U1 Topic 3: Chemical reactions  reactants, products and energy change
Chemical reactions
• recall that chemical reactions and phase changes involve energy changes commonly observable as changes in the temperature of the surroundings and/or the emission of light
• deduce and construct balanced chemical equations when reactants and products are specified and apply state symbols (s), (l), (g) and (aq).
• Balancing equations should cover a variety of reactions, e.g. single displacement, doubledisplacement, acidbase, combustion, combination,
decomposition and simple redox reactions.
• Names of the change of states should be covered: melting, freezing, vaporisation (evaporation and boiling) condensation, sublimation
and deposition.
• recall that chemical reactions and phase changes involve energy changes commonly observable as changes in the temperature of the surroundings and/or the emission of light
• deduce and construct balanced chemical equations when reactants and products are specified and apply state symbols (s), (l), (g) and (aq).
• Balancing equations should cover a variety of reactions, e.g. single displacement, doubledisplacement, acidbase, combustion, combination,
decomposition and simple redox reactions.
• Names of the change of states should be covered: melting, freezing, vaporisation (evaporation and boiling) condensation, sublimation
and deposition.



Practice Soichiometry Question Sets: (Taken from my IB website)
3. Balancing chemical equations:
Balancing Set 1
Balancing Set 2
Balancing Set 3
3b. Writing chemical equations:
Writing Set 1
Writing Set 2
Writing Set 3
3. Balancing chemical equations:
Balancing Set 1
Balancing Set 2
Balancing Set 3
3b. Writing chemical equations:
Writing Set 1
Writing Set 2
Writing Set 3
Exothermic and endothermic reactions
• explain how endothermic and exothermic reactions relate to the law of conservation of energy and the breaking and reforming of bonds; understand
that heat energy is released or absorbed by the system to or from the surrounds
• understand that heat is a form of energy and that temperature is a measure of the average kinetic energy of the particles
• apply the relationship between temperature and enthalpy changes to identify thermochemical reactions as exothermic or endothermic; deduce from enthalpy level diagrams and thermochemical equations the relative stabilities of reactants and products, and the sign of the enthalpy change (ΔH) for a reaction
• explain, in terms of average bond enthalpies, why reactions are exothermic or endothermic
• construct and use appropriate representations (including chemical symbols and formulas, and chemical and thermochemical equations) to communicate conceptual understanding, solve problems and make predictions
• calculate the heat change for a substance given the mass, specific heat capacity and temperature change
• use data to calculate the enthalpy change (ΔH) for a reaction.
• Mandatory practical: Conduct a calorimetry experiment to measure the enthalpy of a reaction.
• Students should be aware of the limitations of using average bond enthalpies to calculate enthalpy change.
• Consider reactions in aqueous solutions and combustion reactions.
• Formulas: ΔH = H (products) – H (reactants), ΔH = energy required to break bonds – energy released when bonds are formed, 𝑞 = 𝑚𝑐Δ𝑇
• explain how endothermic and exothermic reactions relate to the law of conservation of energy and the breaking and reforming of bonds; understand
that heat energy is released or absorbed by the system to or from the surrounds
• understand that heat is a form of energy and that temperature is a measure of the average kinetic energy of the particles
• apply the relationship between temperature and enthalpy changes to identify thermochemical reactions as exothermic or endothermic; deduce from enthalpy level diagrams and thermochemical equations the relative stabilities of reactants and products, and the sign of the enthalpy change (ΔH) for a reaction
• explain, in terms of average bond enthalpies, why reactions are exothermic or endothermic
• construct and use appropriate representations (including chemical symbols and formulas, and chemical and thermochemical equations) to communicate conceptual understanding, solve problems and make predictions
• calculate the heat change for a substance given the mass, specific heat capacity and temperature change
• use data to calculate the enthalpy change (ΔH) for a reaction.
• Mandatory practical: Conduct a calorimetry experiment to measure the enthalpy of a reaction.
• Students should be aware of the limitations of using average bond enthalpies to calculate enthalpy change.
• Consider reactions in aqueous solutions and combustion reactions.
• Formulas: ΔH = H (products) – H (reactants), ΔH = energy required to break bonds – energy released when bonds are formed, 𝑞 = 𝑚𝑐Δ𝑇







ATAR Notes Chemistry  Chemical reactions
Measurement uncertainty and error
• distinguish between precision and accuracy and appreciate that all measurements have limits to their precision and accuracy that must be
considered when evaluating experimental results
• distinguish between qualitative and quantitative data; appreciate that quantitative data obtained from measurements is always associated with random error/measurement uncertainties
• communicate measurement uncertainties as a range (±) to an appropriate precision
• understand that propagation of random error in data processing shows the impact of measurement uncertainties on the final result
• calculate the measurement uncertainties in processed data, including the use of absolute uncertainties and percentage uncertainties
• construct and use appropriate graphical representations to communicate understanding, solve problems and make predictions; interpret graphs in terms of the relationship between dependent and independent variables; draw and interpret bestfit lines or curves through data points, including evaluating when it can and cannot be considered as a linear function
• calculate the percentage error when the experimental result can be compared with a theoretical or accepted result (value)
• distinguish between random and systematic errors; understand that experimental design and procedure usually leads to systematic errors in
measurement, which causes a deviation in a direction and appreciate that repeated trials and measurements will reduce random error but not systematic error
• analyse the impact of random error/measurement uncertainties and systematic errors in experimental work and evaluate how these errors/measurement uncertainties can be reduced
• understand that the number of significant figures in a result is based on the figures given in the data and determine results of calculations to the appropriate number of significant figures.
• Only a simple treatment of error analysis is required. For functions such as addition or subtraction, absolute uncertainties should be added. For
multiplication, division and powers, percentage uncertainties can be added.
• Formula: 𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦 (%) = 𝑎𝑏𝑠𝑜𝑙𝑢𝑡𝑒 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦/𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡 × 100
• When adding or subtracting, the final answer should be given to the least number of decimal places. When multiplying or dividing, the final answer
should be given to the least number of significant figures.
• Formula: 𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑒𝑟𝑟𝑜𝑟 (%) =  𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝑣𝑎𝑙𝑢𝑒−𝑡𝑟𝑢𝑒 𝑣𝑎𝑙𝑢𝑒  / 𝑡𝑟𝑢𝑒 𝑣𝑎𝑙𝑢𝑒 × 100
• distinguish between precision and accuracy and appreciate that all measurements have limits to their precision and accuracy that must be
considered when evaluating experimental results
• distinguish between qualitative and quantitative data; appreciate that quantitative data obtained from measurements is always associated with random error/measurement uncertainties
• communicate measurement uncertainties as a range (±) to an appropriate precision
• understand that propagation of random error in data processing shows the impact of measurement uncertainties on the final result
• calculate the measurement uncertainties in processed data, including the use of absolute uncertainties and percentage uncertainties
• construct and use appropriate graphical representations to communicate understanding, solve problems and make predictions; interpret graphs in terms of the relationship between dependent and independent variables; draw and interpret bestfit lines or curves through data points, including evaluating when it can and cannot be considered as a linear function
• calculate the percentage error when the experimental result can be compared with a theoretical or accepted result (value)
• distinguish between random and systematic errors; understand that experimental design and procedure usually leads to systematic errors in
measurement, which causes a deviation in a direction and appreciate that repeated trials and measurements will reduce random error but not systematic error
• analyse the impact of random error/measurement uncertainties and systematic errors in experimental work and evaluate how these errors/measurement uncertainties can be reduced
• understand that the number of significant figures in a result is based on the figures given in the data and determine results of calculations to the appropriate number of significant figures.
• Only a simple treatment of error analysis is required. For functions such as addition or subtraction, absolute uncertainties should be added. For
multiplication, division and powers, percentage uncertainties can be added.
• Formula: 𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦 (%) = 𝑎𝑏𝑠𝑜𝑙𝑢𝑡𝑒 𝑢𝑛𝑐𝑒𝑟𝑡𝑎𝑖𝑛𝑡𝑦/𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑚𝑒𝑛𝑡 × 100
• When adding or subtracting, the final answer should be given to the least number of decimal places. When multiplying or dividing, the final answer
should be given to the least number of significant figures.
• Formula: 𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑒𝑟𝑟𝑜𝑟 (%) =  𝑚𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝑣𝑎𝑙𝑢𝑒−𝑡𝑟𝑢𝑒 𝑣𝑎𝑙𝑢𝑒  / 𝑡𝑟𝑢𝑒 𝑣𝑎𝑙𝑢𝑒 × 100






ATAR Notes Chemistry  Errors
Fuels
• compare fuels, including fossil fuels and biofuels, in terms of their energy output, and evaluate their suitability for purpose, and the nature of products
of combustion.
• compare fuels, including fossil fuels and biofuels, in terms of their energy output, and evaluate their suitability for purpose, and the nature of products
of combustion.
(no video)
• compare fuels, including fossil fuels and biofuels, in terms of their energy output, and evaluate their suitability for purpose, and the nature of products
of combustion.
• compare fuels, including fossil fuels and biofuels, in terms of their energy output, and evaluate their suitability for purpose, and the nature of products
of combustion.
(no video)
Mole concept and law of conservation of mass
• recognise that a mole is a precisely defined quantity of matter equal to Avogadro’s number of particles
• appreciate the law of conservation of mass and understand that the mole concept relates mass, moles and molar mass
• understand that the empirical formula expresses the simplest whole number ratio of elements in a compound
• use the appropriate stoichiometric ratio to determine that reactants can be limiting
• appreciate that experimental yield can be different from theoretical yield
• use appropriate mathematical representation to solve problems and make predictions, including using the mole concept to calculate the mass of reactants and products; amount of substance in moles; number of representative particles; and molar mass of atoms, ions, molecules and formula units
• use appropriate mathematical representation to solve problems and make predictions, including determining the percentage composition from relative
atomic masses; empirical formula of a compound from the percentage composition by mass; and molecular formula of a compound from its empirical
formula and molar mass
• calculate percentage yield from experimental or given data.
• Mandatory practical: Derive the empirical formula of a compound from reactions involving mass changes.
• Formula: 𝑝𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 (%) = 𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑦𝑖𝑒𝑙𝑑/𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 × 100
• recognise that a mole is a precisely defined quantity of matter equal to Avogadro’s number of particles
• appreciate the law of conservation of mass and understand that the mole concept relates mass, moles and molar mass
• understand that the empirical formula expresses the simplest whole number ratio of elements in a compound
• use the appropriate stoichiometric ratio to determine that reactants can be limiting
• appreciate that experimental yield can be different from theoretical yield
• use appropriate mathematical representation to solve problems and make predictions, including using the mole concept to calculate the mass of reactants and products; amount of substance in moles; number of representative particles; and molar mass of atoms, ions, molecules and formula units
• use appropriate mathematical representation to solve problems and make predictions, including determining the percentage composition from relative
atomic masses; empirical formula of a compound from the percentage composition by mass; and molecular formula of a compound from its empirical
formula and molar mass
• calculate percentage yield from experimental or given data.
• Mandatory practical: Derive the empirical formula of a compound from reactions involving mass changes.
• Formula: 𝑝𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 (%) = 𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑦𝑖𝑒𝑙𝑑/𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 × 100









ATAR Notes Chemistry  Mass Stoichiometry
Practice Soichiometry Question Sets: (Taken from my IB website)
4. Calculating molecular mass (covalent) and formula mass (ionic):
Formula mass Set 1
Formula mass Set 2
Formula mass Set 3
6. Converting between mass and moles:
Mass and mole conversions Set 1
Mass and mole conversions Set 2
Mass and mole conversions Set 3
7. Determining empirical and molecular formula from experimental data:
Empirical formula Set 1
Empirical formula Set 2
Empirical formula Set 3
8. Determining the percentage composition within a compound:
Percent composition Set 1
Percent composition Set 2
Percent composition Set 3
9. Calculating stoichiometric problems involving mass:
Mass stoichiometric problems Set 1
Mass stoichiometric problems Set 2
Mass stoichiometric problems Set 3
10. Determining which compound is the limiting reagent:
Limiting reagent Set 1
Limiting reagent Set 2
Limiting reagent Set 3
11. Determining the percentage yield in experimental data:
Percentage yield Set 1
Percentage yield Set 2
4. Calculating molecular mass (covalent) and formula mass (ionic):
Formula mass Set 1
Formula mass Set 2
Formula mass Set 3
6. Converting between mass and moles:
Mass and mole conversions Set 1
Mass and mole conversions Set 2
Mass and mole conversions Set 3
7. Determining empirical and molecular formula from experimental data:
Empirical formula Set 1
Empirical formula Set 2
Empirical formula Set 3
8. Determining the percentage composition within a compound:
Percent composition Set 1
Percent composition Set 2
Percent composition Set 3
9. Calculating stoichiometric problems involving mass:
Mass stoichiometric problems Set 1
Mass stoichiometric problems Set 2
Mass stoichiometric problems Set 3
10. Determining which compound is the limiting reagent:
Limiting reagent Set 1
Limiting reagent Set 2
Limiting reagent Set 3
11. Determining the percentage yield in experimental data:
Percentage yield Set 1
Percentage yield Set 2