HSC Exam Practice

Sample response. Electronegativity is a dimensionless measure of an atom’s tendency to attract bonding electrons toward itself when it is part of a covalent bond; measured on the Pauling scale. Pauling values: F = 4.0 (highest, most electronegative); Cs ≈ 0.7 (lowest). Fluorine is the most electronegative because it is in Period 2, Group 17: it has a very high effective nuclear charge (9 protons, 2 inner electrons giving Z_eff ≈ 7), a very small atomic radius, and its valence shell is close to the nucleus → extremely strong attraction for bonding electrons.

Marking notes. 1 mark for a correct definition of electronegativity (bonding electrons, attraction, covalent bond). 1 mark for correct Pauling values for F (4.0) and Cs (~0.7 or ~0.8). 1 mark for explaining F’s high EN using Z_eff, atomic radius, or period/group position.

Sample response. Across a period (left to right): Electronegativity increases. Each successive element has one more proton but electrons are added to the same principal shell, so shielding is roughly constant. Z_eff increases → stronger nuclear pull on bonding electrons → higher electronegativity. Atomic radius decreases across a period (bonding electrons are pulled closer to the nucleus), reinforcing the trend. Down a group: Electronegativity decreases. Each successive element has an additional electron shell, increasing atomic radius and inner-shell shielding. Bonding electrons are further from the nucleus and more shielded → weaker nuclear attraction for bonding electrons → lower electronegativity.

Marking notes. 1 mark for stating EN increases across a period. 1 mark for correct explanation referencing Z_eff increasing (or constant shielding with increasing protons). 1 mark for stating EN decreases down a group. 1 mark for correct explanation referencing increased atomic radius and/or increased shielding.

Sample response. (a) O–H: |Δχ| = 3.4 − 2.2 = 1.2. Since 0.4 ≤ 1.2 < 1.7 → polar covalent. δ− on O. (b) Na–Cl: |Δχ| = 3.2 − 0.9 = 2.3. Since 2.3 ≥ 1.7 → ionic. Full charge transfer: Na⁺ and Cl⁻. (c) Cl–Cl: |Δχ| = 3.2 − 3.2 = 0. Since 0 < 0.4 → non-polar covalent. No partial charges.

Marking notes. 1 mark per bond — must show |Δχ| calculation and correct classification. Award the mark even if the student omits δ− direction as long as classification is correct. Note: award 3 marks for the three bonds, final mark for correctly identifying Cl–Cl as non-polar due to |Δχ| = 0 (same atom).

Sample response. Going down a group, atomic radius increases (new electron shells added) and inner-shell shielding increases. For metals: metals react by losing valence electrons. The increased atomic radius and shielding reduce the effective nuclear charge experienced by the valence electron, decreasing ionisation energy → the valence electron is more easily lost → metallic reactivity increases down the group. For non-metals: non-metals react by gaining electrons. The increased atomic radius means an incoming electron would enter a valence shell that is further from the nucleus and more shielded → weaker nuclear attraction for the incoming electron → harder to gain an electron → non-metallic reactivity decreases down the group. Same structural change (larger, more shielded atom) — but opposite outcomes because one type loses and the other gains electrons.

Marking notes. 1 mark for correct metal trend (reactivity increases down group) with ionisation energy reference. 1 mark for structural reason for metal trend (atomic radius increases, valence electron more loosely held). 1 mark for correct non-metal trend (reactivity decreases down group) with electron gain reference. 1 mark for structural reason for non-metal trend (incoming electron enters more distant, more shielded shell).

Sample response. The student is incorrect because electronegativity and electron affinity, while related, measure different things. Electronegativity is a dimensionless relative measure of an atom’s ability to attract bonding electrons toward itself in a covalent bond, measured on the Pauling scale (0–4). Electron affinity is the energy change (in kJ/mol) when an isolated gaseous atom gains one electron to form an anion; it is a measurable thermodynamic quantity. The key difference is that electronegativity refers to behaviour in a bond (bonding electrons), whereas electron affinity describes the isolated atom gaining a free electron.

Marking notes. 1 mark for correct definition of electronegativity (bonding electrons, relative/Pauling scale). 1 mark for correct definition of electron affinity (energy change when isolated atom gains an electron; kJ/mol or thermodynamic quantity). 1 mark for clearly stating one key difference (bonding context vs isolated atom; relative scale vs measured energy; dimensionless vs energy units).

Sample response. In the O–H bond of water, |Δχ| = 3.4 − 2.2 = 1.2 (polar covalent). Oxygen is more electronegative (χ = 3.4) than hydrogen (χ = 2.2), so the bonding electrons are shifted toward O → oxygen carries δ− and hydrogen carries δ+. For Teflon (C–F bonds): χ(F) = 4.0, χ(C) = 2.6, so C–F is a highly polar covalent bond with F carrying δ−. Despite this polarity, Teflon is unreactive because F has the highest electronegativity of all elements and holds the bonding electrons so tightly that it is almost impossible for any other atom to attract them away — the C–F bond is very strong and F has no tendency to react further since it has already achieved a near-complete electron share.

Marking notes. 1 mark for identifying O as the atom carrying δ− in O–H, justified by O having higher χ (3.4 vs 2.2). 1 mark for explaining C–F polarity correctly (δ− on F). 1 mark for explaining Teflon’s unreactivity: F’s extreme electronegativity (highest, 4.0) holds bonding electrons so tightly that the C–F bond is very strong and F does not react further (accept: very strong bond, no tendency for F to attract more electrons once bonded, or F has a full valence shell in the bond).

Sample response (a) — Ranking (3 marks). Decreasing reactivity: Cl₂ > Br₂ > I₂ [1]. Evidence: Cl₂(aq) produces a colour change in both KBr and KI solutions (displaces both Br⁻ and I⁻), confirming it is the most reactive [1]. Br₂(aq) only produces a colour change in KI solution (displaces I⁻ but not Cl⁻), confirming it is intermediate in reactivity. I₂(aq) produces no colour change in any halide solution, confirming it is the least reactive — it cannot displace Cl⁻ or Br⁻ [1].

Sample response (b) — Periodic trend explanation (3 marks). Cl, Br, and I are all in Group 17. Cl is Period 3, Br is Period 4, and I is Period 5. Going down Group 17, atomic radius increases [1] because additional electron shells are added, increasing the distance between the nucleus and the outermost valence shell. Shielding by inner electrons also increases [1], so the effective nuclear charge experienced at the valence shell decreases. For halogens to react (gain an electron), the nucleus must attract an incoming electron; as the radius and shielding increase, this attraction weakens → Cl₂ gains an electron most readily (most reactive) and I₂ gains an electron least readily (least reactive) [1].

Sample response (c) — Mixed halide prediction (2 marks). I⁻ would be displaced first [1]. Cl₂ is a more powerful oxidising agent than both Br₂ and I₂. I⁻ is more easily oxidised than Br⁻ because iodide comes from the least reactive halogen (I₂); I⁻ has a greater tendency to lose an electron than Br⁻. Therefore, Cl₂ would first displace I⁻ to form I₂ (purple colour), and then (once I⁻ is consumed) would displace Br⁻ to form Br₂ (orange-brown). The purple colour of I₂ would appear first [1].

Marking notes. (a) 1 mark for correct ranking; 1 mark for evidence referencing Cl₂ displacing two halides; 1 mark for evidence referencing I₂ failing to displace any halide. (b) 1 mark for identifying increasing atomic radius down Group 17; 1 mark for identifying increasing shielding; 1 mark for linking to weaker attraction for incoming electron. (c) 1 mark for predicting I⁻ displaced first; 1 mark for correct reasoning (I⁻ more easily oxidised / Cl₂ preferentially oxidises the least reactive halide ion).

Sample response. The |Δχ| thresholds (non-polar: < 0.4; polar covalent: 0.4–1.7; ionic: ≥ 1.7) are useful, practical tools for predicting bond character in a wide range of everyday and industrial chemical contexts. For example, the Na–Cl bond in common table salt (|Δχ| = 2.3) is correctly classified as ionic, consistent with NaCl forming a crystal lattice, dissolving in water to give Na⁺ and Cl⁻ ions, and conducting electricity in solution — all properties expected of an ionic compound. Similarly, the O–H bond in water (|Δχ| = 1.2) is correctly classified as polar covalent, which accounts for water’s high boiling point, surface tension, and solvent properties that are central to the function of Australian desalination plants (such as the Perth Seawater Desalination Plant), where the polarity of water molecules drives the reverse osmosis process. The thresholds thus provide quick, reliable predictions for the large majority of inorganic compounds without requiring sophisticated quantum mechanical calculations.

However, the thresholds have significant limitations that prevent them from being described as “precise.” First, bond character is a continuum from fully covalent to fully ionic, not a set of discrete categories. The |Δχ| = 1.7 threshold is particularly problematic: hydrogen fluoride (HF, |Δχ| = 1.8) is classified as ionic by the threshold, yet HF is a discrete molecular compound, exists as a gas at room temperature, does not dissociate completely into ions in water (it is a weak acid), and clearly has a covalent H–F bond. Treating HF as ionic would lead to incorrect predictions about its physical state and reactivity. Second, the thresholds do not account for molecular geometry or resonance. Benzene (C–H and C–C bonds) has |Δχ| values suggesting non-polar to mildly polar covalent, but its delocalised electron system gives it very different reactivity (electrophilic aromatic substitution rather than addition reactions) that the |Δχ| approach cannot predict. Third, the thresholds apply to individual bonds, not entire molecules — CCl₄ has four polar C–Cl bonds but is a non-polar molecule due to its symmetric tetrahedral geometry, which the |Δχ| approach alone does not reveal. In summary, the |Δχ| thresholds are a useful, accessible, and broadly reliable first approximation for bond classification, especially for clear-cut ionic compounds and moderately polar covalent molecules. However, they are approximate guides rather than precise rules: they fail near the 1.7 boundary, cannot account for molecular geometry or resonance, and must always be applied with the acknowledgment that bond polarity is a continuum. Their utility is greatest when combined with other evidence (physical state, conductivity, melting point) to reach a well-supported classification.

Marking criteria (9 marks). 1 = states that thresholds are useful for a majority of compounds (general claim supported). 1 = provides a specific correct example of useful application, referencing an everyday or Australian industrial context (e.g. NaCl → table salt / desalination; H₂O / desalination or agriculture; HCl / industrial). 1 = provides a second specific, correctly analysed example. 1 = identifies the limitation that bond character is a continuum, not a set of discrete categories. 1 = uses H–F (or similar boundary compound) as evidence for the limitation of the 1.7 threshold with correct |Δχ| calculation. 1 = identifies at least one additional type of limitation (geometry/CCl₄; resonance/benzene; molecular vs bond polarity distinction). 1 = correctly analyses the additional limitation example. 1 = reaches an explicit, balanced evaluative judgement integrating both the usefulness and limitations. 1 = uses precise chemical terminology consistently throughout (ionic, polar covalent, non-polar covalent, |Δχ|, electronegativity, bond character continuum, δ+/δ−).