HSC Exam Practice

Sample response. The Aufbau (building-up) principle states that electrons occupy the lowest available energy subshell first when filling an atom from the ground state. Filling order to 4p: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p. Note that 4s fills before 3d because 4s has lower energy for most elements.

Marking notes. 1 mark for a correct definition of the Aufbau principle (lowest energy first); 1 mark for the correct sequence up to 4s (1s 2s 2p 3s 3p 4s); 1 mark for correctly placing 3d after 4s and before 4p in the sequence.

Sample response. s: 2 electrons (1 orbital × 2). p: 6 electrons (3 orbitals × 2). d: 10 electrons (5 orbitals × 2). f: 14 electrons (7 orbitals × 2). The limits arise because each orbital holds a maximum of 2 electrons with opposite spins (Pauli exclusion principle), and the number of orbitals per subshell is 2l + 1, where l = 0 (s), 1 (p), 2 (d), or 3 (f).

Marking notes. 1 mark per correct subshell capacity (s=2, p=6, d=10, f=14). Award the 4th mark for a correct explanation referencing the Pauli exclusion principle or the 2l+1 orbital count. Do not award the explanation mark for just listing numbers without a reason.

Sample response. The Pauli exclusion principle governs the capacity of each individual orbital: no two electrons in the same atom can share identical quantum numbers, so each orbital holds at most two electrons and they must have opposite spins (↑↓). Hund’s rule governs how electrons distribute among the orbitals of a subshell: electrons fill each orbital singly with parallel spins before any orbital is doubly occupied, minimising electron–electron repulsion. For carbon (Z=6), the configuration is 1s²2s²2p²; the two 2p electrons occupy two separate 2p orbitals singly (↑ ↑ empty) by Hund’s rule, not both in the same orbital (↑↓ empty empty) which would violate Hund’s rule.

Marking notes. 1 mark for correctly defining the Pauli exclusion principle (orbital capacity, opposite spins); 1 mark for correctly defining Hund’s rule (fill singly before pairing, parallel spins); 1 mark for writing the correct configuration for carbon 1s²2s²2p²; 1 mark for correctly describing the 2p orbital arrangement (two singly-occupied p orbitals with parallel spins).

Sample response. The filling order is the sequence in which subshells accept electrons, determined by the Aufbau principle: 4s fills before 3d. The written order is the convention for expressing the final configuration: subshells are listed in ascending shell (n) number, so 3d is written before 4s in the notation. For iron (Z=26): filling order has 4s filled before 3d, but the written configuration is 1s²2s²2p⁶3s²3p⁶3d⁶4s² (abbreviated [Ar]3d⁶4s²), where 3d appears before 4s in the written text because shell 3 comes before shell 4.

Marking notes. 1 mark for explaining that 4s fills before 3d (filling order); 1 mark for explaining the written convention (numerical shell order, so 3d is written before 4s); 1 mark for the correct written configuration of iron (full or abbreviated form).

Sample response. Chromium (Z=24): Expected [Ar]3d⁴4s²; actual [Ar]3d⁵4s¹. One electron is promoted from 4s to 3d, achieving a half-filled d subshell where all five 3d orbitals are singly occupied (↑ ↑ ↑ ↑ ↑). The extra stability arises from exchange energy — the increased number of same-spin electron pairs in the half-filled configuration lowers the total energy, outweighing the cost of the promotion. Copper (Z=29): Expected [Ar]3d⁹4s²; actual [Ar]3d¹⁰4s¹. One electron is promoted from 4s to 3d, achieving a fully-filled d subshell (3d¹⁰), which has similarly enhanced exchange stability.

Marking notes. 1 mark for stating the expected and actual configuration of Cr; 1 mark for explaining Cr anomaly using half-filled d stability / exchange energy; 1 mark for stating the expected and actual configuration of Cu; 1 mark for explaining Cu anomaly using fully-filled d stability / exchange energy.

Sample response. The period number equals the highest principal quantum number (n) of occupied subshells in the ground-state configuration. The group number equals the number of valence electrons (for main-group elements). Silicon (Z=14): 1s²2s²2p⁶3s²3p². Highest n = 3, so Period 3. Valence electrons = 3s²3p² = 4 electrons, so Group 14 (or Group IV in old notation). Silicon is in the p-block because its highest-energy electrons are in a p subshell.

Marking notes. 1 mark for correctly stating that period = highest n value; 1 mark for correctly stating that group = number of valence electrons (main group); 1 mark for correctly applying both rules to silicon (Period 3, Group 14) with a written configuration.

Sample response (a) — Ranking: The energy of the emitted photon equals the energy difference between the two levels. From the diagram: A (Z→W) spans the largest gap (highest to lowest level) = highest energy photon. B (Z→X) spans the Z-to-X gap = second highest. C (Y→W) spans the Y-to-W gap. D (Y→X) spans the smallest gap (adjacent levels) = lowest energy photon. Ranking in decreasing photon energy: A > B > C > D. (Note: exact relative ordering of B and C depends on the actual energy spacings shown; from the diagram spacings, Z→W > Z→X > Y→W > Y→X is consistent with the diagram.) [1 mark for correct identification of A as largest energy; 1 mark for D as smallest; 1 mark for a justified complete ranking].

Sample response (b): Transition A (Z→W) produces UV light, which has high frequency and high energy (E = hf). This indicates the energy gap between levels Z and W is large. Transition D (Y→X) produces infrared light, which has lower frequency and lower energy. This indicates the energy gap between levels Y and X is smaller than the Z→W gap. [1 mark for correctly linking UV to large energy gap Z→W; 1 mark for linking infrared to small energy gap Y→X].

Sample response (c): Sodium (Z=11) has the ground-state configuration 1s²2s²2p⁶3s¹. When a sodium atom is excited (e.g. in a flame), the single 3s valence electron is promoted to higher energy subshells (3p, 4s, etc.). As the electron falls back to the ground state, photons of specific wavelengths are emitted corresponding to the fixed energy differences between sodium’s subshells. These wavelengths form the unique sodium emission spectrum (notably two bright yellow-orange D lines at ~589 nm). Because each element has a unique set of energy levels determined by its electron configuration, each element produces a unique spectral fingerprint. [1 mark for linking emission to electron transitions between specific energy levels; 1 mark for explaining uniqueness of spectrum reflects unique energy-level spacing / electron configuration].

Sample response. The three rules governing electron configuration — the Aufbau principle, Pauli exclusion principle, and Hund’s rule — are powerful predictive tools that correctly describe the electron arrangements of the vast majority of elements and provide the theoretical basis for the structure of the periodic table. The Aufbau principle correctly predicts the filling order (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p…) and explains why periods 1, 2, and 3 of the periodic table contain 2, 8, and 8 elements respectively: period 1 fills the 1s subshell (2 elements), period 2 fills 2s and 2p (8 elements), period 3 fills 3s and 3p (8 elements), and period 4 fills 4s, 3d, and 4p (18 elements). The Pauli exclusion principle correctly limits each orbital to two electrons and predicts that paired electrons have opposite spins, consistent with experimental spectroscopic data. Hund’s rule correctly predicts that nitrogen (Z=7) has three singly-occupied 2p orbitals rather than one doubly-occupied orbital, which is confirmed by magnetic measurements showing N is paramagnetic with three unpaired electrons. Together, the three rules correctly predict electron configurations for the majority of elements (Z=1 to 18) without exception. However, the rules have significant limitations for transition metals. The Aufbau principle alone predicts [Ar]3d⁴4s² for chromium and [Ar]3d⁹4s² for copper, but the actual configurations are [Ar]3d⁵4s¹ and [Ar]3d¹⁰4s¹ respectively. These anomalies arise because the exchange energy gained from having a half-filled (Cr) or fully-filled (Cu) d subshell exceeds the energy cost of promoting one electron from 4s to 3d — a quantum-mechanical effect that the simple Aufbau building-up picture does not capture. Similar anomalies occur in the second- and third-row transition metals and in the lanthanides and actinides, where 4f and 5f subshells introduce additional complexity. A further limitation is that the rules describe isolated, neutral ground-state atoms; they do not directly predict configurations of ions (for which 4s electrons are removed before 3d in transition metal cations, reversing the apparent filling priority). Despite these limitations, the three rules remain the standard HSC framework because they predict the correct configuration for 90%+ of elements in the first four periods and provide an elegant mechanistic link between quantum mechanics (quantum numbers, spin) and the observable macroscopic structure of the periodic table. In summary, the rules are highly useful tools with clear predictive power across most of the periodic table, but students must be aware of their boundaries — particularly the anomalous configurations of Cr and Cu — to apply them accurately in examination contexts.

Marking criteria (7 marks). 1 = correctly states and applies the Aufbau principle with the filling order including 4s before 3d, linking it to period lengths in the periodic table. 1 = correctly states and applies the Pauli exclusion principle (orbital capacity, opposite spins) with at least one specific example. 1 = correctly states and applies Hund’s rule (singly fill before pairing) with at least one specific element example showing the orbital diagram rationale. 1 = identifies the anomalous configurations of Cr and Cu by name, states both expected and actual configurations, and identifies the type of stability involved (half-filled / fully-filled d subshell, exchange energy). 1 = discusses at least one limitation of the rules beyond Cr and Cu (e.g. ions, heavier elements, lanthanides, or bulk vs atomic behaviour). 1 = explicitly evaluates the rules as useful but imperfect — must include a comparative judgement (e.g. “highly useful for Z=1–18… but fail for Cr and Cu”). 1 = uses precise chemical/quantum terminology throughout (Aufbau, Pauli exclusion, Hund’s rule, exchange energy, principal quantum number, subshell, orbital, spin, anomalous, paramagnetic or related terms).