Questions: Ionic Bonding: Electron Transfer and Electrostatic Forces

Question 1

Which pair would most likely form an ionic compound?

• A. Carbon and oxygen (electronegativities 2.5 and 3.4, difference 0.9)
• B. Hydrogen and chlorine (electronegativities 2.1 and 3.2, difference 1.1)
• C. Potassium and fluorine (electronegativities 0.8 and 4.0, difference 3.2)
• D. Nitrogen and oxygen (electronegativities 3.0 and 3.4, difference 0.4)

Answer: (C)

Ionic bonding requires a large electronegativity difference, conventionally greater than ~1.7. Potassium–fluorine has a difference of 3.2, far above the threshold, so fluorine pulls the electron completely off potassium to form K⁺F⁻. The other pairs have differences below 1.7 and form polar covalent bonds. This illustrates that ionic bonding is the extreme case of electronegativity difference, not simply the pairing of any metal with any nonmetal.

Question 2

A student argues that solid NaCl conducts electricity because Na⁺ and Cl⁻ ions are present in the crystal. What is wrong with this reasoning?

• A. NaCl does not contain ions — it is a covalent compound in the solid state
• B. Solid NaCl is an insulator because the ions are locked in the crystal lattice and cannot move; conduction requires free charge carriers
• C. The ions are present but cancel each other out electrically, so there is no net charge to carry
• D. The student is correct — solid NaCl does conduct electricity

Answer: (B)

The presence of ions is necessary but not sufficient for electrical conduction — the ions must also be free to move. In solid NaCl, each ion is locked in a rigid crystal lattice by electrostatic attraction on all sides. NaCl only conducts when dissolved in water (ions separate and move freely) or when melted (the lattice breaks down). This distinguishes ionic compounds from metals, where electrons (not ions) are the free charge carriers.

Question 3

Ionic compounds consist of discrete molecules — the formula NaCl represents one Na⁺ bonded to one Cl⁻, just as H₂O represents two H atoms bonded to one O.

• T. True
• F. False

Answer: False

Ionic compounds do not form discrete molecules. Each Na⁺ attracts every surrounding Cl⁻ and vice versa, forming an extended three-dimensional crystal lattice. The formula NaCl is an empirical formula representing the simplest ratio of ions, not a single bonded pair. There is no specific 'Na–Cl bond' — just the totality of electrostatic attractions throughout the lattice. This is fundamentally different from covalent molecules like H₂O, where specific atoms bond to specific other atoms.

Question 4

Ionic crystals are hard because the electrostatic bonds are strong, but they are also flexible because ions can slide past each other while maintaining attraction.

• T. True
• F. False

Answer: False

Ionic crystals are hard but brittle, not flexible. When a layer of ions is displaced, ions of the same charge end up adjacent, causing powerful electrostatic repulsion that shatters the crystal along cleavage planes. Unlike metals (where electron clouds allow layers to slide), ionic crystals cannot deform without fracturing. Brittleness and hardness are both direct consequences of the electrostatic lattice structure.

Question 5

Why do ionic compounds have high melting points? Explain the reasoning from the atomic-level structure rather than simply stating 'ionic bonds are strong.'

Model answer: In an ionic crystal lattice, each ion is held in place by electrostatic attraction to all its neighbors simultaneously — not just one partner but many. To melt the crystal, every ion must gain enough energy to escape its lattice position by overcoming these many simultaneous Coulomb attractions. This is a collective process requiring large energy input. The Coulomb force is strong, acts in all directions from each ion, and must be disrupted cooperatively across the entire lattice — explaining the high melting temperature.

The key is the extended lattice structure: unlike a covalent molecule where melting just requires overcoming intermolecular forces, ionic melting requires dismantling the entire periodic array of strong electrostatic attractions. Both the strength of individual Coulomb forces and the cooperative, three-dimensional nature of the lattice contribute to the high melting point.