HSC Exam Practice

Sample response. An ionic bond is the electrostatic attraction between oppositely charged ions formed by the transfer of electrons from a metal atom (which loses electrons to become a cation) to a non-metal atom (which gains electrons to become an anion). An ionic lattice is a regular, repeating three-dimensional arrangement of cations and anions in which every cation is surrounded by anions and every anion is surrounded by cations. There are no discrete molecules in NaCl because each Na⁺ ion is attracted equally to all surrounding Cl⁻ ions; the formula NaCl simply represents the simplest whole-number ratio of ions in the giant structure.

Marking notes. 1 mark for a correct definition of ionic bond (electrostatic attraction, electron transfer, cation/anion); 1 mark for a correct definition of ionic lattice (regular, 3D, repeating, alternating ions); 1 mark for explaining why no discrete molecules exist (each ion attracted to all nearest neighbours, forming a giant structure; formula = simplest ratio only).

Sample response. Solid NaCl does not conduct electricity because the Na⁺ and Cl⁻ ions are fixed in rigid positions within the ionic lattice and cannot move. Since electrical conductivity requires mobile charge carriers, and the immobile ions cannot carry charge, solid NaCl is a non-conductor. When NaCl is heated above 801 °C and melts, the ionic lattice breaks down. The Na⁺ and Cl⁻ ions are then free to move independently through the liquid. Under an applied voltage, Na⁺ migrates toward the cathode and Cl⁻ migrates toward the anode, constituting an electrical current — hence excellent conductivity.

Marking notes. 1 mark for correctly identifying that ions are fixed/immobile in the solid lattice as the reason for non-conductivity; 1 mark for stating that the lattice breaks down on melting, releasing mobile ions; 1 mark for explaining that mobile ions carry charge to the appropriate electrodes, producing conductivity.

Sample response. Hardness refers to resistance to scratching or deformation. Ionic compounds are hard because the rigid, strongly bonded lattice resists displacement of ions from their equilibrium positions — large energy is required to push ions out of their regular sites. Brittleness refers to the tendency to shatter under impact rather than deforming. Ionic compounds are brittle because when an external force shifts one layer of ions relative to another, like-charged ions come into alignment. The strong repulsion between like charges (e.g. Na⁺ aligned with Na⁺) causes the crystal to split apart violently rather than deforming plastically. These two properties coexist: the lattice is hard (resists gentle deformation) but brittle (fails catastrophically under impact).

Marking notes. 1 mark for defining hardness and linking it to the rigid, strongly bonded lattice resisting displacement; 1 mark for explaining hardness in terms of the energy required to move ions from equilibrium positions; 1 mark for defining brittleness and linking it to lattice cleavage on impact; 1 mark for explaining brittleness using the like-charge repulsion mechanism when layers shift.

Sample response. NaCl contains Na⁺ (+1) and Cl⁻ (−1) ions; MgO contains Mg²⁺ (+2) and O²⁻ (−2) ions. According to Coulomb’s Law, the electrostatic force between two ions is proportional to the product of their charges divided by the square of the distance between them. For NaCl: force ∝ (1 × 1) = 1. For MgO: force ∝ (2 × 2) = 4. The electrostatic attraction in MgO is approximately four times stronger than in NaCl (assuming similar ionic radii). This means the lattice energy of MgO is much higher than that of NaCl. Melting requires enough thermal energy to overcome the electrostatic forces holding ions in the lattice; because MgO has ~4× stronger forces, it requires much more energy to melt — hence its melting point of 2852 °C versus 801 °C for NaCl.

Marking notes. 1 mark for correctly identifying the ion charges in both compounds (±1 for NaCl, ±2 for MgO); 1 mark for applying Coulomb’s Law quantitatively (force proportional to product of charges: 1 vs 4); 1 mark for linking the stronger force to higher lattice energy in MgO; 1 mark for connecting higher lattice energy to a higher temperature required to melt — hence higher melting point.

Sample response. The coordination number of Na⁺ in the NaCl lattice is 6. The six surrounding Cl⁻ ions are arranged in an octahedral geometry — one above, one below, and four around the equatorial plane of the Na⁺ ion. This demonstrates that NaCl is a giant structure rather than discrete molecules because there is no preferential “partner” Cl⁻ ion; each Na⁺ is equally attracted to all six nearest Cl⁻ neighbours, and the entire crystal is one continuous three-dimensional network. The formula NaCl only gives the simplest integer ratio of ions.

Marking notes. 1 mark for coordination number = 6; 1 mark for describing the octahedral arrangement (6 Cl⁻ surrounding each Na⁺: above, below, and four around); 1 mark for explaining that each ion is equally bonded to all nearest neighbours — no single partner — making it a giant structure, not discrete molecules.

Sample response. The student’s explanation is incorrect for ionic compounds. While it is true that ionic solids have no free electrons, this is not the reason they do not conduct: the correct explanation is that the ions are fixed in the ionic lattice and cannot move. Electrical conductivity requires mobile charge carriers; in solid ionic compounds, the potential charge carriers (ions) are immobile. The “no free electrons” explanation does apply to covalent (molecular) compounds such as glucose or wax, which neither have free electrons nor mobile ions and therefore cannot conduct in any state.

Marking notes. 1 mark for identifying the flaw: “no free electrons” is the wrong explanation for ionic compounds (even though the statement is true); 1 mark for stating the correct explanation: ions are fixed/immobile in the lattice and cannot move to carry charge; 1 mark for correctly naming the compound type to which “no free electrons” does apply (covalent molecular compounds, with a named example).

Sample response (a) — Classification. NaCl: ionic compound — no solid conductivity (ions fixed in lattice), excellent conductivity when molten (lattice broken, ions mobile) and when dissolved (ions hydrated and mobile) [1]. Glucose: covalent molecular compound — no conductivity in any state because it dissolves as neutral molecules, not ions, so there are no mobile charge carriers regardless of state [1]. Magnesium: metallic element — conducts as both solid and liquid because it has a sea of delocalised electrons that are always mobile, regardless of state [1]. Calcium bromide: ionic compound — same pattern as NaCl (no solid conductivity; conducts when molten and when dissolved), consistent with ions fixed in solid lattice and mobile when lattice is disrupted [1].

Sample response (b) — Mg vs NaCl solid conductivity. Magnesium conducts as a solid because its metallic lattice contains delocalised (free) electrons that can move throughout the lattice even when the metal is solid — electrons are always mobile in a metal [1]. Solid NaCl does not conduct because its charge carriers are ions, not electrons; the Na⁺ and Cl⁻ ions are locked in the ionic lattice and cannot move [1]. Therefore the charge carrier in solid Mg is electrons; in molten or dissolved NaCl the charge carrier is mobile ions (Na⁺ and Cl⁻) [1].

Sample response (c) — Glucose solution. A glucose solution does not conduct electricity because glucose dissolves as neutral molecules (C₆H₁₂O₆), not as ions. There are no mobile charge carriers in the solution [1]. To distinguish a glucose solution from a NaCl solution without a conductivity meter: taste is unscientific; an alternative experimental approach is to measure the freezing point depression — equal molar concentrations of NaCl depress the freezing point nearly twice as much as glucose (because NaCl dissociates into two ions per formula unit, while glucose remains as one particle); or use a flame test (NaCl solution gives a yellow flame due to Na⁺ emission; glucose solution does not). Accept any valid chemical distinction [1].

Marking notes. Part (a): 1 mark per correctly classified substance with a conductivity-based justification (4 marks total). Part (b): 1 mark for explaining metallic solid conductivity (delocalised electrons, always mobile); 1 mark for explaining ionic solid non-conductivity (ions fixed in lattice); 1 mark for correctly naming both charge carriers. Part (c): 1 mark for predicting non-conductivity with correct ionic/molecular reasoning; 1 mark for any valid distinguishing experimental method (freezing point, flame test, mass spectrometry — accept reasonable Year 11 method).

Sample response. The physical properties of ionic compounds arise directly from the structure of the ionic lattice — a giant, three-dimensional, regular arrangement of oppositely charged cations and anions held together by strong electrostatic (ionic) forces in all directions simultaneously. High melting point is the direct result of these strong forces. Melting requires enough thermal energy to overcome the electrostatic attractions and allow ions to move freely. The lattice energy — the energy holding the lattice together — depends on both ion charge and ion size (Coulomb’s Law: F ∝ q₁q₂/r²). Higher ion charges produce stronger forces: MgO (Mg²⁺/O²⁻, ±2 charges; force ∝ 4) melts at 2852 °C, while NaCl (Na⁺/Cl⁻, ±1; force ∝ 1) melts at only 801 °C. Smaller ions pack closer together, further increasing the force: LiF (small Li⁺, F⁻) has a higher MP than CsI (large Cs⁺, I⁻) despite identical ±1 charges. Hardness similarly arises from the rigid, strongly bonded lattice: displacing ions from their equilibrium sites requires overcoming many simultaneous ionic bonds, making the crystal resistant to scratching. However, the same lattice structure makes ionic compounds brittle: if a mechanical force shifts one layer of ions relative to the next, like-charged ions align; the strong repulsion between like charges (e.g. Na⁺-Na⁺) cleaves the crystal rather than allowing it to deform. This distinguishes ionic compounds from metals, which can deform because their non-directional metallic bonding allows layers to slide without producing repulsion. State-dependent conductivity is explained by ion mobility: in the solid state, all ions are fixed in lattice positions — no mobile charge carriers, no conductivity. When melted or dissolved in water, the lattice breaks down and ions become mobile, enabling conduction. In water, polar water molecules hydrolyse the lattice through hydration, releasing mobile ions. In an Australian industrial context, MgO is used as a refractory ceramic lining in high-temperature furnaces (e.g. electric arc furnaces at steel plants such as BlueScope Steel in Port Kembla) because its very high melting point (2852 °C) ensures it does not melt under operating conditions, its hardness resists mechanical wear, and its non-conductivity in the solid state means it can be used as electrical insulation without risk of short circuits. Al₂O₃ serves similar roles in aluminium smelting. In both cases, the high lattice energies produced by high-charge, small ions make these compounds uniquely suited to extreme industrial conditions. In summary, every observable property of an ionic compound — high MP, hardness, brittleness, and state-dependent conductivity — traces back to the strength, directionality, and rigidity of the ionic lattice, which is itself governed by ion charge and ion size through Coulomb’s Law.

Marking criteria (7 marks). 1 = high melting point explained via strong electrostatic forces and lattice energy, linked to charge and/or size (Coulomb’s Law). 1 = at least one quantitative comparison of MP using specific compounds (NaCl vs MgO, or LiF vs CsI), correctly linked to charge or size. 1 = hardness correctly attributed to rigid lattice and energy needed to displace ions. 1 = brittleness correctly attributed to like-charge alignment when layers shift, producing cleavage rather than plastic deformation. 1 = state-dependent conductivity correctly explained for all three states: solid (immobile ions), molten (mobile ions), dissolved (hydrated mobile ions). 1 = named Australian industrial context (Port Kembla / BlueScope / Snowy Scheme / Olympic Dam) with a specific compound used and a specific property that makes it suitable. 1 = explicit evaluative statement connecting ion charge/size to lattice energy to macroscopic property, integrating at least three properties in a single cohesive argument.