Understand	Apply their knowledge to
When metal atoms lose electrons, they form positive ions called cations (cations are positive ions that contain more protons than electrons. They can also be formed by neutral molecules, such as ammonia, gaining a proton. When non-metal atoms gain electrons, they form negative ions called anions. (anions are negative ions that contain more electrons than protons.) Ionic bonds are due to electrostatic attractions between oppositely charged ions. Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “-ide”. Under normal conditions, ionic compounds exist as three-dimensional lattice structures represented by empirical formulas. Lattice enthalpy is a measure of the strength of the ionic bond in different compounds, influenced by ion radius and charge.	Predict the charge of an ion from the electron configuration of the atom. Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Interconvert names and formulas of binary ionic compounds. Explain the physical properties (volatility, electrical conductivity and solubility) of ionic compounds in terms of their structure
IB Clarification Notes
The formation of ions having different charges from a transition metal element should be included. Common ions of the transition metals are given in the data booklet. Students should be familiar with formulas and the names of the following polyatomic ions: ammonium NH₄⁺, hydroxide OH^–, nitrate NO₃^–, hydrogencarbonate HCO₃^–, carbonate CO₃^2–, sulfate SO₄^2– and phosphate PO₄^3–. Include lattice enthalpy as a measure of the strength of the ionic bond in different compounds, influenced by ion radius and charge.

Structure 2.2—The covalent model

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Guiding question

What determines the covalent nature and properties of a substance?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
covalent bonds are formed by the electrostatic attraction between shared pairs of electrons and the positively charged nuclei on either side of the electrons. the “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons. single, double and triple covalent bonds involve one, two and three shared pairs of electrons respectively. a coordination bond is a covalent bond in which both the electrons of the shared pair originate from the same atom. bond strength increases and bond length decreases as the number of shared pairs of electrons increases. bond polarity is due to the difference in the electronegativities of the bonded atoms. the shapes of molecules and simple ions can be predicted by The Valence Shell Electron Pair Repulsion (VSEPR) model which considers the repulsion of electron domains around a central atom. molecular polarity depends on both bond polarity and molecular geometry. carbon and silicon form covalent network structures. the nature of the force that exists between molecules is determined by the size and polarity of these molecules. London (dispersion) forces, dipole-dipole forces and hydrogen bonding are all included in intermolecular forces. in terms of their relative strengths for molecules with comparable molar mass the order of these interactive forces is London (dispersion) forces < dipole-dipole forces < hydrogen bonds. 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.	deduce the Lewis formula (electron dot or Lewis structure) of molecules and ions for up to four electron pairs on each atom. explain the relationship between the number of bonds, bond length and bond strength. recognize coordination bonds in compounds. deduce the polar nature of a covalent bond using electronegativity values. use VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains. predict bond angles from molecular geometry and the presence of non-bonding pairs of electrons. deduce molecular polarity, i.e. the net dipole of a molecule or ion, from bond polarity and geometry. deduce using VSEPR theory the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles. describe the structures and explain the properties of silicon, silicon dioxide and carbon’s allotropes: diamond, graphite, fullerenes and graphene. deduce the types of intermolecular force present in substances, based on the structural features of covalent molecules. explain the physical properties of covalent compounds (to include volatility, electrical conductivity and solubility) in terms of their structure and intermolecular forces. explain, calculate and interpret the retardation factor, R_F values.
AHL	AHL
resonance structures occur when there is more than one possible position for a double bond in a molecule. benzene, C₆H₆, is an important example of a molecule which has resonance. some atoms can form molecules in which they have an expanded octet of electrons. formal charge values can be calculated for each atom in a species and used to determine which of several possible Lewis formulas is preferred. sigma (σ) bonds form by the head-on overlap of atomic orbitals. Electron density is concentrated along the bond axis. pi (π) bonds form by the lateral overlap of p-orbitals where the electron density is concentrated on opposite sides of the bond axis. hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals for bonding	deduce resonance structures of molecules and ions. discuss the structure of benzene from physical and chemical evidence. represent Lewis formulas for species with five and six electron domains around the central atom. predict the electron domain geometry and the molecular geometry for these species (covered in S2.2 The covalent model (2). apply formal charge to determine a preferred Lewis formula from different Lewis formulas for a species. deduce the presence of sigma bonds and pi bonds in molecules. analyse the hybridization and bond formation in molecules and ions. Identify the relationships between Lewis structures, electron domains, molecular geometry and type of hybridization. predict the shape around an atom from its hybridization, and vice versa.
IB Clarification Notes
Lewis formulas show all the valence electrons (bonding and non-bonding pairs) in a covalently bonded species. Electron pairs in a Lewis formula can be shown as dots, crosses and/or dashes. Organic and inorganic examples should be used. Molecules with atoms having fewer than an octet of electrons should be covered. Partial charges, dipoles or vectors can all be used to show bond polarity. Electronegativity values (according to Pauling) are given in the data booklet. Include prediction of how non-bonding pairs and multiple bonds affect bond angles. Examples of molecular polarity should include species in which bond dipoles do and do not cancel each other. AHL only: The shapes and molecular polarities of geometries corresponding to five and six electron domains should also be covered. Allotropes of the same element have different bonding and structural patterns, and so have different chemical and physical properties. The term “van der Waals’ forces” should be used as an inclusive term to include dipole-dipole, dipole-induced dipole, and London (dispersion) forces. London (dispersion) forces” describes the instantaneous induced dipole – induced dipole forces that exist between any atoms or groups of atoms and should be used for non-polar entities. Hydrogen bonds occur when hydrogen covalently bonded to an electronegative atom has an attractive interaction on a neighbouring electronegative atom with evidence of bond formation. A mixture can be separated using paper chromatography or thin layer chromatography (TLC). The use of locating agents is not required. AHL Include the term delocalization. Both organic and inorganic examples should be included. Only species (organic and inorganic) with sp³, sp² and sp hybridization need to be considered.

Structure 2.3—The metallic model

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Guiding question

What determines the metallic nature and properties of an element

Learning outcomes

After studying this topic students should be able to:

Understand	Apply their knowledge to
A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons. The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.	Explain the electrical and thermal conductivity and malleability of metals. Relate characteristic properties of metals to their uses Explain the trends in melting points of s and p block metals.
AHL	AHL
Transition metals have delocalized d-electrons.	Explain the high melting point and electrical conductivity of transition metals.
IB Clarification Notes
Relate characteristic properties of metals to their uses. A simple treatment in terms of charge of cations and electron density is required. Some chemical properties of transition metals are covered in R3.4.

Structure 2.4—From models to materials

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Guiding question

What role do bonding and structure have in the design of materials?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
Bonding is best described as a continuum between the ionic, covalent and metallic models, and can be represented by a bonding triangle. The position of a compound in the bonding triangle is determined by the relative contributions of the three bonding types to the overall bond. Alloys are mixtures of a metal and other metals or non-metals. They have enhanced properties. Polymers are large molecules, or macromolecules, made from repeating sub-units called monomers. Addition polymers form by the breaking of a double bond in each monomer.	Use bonding models to explain the properties of a material. 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. Explain the properties of alloys in terms of non-directional bonding. Describe the common properties of plastics in terms of their structure. Represent the repeating unit of an addition polymer from given monomer structures.
AHL	AHL
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. describe the hydrolysis and condensation reactions that break down and form biological molecules
IB Clarification Notes
A triangular bonding diagram is provided in the data booklet. To illustrate the relationship between bonding type and properties include examples of materials with varying percentage bonding character. Only binary compounds need to be considered. Calculations of percentage ionic character are not required. Electronegativity data are given in the data booklet. Illustrate alloys with common examples such as bronze, brass, and stainless steel. Include other examples of choice. Specific examples of alloys do not have to be learned. Examples of natural and synthetic polymers should be discussed. Examples of addition polymers should include polymerisation reactions of alkenes. Structures of monomers do not have to be learned but will be provided or will need to be deduced from the polymer. All biological macromolecules form by condensation reactions and break down by hydrolysis.

Structure 3. Classification of matter

Structure 3.1 - The Periodic Table

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Guiding question

How does the periodic table help us to predict patterns and trends in the properties of the elements?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
the periodic table consists of periods, groups and blocks the period number shows the outer energy level that is occupied by electrons elements in a group have a common number of valence electrons periodicity refers to trends in properties of elements across a period and down a group 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 metallic and non-metallic properties show a continuum; this includes the trend from basic metal oxides through amphoteric to acidic non-metal oxides 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	identify the positions of metals, metalloids and non-metals in the periodic table deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa 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 deduce the oxidation state of an atom in an ion or compound explain why the oxidation state of an element is zero
AHL	AHL
discontinuities occur in the trend of increasing first ionization energy across a period transition elements have incomplete d-sublevels that give them characteristic properties the formation of variable oxidation states in transition elements can be explained by the fact that their successive ionization energies are close in value 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	explain how discontinuities in first ionization energy across a period provide evidence for the existence of sublevels recognize properties of transition elements, including: variable oxidation state, high melting points, magnetic properties, catalytic properties, formation of coloured compounds and formation of complex ions with ligands deduce the electron configurations of ions of the first-row transition elements 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

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Guiding Question

How does the classification of organic molecules help you to predict their properties?

Learning Outcomes
After studying this topic you should be able to:

Understand	Apply your knowledge to
organic compounds can be represented by different types of formulas. These include empirical, molecular, structural (full and condensed), stereochemical and skeletal. 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. 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. successive members of a homologous series show a trend in physical properties “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. structural isomers are molecules which have the same molecular formula but different connectivities.	identify different formulas, and interconvert between molecular, skeletal and structural formulas. construct 3-D models (real or virtual) of organic molecules. identify the following functional groups by name and structure: halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester, phenyl. recognize 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. Recognize isomers, including branched, straight-chain, position and functional group isomers. Apply IUPAC nomenclature to saturated or mono-unsaturated compounds having up to six carbon atoms in the parent chain and containing one type of the following functional groups: halogeno, hydroxyl, carbonyl, carboxyl (to include straight-chain and branched-chain isomers).
AHL	AHL
stereoisomers have the same constitution (atom identities, connectivities and bond multiplicities) but different spatial arrangements of atoms. mass spectrometry of organic compounds can cause fragmentation of molecules. infrared (IR) spectra can be used to identify the type of bond present in a molecule. proton nuclear magnetic resonance spectroscopy (1H NMR) gives information on the different chemical environments of hydrogen atoms in a molecule. individual NMR signals can be split into clusters of peaks. data from different techniques are often combined in structural analysis.	describe and explain the features which give rise to cis–trans isomerism; recognition in non-cyclic alkenes and C3 and C4 cycloalkanes. drawing stereochemical formulas showing the tetrahedral arrangement around a chiral carbon. describe and explain chiral carbon atom giving rise to stereoisomers with different optical properties. recognize of a pair of enantiomers as non-superimposable mirror images from 3-D modelling (real or virtual). deduce information about the structural features of a compound from specific MS fragmentation patterns. interpret the functional group region of an IR spectrum, using a table of characteristic frequencies (wavenumber / cm⁻¹). 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). Interpret 1H NMR spectra from splitting patterns showing singlets, doublets, triplets and quartets to deduce greater structural detail. interpret a variety of data including analytical spectra to determine the structure of a molecule.

Reactivity

Reactivity refers to how and why chemical reactions occur

Reactivity 1. What drives chemical reactions?

Reactivity 1.1—Measuring enthalpy changes.

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Guiding Question

What can be deduced from the temperature change that accompanies chemical or physical change?

Learning outcomes.
After studying this topic students should be able to:

Understand	Apply their knowledge to
Chemical reactions involve a transfer of energy between the system and the surroundings, while total energy is conserved. Reactions are described as endothermic or exothermic, depending on the direction of energy transfer between the system and the surroundings. Temperature changes accompany endothermic and exothermic reactions The relative stability of reactants and products determines whether reactions are endothermic or exothermic. 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.	Distinguish between heat and temperature. Understand the temperature change (decrease or increase) that accompanies endothermic and exothermic reactions, respectively. Sketch and interpret potential energy profiles for endothermic and exothermic reactions. apply the equations Q = mcΔT and ΔH^⦵ = − Q/n in the calculation of the enthalpy change of a reaction.
IB Clarification Notes
Axes for energy profiles should be labelled as reaction coordinate (x), potential energy (y). The units of ΔH^⦵ are kJ mol^-1. The equation Q = mcΔT and the value of c, the specific heat capacity of water, are given in the data booklet There is no higher-level only material in R1.1.

Reactivity 1.2—Energy cycles in reactions

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Guiding Question

How does application of the law of conservation of energy help us to predict energy changes during reactions?

Learning outcomes.
After studying this topic students should be able to:

Understand	Apply their knowledge to
Bond-breaking absorbs and bond-forming releases energy. Hess’s law states that the enthalpy change for a reaction is independent of the pathway between the initial and final states.	Calculate the enthalpy change of a reaction from given average bond enthalpy data. Apply hess’s law to calculate enthalpy changes in multi-step reactions.
AHL	AHL
Standard enthalpy changes of combustion, ΔH_c^⦵ and formation, ΔH_f^⦵, data are used in thermodynamic calculations. An application of Hess’s Law uses enthalpy of formation and/or enthalpy of combustion to calculate the enthalpy change of a reaction. A Born-Haber cycle is an application of Hess’s law, used to show energy changes in the formation of an ionic compound.	Deduce equations and solutions to problems involving these terms. Calculate enthalpy changes of a reaction using ΔH_f^⦵ data or ΔH_c^⦵ data: ΔH^⦵ = Σ ΔH_f^⦵_products − Σ ΔH_f^⦵_reactants ΔH^⦵ = Σ ΔH_c^⦵_reactants − Σ ΔH_c^⦵_products Interpret and determine values from a Born–Haber cycle for compounds composed of univalent and divalent ions.
IB Clarification Notes
Enthalpy of combustion and formation data are given in the data booklet. The equations to determine the enthalpy change of a reaction using data are given in the data booklet ΔH^⦵ = Σ ΔH_f^⦵_products − Σ ΔH_f ^⦵_reactants ΔH^⦵ = Σ ΔHc^⦵_reactants − Σ ΔH_c^⦵_products The Born-Haber cycle includes ionization energies, enthalpy of atomization (using sublimation and/or bond enthalpies), electron affinities, lattice enthalpy, enthalpy of formation. The construction of a complete Born-Haber cycle will not be assessed.

Reactivity 1.3—Energy from fuels

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Guiding question

What are the challenges of using chemical energy to address our energy needs?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
Reactive metals, non-metals and organic compounds undergo combustion reactions when heated in oxygen. Incomplete combustion of organic compounds, especially hydrocarbons, leads to the production of carbon monoxide and carbon. Fossil fuels, which include coal, crude oil and natural gas, have different advantages and disadvantages. Understand the link between carbon dioxide levels and the greenhouse effect. 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. A fuel cell can be used to convert chemical energy from a fuel directly to electrical energy.	Deduce equations for reactions of combustion, including hydrocarbons and alcohols. Deduce equations for the incomplete combustion of hydrocarbons and alcohols. Evaluate the amount of carbon dioxide added to the atmosphere when different fuels burn. Consider the advantages and disadvantages of biofuels. Deduce the half-equations for the electrode reactions in a fuel cell.
IB Clarification Notes
The tendency for incomplete combustion and energy released per unit mass should be covered. The reactants and products of photosynthesis should be known. Hydrogen and methanol should be considered as fuels for fuel cells. The use of proton exchange membranes will not be assessed. There is no higher-level only material in R1.3.

Reactivity 1.4—Entropy and spontaneity (AHL)

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Guiding Question

What determines the direction of chemical change?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
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, entropy of gas > liquid > solid. 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. At constant pressure, a change is spontaneous if the change in Gibbs energy, ΔG, is negative. As a reaction approaches equilibrium, ΔG becomes less negative and finally reaches zero.	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^⦵. 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. Perform calculations using the equation ΔG = ΔG^⦵ + RT ln Q and its application to a system at equilibrium ΔG^⦵ = − RT ln K_c.
IB Clarification Notes
Standard entropy values are given in the data booklet. Thermodynamic data values are given in the data booklet. Note the units: ΔH kJ mol⁻¹; ΔS J K^-1 mol⁻¹; ΔG kJ mol⁻¹ ΔG considers the direct entropy change resulting from the transformation of the chemicals and the indirect entropy change of the surroundings as a result of the transfer of heat energy. The equations ΔG = ΔG^⦵ + RT ln Q and ΔG^⦵ = − RT ln K_c are given in the data booklet.

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

Reactivity 2.1—How much? The amount of chemical change

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Guiding Question

How are chemical equations used to calculate reacting ratios?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
Chemical equations show the ratio of reactants and products in a reaction. 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. The limiting reactant determines the theoretical yield. The percentage yield is calculated from the ratio of experimental yield to theoretical yield. The atom economy is a measure of efficiency in Green Chemistry.	Deduce chemical equations when reactants and products are specified. Calculate reacting masses and/or volumes and concentrations of reactants and products. 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. Calculate the atom economy from the stoichiometry of a reaction.
IB Clarification Notes
Include the use of state symbols in chemical equations. Avogadro’s law and definitions of molar concentration are covered in Structure 1.4. The values for A_r given in the data booklet to two decimal places should be used in calculations. Include discussion of the inverse relationship between atom economy and wastage in industrial processes. The equation for calculation of the atom economy is given in the data booklet. There is no higher-level only material in R2.1

Reactivity 2.2—How fast? The rate of chemical change

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Guiding question

How can the rate of a reaction be controlled?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
the rate of reaction is expressed as the change in concentration of a particular reactant/product per unit time. species react as a result of collisions of sufficient energy and proper orientation. factors that influence the rate of a reaction include pressure/concentration, surface area, temperature and the presence of a catalyst. activation energy, E_a, is the minimum energy that colliding particles need for a successful collision leading to a reaction. catalysts increase the rate of reaction by providing an alternative reaction pathway with lower E_a.	determine rates of reaction explain the relationship between the kinetic energy of the particles and the temperature in kelvin, and the role of collision geometry. predict and explain the effects of changing conditions on the rate of a reaction. construct Maxwell-Boltzmann energy distribution curves to explain the effect of temperature on the probability of successful collisions. 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 E_a on the probability of successful collisions affecting these, including the effect of a catalyst.
AHL	AHL
many reactions occur in a series of elementary steps, and the slowest step determines the rate of the reaction. energy profiles can be used to show the activation energy and transition state of the rate determining step in a multi-step reaction. the molecularity of an elementary step is the number of reacting particles taking part in that step. rate equations depend on the mechanism of the reaction and can only be determined experimentally. 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. the rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction. the Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy. the Arrhenius factor, A, takes into account the frequency of collisions with proper orientations.	evaluate proposed reaction mechanisms and recognise reaction intermediates. distinguish between intermediates and transition states, and recognise both in energy profiles of reactions. construct and interpret energy profiles from kinetic data. interpret the terms unimolecular, bimolecular and termolecular. deduce the rate equation for a reaction from experimental data. sketch, identify and analyse graphical representations of zero, first and second order reactions. solve problems involving the rate equation, including the units of k. analyse graphical representations of the Arrhenius equation, including its linear form. determine the activation energy and the Arrhenius factor from experimental data.
IB Clarification notes
Calculation of reaction rates from tangents of graphs of concentration, volume or mass against time should be covered. Biological catalysts are called enzymes. The different mechanisms of homogeneous and heterogeneous catalysts will not be assessed. AHL include examples where the rate-determining step is not the first step. Proposed reaction mechanisms must be consistent with kinetic and stoichiometric data. Concentration-time and rate-concentration graphs should be included. Only integer values for order of reaction will be assessed.

Reactivity 2.3—How far? The extent of chemical reaction

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Guiding question

How can the extent of a reversible reaction be influenced?

Learning outcomes
After studying this topic students should be able to:

Understand	Apply their knowledge to
a state of dynamic equilibrium is reached in a closed system when the rates of forward and backward reactions are equal the equilibrium law describes how the equilibrium constant, K, can be determined from the stoichiometry of a reaction the magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent Le Châtelier’s principle enables the prediction of the qualitative effects of changes in concentration, temperature and pressure on a system at equilibrium	describe the characteristics of a physical and chemical system at equilibrium. deduce the equilibrium constant expression from an equation for a homogeneous reaction. determine the relationships between K values for reactions that are the reverse of each other at the same temperature. apply Le Châtelier’s principle to predict and explain responses to changes of systems at equilibrium.
AHL	AHL
the reaction quotient, Q, is calculated using the equilibrium expression with non-equilibrium concentrations of reactants and products. the equilibrium law is the basis for quantifying the composition of an equilibrium mixture. the equilibrium constant and Gibbs energy change, ΔG, can both be used to measure the position of an equilibrium reaction	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.
IB Clarification Notes
Include the extent of reaction for: K<<1, K<1, K=1, K>1, K>>1 Include the effects of K on the equilibrium position. LeChatlier’s principle can be applied to heterogenous equilibria such as: X(g)⇌X(aq)

Reactivity 3. What are the mechanisms of chemical change?

Reactivity 3.1—Proton transfer reactions

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Guiding question

What happens when protons are transferred?

Learning outcomes
After studying this topic you should be able to:

Understand	Apply your knowledge to
a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor. a pair of species differing by a single proton is called a conjugate acid-base pair. some species can act as both Brønsted-Lowry acids and bases. the pH scale is a convenient means to describe the [H+] of a solution; pH = − log₁₀[H⁺]; [H⁺] = 10^−pH. the ionic product constant of water, K_w, shows an inverse relationship between [H⁺] and [OH⁻]; K_w = [H⁺][OH⁻]. strong and weak acids and bases differ in the extent of ionization. acids react with bases in neutralization reactions. pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features.	deduce the Brønsted-Lowry acid and base in a reaction. deduce the formula of the conjugate acid or base of any Brønsted-Lowry base or acid. interpret and formulate equations to show acid-base reactions of these species. perform calculations involving the logarithmic relationship between pH and [H⁺]. recognise solutions as acidic, neutral and basic by the relative values of [H⁺] and [OH⁻] recognise that acid-base equilibria lie in the direction of the weaker conjugate formulate equations for the reaction between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates. sketch and interpret the general shape of the pH curve for the neutralisation reaction between a strong acid and a strong base
AHL	AHL
the pOH scale describes the [OH⁻] of a solution. pOH = − log₁₀[OH⁻]; [OH⁻] = 10^−pOH the strengths of weak acids and bases are described by their K_a, K_b, pK_aor pK_bvalues. for a conjugate acid-base pair, the relationship K_ax K_b= Kw can be derived from the expressions for K_aand K_b. the pH of a salt solution depends on the relative strengths of the parent acid and base. pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features. 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 pK_a value. an appropriate indicator for an acid-base titration has an end point range that coincides with the pH at the equivalence point. a buffer solution is one which resists change in pH on the addition of small amounts of acid or alkali. the pH of a buffer solution depends on both: → the pK_a or pK_b of its acid or base. → the ratio of the concentration of acid or base to the concentration of the conjugate base or acid.	interconvert between [H⁺], [OH⁻], pH and pOH values. interpret the relative strengths of acids and bases from K_a, K_b, pK_aand pK_b data. solve problems involving K_a, K_b, K_w, pK_a and pK_b values construct equations for the hydrolysis of ions in a salt, and the predict the effect of each ion on the pH of the salt solution. interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases. construct equilibria expressions to show why the colour of an indicator changes with pH. identify an appropriate indicator for a titration from the identity of the salt and the pH range of the indicator. describe the composition of acidic and basic buffers and explain their action. solve problems involving the composition and pH of a buffer solution, using the equilibrium constant.

Reactivity 3.2— Electron transfer reactions

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Required Course Content
LEARNING OBJECTIVE
6.2.A Represent a chemical or physical transformation with an energy diagram.
ESSENTIAL KNOWLEDGE
6.2.A.1 A physical or chemical process can be described with an energy diagram that shows the endothermic or exothermic nature of that process.

Reactivity 3.3— Electron sharing reactions

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Required Course Content
LEARNING OBJECTIVE
6.3.A Explain the relationship between the transfer of thermal energy and molecular collisions.
ESSENTIAL KNOWLEDGE
6.3.A.1 The particles in a warmer body have a greater average kinetic energy than those in a cooler body.
6.3.A.2 Collisions between particles in thermal contact can result in the transfer of energy. This process is called “heat transfer,” “heat exchange,” or “transfer of energy as heat.”
6.3.A.3 Eventually, thermal equilibrium is reached as the particles continue to collide. At thermal equilibrium, the average kinetic energy of both bodies is the same, and hence, their temperatures are the same.

Reactivity 3.4— Electron-Pair Sharing Reactions

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Required Course Content
LEARNING OBJECTIVE
6.4.A Calculate the heat q absorbed or released by a system undergoing heating/cooling based on the amount of the substance, the heat capacity, and the change in temperature.
ESSENTIAL KNOWLEDGE
6.4.A.1 The heating of a cool body by a warmer body is an important form of energy transfer between two systems. The amount of heat transferred between two bodies may be quantified by the heat transfer equation:
EQN: q = mcΔT.
Calorimetry experiments are used to measure the transfer of heat.
6.4.A.2 The first law of thermodynamics states that energy is conserved in chemical and physical processes.
6.4.A.3 The transfer of a given amount of thermal energy will not produce the same temperature change in equal masses of matter with differing specific heat capacities.
6.4.A.4 Heating a system increases the energy of the system, while cooling a system decreases the energy of the system.
6.4.A.5 The specific heat capacity of a substance and the molar heat capacity are both used in energy calculations.
6.4.A.6 Chemical systems change their energy through three main processes: heating/cooling, phase transitions, and chemical reactions.
6.4.A.7 In calorimetry experiments involving dissolution, temperature changes of the mixture within the calorimeter can be used to determine the direction of energy flow. If the temperature of the mixture increases, thermal energy is released by the dissolution process (exothermic). If the temperature of the mixture decreases, thermal energy is absorbed by the dissolution process (endothermic).

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