Structure 1. Models of the particulate nature of matter

Structure 1.1—Introduction to the particulate nature of matter

• Structure 1.1.1—Elements are the primary constituents of matter, which cannot be chemically broken down into simpler substances. (SL and HL)
• Structure 1.1.2—The kinetic molecular theory is a model to explain physical properties of matter (solids, liquids and gases) and changes of state. (SL and HL)
• Structure 1.1.3—The temperature, T, in Kelvin (K) is a measure of average kinetic energy E_k of particles (SL and HL)

Structure 1.2—The nuclear atom

• Structure 1.2.2—Isotopes are atoms of the same element with different numbers of neutrons. (SL and HL)

Structure 1.3—Electron configurations

• Structure 1.3.1—Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels. (SL and HL)
• Structure 1.3.2—The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. (SL and HL)
• Structure 1.3.6—In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization (AHL)

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. (SL and HL)
• 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. (SL and HL)
• Structure 1.4.5—The molar concentration is determined by the amount of solute and the volume of solution. (SL and HL)
• 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. (SL and HL)

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. (SL and HL)
• Structure 1.5.2—Real gases deviate from the ideal gas model, particularly at low temperature and high pressure. (SL and HL)
• Structure 1.5.3—The molar volume of an ideal gas is a constant at a specific temperature and pressure. (SL and HL)
• 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 P1V1/T1=P2V2/T2 (SL and HL)

Structure 2. Models of bonding and structure

Structure 2.1—The ionic model

• Structure 2.1.2—The ionic bond is formed by electrostatic attractions between oppositely charged ions. (SL and HL)

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. (SL and HL)
• 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. (SL and HL)
• Structure 2.2.9—Given comparable molar mass, the relative strengths of intermolecular forces are generally: London (dispersion) forces < dipole–dipole forces < hydrogen bonding. (SL and HL)
• 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 (SL and HL)

Structure 3. Classification of matter

Structure 3.2—Functional groups: Classification of organic compounds

• Structure 3.2.4—Successive members of a homologous series show a trend in physical properties. (SL and HL)
• Structure 3.2.12—Data from different techniques are often combined in structural analysis. (AHL)

Reactivity 1. What drives chemical reactions?

Reactivity 1.1—Measuring enthalpy changes

• Reactivity 1.1.3—The relative stability of reactants and products determines whether reactions are endothermic or exothermic. (SL and HL)
• 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. (SL and HL)

Reactivity 1.2—Energy cycles in reactions

• Reactivity 1.2.3—Standard enthalpy changes of combustion, ΔH_c⦵, and formation, ΔH_f⦵, data are used in thermodynamic calculations. (AHL)
• 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. (AHL)

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

Reactivity 2.1—How much? The amount of chemical change

• 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. (SL and HL)
• Reactivity 2.1.4—The percentage yield is calculated from the ratio of experimental yield to theoretical yield. (SL and HL)

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. (SL and HL)
• Reactivity 2.2.3—Factors that influence the rate of a reaction include pressure, concentration, surface area, temperature and the presence of a catalyst. (SL and HL)
• Reactivity 2.2.4—Activation energy, Ea, is the minimum energy that colliding particles need for a successful collision leading to a reaction. (SL and HL)
• 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. (AHL)
• Reactivity 2.2.11—The rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction. (AHL)
• Reactivity 2.2.12—The Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy. (AHL)

Reactivity 3. What are the mechanisms of chemical change?

Reactivity 3.2—Electron transfer reactions

• Reactivity 3.2.5—Oxidation occurs at the anode and reduction occurs at the cathode in electrochemical cells. (SL and HL)
• Reactivity 3.2.6—A primary (voltaic) cell is an electrochemical cell that converts energy from spontaneous redox reactions to electrical energy. (SL and HL)
• Reactivity 3.2.7—Secondary (rechargeable) cells involve redox reactions that can be reversed using electrical energy. (SL and HL)