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.
To be used for the first time in 2008.
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
1: 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.
S1.5 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.
S2.0 Materials can be categorised 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.
S2.2 The macroscopic properties are related to their microscopic
properties.
S2.3 Pairs of atoms may be bound together by the sharing of
electrons between them in a covalent bond.
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.
S2.5 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.
S2.8 Forces weaker than covalent bonding exist between molecules.
S2.9 The structure of a metal involves positive ions embedded in
a sea of electrons.
S2.10 Materials may be elements, compounds or mixtures.
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.
2: 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.
R1.2 Precipitation reactions result in the appearance of a
solid from reactants in aqueous solution.
R1.3 Acid-base reactions involve transfer of protons
from donors to acceptors.
R1.4 Polymerisation reactions produce large molecules
with repeating units.
R2.0 Chemical reactions involve energy changes.
R2.1 All chemical reactions involve energy
transformations.
R2.2 The spontaneous directions of chemical
reactions are towards lower energy and greater
randomness.
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.
R3.2 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.
R3.3 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.
R3.5 The ideal gas equation may be used to relate the
volume of a gas at defined temperature and pressure
to its quantity in moles.
R4.0 Specialised qualitative and quantitative
techniques are used to determine the quantity,
composition and type of reaction.
R4.1 Techniques such as volumetric and gravimetric
analysis are used to determine amounts of reactants
and products.
R4.2 Specialised techniques and instrumentation are
used in chemical analysis.
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.
R5.2 Life is maintained by chemical reactions,
especially those catalysed by large molecules called
enzymes.
R5.3 Chemical reactions may be reversible.
R5.4 Reversible chemical reactions may reach a state
of dynamic balance known as equilibrium which,
when disturbed, will be re- established.
S1 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 categorised 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, ionisation
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 and
atomic 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 vaporisation, 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
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 oxidising 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 (HCl, HNO3, H2SO4) and weak (HF, CH3CO2H)
acids and strong (NaOH, KOH) and weak (NH3) bases
R1.3.5definitions 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 Polymerisation reactions produce large molecules with repeating
units.
R1.4.1 explain the terms monomer, polymer, polymerisation, 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,
vulcanisation, 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
neutralisation
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, Charles's, Gay-Lusac's and Avogadro's
hypothesis
R3.5.6 combined gas equation and ideal gas equation
R4.0 Specialised 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 Specialised 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
Sample course of study
Unit titles
1. 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
2. Fertilisers and explosives
The chemistry and technological difficulties of fixing nitrogen,
the concept of how equilibrium applies to the production of
fertilisers, and describes nitrogen's role in the explosives
industry
Analysis of Haber Process software
2a. Forensic chemistry and big instruments
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
3. Petroleum
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
4. Cold packs
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.
5. The ocean is deep
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
6. Party drinks
Drinks and their contents are frequently an issue in our 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 recognise 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
6. 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
7. Flavours
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
8. 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
9. 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", % per
volume in beverage, mG/litre in blood
Volumetric procedures
10. 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
11. Names and structures of various food acids, and uses
Definitions of acids, bases
Properties of acids and bases, pH, neutralisation, strong and
weak acids/bases, buffers, indicators, titrations
Stoichiometry
12. 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
13. 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
14. 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
15. Oops! Ka boom!
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