Questions: Classification of Bonds: Ionic, Covalent, and Metallic

Question 1

Sodium chloride (NaCl) conducts electricity when dissolved in water but not in the solid state. Which explanation correctly follows from bond classification?

• A. In solution, water molecules break ionic bonds and release free electrons that carry current; in the solid, no free electrons exist
• B. In solution, the ions dissociate and become mobile, carrying charge; in the solid, ions are locked in the crystal lattice and cannot migrate
• C. Water itself is a good conductor and carries the current on behalf of the dissolved salt
• D. Ionic bonds break down at elevated temperatures like dissolution, releasing electrons

Answer: (B)

Electrical conductivity requires mobile charges. In solid NaCl, Na⁺ and Cl⁻ are locked in a rigid electrostatic lattice — they cannot migrate, so no current flows. Dissolving in water breaks the lattice, freeing the ions to move through the solution and carry charge. No free electrons are involved (options A and D are wrong — ionic conduction is via ion movement, not electrons). Water does not carry the current for the ions (option C). This prediction flows directly from ionic bond classification: ionic compounds conduct when ions are free to move.

Question 2

Nitrogen trifluoride (NF₃) has a ΔEN of approximately 1.0. A student concludes it must have ionic bonds because 'the ΔEN is substantial.' What is wrong with this reasoning?

• A. Electronegativity differences don't apply to compounds involving fluorine
• B. A ΔEN of 1.0 falls below the ~1.7 ionic threshold — it predicts polar covalent bonding where electrons are shared unequally, not transferred
• C. NF₃ is a gas at room temperature, which proves it must be covalent regardless of ΔEN
• D. Nitrogen and fluorine are both nonmetals, so they always form covalent bonds regardless of ΔEN

Answer: (B)

The student is confusing 'non-zero ΔEN' with 'ionic bonding.' The threshold for ionic character is roughly ΔEN > 1.7. At ΔEN = 1.0, bonding is polar covalent: fluorine pulls the shared electrons closer, creating partial charges (δ⁺ on N, δ⁻ on F), but the electrons are not fully transferred. Option D contains a useful heuristic but is not a principled explanation — the ΔEN framework is the correct approach, not element-type rules. The 1.7 threshold itself is a guideline, not a sharp boundary.

Question 3

The boundary between ionic and covalent bonding is sharp: a bond is either ionic or covalent, seldom intermediate.

• T. True
• F. False

Answer: False

Bonding is a continuous spectrum based on ΔEN, not three discrete categories. The 1.7 threshold is a guideline, not a sharp boundary. Real bonds near this value have partial ionic and partial covalent character — HCl, for example, has polar covalent bonding with measurable partial charges but is not fully ionic. As ΔEN increases from 0 (pure covalent, like H₂) toward large values (approaching pure ionic, like CsF at ΔEN ≈ 3.2), bonding character transitions gradually. The spectrum model is more accurate than a three-bin classification.

Question 4

Metals conduct electricity in the solid state because their atoms share localized electron pairs in covalent bonds, and those bonded electrons can move when a voltage is applied.

• T. True
• F. False

Answer: False

Metallic bonding is not localized electron-pair sharing — that is covalent bonding. In metallic bonding, valence electrons are delocalized across the entire lattice, forming a 'sea' of electrons not associated with any particular atom. These delocalized electrons move freely under an applied voltage, producing conductivity. Covalent compounds (diamond, plastic) do not conduct precisely because their electrons are localized in bonds and cannot migrate. The delocalized electron model explains not just conductivity but also malleability (ion layers can slide without breaking the electron sea) and metallic luster.

Question 5

Explain how identifying a compound's bond type (ionic, covalent, or metallic) allows you to predict its macroscopic properties without memorizing each compound individually.

Model answer: Bond type reflects how electrons are distributed: transferred (ionic), shared between specific atoms (covalent), or delocalized across a lattice (metallic). Each electron arrangement produces predictable macroscopic consequences. Ionic compounds form rigid lattices with high melting points and conduct when ions become mobile. Covalent compounds form discrete molecules with weak intermolecular forces (lower melting points) and no free charges. Metals conduct in the solid state due to their electron sea and are malleable because cation layers can slide without disrupting delocalized bonding.

This predictive power is the central payoff of bond classification. A new compound can be characterized by its ΔEN, its bonding type identified, and its properties predicted without memorization. The logic runs from electron behavior (how electrons are distributed) to atomic-scale structure (lattice vs. molecule vs. metallic array) to macroscopic properties (melting point, conductivity, solubility, malleability). Each step follows from the previous one. This is why bond classification is taught as a framework rather than a list — the framework generates the properties rather than requiring them to be individually memorized.