A-level and IB · Sciences

IB Chemistry: structure and reactivity

IB Chemistry flashcards organized by structure and reactivity: atoms, bonding, energetics, kinetics, equilibrium, acids, redox and organic mechanisms.

36 cards5 topicsAbout 15 minutes a roundIB Diploma Programme Chemistry (SL and HL) style content, comparable to Grade 12 chemistry plus first-year university general chemistry

Independent study material written by StudyBrix, not produced or reviewed by any exam board or by the International Baccalaureate Organization. Check your own specification or subject guide.

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About this deck

This deck follows the organization of the IB Diploma Chemistry course first assessed in 2025, which splits content into two strands: structure (models of the particulate nature of matter, bonding and structure, and classification of matter) and reactivity (what drives reactions, how much, how fast and how far, and the mechanisms of chemical change). Tags reflect those strands so you can link a property to the structure that explains it.

The course is assessed by two papers plus an internal scientific investigation worth 20%. If you also look at A-level material, remember that specifications differ between exam boards (AQA, Edexcel, OCR) and from the IB, so check your own guide or specification and the data booklet you are allowed to use.

You should be confident with MYP or Grade 10 chemistry: atomic structure, ionic and covalent bonding, balancing equations and basic mole ideas. Work the numerical examples with significant figures and units, then explain each idea in terms of particles. The IB reactivity strand groups reactions by type of particle transferred (protons, electrons or electron pairs), so ask yourself which type each reaction is.

Common mistakes: confusing intermolecular forces with covalent bonds when explaining boiling points; forgetting to convert kJ and J when using entropy; using the wrong sign convention for enthalpy (exothermic is negative); and writing equilibrium expressions that include solids or pure liquids. For curly arrows, start at an electron pair.

Before you start: MYP or Grade 10 chemistry and algebra: atomic structure, bonding, balancing equations, rearranging formulae and logarithms (for pH).

All 36 cards

Cover the definitions with a sheet of paper and test yourself, or use the study modes above.

Structure: particles and atoms

Kinetic molecular theory: Matter is made of particles in constant motion; their average kinetic energy is proportional to absolute temperature.Example: Doubling the kelvin temperature of a gas doubles the average kinetic energy of its particles.
Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons, so they have different masses but very similar chemistry.Example: Carbon-12 and carbon-14 both have 6 protons; carbon-14 has 8 neutrons.
Emission spectrum: A pattern of coloured lines produced when excited electrons fall to lower energy levels; each line matches one energy difference.Example: Hydrogen's lines converge at higher frequency, showing levels get closer together further from the nucleus.
Molar volume of a gas: One mole of any ideal gas occupies the same volume at a given temperature and pressure: 22.7 dm^3 at 273 K and 100 kPa.Example: 0.25 mol of O2 at those conditions takes up 0.25 * 22.7 = 5.68 dm^3.
Limiting reactant: The reactant that runs out first, so it sets the maximum amount of product.Example: 2H2 + O2 -> 2H2O with 3 mol H2 and 2 mol O2: H2 needs only 1.5 mol O2, so H2 is limiting and 3 mol water forms.
Successive ionization energies: The energies to remove electrons one at a time; a big jump shows an electron is being taken from a new, inner shell.Example: Magnesium shows a large jump between the second and third values, so it has 2 outer electrons.

Structure: bonding and classification

Lewis formula: A diagram showing all valence electrons in a molecule as bonding pairs and lone pairs.Example: CO2: each oxygen forms a double bond to carbon and keeps two lone pairs.
VSEPR theory: Electron domains around a central atom arrange to be as far apart as possible, which predicts molecular shape and bond angle.Example: Four domains with one lone pair (NH3) give a trigonal pyramidal shape near 107 degrees.
Bond polarity versus molecular polarity: A bond is polar when electronegativities differ; a molecule is polar only if the bond dipoles do not cancel because of its shape.Example: CO2 has polar bonds but is non-polar because it is linear; H2O is bent and polar.
London dispersion forces: Weak attractions between temporary induced dipoles; they grow with the number of electrons and the surface area of molecules.Example: Pentane boils higher than its branched isomer 2,2-dimethylpropane because it has more surface contact.
Metallic bonding: Electrostatic attraction between a lattice of positive ions and a sea of delocalized electrons.Example: Delocalized electrons explain why metals conduct electricity when solid.
Giant covalent structure: A network of atoms joined by covalent bonds in all directions, giving very high melting points.Example: Diamond has each carbon bonded to four others; graphite has layers with delocalized electrons.
Homologous series: A family of organic compounds with the same functional group and general formula, differing by CH2 each time, with gradually changing physical properties.Example: Alkanes CnH2n+2: methane, ethane, propane; boiling point rises with chain length.

Reactivity: energy

Standard enthalpy of combustion: The enthalpy change when one mole of a substance burns completely in oxygen under standard conditions.Example: Methane: about -890 kJ/mol.
Calorimetry calculation: Heat transferred q = mc delta T; divide by moles of the limiting reactant and give a negative sign if the temperature rose.Example: 100 g solution rising 3.5 K: q = 100 * 4.18 * 3.5 = 1463 J. For 0.050 mol: delta H = -29.3 kJ/mol.
Hess's law cycle: Because enthalpy is a state function, an unknown enthalpy change can be found by adding known changes along another route.Example: From combustion data: delta H(reaction) = sum of delta H(c) of reactants minus sum of delta H(c) of products.
Average bond enthalpy method: Estimate delta H as the energy to break bonds in reactants minus the energy released forming bonds in products; values are averages, so results are approximate.Example: CH4 + 2O2 -> CO2 + 2H2O: 4(414) + 2(498) - 2(804) - 4(463) = -808 kJ/mol, near the measured -890.
Entropy: A measure of how many ways energy and particles can be arranged; gases have much higher values than liquids or solids.Example: CaCO3(s) -> CaO(s) + CO2(g) has a positive entropy change because a gas forms.
Gibbs energy and spontaneity: delta G = delta H - T delta S; a negative value means the reaction is thermodynamically favourable at that temperature.Example: delta H = -92 kJ/mol and delta S = -0.199 kJ/(K mol) give delta G = -92 + 298 * 0.199 = -32.7 kJ/mol at 298 K.

Reactivity: rates and equilibrium

Collision theory: Reactions happen when particles collide with enough energy (at least the activation energy) and the correct orientation.Example: Crushing a solid increases surface area and so the frequency of successful collisions.
Rate-determining step: The slowest step in a reaction mechanism; it controls the overall rate and appears in the rate equation.Example: If the slow step involves one A and one B, the rate equation is rate = k[A][B].
Reaction quotient Q: The same ratio as the equilibrium constant but using current concentrations; comparing it with K shows which way the reaction will go.Example: If Q < K, the forward reaction is favoured until Q equals K.
Equilibrium constant and temperature: Only temperature changes the value of K. For an exothermic forward reaction, raising temperature lowers K.Example: The Haber process has a smaller K at 700 K than at 500 K.
Percentage yield: Actual mass of product divided by theoretical mass, times 100.Example: Theoretical 8.0 g, actual 6.2 g: yield = 77.5%.
Catalyst: A substance that speeds up a reaction by providing an alternative pathway with lower activation energy, without being used up overall.Example: Iron speeds up ammonia production in the Haber process; it does not change the value of K.
Order of reaction from data: The power of a concentration term in the rate equation, found by changing one concentration at a time and watching the rate.Example: Tripling [A] triples the rate: first order in A. Tripling [B] leaves the rate unchanged: zero order in B.

Reactivity: mechanisms

Proton transfer reaction: A reaction in which a hydrogen ion moves from a Brønsted-Lowry acid to a base, forming a conjugate base and conjugate acid.Example: NH3 + H2O <=> NH4+ + OH-: water donates a proton to ammonia.
pH and pOH: pH = -log10[H+] and pOH = -log10[OH-]; at 298 K they add to 14.Example: [OH-] = 0.0010 mol/dm^3: pOH = 3, so pH = 11.
Strong versus weak acid: A strong acid dissociates almost completely in water; a weak one only partly, so at equal concentration it has a higher pH and conducts less.Example: 0.10 mol/dm^3 HCl has pH 1.0; 0.10 mol/dm^3 ethanoic acid has pH about 2.9.
Electron transfer reaction: A redox process: the species losing electrons is oxidized and acts as the reducing agent; the one gaining electrons is reduced.Example: Zn + Cu2+ -> Zn2+ + Cu: zinc is oxidized, copper ions are reduced.
Electrolysis: Using electrical energy to drive a non-spontaneous redox reaction; oxidation at the anode, reduction at the cathode.Example: Molten NaCl gives sodium at the cathode and chlorine at the anode.
Radical: A species with an unpaired electron, formed by homolytic fission; highly reactive.Example: UV light splits Cl2 into two Cl atoms that start a chain reaction with methane.
Nucleophile: An electron-rich species that donates a pair of electrons to an electron-poor carbon to form a new covalent bond.Example: OH- attacks the carbon in bromoethane, replacing bromine to form ethanol.
Lewis acid: A species that accepts an electron pair to form a coordinate bond.Example: BF3 accepts the lone pair on NH3, forming F3B-NH3.
Electrophilic addition to alkenes: An electron-pair sharing reaction in which the C=C bond attacks an electron-poor species, breaking the pi bond.Example: Ethene decolourizes bromine water as bromine adds across the double bond.
SN1 versus SN2: SN2: one concerted step, rate depends on both reactants, typical of primary halogenoalkanes. SN1: carbocation intermediate, rate depends on the halogenoalkane only, typical of tertiary ones.Example: 2-bromo-2-methylpropane hydrolyses mainly by SN1; bromoethane mainly by SN2.