HSC Exam Practice

Sample response. The atomic number (Z) is the number of protons in the nucleus; it defines the element and is unique to each element. The mass number (A) is the total number of protons plus neutrons in the nucleus. The number of neutrons is found by subtracting the atomic number from the mass number: neutrons = A − Z.

Marking notes. 1 mark for correct definition of atomic number (number of protons, defines the element); 1 mark for correct definition of mass number (protons + neutrons); 1 mark for the correct neutron relationship (neutrons = A − Z).

Sample response. Thomson discovered the electron. He used a cathode ray tube, in which an electric current through a sealed low-pressure gas tube produced a beam of particles (cathode rays). He showed these rays were deflected by both electric and magnetic fields, proving they were charged particles. Key feature: electrons carry a negative charge (relative charge = −1) and have negligible mass (approximately 1/1836 of a proton mass).

Marking notes. 1 mark for identifying the electron; 1 mark for describing the cathode ray tube technique (deflection by E/B fields); 1 mark for one correct property of the electron (negative charge, or negligible mass, or specific charge ratio).

Sample response. In Thomson’s plum pudding model, the positive charge is spread uniformly throughout the atom in a diffuse sphere. If alpha particles (which are positively charged) passed through this uniform positive sphere, the repulsive force would be spread over the entire volume of the atom, producing only small, continuous deflections for all particles — no particle should be strongly deflected. The observation that a small fraction was deflected at angles greater than 90° (back-scattered) is impossible if positive charge is diffuse: such a reversal requires encountering an intensely concentrated positive region. The plum pudding model cannot generate the strong, localised repulsion needed to deflect a fast, massive alpha particle backwards.

Marking notes. 1 mark for explaining that the plum pudding model predicts only small/uniform deflections (diffuse positive charge); 1 mark for identifying that back-scatter (>90°) requires concentrated positive charge (not diffuse); 1 mark for explicitly stating why the plum pudding model cannot account for large-angle deflection (no localised repulsive force).

Sample response. (a) Electron arrangement: In Rutherford’s model, electrons orbit the nucleus at large distances in no specific pattern, like planets orbiting the sun, with no restriction on their energies. In Bohr’s model, electrons are restricted to fixed, discrete circular orbits (shells) at specific, quantised energy levels; an electron can only exist at these specific distances from the nucleus and cannot exist between levels. (b) Hydrogen emission spectrum: Rutherford’s model could not explain why hydrogen emits only specific, discrete wavelengths (lines) rather than a continuous spectrum — an electron orbiting at any distance would emit light of varying frequencies as it radiated. Bohr’s model explained the discrete lines exactly: only specific electron transitions between fixed energy levels are possible, so only specific photon energies (specific wavelengths) can be emitted, matching the observed hydrogen Balmer series wavelengths.

Marking notes. 1 mark for Rutherford electron arrangement (no fixed orbits / any distance); 1 mark for Bohr electron arrangement (fixed circular orbits at discrete quantised energy levels); 1 mark for Rutherford’s inability to explain discrete spectral lines; 1 mark for Bohr’s explanation of discrete lines via quantised transitions.

Sample response. (a) Protons = Z = 16. (b) Neutrons = A − Z = 32 − 16 = 16. (c) S²⁻ is an anion that has gained 2 electrons: electrons = 16 + 2 = 18. Check: 16 protons − 18 electrons = −2 charge (confirms 2−).

Marking notes. 1 mark for protons = 16; 1 mark for neutrons = 16 (with working A − Z shown); 1 mark for electrons = 18 with explanation that S²⁻ gains 2 electrons.

Sample response. The critical limitation was that classical physics (Maxwell’s electromagnetism) predicts that any electrically charged particle undergoing acceleration (including circular orbital motion) must continuously radiate electromagnetic energy. An orbiting electron would therefore lose energy continuously, spiralling inward and collapsing into the nucleus within a fraction of a nanosecond — atoms would be inherently unstable and could not exist, which contradicts observation. Rutherford’s model offered no explanation for why electrons remain in stable orbits. Niels Bohr (1913) addressed this by postulating that electrons occupy fixed, quantised energy levels in which they do not radiate energy; energy is only emitted or absorbed when an electron jumps between levels.

Marking notes. 1 mark for identifying the limitation (accelerating electrons radiate energy / electrons spiral into nucleus / atoms would be unstable); 1 mark for explaining the mechanism (classical physics requires radiating charged particle to lose energy); 1 mark for identifying Bohr (1913) as the scientist who resolved this, with reference to quantised energy levels.

Sample response (a). The red line at 656 nm corresponds to the electron transition from n = 3 to n = 2 (the smallest energy drop in the Balmer series). The 434 nm violet-blue line corresponds to the transition from n = 5 to n = 2 (a larger energy drop). The 434 nm line has a higher frequency than the 656 nm line because frequency and wavelength are inversely related (f = c/λ): a shorter wavelength means higher frequency. Higher frequency corresponds to greater photon energy (E = hf). The n = 5 → n = 2 transition spans a larger energy gap than n = 3 → n = 2; therefore the emitted photon has greater energy and higher frequency, corresponding to the shorter 434 nm wavelength.

Sample response (b). If electrons could exist at any energy (as classical physics assumed), they could transition between any two energies and emit photons of any frequency, producing a continuous spectrum. Instead, hydrogen emits only four specific visible wavelengths, proving that only specific electron transitions are permitted. This is only possible if the energy levels themselves are discrete (quantised) — electrons can only occupy certain fixed energy values, so only certain energy differences (and thus only certain photon frequencies) can be emitted. The discrete lines are direct experimental evidence for quantised energy levels, exactly as Bohr proposed.

Sample response (c). Bohr’s model assumed that electron orbits are simple circular paths and only worked for hydrogen (one electron). In multi-electron atoms, the model cannot accurately account for interactions between electrons, so its energy level predictions are inaccurate. The model also could not explain the fine structure (splitting) of spectral lines observed in multi-electron atoms, as described in the lesson. As a result, a more advanced quantum mechanical model was needed.

Marking notes. Part (a): 1 mark for correctly identifying n = 3 → n = 2 for the 656 nm line; 1 mark for explaining higher frequency = shorter wavelength; 1 mark for linking higher frequency to larger energy gap (n = 5 → n = 2 vs n = 3 → n = 2). Part (b): 1 mark for explaining that continuous spectrum would result if all energies were possible; 1 mark for linking discrete lines to fixed/quantised energy levels; 1 mark for stating this supports Bohr’s quantised model. Part (c): 1 mark for identifying electron–electron repulsion as missing from Bohr’s model; 1 mark for explaining this causes inaccurate energy level predictions in multi-electron atoms.

Sample response. Science does not advance by accumulating correct truths; it advances by building progressively better models that explain more evidence while retaining what worked before. The history of atomic models from Dalton to Bohr exemplifies this perfectly. Dalton’s solid sphere model (1803) was the first quantitative atomic theory. It explained the law of definite proportions (compounds form from fixed mass ratios) and conservation of mass (atoms are neither created nor destroyed in reactions). For 90 years it was entirely adequate because no evidence of internal atomic structure existed. The discovery of electrons by Thomson (1897) through cathode ray experiments created evidence that Dalton’s model could not accommodate: if atoms were indivisible, where did the electrons come from? Thomson revised the model by embedding electrons in a positive sphere (plum pudding, 1904). This retained Dalton’s key idea — atoms are overall neutral — while adding the concept of internal structure. Thomson’s model held until Rutherford’s gold foil experiment (1909–1911). The plum pudding model predicted only minor, uniform deflections of alpha particles from diffuse positive charge; instead, a small fraction deflected at angles greater than 90°, with roughly 1 in 20,000 reflecting almost straight back. A diffuse charge cannot produce these deflections; the data demanded a concentrated positive nucleus. Rutherford’s nuclear model retained Thomson’s ideas (electrons exist, atom is overall neutral) but restructured the atom around a dense nucleus surrounded by mostly empty space. Yet Rutherford’s model had its own fatal limitation: accelerating orbiting electrons should radiate energy continuously (classical electromagnetism), causing them to spiral inward — atoms should collapse in nanoseconds. Furthermore, hydrogen emits only discrete spectral lines, not a continuous spectrum; Rutherford’s model could not explain this. Bohr (1913) resolved both problems by postulating quantised energy levels. Electrons do not radiate when in fixed orbits; energy is only emitted or absorbed as photons when electrons jump between levels. This retained the nuclear structure from Rutherford while adding quantisation, and precisely predicted the wavelengths of hydrogen’s Balmer series. Each revision preserved working explanatory elements while adding new mechanisms to account for new evidence. This is precisely how science progresses: models are not discarded wholesale; they are improved. The provisional nature of scientific models is not a weakness — it is the mechanism by which understanding deepens when the frontier of evidence advances.

Marking criteria (7 marks). 1 = identifies and explains the evidence base that prompted the transition from Dalton to Thomson (electron discovery / cathode ray experiment). 1 = explains what Thomson’s model retained from Dalton and what it added. 1 = identifies and explains the specific gold foil evidence that Thomson’s model could not explain (large-angle deflection / back-scatter + reason: diffuse charge cannot produce it). 1 = explains how Rutherford’s nuclear model addressed the gold foil evidence while retaining prior concepts. 1 = identifies the specific limitation of Rutherford’s model that required revision (electron instability / discrete spectral lines) and explains why. 1 = explains how Bohr’s model addressed Rutherford’s limitation (quantised levels, no radiation in orbit, discrete transitions explain spectral lines). 1 = reaches an explicit evaluative statement that models are provisional, retained when explanatory, and revised when new evidence is anomalous — using the Dalton-to-Bohr progression as evidence.