HSC Exam Practice

Sample response. A period is a horizontal row in the periodic table; all elements in the same period have the same number of occupied electron shells (e.g. all Period 3 elements have 3 shells). A group is a vertical column; all elements in the same group have the same number of valence electrons and therefore similar chemical properties (e.g. all Group 1 elements have 1 valence electron).

Marking notes. 1 mark for correct definition of period (horizontal row); 1 mark for correctly identifying same number of occupied electron shells as the shared structural feature; 1 mark for correct definition of group (vertical column); 1 mark for correctly identifying same number of valence electrons as the shared structural feature.

Sample response. (a) Mg (Z = 12): Group 2, Period 3, s-block, 2 valence electrons. (b) Cl (Z = 17): Group 17, Period 3, p-block, 7 valence electrons. (c) Fe (Z = 26): Group 8, Period 4, d-block, 2 valence electrons (d-block exception — accept 2 or 8 with valid reasoning).

Marking notes. 2 marks per element: 1 mark for correct group + period, 1 mark for correct block + valence electrons. Accept Fe with either 2 valence electrons (s-block convention) or note of d-block complexity.

Sample response. Mendeleev’s table organised elements in order of increasing relative atomic mass, whereas the modern periodic table organises by increasing atomic number (Z). In Mendeleev’s arrangement, tellurium (Ar = 127.6) would appear after iodine (Ar = 126.9) based on mass; however, their chemical properties require Te (Group 16) to come before I (Group 17). The modern table places Te (Z = 52) before I (Z = 53) correctly.

Marking notes. 1 mark for correctly identifying the two different organising principles (mass vs Z); 1 mark for a valid example of different ordering (Te/I, Co/Ni accepted); 1 mark for explaining why the modern arrangement is correct for that example.

Sample response. Going down Group 1, atomic radius increases because each element gains an additional electron shell. The single valence electron is progressively further from the nucleus and shielded by more inner shells, so the effective nuclear attraction holding it decreases. This makes the valence electron easier to lose in reactions — hence reactivity increases (e.g. Cs > K > Na > Li). For Group 17 halogens, reactivity depends on the ability to gain an electron. Going down the group, atoms grow larger; the valence shell is further from the nucleus and more shielded, so the nuclear attraction for an incoming electron is weaker. This makes it harder to gain an electron — hence reactivity decreases (e.g. F > Cl > Br > I).

Marking notes. 1 mark for Group 1 trend direction (increases) with reference to atomic radius; 1 mark for Group 1 explanation (larger radius + shielding → easier to lose valence electron); 1 mark for Group 17 trend direction (decreases) with reference to atomic radius; 1 mark for Group 17 explanation (larger radius + shielding → harder to attract incoming electron).

Sample response. The student is incorrect: noble gases have the same number of electrons as any other element. Their valence shells are complete — noble gases (except He) have 8 valence electrons (helium has 2), giving them a full outermost shell. A complete valence shell confers exceptional stability: there is no tendency to gain, lose, or share electrons to achieve a more stable configuration. It is this full-shell configuration, not the absence of electrons, that makes noble gases essentially unreactive.

Marking notes. 1 mark for identifying the flaw (noble gases do have electrons; statement is wrong); 1 mark for correctly stating the valence electron configuration (full valence shell: 8, or 2 for He); 1 mark for explaining why a full shell = low reactivity (no tendency to gain, lose, or share electrons).

Sample response. Elements in the same group have the same number of valence electrons. Because chemical reactions involve the interactions of valence electrons (gaining, losing, or sharing), elements with identical valence electron counts react in analogous ways, forming the same types of bonds and ions. Example: sodium (Na, Group 1, Period 3) and potassium (K, Group 1, Period 4) both have 1 valence electron. Both react vigorously with water to produce a metal hydroxide and hydrogen gas: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g); 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g).

Marking notes. 1 mark for the core explanation (same group = same valence electron count = similar bonding / reactivity behaviour); 1 mark for naming a valid pair of elements in the same group; 1 mark for describing a correct shared chemical property (reaction type, ion formed, bonding behaviour — must match the named pair).

Sample response (a) — Panel A trend and explanation. Atomic radius decreases from Na (186 pm) to Cl (99 pm) across Period 3, almost halving across the period [1 — describe trend with data]. This occurs because all Period 3 elements have the same number of electron shells (3). However, as atomic number increases from Na (Z = 11) to Cl (Z = 17), the nuclear charge increases while the amount of shielding by inner shells remains roughly constant. The increased effective nuclear charge (Z_eff) pulls the valence electrons closer to the nucleus, decreasing atomic radius [1 — Z_eff explanation]. Award a third mark for quantitative reference to the data in the explanation [1].

Sample response (b) — Panel B trend and explanation. Atomic radius increases from Li (152 pm) to Cs (265 pm) going down Group 1 [1 — describe trend with data]. This occurs because each successive element down the group has an additional electron shell; the outermost valence electron occupies a shell that is further from the nucleus [1 — electron shells explanation]. Additionally, the increasing number of inner electron shells provides greater shielding of the nuclear charge, reducing Z_eff on the outer electron, which allows the electron cloud to expand [1 — shielding explanation].

Sample response (c) — connection to reactivity. For metals, a larger atomic radius means the valence electron is further from the nucleus and more easily lost. From Panel B, Cs (265 pm) has the largest atomic radius and is therefore the most reactive Group 1 metal; K (227 pm) is more reactive than Na (186 pm). From Panel A, Na (186 pm) has a much larger radius than Cl (99 pm); Na readily loses its outermost electron (metal reactivity), while Cl, with its small radius and high Z_eff, tends to gain an electron (non-metal reactivity) [1 per valid point, max 2].

Marking notes. Part (a): 3 marks as per criteria above. Part (b): 3 marks as above. Part (c): 2 marks for correctly linking larger radius to lower ionisation energy / easier electron loss and making a valid comparison using a named element from each panel.

Sample response. The claim that the periodic table is “just a way to organise known elements” significantly undervalues its scientific power. The table emerged from Mendeleev’s observation of periodicity — the repeating pattern of properties when elements are arranged in order of increasing atomic mass (later refined to atomic number). This pattern was powerful enough to predict the existence and properties of undiscovered elements from blank spaces in the table. For example, Mendeleev predicted “eka-aluminium” (later gallium, discovered 1875) and “eka-silicon” (germanium, 1886), accurately forecasting their atomic mass, density, oxide and chloride formulas, and boiling points of their chlorides — before anyone had observed these elements. The close agreement between prediction and measurement (e.g. germanium density: predicted 5.5, measured 5.35 g cm⁻³) demonstrates that the table captures a deeper structural reality. This predictive power is explained by electron configuration: atomic number determines the electron configuration of an atom, which in turn determines all chemical and many physical properties. Group position predicts valence electron count and therefore reactivity type; period position predicts shell count and atomic radius. Using these rules, properties of any element — including undiscovered ones — can be forecast. Periodic trends such as ionisation energy, electronegativity, and atomic radius all follow predictable patterns from group and period position. For unknown elements (e.g. hypothetical Z = 119 would be in Group 1, Period 8, s-block — a highly reactive alkali metal more reactive than Cs), the table generates precise, testable predictions without experimental data. In conclusion, the periodic table is far more than an organisational chart. It is a theory-driven predictive framework grounded in atomic theory. The claim is false: the periodic table’s power lies precisely in its ability to predict the properties of elements not yet known — a standard of scientific utility that few other single frameworks in science can match.

Marking criteria (7 marks). 1 = correctly identifies and explains periodicity as the underlying pattern (repeating properties with Z). 1 = historical evidence: names at least one correct predicted element (gallium or germanium) with a specific matched property. 1 = second piece of historical evidence or a quantitative comparison of prediction vs measurement for the named element. 1 = explains how group and period position predicts properties of unknown elements (references valence electrons, shell count, or a trend). 1 = applies reasoning to a concrete example of an unknown/undiscovered element with a valid prediction (e.g. Z = 119). 1 = evaluates the original claim explicitly — reaches a clear judgement that the table is a predictive framework, not merely organisational. 1 = consistent use of precise terminology: periodicity, atomic number, electron configuration, valence electrons, effective nuclear charge, group, period (at least 4 terms used correctly).