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IB Chemistry HL

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Chemistry HL

6 topics64 outcomesLocal syllabus version: 2025

S1: Models of the Particulate Nature of Matter

9 outcomes

Open
  1. S1.1.1Describe the nuclear model of the atom: protons and neutrons in a nucleus surrounded by electrons; define atomic number, mass number, and isotopes
  2. S1.1.2Calculate relative atomic mass from isotopic masses and abundances; describe how a mass spectrometer separates ions by mass-to-charge ratio
  3. S1.2.1Describe the electromagnetic spectrum and explain how atomic emission and absorption line spectra provide evidence for discrete electron energy levels
  4. S1.2.2Write electron configurations for elements and ions up to Z = 36 using sublevel (s, p, d) notation; apply the Aufbau principle, Hund's rule, and Pauli exclusion principle
  5. S1.2.3Explain trends in first ionisation energy across periods and down groups in terms of nuclear charge, shielding, and atomic radius HL only
  6. S1.2.4Describe the shapes and relative energies of s, p, and d orbitals; explain hybridisation (sp, sp², sp³) and its application to molecular geometry HL only
  7. S1.3.1Apply the mole concept: use the Avogadro constant (6.022 × 10²³ mol⁻¹) to interconvert between amount of substance, number of particles, and molar mass
  8. S1.3.2Determine empirical and molecular formulae from percentage composition data and combustion analysis results
  9. S1.4.1Apply the ideal gas law PV = nRT to solve problems involving pressure, volume, temperature, and amount of gas; describe deviations of real gases from ideal behaviour HL only

S2: Models of Bonding and Structure

10 outcomes

Open
  1. S2.1.1Describe ionic bonding as electrostatic attraction between oppositely charged ions; explain lattice structure and relate properties (melting point, conductivity) to ionic bonding
  2. S2.1.2Apply the Born-Haber cycle to calculate lattice enthalpies of ionic compounds, using Hess's law and standard enthalpy data HL only
  3. S2.2.1Describe covalent bonding as the sharing of electron pairs; draw Lewis (electron dot) structures for molecules and polyatomic ions
  4. S2.2.2Predict and explain the shapes and bond angles of molecules and ions (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral) using VSEPR theory
  5. S2.2.3Distinguish between polar and non-polar bonds using electronegativity differences; determine the polarity of molecules from their geometry
  6. S2.2.4Describe sigma (σ) and pi (π) bonds in terms of orbital overlap; explain resonance and delocalisation using formal charge analysis HL only
  7. S2.2.5Apply molecular orbital theory to describe bonding and antibonding orbitals in diatomic molecules; predict bond order and magnetic properties HL only
  8. S2.3.1Describe metallic bonding as the electrostatic attraction between a lattice of cations and a sea of delocalised electrons; relate to physical properties of metals
  9. S2.4.1Compare the four types of solid structure (ionic, metallic, covalent network, molecular) and relate their structures to physical properties
  10. S2.4.2Describe London dispersion forces, dipole-dipole attractions, and hydrogen bonding; explain their influence on physical properties such as boiling points and solubility

S3: Classification of Matter

7 outcomes

Open
  1. S3.1.1Explain the organisation of the periodic table by increasing atomic number; identify patterns in physical and chemical properties across periods and down groups
  2. S3.1.2Explain trends in atomic radius, ionic radius, electronegativity, electron affinity, and ionisation energies across periods and down groups
  3. S3.1.3Describe the characteristic properties of transition metals: variable oxidation states, coloured compounds, catalytic activity, and complex ion formation HL only
  4. S3.1.4Explain the colour of transition metal complexes using crystal field theory (d-orbital splitting); describe the effect of ligand type on colour and stability HL only
  5. S3.2.1Name and draw structural formulae for the major homologous series: alkanes, alkenes, alkynes, arenes, halogenoalkanes, alcohols, aldehydes, ketones, carboxylic acids, esters, amines, and amides
  6. S3.2.2Identify structural isomers (chain, position, and functional group) and stereoisomers (cis/trans and optical) for organic compounds
  7. S3.2.3Predict physical and chemical properties of organic compounds from their functional groups and carbon skeleton

R1: What Drives Chemical Reactions?

8 outcomes

Open
  1. R1.1.1Define enthalpy change (ΔH) and distinguish between exothermic and endothermic reactions; calculate ΔH from calorimetric data using q = mcΔT
  2. R1.1.2Calculate standard enthalpy changes using Hess's law and standard enthalpies of formation (ΔfH°) or combustion (ΔcH°)
  3. R1.1.3Estimate enthalpy changes using average bond enthalpies and explain why calculated values differ from experimental values
  4. R1.2.1Apply Hess's law in energy cycles; calculate lattice enthalpies and enthalpies of solution using Born-Haber and energy cycles HL only
  5. R1.3.1Compare energy density and combustion products of different fuels (hydrogen, hydrocarbons, biofuels) and evaluate their suitability as energy sources
  6. R1.4.1Define entropy (S) and explain how dispersal of energy and matter contribute to an increase in entropy in physical and chemical processes
  7. R1.4.2Calculate Gibbs free energy change (ΔG = ΔH − TΔS) and use its sign to predict whether a reaction is spontaneous under given conditions
  8. R1.4.3Relate ΔG° to the equilibrium constant K using ΔG° = −RT ln K, and interpret the relationship between thermodynamics and the position of equilibrium HL only

R2: How Much, How Fast, and How Far?

13 outcomes

Open
  1. R2.1.1Balance chemical equations and use stoichiometric mole ratios to calculate masses, volumes, and amounts of reactants and products
  2. R2.1.2Identify the limiting reagent in a reaction; calculate theoretical yield, actual yield, and percentage yield
  3. R2.1.3Perform solution stoichiometry calculations involving concentration (mol dm⁻³), dilution, and acid-base titrations
  4. R2.2.1Define rate of reaction; describe experimental methods for monitoring reaction rate (gas volume, mass loss, colour change, conductivity)
  5. R2.2.2Explain collision theory and the effect of concentration, temperature, surface area, pressure (for gases), and catalysts on reaction rate using Maxwell-Boltzmann distributions
  6. R2.2.3Derive the rate expression (rate law) from experimental data; determine reaction orders and the rate constant k HL only
  7. R2.2.4Use the Arrhenius equation (k = Ae^(−Ea/RT)) to calculate activation energy from rate constant data at different temperatures HL only
  8. R2.2.5Identify the rate-determining step from a proposed reaction mechanism and verify consistency with the observed rate expression HL only
  9. R2.3.1Define dynamic equilibrium; describe the characteristics of a system at equilibrium in terms of equal forward and reverse rates
  10. R2.3.2Apply Le Chatelier's principle to predict the effect of changes in concentration, pressure/volume, temperature, and addition of a catalyst on the position of equilibrium
  11. R2.3.3Write the equilibrium constant expression Kc for homogeneous equilibria and calculate Kc from equilibrium concentrations
  12. R2.3.4Write and use Kp expressions for gaseous equilibria; interconvert Kc and Kp HL only
  13. R2.3.5Use the reaction quotient Q to predict the direction in which a reaction will proceed towards equilibrium HL only

R3: What Are the Mechanisms of Chemical Change?

17 outcomes

Open
  1. R3.1.1Apply the Brønsted-Lowry theory to identify acids (proton donors), bases (proton acceptors), and conjugate acid-base pairs in reactions
  2. R3.1.2Distinguish between strong and weak acids/bases; calculate pH and pOH from [H⁺] and [OH⁻] concentrations using Kw = 1.0 × 10⁻¹⁴ at 25 °C
  3. R3.1.3Calculate pH of weak acids and bases using Ka and Kb values; derive the relationship pKa + pKb = pKw
  4. R3.1.4Describe acid-base titration curves for strong/weak acid and strong/weak base combinations; identify the equivalence point and choose an appropriate indicator
  5. R3.1.5Explain the composition and action of buffer solutions; calculate buffer pH using the Henderson-Hasselbalch equation HL only
  6. R3.1.6Apply the Lewis theory to identify acids (electron pair acceptors) and bases (electron pair donors), extending to reactions not covered by Brønsted-Lowry HL only
  7. R3.2.1Define oxidation and reduction in terms of electron transfer and changes in oxidation state; identify oxidising and reducing agents and balance redox equations
  8. R3.2.2Describe the operation of a voltaic (galvanic) cell including electrodes, salt bridge, and direction of electron and ion flow
  9. R3.2.3Calculate cell potential E°cell = E°cathode − E°anode using standard electrode potentials; predict spontaneity of redox reactions
  10. R3.2.4Apply ΔG° = −nFE° to relate thermodynamics and electrochemistry; use Faraday's laws to calculate mass deposited or gas evolved during electrolysis HL only
  11. R3.2.5Describe the products of electrolysis of aqueous and molten electrolytes; explain the factors that determine electrode products HL only
  12. R3.3.1Describe free-radical reactions (homolytic bond cleavage), including the initiation, propagation, and termination steps of the halogenation of alkanes
  13. R3.4.1Describe electrophilic addition reactions of alkenes with H₂, X₂, HX (Markovnikov's rule), and H₂O; explain the mechanism using curly arrows
  14. R3.4.2Describe nucleophilic substitution of halogenoalkanes (SN1 and SN2); explain factors affecting which mechanism predominates HL only
  15. R3.4.3Describe electrophilic aromatic substitution (nitration, halogenation, Friedel-Crafts) of benzene and explain the role of the activating/deactivating substituent effects HL only
  16. R3.4.4Design multi-step organic synthesis routes, using retrosynthetic analysis, for target molecules from given starting materials HL only
  17. R3.4.5Interpret IR, ¹H NMR, and mass spectra to identify functional groups, molecular fragments, and determine structures of organic compounds HL only