Guiding question

How can we model the particulate nature of matter?

Understand:

Apply their knowledge to:

• 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.

• The kinetic molecular theory is a model to explain physical properties of matter (solids, liquids, and gases) and changes of state.

• Temperature (in K) is a measure of average kinetic energy of particles.

• Distinguish between the properties of elements, compounds, and mixtures.

• Distinguish between the different states of matter.

• Use of state symbols (s, l, g and aq) in chemical equations.

• Interpret observable changes in physical properties and temperature during changes of state.

• Convert between values in the Celsius and kelvin scales.

IB Clarification Notes

• • Solvation, filtration, recrystallization, evaporation, distillation, and paper chromatography should be covered.

  • The differences between homogeneous and heterogeneous mixtures should be understood.
    

  • Names of the changes of state should be covered: melting, freezing, vaporization (evaporation and boiling), condensation, sublimation, and deposition.
    ​

  • The kelvin (K) is the SI unit of temperature and has the same incremental value as the Celsius degree (°C)

• • Note: There is no higher-level only content in S1.1.

Guiding question

How do the nuclei of atoms differ?

Understand:

Apply their knowledge to:

• Atoms contain a positively charged, dense nucleus which (except for hydrogen) is composed of protons and neutrons (nucleons).

• Negatively charged electrons occupy the space outside the nucleus.

• subatomic particles have different masses and charges

• The concept of relative atomic mass.

• Isotopes are atoms of the same element with different numbers of neutrons.

• Deduce the number of protons, neutrons and electrons in atoms and ions by using the nuclear symbol ^A_Z X.

• Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data.

AHL

AHL

• 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.

IB Clarification Notes

• ​Relative masses and charges of the subatomic particles should be known, actual values are given in the data booklet. The mass of the electron can be considered negligible.

• Differences in the physical properties of isotopes should be understood.

• Specific examples of isotopes do not need to be learned.

• The operational details of the mass spectrometer will not be assessed.

Guiding question

How can we model the energy states of electrons in atoms?

Understand:

Apply their knowledge to:

• Emission spectra are produced when photons are emitted from atoms as electrons in excited states return to lower energy levels.

• The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.

• The main energy level is given an integer number, n, and can hold a maximum of 2n² electrons.

• 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.

• Each orbital has a defined energy state for a given electron configuration and chemical environment and can hold two electrons of opposite spin.

• Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.

• Describe qualitatively the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.

• Distinguish between a continuous spectrum and a line spectrum.

• 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.

• Deduce the maximum number of electrons that can occupy each energy level.

• Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals.

• Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.

AHL

AHL

• In a hydrogen emission spectrum, the limit of convergence at higher frequency corresponds to ionization.

• Successive ionization energy data for an element give information about its electron configuration.

• explain the trends and discontinuities in first ionization energy across a period and down a group.

• Calculate the value of the first ionization energy from spectral data given the wavelength or frequency of the convergence limit.

• Explain how these discontinuities provide evidence for the existence of energy sub-levels.

• Deduce the group of an element from its successive ionization data.

IB Clarification Notes

• Details of the electromagnetic spectrum are given in the data booklet.

• The names of the different series in the hydrogen emission spectrum will not be assessed.

• Full electron configurations and condensed electron configurations using the noble gas core should be covered. Orbital diagrams, i.e. arrow-in-box diagrams, should be used to represent the filling and relative energy of orbitals.

• The electron configurations of Cr and Cu as exceptions should be covered.

• The value of the Planck constant h and the equations E = h f and c = λ f are given in the data booklet.

Guiding question

How do we quantify matter on the atomic scale?

1. 1. 6. Understand

Apply your knowledge to

• The mole (mol) is the SI unit of amount of substance.

• One mole contains exactly the Avogadro constant elementary entities.

• Masses of atoms are compared on a scale relative to ¹²C and are expressed as relative atomic mass (A_r) and relative formula mass (M_r).

• Molar mass, M, has the units g mol^-1.

• The empirical formula of a compound gives the simplest ratio of the number of atoms of each element present in one molecule of the compound.

• The molecular formula gives the actual number of atoms of each element present in one molecule of the compound.

• the molar concentration is determined by the amount of solute and the volume of solution.

• Convert amount of substance, n, to number of specified elementary entities.

• Calculate the molar masses of atoms, ions, molecules and formula units.

• Solve problems involving the relationships between the number of particles, the amount of substance in mol and the mass in grams.

• Use molar masses with chemical equations to determine the masses of the products of a reaction

• Determine the empirical formula of a compound from its percentage composition by mass and vice-versa.

• Determine empirical formula from experimental data on mass changes in combustion reactions.

• Determine the molecular formula of a compound from its empirical formula and its molar mass.

• Solve problems involving the molar concentration, amount of solute and volume of solution

• Interconvert molar concentration and concentration in g dm⁻³

• Solve problems involving the mole ratio of reactants and/or products and the volume of gases.​

• Avogadro’s law states that equal volumes of all gases measured under the same conditions of

temperature and pressure contain equal numbers of molecules.

• Avogadro’s law applies to ideal gases.

IB Clarification Notes

• An elementary entity may be an atom, a molecule, an ion, an electron, any other particle, or a specified group of particles. The Avogadro constant NA is given in the data booklet. It has the units mol^–1 .

• Relative atomic mass and relative formula mass have no units. The values of relative atomic masses given to two decimal places in the data booklet should be used in calculations.

• The relationship n = m/M is given in the data booklet.

• The use of square brackets to represent molar concentration is required. Units of concentration should include g dm^–3 and mol dm^–3 and conversion between these. The relationship n = CV is given in the data booklet.

• Note: There is no higher-level only content in S1.4

Guiding question

How does the model of ideal gas behaviour help us to predict the behaviour of real gases?

Understand

Apply your knowledge to

• An ideal gas consists of moving particles with negligible volume and no intermolecular forces. All collisions between particles are considered elastic.

• Real gases deviate from the ideal gas model, particularly at low temperature and high pressure.

• The molar volume of an ideal gas is a constant at specified temperature and pressure.

• 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 = P₂T₂/V₂.

• Recognize the key assumptions in the ideal gas model.

• Explain the limitations of the ideal gas model.

• Investigate the relationship between temperature, pressure, and volume for a fixed mass of an ideal gas and analysis of graphs relating these variables.

• Solve problems relating to the ideal gas equation.

• Use the ideal gas equation to calculate the molar mass of a gas from experimental data.

IB Clarification Notes

• No mathematical coverage is required for limitations of the ideal gas model.

• The names of specific gas laws will not be assessed.

• Units of volume and pressure should be SI only. The value of the gas constant R, the ideal gas equation, and the combined gas law, are given in the data booklet

• Note: There is no higher-level only content in S1.5

Guiding question

What determines the ionic nature and properties of a compound?

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.

Guiding question

What determines the covalent nature and properties of a substance?

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.

Guiding question

What determines the metallic nature and properties of an element

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.

Guiding question

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

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.

Guiding question

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

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

Guiding Question

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

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.

Guiding Question

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

Guiding Question

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

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.

Guiding question

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

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.

Guiding Question

What determines the direction of chemical change?

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.

Guiding Question

How are chemical equations used to calculate reacting ratios?

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

Guiding question

How can the rate of a reaction be controlled?

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.

Guiding question

How can the extent of a reversible reaction be influenced?

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)

Guiding question

What happens when protons are transferred?

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_a or pK_b values.

• for a conjugate acid-base pair, the relationship K_a x K_b = Kw can be derived from the expressions for K_a and 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_a and 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.

Guiding Question

What happens when electrons are transferred?

Understand

Apply your knowledge to

• oxidation and reduction can be considered in terms of electron transfer, change in oxidation state, oxygen gain/loss or hydrogen loss/gain.

• half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons.

• 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.

• oxidation occurs at the anode and reduction occurs at the cathode in electrochemical cells.

• a primary (voltaic) cell is an electrochemical cell that converts energy from spontaneous redox reactions to electrical energy.

• secondary (rechargeable) cells involve redox reactions that can be reversed using electrical energy.

• an electrolytic cell is an electrochemical cell that converts electrical energy to chemical energy, by bringing about non-spontaneous reactions.

• functional groups in organic compounds may undergo oxidation.

• functional groups in organic compounds may undergo reduction.

• reduction of unsaturated compounds by the addition of hydrogen lowers the degree of unsaturation

• deduce the oxidation states of an atom in a compound or an ion.

• identify the species oxidized and reduced and the oxidizing and reducing agents in a chemical reaction.

• deduce the redox half-equations and equations in acidic or neutral solutions.

• predict the relative ease of oxidation of metals.

• predict the relative ease of reduction of halogens.

• interpret data regarding metal and metal ion reactions.

• deduce equations for the reactions of reactive metals with dilute HCl and H₂SO₄.

• label 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.

• explain the direction of electron flow from anode to cathode in the external circuit, and ion movement across the salt bridge.

• deduce the reactions of the charging process from given electrode reactions for discharge, and vice versa.

• explain how current is conducted in an electrolytic cell.

• deduce the products of the electrolysis of a molten salt.

• 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.

• deduce equations to show reduction of carboxylic acids to primary alcohols via the aldehyde, and reduction of ketones to secondary alcohols.

• deduce the products of the reactions of hydrogen with alkenes and alkynes.

AHL

AHL

• the hydrogen half-cell, H_2(g) ⇌ 2H⁺_(aq) + 2e⁻ is assigned a standard reduction potential of zero by convention. It is used in the measurement of standard electrode potential, E^⦵.

• standard cell potential, E^⦵_cell, can be calculated from standard reduction potentials. E^⦵cell has a positive value for a spontaneous reaction.

• the equation ΔG^⦵ = − nFE^⦵_cell shows the relationship between standard change in Gibbs energy and standard reduction potential for a reaction.

• during the electrolysis of aqueous solutions, competing reactions can occur at the anode and cathode, including the oxidation and reduction of water.

• electroplating involves the electrolytic coating of an object with a metallic thin layer.

• interpret standard reduction potential data in terms of ease of oxidation/reduction.

• predict whether a reaction is spontaneous in the forward or reverse direction from E^⦵ data.

• determine the value for ΔG^⦵ from E^⦵ data.

• deduce from standard reduction potentials the products from the electrolysis of aqueous solutions.

• deduce equations for the electrode reactions during electroplating.

Guiding question

What happens when a species has an unpaired electron?

Understand

Apply your knowledge to

• a radical is a chemical entity that has an unpaired electron.

• radicals are highly reactive.

• radicals are produced by homolytic fission, e.g. of halogens, in the presence of UV light or heat.

• radicals take part in substitution reactions with alkanes, producing a mixture of products.

• identify and represent radicals e.g. •CH₃ and •Cl.

• explain, including equations, the homolytic fission of halogens, known as the initiation step in a chain reaction.

• explain, using equations, the propagation and termination steps in the reactions between alkanes and halogens.

• There is no higher-level only material in R3.3.

Guiding Question

What happens when reactants share their electron pairs with others?

Understand

Apply your knowledge to

• a nucleophile is a reactant that forms a bond to its reaction partner (the electrophile) by donating both bonding electrons

• in a nucleophilic substitution reaction, a nucleophile donates an electron pair to form a new bond, as another bond breaks producing a leaving group

• heterolytic fission is the breakage of a covalent bond when both bonding electrons remain with one of the two fragments formed

• an electrophile is a reactant that forms a bond to its reaction partner (the nucleophile) by accepting both bonding electrons from that reaction partner

• alkenes are susceptible to electrophilic attack because of the high electron density of the carbon-carbon double bond. These reactions lead to electrophilic addition

• recognize nucleophiles and electrophiles in chemical reactions

• deduce equations for nucleophilic substitution reactions

• describe and explain the movement of electron pairs in nucleophilic substitution reactions

• explain, with equations, the formation of ions by heterolytic fission

• deduce equations for the reactions of alkenes with water, halogens, and hydrogen halides

AHL

AHL

• a Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor

• when a Lewis base reacts with a Lewis acid, a coordination bond is formed; nucleophiles are Lewis bases and electrophiles are Lewis acids

• coordination bonds are formed when ligands donate an electron pair to transition element cations, forming complex ions

• nucleophilic substitution reactions include the reactions between halogenoalkanes and nucleophiles

• the rate of substitution reactions is influenced by the identity of the leaving group

• alkenes readily undergo electrophilic addition reactions

• the relative stability of carbocations in the addition reactions between hydrogen halides and unsymmetrical alkenes can be used to explain the reaction mechanism

• electrophilic substitution reactions include the reactions of benzene with electrophiles

• apply Lewis acid–base theory to inorganic and organic chemistry to identify the role of the reacting species

• draw and interpret Lewis formulas of reactants and products to show coordination bond formation in Lewis acid–base reactions

• deduce the charge on a complex ion, given the formula of the ion and ligands present

• describe and explain the mechanisms of the reactions of primary and tertiary halogenoalkanes with nucleophiles

• predict and explain the relative rates of the substitution reactions for different halogenoalkanes

• describe and explain the mechanisms of the reactions between symmetrical alkenes and halogens, water and hydrogen halides

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

• describe and explain the mechanism of the reaction between benzene and a charged electrophile, E⁺