DP Chemistry (first assessment 2025)

Syllabus sections

Structure 1. Models of the particulate nature of matter
Structure 1.1—Introduction to the particulate nature of matter
Structure 1.1.1- Elements, compounds and mixtures. Elements are the primary constituents of matter, which cannot be chemically broken down into simpler substances. Compounds consist of atoms of different elements chemically bonded together in a fixed ratio. Mixtures contain more than one element or compound in no fixed ratio, which are not chemically bonded and so can be separated by physical methods. Distinguish between the properties of elements, compounds and mixtures.
Structure 1.1.2—Kinetic molecular theory. The kinetic molecular theory is a model to explain physical properties of matter (solids, liquids and gases) and changes of state. Distinguish the different states of matter. Use state symbols (s, , g and aq) in chemical equations.
Structure 1.1.3 - Temperature and average kinetic energy. <br>The temperature, T, in Kelvin (K) is a measure of average kinetic energy Ek of particles. Interpret observable changes in physical properties and temperature during changes of state. Convert between values in the Celsius and Kelvin scales.
Structure 1.2—The nuclear atom
Structure 1.2.1 - Atoms and negatively charged electrons. Atoms contain a positively charged, dense nucleus composed of protons and neutrons (nucleons). Negatively charged electrons occupy the space outside the nucleus. Use the nuclear symbol ZXA to deduce the number of protons, neutrons and electrons in atoms and ions.
Structure 1.2.2—Isotopes. Isotopes are atoms of the same element with different numbers of neutrons. Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data.
Structure 1.2.3—Mass spectra. Mass spectra are used to determine the relative atomic masses of elements from their isotopic composition. Interpret mass spectra in terms of identity and relative abundance of isotopes.
Structure 1.3—Electron configurations
Structure 1.3.1—Emission spectra. Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels. Qualitatively describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum. Distinguish between a continuous and a line spectrum.
Structure 1.3.2—The line emission spectrum of the hydrogen atom. The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. 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.
Structure 1.3.3—The main energy level. The main energy level is given an integer number, n, and can hold a maximum of 2n2 electrons. Deduce the maximum number of electrons that can occupy each energy level.
Structure 1.3.4—Detailed model of the atom. A more detailed model of the atom describes the division of the main energy level into s, p, d and f sublevels of successively higher energies. Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals.
Structure 1.3.5—Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin. Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.
Structure 1.3.6—In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization. Explain the trends and discontinuities in first ionization energy (IE) across a period and down a group. Calculate the value of the first IE from spectral data that gives the wavelength or frequency of the convergence limit.
Structure 1.3.7—Successive ionization energy (IE) data for an element give information about its electron configuration. Deduce the group of an element from its successive ionization data.
Structure 1.4—Counting particles by mass: The mole
Structure 1.4.1—The mole (mol) is the SI unit of amount of substance. One mole contains exactly the number of elementary entities given by the Avogadro constant. Convert the amount of substance, n, to the number of specified elementary entities.
Structure 1.4.2—Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass Ar and relative formula mass Mr . Determine relative formula masses Mr from relative atomic masses Ar .
Structure 1.4.3—Molar mass M has the units g mol–1. Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
Structure 1.4.4—The empirical formula of a compound gives the simplest ratio of atoms of each element present in that compound. The molecular formula gives the actual number of atoms of each element present in a molecule. Interconvert the percentage composition by mass and the empirical formula. Determine the molecular formula of a compound from its empirical formula and molar mass.
Structure 1.4.5—The molar concentration is determined by the amount of solute and the volume of solution. Solve problems involving the molar concentration, amount of solute and volume of solution.
Structure 1.4.6—Avogadro’s law states that equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules. Solve problems involving the mole ratio of reactants and/or products and the volume of gases.
Structure 1.5—Ideal gases
Structure 1.5.1—An ideal gas consists of moving particles with negligible volume and no intermolecular forces. All collisions between particles are considered elastic. Recognize the key assumptions in the ideal gas model.
Structure 1.5.2—Real gases deviate from the ideal gas model, particularly at low temperature and high pressure. Explain the limitations of the ideal gas model.
Structure 1.5.3—The molar volume of an ideal gas is a constant at a specific temperature and pressure. Investigate the relationship between temperature, pressure and volume for a fixed mass of an ideal gas and analyse graphs relating these variables.
Structure 1.5.4—The relationship between the pressure, volume, temperature and amount of an ideal gas is shown in the ideal gas equation PV = nRT and the combined gas law P1V1T1= P2V2T2. Solve problems relating to the ideal gas equation.
Structure 2. Models of bonding and structure
Structure 2.1—The ionic model
Structure 2.1.1—When metal atoms lose electrons, they form positive ions called cations. When non-metal atoms gain electrons, they form negative ions called anions. Predict the charge of an ion from the electron configuration of the atom.
Structure 2.1.2—The ionic bond is formed by electrostatic attractions between oppositely charged ions. Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “ide”. Interconvert names and formulas of binary ionic compounds.
Structure 2.1.3—Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas. Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility.
Structure 2.2—The covalent model
Structure 2.2.1—A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei. The octet rule refers to the tendency of atoms to gain a valence shell with a total of 8 electrons. Deduce the Lewis formula of molecules and ions for up to four electron pairs on each atom.
Structure 2.2.2—Single, double and triple bonds involve one, two and three shared pairs of electrons respectively. Explain the relationship between the number of bonds, bond length and bond strength.
Structure 2.2.3—A coordination bond is a covalent bond in which both the electrons of the shared pair originate from the same atom. Identify coordination bonds in compounds.
Structure 2.2.4—The valence shell electron pair repulsion (VSEPR) model enables the shapes of molecules to be predicted from the repulsion of electron domains around a central atom. Predict the electron domain geometry and the molecular geometry for species with up to four electron domains.
Structure 2.2.5—Bond polarity results from the difference in electronegativities of the bonded atoms. Deduce the polar nature of a covalent bond from electronegativity values.
Structure 2.2.6—Molecular polarity depends on both bond polarity and molecular geometry. Deduce the net dipole moment of a molecule or ion by considering bond polarity and molecular geometry.
Structure 2.2.7—Carbon and silicon form covalent network structures. Describe the structures and explain the properties of silicon, silicon dioxide and carbon’s allotropes: diamond, graphite, fullerenes and graphene.
Structure 2.2.8—The nature of the force that exists between molecules is determined by the size and polarity of the molecules. Intermolecular forces include London (dispersion), dipole-induced dipole, dipole–dipole and hydrogen bonding. Deduce the types of intermolecular force present from the structural features of covalent molecules.
Structure 2.2.9—Given comparable molar mass, the relative strengths of intermolecular forces are generally: London (dispersion) forces < dipole–dipole forces < hydrogen bonding. Explain the physical properties of covalent substances to include volatility, electrical conductivity and solubility in terms of their structure.
Structure 2.2.10—Chromatography is a technique used to separate the components of a mixture based on their relative attractions involving intermolecular forces to mobile and stationary phases. Explain, calculate and interpret the retardation factor values, RF.
Structure 2.2.11—Resonance structures occur when there is more than one possible position for a double bond in a molecule. Deduce resonance structures of molecules and ions.
Structure 2.2.12—Benzene, C6H6, is an important example of a molecule that has resonance. Discuss the structure of benzene from physical and chemical evidence.
Structure 2.2.13—Some atoms can form molecules in which they have an expanded octet of electrons. Visually represent Lewis formulas for species with five and six electron domains around the central atom. Deduce the electron domain geometry and the molecular geometry for these species using the VSEPR model.
Structure 2.2.14—Formal charge values can be calculated for each atom in a species and used todetermine which of several possible Lewis formulas is preferred. Apply formal charge to determine a preferred Lewis formula from different Lewis formulas for a species.
Structure 2.2.15—Sigma bonds σ form by the head-on combination of atomic orbitals where the electron density is concentrated along the bond axis. Pi bonds π form by the lateral combination of p-orbitals where the electron density is concentrated on opposite sides of the bond axis. Deduce the presence of sigma bonds and pi bonds in molecules and ions.
Structure 2.2.16—Hybridization is 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.
Structure 2.3—The metallic model
Structure 2.3.1—A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons. Explain the electrical conductivity, thermal conductivity and malleability of metals.
Structure 2.3.2—The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion. Explain trends in melting points of s and p block metals.
Structure 2.3.3—Transition elements have delocalized d-electrons. Explain the high melting point and electrical conductivity of transition elements.
Structure 2.4—From models to materials
Structure 2.4.1—Bonding is best described as a continuum between the ionic, covalent and metallic models, and can be represented by a bonding triangle. Use bonding models to explain the properties of a material.
Structure 2.4.2—The position of a compound in the bonding triangle is determined by the relative contributions of the three bonding types to the overall bond. Determine the position of a compound in the bonding triangle from electronegativity data. Predict the properties of a compound based on its position in the bonding triangle.
Structure 2.4.3—Alloys are mixtures of a metal and other metals or non-metals. They have enhanced properties. Explain the properties of alloys in terms of non-directional bonding.
Structure 2.4.4—Polymers are large molecules, or macromolecules, made from repeating subunits called monomers. Describe the common properties of plastics in terms of their structure.
Structure 2.4.5—Addition polymers form by the breaking of a double bond in each monomer. Represent the repeating unit of an addition polymer from given monomer structures.
Structure 2.4.6—Condensation polymers form by the reaction between functional groups in each monomer with the release of a small molecule. Represent the repeating unit of polyamides and polyesters from given monomer structures.
Structure 3. Classification of matter
Structure 3.1—The periodic table: Classification of elements
Structure 3.1.1—The periodic table consists of periods, groups and blocks. Identify the positions of metals, metalloids and non-metals in the periodic table.
Structure 3.1.2—The period number shows the outer energy level that is occupied by electrons. Elements in a group have a common number of valence electrons. Deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa.
Structure 3.1.3—Periodicity refers to trends in properties of elements across a period and down a group.
Structure 3.1.4—Trends in properties of elements down a group include the increasing metallic character of group 1 elements and decreasing non-metallic character of group 17 elements. Describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions.
Structure 3.1.5—Metallic and non-metallic properties show a continuum. This includes the trend from basic metal oxides through amphoteric to acidic non-metal oxides. Deduce equations for the reactions with water of the oxides of group 1 and group 2 metals, carbon and sulfur.
Structure 3.1.6—The oxidation state is a number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that atom would have if the compound were composed of ions. Deduce the oxidation states of an atom in an ion or a compound.
Structure 3.1.7—Discontinuities occur in the trend of increasing first ionization energy across a period. Explain how these discontinuities provide evidence for the existence of energy sublevels.
Structure 3.1.8—Transition elements have incomplete d-sublevels that give them characteristic properties. Recognize properties, including: variable oxidation state, high melting points, magnetic properties, catalytic properties, formation of coloured compounds and formation of complex ions with ligands.
Structure 3.1.9—The formation of variable oxidation states in transition elements can be explained by the fact that their successive ionization energies are close in value. Deduce the electron configurations of ions of the first-row transition elements.
Structure 3.1.10—Transition element complexes are coloured due to the absorption of light when an electron is promoted between the orbitals in the split d-sublevels. The colour absorbed is complementary to the colour observed. Apply the colour wheel to deduce the wavelengths and frequencies of light absorbed and/or observed.
Structure 3.2—Functional groups: Classification of organic compounds
Structure 3.2.1—Organic compounds can be represented by different types of formulas. These include empirical, molecular, structural (full and condensed), stereochemical and skeletal. Identify different formulas and interconvert molecular, skeletal and structural formulas. Construct 3D models (real or virtual) of organic molecules.
Structure 3.2.2—Functional groups give characteristic physical and chemical properties to a compound. Organic compounds are divided into classes according to the functional groups present in their molecules. Identify the following functional groups by name and structure: halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester, phenyl.
Structure 3.2.3—A homologous series is a family of compounds in which successive members differ by a common structural unit, typically CH2. Each homologous series 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.
Structure 3.2.4—Successive members of a homologous series show a trend in physical properties. Describe and explain the trend in melting and boiling points of members of a homologous series.
Structure 3.2.5—“IUPAC nomenclature” refers to a set of rules used by the International Union of Pure and Applied Chemistry to apply systematic names to organic and inorganic compounds. 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.
Structure 3.2.6—Structural isomers are molecules that have the same molecular formula but different connectivities. Recognize isomers, including branched, straight-chain, position and functional group isomers.
Structure 3.2.7—Stereoisomers have the same constitution (atom identities, connectivities and bond multiplicities) but different spatial arrangements of atoms. Describe and explain the features that give rise to cis-trans isomerism; recognize it in non-cyclic alkenes and C3 and C4 cycloalkanes. Draw stereochemical formulas showing the tetrahedral arrangement around a chiral carbon. Describe and explain a chiral carbon atom giving rise to stereoisomers with different optical properties. Recognize a pair of enantiomers as non-superimposable mirror images from 3D modelling (real or virtual).
Structure 3.2.8—Mass spectrometry (MS) of organic compounds can cause fragmentation of molecules. Deduce information about the structural features of a compound from specific MS fragmentation patterns.
Structure 3.2.9—Infrared (IR) spectra can be used to identify the type of bond present in a molecule. Interpret the functional group region of an IR spectrum, using a table of characteristic frequencies (wavenumber/cm–1).
Structure 3.2.10—Proton nuclear magnetic resonance spectroscopy (1H NMR) gives information on the different chemical environments of hydrogen atoms in a molecule. Interpret 1H NMR spectra to deduce the structures of organic molecules from the number of signals, the chemical shifts, and the relative areas under signals (integration traces).
Structure 3.2.11—Individual signals can be split into clusters of peaks. Interpret 1H NMR spectra from splitting patterns showing singlets, doublets, triplets and quartets to deduce greater structural detail. Data for interpretation of 1H NMR spectra are given in the data booklet.
Structure 3.2.12—Data from different techniques are often combined in structural analysis. Interpret a variety of data, including analytical spectra, to determine the structure of a molecule.
Reactivity 1. What drives chemical reactions?
Reactivity 1.1—Measuring enthalpy changes
Reactivity 1.1.1—Chemical reactions involve a transfer of energy between the system and the surroundings, while total energy is conserved. Understand the difference between heat and temperature.
Reactivity 1.1.2—Reactions are described as endothermic or exothermic, depending on the direction of energy transfer between the system and the surroundings. Understand the temperature change (decrease or increase) that accompanies endothermic and exothermic reactions, respectively.
Reactivity 1.1.3—The relative stability of reactants and products determines whether reactions are endothermic or exothermic. Sketch and interpret energy profiles for endothermic and exothermic reactions.
Reactivity 1.1.4—The standard enthalpy change for a chemical reaction, ΔH⦵, refers to the heat transferred at constant pressure under standard conditions and states. It can be determined from the change in temperature of a pure substance. Apply the equations Q = mcΔT and ΔH = − Qn in the calculation of the enthalpy change of a reaction.
Reactivity 1.2—Energy cycles in reactions
Reactivity 1.2.1—Bond-breaking absorbs and bond-forming releases energy. Calculate the enthalpy change of a reaction from given average bond enthalpy data.
Reactivity 1.2.2—Hess’s law states that the enthalpy change for a reaction is independent of the pathway between the initial and final states. Apply Hess’s law to calculate enthalpy changes in multistep reactions.
Reactivity 1.2.3—Standard enthalpy changes of combustion, ΔHc⦵, and formation, ΔHf⦵, data are used in thermodynamic calculations. Deduce equations and solutions to problems involving these terms.
Reactivity 1.2.4—An application of Hess’s law uses enthalpy of formation data or enthalpy of combustion data to calculate the enthalpy change of a reaction. Calculate enthalpy changes of a reaction using ΔHf⦵ data or ΔHc⦵ data: ΔH⦵ = Σ ΔHf⦵products − Σ ΔHf⦵reactants ΔH⦵ = Σ ΔHc, ⦵reactants − Σ ΔHc. ⦵products
Reactivity 1.2.5—A Born–Haber cycle is an application of Hess’s law, used to show energy changes in the formation of an ionic compound. Interpret and determine values from a Born–Haber cycle for compounds composed of univalent and divalent ions.
Reactivity 1.3—Energy from fuels
Reactivity 1.3.1—Reactive metals, non-metals and organic compounds undergo combustion reactions when heated in oxygen. Deduce equations for reactions of combustion, including hydrocarbons and alcohols.
Reactivity 1.3.2—Incomplete combustion of organic compounds, especially hydrocarbons, leads to the production of carbon monoxide and carbon. Deduce equations for the incomplete combustion of hydrocarbons and alcohols.
Reactivity 1.3.3—Fossil fuels include coal, crude oil and natural gas, which have different advantages and disadvantages. Evaluate the amount of carbon dioxide added to the atmosphere when different fuels burn. Understand the link between carbon dioxide levels and the greenhouse effect.
Reactivity 1.3.4—Biofuels are produced from the biological fixation of carbon over a short period of time through photosynthesis. Understand the difference between renewable and non-renewable energy sources. Consider the advantages and disadvantages of biofuels.
Reactivity 1.3.5—A fuel cell can be used to convert chemical energy from a fuel directly to electrical energy. Deduce half-equations for the electrode reactions in a fuel cell.
Reactivity 1.4—Entropy and spontaneity (Additional higher level)
Reactivity 1.4.1—Entropy, S, is a measure of the dispersal or distribution of matter and/or energy in a system. The more ways the energy can be distributed, the higher the entropy. Under the same conditions, the entropy of a gas is greater than that of a liquid, which in turn is greater than that of a solid. 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⦵.
Reactivity 1.4.2—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.
Reactivity 1.4.3—At constant pressure, a change is spontaneous if the change in Gibbs energy, ΔG, is negative. Interpret the sign of ΔG calculated from thermodynamic data. Determine the temperature at which a reaction becomes spontaneous.
Reactivity 1.4.4—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.
Reactivity 2. How much, how fast and how far?
Reactivity 2.1—How much? The amount of chemical change
Reactivity 2.1.1—Chemical equations show the ratio of reactants and products in a reaction. Deduce chemical equations when reactants and products are specified.
Reactivity 2.1.2—The mole ratio of an equation can be used to determine: • the masses and/or volumes of reactants and products • the concentrations of reactants and products for reactions occurring in solution. Calculate reacting masses and/or volumes and concentrations of reactants and products.
Reactivity 2.1.3—The limiting reactant determines the theoretical yield. Identify the limiting and excess reactants from given data.
Reactivity 2.1.4—The percentage yield is calculated from the ratio of experimental yield to theoretical yield. Solve problems involving reacting quantities, limiting and excess reactants, theoretical, experimental and percentage yields.
Reactivity 2.1.5—The atom economy is a measure of efficiency in green chemistry. Calculate the atom economy from the stoichiometry of a reaction.
Reactivity 2.2—How fast? The rate of chemical change
Reactivity 2.2.1—The rate of reaction is expressed as the change in concentration of a particular reactant/product per unit time.
Reactivity 2.2.2—Species react as a result of collisions of sufficient energy and proper orientation. Explain the relationship between the kinetic energy of the particles and the temperature in kelvin, and the role of collision geometry.
Reactivity 2.2.3—Factors that influence the rate of a reaction include pressure, concentration, surface area, temperature and the presence of a catalyst. Predict and explain the effects of changing conditions on the rate of a reaction.
Reactivity 2.2.4—Activation energy, Ea, is the minimum energy that colliding particles need for a successful collision leading to a reaction. Construct Maxwell–Boltzmann energy distribution curves to explain the effect of temperature on the probability of successful collisions.
Reactivity 2.2.5—Catalysts increase the rate of reaction by providing an alternative reaction pathway with lower Ea. Sketch and explain energy profiles with and without catalysts for endothermic and exothermic reactions. Construct Maxwell–Boltzmann energy distribution curves to explain the effect of different values for Ea on the probability of successful collisions.
Reactivity 2.2.6—Many reactions occur in a series of elementary steps. The slowest step determines the rate of the reaction. Evaluate proposed reaction mechanisms and recognize reaction intermediates. Distinguish between intermediates and transition states, and recognize both in energy profiles of reactions.
Reactivity 2.2.7—Energy profiles can be used to show the activation energy and transition state of the rate-determining step in a multistep reaction. Construct and interpret energy profiles from kinetic data.
Reactivity 2.2.8—The molecularity of an elementary step is the number of reacting particles taking part in that step. Interpret the terms “unimolecular”, “bimolecular” and “termolecular”.
Reactivity 2.2.9—Rate equations depend on the mechanism of the reaction and can only be determined experimentally. Deduce the rate equation for a reaction from experimental data.
Reactivity 2.2.10—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. The order with respect to a reactant can describe the number of particles taking part in the rate determining step. 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.
Reactivity 2.2.11—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, including the units of k.
Reactivity 2.2.12—The Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy. Describe the qualitative relationship between temperature and the rate constant. Analyse graphical representations of the Arrhenius equation, including its linear form.
Reactivity 2.2.13—The Arrhenius factor, A, takes into account the frequency of collisions with proper orientations. Determine the activation energy and the Arrhenius factor from experimental data.
Reactivity 2.3—How far? The extent of chemical change
Reactivity 2.3.1—A state of dynamic equilibrium is reached in a closed system when the rates of forward and backward reactions are equal. Describe the characteristics of a physical and chemical system at equilibrium.
Reactivity 2.3.2—The equilibrium law describes how the equilibrium constant, K, can be determined from the stoichiometry of a reaction. Deduce the equilibrium constant expression from an equation for a homogeneous reaction.
Reactivity 2.3.3—The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent. Determine the relationships between K values for reactions that are the reverse of each other at the same temperature.
Reactivity 2.3.4—Le Châtelier’s principle enables the prediction of the qualitative effects of changes in concentration, temperature and pressure to a system at equilibrium. Apply Le Ch.telier’s principle to predict and explain responses to changes of systems at equilibrium.
Reactivity 2.3.5—The reaction quotient, Q, is calculated using the equilibrium expression with nonequilibrium concentrations 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.
Reactivity 2.3.6—The equilibrium law is the basis for quantifying the composition of an equilibrium mixture. Solve problems involving values of K and initial and equilibrium concentrations of the components of an equilibrium mixture.
Reactivity 2.3.7—The equilibrium constant and Gibbs energy change, ΔG, can both be used to measure the position of an equilibrium reaction.
Reactivity 3. What are the mechanisms of chemical change?
Reactivity 3.1—Proton transfer reactions
Reactivity 3.1.1—Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor. Deduce the Br.nsted–Lowry acid and base in a reaction.
Reactivity 3.1.2—A pair of species differing by a single proton is called a conjugate acid–base pair. Deduce the formula of the conjugate acid or base of any Br.nsted–Lowry base or acid.
Reactivity 3.1.3—Some species can act as both Brønsted–Lowry acids and bases. Interpret and formulate equations to show acid–base reactions of these species.
Reactivity 3.1.4—The pH scale can be used to describe the [H+] of a solution: pH = –log10[H+]; [H+] = 10–pH. Perform calculations involving the logarithmic relationship between pH and [H+].
Reactivity 3.1.5—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–].
Reactivity 3.1.6—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. and are strong acids, and group 1 hydroxides are strong bases.
Reactivity 3.1.7—Acids react with bases in neutralization reactions. Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates.
Reactivity 3.1.8—pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features. Sketch and interpret the general shape of the pH curve.
Reactivity 3.1.9—The pOH scale describes the [OH–] of a solution. pOH = –log10[OH–]; [OH–] = 10–pOH. Interconvert [H+], [OH–], pH and pOH values.
Reactivity 3.1.10—The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values. Interpret the relative strengths of acids and bases from these data.
Reactivity 3.1.11—For a conjugate acid–base pair, the relationship Ka Å~ Kb = Kw can be derived from the expressions for Ka and Kb. Solve problems involving these values.
Reactivity 3.1.12—The pH of a salt solution depends on the relative strengths of the parent acid and base. Construct equations for the hydrolysis of ions in a salt, and predict the effect of each ion on the pH of the salt solution.
Reactivity 3.1.13—pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features. Interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases.
Reactivity 3.1.14—Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. The pH of the end point of an indicator, where it changes colour, approximately corresponds to its pKa value. Construct equilibria expressions to show why the colour of an indicator changes with pH.
Reactivity 3.1.15—An appropriate indicator for a titration has an end point range that coincides with the pH at the equivalence point. Identify an appropriate indicator for a titration from the identity of the salt and the pH range of the indicator.
Reactivity 3.1.16—A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali. Describe the composition of acidic and basic buffers and explain their actions.
Reactivity 3.1.17—The pH of a buffer solution depends on both: • the pKa or pKb of its acid or base • the ratio of the concentration of acid or base to the concentration of the conjugate base or acid. Solve problems involving the composition and pH of a buffer solution, using the equilibrium constant.
Reactivity 3.2—Electron transfer reactions
Reactivity 3.2.1—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.
Reactivity 3.2.2—Half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons. Deduce redox half-equations and equations in acidic or neutral solutions.
Reactivity 3.2.3—The relative ease of oxidation and reduction of an element in a group can be predicted from its position in the periodic table. The reactions between metals and aqueous metal ions demonstrate the relative ease of oxidation of different metals. Predict the relative ease of oxidation of metals. Predict the relative ease of reduction of halogens. Interpret data regarding metal and metal ion reactions.
Reactivity 3.2.4—Acids react with reactive metals to release hydrogen. Deduce equations for reactions of reactive metals with dilute and .
Reactivity 3.2.5—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.
Reactivity 3.2.6—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.
Reactivity 3.2.7—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.
Reactivity 3.2.8—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.
Reactivity 3.2.9—Functional groups in organic compounds may undergo oxidation. 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.
Reactivity 3.2.10—Functional groups in organic compounds may undergo reduction. Deduce equations to show reduction of carboxylic acids to primary alcohols via the aldehyde, and reduction of ketones to secondary alcohols.
Reactivity 3.2.11—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.
Reactivity 3.2.12—The hydrogen half-cell H+(aq) + e− ⇌12 H2(g) is assigned a standard electrode potential of zero by convention. It is used in the measurement of standard electrode potential, E⦵. Interpret standard electrode potential data in terms of ease of oxidation/reduction.
Reactivity 3.2.13—Standard cell potential, E⦵cell, can be calculated from standard electrode potentials. E⦵cell has a positive value for a spontaneous reaction. Predict whether a reaction is spontaneous in the forward or reverse direction from E⦵ data.
Reactivity 3.2.14—The equation ΔG⦵ = − nFE⦵ cell shows the relationship between standard change in Gibbs energy and standard cell potential for a reaction. Determine the value for ΔG⦵ from E⦵ data.
Reactivity 3.2.15—During electrolysis of aqueous solutions, competing reactions can occur at the anode and cathode, including the oxidation and reduction of water. Deduce from standard electrode potentials the products of the electrolysis of aqueous solutions.
Reactivity 3.2.16—Electroplating involves the electrolytic coating of an object with a metallic thin layer. Deduce equations for the electrode reactions during electroplating.
Reactivity 3.3—Electron sharing reactions
Reactivity 3.3.1—A radical is a molecular entity that has an unpaired electron. Radicals are highly reactive. Identify and represent radicals, e.g. and .
Reactivity 3.3.2—Radicals are produced by homolytic fission, e.g. of halogens, 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.
Reactivity 3.3.3—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.
Reactivity 3.4—Electron-pair sharing reactions
Reactivity 3.4.1—A nucleophile is a reactant that forms a bond to its reaction partner (the electrophile) by donating both bonding electrons. Recognize nucleophiles in chemical reactions.
Reactivity 3.4.2—In a nucleophilic substitution reaction, a nucleophile donates an electron pair to form a new bond, as another bond breaks producing a leaving group. Deduce equations with descriptions and explanations of the movement of electron pairs in nucleophilic substitution reactions.
Reactivity 3.4.3—Heterolytic fission is the breakage of a covalent bond when both bonding electrons remain with one of the two fragments formed. Explain, with equations, the formation of ions by heterolytic fission.
Reactivity 3.4.4—An electrophile is a reactant that forms a bond to its reaction partner (the nucleophile) by accepting both bonding electrons from that reaction partner. Recognize electrophiles in chemical reactions.
Reactivity 3.4.5—Alkenes are susceptible to electrophilic attack because of the high electron density of the carbon–carbon double bond. These reactions lead to electrophilic addition. Deduce equations for the reactions of alkenes with water, halogens, and hydrogen halides.
Reactivity 3.4.6—A Lewis acid is 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.
Reactivity 3.4.7—When a Lewis base reacts with a Lewis acid, a coordination bond is formed. 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.
Reactivity 3.4.8—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.
Reactivity 3.4.9—Nucleophilic substitution reactions include the reactions between halogenoalkanes and nucleophiles. Describe and explain the mechanisms of the reactions of primary and tertiary halogenoalkanes with nucleophiles.
Reactivity 3.4.10—The rate of the substitution reactions is influenced by the identity of the leaving group. Predict and explain the relative rates of the substitution reactions for different halogenoalkanes.
Reactivity 3.4.11—Alkenes readily undergo electrophilic addition reactions. Describe and explain the mechanisms of the reactions between symmetrical alkenes and halogens, water and hydrogen halides.
Reactivity 3.4.12—The relative stability of carbocations in the addition reactions between hydrogen halides and unsymmetrical alkenes can be used to explain the reaction mechanism. Predict and explain the major product of a reaction between an unsymmetrical alkene and a hydrogen halide or water.
Reactivity 3.4.13—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+.
Tools
Tool 1: Experimental techniques
Tool 2: Technology
Tool 3: Mathematics
Inquiry
Inquiry 1: Exploring and designing
Inquiry 2: Collecting and processing data
Inquiry 3: Concluding and evaluating
Nature of science