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The National Curriculum of England has a clearly defined series of academic and other objectives at every level. mydrasa focuses on Key stage 3 (Year 7-9), Key stage 4 IGCSE/GCSE (Year 10-11) and Key stage 5 A-Level (Year 12-13).

mydrasa added subjects related to Key stage 4 to Year 9, and added subjects related to Key stage 5 to Year 11 for student preparation.

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SUbjects

Subjects

Edexcel - Chemistry - 9CH0

  • Overview
  • Chapters

Chemistry 9CH0 is the Pearson Edexcel Level 3 Advanced GCE in Chemistry.

The Pearson Edexcel Level 3 Advanced GCE in Chemistry consists of three externally examined papers and the Science Practical Endorsement.

Students are expected to carry out the sixteen core practical experiments that are identified in the topics.

Component 1: Advanced Inorganic and Physical Chemistry.

Component 2: Advanced Organic and Physical Chemistry.

Component 3: General and Practical Principles in Chemistry.

Component 4: Science Practical Endorsement.



  • 1: Atomic Structure and the Periodic Table
    1.1: Atomic Structure and the Periodic Table
    1.1.1: The structure of an atom
    1.1.2: The relative mass and relative charge of protons, neutrons and electrons
    1.1.3: Atomic (proton) number’ and ‘mass number’
    1.1.4: The number of each type of sub-atomic particle in an atom, molecule or ion
    1.1.5: Understand the term ‘isotopes’
    1.1.6: ‘relative isotopic mass’ and ‘relative atomic mass’
    1.1.7: ‘relative molecular mass’ and ‘relative formula mass’
    1.1.8: Analyse and interpret data from mass spectrometry
    1.1.9: Predict the mass spectra
    1.1.10: Determine the relative molecular mass of a molecule
    1.1.11: ‘first ionisation energy’ and ‘successive ionisation energies’
    1.1.12: Ionisation energies are influenced by the number of protons
    1.1.13: Reasons for the general increase in first ionisation energy across a period
    1.1.14: Reasons for the decrease in first ionisation energy down a group
    1.1.15: Ideas about electronic configuration
    1.1.16: know the number of electrons that can fill the first four quantum shells
    1.1.17: Orbital is a region within an atom that can hold up to two electrons
    1.1.18: know the shape of an s-orbital and a p-orbital
    1.1.19: The number of electrons that occupy s, p and d-subshells
    1.1.20: Electrons fill sub-shells singly, before pairing up
    1.1.21: Predict the electronic configurations
    1.1.22: Elements can be classified as s, p and d-block elements
    1.1.23: Electronic configuration determines the chemical properties of an element
    1.1.24: Periodicity in terms of a repeating pattern across different periods
    1.1.25: The elements from periods 2 and 3 of the Periodic Table
    1.1.26: Illustrate periodicity using data
  • 2: Bonding and Structure
    2.1: Bonding
    2.1.1: Ionic bonding is the strong electrostatic attraction
    2.1.2: The effects that ionic radius and ionic charge
    2.1.3: The formation of ions
    2.1.4: Draw electronic configuration diagrams of cations and anions
    2.1.5: Reasons for the trends in ionic radii down a group
    2.1.6: The physical properties of ionic compounds
    2.1.7: Covalent bond is the strong electrostatic attraction between two nuclei
    2.1.8: Dot-and-cross diagrams
    2.1.9: The relationship between bond lengths and bond strengths for covalent bonds
    2.1.10: The repulsion between the electron pairs that surround a central atom
    2.1.11: Shapes of simple molecules and ions with up to six outer pairs of electrons
    2.1.12: Predict the shapes of, and bond angles in, simple molecules and ions
    2.1.13: Electronegativity is the ability of an atom to attract the bonding electrons
    2.1.14: Ionic and covalent bonding are the extremes of a continuum of bonding type
    2.1.15: Molecules with polar bonds may not be polar molecules
    2.1.16: The nature of intermolecular forces
    2.1.17: The interactions in molecules
    2.1.18: Properties of water resulting from hydrogen bonding
    2.1.19: Predict the presence of hydrogen bonding in molecules
    2.1.20: Intermolecular forces and physical properties
    2.1.21: Factors that influence the choice of solvents
    2.1.22: Metallic bonding is the strong electrostatic attraction between metal ions
    2.2: Structure
    2.2.1: Giant lattices
    2.2.2: The structure of covalently bonded substances
    2.2.3: The different structures formed by carbon atoms
    2.2.4: Predict the type of structure and bonding
    2.2.5: Predict the physical properties of a substance
  • 3: Redox I
    3.1: Redox I
    3.1.1: know what is meant by the term ‘oxidation number’
    3.1.2: The oxidation number of elements in compounds and ions
    3.1.3: Oxidation and reduction
    3.1.4: Oxidation and reduction in terms of electron loss or electron gain
    3.1.5: Oxidising agents gain electrons
    3.1.6: Reducing agents lose electrons
    3.1.7: Disproportionation reaction involves an element in a single species
    3.1.8: Oxidation number is a useful concept
    3.1.9: Indicate the oxidation number of an element in a compound or ion
    3.1.10: Formulae given oxidation numbers
    3.1.11: Metals form positive ions by loss of electrons
    3.1.12: Non-metals form negative ions by gain of electrons
    3.1.13: Ionic half-equations
  • 4: Inorganic Chemistry and the Periodic Table
    4.1: The elements of Groups 1 and 2
    4.1.1: Reasons for the trend in ionisation energy down Group 2
    4.1.2: Reasons for the trend in reactivity of the Group 2 elements down the group
    4.1.3: know the reactions of the elements Mg to Ba
    4.1.4: The reactions of the oxides of Group 2 elements with water and dilute acid
    4.1.5: The trends in solubility of the hydroxides
    4.1.6: Reasons for the trends in thermal stability of the nitrates and the carbonates
    4.1.7: The formation of characteristic flame colours by Group 1 and 2 compounds
    4.1.8: Understand experimental procedures
    4.1.9: Trends in melting and boiling temperatures
    4.1.10: The trend in reactivity of Group 7 elements
    4.1.11: The redox reactions of Cl2, Br2 and I2 with halide ions in aqueous solution
    4.1.12: Reactions of the halogens
    4.1.13: Understand the reactions of halides
    4.1.14: Predictions about fluorine and astatine and their compound
    4.2: The elements of Group 7 (halogens)
    4.3: Analysis of inorganic compounds
    4.3.1: Ionic equations
  • 5: Formulae, Equations and Amounts of Substance
    5.1: Formulae, Equations and Amounts of Substance
    5.1.1: The mole (mol) is the unit for amount of a substance
    5.1.2: Use the Avogadro constant
    5.1.3: The molar mass of a substance is the mass per mole of the substance in g mol-1
    5.1.4: ‘empirical formula’ and ‘molecular formula’
    5.1.5: Use experimental data
    5.1.6: Write balanced full and ionic equations
    5.1.7: Calculate amounts of substances (in mol) in reactions
    5.1.8: Calculate reacting masses from chemical equations
    5.1.9: calculating reacting volumes of gases from chemical equations
    5.1.10: Molar volume of gases
    5.1.11: Measure the molar volume of a gas
    5.1.12: Solution concentrations
    5.1.13: Prepare a standard solution from a solid acid
    5.1.14: The concentration of a solution of hydrochloric acid
    5.1.15: Calculate measurement uncertainties and measurement errors
    5.1.16: Minimise the percentage error and percentage uncertainty in experiments
    5.1.17: Calculate percentage yields and percentage atom economies
    5.1.18: Relate ionic and full equations to observations from simple test tube reactions
    5.1.19: Risks and hazards in practical procedures
  • 6: Organic Chemistry I
    6.1: Introduction to organic chemistry
    6.1.1: hydrocarbon is a compound of hydrogen and carbon only
    6.1.2: Represent organic molecules
    6.1.3: Homologous series’ and ‘functional group’
    6.1.4: The rules of International Union of Pure and Applied Chemistry (IUPAC)
    6.1.5: Addition, elimination, substitution, oxidation, reduction, hydrolysis, etc
    6.1.6: ‘structural isomerism’
    6.1.7: ‘stereoisomerism’
    6.2: Alkanes
    6.2.1: The general formula for alkanes
    6.2.2: Alkanes and cycloalkanes are saturated hydrocarbons
    6.2.3: Alkane fuels are obtained from crude oil
    6.2.4: pollutants are formed during the combustion of alkane fuels
    6.2.5: The problems arising from pollutants from the combustion of fuels
    6.2.6: The use of a catalytic converter solves some problems caused by pollutants
    6.2.7: The use of alternative fuels
    6.2.8: A radical
    6.2.9: The reactions of alkanes
    6.2.10: The use of radical substitution reactions in the synthesis of organic molecules
    6.3: Alkenes
    6.3.1: know the general formula for alkenes
    6.3.2: Alkenes and cycloalkenes
    6.3.3: The bonding in alkenes in terms of σ- and π- bonds
    6.3.4: ‘electrophile’
    6.3.5: The addition reactions of alkenes
    6.3.6: Heterolytic bond fission of a covalent bond results in the formation of ions
    6.3.7: The mechanism of the electrophilic addition reactions between alkenes
    6.3.8: The qualitative test for a C=C double bond using bromine or bromine water
    6.3.9: Alkenes form polymers through addition polymerisation
    6.3.10: Waste polymers can be separated into specific types of polymer
    6.3.11: The use of energy and resources over the life cycle of polymer products
    6.3.12: How chemists limit the problems caused by polymer disposal
    6.4: Halogenoalkanes
    6.4.1: Halogenoalkanes can be classified as primary, secondary or tertiary
    6.4.2: Nucleophile
    6.4.3: The reactions of halogenoalkanes
    6.4.4: Compare the relative rates of hydrolysis
    6.4.5: The rates of hydrolysis of some halogenoalkanes
    6.4.6: The trend in reactivity of primary, secondary and tertiary halogenoalkanes
    6.4.7: The trend in reactivity of chloro-, bromo-, and iodoalkanes
    6.4.8: The mechanisms of the nucleophilic substitution reactions
    6.5: Alcohols
    6.5.1: Alcohols can be classified as primary, secondary or tertiary
    6.5.2: The reactions of alcohols
    6.5.3: The preparation and purification of a liquid organic compound
    6.5.4: The oxidation of ethanol
    6.5.5: Chlorination of 2-methylpropan-2-ol using concentrated hydrochloric acid
  • 7: Modern Analytical Techniques I
    7.1: Mass spectrometry
    7.1.1: Use data from a mass spectrometer
    7.2: Infrared (IR) spectroscopy
    7.2.1: Deduce functional groups and to predict infrared absorptions
    7.2.2: Analysis of some inorganic and organic unknowns
  • 8: Energetics I
    8.1: Energetics I
    8.1.1: Standard conditions are 100 kPa and a specified temperature, usually 298 K
    8.1.2: The enthalpy change is the heat energy change measured at constant pressure
    8.1.3: Construct and interpret enthalpy level diagrams
    8.1.4: Standard enthalpy change
    8.1.5: Experiments to measure enthalpy changes
    8.1.6: Calculate enthalpy changes in kJ mol-1
    8.1.7: Construct enthalpy cycles using Hess’s Law
    8.1.8: Calculate enthalpy changes from data using Hess’s Law
    8.1.9: The enthalpy change of a reaction using Hess’s Law
    8.1.10: ‘bond enthalpy’ and ‘mean bond enthalpy’
    8.1.11: Enthalpy change of reaction using mean bond enthalpies
    8.1.12: Mean bond enthalpies from enthalpy changes of reaction
  • 9: Kinetics I
    9.1: Kinetics I
    9.1.1: The effect on the rate of a chemical reaction
    9.1.2: Activation energy
    9.1.3: The rate of a reaction
    9.1.4: Changes in temperature affect the rate of a reaction
    9.1.5: The role of catalysts in providing alternative reaction routes
    9.1.6: The reaction profiles for uncatalysed and catalysed reactions
    9.1.7: The Maxwell-Boltzmann distribution of molecular energies
    9.1.8: The use of a solid (heterogeneous) catalyst for industrial reactions
    9.1.9: The economic benefits of the use of catalysts in industrial reactions
  • 10: Equilibrium I
    10.1: Equilibrium I
    10.1.1: Many reactions can reach a state of dynamic equilibrium
    10.1.2: The qualitative effect of a change in temperature, concentration or pressure
    10.1.3: Compromise between the yield and the rate of reaction
    10.1.4: Equilibrium concentrations
  • 11: Equilibrium II
    11.1: Equilibrium II
    11.1.1: Expression for Kp, for homogeneous and heterogeneous systems
    11.1.2: Value for the equilibrium constant for homogeneous and heterogeneous reactions
    11.1.3: The effect of changing temperature on the equilibrium constant
    11.1.4: The effect of temperature on the position of equilibrium
    11.1.5: The value of the equilibrium constant
  • 12: Acid-base Equilibria
    12.1: Acid-base Equilibria
    12.1.1: Brønsted–Lowry acid and base
    12.1.2: Acid-base reactions involve the transfer of protons
    12.1.3: Brønsted–Lowry conjugate acid-base pairs
    12.1.4: The term ‘pH’
    12.1.5: Calculating pH from hydrogen ion concentration
    12.1.6: The concentration of hydrogen ions in a solution from its pH
    12.1.7: The difference between a strong acid and a weak acid
    12.1.8: The pH of a strong acid
    12.1.9: The expression for the acid dissociation constant
    12.1.10: The pH of a weak acid
    12.1.11: The ionic product of water
    12.1.12: The pH of a strong base
    12.1.13: ‘pKa’ and ‘pKw’
    12.1.14: Analyse data from the experiments
    12.1.15: Ka for a weak acid
    12.1.16: Draw and interpret titration curves
    12.1.17: Suitable indicator
    12.1.18: ‘buffer solution’
    12.1.19: The action of a buffer solution
    12.1.20: The pH of a buffer solution
    12.1.21: Concentrations of solutions required to prepare a buffer solution of a given pH
    12.1.22: Using a weak acid–strong base titration curve
    12.1.23: Enthalpy changes of neutralisation values for strong and weak acids
    12.1.24: The roles of carbonic acid molecules and hydrogencarbonate ions
    12.1.25: Finding the Ka value for a weak acid
  • 13: Energetics II
    13.1: Lattice energy
    13.1.1: Lattice energy
    13.1.2: Enthalpy change of atomisation and electron affinity
    13.1.3: Born-Haber cycles
    13.1.4: Lattice energy provides a measure of ionic bond strength
    13.1.5: The degree of covalent bonding
    13.1.6: Polarisation as applied to ions
    13.1.7: The polarising power of a cation depends on its radius and charge
    13.1.8: Polarisability of an anion depends on its radius and charge
    13.1.9: 'enthalpy change of solution, ΔsolH’, and ‘enthalpy change of hydration, ΔhydH’
    13.1.10: Use energy cycles and energy level diagrams
    13.1.11: The effect of ionic charge and ionic radius on the values
    13.2: Entropy
    13.2.1: Enthalpy changes alone do not control whether reactions occur
    13.2.2: Entropy is a measure of the disorder of a system
    13.2.3: Why entropy changes occur
    13.2.4: The total entropy change in any reaction
    13.2.5: The entropy change for the system
    13.2.6: The entropy change in the surroundings
    13.2.7: The balance between the entropy change and the enthalpy change
    13.2.8: The equation ΔG = ΔH − TΔSsystem
    13.2.9: The equation ΔG = −RT ln K
    13.2.10: A reaction for which the ΔG value is negative may not occur in practice
    13.2.11: Thermodynamically feasible reactions may be inhibited by kinetic factors
  • 14: Redox II
    14.1: Redox II
    14.1.1: ‘oxidation’ and ‘reduction’ in terms of electron transfer
    14.1.2: ‘oxidation’ and ‘reduction’ in terms of changes in oxidation number
    14.1.3: Standard electrode potential, Eo
    14.1.4: The standard electrode potential, Eo, refers to conditions
    14.1.5: The features of the standard hydrogen electrode
    14.1.6: Different methods are used to measure standard electrode potentials
    14.1.7: Investigating some electrochemical cells
    14.1.8: Standard emf, Eo cell
    14.1.9: Write cell diagrams using the conventional representation of half-cells
    14.1.10: The importance of the conditions when measuring the electrode potential, E
    14.1.11: The thermodynamic feasibility of a reaction
    14.1.12: Eo cell is directly proportional to the total entropy change and to ln K
    14.1.13: The limitations of predictions made using standard electrode potentials
    14.1.14: Standard electrode potentials can be listed as an electrochemical series
    14.1.15: Disproportionation reactions relate to standard electrode potentials
    14.1.16: The application of electrode potentials to storage cells
    14.1.17: Energy released on the reaction of a fuel with oxygen
    14.1.18: The electrode reactions that occur in a hydrogen-oxygen fuel cell
    14.1.19: Structured and non-structured titration calculations
    14.1.20: The methods used in redox titrations
    14.1.21: Redox titration
  • 15: Transition Metals
    15.1: Principles of transition metal chemistry
    15.1.1: The electronic configurations of atoms and ions of the d-block elements
    15.1.2: Transition metals are d-block elements
    15.1.3: Transition metals show variable oxidation number
    15.1.4: ‘ligand’
    15.1.5: dative (coordinate) bonding is involved in the formation of complex ions
    15.1.6: A complex ion is a central metal ion surrounded by ligands
    15.1.7: Transition metals form coloured ions in solution
    15.1.8: The colour of aqueous ions and other complex ions
    15.1.9: A lack of colour in some aqueous ions and other complex ions
    15.1.10: Colour changes in transition metal ions
    15.1.11: ‘coordination number’
    15.1.12: H2O, OH− and NH3 act as monodentate ligands
    15.1.13: Complexes with six-fold coordination have an octahedral shape
    15.1.14: Transition metal ions may form tetrahedral complexes
    15.1.15: Square planar complexes are also formed by transition metal ions
    15.1.16: Cis-platin used in cancer treatment is supplied as a single isomer
    15.1.17: Identify bidentate ligands and multidentate ligands
    15.1.18: Haemoglobin is an iron(II) complex containing a multidentate ligand
    15.1.19: Ligand exchange reaction
    15.2: Reactions of transition metal elements
    15.2.1: The colours of the oxidation states of vanadium
    15.2.2: Redox reactions for the interconversion of the oxidation states of vanadium
    15.2.3: Dichromate(VI) ion, Cr2O72− can be reduced and can be produced by oxidation
    15.2.4: Dichromate(VI) ion, Cr2O72−, can be converted into chromate(VI) ions
    15.2.5: The reactions of Cr3+(aq), Fe2+(aq), Fe3+(aq), Co2+(aq) and Cu2+(aq)
    15.2.6: ligand exchange and amphoteric behaviour
    15.2.7: Ligand exchange, and an accompanying colour change
    15.2.8: The substitution of small, uncharged ligands by larger, charged ligands
    15.2.9: The substitution of a monodentate ligand by a bidentate or multidentate ligand
    15.2.10: Heterogeneous and homogeneous catalysts
    15.2.11: A heterogeneous catalyst is in a different phase from the reactants
    15.2.12: V2O5 acts as a catalyst in the contact process
    15.2.13: Catalytic converter decreases carbon monoxide and nitrogen monoxide emissions
    15.2.14: A homogeneous catalyst is in the same phase as the reactants
    15.2.15: The role of Fe2+ ions in catalysing the reaction between I− and S2O82− ions
    15.2.16: The role of Mn2+ ions in autocatalysing the reaction
    15.2.17: The preparation of a transition metal complex
  • 16: Kinetics II
    16.1: Kinetics II
    16.1.1: Understand the terms
    16.1.2: rate equations of the form: rate = k[A]m[B]n
    16.1.3: Rate data for a given reaction
    16.1.4: Experiments that can be used to investigate reaction rates
    16.1.5: The rate of reaction and the half-life of a first-order reaction
    16.1.6: Deduce the order (0, 1 or 2) with respect to a substance in a rate equation
    16.1.7: Deduce the order (0, 1 or 2) with respect to a substance in a rate equation
    16.1.8: Obtain data to calculate the order and use these data to make predictions
    16.1.9: Deduce a rate-determining step from a rate equation and vice versa
    16.1.10: deduce a reaction mechanism
    16.1.11: Evidence for SN1 or SN2 mechanisms for halogenoalkane hydrolysis
    16.1.12: The activation energy for a reaction from experimental data
    16.1.13: Following the rate of the iodine-propanone reaction by a titrimetric method
    16.1.14: Finding the activation energy of a reaction
  • 17: Organic Chemistry II
    17.1: Chirality
    17.1.1: Optical isomerism
    17.1.2: Optical isomerism results from chiral centre(s) in a molecule
    17.1.3: Optical activity
    17.1.4: Racemic mixture
    17.1.5: Optical activity of reactants and products
    17.2: Carbonyl compounds
    17.2.1: The aldehyde and ketone functional groups
    17.2.2: Aldehydes and ketones
    17.2.3: The reactions of carbonyl compounds
    17.2.3: The reactions of carbonyl compounds
    17.3: Carboxylic acids
    17.3.1: The carboxylic acid functional group
    17.3.2: Hydrogen bonding affects the physical properties of carboxylic acids
    17.3.3: Carboxylic acids can be prepared by the oxidation of alcohols or aldehydes
    17.3.4: The reactions of carboxylic acids
    17.3.5: The acyl chloride and ester functional groups
    17.3.6: The reactions of acyl chlorides
    17.3.7: The hydrolysis reactions of esters, in acidic and alkaline solution
    17.3.8: Polyesters are formed by condensation polymerisation reactions
  • 18: Organic Chemistry III
    18.1: Arenes - benzene
    18.1.1: The bonding in benzene has been represented using the Kekulé
    18.1.2: Enthalpy changes of hydrogenation and carbon-carbon bond lengths
    18.1.3: Benzene is resistant to bromination, compared with alkenes
    18.1.4: The reactions of benzene
    18.1.5: The mechanism of the electrophilic substitution reactions of benzene
    18.1.6: The reaction of phenol with bromine water
    18.1.7: Reasons for the relative ease of bromination of phenol, compared to benzene
    18.2: Amines, amides, amino acids and proteins
    18.2.1: Amine and amide functional groups
    18.2.2: The reactions of primary aliphatic amines
    18.2.3: Difference in basicity of ammonia, primary aliphatic and primary aromatic amines
    18.2.4: Reagents and general reaction conditions
    18.2.5: Aromatic nitro-compounds can be reduced
    18.2.6: Amides can be prepared from acyl chlorides
    18.2.7: The formation of a polyamide is a condensation polymerisation reaction
    18.2.8: The structural formulae of the repeat units of condensation polymers
    18.2.9: The properties of 2-amino acids
    18.2.10: The peptide bond in proteins
    18.2.11: Inorganic and organic unknowns
    18.3: Organic Synthesis
    18.3.1: The empirical formulae, molecular formulae and structural formulae of compounds
    18.3.2: Reaction schemes to form both familiar and unfamiliar compounds
    18.3.3: Methods of increasing the length of the carbon chain in a molecule
    18.3.4: Reactions involving compounds with functional group
    18.3.5: The preparation and purification of organic compounds
    18.3.6: The preparation of aspirin
  • 19: Modern Analytical Techniques II
    19.1: Mass spectrometry
    19.1.1: Use data from mass spectra
    19.2: Nuclear magnetic resonance (NMR)
    19.2.1: The positions of 13C atoms in a molecule
    19.2.2: Use data from 13C NMR spectroscopy
    19.2.3: The positions of 1H atoms in a molecule
    19.2.4: Use data from high resolution 1H NMR spectroscopy
    19.3: Chromatography
    19.3.1: Chromatography separates components of a mixture
    19.3.2: Calculate Rf values from one-way chromatograms
    19.3.3: High performance liquid chromatography, HPLC, and gas chromatography, GC

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