S1.3
Electron configurations
3 SL · +3 HL (SL 1.3.1-1.3.5) (HL 1.3.6-1.3.7)
Unverified
← View on the mapParent topic: Models of the particulate nature of matter
Guiding question
Guiding questionHow can we model the energy states of electrons in atoms?
The one big idea
Electrons occupy discrete energy levels and orbitals, not orbits. The Aufbau principle, Hund's rule and the Pauli exclusion principle determine configurations up to Z = 36. At HL, ionisation energy data and the convergence limit confirm the model experimentally.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS1.2 The nuclear atomElectron configuration is built on the nuclear atom (protons, neutrons, electrons).
What rests on this
- EditorialS2.1 The ionic model
- EditorialS2.2 The covalent model
- EditorialS3.1 The periodic table: Classification of elements
- EditorialR1.2 Energy cycles in reactions
- EditorialR3.1 Proton transfer reactions
- EditorialR3.2 Electron transfer reactions
Core concepts that must be mastered
- Emission spectra and discrete energy levels(1.3.1, 1.3.2)Electrons do not orbit like planets. They occupy discrete energy levels, and emission spectra are the evidence. When an excited electron returns to a lower level, it emits a photon whose energy equals the difference between the two levels. A continuous spectrum (like a rainbow) has all wavelengths; a line spectrum has only specific wavelengths, proving that only specific energy transitions are allowed. The hydrogen emission spectrum shows several series of lines converging at higher frequency — each series corresponds to transitions to a particular final level (n = 1, 2, 3…). The convergence happens because the energy gaps between adjacent levels shrink as n increases, so transitions from very high levels produce lines that crowd together. The names of the series (Lyman, Balmer, etc.) are not assessed.
- Energy levels, sublevels and orbitals(1.3.3–1.3.5)The main energy level n (n = 1, 2, 3…) holds a maximum of 2n² electrons. Each level divides into sublevels — s, p, d, f — of successively higher energy. The s sublevel has one orbital (spherical), the p sublevel has three (perpendicular dumbbells along x, y, z axes). Each orbital holds two electrons of opposite spin. An orbital is a region of high probability of finding an electron — it has a shape and an energy, not a trajectory. The three principles that determine filling order are the Aufbau principle (fill lowest energy first), Hund's rule (maximise unpaired electrons before pairing) and the Pauli exclusion principle (no two electrons share all four quantum numbers). You must deduce configurations up to Z = 36, in both full and condensed (noble-gas core) forms, with orbital arrow-in-box diagrams. The two exceptions to memorise are Cr ([Ar]3d⁵4s¹) and Cu ([Ar]3d¹⁰4s¹): the half-filled and fully-filled d sublevels are more stable than the naive Aufbau prediction.
- HL: Convergence limit and ionisation energy(1.3.6)At HL, the convergence limit at higher frequency in the hydrogen spectrum corresponds to ionisation — the electron is removed entirely. You can calculate the first ionisation energy from the wavelength or frequency of the convergence limit using E = hf and c = λf (both in the data booklet). You must explain both trends and discontinuities in first ionisation energy across a period and down a group. Across a period, IE generally increases because nuclear charge rises while electrons enter the same main level. The discontinuities are the key evidence for sublevels: the group 2 → 13 dip occurs because the group 13 outer electron is in a higher-energy p orbital (easier to remove); the group 15 → 16 dip occurs because the group 16 electron is the first to pair up in a p orbital, and the repulsion between the paired electrons makes it easier to remove. Down a group, IE decreases because the outer electron is in a higher main level, farther from the nucleus and shielded by more inner electrons.
- HL: Successive ionisation energy(1.3.7)Successive ionisation energy data reveal the group and electron configuration of an element. Each successive IE is larger than the last because the electron is removed from an increasingly positive ion. The large jumps mark the removal of a core electron — when you cross into a new inner shell, the IE jumps sharply because the electron is closer to the nucleus and less shielded. From the position of these jumps you can deduce the group: sodium shows a big jump after the first IE (group 1), magnesium after the second (group 2), and so on. This data is direct experimental evidence for the shell model and for the existence of energy sublevels.
- How this sub-topic connectsElectron configurations explain the periodic table (S3.1): period number equals the outer occupied energy level, group number equals the valence electron count, and the block (s, p, d, f) equals the sublevel being filled. They explain bonding (S2): ions form by losing or gaining electrons to achieve a stable configuration, and covalent bonds form by sharing electrons. At HL, the 4s/3d energy ordering (fill 4s before 3d, remove 4s before 3d) explains transition element chemistry (S3.1, R3.4). Ionisation energy trends connect to metallic character, reactivity of group 1 and group 17, and electrode potentials (R3.2).
Quantitative non-negotiables
- Deduce electron configurations of atoms and ions up to Z = 36, including full and condensed (noble-gas core) forms and orbital arrow-in-box diagrams.
- Recognise the Cr ([Ar]3d⁵4s¹) and Cu ([Ar]3d¹⁰4s¹) exceptions.
- At HL: calculate first ionisation energy from the convergence limit using E = hf and c = λf.HL
- At HL: deduce the group of an element from successive ionisation energy data.HL
Common failure modes
M-01 — Electrons orbit the nucleus in circles like planets.
Confidence: verified
Why it’s wrong: The Bohr model is a superseded approximation; electrons occupy orbitals, which are probability distributions with no defined trajectory.
Correction: An orbital is a region where there is a high probability of finding an electron; it has a shape and an energy, not a path.
M-02 — The 4s subshell is always lower in energy than 3d.
Confidence: verified
Why it’s wrong: It is lower before the 3d starts to fill, which is why it fills first; but once d electrons are present the ordering effectively inverts, which is why 4s electrons are removed first on ionisation.
Correction: Fill 4s before 3d; remove 4s before 3d. Fe²⁺ = [Ar]3d⁶.
- Caveat (uncertain) The preferred wording for the 4s/3d energy ordering: This is a genuinely subtle point and different textbooks phrase the justification differently — teacher should confirm the phrasing IB expects.
M-03 — Chromium is [Ar]3d⁴4s².
Confidence: verified
Why it’s wrong: Cr is [Ar]3d⁵4s¹ and Cu is [Ar]3d¹⁰4s¹.
Correction: Memorise these two exceptions; the half-filled and filled d sublevels are lower in energy than the naive aufbau prediction.
M-04 — Ionisation energy increases smoothly across a period.
Confidence: verified
Why it’s wrong: There are dips at group 2→13 and group 15→16.
Correction: The group 13 dip is because the outer electron is in a higher-energy p orbital; the group 16 dip is because of repulsion between the first pair of paired p electrons.
M-06 — Shells and energy levels and subshells are the same thing.
Confidence: likely
Why it’s wrong: These are three nested levels of structure that students conflate.
Correction: Main energy level (n) → sublevel (s, p, d, f) → orbital (each holding 2 electrons). Use the precise word; examiners penalise conflation.
What "HL standard" actually looks like
HLHL adds the convergence limit (ionisation from spectral data), trends and discontinuities in first ionisation energy across a period and down a group, and successive ionisation energy data as evidence for electron configuration and group number.
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
What determines the different chemical properties of atoms?
- Official IB
How do emission spectra provide evidence for the existence of different elements?
- Official IB
How does an element's highest main energy level relate to its period number in the periodic table?
- Official IB
What is the relationship between energy sublevels and the block nature of the periodic table?
- Official IB
How does the trend in IE values across a period and down a group explain the trends in properties of metals and non-metals?
- Official IB
Why are log scales useful when discussing [H⁺] and IEs?
- Official IB
How do patterns of successive IEs of transition elements help to explain the variable oxidation states of these elements?
- Official IB
How does the trend in successive ionization energies of transition elements explain their variable oxidation states?
- Official IB
Why do noble gases form covalent bonds less readily than other elements?
- Official IB
Why are oxygen and ozone dissociated by different wavelengths of light?
- Official IB
How can titration be used to calculate the concentration of an acid or base in solution?
1 further official linking question target a Tool, Inquiry or Nature of Science strand and are not drawn as edges on the map:
Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.
- Official IB
In the study of emission spectra from gaseous elements and of light, what qualitative and quantitative data can be collected from instruments such as gas discharge tubes and prisms?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.