come


Pathways 2polymers.6. Identify that some polymers are formed by the process of condensation polymerisation. Write structural equations to show the process of condensation polymerisation and name some common condensation polymers (PET and nylon).7. Describe cellulose as a biopolymer and a natural condensation polymer formed from glucose monomers. Draw a structure for a section of the cellulose chain.8. Identify cellulose as a major component of biomass that has the basic carbon chain structure to allow the manufacture of alternatives to petrochemicals.9. Assess current developments in the use of biopolymers including the production of energy and biodegradable plastics.10. Identify ethanol as a biofuel which can be produced by yeast fermentation of sugars and summarise the fermentation process.11. Explain that ethanol can be produced commercially from ethene by the acid catalysed addition of water. The ethanol formed can be dehydrated to produce ethanol in the presence of concentrated sulfuric acid.12. Describe ethanol as a useful solvent and explain that its physical properties are related to its polar nature and its ability to hydrogen bond.13. Discuss the advantages and disadvantages of ethanol as an alternative biofuel and renewable resource.14. Investigate the combustion of ethanol and determine a value for the molar heat of combustion.

as PLA. (p33)A: Use model kits to simulate the addition of water to ethene and the dehydration of ethanol.P: Investigate the solubility of solutes in ethanol.A: Use secondary sources to summarise the process of sugar fermentation in industry and to write a balanced equation for the process. (p35)P: Measure mass changes and CO2 production during the yeast fermentation of sucrose.A: Use secondary sources to discuss ethanol as an alternative fuel. (p35)P: Investigate the heat of combustion of three alkanols. (p39)

6. p24–26

7. p26–29

8. p29–30

9. p31–35

10. p35

11. p35–36

12. p36–37

13. p38–39

14. p39–40


Out-come


Pathways 2H4H5H6H7H8H11H12H13H14H15H16

6–7 Electrochemical Energy1. Describe oxidation-reduction in terms of electron transfer and changes in oxidation state.2. Identify the relationship between metal displacement reactions and the activity series of metals. Relate these metal displacement reactions to electron transfer half-equations.3. Outlines the construction of galvanic cells and identify the terminology used in describing these cells.4. Describe and explain galvanic cells in terms of electron transfer reactions.5. Solve problems using a table of standard reduction potentials. 6. Discuss the structure, chemistry and uses of a dry cell, lead-acid cell and compare them with other types of batteries such as button cells or the vanadium redox cell.7. Describe the use of electrolysis in industry to electro-refine copper. Identify the half-reactions and the conditions required. 8. Identify an example of electroplating and explain the processes involved.

P: Investigate the construction of simple galvanic cells and to measure the cell potentials. (p72)

A: Evaluate the chemistry, cost and environmental impact of the use of dry cells, lead-acid cells and compare them to button cells and vanadium redox cells. (p63–70)

A: Use tables of standard reduction potentials to solve problems and analyse information about named electrochemical processes. (p61–63)

1. p43–49

2. p51–54

3. p55–56

4. p56–59

5. p60–63

6. p63–70

7. p71

8. p72

H4H5H6H7H8H13H14H16

8–9 Nuclear Chemistry1. Identify instruments and processes that detect radiation and explain how the instrument or process works.2. Describe how transuranic elements and commercial radioisotopes are produced in nuclear reactors and accelerators.3. Identify uses of named radioisotopes in industry, medicine and in determining reaction mechanisms. Relate their use to the chemical properties of the radioisotope.4. Distinguish between stable and radioactive isotopes and

A: Process second-hand data on the recent discovery of elements such as Z = 118. (p83)

A: Process second-hand data to describe the benefits and problems associated with the use of radioactive isotopes in an identified industry and in medicine. (p85)

A: Trace the sequence of fission products in the

1. p77–80

2. p80–84

3. p85–90

4. p90–91


Out-come


Pathways 2describe the conditions under which a nucleus is unstable in terms of its neutron to proton ratio.5. Use flow charts to trace the sequence of nuclear events when uranium decays.

decay series for natural uranium. (p91)

5. p91–92

School: 16


HSC MODULE 2: 9.3 THE ACIDIC ENVIRONMENT (9 WEEKS)Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come


Pathways 2H1 H2 H6H7 H8H9 H10 H11H12H13H14

1–3 Acids, Bases and Equilibrium1. Recall that chemical reactions are reversible. Identify factors that can affect the equilibrium in a reversible reaction.2. Define Le Chatelier’s principle and relate it to changes in equilibrium systems.3. Classify common substances as acidic, basic or neutral using indicators.4. Identify that various indicators (litmus, methyl orange, phenolphthalein and bromothymol blue) can be used to determine acidity/basicity over a range based on colour changes.5. Outline the historical development of ideas about acids including those of Lavoisier, Davy and Arrhenius.6. Outline the Bronsted-Lowry theory of acids and bases.7. Describe the relationship between an acid and its conjugate base and a base and its conjugate acid. Identify conjugate acid/base pairs.8. Identify a range of salts which form acidic, basic or neutral solutions and explain this behaviour in terms of the Bronsted-Lowry theory.9. Identify amphiprotic substances and write equations to describe their behaviour in acidic and basic solutions.10. Outline the Lewis definition of an acid.11. Assess the importance of each definition of an acid in terms of understanding.12. Define acids as proton donors and describe the ionisation

P: Prepare and test a natural indicator (eg from red cabbage or petals).A: Gather second-hand data about the colour changes of indicators. (p122,148)P: Use indicators to test the acidity/basicity of soil and other household substances.

A: Use second-hand data to trace the development of our understanding of acid/base reactions and explain the expanded view of acids developed by Lewis. (p124–133)

A: Write ionic equations to represent the ionisation of acids (p133)

1. p110–115

2. p115–119

3. p122

4. p122–123

5. p124–127

6. p127–1287. p129–130

8. p130

9. p130–131

10. p13111. p124–131

12. p133

School: HSC Course Module 2: The Acidic Environment 17

Out-come


Pathways 2of acids in water.13. Describe acids and their solutions with appropriate use of the terms strong, weak, concentrated and dilute.14. Compare the relative strengths of equal concentration of citric, acetic and hydrochloric acids and relate this to the degree of ionisation of their molecules.15. Describe the difference between a strong and a weak acid in terms of equilibrium between the molecules and its ions.

P: Use molecular models to model the ionisation of a strong and weak acid. 13. p135–136

14. p134–138

15. p136

H1H2H7H8H9H10H11H13H14H16

4–6 Acids and the Environment1. Describe the ionisation of water and define pH as –log10[H+] and explain that a change in pH of 1 unit means a tenfold change in [H+].2. Describe the use of the pH scale in comparing the concentration of acids and alkalis.3. Relate indicator colour changes to pH.4. Explain the formation and effects of acid rain.5. Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids.6. Identify natural and industrial sources of sulfur dioxide and sulfur trioxide.7. Describe using equations examples of chemical reactions which release sulfur dioxide and reactions which release oxides of nitrogen.8. Analyse the periodic table to locate non-metals that form acidic oxides and generalise about the relationships between position of the elements in the periodic table and the acidity/basicity of the oxides.9. Assess the evidence that indicates increases in atmospheric concentrations of the oxides of sulfur and

P: Use pH meters to distinguish between acidic, basic and neutral chemicals including the pH of a range of salt solutions.P: Use a pH meter to measure the pH of strong and weak acids of the same molarity. (p147)A: Gather second-hand data on the uses of acidic oxides such as SO2 as food additives. (p150;159)A: Process second-hand data on the properties and origin of sulfur dioxide and oxides of nitrogen. Discuss environmental concerns. (p158–160)

A: Gather and process second-hand data to identify examples of naturally occurring acids and

1. p144–145

2. p146–148

3. p148–1494. p150–1515. p150

6. p150

7. p150

8. p155–158

9. p153


Out-come


Pathways 2nitrogen.10. Calculate volumes and masses of gases at STP and SLC.11. Identify naturally occurring acids such as acetic acid, citric acid and hydrochloric acid. Identify sulfuric acid and hydrobromic acid as manufactured acids.12. Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and relate this to Le Chatelier’s principle (eg limestone caves and soft-drinks).13. Qualitatively describe the effects of buffers with reference to a specific example in a natural system (eg blood; stomach acidity).

bases including their chemical composition and pH.(p161–170)P: Measure the mass change in the decarbonation of a soft drink and calculate the volume of gas released at SLC. A: Process information from secondary sources to visualise the rearrangement of particles and changes in electrical conductivity and pH which occurs during a neutralisation reaction.(p170)

10. p159–16111. p161–164

12. p165

13. p167–170

H2H6H7H8H9H10H11H12H13H14H15

7–9 Volumetric Analysis and Esterification1. Describe the correct technique for conducting titrations and preparation of standard solutions. Explain the need for accuracy during a titration.2. Identify neutralisation as a proton transfer reaction that is exothermic.3. Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds.4. Explain the difference in melting points and boiling point caused by the alkanoic acid and alkanol functional groups.5. Identify esterification as the reaction between an alkanoic acid and an alkanol and describe (using equations) examples of esterification.6. Describe the purpose of using concentrated sulfuric acid in esterification for catalysis and absorption of water.7. Explain the need for refluxing during esterification.8. Outline some examples of the occurrence, production and

A: Analyse information from secondary sources to predict combing volumes and masses of acids and bases and products in neutralisation reactions. P: Prepare a standard solution (eg sodium carbonate). (p178)P: Use titration techniques to solve problems (eg to determine the concentration of acetic acid in vinegar). (p180–182, 193)P: Measure temperature changes during neutralisation and use calorimetry data to calculate the molar heat of neutralisation. (p184, 194)A: Analyse secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills.

P: Prepare an ester using reflux. (p190, 196)

1. p175–182

2. p184–186

3. p21–23

4. p188–189

5. p190–191

6. p191

7. p1918. p192


Out-come


Pathways 2uses of esters. A: Use secondary sources to identify the uses of

esters in flavours, perfumes, cosmetics and in processed foods. (p192)

School: 20


HSC MODULE 3: 9.4 CHEMICAL MONITORING AND MANAGEMENT (8 WEEKS)Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come


Pathways 2H3H7H8H9H10

1–2 Chemical Processes in Industry1. Outline the role of an industrial chemist employed in a power station. Describe the principles of their research and the plant-monitoring processes for which they would be responsible.2. Compare complete and incomplete combustion and describe the role of an industrial chemist in monitoring combustion reactions in industry.3. Identify that ammonia can be synthesised industrially from nitrogen and hydrogen in the Haber process.4. Describe the synthesis of ammonia as a reversible exothermic reaction that will reach equilibrium.5. Explain why the rate of ammonia synthesis is increased by high temperatures and use Le Chatelier’s principle to explain why the yield of ammonia is reduced at higher temperatures. Thus, explain why the Haber process is a delicate balancing act involving reaction energy, reaction rate and equilibrium.6. Explain that the use of a suitable catalyst will allow lower temperatures to be used in the Haber process. Name the catalysts used. 7. Use Le Chatelier’s principle to analyse the impact of increased pressure on the system in the Haber process.8. Explain why monitoring the reaction vessel used in the Haber process is crucial and discuss the monitoring required.

A: Report on the variety of chemical occupations and describe the role of an industrial chemist in detail. (p214–218)

A: Report on the historical development of the Haber process and its significance in world history. (p219–221)

A: Solve problems involving equilibrium constant calculations of the effect of volume, temperature and concentration changes on the yield of ammonia. (p223–226)

1. p213–214

2. p215–218

3. p219–221

4. p219

5. p219–220

6. p219–220

7. p220

8. p221

School: HSC Course Module 3: Chemical Monitoring and Management 21

Out-come


Pathways 2H2H3H10H11H12H13H14H15

3–4 Analytical Chemistry1. Describe chemical tests to identify the following anions and cations: ∑ Anions: phosphate; sulfate; carbonate; chloride∑ Cations: barium, calcium, lead, copper, iron

2. Deduce the ions present in unknown samples as a result of chemical tests.3. Describe how atomic absorption spectrophotometry (AAS) is used to detect the concentration of metal ions in solution and assess its impact on the scientific understanding of the effects of trace elements.4. Devise a procedure that could be used to determine the water content, alcohol content and aspirin content of a medicine.

P: Practically identify the following ions using a range of tests including flame tests:∑ Anions: phosphate; sulfate; carbonate;

chloride∑ Cations: barium, calcium, lead, copper, iron

A: Report on the need to monitor the levels of lead and mercury in substances used by society and interpret supplied AAS data to evaluate the effectiveness of pollution control.

P: Experimentally determine:a. The phosphate content in detergent (p242)b. The nitrogen and sulfate content in a fertiliser (p240; 245)c. The ethanol content of wine (p249) d. The citric acid content of orange juice (p250)

For one of these experiments, discuss the errors associated with the determination.

1. p227–231

2. p232

3. p234–236

4. p236–240

H2H3H6H7H8H9H10

5–6 Atmospheric Chemistry 1. Describe the composition and layered structure of the atmosphere.2. Identify the main pollutants found in the atmosphere and their sources. 3. Use electron dot structures to show the formation of coordinate bonds in molecules such as carbon monoxide and

A: Solve problems to compare the amounts of atmospheric components by moles, mass, volume, ppm. (p252)

A: Write equations to show the reactions involving CFCs and ozone. (p262)

1. p251–253

2. p253–255

3. p256–257


Out-come


Pathways 2ozone. 4. Compare the properties in terms of bonding and structure of the oxygen allotropes, O2 and O3.5. Compare the properties of O2, O2- and O radicals.6. Explain the effects of UV light on living material including natural polymers. Explain the importance of the ozone layer in filtering out harmful UV radiation.7. Discuss how the evidence for the thinning of the ozone layer by halogens, halons and CFCs was obtained. Identify the origins of these pollutants and write balanced equations to show the key reactions involved. 8. Identify, draw and systematically name straight chain haloalkanes, haloalkenes and their isomers. 9. Assess the steps that are being taken to alleviate the problems of ozone destruction in the stratosphere.

P: Use model kits to construct models of haloalkanes and their isomers. (p267)

A: Report on the effect of UV light on biological molecules. Use diagrams as part of this report.

A: Use the Internet or other resources to report on replacements for CFCs and evaluate their effectiveness. (p266)

4. p257–259

5. p2586. p260–261

7. p261–265

8. p266–268

9. p264

H3H8H10H11H12H13H14H15H16

7–8 Managing Water Quality1. Identify and describe tests that can be used to determine the quality of water samples including:

a. Determining the concentration of ions (eg AS/ gravimetric)b. Total dissolved solids (eg evaporation/ gravimetric or electrical conductivity)c. Hardness (eg titration)d. Turbidity (eg turbidimetry)e. Acidity (eg pH electrodes)f. Dissolved oxygen (eg titration or oxygen electrode)g. Biochemical oxygen Demand (BOD), (eg titration)h. Nitrogen: phosphorus ratio (titration/colorimetry)

2. Identify factors (eg rain and its pH, floods, evaporation, weathering, discharge of effluents from industry, human

P: Practically analyse and compare the water quality of different water samples.

A: Report on the range and chemistry of the tests used to:∑ Identify heavy metal pollution in water∑ Monitor possible eutrophication in waterways

(p286)

P: Investigate the amount of solids in water samples from the local environment. (p281)

A: Report on the following features of the

1. p275–288

2. p293–296


Out-come


Pathways 2activity) that affect the concentration of various ions in solution in natural bodies of water such as rivers/oceans. 3. Assess the effectiveness of determining total dissolved solids in water samples by evaporation and gravimetric analysis versus electrical conductivity measurements (assuming that most dissolved solids are ionic). 4. Describe and assess the effectiveness of methods used to purify and sanitise mass water supplies.5. Describe the design and composition of microscopic membrane filters (eg NF and UF) and explain how they purify contaminated water.6. Identify the need for collaboration between chemists as they collect and analyse data.

Sydney water supply:∑ Catchment area∑ Possible sources of contamination in this

catchment∑ Chemical tests used to determine levels and

types of pollutants∑ Physical and chemical processes used to purify

the water∑ Chemical additives in the water and the reason

for their presence (p288–292)

3. p278

4. p288–292

5. p292–293

6. p294

School: 24


HSC MODULE: OPTION – SHIPWRECKS AND SALVAGE (8 WEEKS)Key to compulsory experiences – P: First-hand practical activity; A: Second-hand data processing/collection activity.

Out-come


Pathways 2H3H7H8

1 Electrolytes and Electron Transfer1. Identify the origins of minerals in the ocean as:

a. leaching by rainwater from terrestrial environmentsb. hydrothermal vents in mid-ocean ridges

2. Outline the role of electron transfer in oxidation-reduction reactions.3. Identify that oxidation-reduction reactions can occur when ions are free to move in solid and liquid electrolytes. 4. Describe the work of Galvani, Volta, Davy and Faraday in increasing our understanding of electron transfer reactions.

A: Report on the impact of the work of Galvani, Volta, Davy and Faraday on our understanding of electron transfer reactions.(p389–392)

1. p384–386

2. p389

3. p387–388

4. p389–392

H7H8H10H11H12H13H14

2–3 Galvanic and Electrolytic Cells1. Define the terms ‘galvanic cell’ and ‘electrolytic cell’ and distinguish between each using various examples. 2. Describe, using half-equations, what happens at the anode and cathode during the electrolysis of selected aqueous solutions.3. Describe factors that affect an electrolysis reaction:

a. Effect of electrode materialb. Effect of different electrolytesc. Effect of changes in concentration

4. Define Faraday’s first law of electrolysis.

P: Construct galvanic cells between pairs of metals to determine the difference in reactivity between the two metals.A: Calculate the cell potential given data on the two half-equations and half-cell potentials(p393)A: Represent galvanic cells using shorthand notation. (p392) P: Electrolyse a variety of solutions and investigate the reactions at each electrode and the factors that affect the rate of the electrolysis reaction. (p394–402, 407–8)

1. p392–294

2. p394–402

3. p394–398p398–402p402

4. p403–405

School: HSC Course Module: Option – Shipwrecks and Salvage 25

Out-come


Pathways 2

H3H5H6H7H8H10H11H12H13H14H15H16

4–5 Corrosion of Metals1. Analyse data to explain the conditions under which the rusting of iron occurs and explain the process of rusting. 2. Use a table of standard potentials to predict which metal corrodes when two metals form a galvanic cell. 3. Account for the differences in corrosion of active and passivating metals.4. Identify the composition of steel and explain how the percentage composition of steel can determine its properties.5. Identify iron and steel as the main metals used in ships.6. Outline the process of cathodic protection, describing examples of its use in both marine and wet terrestrial environments.7. Use gathered data to describe the process of cathodic protection in selected examples in terms of the oxidation/reduction chemistry involved. 8. Identify the ways in which a metal hull may be protected including:

a. Corrosion-resistant metalsb. Development of surface alloysc. New paints

P: Experimentally investigate the rate of corrosion of pure iron, an identified form of steel and a variety of other metals and modern alloys. Identify the metals that would be best suited for use in marine vessels. (p417)

A: Report on the composition, properties and uses of a range of steels. (p417)

A: Report on the historical developments in the design and construction of ocean-going vessels with a focus on the metals used. (p426)

P: Compare the effectiveness in using different metal coatings to protect iron from rusting. (p423)

1. p414–415

2. p60–63

3. p415–416

4. p416–417

5. p418 p4226. p418–420

7. p423–424

8. p424p425–426p418, p425

H5H6H7H8H10H11H12H13

6–7 Corrosion of Shipwrecks1. Outline the effect of temperature and pressure on the solubility of gases and salts in water.2. Identify the gases that are normally dissolved in the oceans and compare their concentrations in the oceans to their concentration in the atmosphere.3. Compare and explain the solubility of selected gases at increasing depth in the oceans.

P: Experimentally investigate the corrosion of materials in different environments including:∑ High and low oxygen levels∑ High and low temperatures∑ High and low salt concentrations∑ High and neutral solutions (p414)

1. p428

2. p429

3. p430

School: HSC Course Module: Option – Shipwrecks and Salvage 26

Out-come


Pathways 2H14H15

4. Predict the effect of low temperature at great depths on the rate of corrosion of a metal such as iron. Give reasons for the prediction made. 5. Explain that shipwrecks at great depth are not corroded by electrochemical reactions but rather by anaerobic bacteria.6. Describe the anaerobic bacteria as sulfur-reducing species whose wastes produce acidic environments around deep wrecks.7. Explain that acidic environments accelerate corrosion in non-passivating metals.

4. p428

5. p430

6. p431

7. p431–432

H3H5H6H7H8H11H12H13H14H16

8 Salvage and Restoration1. Explain that artefacts from long-submerged wrecks will be saturated with dissolved chlorides and sulfates.2. Describe the processes that occur when a saturated solution evaporates and relate this to the potential damage to drying artefacts.3. Identify the use of electrolysis as a means of cleaning and stabilising strong metal artefacts.4. Discuss the range of chemical procedures which can be used to clean, preserve and stabilise artefacts from wrecks and provide an example of the use of each procedure.

P: Experimentally model the evaporation of water from salt saturated wood, leather and textiles and describe the change in the material as drying occurs. (p435)

A: Use Internet or text resources to identify appropriate electrolytic treatment methods to clean and remove the salt from chloride or sulfate saturated artefacts. (p432–441)

1. p432

2. p433

3. p433

4. p433–441

macmillan chemistry pathways 2001 update · web viewanalyse data to explain the conditions under...

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