Topic 2.3: Structure of Ionic Solids Quiz

Prepare for AP Chemistry Topic 2.3 with this free practice quiz. Test your understanding of crystal lattices, ionic radii, and cleavage planes to maximize your AP score.

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Structure of Ionic Solids Quiz

21 MCQs

Topics Covered: Crystal Lattice, Ionic Radii, Cleavage Planes, Brittleness, Conductivity of Ionic Compounds

Description: This Topic 2.3 quiz tests your knowledge of how ions pack together in a solid state. Practice predicting macroscopic properties like brittleness and conductivity based on the microscopic ionic structure.

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1. A particular ionic compound is composed of cations (M⁺) and anions (X⁻). If the radius of X⁻ is significantly increased by substituting it with a heavier halogen, what is the expected effect on the crystal's melting point?

The melting point will increase due to greater London dispersion forces between the larger ions.

The melting point will increase because larger ions contain more electrons, increasing the Coulombic attraction.

The melting point will remain unchanged since the charges of the ions remain +1 and -1.

The melting point will decrease because the increased internuclear distance weakens the Coulombic forces of attraction.

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2. Which of the following particulate-level descriptions correctly explains why an ionic solid shatters when struck by a hammer?

The applied force breaks the localized covalent bonds between the atoms.

The applied force causes the delocalized electrons to shift, creating localized areas of positive charge that repel.

The applied force displaces a layer of ions, aligning ions of the same charge next to each other and causing electrostatic repulsion.

The applied force overcomes the weak London dispersion forces holding the macro-structure together.

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3. What explains the difference in physical properties between an ionic solid like NaCl and a metal like Ag, given both are crystalline?

NaCl conducts electricity in the solid phase because its ions vibrate, while Ag does not.

NaCl consists of discrete molecules held by intermolecular forces, while Ag is a continuous network.

Ag has higher lattice energy due to the +1 charge on every atom, making it harder than NaCl.

The ions in NaCl are held by non-directional electrostatic forces that lock them in place, making it brittle, while the delocalized electrons in Ag allow planes of atoms to slide without breaking bonds.

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4. A chemist analyzes an ionic solid and finds that applying pressure to one side of the crystal causes it to instantly shatter rather than bend. Which of the following describes the particulate-level mechanism of this shattering?

Pressure causes the localized electron pairs in the covalent bonds to decouple.

Pressure squeezes the electron clouds too close together, causing quantum mechanical tunneling that destabilizes the nucleus.

Pressure lowers the melting point of the solid until it sublimates directly into a gas.

Pressure shifts the rows of alternating ions so that cations align with cations and anions with anions, creating sudden, overwhelming repulsive forces.

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5. A student hypothesizes that solid CaCl₂ will conduct electricity better than solid KCl because calcium has a +2 charge. Why is this hypothesis incorrect?

Calcium chloride is covalent in the solid state and only becomes ionic when dissolved in water.

The calcium ion is smaller than the potassium ion, restricting its movement in the lattice.

Neither compound has mobile valence electrons or mobile ions in the solid state to carry an electrical current.

KCl has a more exothermic lattice energy, allowing electrons to move more freely.

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6. A beaker contains molten MgCl₂. When electrodes connected to a battery and lightbulb are placed in the molten liquid, the bulb lights up. The moving charge carriers in the liquid are:

Mg²⁺ and Cl⁻ ions.

Intact MgCl₂ molecules.

Water molecules that have ionized due to the heat.

Delocalized valence electrons.

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7. Which of the following best explains why the melting point of MgO (2852 °C) is significantly higher than that of NaF (993 °C)?

The Mg²⁺ and O²⁻ ions have larger ionic radii than Na⁺ and F⁻ ions.

Oxygen is more electronegative than fluorine, creating a more polarized bond.

The Mg²⁺ and O²⁻ ions possess greater magnitudes of charge than Na⁺ and F⁻ ions, resulting in stronger Coulombic attractions.

MgO contains covalent network bonds, whereas NaF consists of purely ionic bonds.

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8. Which of the following trends is most likely observed regarding the solubility of simple Group 1 chloride salts (e.g., LiCl, NaCl, KCl) in water?

The lattice energy increases down the group, making it harder for water to hydrate the ions.

The salts become increasingly brittle and therefore harder to dissolve.

Solubility is impossible because the Coulombic attractions in the solid are always stronger than ion-dipole attractions.

The energy required to break the ionic lattice generally decreases down the group as the ionic radius of the cation increases.

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9. The lattice energy of CaO(s) is -3414 kJ/mol, while the lattice energy of BaO(s) is -3029 kJ/mol. Which statement best explains this data?

Oxygen in BaO has a smaller radius than oxygen in CaO.

Ca²⁺ has a higher charge than Ba²⁺, leading to stronger attractions in CaO.

Ba²⁺ is larger than Ca²⁺, resulting in a greater internuclear distance and weaker Coulombic attractions in BaO.

Ba has a higher electronegativity than Ca, leading to less ionic character in BaO.

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10. How does the arrangement of ions in a crystal lattice maximize the stability of an ionic compound?

It aligns all the dipoles in a single direction to generate a macroscopic magnetic field.

It maximizes the attractive forces between oppositely charged ions while minimizing the repulsive forces between similarly charged ions.

It allows every ion to share electrons equally with all adjacent ions.

It creates a sea of delocalized electrons that cement the positive nuclei together.

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11. A substance has the following properties: melting point = 2050 °C, boils at 2980 °C, rigid, brittle, and non-conductive as a solid but highly conductive as a liquid. The substance is most likely:

C₁₂H₂₂O₁₁

SiO₂

Al₂O₃

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12. The standard enthalpy of formation for solid NaCl from elemental sodium and chlorine includes several energy steps (Born-Haber cycle). Which of these steps corresponds strictly to the formation of the ionic crystal lattice from its constituent gaseous ions?

Na⁺(aq) + Cl⁻(aq) → NaCl(s)

NaCl(s) → Na⁺(g) + Cl⁻(g)

Na(s) + ½ Cl₂(g) → NaCl(s)

Na⁺(g) + Cl⁻(g) → NaCl(s)

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13. Scandium oxide (Sc₂O₃) and titanium dioxide (TiO₂) are both solid ionic compounds. If the lattice energy of TiO₂ is greater in magnitude than that of Sc₂O₃, which of the following best explains this observation?

TiO₂ has a larger molar mass, leading to stronger London dispersion forces.

Sc₂O₃ contains more atoms per formula unit than TiO₂, spreading the lattice energy over more bonds.

The oxide ions in TiO₂ have a smaller radius than the oxide ions in Sc₂O₃.

The Ti⁴⁺ ion has a higher charge than the Sc³⁺ ion, resulting in a stronger Coulombic attraction with the oxide ions.

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14. Which set of ions is arranged in order of decreasing ionic radius?

Mg²⁺, Na⁺, F⁻, O²⁻, N³⁻

N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺

N³⁻, Na⁺, O²⁻, Mg²⁺, F⁻

F⁻, O²⁻, N³⁻, Mg²⁺, Na⁺

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15. Which experimental technique would provide the best macroscopic evidence for the presence of distinct cleavage planes in an ionic crystalline solid?

Dissolving the solid in water and testing for electrical conductivity.

Applying a sharp, localized mechanical stress to the crystal and observing the geometry of the resulting fragments.

Placing the crystal in a mass spectrometer to determine the isotopic masses.

Heating the solid in a crucible and recording the temperature at which it melts.

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16. Consider the ionic compounds SrS and RbCl. Which is expected to be more brittle and have a higher melting point, and why?

SrS, because the +2 and -2 charges create a much stronger Coulombic attraction, locking the ions more rigidly into the lattice.

SrS, because strontium is a transition metal and forms complex crystal structures.

RbCl, because the larger total mass of RbCl relative to SrS increases the gravitational attraction in the lattice.

RbCl, because the +1 and -1 charges allow the lattice to slip more easily, building up kinetic energy.

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17. A solid sample is brittle, does not conduct electricity as a solid, but conducts electricity highly when dissolved in water. Which of the following types of structures does the solid most likely have?

Molecular solid

Metallic crystal lattice

Covalent network

Ionic crystal lattice

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18. When a student tests a crystalline solid, it shatters cleanly along flat surfaces. This macroscopic property, known as a cleavage plane, is a direct result of:

The mobility of valence electrons in the solid phase.

The random, amorphous arrangement of the constituent ions.

The presence of polar covalent bonds that snap under mechanical stress.

The highly regular, repeating three-dimensional arrangement of ions in the crystal lattice.

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19. In an experiment, an unknown solid substance is found to be highly soluble in water. The resulting solution conducts electricity, but the solid does not. Which combination of subatomic interactions is primarily responsible for the lack of conductivity in the solid state?

The covalent sharing of electrons between atoms prevents any charge mobility.

The metallic bonding creates localized bands of electrons that cannot traverse the lattice.

The dispersion forces are too weak to allow electrons to hop from one molecule to another.

The strong electrostatic attractions lock the oppositely charged ions in a rigid three-dimensional matrix.

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20. Why do transition metals frequently form ionic compounds with varying colors and crystal structures, compared to Group 1 and 2 metals?

Transition metals form primarily covalent network solids rather than ionic lattices.

Transition metal cations are larger than Group 1 anions, disrupting normal packing.

Transition metals possess lower effective nuclear charges than alkali metals.

Transition metals can lose different numbers of d-orbital electrons, forming cations of varying charges and radii.

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21. Based on periodic trends and Coulomb's law, which of the following ionic compounds would be expected to have the most negative (most exothermic) lattice energy?

LiF

LiCl

RbI

KBr

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The Pre-Quiz Review

Crystal Lattice

An ionic solid consists of positive ions (cations) and negative ions (anions) arranged in a repeating three-dimensional structure called a crystal lattice.

The lattice is held together by electrostatic attractions between oppositely charged ions. Each ion is surrounded by ions of the opposite charge, so the attraction is not limited to one cation-anion pair.

The strength of ionic interactions depends mainly on:

Magnitude of ionic charges: higher charges produce stronger electrostatic attraction.
Distance between ion centers: smaller ions allow ions to approach more closely, increasing attraction.
Arrangement of ions: the three-dimensional lattice determines how many oppositely charged ions surround each ion and how efficiently the structure packs.

For ions with charges q₁ and q₂, electrostatic potential energy follows the relationship:

E ∝ q₁q₂/r

where r is the distance between the ion centers. For oppositely charged ions, the interaction is attractive.

This explains why MgO has much stronger ionic attractions than NaCl: Mg²⁺ and O²⁻ have larger charge magnitudes than Na⁺ and Cl⁻.

A crystal lattice must maintain overall electrical neutrality. The formula of an ionic compound represents the simplest whole-number ratio of ions that gives a neutral compound. For example:

NaCl → 1 Na⁺ : 1 Cl⁻
MgCl₂ → 1 Mg²⁺ : 2 Cl⁻
Al₂O₃ → 2 Al³⁺ : 3 O²⁻

The formula does not represent discrete molecules of the compound. It represents the ratio of ions throughout the extended lattice.

Lattice Energy

Lattice energy describes the energy associated with separating or forming an ionic solid from its gaseous ions. AP Chemistry questions may use either convention, so always check how the term is defined in the question.

For the formation process:

gaseous ions → ionic solid

energy is released.

For the reverse process:

ionic solid → gaseous ions

energy is required.

Lattice energy becomes greater in magnitude when:

Ion charges increase.
Ionic radii decrease.
The distance between oppositely charged ions decreases.

Thus, comparing compounds with similar structures:

MgO has a greater-magnitude lattice energy than NaF because the ions have higher charges.

For compounds with the same charges, smaller ions generally produce stronger attractions. LiF therefore has stronger ionic attractions than CsF because Li⁺ is much smaller than Cs⁺.

Lattice energy is related to, but not identical to, melting point, boiling point, and solubility. Stronger ionic attractions generally require more energy to disrupt the lattice, but physical properties depend on the entire structure and interactions involved.

Ionic Radii

Ionic radius is the effective size of an ion within a crystal structure.

Cations are smaller than their neutral atoms because losing electrons reduces electron-electron repulsion and may remove an entire outer energy level.

For example:

Na > Na⁺

Anions are larger than their neutral atoms because gaining electrons increases electron-electron repulsion.

For example:

Cl⁻ > Cl

Periodic Trends in Ionic Radius

For ions with the same charge type within a period, radius generally decreases from left to right as effective nuclear charge increases.

Down a group, ionic radius generally increases because additional principal energy levels are occupied.

Examples:

Li⁺ < Na⁺ < K⁺ < Rb⁺

F⁻ < Cl⁻ < Br⁻ < I⁻

Isoelectronic Ions

Isoelectronic ions have the same number of electrons.

For an isoelectronic series, ionic radius decreases as nuclear charge increases.

For example:

O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺

All have 10 electrons.

Their radii decrease in this order:

O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺

The nucleus with more protons attracts the same electron cloud more strongly.

This is a frequent basis for AP Chemistry comparison questions.

Ionic Size and Crystal Structure

Ionic radii influence how ions fit into a lattice. The relative sizes of cations and anions affect coordination and the geometry of the crystal structure.

A smaller cation can fit into spaces between larger anions. The ratio of cation radius to anion radius can therefore help determine the coordination environment in an ionic solid.

AP questions may provide ionic radii and ask you to compare structures or explain why changing one ion changes the arrangement.

When comparing ionic attraction, remember that the relevant distance is approximately related to the sum of the ionic radii:

r ≈ r₊ + r₋

Smaller ions generally permit a smaller separation between ion centers, increasing electrostatic attraction.

Cleavage Planes

Ionic crystals have ordered layers of ions. This regular structure produces characteristic planes along which the crystal can break, called cleavage planes.

A cleavage plane forms where shifting layers of ions causes ions with the same charge to become adjacent.

For example, consider two neighboring layers containing alternating cations and anions. If the layers shift slightly, a cation can be brought next to another cation, while an anion can be brought next to another anion.

Like charges repel strongly.

The resulting electrostatic repulsion can cause the crystal to split along a particular plane.

This explains why ionic solids are generally brittle rather than malleable.

The key structural sequence is:

Ionic solid contains an ordered lattice.
Mechanical force shifts layers.
Like charges can become aligned.
Strong electrostatic repulsion develops.
The lattice fractures along a plane.

This is fundamentally different from metallic bonding. In metals, layers of atoms can shift while metallic bonding remains relatively intact because the valence electrons are delocalized. Ionic lattices do not have this same ability to accommodate layer movement.

Brittleness of Ionic Solids

Ionic solids are hard but brittle.

Their hardness results from strong electrostatic attractions throughout the crystal lattice. A significant amount of energy is required to disrupt the ordered structure.

Their brittleness results from the effect of shifting lattice layers. A small displacement can place ions of the same charge next to one another, producing strong repulsion and causing the crystal to fracture.

Therefore:

Strong ionic attractions → rigid, hard structure.
Layer displacement → like charges become adjacent.
Like-charge repulsion → fracture.

When an AP question asks why an ionic solid shatters rather than bends, focus on the arrangement and movement of ions in the lattice.

Conductivity of Ionic Compounds

Electrical conductivity requires mobile charged particles.

In a solid ionic compound, the ions are present but fixed in their lattice positions. They cannot move freely through the crystal.

Therefore, a solid ionic compound generally does not conduct electricity.

When the ionic compound is melted, the lattice structure breaks down sufficiently for the ions to move. The ions can then transport charge through the liquid.

Therefore:

Solid ionic compound → ions fixed → poor electrical conductor.
Molten ionic compound → ions mobile → conducts electricity.
Aqueous ionic solution → ions dispersed and mobile → conducts electricity.

The conductivity of an ionic solution depends on the presence and concentration of mobile ions.

For example:

NaCl(s) → poor conductor

NaCl(l) → conducts

NaCl(aq) → conducts

In aqueous solution, dissolved ions are hydrated by water molecules and can move independently through the solution.

Dissociation in Water

When a soluble ionic compound dissolves, its crystal lattice is disrupted and the ions become surrounded by water molecules.

For example:

NaCl(s) → Na⁺(aq) + Cl⁻(aq)

For MgCl₂:

MgCl₂(s) → Mg²⁺(aq) + 2Cl⁻(aq)

The ions, rather than intact NaCl or MgCl₂ molecules, are responsible for electrical conductivity.

The ability of an ionic compound to dissolve is determined by competition between ion-ion attractions in the lattice and ion-dipole attractions between the ions and water.

Water is polar. Its oxygen atom has a partial negative charge and its hydrogen atoms have partial positive charges.

Therefore:

Cations are attracted to the oxygen end of water.
Anions are attracted to the hydrogen end of water.

Dissolution is favorable when the interactions formed between ions and water sufficiently compensate for the energy required to separate the ions in the lattice.

Solubility and conductivity should therefore be treated as related but distinct properties. A substance must produce mobile charged particles to conduct, but the extent of conductivity depends on how many mobile ions are present and how freely they move.

Comparing Ionic Compounds

When comparing ionic solids, identify the variables that control electrostatic attraction.

For two compounds with similarly structured lattices:

Higher ionic charges → stronger attraction.
Smaller ionic radii → stronger attraction.
Smaller ion-to-ion distance → stronger attraction.

For example:

NaCl contains Na⁺ and Cl⁻.

MgO contains Mg²⁺ and O²⁻.

Because MgO contains ions with charges of ±2, its electrostatic attractions are substantially stronger than those between Na⁺ and Cl⁻.

If the charges are the same, compare ion size.

For example:

LiF vs. CsF

Both contain +1 and −1 ions, so the smaller Li⁺ allows closer approach to F⁻. This produces stronger electrostatic attraction.

Reading Particle Diagrams

AP Chemistry questions frequently use particle diagrams to represent ionic solids.

When interpreting a diagram, look for:

Alternating positive and negative ions.
Repeating three-dimensional arrangement.
Relative ion sizes.
Displacement of lattice layers.
Whether particles are fixed or mobile.
Formation of separated ions in solution.

A regular repeating arrangement indicates a crystalline solid.

A diagram showing ions locked into fixed positions represents a solid ionic lattice.

A diagram showing separated ions moving through the sample can represent a molten ionic compound or an aqueous ionic solution, depending on the surrounding particles.

For cleavage questions, look specifically at what happens after layers shift. If like charges are brought into neighboring positions, repulsion increases and the crystal can fracture.

For conductivity questions, focus on particle mobility rather than simply asking whether ions are present. Ions exist in a solid ionic crystal, but they are not freely mobile.

AP Chemistry Comparison Framework

For most structure-of-ionic-solid questions, analyze the structure in this order:

Identify the ions and their charges.
Compare their ionic radii.
Estimate the distance between ion centers.
Compare electrostatic attraction.
Relate the attraction to lattice stability or energy required to disrupt the structure.
For conductivity, determine whether the ions are mobile.
For brittleness, determine whether layer movement brings like charges together.

The central connection is:

ionic charge + ionic size → electrostatic attraction → lattice structure and stability → observable properties

Properties commonly connected to ionic structure include hardness, brittleness, melting point, lattice energy, and electrical conductivity under different physical conditions.

Report a question

The Pre-Quiz Review

Crystal Lattice

Lattice Energy

Ionic Radii

Periodic Trends in Ionic Radius

Isoelectronic Ions

Ionic Size and Crystal Structure

Cleavage Planes

Brittleness of Ionic Solids

Conductivity of Ionic Compounds

Dissociation in Water

Comparing Ionic Compounds

Reading Particle Diagrams

AP Chemistry Comparison Framework

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