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IB DP Chemistry 2025 Syllabus

IB DP Chemistry 2025 Syllabus

The new DP chemistry course was launched in February 2023 for first teaching in August 2023. First assessment will take place in May 2025. IBDP Chemistry courses are here as per new syllabus and guidelines provided by the board. Details of changes in pattern and syllabus are also mentioned on this page. 

Structure 1. Models of the particulate nature of matter

1.1 Introduction to the particulate nature of matter

  • Distinguish between the properties of elements, compounds and mixtures.
    • Solvation, filtration, recrystallization, evaporation,
    • distillation and paper chromatography .
    • The differences between homogeneous and heterogeneous mixtures.
  • Distinguish the different states of matter. Use state symbols (s, l, g and aq) in chemical equations.
    • Names of the changes of state should be covered:
    • melting, freezing, vaporization (evaporation and boiling), condensation, sublimation and deposition.
  • Interpret observable changes in physical properties and temperature during changes of state.
    • Convert between values in the Celsius and Kelvin scales.

1.2 The nuclear atom

  • Use the nuclear symbol ZXA to deduce the number of protons, neutrons and electrons in atoms and ions..
    • Relative masses and charges of the subatomic particles
  • Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data..
    • Differences in the physical properties of isotopes:
    • Specific examples of isotopes.

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  • Interpret mass spectra in terms of identity and relative abundance of isotopes.

1.3 Electron configurations

  • Emission spectra
    • Qualitatively describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.
    • Distinguish between a continuous and a line spectrum.
  • Line emission spectrum of hydrogen
    • Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.
  • Energy levels
    • Deduce the maximum number of electrons that can occupy each energy level.
  • Energy sublevels
    • Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals.
  • Electron Configuration
    • Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.
    • electron configurations and condensed electron configurations using the noble gas core.
    • Orbital diagrams, i.e. arrow-in-box diagrams,
    • The electron configurations of Cr and Cu

AHL

  • First ionization energy (IE)
    • Trends and discontinuities in first ionization energy (IE) across a period and down a group.
  • Successive ionization energy (IE)
    • Deduce the group of an element from its successive ionization data.

1.4 The mole

  • Convert the amount of substance, n, to the number of specified elementary entities.
    • An elementary entity may be an atom, a molecule, an ion, an electron, any other particle or a specified group of particles.
  • Determine relative formula masses Mr from relative atomic masses Ar .
  • Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
  • Interconvert the percentage composition by mass and the empirical formula.
  • Determine the molecular formula of a compound from its empirical formula and molar mass.
  • Solve problems involving the molar concentration, amount of solute and volume of solution.
    • The use of square brackets to represent molar concentration
    • Units of concentration should include g dm–3 and mol dm–3 and conversion between these.
  • Avogadro’s law
    • Solve problems involving the mole ratio of reactants and/or products and the volume of gases.

1.5 Ideal gases

  • key assumptions in the ideal gas model.
  • limitations of the ideal gas model.
  • relationship between temperature, pressure and volume for a fixed mass of an ideal gas and analyse graphs relating these variables.
  • ideal gas equation
    • Solve problems relating to the ideal gas equation.

Structure 2. Models of bonding and structure

2.1 Ionic bonds & structure

    • Anions and Cations.
      • Predict the charge of an ion from the electron configuration of the atom.
    • Formation of Ionic Bond
      • Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions.
      • Binary ionic compounds
      • Interconvert names and formulas of binary ionic compounds.
      • name and formula: ammonium NH4+, hydroxide OH–, nitrate NO3–, hydrogencarbonate HCO3–, carbonate CO32–, sulfate SO42–, phosphate PO43–.
    • Physical properties of ionic compounds
      • Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility.
      • lattice enthalpy

2.2 The covalent model

  • Covalent Bond and Lewis Structures

    • Define a covalent bond.
    • Deduce the Lewis formula for molecules and ions with up to four electron pairs per atom.
    • Apply the octet rule in organic and inorganic molecules.

  • Bond Types and Properties

    • Differentiate between single, double, and triple bonds.
    • Explain the relationship between bond number, length, and strength.

  • Coordination Bonds

    • Define and identify coordination bonds, including in transition element complexes.

  • VSEPR Theory

    • Use the VSEPR model to predict electron domain and molecular geometry.
    • Discuss the influence of non-bonding pairs and multiple bonds on bond angles.

  • Bond Polarity

    • Determine the polar nature of covalent bonds using electronegativity values.
    • Understand molecular polarity and deduce net dipole moments.

  • Molecular Structures and Properties

    • Describe and explain properties of silicon, silicon dioxide, and carbon allotropes (diamond, graphite, fullerenes, graphene).

  • Intermolecular Forces

    • Understand key intermolecular forces (London dispersion, dipole–dipole, hydrogen bonding).
    • Deduce types of intermolecular forces based on molecular structure.

  • Intermolecular Force Comparison

    • Compare the strengths of intermolecular forces and their impact on the physical properties of covalent substances.

  • Chromatography

    • Explain chromatography and interpret retardation factor values (RF) based on intermolecular forces.
  • Resonance Structures
    • Understand and deduce resonance structures in molecules and ions.
  • Benzene Structure
    • Discuss benzene’s structure based on physical and chemical evidence.
  • Expanded Octets
    • Represent Lewis structures for species with expanded octets.
  • Sigma and Pi Bonds
    • Define and identify sigma (σ) and pi (π) bonds in various examples.
  • Hybridization
    • Define hybridization and analyze bond formation.
    • Predict geometry from hybridization.

2.3 The metallic model

  • Metallic Bond and Properties
      • Define a metallic bond.
      • Explain the electrical and thermal conductivity of metals.
      • Discuss the malleability of metals and relate their properties to experimental uses.

    • Factors Affecting Metallic Bond Strength

      • Understand how ion charge and metal ion radius influence bond strength.
      • Explain trends in melting points for s and p block metals.
    • Transition Elements
        • Understand delocalized d-electrons in transition elements.
        • Explain their high melting point and electrical conductivity.

2.4 From models to materials

  • Bonding Continuum and Bonding Triangle
    • Understand bonding as a continuum between ionic, covalent, and metallic models.
    • Use bonding models to explain material properties.
  • Bonding Triangle and Material Properties
    • Determine a compound’s position in the bonding triangle using electronegativity data.
    • Predict material properties based on this position.
  • Alloys
    • Define alloys as mixtures of metals with other elements.
    • Explain alloy properties, including examples like bronze, brass, and stainless steel.
  • Polymers
    • Define polymers as large molecules made from repeating monomers.
    • Describe the properties of plastics based on their structure.
    • Understand and represent addition polymers from monomer structures. 

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  • Formation of Condensation Polymers
    • Understand the formation reactions of condensation polymers.
    • Represent repeating units of polyamides and polyesters from monomers

Structure 3. Classification of matter

3.1 Periodic Table & Trends

  • Classification and Electron Configuration
    • Identify positions of metals, metalloids, and non-metals in the periodic table.
    • Deduce electron configurations up to Z = 36 and classify elements as alkali metals, halogens, transition elements, or noble gases.
  • Periodicity
    • Define periodicity and explain trends in atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity.
  • Group Reactions
    • Describe reactions of Group 1 metals with water and Group 17 elements with halide ions.
    • Deduce equations for reactions involving oxides of Group 1, Group 2 metals, carbon, and sulfur.
  • Oxidation States
    • Understand oxidation states and determine oxidation numbers in compounds, including special cases like hydrogen in metal hydrides and oxygen in peroxides.

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  • Ionization Energy
    •  Explain discontinuities in first ionization energy trends across a period and their evidence for energy sublevels.
  • Transition Elements
    • Recall properties of transition elements, including variable oxidation states, high melting points, magnetic properties, and the formation of colored compounds.
  • Oxidation States and Color
    • Explain the formation of variable oxidation states and why transition element complexes are colored. Apply the color wheel to deduce absorbed/observed light.

3.2 Functional Groups

  • Chemical Formulas
    • Use and interconvert between empirical, molecular, structural, stereochemical, and skeletal formulas. Construct 3D models of organic molecules.
  • Functional Group Identification
    • Identify functional groups (e.g., halogeno, hydroxyl, carbonyl) and understand terms like “saturated” and “unsaturated.”
  • Homologous Series
    • Define and identify homologous series (e.g., alkanes, alkenes) and describe trends in melting and boiling points.
  • IUPAC Nomenclature
    • Apply IUPAC rules to name saturated or mono-unsaturated compounds with up to six carbon atoms, including straight-chain and branched-chain isomers.
  • Structural Isomers
    • Define structural isomers and recognize different types (branched, straight-chain, position, functional group isomers).

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  • Stereoisomers
    • Define stereoisomers, describe cis-trans isomerism, and recognize enantiomers. Understand and use terms like “chiral,” “optical activity,” “enantiomer,” and “racemic mixture.”
  • Structural Analysis
    • Deduce structural features from mass spectrometry (MS) and interpret IR and 1H NMR spectra to determine molecular structures.

Reactivity 1. What drives chemical reactions?

1.1 Enthalpy

  • Energy Transfer
    • Understand the transfer of energy in chemical reactions, differentiating between heat and temperature.
  • Endothermic vs. Exothermic
    • Classify reactions as endothermic or exothermic, and understand associated temperature changes.
  • Energy Profiles
    • Sketch and interpret energy profiles for endothermic and exothermic reactions.
  • Calculations
    • Understand and apply the standard enthalpy change for reactions (ΔH⦵). Use equations \(Q = mcΔT\) and \(ΔH = -\frac{Q}{n}\) in enthalpy calculations.

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  • Standard Enthalpy Changes
    • Understand and calculate standard enthalpy changes of combustion (ΔHc⦵) and formation (ΔHf⦵).
  • Born–Haber Cycle
    • Understand and interpret a Born–Haber cycle for ionic compounds.

1.2 Cycles of Energy

  • Bond Energy
    • Understand that bond-breaking absorbs energy and bond-forming releases energy. Calculate enthalpy changes using average bond enthalpy data.
  • Hess’s Law
    • Define and apply Hess’s law to calculate enthalpy changes in multi-step reactions.

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  • Entropy
    • Define entropy (S) as a measure of matter/energy dispersal in a system. Predict changes in entropy and calculate standard entropy changes (ΔS⦵).
  • Gibbs Energy
    • Relate Gibbs energy change (ΔG) to enthalpy, entropy, and temperature using \(ΔG⦵ = ΔH⦵ – TΔS⦵\).
  • Spontaneity & Equilibrium
    • Interpret ΔG values to determine reaction spontaneity and equilibrium conditions. Apply equations \(ΔG = ΔG⦵ + RT \ln Q\) and \(ΔG⦵ = -RT \ln K\).

1.3 Fuels

  • Combustion Reactions
    • Deduce equations for combustion and incomplete combustion of hydrocarbons and alcohols.
  • Fossil Fuels & Environment
    • Recall fossil fuels like coal, crude oil, and natural gas. Evaluate CO₂ emissions and understand the greenhouse effect.
  • Biofuels
    • Understand biofuel production and the differences between renewable and non-renewable energy sources. Discuss advantages and disadvantages of biofuels.
  • Fuel Cells
    • Deduce half-equations for fuel cell electrode reactions using hydrogen and methanol as fuels.

Reactivity 2. How much, how fast and how far?

2.1 Amount of Change

  • Chemical Equations

    • Write and balance chemical equations, including state symbols.

  • Stoichiometry

    • Use mole ratios to calculate masses, volumes, and concentrations. Understand Avogadro’s law and molar concentration.

  • Limiting Reactants

    • Identify limiting and excess reactants. Calculate theoretical yield and compare it with experimental yield.

  • Atom Economy

    • Calculate atom economy and understand its importance in green chemistry.

2.2 Rate of Change

  • Rate of Reaction

    • Define and calculate reaction rates. Understand the relationship between kinetic energy, temperature, and reaction rates.

  • Factors Influencing Rate

    • Explain how factors like temperature, concentration, and catalysts affect reaction rates.

  • Activation Energy

    • Define activation energy (Ea). Use Maxwell–Boltzmann distribution curves to explain successful collisions and the role of catalysts.

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  • Reaction Mechanisms

    • Analyze proposed mechanisms, identify intermediates, and construct energy profiles.

  • Molecularity & Rate Equation

    • Define molecularity and deduce rate equations from experimental data.

  • Order of Reaction

    • Understand reaction order, interpret graphical representations, and solve problems involving rate equations.

  • Arrhenius Equation

    • Understand and use the Arrhenius equation to determine activation energy and the rate constant.

2.3 Extent of Change

  • Dynamic Equilibrium

    • Understand dynamic equilibrium in closed systems. Describe characteristics of equilibrium in physical and chemical systems.

  • Equilibrium Law

    • Deduce equilibrium constant expressions for homogeneous reactions and analyze their significance.

  • Le Châtelier’s Principle

    • Predict the effect of changes in conditions on equilibrium using Le Châtelier’s principle.

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  • Reaction Quotient (Q)

    • Calculate the reaction quotient and predict the direction of reaction progress.

  • Equilibrium Calculations

    • Solve problems involving K values and equilibrium concentrations.

  • Gibbs Energy and Equilibrium

    • Use the relationship between Gibbs energy change (ΔG) and equilibrium constant (K) to analyze reaction equilibria.

Reactivity 3. What are the mechanisms of chemical change?

3.1 Proton Transfer

  • Brønsted–Lowry Acids and Bases:
    • Acid: A proton donor.
    • Base: A proton acceptor.
    • Identification in Reactions: Identify which species donates and which accepts a proton.
    • Base vs. Alkali:
    • Base: Any proton acceptor.
    • Alkali: A base that dissolves in water to produce OH⁻ ions (e.g., NaOH).
  • Conjugate Acid-Base Pairs:
    • Definition: Pairs of species differing by a single proton.
    • Deduction: Given a base, its conjugate acid is formed by adding H⁺; given an acid, its conjugate base is formed by removing H⁺.
  • Amphoteric Species:
    • Definition: Species that can act as both an acid and a base (e.g., H₂O).
    • Equations: Write reactions showing how such species can donate or accept protons.
  • pH Scale:
    • Definition: \( \text{pH} = -\log_{10}[\text{H}^+] \); \( [\text{H}^+] = 10^{-\text{pH}} \).
    • Calculations: Perform pH-related calculations and understand how to estimate pH using indicators or measure it precisely with pH meters.
  • Ion Product of Water (Kw):
    • Expression: \( \text{Kw} = [\text{H}^+] [\text{OH}^-] \).
  • Acidity/Basicity:
    • Acidic: \( [\text{H}^+] > [\text{OH}^-] \)
    • Neutral: \( [\text{H}^+] = [\text{OH}^-] \)
    • Basic: \( [\text{H}^+] < [\text{OH}^-] \)
  • Strong vs. Weak Acids/Bases:
    • Strong Acids/Bases: Fully ionize in solution (e.g., HCl).
    • Weak Acids/Bases: Partially ionize (e.g., CH₃COOH).
    • Equilibria: Lies toward the weaker conjugate pair.
  • Neutralization Reactions:
    • Definition: Acids react with bases to form water and salts.
    • Equations: Formulate equations for acid reactions with metal oxides, hydroxides, hydrogencarbonates, and carbonates.
  • pH Curves:
    • Understanding: Recognize, sketch, and interpret the pH curves for strong acid-strong base neutralization reactions.

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  • pOH Scale:
    • Definition: \( \text{pOH} = -\log_{10}[\text{OH}^-] \); \( [\text{OH}^-] = 10^{-\text{pOH}} \).
    • Interconversion: Convert between pH, pOH, [H⁺], and [OH⁻].
  • Strengths of Weak Acids/Bases:
    • Described by: \( K_a \), \( K_b \), \( pK_a \), \( pK_b \) values.
    • Equilibrium Relationships: \( K_a \times K_b = K_w \).
  • pH of Salt Solutions:
    • Dependence: On the relative strengths of the parent acid and base.
    • Hydrolysis: Write equations showing how ions in salts affect the pH.
  • pH Curves for Different Combinations:
    • Combinations: Strong/Weak Acid with Strong/Weak Base.
    • Key Points: Identify buffer regions, equivalence points, and the significance of \( pK_a \) and \( pK_b \).
  • Acid-Base Indicators:
    • Definition: Weak acids where the conjugate forms have different colors.
    • Equilibria: Show why the color changes with pH.
    • Universal Indicator: A mixture with a wide range of color changes.
  • Titration Indicators:
    • Selection: Based on the identity of the salt and the indicator’s pH range.
    • End Point vs. Equivalence Point: Differentiate between these in titration.
  • Buffer Solutions:
    • Definition: Resist changes in pH when small amounts of acid/alkali are added.
    • Composition: Typically consist of a weak acid and its conjugate base or vice versa.
  • pH of Buffers:
    • Dependence: On \( pK_a \)/\( pK_b \) and the concentration ratio.
    • Calculations: Use equilibrium constants to solve for buffer pH.
    • Dilution Effect: Explain how dilution impacts buffer pH.

3.2 Electron Transfer

  • Oxidation and Reduction:
    • Oxidation: Loss of electrons, gain of oxygen, or loss of hydrogen.
    • Reduction: Gain of electrons, loss of oxygen, or gain of hydrogen.
    • Oxidation States: Determine the oxidation state of atoms and identify the oxidizing/reducing agents in reactions.
  • Redox Equations:
    • Half-Equations: Deduce redox half-equations in acidic or neutral conditions.
    • Overall Equations: Combine half-equations to get the full redox reaction.
  • Ease of Oxidation/Reduction:
    • Metals: Predict oxidation ease based on reactivity.
    • Halogens: Predict reduction ease based on their position in the periodic table.
  • Metal-Acid Reactions:
    • Reaction: Acids react with reactive metals to release hydrogen gas.
    • Equations: Write balanced equations for reactions with HCl and H₂SO₄.
  • Electrochemical Cells:
    • Voltaic Cells: Oxidation at the anode, reduction at the cathode.
    • Electrode Identification: Assign anode/cathode roles and their polarities in voltaic and electrolytic cells.
  • Primary (Voltaic) Cells:
    • Function: Convert energy from spontaneous redox reactions into electrical energy.
    • Construction: Understand components like half-cells, anode, cathode, electric circuit, and salt bridge.
  • Secondary (Rechargeable) Cells:
    • Redox Reactions: Can be reversed with electrical energy.
    • Comparison: Assess pros and cons of fuel cells, primary cells, and secondary cells.
  • Electrolytic Cells:
    • Function: Convert electrical energy to chemical energy via non-spontaneous reactions.
    • Electrolysis Products: Predict products of molten salt electrolysis.
  • Oxidation of Alcohols:
    • Primary Alcohols: Oxidized to aldehydes, then to carboxylic acids.
    • Experimental Set-up: Understand distillation and reflux techniques.
  • Reduction of Carboxylic Acids and Ketones:
    • Carboxylic Acids: Reduced to primary alcohols via aldehydes.
    • Ketones: Reduced to secondary alcohols, typically using hydride ions.
  • Reduction of Unsaturated Compounds:
    • Alkenes/Alkynes: Reduction with hydrogen reduces the degree of unsaturation.

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  • Standard Hydrogen Electrode:
    • Electrode Potential: Assigned a value of zero for standardization.
    • Interpretation: Compare ease of oxidation/reduction using standard electrode potentials.
  • Standard Cell Potential (E°cell):
    • Calculation: From standard electrode potentials.
    • Spontaneity: Positive E°cell indicates a spontaneous reaction.
  • Relationship Between ΔG° and E°cell:
    • Equation: ΔG° = -nFE°cell.
    • Determination: Calculate ΔG° using E°cell data.
  • Electrolysis of Aqueous Solutions:
    • Competing Reactions: Predict anode/cathode products based on electrode potentials.
    • Factors: Concentration and electrode nature impact products.
  • Electroplating:
    • Process: Electrolytic deposition of a thin metal layer on an object.
    • Electrode Reactions: Write equations for reactions during electroplating.

3.3 Electron Sharing

  • Radicals:

    • Definition: A radical is a molecular entity with an unpaired electron, making it highly reactive.
    • Identification: Represent radicals with a dot (•) to denote the unpaired electron.

  • Formation of Radicals:

    • Homolytic Fission: Radicals are produced by homolytic fission, where a covalent bond breaks evenly, splitting electrons between atoms.
    • Conditions: This process is typically initiated by ultraviolet (UV) light or heat.
    • Example: Homolytic fission of halogens (e.g., Cl₂ → 2Cl•) is the initiation step in radical chain reactions.

  • Radical Substitution Reactions:

    • Reactivity with Alkanes: Radicals react with alkanes in substitution reactions, often leading to a mixture of products.
    • Mechanism:
      • Propagation: Radicals continue to react with molecules, creating new radicals and sustaining the reaction.
      • Termination: The reaction ends when two radicals combine, forming a stable product.
    • Alkane Stability: Alkanes are relatively stable due to the strength of C–C and C–H bonds and their non-polar nature.

3.4 Electron-Pair Sharing

  • Nucleophiles:

    • Definition: A nucleophile is a reactant that donates an electron pair to form a bond with an electrophile.
    • Recognition: Nucleophiles can be neutral or negatively charged species.

  • Nucleophilic Substitution:

    • Reactions: In nucleophilic substitution, a nucleophile replaces a leaving group in a molecule. Electron pair movement is depicted with arrows in reaction mechanisms.

  • Heterolytic Fission:

    • Definition: The breaking of a covalent bond where both electrons stay with one atom, forming ions.

  • Electrophiles:

    • Definition: An electrophile accepts an electron pair from a nucleophile to form a bond.
    • Recognition: Electrophiles can be neutral or positively charged species.

  • Electrophilic Attack on Alkenes:

    • Reactivity: Alkenes are susceptible to electrophilic attack due to the electron-rich double bond.
    • Reactions: Alkenes react with water, halogens, and hydrogen halides, where the double bond opens up to form new products.
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  • Lewis Acids and Bases:

    • Definition: A Lewis acid accepts an electron pair, and a Lewis base donates one. Nucleophiles are Lewis bases, and electrophiles are Lewis acids.
    • Coordination Bonds: Formed when a Lewis base donates an electron pair to a Lewis acid.

  • Coordination Complexes:

    • Formation: Ligands donate electron pairs to transition metal cations, forming complex ions. The charge on the complex ion can be deduced from the ion and ligand formula.

  • Mechanisms of Halogenoalkane Reactions:

    • SN1 and SN2 Reactions: Primary halogenoalkanes undergo one-step SN2 reactions, while tertiary halogenoalkanes undergo two-step SN1 reactions. SN2 reactions are stereospecific.
    • Rate Predictions: Relative rates of substitution for different halogenoalkanes can be predicted.

  • Mechanisms of Alkene Reactions:

    • Symmetrical Alkenes: Reactions with halogens, water, and hydrogen halides involve electrophilic addition, leading to specific products.
    • Unsymmetrical Alkenes: The major product can be predicted based on the stability of the carbocation intermediate formed during the reaction.

  • Electrophilic Substitution in Benzene:

    • Mechanism: Benzene undergoes electrophilic substitution with electrophiles, where the aromatic ring is preserved, and a new substituent replaces a hydrogen atom.
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