Structure 1 -- Models of the Particulate Nature of Matter: Revision Notes

Condensed for quick recall of Structure 1 – Models of the particulate nature of matter, the first content component of IB Diploma Programme Chemistry (17 guided hours at SL, 21 at HL). For the full six-component syllabus map, use the IB DP Chemistry syllabus guide.

What Structure 1 covers

Structure 1 builds the particle-level model that every later Structure and Reactivity component assumes: what matter is made of, how atoms are arranged and counted, and how gases behave. It has five official sub-topics.

1.1 Introduction to the particulate nature of matter

• The three classical states (solid, liquid, gas) and the particle-kinetic explanation of state changes (melting, freezing, boiling, condensing, sublimation, deposition).
• Physical vs chemical change: a physical change does not alter chemical composition; a chemical change does, via bond-breaking and bond-forming.
• Pure substances (elements and compounds) vs mixtures, and how to represent them at the particulate level using simple particle diagrams.

1.2 The nuclear atom

• Atomic structure: protons, neutrons and electrons, their relative mass and charge, and their location (nucleus vs electron cloud).
• Atomic number (Z) and mass number (A); isotopes as atoms of the same element with different neutron numbers, and how isotopic abundance is used to calculate relative atomic mass. (The periodic table in the data booklet supplies each element’s relative atomic mass to two decimal places – this does not need to be memorised.)
• Ion formation by electron loss or gain, and how this changes the particle’s overall charge without changing Z.

1.3 Electron configurations

• Emission spectra: exciting an atom’s electrons (for example, by heating) lets them absorb energy and jump to a higher energy level; when an electron then falls back down, it emits a photon whose energy (and so frequency) equals the energy gap between the two levels. A line emission spectrum – discrete coloured lines rather than a continuous band – is the direct evidence that electron energy levels are quantised rather than continuous.
• The hydrogen line emission spectrum specifically: the lines converge (get closer together) at higher frequency, and the convergence limit corresponds to the electron being removed from the atom entirely (ionisation) – this convergence limit is the HL route to calculating ionisation energy directly from spectral data.
• Energy levels, sublevels (s, p, d and f) and orbitals, and the aufbau principle for filling them in order of increasing energy.
• Filling orbitals correctly requires three rules together, shown with orbital diagrams (boxes and arrows): the aufbau principle (fill the lowest-energy sub-level first), the Pauli exclusion principle (an orbital holds a maximum of two electrons, which must have opposite spins), and Hund’s rule (electrons fill separate orbitals within the same sub-level singly, with parallel spins, before any orbital is doubly occupied).
• Writing full electron configurations (e.g. 1s² 2s² 2p⁶) for the first several periods of the periodic table, and condensed/noble-gas-core notation.
• The chromium and copper exceptions: both gain extra stability from a half-filled (Cr, 3d⁵4s¹) or fully-filled (Cu, 3d¹⁰4s¹) 3d sub-shell, so one electron moves from 4s into 3d rather than following the 3d⁴4s² / 3d⁹4s² pattern the aufbau order alone would predict.
• The link between an element’s electron configuration – specifically its valence electrons – and its position in the periodic table (group and period).
• Ionisation energy (SL): the energy needed to remove one mole of electrons from one mole of gaseous atoms (or ions) in their ground state. First ionisation energy generally increases across a period (greater nuclear charge, similar shielding) and decreases down a group (outer electron further from the nucleus, more shielding) – the same reasoning used for reactivity trends elsewhere in the syllabus.

1.4 Counting particles by mass: The mole

• The mole as the SI unit for amount of substance, defined via the Avogadro constant (6.02 × 10²³ mol⁻¹) – supplied in the data booklet, so it does not need to be memorised.
• Molar mass (M) and its use to convert between mass, moles and number of particles.
• Empirical and molecular formulae, and determining them from percentage composition or combustion data.
• Concentration of solutions (mol dm⁻³) and molar volume of a gas at standard conditions.

1.5 Ideal gases

• The ideal gas model’s assumptions (negligible particle volume, no intermolecular forces, elastic collisions) and where real gases deviate from it.
• The combined and ideal gas equations (pV = nRT), and using them to relate pressure, volume, temperature and moles of gas.
• Standard temperature and pressure (STP) conventions used in gas calculations.

Worked example: mass to moles to particles

A typical 1.4 calculation chain: given 4.40 g of CO2 (M = 44.01 g mol-1), find the number of molecules present.

n = m / M = 4.40 / 44.01 = 0.0999... mol
number of molecules = n x Avogadro constant = 0.0999... x 6.02 x 10^23 ~= 6.02 x 10^22 molecules

The same three-step chain (mass -> moles -> particles, or the reverse) underlies most 1.4 exam questions; the only thing that changes is which quantity you’re solving for.

Common exam pitfalls

• Confusing mass number (A) with atomic number (Z) when working out neutron count – neutrons = A minus Z, not A alone.
• Writing electron configurations in filling-order sequence (1s, 2s, 2p, 3s, 3p, 4s, 3d…) but forgetting to re-sort into energy-level order when asked for the “configuration”, which for transition metals is a genuinely separate skill from filling order.
• Treating relative atomic mass as if it were the mass of the most common isotope, rather than the isotopic-abundance-weighted average the syllabus actually defines it as.
• Confusing the molar volume at STP with the molar volume at room temperature and pressure (RTP) used in some other courses – this course and its data booklet give only the STP value, so use that value unless a question states otherwise.
• In ideal gas calculations, forgetting to convert Celsius to kelvin before substituting into pV = nRT – a very common arithmetic slip that produces an answer that looks plausible but is wrong by a fixed offset.

Self-test

1. What three assumptions define an ideal gas?
2. Explain, in terms of sub-atomic particles, why isotopes of the same element have identical chemical properties but different physical properties such as density.
3. What is the difference between an empirical formula and a molecular formula?
4. Why does relative atomic mass often appear as a non-integer value on the periodic table?

Answers: 1. Negligible particle volume compared with the container, no intermolecular forces between particles, and perfectly elastic collisions between particles (and with the container walls). 2. Isotopes have the same number of protons and electrons, so their electron configuration – and therefore their chemical bonding behaviour – is identical; they differ only in neutron number, which changes mass (and so physical properties like density) without changing charge or electron arrangement. 3. An empirical formula gives the simplest whole-number ratio of atoms of each element in a compound; a molecular formula gives the actual number of atoms of each element in one molecule, which may be a whole-number multiple of the empirical formula. 4. Because it is a weighted average of the masses of all naturally occurring isotopes of the element, weighted by their relative abundance – not the mass of a single atom.

How to approach it

Structure 1 is mostly SL-depth content shared between SL and HL students, but it does carry 4 hours of HL-only extension (17 hours SL rising to 21 at HL): one additional hour in 1.2, on interpreting mass spectra, and three additional hours in 1.3, on ionisation energy from the convergence limit of a series of spectral lines and on successive ionisation energies. Do not assume the whole component is SL-only just because most of it is shared. The mole (1.4) is the sub-topic most exam answers actually lean on later in the course – stoichiometric calculations in Reactivity 1 and 2 assume fluent mole arithmetic, so treat 1.4 as a skill to over-learn rather than a topic to revise once and move past. Electron configuration (1.3) is best practised by writing out configurations for unfamiliar elements under timed conditions rather than memorising a fixed list, since Paper 1 multiple-choice questions frequently test the underlying aufbau logic with elements you won’t have seen drilled in class.

Official syllabus

International Baccalaureate Organization, Chemistry guide, Diploma Programme, first assessment 2025 (published February 2023, updated May 2023) – copy consulted via a school-hosted mirror, not ibo.org.