IB Chemistry Curriculum Specifications

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• Everything is built from elements, which cannot be split into simpler substances by chemical means, and compounds, in which atoms of different elements are chemically bonded in a fixed ratio.
• A mixture is different: its parts are not bonded, can be present in any proportion and can be separated physically, for example by filtration, evaporation, distillation, recrystallisation or paper chromatography; dissolving also involves solvation, where solvent particles surround the dissolved particles.
• Mixtures may be homogeneous (uniform throughout) or heterogeneous.
• The kinetic molecular theory pictures solids, liquids and gases as moving particles and explains their properties and the changes of state: melting, freezing, vaporisation (evaporation and boiling), condensation, sublimation and deposition.
• Temperature measures the average kinetic energy of the particles.
• The kelvin has the same step size as the Celsius degree, so converting means adding or subtracting 273.
• Use the state symbols (s), (l), (g) and (aq) in equations, and be able to link observations such as temperature plateaus to changes of state.

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• An atom has a tiny, dense, positively charged nucleus made of protons and neutrons (together called nucleons; ordinary hydrogen has no neutron).
• Negatively charged electrons occupy the space around it.
• Know the relative masses and charges of the three particles; the electron's mass is negligible.
• Using the nuclear symbol, with mass number above and atomic number below, you can deduce the numbers of protons, neutrons and electrons in atoms and ions.
• Isotopes are atoms of the same element with different numbers of neutrons.
• They behave the same chemically but differ in physical properties such as density.
• Relative atomic mass is a weighted average over the isotopes, so you should be able to calculate it from abundances (and work backwards from it).
• You do not need to memorise specific isotopes.

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• When an excited electron drops to a lower energy level it emits a photon, and the set of photon energies forms an emission spectrum.
• The line spectrum of hydrogen, with lines that crowd together towards higher energy, is the evidence that electrons occupy discrete energy levels.
• Be able to relate colour, wavelength, frequency and energy across the electromagnetic spectrum (E = hf and c = λf are given), tell a continuous spectrum from a line spectrum, and describe the hydrogen transitions to levels 1, 2 and 3 (the names of the series are not needed).
• Each main level, numbered n, holds up to 2n² electrons.
• Levels divide into s, p, d and f sub-levels of rising energy, made of orbitals that each hold two electrons of opposite spin.
• Recognise the shape of an s orbital and the three p orbitals.
• Use the Aufbau principle, Hund's rule and the Pauli exclusion principle to write configurations for atoms and ions up to Z = 36, in full and condensed (noble-gas core) form and as orbital (arrow-in-box) diagrams, and know the exceptions chromium and copper.

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• The mole is the SI unit of amount of substance: one mole contains the Avogadro constant of elementary entities, which may be atoms, molecules, ions, electrons or any specified particle.
• Relative atomic mass (Ar) and relative formula mass (Mr) compare masses with carbon-12 and have no units; molar mass (g mol⁻¹) is the mass of one mole, and n = m/M.
• Convert freely between mass, moles and number of particles, using Ar values to two decimal places.
• The empirical formula is the simplest ratio of atoms; the molecular formula gives the actual numbers.
• Find an empirical formula from percentage composition or from combustion mass data, and a molecular formula from the empirical formula and the molar mass.
• For solutions, concentration c = n/V in mol dm⁻³ (written with square brackets), and you should convert to and from g dm⁻³.
• Use mole ratios from equations to find product masses and gas volumes, applying Avogadro's law: equal volumes of ideal gases at the same temperature and pressure contain equal numbers of molecules.

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• The ideal gas model treats gas particles as moving points with negligible volume, no intermolecular forces and perfectly elastic collisions.
• Real gases depart from it, most at low temperature and high pressure, where particles are close together and attractions matter (no calculations are needed for this).
• At a given temperature and pressure, the molar volume of an ideal gas is a constant.
• Use PV = nRT and the combined gas law with SI units only (R is given), including to find the molar mass of a gas from experimental data.
• Be able to investigate and graph how pressure, volume and temperature relate for a fixed mass of gas.
• The names of the individual gas laws are not assessed.

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• Metal atoms lose electrons to form cations; non-metal atoms gain electrons to form anions (a cation can also form when a molecule such as ammonia gains a proton).
• Predict an ion's charge from its electron configuration, and remember that transition metals can form ions of more than one charge.
• Ionic bonding is the electrostatic attraction between oppositely charged ions, which pack into a three-dimensional lattice and are written as empirical formulas.
• Name binary ionic compounds with the cation first and the anion ending in -ide, and be fluent with names and formulas of ammonium, hydroxide, nitrate, hydrogencarbonate, carbonate, sulfate and phosphate.
• Lattice enthalpy measures bond strength and is larger for higher charges and smaller ions.
• Explain the properties of ionic compounds from structure: low volatility and high melting points because the lattice is strong, conduction only when molten or dissolved because ions must be free to move, and solubility in water through ion–water attractions.

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• A covalent bond is the attraction between a shared pair of electrons and the nuclei either side of it.
• Atoms tend towards eight valence electrons (the octet rule), though some species have fewer.
• Single, double and triple bonds share one, two and three pairs, and more shared pairs mean shorter, stronger bonds.
• In a coordination bond both electrons come from one atom.
• Draw Lewis formulas for molecules and ions with up to four electron pairs on each atom.
• Bond polarity comes from the electronegativity difference between the atoms (Pauling values are given), and can be shown with partial charges, dipoles or vectors.
• Lewis formulas may use dots, crosses or dashes for electron pairs, and you should meet both organic and inorganic examples and species with fewer than eight electrons.
• VSEPR theory says electron domains around a central atom repel one another.
• Use it to predict electron-domain geometry, molecular geometry and bond angles for two, three and four domains, allowing for lone pairs and multiple bonds squeezing angles.
• A molecule is polar only if its bond dipoles do not cancel, so shape matters.
• Carbon and silicon form giant covalent structures: describe silicon, silicon dioxide, diamond, graphite, fullerenes and graphene, and why allotropes differ in properties.
• Intermolecular forces depend on molecular size and polarity.
• “Van der Waals' forces” is the umbrella term for dipole–dipole, dipole–induced dipole and London forces.
• London (dispersion) forces, which come from instantaneous induced dipoles, act between all particles and are the main force between non-polar ones, dipole–dipole forces between polar molecules, and hydrogen bonds when hydrogen bonded to a highly electronegative atom is attracted to a lone pair on a neighbour.
• For similar molar masses the strength order is London < dipole–dipole < hydrogen bonding.
• Use this to explain volatility, conductivity and solubility.
• Chromatography (paper or TLC) separates a mixture by how strongly each component is attracted to the stationary and mobile phases; calculate and interpret Rf values.

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• A metallic bond is the attraction between a lattice of positive ions and a sea of delocalised electrons.
• Its strength depends on the charge and radius of the ions, which is all the detail you need (charge and electron density).
• Use this to explain electrical and thermal conductivity and malleability, to relate properties of metals to their uses, and to explain the trends in melting points of s-block and p-block metals.

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• Real bonding is a continuum between the ionic, covalent and metallic models, shown on a bonding triangle.
• For a binary compound, use electronegativity data to place it on the triangle and predict its properties from its position (percentage ionic character is not calculated).
• Alloys are mixtures of a metal with other metals or non-metals; because metallic bonding is non-directional, layers can slide without breaking the bonding, which gives alloys properties different from the pure metal (bronze, brass and stainless steel are common examples).
• Polymers are macromolecules built from repeating monomer units, and the common properties of plastics follow from their structure.
• Addition polymers form when the C=C double bond in each alkene monomer opens up; draw the repeating unit from a monomer's structure.
• Be ready to discuss natural and synthetic polymers.

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• The table is organised into periods (rows), groups (columns) and blocks.
• The period number shows the outer occupied energy level, and elements in a group share the same number of valence electrons.
• Identify metals, metalloids and non-metals, and deduce an electron configuration up to Z = 36 from an element's position and the reverse.
• Periodicity means repeating trends.
• Explain the trends in atomic radius, ionic radius, ionisation energy, electron affinity and electronegativity across periods and down groups.
• Metallic character rises down group 1 and non-metallic character falls down group 17.
• Describe and explain the reactions of group 1 metals with water and of group 17 elements with halide ions.
• Oxides move from basic (metals) through amphoteric to acidic (non-metals); write equations for the oxides of groups 1 and 2, carbon and sulfur with water.
• The oxidation state is the charge an atom would have if the compound were fully ionic.
• Deduce it for atoms in compounds and ions, and explain why it is zero for an element.

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• Organic compounds can be shown as empirical, molecular, structural (full and condensed), stereochemical and skeletal formulas; be able to convert between them and build 3D models.
• Functional groups give a compound its characteristic physical and chemical properties, so compounds are sorted into classes by them.
• Know the halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester and phenyl groups, and recognise the series: alkanes, alkenes, alkynes, halogenoalkanes, alcohols, aldehydes, ketones, carboxylic acids, ethers, amines, amides and esters.
• A homologous series has a general formula and successive members differ by a common unit, usually CH₂, giving trends in melting and boiling points.
• Structural isomers share a molecular formula but differ in connectivity (chain, position and functional-group isomers).
• Apply IUPAC nomenclature to saturated or mono-unsaturated compounds with up to six carbons in the parent chain and one type of halogeno, hydroxyl, carbonyl or carboxyl group, including branched isomers.

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• Reactions transfer energy between the system and its surroundings, and total energy is conserved.
• Exothermic reactions release energy and the surroundings warm up; endothermic reactions absorb it and the surroundings cool.
• Which one occurs depends on the relative stability of reactants and products.
• Distinguish heat (energy transferred) from temperature (average kinetic energy).
• Sketch and interpret energy profiles, with reaction coordinate on the x-axis and potential energy on the y-axis.
• The standard enthalpy change ΔH⦵ is the heat transferred at constant pressure under standard conditions, in kJ mol⁻¹.
• Find it from the temperature change of a pure substance using Q = mcΔT and ΔH⦵ = −Q/n.

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• Breaking bonds absorbs energy and forming bonds releases it, so you can estimate a reaction's enthalpy change from average bond enthalpies.
• Hess's law says the enthalpy change depends only on the start and end states, not on the route; use it to combine multi-step reactions and calculate unknown values.

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• Reactive metals, non-metals and organic compounds burn in oxygen.
• Write equations for complete combustion of hydrocarbons and alcohols, and for incomplete combustion, which produces carbon monoxide and carbon; know which fuels are most prone to it.
• Compare fuels by energy released per unit mass and by the carbon dioxide they add to the atmosphere, and link CO₂ levels to the greenhouse effect.
• Fossil fuels (coal, crude oil, natural gas) each have advantages and disadvantages.
• Biofuels come from carbon recently fixed by photosynthesis (know its reactants and products), and you should weigh their pros and cons and distinguish renewable from non-renewable sources.
• A fuel cell converts a fuel's chemical energy directly into electricity; write half-equations for hydrogen and methanol fuel cells (proton-exchange membranes are not assessed).

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• A balanced equation shows the mole ratio of reactants and products.
• Use it to find masses, gas volumes and solution concentrations, and write equations with state symbols when reactants and products are given.
• The limiting reactant runs out first and sets the theoretical yield; identify the limiting and excess reactants from data.
• Percentage yield compares the experimental yield with the theoretical yield.
• Atom economy measures how much of the reactants ends up in the desired product, a key idea in green chemistry (the formula is given).
• High atom economy means less waste, so discuss the inverse relationship between atom economy and wastage in industry.

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• Rate is the change in concentration of a reactant or product per unit time; calculate it from tangents to graphs of concentration, volume or mass against time.
• Reactions happen when particles collide with enough energy (at least the activation energy, Ea) and the right orientation.
• Rate rises with concentration or pressure, surface area, temperature and a catalyst.
• Temperature raises the particles' kinetic energy (in kelvin), so a larger fraction of collisions are successful; use Maxwell–Boltzmann curves to show this and the effect of a different Ea.
• A catalyst offers an alternative pathway with lower Ea; sketch energy profiles with and without a catalyst for endothermic and exothermic reactions.
• Enzymes are biological catalysts.
• The different mechanisms of homogeneous and heterogeneous catalysts are not assessed.

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• In a closed system, dynamic equilibrium is reached when forward and backward rates are equal and concentrations stay constant.
• Write the equilibrium constant expression K for a homogeneous reaction from its equation.
• The size of K shows how far the reaction goes (from K much less than 1 to much greater than 1) and depends on temperature; the K for the reverse reaction is the reciprocal.
• Le Châtelier's principle predicts qualitatively how a system responds to changes in concentration, temperature and pressure, including heterogeneous cases such as X(g) ⇌ X(aq).

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• A Brønsted–Lowry acid donates a proton and a base accepts one; two species differing by one proton are a conjugate pair, and some species can act as both. pH = −log[H⁺] and Kw = [H⁺][OH⁻], so you can classify solutions as acidic, neutral or basic and calculate with the logarithmic scale.
• Strong and weak acids and bases differ in how far they ionise, and equilibria lie towards the weaker conjugate.
• Write equations for acids reacting with metal oxides, metal hydroxides, hydrogencarbonates and carbonates, and sketch and interpret the pH curve for a strong acid with a strong base.

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• Redox can be described as electron transfer, change in oxidation state, or oxygen and hydrogen gain or loss.
• Deduce oxidation states, identify what is oxidised and reduced and the oxidising and reducing agents, and write half-equations and full equations in acidic or neutral solution.
• Predict the relative ease of oxidation of metals and reduction of halogens from their positions in the table, interpret metal/metal-ion displacement data, and write equations for reactive metals with dilute HCl and H₂SO₄.
• In electrochemical cells oxidation occurs at the anode and reduction at the cathode.
• Voltaic (primary) cells turn spontaneous redox energy into electricity; label electrodes and polarities, explain electron flow in the external circuit and ion movement across the salt bridge.
• Rechargeable (secondary) cells reverse their reactions on charging.
• Electrolytic cells use electrical energy to force non-spontaneous reactions, with current carried by ions; deduce the products of electrolysing a molten salt.
• In organic chemistry, primary alcohols oxidise in two steps via an aldehyde and secondary alcohols oxidise to ketones; carboxylic acids reduce to primary alcohols via the aldehyde and ketones to secondary alcohols; adding hydrogen to alkenes and alkynes lowers unsaturation.

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• A radical is a species with an unpaired electron, written with a dot (for example •CH₃ and •Cl), and radicals are highly reactive.
• Homolytic fission of a halogen by UV light or heat produces them, which is the initiation step of a chain reaction.
• Radicals substitute into alkanes with halogens, giving a mixture of products; write equations for the initiation, propagation and termination steps.

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• A nucleophile forms a bond to its partner (the electrophile) by donating both bonding electrons.
• In nucleophilic substitution the nucleophile donates a pair as another bond breaks and a leaving group departs; deduce the equations and describe the movement of electron pairs.
• Heterolytic fission breaks a covalent bond so both electrons stay with one fragment, forming ions.
• An electrophile accepts a pair of electrons; alkenes are electron-rich at the C=C bond and are attacked by electrophiles, giving electrophilic addition.
• Write equations for alkenes reacting with water, halogens and hydrogen halides.