11.1 Acids and bases 11.2 A closer look at acids and alkalis 11.3 The reaction of acids and bases 11.4 A closer look at neutralisation 11.5 Oxides 11.6 Making Salts 11.7 Making insoluble salt by precipitation 11.8 Finding the concentration by titration

Compounds, mixtures, and chemical change Why do atoms form bonds? The ionic bond More about ions The covalent bond Covalent compounds Comparing ionic and covalent compounds Giant covalent structures The bonding in metals

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. Reactivity 3.1.2—A pair of species differing by a single proton is called a conjugate acid–base pair. Reactivity 3.1.3—Some species can act as both Brønsted–Lowry acids and bases. Reactivity 3.1.4—The pH scale can be used to describe the [H+] of a solution Reactivity 3.1.5—The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH–]. Kw = [H+] [OH–] Reactivity 3.1.6—Strong and weak acids and bases differ in the extent of ionization. Reactivity 3.1.7—Acids react with bases in neutralization reactions. Reactivity 3.1.8—pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features. Reactivity 3.1.9—The pOH scale describes the [OH–] of a solution Reactivity 3.1.10—The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values. Reactivity 3.1.11—For a conjugate acid–base pair, the relationship Ka × Kb = Kw can be derived from the expressions for Ka and Kb. Reactivity 3.1.12—The pH of a salt solution depends on the relative strengths of the parent acid and base. Reactivity 3.1.13—pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features. 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. Reactivity 3.1.15—An appropriate indicator for a titration has an end point range that coincides with the pH at the equivalence point. Reactivity 3.1.16—A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali. Reactivity 3.1.17—The pH of a buffer solution

9.1 Introducing reaction rates 9.2 Measuring the rate of reaction 9.3 / 9.4 Changing the rate 9.5 Explaining rate changes 9.6 Catalyst

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. Reactivity 2.3.2—The equilibrium law describes how the equilibrium constant, K, can be determined from the stoichiometry of a reaction. Reactivity 2.3.3—The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent. 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. Reactivity 2.3.5—The reaction quotient, Q, is calculated using the equilibrium expression with non- equilibrium concentrations of reactants and products. Reactivity 2.3.6—The equilibrium law is the basis for quantifying the composition 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 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. Reactivity 2.2.3—Factors that influence the rate of a reaction include pressure, concentration, surface area, temperature and the presence of a catalyst. Reactivity 2.2.4—Activation energy, Ea, is the minimum energy that colliding particles need for a successful collision leading to a reaction. Reactivity 2.2.5—Catalysts increase the rate of reaction by providing an alternative reaction pathway with lower Ea. Reactivity 2.2.6—Many reactions occur in a series of elementary steps. The slowest step determines the rate of the reaction. 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. Reactivity 2.2.8—The molecularity of an elementary step is the number of reacting particles taking part in that step. Reactivity 2.2.9—Rate equations depend on the mechanism of the reaction and can only be determined experimentally. 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. Reactivity 2.2.11—The rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction. Reactivity 2.2.12—The Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy. Reactivity 2.2.13—The Arrhenius factor, A, takes into account the frequency of collisions with proper orientations.

Energy changes in reactions Describing exothermic reactions Describing endothermic reactions A closer look at energy changes Reaction pathway diagram Activation Energy and Enthalpy Calculation Enthalpy changes The hydrogen-oxygen fuel cell

2.1 - Meet the elements 2.2 – More about atoms 2.3 – How electrons are arranged 2.4 –Isotops and Ar

Reactivity 1.2 - Energy cycles in reactions Reactivity 1.2.1 - Bond-breaking absorbs and bond-forming releases energy. 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. Reactivity 1.2.3 - Standard enthalpy changes of combustion, ΔHc ⦵, and formation, ΔHf ⦵, data are used in thermodynamic calculations. 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. 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.

5 - usings moles 5.1 The mole 5.2 Calculations from equations 5.3 Reactions involving gases 5.4 The concentration of a solution 5.5 Finding the empirical formula 5.6 From empirical to final formula 5.7 Finding % yield and % purity

1-States of matter 1-1 Everythings is made of particles 1-2 Solids, liquids, and gases 1.3 The particles in solids, Liquids, and gases 1.4 Heating and cooling curves 1.5 A closer look at gases

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. Reactivity 1.1.2—Reactions are described as endothermic or exothermic, depending on the direction of energy transfer between the system and the surroundings. Reactivity 1.1.3—The relative stability of reactants and products determines whether reactions are endothermic or exothermic. 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.

Structure 3.1—The periodic table: Classification of elements Structure 3.1.1—The periodic table consists of periods, groups and blocks. 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. 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. 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. 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. 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. 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. 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. Structure 3.2.4—Successive members of a homologous series show a trend in physical properties. 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. Structure 3.2.6—Structural isomers are molecules that have the same molecular formula but different connectivities.

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. Structure 2.3.2 : The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion. Structure 2.3.3 : Transition elements have delocalized d-electrons. 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. 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. Structure 2.4.3 : Alloys are mixtures of a metal and other metals or non-metals. They have enhanced properties. Structure 2.4.4 : Polymers are large molecules, or macromolecules, made from repeating subunits called monomers. Structure 2.4.5 : Addition polymers form by the breaking of a double bond in each monomer.Structure 2.4.6 : Condensation polymers form by the reaction between functional groups in each monomer with the release of a small molecule.

Structure 2 / IB Chemistry / Structure 2.2 (lesson / Worksheets / Tests/ Tables / Figures) 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. Structure 2.2.2—Single, double and triple bonds involve one, two and three shared pairs of electrons respectively. 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. 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. Structure 2.2.5—Bond polarity results from the difference in electronegativities of the bonded atoms. Structure 2.2.6—Molecular polarity depends on both bond polarity and molecular geometry. Structure 2.2.7—Carbon and silicon form covalent network structures. 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. Structure 2.2.9—Given comparable molar mass, the relative strengths of intermolecular forces are generally: London (dispersion) forces < dipole–dipole forces < hydrogen bonding. 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.

Structure 2 / IB Chemistry / Structure 2.1 +worksheets +Formulae of common ions / ionic compounds Structure 2. Models of bonding and structure 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. Structure 2.1.2 — The ionic bond is formed by electrostatic attractions between oppositely charged ions. Structure 2.1.3—Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas.

Structure 1 / IB Chemistry / Structure 1.5 (Including worksheets) 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. Structure 1.5.2 - Real gases deviate from the ideal gas model, particularly at low temperature and high pressure. Structure 1.5.3 - The molar volume of an ideal gas is a constant at a specific temperature and pressure. 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

Structure 1 / IB Chemistry / Structure 1.4 Structure 1.4—Counting particles by mass: The mole Structure 1.4.1 Structure 1.4.2 Structure 1.4.3 Structure 1.4.4 Structure 1.4.5 Structure 1.4.6

Structure 1.3 : Electron configurations Structure 1.3.1 : 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 : 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 : Deduce the maximum number of electrons that can occupy each energy level. Structure 1.3.4 : Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals. Structure 1.3.5 : 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 : 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 : Deduce the group of an element from its successive ionization data.

Structure 1 / IB Chemistry / Structure 1.2 Structure 1.2 : The nuclear atom Structure 1.2.1 : Use the nuclear symbol A Z X to deduce the number of protons, neutrons and electrons in atoms and ions. Structure 1.2.2 : Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data. Structure 1.2.3 : Interpret mass spectra in terms of identity and relative abundance of isotopes.