Learning Goal Breakdown

Unit 3: Equilibrium, acids and redox reactions

Topic 1: Chemical equilibrium systems

Chemical Equilibrium

• Recognise that chemical systems may be open (allowing matter and energy to be exchanged with the surroundings) or closed (allowing energy, but not matter, to be exchanged with the surroundings).
• Understand that physical changes are usually reversible, whereas only some chemical reactions are reversible.
• Appreciate that observable changes in chemical reactions and physical changes can be described and explained at an atomic and molecular level.
• Symbolise equilibrium equations by using ⇌ in balanced chemical equations.
• Understand that, over time, physical changes 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 experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration of reactants and products against time, to identify the position of equilibrium.

Factors That Affect Equilibrium

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

Equilibrium Constants

• Understand that equilibrium law expressions can be written for homogeneous and heterogeneous systems and that the equilibrium constant (Kc), at any given temperature, indicates the relationship between product and reactant concentrations at equilibrium.
• Deduces the equilibrium law expression from the equation for a homogeneous reaction and use equilibrium constants (Kc) to predict qualitatively the relative amounts of reactants and products (equilibrium position).
• Deducing the extent of a reaction from the magnitude of the equilibrium constant.
• Use appropriate mathematical representation to solve problems, including calculating equilibrium constants and the concentration of reactants and products.
• Formula: \( K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b} \) for the reaction: \( aA + bB ⇌ cC + dD \).
• Students should state when assumptions are used, such as assuming [reactants]initial ≈ [reactants]equilibrium when Kc is very small.

Properties of Acids and Bases

• Understand that acids are substances that can act as proton (hydrogen ion) donors and can be classified as monoprotic or polyprotic depending on the number of protons donated by each molecule of the acid.
• Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water, and electrical conductivity.
• Distinguish between the terms strong and concentrated for acids and bases.

pH Scale

• Understand that water is a weak electrolyte and the self-ionisation of water is represented by \( K_w = [H^+][OH^-] \); Kw can be used to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution.
• Understand that the pH scale is a logarithmic scale and the pH of a solution can be calculated from the concentration of hydrogen ions using the relationship \( pH = -\log_{10} [H^+] \).
• Use appropriate mathematical representation to solve problems for hydrogen ion concentration \([H^+ (aq)]\), pH, hydroxide ion concentrations \([OH^- (aq)]\), and pOH.
• Formulas:
  • \( K_w = [H^+][OH^-] \)
  • \( pH = -\log_{10}[H^+] \)
  • \( pOH = -\log_{10}[OH^-] \)

Brønsted-Lowry Model

• Recognise that the relationship between acids and bases in equilibrium systems can be explained using the Brønsted-Lowry model and represented using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs.
• Recognise that amphiprotic species can act as Brønsted-Lowry acids and bases.
• Identify and deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid).
• Appreciate 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; Le Châtelier’s principle can be applied to predict how buffer solutions respond to the addition of hydrogen ions and hydroxide ions.
• Buffer calculations are not required.

Dissociation Constants

• Recognise that the strength of acids is explained by the degree of ionisation at equilibrium in aqueous solution, which 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.
• Analyse experimental data to determine and compare the relative strengths of acids and bases.
• Use appropriate mathematical representation to solve problems, including calculating dissociation constants (Ka and Kb) and the concentration of reactants and products.
• Students should consider:
  • Hydrochloric, nitric, and sulfuric acids as examples of strong acids.
  • Carboxylic and carbonic acids (aqueous carbon dioxide) as weak acids.
  • All group 1 hydroxides and barium hydroxide as strong bases.
  • Ammonia and amines as weak bases.
• Formulas:
  • Kw = Ka × Kb
  • Ka = [H₃O⁺][A^–] / [HA]
  • Kb = [BH⁺][OH^–] / [B]

Acid-Base Indicators

• Understand that an acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours; the acidic form is of a different colour to the basic form.
• Explain the relationship between the pH range of an acid-base indicator and its pKa value.
• Recognise that indicators change colour when the pH = pKa and identify an appropriate indicator for a titration, given the equivalence point of the titration and pH range of the indicator.
• For an indicator that is a weak acid:
  • HIn(aq) ⇌ H⁺ (aq) + In^–(aq), Colour A → Colour B
• For an indicator that is a weak base:
  • BOH(aq) ⇌ B⁺(aq) + OH^–(aq), Colour A → Colour B
• The colour change can be considered to take place over a range of pKa ± 1.
• Examples of indicators and their pKa values are listed in the Chemistry formula and data booklet.

Volumetric Analysis

• Distinguish between the terms end point and equivalence point.
• Recognise that acid-base titrations rely on the identification of an equivalence point by measuring the associated change in pH, using chemical indicators or pH meters, to reveal an observable end point.
• Sketch the general shapes of graphs of pH against volume (titration curves) involving strong and weak acids and bases, identifying and explaining their important features:
  • Intercept with pH axis.
  • Equivalence point.
  • Buffer region.
  • Points where pKa = pH or pKb = pOH.
• Use appropriate mathematical representations and analyse experimental data and titration curves to solve problems and make predictions, including using the mole concept to calculate moles, mass, volume, and concentration from volumetric analysis data.
• Mandatory practical: Acid-base titration to calculate the concentration of a solution with reference to a standard solution.
• Titration of weak acid to weak base is not required.

Topic 2: Oxidation and Reduction

Redox Reactions

• Recognise that a range of reactions, including 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.
• Understand that the ability of an atom to gain or lose electrons can be predicted from the atom’s position in the periodic table, and explained with reference to valence electrons, energy considerations, and overall stability.
• Identify the species oxidised and reduced, and the oxidising and reducing agents in redox reactions.
• Understand that:
  • Oxidation can be modelled as the loss of electrons from a chemical species.
  • Reduction can be modelled as the gain of electrons by a chemical species.
• Represent these processes using balanced half-equations and redox equations (acidic conditions only).
• Deduces oxidation states of atoms in ions or compounds and name transitional metal compounds using oxidation numbers (Roman numerals).
• Use appropriate representations, including half-equations and oxidation numbers, to communicate understanding, solve problems, and make predictions.

Mandatory practical: Perform single displacement reactions in aqueous solutions.

Additional Notes:

• Oxidation numbers and oxidation states are often interchanged. Use Roman numerals for oxidation numbers.
• Oxidation states should be represented with the sign before the number (e.g., +2, not 2+).
• Cover specific cases such as hydrogen in metal hydrides (-1) and oxygen in peroxides (-1).
• A simple activity series is provided in the Chemistry formula and data booklet.

Electrochemical Cells

• Understand 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) to the cathode (reduction).

Galvanic Cells

• Understand that galvanic cells, including fuel cells, generate an electrical potential difference from a spontaneous redox reaction, represented as cell diagrams with anode and cathode half-equations.
• Recognise:
  • Oxidation occurs at the negative electrode (anode).
  • Reduction occurs at the positive electrode (cathode).
• Explain how two half-cells are connected by a salt bridge to create a voltaic cell (examples include Mg, Zn, Fe, and Cu with their ion solutions).
• Describe, using a diagram, the essential components of a galvanic cell, including:
  • Oxidation and reduction half-cells.
  • Positive and negative electrodes and their ion solutions.
  • Electron flow and ion movement.
  • The salt bridge.

Mandatory practical: Construct a galvanic cell using two metal/metal-ion half-cells.

Standard Electrode Potential

• Determine the relative strength of oxidising and reducing agents by comparing standard electrode potentials.
• Recognise that cell potentials at standard conditions can be calculated from standard electrode potentials, and these values can be used to compare cells constructed from different materials.
• Recognise limitations associated with standard reduction potentials.
• Use appropriate mathematical representations to solve problems and make predictions about spontaneous reactions, including calculating cell potentials under standard conditions.

Note: A table of standard reduction potentials is provided in the Chemistry formula and data booklet.

Electrolytic Cells

• Understand that electrolytic cells use an external electrical potential difference to drive a non-spontaneous redox reaction.
• Appreciate that electrolytic cells have industrial applications, including metal plating and copper purification.
• Predict and explain the products of the electrolysis of:
  • Molten salts.
  • Aqueous solutions of sodium chloride and copper sulphate.
• Explanations should consider:
  • E° values.
  • The nature and concentration of the electrolyte.
• Describe, using a diagram, the essential components of an electrolytic cell, including:
  • Source of electric current and conductors.
  • Positive and negative electrodes.
  • The electrolyte.

Note: Products of dilute and concentrated solutions of sodium chloride and copper sulphate should be considered.

Unit 4: Structure, synthesis and design

Topic 1: Properties and Structure of Organic Materials

Structure of Organic Compounds

• Recognise that organic molecules have a hydrocarbon skeleton and can contain functional groups, including alkenes, alcohols, aldehydes, ketones, carboxylic acids, haloalkanes, esters, nitriles, amines, and amides.
• Use structural formulas (condensed and extended) to show the arrangement of atoms and bonding in organic molecules.
• Deduces structural formulas and applies IUPAC rules in the nomenclature of organic compounds (parent chain up to 10 carbon atoms) with simple branching for alkanes, alkenes, alkynes, alcohols, aldehydes, ketones, carboxylic acids, haloalkanes, esters, nitriles, amines, and amides.
• Identify structural isomers as compounds with the same molecular formula but different arrangements of atoms; deduce structural formulas and apply IUPAC rules for isomers of non-cyclic alkanes up to C6.
• Identify stereoisomers as compounds with the same structural formula but different arrangements of atoms in space; describe and explain geometrical (cis and trans) isomerism in non-cyclic alkenes.

Mandatory practical: Construct 3D models of organic molecules.

Physical Properties and Trends

• Recognise that organic compounds display characteristic physical properties, including melting point, boiling point, and solubility in water and organic solvents. These properties can be explained in terms of intermolecular forces (dispersion forces, dipole-dipole interactions, and hydrogen bonds) influenced by the nature of the functional groups.
• Predict and explain trends in melting and boiling points for members of a homologous series.
• Discuss the volatility and solubility in water of alcohols, aldehydes, ketones, carboxylic acids, and halides.

Additional guidance: Consider physical properties of hydrocarbons, alcohols, aldehydes, ketones, carboxylic acids, amines, amides, and esters.

Organic Reactions and Reaction Pathways

• Appreciate that each class of organic compound displays characteristic chemical properties and undergoes specific reactions based on the functional group present; these reactions, including acid-base and oxidation reactions, can be used to identify the class of the organic compound.
• Understand that saturated compounds contain single bonds only and undergo substitution reactions, and that unsaturated compounds contain double or triple bonds and undergo addition reactions.
• Determine the primary, secondary, and tertiary carbon atoms in haloalkanes and alcohols and apply IUPAC rules of nomenclature.
• Describe, using equations:
  • Oxidation reactions of alcohols and the complete combustion of alkanes and alcohols.
  • Substitution reactions of alkanes with halogens.
  • Substitution reactions of haloalkanes with halogens, sodium hydroxide, ammonia, and potassium cyanide.
  • Addition reactions of alkenes with water, halogens, and hydrogen halides.
  • Addition reactions of alkenes to form poly(alkenes).
• Recall the acid-base properties of carboxylic acids and explain, using equations, that esterification is a reversible reaction between an alcohol and a carboxylic acid.
• Recognise the acid-base properties of amines and explain, using equations, the reaction with carboxylic acids to form amides.
• Recognise reduction reactions and explain, using equations, the reaction of nitriles to form amines and alkenes to form alkanes.
• The distinction between class and functional group should be made, e.g., for OH, hydroxyl is the functional group whereas alcohol is the class.
• Students are not required to recall reaction mechanisms for substitution and elimination reactions.
• Addition reactions with alkenes:
  • Reactions with H2, Br2, H2O, and HBr (Markovnikov’s rule) should be covered.
• Recognise and explain, using equations, that:
  • Esters and amides are formed by condensation reactions.
  • Elimination reactions can produce unsaturated molecules and explain, using equations, the reaction of haloalkanes to form alkenes.
• Understand that organic reactions can be identified using characteristic observations and recall tests to distinguish between:
  • Alkanes and alkenes using bromine water.
  • Primary, secondary, and tertiary alcohols using acidified potassium dichromate (VI) and potassium manganate (VII).
• Understand that the synthesis of organic compounds often involves constructing reaction pathways that may include more than one chemical reaction.
• Deduce reaction pathways, including reagents, conditions, and chemical equations, given the starting materials and the product.

Organic Materials: Structure and Function

• Appreciate that organic materials, including proteins, carbohydrates, lipids, and synthetic polymers, display properties such as strength, density, and biodegradability that can be explained by considering their primary, secondary, and tertiary structures.
• Describe and explain the primary, secondary (α-helix and β-pleated sheets), tertiary, and quaternary structures of proteins.
• Recognise that enzymes are proteins and describe the characteristics of biological catalysts, including specificity and activity dependent on structure.
• Recognise that monosaccharides contain either an aldehyde group (aldose) or a ketone group (ketose) and several -OH groups, and have the empirical formula CH₂O.
• Distinguish between α-glucose and β-glucose, and compare the structural properties of starch (amylose and amylopectin) and cellulose.

Triglycerides and Fatty Acids

• Recognise that triglycerides (lipids) are esters.
• Describe the difference in structure between saturated and unsaturated fatty acids.
• Describe, using equations, the base hydrolysis (saponification) of fats (triglycerides) to produce glycerol and its long-chain fatty acid salt (soap).
• Explain how the cleaning action and solubility in hard water are related to their chemical structure.

Properties of Polymers

• Explain how the properties of polymers depend on their structural features, including:
  • The degree of branching in polyethene (LDPE and HDPE).
  • The position of the methyl group in polypropene (syndiotactic, isotactic, and atactic).
  • Polytetrafluoroethene (PTFE).

Analytical Techniques

• Explain how proteins can be analysed by chromatography and electrophoresis.
• Select and use data from analytical techniques, including:
  • Mass spectrometry.
  • X-ray crystallography.
  • Infrared spectroscopy.
• Analyse data from spectra, including mass spectrometry and infrared spectroscopy, to:
  • Communicate conceptual understanding.
  • Solve problems.
  • Make predictions.

Topic 2: Chemical synthesis and design

Chemical Synthesis

• Understand that reagents and reaction conditions are chosen to optimise yield and rate for processes, including:
  • Ammonia production (Haber process).
  • Sulfuric acid production (contact process).
  • Biodiesel production (base-catalysed and lipase-catalysed methods).
• Understand that fuels, including biodiesel, ethanol, and hydrogen, can be synthesised through:
  • Addition reactions.
  • Oxidation.
  • Esterification.
• Describe, using equations, the production of ethanol from fermentation and the hydration of ethene.
• Describe, using equations, the transesterification of triglycerides to produce biodiesel.
• Discuss, using diagrams and relevant half-equations, the operation of a hydrogen fuel cell under acidic and alkaline conditions.
• Calculate yields of chemical synthesis reactions by comparing stoichiometric and actual quantities.

Green Chemistry

• Appreciate that green chemistry principles include the design of chemical synthesis processes that use renewable raw materials, limit the use of potentially harmful solvents, and minimise the amount of unwanted products.
• Outline the principles of green chemistry and recognise that the higher the atom economy, the 'greener' the process.
• Calculate atom economy and draw conclusions about the economic and environmental impact of chemical synthesis processes.
• 100% atom economy equates to all the atoms in the reactants being converted to the desired product.

Macromolecules: Polymers, Proteins, and Carbohydrates

• Describe, using equations, how addition polymers can be produced from their monomers including polyethene (LDPE and HDPE), polypropene, and polytetrafluorethene.
• Describe, using equations, how condensation polymers, including polypeptides (proteins), polysaccharides (carbohydrates), and polyesters, can be produced from their monomers.
• Discuss the advantages and disadvantages of polymer use, including strength, density, lack of reactivity, use of natural resources, and biodegradability.
• Describe the condensation reaction of 2-amino acids to form polypeptides (involving up to three amino acids), and understand that polypeptides (proteins) are formed when amino acid monomers are joined by peptide bonds.
• Describe the condensation reaction of monosaccharides to form disaccharides (lactose, maltose, and sucrose) and polysaccharides (starch, glycogen, and cellulose), and understand that polysaccharides are formed when monosaccharide monomers are joined by glycosidic bonds.
• The common names, symbol, structural formula, and pH of the isoelectric point for amino acids are given in the Chemistry data booklet.

Molecular Manufacturing

• Appreciate 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.