QCE Chemistry Syllabus Guide (Sciences)

Atomic structure

• Describe that atoms can be modelled as a nucleus surrounded by electrons in distinct energy levels.
• Discriminate between the terms atomic number (Z), mass number (A) and isotopes of an element.
• Apply the nuclear symbol notation to determine the number of protons, neutrons and electrons in atoms, ions and isotopes.
• State the relative energies of the s, p and d orbitals.
• Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.
• Determine full and condensed electron configurations for atoms and ions up to Z = 36, e.g. 1s² 2s² 2p⁶ 3s² 3p⁵ and [Ne]3s² 3p⁵.
• Identify the electron configuration of Cr and Cu as exceptions.
• Explain how successive ionisation energy data is related to the electron configuration of an atom.

Isotopes

• Describe that isotopes are atoms of the same element that have different numbers of neutrons.
• State that isotopes can be represented in the form ᴬX (IUPAC) or X-A.
• Identify that isotopes of an element have the same electron configuration and possess similar chemical properties but have different physical properties.
• Explain that the relative atomic mass of an element is the ratio of the weighted average mass per atom of the naturally occurring form of the element to 1/12 the mass of an atom of carbon-12.

Analytical techniques

• State that mass spectrometry involves the ionisation of substances and the separation and detection of the resulting ions. (The operation of the mass spectrometer is not required.)
• Analyse mass spectrometry spectra, to determine the isotopic composition of elements, the relative atomic mass of an element and percentage abundances of the isotopes of an element.
• Discriminate between absorption and emission line spectra.
• Explain that flame tests and atomic absorption spectroscopy (AAS) rely on electron transfer between atomic energy levels.
• Explain that the emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels (Bohr model), which converge at higher energies.
• Analyse flame tests and atomic absorption spectroscopy (AAS) to identify elements and determine the concentration of metallic ions in solution.

Periodic table and trends

• State that elements are represented by symbols.
• Identify that the structure of the periodic table based on increasing atomic number.
• Identify that the periodic table is arranged into four blocks associated with the four sub-levels — s, p, d and f.
• Describe the relationship between the structure of the periodic table and the electronic configuration of atoms.
• Explain that elements of the periodic table show trends in chemical and physical properties across periods and down groups as exemplified by groups 1, 2, 13–18 and period 3.
• Compare the metallic and non-metallic behaviours of elements, including group trends and the reactivity for the alkali metals (Li–Cs) and the halogens (F–I).
• Identify that oxides change from basic through amphoteric to acidic across period 3.
• Analyse data for atomic radii, valencies, ionic radii, 1st ionisation energy and electronegativities to determine periodic trends, patterns and relationships.

Introduction to bonding

• Explain that the ability of atoms to form chemical bonds, is related to the arrangement of electrons in the atom and the stability of the valence electron shell.
• Identify that the number of electrons lost, gained or shared is determined by the electron configuration of the atom.
• State that transition elements can form more than one ion.
• Explain that ions are atoms or groups of atoms that are electrically charged due to an imbalance in the number of electrons and protons.
• Explain that chemical bonds are caused by electrostatic attractions that arise because of the sharing or transfer of electrons.
• Identify that the valency is a measure of the number of bonds that an atom can form.
• Determine the formula and IUPAC name of ionic and molecular compounds.
• Discriminate between the terms empirical formula, molecular formula and the formula unit.
• Determine Lewis (electron dot) structure of molecules and ions showing all valence electrons for up to four electrons pairs for each atom.
• Identify the numbers of bonding and lone pairs of electrons around each atom in a molecule.

Science as a Human Endeavour (SHE)

• Appreciate that experiments provided evidence that enabled scientists to develop models of the atom.
• Consider the role Geiger-Marsden’s gold foil experiments and Maria Goeppert Mayer’s nuclear shell model played in the development of atomic theory.
• Appreciate that radioisotopes require careful evaluation and monitoring because of the potential harmful effects to humans and/or the environment.
• Explore the use of radioisotopes for carbon dating and radiotherapy and Marie Curie’s contribution to research on radioactivity.
• Appreciate that analysis of the distribution of elements in living things, Earth and the universe has informed a wide range of scientific understandings.
• Explore the composition of stars and Cecilia Payne-Gaposchkin’s contribution to astrophysics.

Science inquiry

• Investigate: flame tests to identify elements
• Investigate: mass spectra and isotopes
• Investigate: atomic absorption spectroscopy (AAS) and the concentration of aqueous metallic ions

Compounds and mixtures

• State that pure substances may be elements or compounds.
• Identify that pure substances have distinct measurable properties (e.g. melting and boiling point, reactivity, strength, density) and mixtures have properties dependent on the identity and relative amounts of the substances that make them up.
• Discriminate between heterogeneous and homogeneous mixtures.
• Analyse data to determine the physical properties of pure substances and mixtures.

Bonding and properties

• Describe the properties of ionic, covalent and metallic compounds, e.g. melting and boiling point, thermal and electrical conductivity, strength and hardness.
• Explain that the type of bonding within ionic, metallic and covalent substances determines their physical properties.
• Explain the properties of ionic compounds by modelling ionic bonding as ions arranged in a crystalline lattice structure with strong electrostatic forces of attraction between oppositely charged ions.
• Discriminate between ionic and metallic bonding.
• Explain the properties of covalent compounds by modelling covalent bonding as the sharing of an electron pair in the region between two nuclei with a strong electrostatic force of attraction between both nuclei.
• Discriminate between covalent molecules, giant covalent networks and allotropes of carbon.
• Explain that hydrocarbons, including alkanes (saturated), alkenes (unsaturated) and benzene, have different chemical properties that are determined by the nature of the bonding within the molecules.
• Analyse data to determine the properties, structure and bonding of ionic, covalent and metallic compounds.

Science as a Human Endeavour (SHE)

• Appreciate that the development of nanomaterials is important to meet a range of contemporary needs and have specific properties related to the size of the particles (1–100 nm).
• Consider the benefits and potential risks associated with the use of nanomaterials in consumer products, health care, transportation, energy and agriculture.
• Appreciate that impurities can affect the physical and chemical properties of substances, resulting in inefficient or unwanted chemical reactions.
• Explore Ellen Swallow Richard’s contribution to developing water quality standards and the new discipline of home economics.
• Appreciate that carbon has a range of properties that allow a variety of interactions which are pivotal to the formation of biochemical molecules such as carbohydrates, proteins and DNA.
• Consider whether life exists elsewhere in the universe and if it could be carbon-based as it is on Earth.
• Explore Millicent Goldschmidt’s contribution to the development of astrobiology.

Science inquiry

• Investigate: the separation of mixtures based on physical properties
• Investigate: the properties of ionic, metallic, and covalent compounds
• Investigate: tests to distinguish alkanes and alkenes

Chemical reactions

• Identify 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.
• Determine balanced chemical equations, including state symbols (s), (l), (g) and (aq), for a variety of reactions, e.g. single displacement, double-displacement, acid-base, combustion, combination, decomposition and simple redox reactions.

Exothermic and endothermic reactions

• State that heat is a form of energy, and that temperature is a measure of the average kinetic energy of the particles.
• Explain how endothermic and exothermic reactions relate to the law of conservation of energy and the breaking and reforming of bonds.
• Discriminate between exothermic and endothermic reactions.
• Sketch enthalpy level diagrams for exothermic and endothermic reactions.
• Analyse enthalpy level diagrams and thermochemical equations to determine 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.
• Identify the limitations of using average bond enthalpies to calculate enthalpy change.
• Calculate the heat change (Q) for a substance given the mass, specific heat capacity and temperature change. (Formula: Q = mcΔT)
• Calculate the enthalpy change (ΔH) for a reaction given temperature changes, quantities of reactants and mass of water. (Formula: ΔH = H(products) − H(reactants))
• Analyse data for heat of combustion, heat of neutralisation and reactions in aqueous solutions to determine heat, mass, specific heat capacity, temperature and enthalpy change.

Mole concept and law of conservation of mass

• State that a mole is a precisely defined quantity of matter equal to Avogadro’s number of particles.
• State the law of conservation of mass.
• Explain that the mole concept relates mass, moles and molar mass.
• Apply 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. (Formula: moles (n) = mass (m) / molar mass (M))
• Determine 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.
• Determine limiting reactants.
• Discriminate between experimental and theoretical yield.
• Analyse data to determine percentage and theoretical yield. (Formula: percentage yield (%) = experimental yield / theoretical yield × 100/1)

Science as a Human Endeavour (SHE)

• Appreciate that chemistry principles can be applied to industrial processes to reduce energy requirements.
• Explore how industries are reducing their energy requirements in order to save money and reduce greenhouse gas emissions.
• Appreciate that bodies rely on the exothermic reaction of respiration to provide us with sufficient energy.
• Explore how cells use food and convert it to energy and Gerty Cori’s contribution to the treatment of diabetes.
• Appreciate that biofuels are more efficient and have less environmental impact than fossil fuels.
• Evaluate fuels, including fossil fuels and biofuels, in terms of their energy output, their suitability for purpose, and the nature of products of combustion.

Science inquiry

• Investigate: types of chemical reactions
• Investigate: limiting reagent/s and percentage yield
• Investigate: the empirical formula of a compound from reactions involving mass change
• Investigate: the enthalpy change of a reaction, e.g. calorimetry or Hess’s Law

Intermolecular forces

• Apply the valence shell electron pair repulsion (VSEPR) theory to determine the shape and bond angles of linear, bent, trigonal planar, tetrahedral and pyramidal molecules. (Hybridization involving d-orbitals are not required.)
• Determine the polarity of molecules using molecular shape, understanding of symmetry, and comparison of the electronegativity of elements.
• Explain the relationship between vapour pressure, melting point, boiling point and solubility, and the nature and strength of intermolecular forces (e.g. dispersion forces, dipole-dipole attractions, and hydrogen bonding) within molecular covalent substances.

Chromatography techniques

• Identify that paper and thin layer chromatography can be used to determine the composition and purity of substances.
• Explain how variations in the strength of the interactions between atoms, molecules or ions in the mobile and stationary phases can be used to separate components.
• Analyse paper and thin layer (TLC) chromatographs to determine the composition and purity of substances, including calculating Rf values.

Gases

• State the relationship between the volume of a gas, number of moles and molar volume at standard temperature and pressure (STP).
• Apply the kinetic theory of gases to explain the relationships between pressure, temperature, and volume of a gas.
• Identify that the kinetic theory of gases applies to ideal gases.
• Apply the ideal gas equation to calculate the mass of chemicals and/or the volume of a gas (STP) involved in a chemical reaction. (Formula: PV = nRT)
• Analyse data to determine the relationships between pressure, temperature, and volume of a gas.

Science as a Human Endeavour (SHE)

• Appreciate that science relies on chemical processes to analyse materials in order to determine the identity, nature or source of the material.
• Explore how chromatography techniques, including gas and high-performance liquid chromatography, can be used to determine the composition and purity of substances.
• Appreciate that safe scuba diving requires knowledge of the behaviour of gases.
• Explore Jacques Cousteau and Emile Gagnan’s role in the invention of SCUBA.
• Appreciate that two- and three-dimensional graphical models have been developed and adopted by chemists to represent and communicate the shapes of molecules.
• Consider the limitations associated with the VSEPR theory.

Science inquiry

• Investigate: Boyle’s law or the molar volume of a gas
• Investigate: the separation of a mixture using paper or thin layer chromatography (TLC)
• Investigate: 3D models of linear, bent, trigonal planar, tetrahedral and pyramidal molecules

Aqueous solutions and molarity

• Explain that the unique properties of water are related to molecular shape and hydrogen bonding between molecules.
• Discriminate between the terms solute, solvent, solution.
• Discriminate between the terms strength and concentration, e.g. acidic/basic solutions.
• State that square brackets ([ ]) are used to denote concentration.
• Discriminate between unsaturated, saturated and supersaturated solutions.
• Apply the mole concept to calculate moles of solute, concentration and volume of a solution. (Formula: Molarity/Concentration (c) = moles of solute (n) / volume of solution (V))

Identifying ions in solution

• Apply ionic and chemical formulas to construct balanced ionic and chemical equations (including states) for precipitation reactions.
• Apply solubility rules to predict if a precipitation will be formed.
• Analyse data, including precipitation and acid-carbonate reactions, to determine the presence of specific ions in solutions.

Solubility

• Compare the solubility of ionic and molecular substance in water, and the intermolecular forces between species in the substances and water molecules.
• Identify that changes in solvent temperature can affect the solubility of solid and gaseous solutes (solids and gases).
• Analyse data, including solubility curves, to determine the solubility of ionic compounds and the concentration of ions in aqueous solutions.

pH

• State that pH is dependent on the concentration of hydrogen ions in solution.
• Identify that the pH scale is a logarithmic scale.
• Apply the pH scale to compare the levels of acidity or alkalinity of aqueous solutions.
• Apply the Arrhenius model to explain the behaviour of strong and weak acids and bases in aqueous solutions.

Reactions of acids

• Determine balanced chemical and ionic equation (including states) for the reactions of acids with bases, metals and carbonates.

Science as a Human Endeavour (SHE)

• Appreciate that most sulfur dioxide released to the atmosphere comes from burning coal or oil in electric power stations.
• Explore the chemistry of acid rain.
• Appreciate that blood plasma is an aqueous solution containing a range of ionic and molecular substances.
• Explore why blood is red and the chemistry of blood types.
• Appreciate that knowledge of the composition of water from different sources informs decisions about how that water is treated and used.
• Evaluate the measurable properties of water that are used to determine the water quality of a local water way.
• Explore the different water treatment methods used to provide safe drinking water.

Science inquiry

• Investigate: precipitation reactions to identify cations and anions
• Investigate: factors that affect solubility in aqueous solutions
• Investigate: reactions of acids with bases, metals and carbonates

Rates of reactions

• Explain how temperature, surface area, pressure (gaseous systems), concentration and the presence of a catalyst can affect the rate of the reaction.
• Apply the collision theory to determine the effect of concentration, temperature, pressure and surface area on the rate of chemical reactions.
• Sketch Maxwell-Boltzmann distribution curves for reactions with and without catalysts.
• Describe activation energy (Ea).
• Explain the relationship between the strength and number of the existing chemical bonds, the magnitude of the activation energy and the rate of a chemical reaction.
• Sketch energy profile diagrams for reactions with and without catalysts.
• Analyse energy profile diagrams for reactions with and without catalysts, to determine the enthalpy change and activation energy.
• Explain how catalysts affect the rate of a chemical reaction.
• Calculate the rate of chemical reactions by measuring the rate of formation of products or the depletion of reactants. (Formula: rate of reaction = increase in product concentration (Δ[P]) / time taken or decrease in reactant concentration (−Δ[R]) / time taken)
• Analyse data and graphical representations of relative changes in the concentration, volume and mass against time to determine rate of reaction. (Order of reaction is not required.)

Science as a Human Endeavour (SHE)

• Appreciate that catalysts work in a variety of ways, and knowledge of the structure of enzyme molecules helps scientists to explain and predict how they are able to lower the activation energy for reactions.
• Explore Mildred Cohn’s use of isotopic tracers and NMR spectroscopy to study the mechanism of enzymatic catalysis.
• Appreciate that most contemporary methods of corrosion prevention rely on knowledge of chemical and electrochemical redox processes.
• Explore the historical theories on corrosion from the introduction of iron in Antiquity through to the impact of air pollution on the life span of modern metal structures.
• Appreciate that collision theory enables chemists to explain and predict the rates of a vast range of chemical reactions in many different contexts.
• Explore the history of collision theory, its uses and its limitations.

Science inquiry

• Investigate: factors that affect the rate of chemical reactions

Chemical equilibrium

• Discriminate between open or closed chemical systems.
• Identify that physical changes are usually reversible, whereas only some chemical reactions are reversible.
• Symbolise equilibrium equations using ⇌ in balanced chemical equations.
• Explain observable properties and the characteristics of physical and chemical systems in a state of equilibrium.
• Explain that, over time, physical change and reversible chemical reactions reach a state of dynamic equilibrium in a closed system, with the relative concentrations of products and reactants defining the position of equilibrium.
• Explain the reversibility of chemical reactions by considering the activation energies of the forward and reverse reactions.
• Analyse data and interpret graphical representations of relative changes in the concentration of reactants and product against time, to determine the position of equilibrium.

Factors that affect equilibrium

• Determine the effect of temperature change on chemical systems at equilibrium by considering the enthalpy change for the forward and reverse reactions.
• Explain the effect of changes of temperature, concentration and pressure on chemical systems at equilibrium by applying collision theory to the forward and reverse reactions.
• Apply Le Châtelier’s principle to determine the effect changes of temperature, concentration of chemicals, pressure and the addition of a catalyst have on the position of equilibrium and on the value of the equilibrium constant.

Equilibrium constants

• Identify that the equilibrium constant (Kc) indicates the relationship between product and reactant concentrations at equilibrium.
• Identify that the solubility product (Ksp) gives a measure of the solubility of an ionic compound.
• Determine the equilibrium law expression for homogeneous and heterogeneous systems.
• Determine the extent of a reaction from the magnitude of the equilibrium constant (Kc).
• Calculate the reaction quotient (Q) for reversible reactions. (Formula: Q = [C]^c[D]^d / [A]^a[B]^b for the reaction aA + bB ⇌ cC + dD)
• Calculate equilibrium constants (Kc) and the concentrations of reactants and products. Assume [reactants]initial = [reactants]equilibrium when Kc is very small and state assumption when used. (Formula: Kc = [C]^c[D]^d / [A]^a[B]^b for the reaction aA + bB ⇌ cC + dD)
• Calculate solubility products (Ksp) and the concentrations of ions in aqueous solutions. (Formula: Ksp = [C]^c[D]^d for the reaction aA(s) ⇌ cC(aq) + dD(aq))
• Infer shifts in equilibrium reactions using equilibrium constants (Kc) and reaction quotients (Q).
• Analyse data to determine reaction quotients (Q), equilibrium constants (Kc), the concentrations of reactants and products and the concentration of ions in aqueous solutions.

Properties of acids and bases

• Identify that acids are substances that can act as proton (hydrogen ion) donors.
• Identify acids as monoprotic, diprotic or polyprotic.
• Identify hydrochloric, nitric and sulfuric acid as strong acids and group 1 hydroxides and barium hydroxide as strong bases.
• Identify carboxylic and carbonic acids as weak acids and ammonia and amines as weak bases.
• Discriminate between the terms strong, weak, concentrated and dilute for acids and bases.
• Discriminate between strong and weak acids and bases in terms of the extent of dissociation, rate of reaction, pH and electrical conductivity.
• Analyse data to determine the strength, concentration, pH and electrical conductivity of acids and bases.

pH

• Identify that water is a weak electrolyte and the self-ionisation of water is represented by Kw. (Formula: Kw = [H⁺][OH⁻])
• Apply Kw to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution.
• Calculate pH, hydrogen ion concentration [H⁺(aq)], pOH and hydroxide ion concentrations [OH⁻(aq)] for strong acids and bases. (Formula: pH = −log₁₀[H⁺] and pOH = −log₁₀[OH⁻])

Brønsted-Lowry model

• Describe acids and bases in equilibrium systems using the Brønsted-Lowry model.
• Explain the Brønsted-Lowry model using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs.
• Identify that amphiprotic species can act as Brønsted-Lowry acid (or base).
• Determine the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid).
• Identify that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added. (Buffer calculations are not required.)
• Apply Le Châtelier’s principle to explain how buffer solutions respond to the addition of hydrogen ions and hydroxide ions.

Dissociation constants

• Explain that the strength of acids is related to the degree of ionisation at equilibrium in aqueous solution.
• Identify that the strength of acids can be represented with chemical equations and equilibrium constants (Ka).
• Determine the expression for the dissociation constant for weak acids (Ka) and weak bases (Kb) from balanced chemical equations.
• Calculate dissociation constants (Ka, Kb, and Kw), pKa, pKb, and the concentrations of reactants and products. (Formula: Ka = [H₃O⁺][A⁻]/[HA]; Kb = [BH⁺][OH⁻]/[B]; Kw = Ka × Kb)
• Analyse data to compare the relative strengths of acids and bases.

Acid-base indicators

• Identify that acid-base indicators are a weak acid or a weak base where the conjugate acid-base pair have different colours and can be represented by HIn(aq) ⇌ H⁺(aq) + In⁻(aq) or BOH(aq) ⇌ B⁺(aq) + OH⁻(aq).
• Identify that indicators change colour when pH = pKa.
• Explain the relationship between the pH range, the end point and the pKa value of an acid-base indicator.
• Analyse data to determine an appropriate indicator given the equivalence point of the titration and the pH range of the indicator (assuming indicators change colour over a range of pKa ± 1).

Volumetric analysis

• Discriminate between the terms end point and equivalence point.
• Sketch the general shapes of conductometric and acid-base titration curves involving strong and weak acids and bases. (Titration of weak acids to weak bases is not required.)
• Interpret acid-base titration curves to determine the intercept with pH axis, equivalence point, buffer region and points where pKa = pH or pKb = pOH.
• Interpret conductometric titration curves to determine the intercept with conductivity axis, equivalence point and volume of titrant.
• Analyse volumetric data, including solubility, conductometric and acid-base titration curves, to determine moles, mass, volume and concentration.
• Analyse titration curves to calculate the concentration of a solution with reference to a standard solution.

Science as a Human Endeavour (SHE)

• Appreciate that the production of wine, along with that of many other food products, relies on the successful control of a range of reversible reactions in order to maintain the required chemical balance within the product.
• Explore the chemistry of wine.
• Appreciate that oceans contribute to the maintenance of steady concentrations of atmospheric carbon dioxide because the gas can dissolve in seawater through a range of reversible processes.
• Explore the absorption of heat by CO₂ and water vapour and the effect that changing amounts of CO₂ in the atmosphere have on climate.
• Explore Eunice Newton Foote’s contribution to understanding climate change in 1856.
• Appreciate that ‘superacids’, such as carborane acids, have been found to be a million times stronger than sulfuric acid when their strength is extrapolated to aqueous solutions the position of equilibrium is considered.
• Explore the composition of ‘superacids’ and what makes them so strong.

Science inquiry

• Investigate: factors that affect equilibrium (Le Châtelier’s principle)
• Investigate: solubility
• Investigate: properties of acids and bases
• Investigate: acid-base or conductometric titrations

Redox reactions

• Identify that displacement reactions of metals, combustion, corrosion and electrochemical processes, can be modelled as redox reactions involving oxidation of one substance and reduction of another substance.
• Determine the species oxidised and reduced, and the oxidising agent and reducing agent, in redox reactions.
• Explain that oxidation can be modelled as the loss of electrons from a chemical species, and reduction can be modelled as the gain of electrons by a chemical species; these processes can be represented using balanced half-equations and redox equations (acidic conditions only).
• Determine the oxidation state (represented with the sign given before the number) of an atom in an ion or compound, e.g. +2.
• Apply oxidation numbers (represented as roman numerals) to name transition metal compounds.
• Apply half-equations and oxidation numbers to balance redox equations (acid conditions only) and to discriminate between the species oxidised and reduced, and the oxidising agent and reducing agent.
• Analyse data, including displacement reactions of metals, combustion, corrosion and electrochemical processes to determine redox reactions.

Electrochemical cells

• Explain that electrochemical cells, including galvanic and electrolytic cells, consist of oxidation and reduction half-reactions connected via an external circuit that allows electrons to move from the anode (oxidation reaction) to the cathode (reduction reaction).
• Discriminate between a galvanic and an electrolytic cell.

Galvanic cells

• Identify that galvanic cells generate an electrical potential difference from a spontaneous redox reaction.
• Explain that galvanic cells can be represented as cell diagrams, including anode and cathode half-equations.
• Explain that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode).
• Explain that two half-cells can be connected by a salt bridge to create a galvanic cell, e.g. Mg, Zn, Fe and Cu and solutions of their ions.
• Identify the essential components of a galvanic cell, including the oxidation and reduction half-cells, the positive and negative electrodes and their solutions of their ions, the flow of electrons and the movement of ions, and the salt bridge.
• Sketch a galvanic cell and label the essential components.

Standard electrode potential

• Describe the standard hydrogen electrode.
• Explain the term standard electrode (reduction) potential, E°.
• Identify the limitations associated with standard electrode (reduction) potentials, E°.
• Calculate cell potential, E°cell. (Formula: E°cell = E°reduction half-cell − E°oxidation half-cell)
• Apply standard electrode potentials to determine the relative strength of oxidising and reducing agents.
• Analyse data, including standard electrode potentials, to make predictions about the spontaneity of a reaction and to compare electrochemical cells.

Electrolytic cells

• Identify that electrolytic cells use an external electrical potential difference to provide the energy to allow a non-spontaneous redox reaction to occur.
• Identify the essential components of an electrolytic cell, including source of electric current and conductors, positive and negative electrodes, and the electrolyte.
• State the factors that affect the products in electrolysis.
• Determine the products of the electrolysis of a molten salt.
• Explain the products of the electrolysis of aqueous solutions, e.g. dilute and concentration sodium chloride(aq) and copper sulfate(aq).
• Describe that electrolytic cells can be used in small-scale and industrial situations, including metal plating and the purification of copper.
• Calculate moles of electrons, current, time, mass of substance or volume of gas produced or used during electrolysis. (Formula: q = n(e⁻) × F or q = I × t).
• Analyse data to determine the relative amounts of product produced at each electrode in electrolysis.

Science as a Human Endeavour (SHE)

• Appreciate that the level of alcohol in the body can be measured by testing breath or blood alcohol concentrations.
• Explore the chemistry of breath tests for alcohol, drugs and disease.
• Appreciate that fuel cells are a potential lower-emission alternative to the internal combustion engine and are already being used to power buses, boats, trains and cars.
• Explore battery technologies that could power the future.
• Appreciate that electrochemistry has a wide range of uses, ranging from industrial scale metal extraction to personal cosmetic treatments.
• Explore the desalination process to produce fresh water.

Science inquiry

• Investigate: displacement reactions
• Investigate: galvanic cells
• Investigate: factors that affect electrolysis
• Investigate: electroplating using an electrolytic cell

Structure of organic compounds

• Identify organic molecules including alkanes, alkenes, alkynes, alcohols, aldehydes, ketones, carboxylic acids, haloalkanes, esters, amines and amides.
• Discriminate between class and functional groups, e.g. for OH, hydroxyl is the functional group and alcohol is the class.
• Describe the features of a homologous series.
• Discriminate between saturated and unsaturated organic molecules.
• Discriminate between empirical, molecular and structural formulas.
• Determine molecular and structural formulas for organic compounds, up to C₁₀, including simple methyl and ethyl branched chains, for: alkanes, alkenes and alkynes; alcohols (primary, secondary and tertiary); aldehydes and ketones; carboxylic acids; amines and amides; haloalkanes (primary, secondary and tertiary); esters.
• Apply IUPAC rules in the nomenclature of organic compounds, up to C₁₀, including simple methyl and ethyl branched chains, for: alkanes, alkenes and alkynes; alcohols (primary, secondary and tertiary); aldehydes and ketones; carboxylic acids; haloalkanes (primary, secondary and tertiary); esters.
• Identify structural and stereoisomers, including geometrical (cis and trans) and optical isomers.
• Deduce the structural formula of geometrical (cis and trans) isomers (non-cyclic alkenes), optical isomers and isomers of the non-cyclic alkanes up to C₆.
• Sketch the structural formula and apply IUPAC rules in the nomenclature for isomers of alkanes (non-cyclic) and alkenes (straight chain) up to C₆, and for the geometrical (cis and trans) isomers of simple alkenes (non-cyclic).
• Determine the structural formula of optical isomers for simple organic compounds.
• Identify chiral carbon atoms.
• Analyse data to determine the structural, molecular and empirical formula of organic compound and the percentage composition of elements in organic compounds.

Physical properties and trends

• Explain the trends (melting point, boiling point, volatility, solubility in water and organic solvents) within and between homologous series (alkanes, alkenes, alcohols, carboxylic acids) in term of intermolecular and intramolecular bonding, e.g. dispersion forces, dipole-dipole interactions and hydrogen bonds.
• Analyse data to determine the physical properties of an homologous series, trends in melting point, boiling point, volatility and the solubility alkanes, alkenes, alcohols and carboxylic acids.

Organic reactions and reaction pathways

• Identify that an organic compound displays characteristic chemical properties and undergoes specific reactions based on the functional group present.
• Determine, using equations, the reaction of: alkanes with halogens (X₂); haloalkanes with halogens (X₂), sodium hydroxide and ammonia; alkenes with water, halogens(X₂), hydrogen (H₂) and hydrogen halides (HX); alcohols with hydrogen halides (HX); carboxylic acid with alcohol to form esters, and with amines to form amides.
• Determine, using equations, reactions including the: oxidation of alcohols; combustion of alkanes and alcohols; addition of alkenes to form poly(alkenes); reduction of alkynes and alkenes to form alkanes; elimination of haloalkanes to form alkenes.
• Identify reactions as addition, elimination, substitution or redox (oxidation-reduction). (Reaction mechanism for substitution and elimination reactions are not required.)
• Determine the primary, secondary and tertiary carbon atoms in haloalkanes and alcohols.
• Describe the acid-base properties of carboxylic acids and amines.
• Explain that esterification is a reversible reaction.
• Discriminate between: alkanes and alkenes using bromine water; primary, secondary and tertiary alcohols using acidified potassium dichromate (VI) and potassium manganate (VII).
• Apply Markovnikov’s rule to determine the products for addition reactions of alkenes with hydrogen halides (HX) and water.
• Determine reaction pathways, including reagents, condition and chemical equations, given the starting materials and the product/s formed.
• Interpret chemical tests to distinguish between alkanes and alkenes; and primary, secondary and tertiary alcohols.

Organic materials: structure and function

• Describe the structural features of: amino acids, tripeptides, monosaccharides and disaccharides; polyethene (LDPE and HDPE), polypropene (syndiotactic, isotactic and atactic) and polytetrafluoroethene (Teflon); polylactic acid (PLA), polyamide (nylon) and polyester.
• Explain how properties, including strength, density and biodegradability of polymers can be related to the structures of the materials.
• Explain the acid-base properties of 2-amino acids, including the formation of zwitterions.

Analytical techniques

• Explain how amino acids can be separated and identified by paper/TLC chromatography, including intermolecular forces/solubility in mobile and stationary phase and retention (RF) values.
• Explain how amino acids can be separated and analysed by electrophoresis, including pH of buffer, isoelectric points, and movement of charged ions.
• Analyse data, including paper/TLC chromatograms and electrophoresis to determine the identity of amino acids and retention factors. (Formula: RF = distance moved by the amino acid / distance moved by the solvent)
• Analyse data from spectra, including mass spectroscopy and infrared to determine the identity and structure of organic molecules.

Science as a Human Endeavour (SHE)

• Appreciate that the developments in computer modelling enabled more accurate visualisation and prediction of three-dimensional organic structures, such as proteins, which is critical in drug design and biotechnology.
• Recognise that enzymes are proteins and explore characteristics of biological catalysts (enzymes), including activity, depends on the structure and the specificity of the enzyme action.
• Consider that triglycerides (lipids) are esters, and that saturated and unsaturated fatty acids have difference structures and properties.
• Appreciate that synthetic polymers often have large ‘ecological footprints’ as they are synthesised from fossil fuels and do not biodegrade. Therefore, sustainable polymers, produced from renewable sources such as plants, waste products and waste gases are ‘greener’.
• Explore the principles of green chemistry and recognise that the higher the atom economy, the ‘greener’ the process.
• Appreciate that organochlorine compounds, such as DDT, chlordane and lindane, were identified as powerful insecticides in the 1950s because their structure makes them chemically unreactive.
• Explore the relationship between the chemical structure of insecticides, their effectiveness as an insecticide, and their persistent and bioaccumulation in the environment.

Science inquiry

• Investigate: properties of homologous series
• Investigate: 3D models of organic molecules
• Investigate: paper/TLC chromatography to separate amino acids
• Investigate: electrophoresis to separate amino acids
• Investigate: mass spectroscopy and infrared spectra

Chemical synthesis

• Explain that reagents and reaction conditions are chosen to optimise the yield and rate for chemical synthesis processes, including the production of ammonia (Haber process) and sulfuric acid (contact process).
• Describe, using equations, the: production of ammonia by the Haber process; production of sulfuric acid using the contact process; production of ethanol from fermentation and the hydration of ethene; operation of a hydrogen fuel cell under acidic and alkaline conditions.
• Calculate the yield of chemical synthesis reactions by comparing stoichiometric quantities with actual quantities and by determining limiting reagents and/or reaction conditions.
• Analyse and interpret data to determine the impact of reagents and reaction conditions on yield and rate of chemical synthesis processes.

Macromolecules: polymers, proteins and carbohydrates

• Describe, using equations, how: addition polymers, including polyethene (LDPE and HDPE), polypropene and polytetrafluoroethene, can be produced from their monomers; condensation polymers, including polysaccharides (carbohydrates), polylactic acid (PLA), polyamide (proteins and nylon) and polyester, can be produced from their monomers.
• Apply amino acid symbols to construct and name tripeptides.
• Identify that tripeptides are formed when amino acid monomers are joined by peptide bonds.
• Identify that disaccharides are formed when monosaccharides monomers are joined by glycosidic bonds.

Science as a Human Endeavour (SHE)

• Appreciate that green chemistry aims to increase the atom economy of chemical processes by designing novel reactions that can maximise the desired products and minimise by-products.
• Explore important developments in sustainable chemical industries, such as new synthetic schemes that can simplify operations in chemical productions and greener solvents that are inherently environmentally and ecologically benign.
• Consider the principles of green chemistry, including the design of chemical synthesis processes that use renewable raw materials, limiting the use of potentially harmful solvents and minimising the amount of unwanted products. Atom economy can be calculated and used to draw conclusions about the economic and environmental impact of chemical synthesis processes.
• Appreciate that dwindling supplies of economically viable sources of fossil fuels and concerns related to carbon emissions have prompted research into the synthesis of biofuels.
• Explore the development of biofuels from plant feedstocks, such as algae, oil seeds and wood waste, or from waste materials, such as food industry waste oils.
• Appreciate that molecular manufacturing (or molecular assembly) involves building objects to atomic precision using robotic mechanisms to position and react molecules and has the potential to quickly develop products (such as stronger materials, and smaller, faster and more energy-efficient computers) and address a range of global issues through provision of vital materials and products at a greatly reduced cost and environmental impact.
• Explore how enzymes can be used on an industrial scale for chemical synthesis to achieve an economically viable rate, including fermentation to produce ethanol and lipase-catalysed transesterification to produce biodiesel.
• Consider that molecular manufacturing processes involve the positioning of molecules to facilitate a specific chemical reaction; such methods have the potential to synthesise specialised products, including proteins, carbon nanotubes, nanorobots and chemical sensors used in medicine.

Science inquiry

• Investigate: the Haber and contact processes
• Investigate: the properties of polymers