School Science Lessons
2012-02-06
Please send comments to: J.Elfick@uq.edu.au
Edited extracts from Chemistry Senior Syllabus
The State of Queensland (Queensland Studies Authority) 2007
This syllabus is approved for general implementation until 2014, unless otherwise stated.
Queensland Studies Authority, PO Box 307, Spring Hill, Queensland, Australia 4004
Phone: (07) 3864 0299
Fax: (07) 3221 2553
Email: office@qsa.qld.edu.au
Website: www.qsa.qld.edu.au
Table of contents
1.0 Chemical structure
2.0 Chemical reactions
3.0 Sample courses of study - Unit titles
Unit 1. Water quality chemistry
Unit 2. Forensic chemistry
Unit 3. Petroleum chemistry
Unit 4. Cold packs chemistry
Unit 5. Ocean chemistry
Unit 6. Carbon dioxide in drinks chemistry
Unit 7. Flavours chemistry
Unit 8. Sugars in food chemistry
Unit 9. Alcohol in beverages chemistry
Unit 10. Acids in drinks chemistry
Unit 12. Fuels for transport chemistry
Unit 13. Motor car fuel chemistry
Unit 14. Rocket fuel chemistry
Unit 15. Fertilizers and explosives chemistry
1.0 Chemical structure
S1.0 All matter is composed of atoms.
S1.1 Matter is composed of atoms which, in turn, contain protons and neutrons in a nucleus, and electrons outside the nucleus.
S1.2 The number of positively charged protons is equal to the number of negatively charged electrons in a neutral atom, and determines all the chemical properties of an atom.
S1.3 An element is a substance in which all atoms have the same number of protons.
S1.4 Atoms of an element may contain different numbers of neutrons, and are known as isotopes.
S1.4.1 mass number
S1.4.2 radioactivity
S1.5 Every element is assigned a unique chemical symbol.
S1.6 The atomic mass of an atom is arbitrarily defined relative to the mass of the isotope carbon-12.
S1.7 In modern theories of atomic structure, electrons are viewed as occupying orbitals which are grouped in electron shells.
Electron configuration
S2.0 Materials can be categorized and represented symbolically and their macroscopic properties can be explained and predicted from understandings about electronic structure and bonding.
S2.1 From theory of electronic structure it is predicted that elements will display periodic variations in their chemical and physical properties, the trends across a period or down a group in the periodic table for properties such as melting or boiling point, reactivity, ionization energy, atomic radius, metallic character, nature of oxides
S2.1.1 terms used to describe groups and periods of the periodic table; alkali metals, alkali earth metals, halogens, noble gases, lanthanides and actinides
S2.1.2 the relationship between the number of valence electrons for an element, its position in the periodic table, and its chemical properties
S2.1.3 properties of an element (e.g. combining power, general reactivity) and relationship to its position in the periodic table
S2.1.4 anomalies in the properties of an element
S2.2 The macroscopic properties are related to their microscopic properties.
S2.2.1 classification of materials in appropriate bonding categories
S2.2.2 common macroscopic properties
S2.2.3 comparison of models of bonding in metallic, ionic, covalent molecular and covalent network substances
S2.2.4 the properties of solids, liquids and gases using the kinetic particle theory and relating the theory to phase changes
S2.2.5 the physical properties of different types of materials
S2.3 Pairs of atoms may be bound together by the sharing of electrons between them in a covalent bond.
S2.3.1 single and multiple covalent bonds
S2.4 Two or more atoms bound together by one or more covalent bonds form a molecule, with definite size, shape and arrangement of bonds.
S2.4.1 how a dipole arises with reference to electronegativity, polar bonds and the effect of molecular shape
S2.4.2 polar and non-polar covalent bonds and molecules
S2.4.3 electron dot diagrams and Lewis valence structures for simple inorganic and organic molecules
S2.4.4 shapes of simple covalent molecules
S2.4.5 VSEPR theory to predict molecular shape
S2.5 An atom or group of atoms covalently bound together may gain or lose one or more electrons to form ions.
S2.6 Ionic bonding occurs when positive and negative ions are held together in a crystal lattice by electrostatic forces.
S2.7 When chemical bonds, whether ionic or covalent, are formed between different elements, a chemical compound is obtained, which can be represented by a chemical formula.
S2.7.1 chemical formulas' interpretation
S2.7.2 naming a molecular compound given its formula and vice versa
S2.7.3 anions and cations and the symbols/formulas and charges on those designated to be learnt
S2.7.4 formulas' deduction for ionic substances
S2.8 Forces weaker than covalent bonding exist between molecules.
S2.8.1 Van der Waal's dispersion forces, dipole-dipole forces, hydrogen bonding and the factors affecting their strength
S2.8.2 the properties of polar and non-polar compounds and models of intermolecular bonding to explain these properties
S2.8.3 molar heats of fusion and vaporization, specific heat capacity, melting point, boiling point, vapour pressure and surface tension, and the relationship these physical properties have with the strength of intermolecular forces
S2.9 The structure of a metal involves positive ions embedded in a sea of electrons.
S2.9.1 the properties of metals (thermal, conductivity, electrical conductivity, lustre, physical state, ductility, malleability) and relationships to structure
S2.9.2 ways in which metals can be modified and the effects of this on their properties
S2.9.3 alloys
S2.10 Materials may be elements, compounds or mixtures.
S2.10.1 elements, mixtures and compounds can be differentiated experimentally
S2.10.2 operation techniques of different types of mixtures
S2.10.3 composition of pure substances and mixtures
S2.10.4 mixtures can be liquid or non-liquid
S2.11 In compounds containing carbon- hydrogen bonds (known as organic compounds), the carbon atoms bind to one another through single, double or triple covalent bonds to form chains or rings.
S2.11.1 the general formulas for alkanes, alkenes, alkynes, alkanols, carboxylic acids, esters, amines, aldehydes and ketones
S2.11.2 IUPAC rules to name alkanes, alkenes, alkynes and simple alkanols, carboxylic acids, carboxylic acids, esters, amines, aldehydes and ketones
S2.11.3 structural isomerism, geometric isomerism, functional groups, homologous series, saturated and unsaturated, substitution, addition, elimination
S2.11.4 simple physical properties of alkanes, alkenes, alkanols, acids and relation of these properties to structure
S2.11.5 simple chemical properties of alkanes, alkenes, alkanols, acids and esters
S2.11.6 structural features and simple chemistry of some biochemical molecules (e.g. amino acids, proteins, fats, carbohydrates, nucleic acids)
S2.11.7 primary, secondary and tertiary alkanols
S2.11.8 the main products of several chemical reactions, such as substitution, addition, and oxidation, of a selection of simple organic compounds
S2.11.9 equations for the formation:
S2.11.9.1 of alkanes, alkenes and alkynes from each other
S2.11.9.2 and oxidation of alkanols
S2.11.9.3 of alkyl halides (halocarbons)
S2.11.9.4 of esters
S2.11.10 the nature of single and multiple covalent bonds in organic molecules
2.0 Chemical reactions
R1.0 Specific criteria can be used to classify chemical reactions.
R1.1 Redox reactions involve a transfer of electrons and a change in oxidation number.
R1.1.1 electron transfer
R1.1.2 oxidation (including rules for assignment)
R1.1.3 oxidation and reduction (redox)
R1.1.4 oxidizing agent (oxidant) and reducing agent (reductant) with common examples
R1.1.5 half reactions and balanced net equations
R1.1.6 electrochemical cells (galvanic and electrolytic)
R1.1.7 electrodes (anode and cathode)
R1.1.8 salt bridge
R1.1.9 notation for half and whole electrochemical cells
R1.1.10 standard reduction potentials (E§)
R1.1.11 reactivity series
R1.1.12 displacement reaction of metals
R1.1.13 commercial cells and batteries
R1.1.14 electroplating
R1.1.15 corrosion of metals
R1.1.16 sacrificial anodes, cathodic protection
R1.1.17 electrolytic refinement of metals
R1.2 Precipitation reactions result in the appearance of a solid from reactants in aqueous solution.
R1.2.1 examples of common precipitates (AgCl, BaSO4)
R1.2.2 simple solubility rules
R1.2.3 concept of Ksp
R1.2.4 complete and net ionic equations
R1.2.5 spectator ions
R1.2.6 applications of precipitation reactions in qualitative and quantitative (gravimetric) inorganic analysis
R1.3 Acid-base reactions involve transfer of protons from donors to acceptors.
R1.3.1 define acids and bases using Bronsted-Lowry theory
R1.3.2 identify acid-base conjugate pairs
R1.3.3 concept of strong and weak acids and bases
R1.3.4 examples of strong acids (HCl, HNO3, H2SO4) and weak acids (HF, CH3COOH), and strong bases (NaOH, KOH) and weak bases (NH3)
R1.3.5 definitions of Kw and pH
R1.3.6 definitions of Ka, Kb
R1.3.7 relate the strength of acids and bases to the strength of their conjugates
R1.3.8 reactions of acids with metals
R1.3.9 safety precautions in handling acids and bases
R1.4 Polymerization reactions produce large molecules with repeating units.
R1.4.1 explain the terms monomer, polymer, polymerization, repeat unit
R1.4.2 recall simple physical properties of addition and condensation polymers and relate these properties to structure
R1.4.3 use the following terms appropriately to describe the structure and properties of polymers: thermoset, thermoplastic, elastomer, vulcanization, amorphous, crystalline
R1.4.4 describe the effects of chain length, side branches and cross-linking on polymer properties
R1.4.5 amino acids and proteins
R2.0 Chemical reactions involve energy changes.
R2.1 All chemical reactions involve energy transformations.
R2.1.1 the law of conservation of energy
R2.1.2 the terms exothermic, endothermic, combustion, enthalpy, entropy, activated complex, activation energy
R2.1.3 DH and identify whether a reaction is exothermic or endothermic given DH values
R2.1.4 potential energy-reaction coordinate diagrams change if a catalyst is present in a reaction
R2.1.5 the origin of heat of reaction in terms of the breaking and forming of bonds and bond energy
R2.1.6 enthalpy changes in a reaction in relationship to bond energies
R2.1.7 the relationship between amount and heat of reaction
R2.1.8 thermochemical equations including heat of reaction in them
R2.1.9 heats and molar heats of formation and combustion and neutralization
R2.1.10 calorimetry and its use in measuring and calculating the heat content of fuels
R2.1.11 specific heat (Q) capacity
R2.1.12 Hess's law of the heats of summation
R2.2 The spontaneous directions of chemical reactions are towards lower energy and greater randomness.
R2.2.1 entropy and enthalpy considerations to explain the spontaneity of reactions
R3.0 The mole concept and stoichiometry enable the determination of quantities in chemical processes.
R3.1 The mole, defined arbitrarily using the isotope carbon-12, is the basic quantity in stoichiometric calculations.
R3.1.1 Avogadro's number, relative atomic mass, relative molecular mass and relative formula mass, molar mass, molar volume, molarity, empirical and molecular formulas, percentage composition
R3.2 Every chemical reaction can be represented by a balanced equation, whose coefficients indicate both the number of reacting particles and the reacting quantities in moles.
R3.2.1 law of conservation of mass
R3.2.2 coefficients, subscripts of state, formulas
R3.2.3 molecular and ionic equations, half and net ionic equations
R3.3.4 A balanced equation can be used when determining whether reagents are limiting or in excess.
R3.4 The use of molarity for expressing concentration allows easy conversions between volume of solution and moles of solute.
R3.4.1 concentrations (molarity, percentage volume, percentage mass, ppm)
R3.4.2 dilution, concentrated, dilute, saturated, solubility
R3.5 The ideal gas equation may be used to relate the volume of a gas at defined temperature and pressure to its quantity in moles.
R3.5.1 gas properties: compressibility, diffusion, solubility
R3.5.2 kinetic theory of particles, temperature and energy, zero and Kelvin temperature scale
R3.5.3 real and ideal gases
R3.5.4 STP and SLC, molar volume
R3.5.5 gas laws: Boyle's law, Charles's law, Gay-Lusac's law and Avogadro's hypothesis
R3.5.6 combined gas equation and ideal gas equation
R4.0 Specialized qualitative and quantitative techniques are used to determine the quantity, composition and type.
R4.1 Techniques such as volumetric and gravimetric analysis are used to determine
amounts of reactants and products.
R4.1.1 equivalence and end-points, titre, aliquot, standard solution, primary standard
R4.1.2 titrations: back, redox, acid base precipitation, excess/limiting reagents
R4.1.3 equipment, procedures and errors
R4.1.4 empirical and molecular formulas
R4.2 Specialized techniques and instrumentation are used in chemical analysis.
R4.2.1 spectroscopy and colorimetry: mass spectrometry, IR spectroscopy, UV spectroscopy, atomic absorption spectroscopy, line spectra, absorption and emission, calibration
R4.2.2 chromatography: thin layer chromatography, gas chromatography, stationary and mobile phases, adsorption, retention time, Rf
R4.3 Qualitative and quantitative testing may be used to determine the composition or type of material.
R5.0 Chemical reactions are influenced by the conditions under which they take place and, being reversible, may reach a state of equilibrium.
R5.1 Chemical reactions occur at different rates, and changing the nature of the reactants, temperature, or concentration, or introducing a catalyst may alter these.
R5.1.1 rate of reaction, collision theory, units, average reaction rate
R5.1.2 factors nature of reactants, concentration or pressure, the surface area, temperature, catalysts/inhibitors
R5.1.3 activation energy, reaction coordinates, PE v. reaction coordinate diagrams, reaction mechanisms, intermediates, Arrhenius distributions
R5.1.4 graphs of rate data
R5.2 Life is maintained by chemical reactions especially those catalysed by large molecules called enzymes.
R5.2.1 enzymes, specificity, substrates, active site, denaturation
R5.3 Chemical reactions may be reversible.
R5.3.1 steady state and dynamic equilibrium
R5.3.2 phase changes, gas phase reactions, redox, acid-base, solubility processes and reactions in aqueous solution precipitation
R5.3.3 saturated, unsaturated, dilute, concentrated, strong electrolyte, weak electrolyte, non-electrolyte, strong and weak acids
R5.3.4 the equilibrium law and application
R5.3.5 equilibrium constants Keq, Ksp, Kw, Ka
R5.3.6 extent of reactions, solubility of salts
R5.3.7 equilibrium concentrations
R5.4 Reversible chemical reactions may reach a state of dynamic balance known as equilibrium which, when disturbed, will be re-established.
R5.4.1 Le Chatelier's principle, position of an equilibrium: concentration, pressure, temperature and presence of catalysts
R5.4.2 reaction quotient Q and the equilibrium constant K
R5.4.3 collision frequency theory
3.0 Sample course of study - Unit titles
Unit titles
Unit 1. Water quality chemistry
Take me to the river
The problems of water quality and water usage
The properties of water and aqueous solutions
Fundamental concepts related to formulas, nomenclature and solubility.
Monitoring water quality in West Creek, field trip
The physical properties of gases, including changes in volume caused by pressure and temperature changes are explored, and the chemical properties of five gases are tested in the laboratory
The concepts of rates of reaction, molarity and energy changes
Task on global warming GOs Category Technique, description and conditions
Unit 2. Forensic chemistry
A range of instrumental methods capable of handling very small samples are examined mass spectrometry, atomic emission spectroscopy, ultraviolet-visible spectroscopy, infrared spectroscopy and x-ray photoelectron spectroscopy are examined and linked to the structure of atoms and molecules
Unit 3. Petroleum chemistry
3.1 to build
3.2 to burn
The properties and chemistry of petroleum and to some of the central concepts of organic chemistry: naming, properties, covalent bonding, isomers and petrochemical synthesis
Unit 4. Cold packs chemistry
Design of a cold pack
Data to be collected and used to prepare a report that answers a research question regarding the best design for a chemical cold pack.
Unit 5. Ocean chemistry
Major chemical ideas such as the molecular-kinetic theory and energy distribution among molecules, enthalpy changes involving ionic compounds, the role of entropy changes in determining the feasibility of chemical reactions and applications of equilibrium to phase changes, solubility and the behaviour of weak acids, as well as redox reactions, shells, corals, rocks, ships and the behaviour of water itself
Sample units of work
Unit 6. Carbon dioxide in drinks chemistry
Where does the fizz come from?
Carry out and examine reactions which produce carbon dioxide, and others
Write balanced equations for these reactions
Discussion of solubility of carbon dioxide in water, equations for formation of carbonic acid
Design a procedure to collect gas from soft drink and use limewater test to identify
Signs of chemical reactions
Writing and balancing chemical equations
Solubility of carbon dioxide in water
Formation of carbonic acid (dynamic equilibrium), test for carbon dioxide, and associated equations
Unit 7. Flavours chemistry
Use computer software to model structure and formation of esters
Make esters
Effect of colour on flavour
Distillation of esters from fruit
Nomenclature, structure, properties, formation and origins of esters
Steam distillation
Unit 8. Sugars in food chemistry
How much sugar?
Influence of sugar/sweeteners on flavour
Structures of sugars common in drinks
Gravimetric analysis on a number of sweet drinks to determine sugar content
Compare results to label values
Error analysis
The need for a systematic way to represent concentration
Structure of sugars
Qualitative vs quantitative analysis
Gravimetric analysis methods
The mole (isotopes, atomic mass), and molarity
Calculations of concentration grams/litre, percentage composition
Unit 9. Alcohol in beverages chemistry
How much alcohol?
Alcohol production/fermentation
Alcohol content of various beverages, and the way its concentration is expressed
Experimentally determine the alcohol content in beverages using volumetric procedures and boiling point depression (use CBLs and probes)
Calculations to express this concentration in a number of forms
Alcohol production
"Standard" drink
Methods of expressing alcohol concentration "Proof", percentage per volume in beverage, mG/litre in blood
Volumetric procedures
Party drinks
Drinks and their contents are frequently an issue in your society
Whether relating alcohol consumption to the possibility of drink driving, examining labels for undesired contents, or watching calorie intake, most students will benefit from undertaking a course of work that will enable them to recognize some of the contents in drinks, their reasons for being there, and understanding information regarding their concentration
Intrinsic subject material What's in a "party drink"
Discuss necessary ingredients in a good "soft" or alcoholic drink
Prepare, sample and analyse a simple fizzy drink made in the laboratory
Party drinks contain sugar, food acid, and possibly carbon dioxide, alcohol and flavours
Unit 10. Acids in drinks chemistry
How much acid?
The variety of acids found in drinks, and reasons for their presence
Use titration equipment to determine acid content of various drinks
Equations to represent the reactions involved, and calculate associated quantities and concentrations
Plot pH curve using CBLs and pH probes
Interpret curve shape
Unit 11. Food acids chemistry
Names and structures of various food acids, and uses
Definitions of acids, bases
Properties of acids and bases, pH, neutralization, strong and weak acids/bases, buffers, indicators, titrations
Stoichiometry
Unit 12. Fuels for transport chemistry
What makes it go?
Modern transport encompasses a wide variety of vehicles with a wide variety of energy sources
The various sources of these energies, their use and their particular requirements necessitate careful consideration of a fuel's property
Chemical structures, combustion reactions, rates of reactions and energetics necessary to fully understand the selection and use of energy sources for transport
Use and investigate a variety of powered model vehicles, e.g. paper match rockets, "put-put" boats, soda bulb/balloon rockets, battery powered car
There are a variety of fuels, and they work in different ways
Unit 13. Motor car fuel chemistry
What makes cars go?
The variety of fuels and their combustion properties
Ferment alcohol from sugar/other, and distil
Test product
Fractional distillation, examine properties of resultant fractions
Examine an engine, identify combustion chamber, pistons, intake and outlet valves
Combustion equations
Types of fuels (petrol, diesel, biodiesel, ethanol), their sources, formation, extraction, chemical composition
Fermentation, fractional distillation
Basic engine functioning
Combustion reactions, exothermic, endothermic
What makes a good fuel?
Measure heat content of a variety of fuels
Compare results to accepted H values
New vehicle fuels
Locate the catalytic converter in a car
Perform experiments demonstrating catalysis
Perform H calculations
Heat content of fuels, calorimetry, laboratory and bomb calorimeters, H, Hess's law, latent heat
Portability
Fuel pollutants and effects acid rain, greenhouse, photochemical smog
Catalytic converters catalysis
Ignition temperatures and flash points (mobile phones)
Safety
Fuel costs
Combustion efficiency
Alternative uses for fuels
Unit 14. Rocket fuel chemistry
What makes rockets go?
Equations for the combustion/reaction of aviation and rocket fuels
Investigate gas laws
Use stoichiometry and gas laws to calculate volume/pressure changes
Experiment soda bulb rockets calculate mass and volume of gas released
Aviation and rocket fuels their chemical composition, combustion and specific requirements the need for oxygen
Basic rocket functioning
Kinetic theory of gases
Gas laws (Boyle's, Charles's, ideal), absolute zero and Kelvin scale
Various units of pressure
Unit 15. Fertilizers and explosives chemistry
The chemistry and technological difficulties of fixing nitrogen, the concept of how equilibrium applies to the production of fertilizers, and describes nitrogen's role in the explosives industry
Analysis of Haber Process software
Investigate a real life explosive accident (space shuttle, sawmills, car LPG cylinders )
Observe and analyse "exploding can" demonstration
Experiment effect of temperature on rate of reaction (use CBL light meters)
Use stoichiometry and gas laws to calculate volume changes in explosions
Rates of reaction/combustion and influencing factors fuel/oxygen concentrations, surface area, temperature
Fuel cells Use of fuel cells in spacecraft
Experiment basic electrochemical cells
Basic redox, electrochemical cells, Eo, fuel cells