IB Chemistry Syllabus Checklist

Predict the charge of an ion from the electron configuration of the atom, including transition elements.

Define an ionic bond and deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Interconvert names and formulas of binary ionic compounds. Recall the following polyatomic ions by name and formula: ammonium NH₄⁺, hydroxide OH^–, nitrate NO₃^–, hydrogencarbonate HCO₃^–, carbonate CO₃^2–, sulfate SO₄^2–, phosphate PO₄^3–.

Describe the structure of ionic compounds and explain their physical properties including: volatility, electrical conductivity and solubility. Describe lattice enthalpy as a measure of the strength of the ionic bond in different compounds, influenced by ion radius and charge.

Understand the occurance of resonance structures. Deduce resonance structures of molecules and ions, including use of the term “delocalization”.

Discuss the structure of benzene from physical and chemical evidence.

Understand that atoms can form molecules in which they have an expanded octet of electrons. Visually represent Lewis formulas for species with 5 and 6 electron domains around the central atom. Deduce the electron domain geometry and the molecular geometry for species with 5 and 6 electron domains.

Apply formal charge to determine a preferred Lewis formula from different Lewis formulas for a species.

Define sigma bonds (σ) and Pi bonds (π). Deduce the presence of sigma bonds and pi bonds in molecules and ions through both organic and inorganic examples.

Define hybridization as the concept of mixing atomic orbitals to form new hybrid orbitals for bonding. Analyse the hybridization and bond formation in molecules and ions. Identify the relationships between Lewis formulas, electron domains, molecular geometry and type of hybridization. Predict the geometry around an atom from its hybridization, and vice versa.

Understand that the transition elements have delocalized d-electrons. Explain the high melting point and electrical conductivity of transition elements.

Understand the reactions involved in the formation of condensation polymers. Represent the repeating unit of polyamides and polyesters from given monomer structures.

Identify the positions of metals, metalloids and non-metals in the periodic table. Identify the four blocks associated with the sublevels s, p, d, and f.

Deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa. Recall the classifications of: “alkali metals”, “halogens”, “transition elements” and “noble gases”.

Define periodicity as trends in properties of elements across a period and down a group. Explain the periodicity of atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.

Describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions.

Deduce equations for the reactions with water of the oxides of group 1 and group 2 metals, carbon, and sulfur. Recall the formation of acid raid by gaseous non-metal oxides, and the acidiciation of oceans by increasing CO₂ levels.

Understand the concept of oxidation states and explain why the oxidation state of an element is zero. Deduce the oxidation states of an atom in an ion or a compound, including the examples of: hydrogen in metal hydrides (–1) and oxygen in peroxides (–1). Understand and use the terms: “oxidation number” and “oxidation state”.

Use of the chemical formulae: empirical, molecular, structural (full and condensed), stereochemical and skeletal. Interconvert molecular, skeletal and structural formulas. Construct 3D models (real or virtual) of organic molecules.

Understand the classification of organic compounds according to the functional groups. Identify the following functional groups by name and structure: halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester, phenyl. Understand and use the terms “saturated” and “unsaturated”.

Define a homologous series as a family of compounds in which successive members differ by a common structural unit, and that can be described by a general formula. Identify the following homologous series: alkanes, alkenes, alkynes, halogenoalkanes, alcohols, aldehydes, ketones, carboxylic acids, ethers, amines, amides and esters.

Describe and explain the trend in melting and boiling points of members of a homologous series.

Apply IUPAC nomenclature to saturated or mono-unsaturated compounds that have up to six carbon atoms in the parent chain and contain one type of the following functional groups: halogeno, hydroxyl, carbonyl, carboxyl. Name straight-chain and branched-chain isomers.

Define structural isomers as molecules that have the same molecular formula but different connectivities. Recognize isomers, including branched, straight-chain, position and functional group isomers, and primary, secondary and tertiary alcohols.

Understand that chemical reactions involve a transfer of energy. Understand the difference between heat and temperature.

Recall the classification of reactions as endothermic or exothermic. Understand the temperature change (decrease or increase) that accompanies endothermic and exothermic reactions, respectively.

Sketch and interpret energy profiles for endothermic and exothermic reactions.

Understand the standard enthalpy change for a chemical reaction, ΔH^⦵. Apply the equations Q = mcΔT and ΔH = -Qn in the calculation of the enthalpy change of a reaction.
 

Understand standard enthalpy changes of combustion, ΔH_c^⦵, and formation, ΔH_f^⦵. Deduce equations and solutions to problems involving these terms.

Calculate enthalpy changes of a reaction using ΔH_f^⦵ data or ΔHc^⦵ data:

ΔH⦵ = Σ ΔH_f^⦵products − Σ ΔH_f^⦵reactants

ΔH⦵ = Σ ΔHc^⦵reactants − Σ ΔH_f^⦵products.

Understand a Born–Haber cycle for an ionic compound. Interpret and determine values from a Born–Haber cycle for compounds composed of univalent and divalent ions, including: ionization energies, enthalpy of atomization, electron affinities, lattice enthalpy, enthalpy of formation. The construction of a complete Born–Haber cycle will not be assessed.

Define entropy, S, as a measure of the dispersal or distribution of matter and/or energy in a system. Predict whether a physical or chemical change will result in an increase or decrease in entropy of a system. Calculate standard entropy changes, ΔS^⦵, from standard entropy values, S^⦵.

Understand the change in Gibbs energy, ΔG, relates the energy that can be obtained from a chemical reaction to the change in enthalpy, ΔH, change in entropy, ΔS, and absolute temperature, T. Apply the equation ΔG^⦵ = ΔH^⦵ −TΔS^⦵ to calculate unknown values of these terms.

Interpret the sign of ΔG calculated from thermodynamic data. Determine the temperature at which a reaction becomes spontaneous.

Understand that as a reaction approaches equilibrium, ΔG becomes less negative and finally reaches zero. Perform calculations using the equation ΔG = ΔG^⦵ + RT lnQ and its application to a system at equilibrium ΔG^⦵ = −RT lnK.

Deduce chemical equations when reactants and products are specified, including the use of state symbols.

Use the mole ratio of an equation to calculate: the masses, volumes, and concentrations of reactants and products. Use Avogadro’s law and definitions of molar concentration.

Understand how the limiting reactant determines the theoretical yield. Identify the limiting and excess reactants from given data. Distinguish between the theoretical yield and the experimental yield.

Solve problems involving reacting quantities, limiting and excess reactants, theoretical, experimental and percentage yields.

Understand that atom economy is a measure of efficiency in green chemistry. Calculate the atom economy from the stoichiometry of a reaction. Understand the inverse relationship between atom economy and wastage in industrial processes.

Evaluate proposed reaction mechanisms and recognize reaction intermediates. Distinguish between intermediates and transition states, and recognize both in energy profiles of reactions.

Construct and interpret energy profiles from kinetic data.

Define the molecularity of an elementary step as the number of reacting particles taking part in that step. Interpret the terms “unimolecular”, “bimolecular” and “termolecular”.

Deduce the rate equation for a reaction from experimental data.

Recall that the order of a reaction with respect to a reactant is the exponent to which the concentration of the reactant is raised in the rate equation. Understand that the overall reaction order is the sum of the orders with respect to each reactant. Sketch, identify and analyse graphical representations of zero, first and second order reactions.

Recall that the rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction. Solve problems involving the rate equation.

Understand and use the Arrhenius equation. Describe the qualitative relationship between temperature and the rate constant. Analyse graphical representations of the Arrhenius equation, including its linear form.

Determine the activation energy and the Arrhenius factor from experimental data.
 

Understand the reaction quotient, Q, as a non-equilibrium expression of the ratio of reactants and products. Calculate the reaction quotient Q from the concentrations of reactants and products at a particular time, and determine the direction in which the reaction will proceed to reach equilibrium.

Solve problems involving values of K and initial and equilibrium concentrations of the components of an equilibrium mixture. Understand the approximation that [reactant]_initial ≈ [reactant]_eqm when K is very small.

Understand how the equilibrium constant and Gibbs energy change, ΔG, can both be used to measure the position of an equilibrium reaction. Perform calculations using the equation ΔG^⦵ = −RT lnK.

Define a Brønsted–Lowry acid as a proton donor and a Brønsted–Lowry base as a proton acceptor. Deduce the Brønsted–Lowry acid and base in a reaction. Distinguish between the terms “base” and “alkali”.

Define a pair of species differing by a single proton as a conjugate acid–base pair. Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.

Understand that some species can act as both Brønsted–Lowry acids and bases. Interpret and formulate equations to show acid–base reactions of these species.

Define the pH scale as pH = –log₁₀[H⁺] and [H⁺] = 10^–pH. Perform calculations involving the logarithmic relationship between pH and [H⁺]. Estimate the pH using universal indicator, and precisely measure pH using a pH meter/probe.

Explain how the ion product constant of water, Kw, shows an inverse relationship between [H⁺] and [OH^–]. Kw = [H⁺] [OH^–]. Recognize solutions as acidic, neutral and basic from the relative values of [H⁺] and [OH^–].

Explain how strong and weak acids and bases differ in the extent of ionization. Recognize that acid–base equilibria lie in the direction of the weaker conjugate. Recall that HCl, HBr, HNO₃, and H₂SO₄ are strong acids, and group 1 hydroxides are strong bases. Distinguish between strong and weak acids or bases and concentrated and dilute reagents.

Recall that acids react with bases in neutralization reactions. Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates. Identify the parent acid and base of different salts.

Recognise, sketch and interpret pH curves for neutralization reactions involving strong acids and bases. Identify appropriate intercepts with the pH axis and equivalence point.

Explain that oxidation and reduction can be described in terms of electron transfer, change in oxidation state, oxygen gain/loss or hydrogen loss/gain. Deduce oxidation states of an atom in a compound or an ion. Identify the oxidized and reduced species and the oxidizing and reducing agents in a chemical reaction. Utilise oxidation numbers in the naming of compounds.

Deduce redox half-equations and equations in acidic or neutral solutions.

Predict the relative ease of oxidation of metals.
Predict the relative ease of reduction of halogens. Interpret data regarding metal and metal ion reactions.

Recall that acids react with reactive metals to release hydrogen. Deduce equations for reactions of reactive metals with dilute HCl and H₂SO₄.

Recall that oxidation occurs at the anode and reduction occurs at the cathode in electrochemical cells. Identify electrodes as anode and cathode, and identify their signs/polarities in voltaic cells and electrolytic cells, based on the type of reaction occurring at the electrode.

Recall that a primary (voltaic) cell is an electrochemical cell that converts energy from spontaneous redox reactions to electrical energy. Explain the direction of electron flow from anode to cathode in the external circuit, and ion movement across the salt bridge. Construction of primary cells including: half- cells containing metal/metal ion, anode, cathode, electric circuit, salt bridge.

Understand the secondary (rechargeable) cells involve redox reactions that can be reversed using electrical energy. Deduce the reactions of the charging process from given electrode reactions for discharge, and vice versa. State advantages and disadvantages of fuel cells, primary cells and secondary cells.

Recall that an electrolytic cell is an electrochemical cell that converts electrical energy to chemical energy by bringing about non-spontaneous reactions. Explain how current is conducted in an electrolytic cell. Deduce the products of the electrolysis of a molten salt. Construct electrolytic cells including: a DC power source, an anode, a cathode, and the electrolyte.

Deduce equations to show changes in the functional groups during oxidation of primary and secondary alcohols, including the two-step reaction in the oxidation of primary alcohols. Explain the experimental set-up for distillation and reflux. The names and formulas of specific oxidizing agents, and the mechanisms of oxidation, will not be assessed.

Deduce equations to show reduction of carboxylic acids to primary alcohols via the aldehyde, and reduction of ketones to secondary alcohols. Include the role of hydride ions in the reduction reaction. The names and formulas of specific reducing agents, and the mechanisms of reduction, will not be assessed.

Recall that the reduction of unsaturated compounds by the addition of hydrogen lowers the degree of unsaturation. Deduce the products of the reactions of hydrogen with alkenes and alkynes.

Define a radical as a molecular entity that has an unpaired electron and understand that radicals are highly reactive. Identify and represent radicals.

Recall that radicals are produced by homolytic fission in the presence of ultraviolet (UV) light or heat. Explain, including with equations, the homolytic fission of halogens, known as the initiation step in a chain reaction.

Recall that radicals take part in substitution reactions with alkanes, producing a mixture of products. Explain, using equations, the propagation and termination steps in the reactions between alkanes and halogens. Explain the stability of alkanes due to the strengths of the C–C and C–H bonds and their essentially non-polar nature.

Define a Lewis acid as an electron-pair acceptor and a Lewis base is an electron-pair donor. Apply Lewis acid–base theory to inorganic and organic chemistry to identify the role of the reacting species.

Understand that when a Lewis base reacts with a Lewis acid, a coordination bond is formed. Recall that nucleophiles are Lewis bases and electrophiles are Lewis acids. Draw and interpret Lewis formulas of reactants and products to show coordination bond formation in Lewis acid–base reactions.

Recall that coordination bonds are formed when ligands donate an electron pair to transition element cations, forming complex ions. Deduce the charge on a complex ion, given the formula of the ion and ligands present.

Describe and explain the mechanisms of the reactions of primary and tertiary halogenoalkanes with nucleophiles. Distinguish between the one-step S_N2 reaction of primary halogenoalkanes and the two- step S_N1 reaction of tertiary halogenoalkanes. Understand the stereospecificity of S_N2 reactions .

Predict and explain the relative rates of the substitution reactions for different halogenoalkanes. The roles of the solvent and the reaction mechanism will not be assessed.

Describe and explain the mechanisms of the reactions between symmetrical alkenes and halogens, water and hydrogen halides.

Predict and explain the major product of a reaction between an unsymmetrical alkene and a hydrogen halide or water.

Recall that electrophilic substitution reactions include the reactions of benzene with electrophiles. Describe and explain the mechanism of the reaction between benzene and a charged electrophile, E⁺. The formation of the electrophile will not be assessed.