College Chemistry Quiz: Properties Of Solids
20 questions · exam conditions
0:00
Properties Of SolidsQuestion 1 of 20

A solid has a melting point of 1545°C, conducts electricity both as a solid and liquid, and can be hammered into thin sheets. However, when dissolved in water, the resulting solution does not conduct electricity. What type of solid is this?

Ionic solid that decomposes upon dissolution
Metallic solid that forms covalent compounds in solution
Covalent network solid with some metallic character
Molecular solid with unusually strong intermolecular forces
Metallic solid that is insoluble in water
← Back to quizzes

College Chemistry Quiz

College Chemistry Quiz: Properties Of Solids

Practice Properties Of Solids in College Chemistry with focused quiz questions that help you check what you know, review explanations, and build confidence with test-style prompts.

What this quiz covers

This quiz focuses on Properties Of Solids, giving you a quick way to practice the rules, question types, and explanations that matter most for College Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

All questions

Question 1

A solid has a melting point of 1545°C, conducts electricity both as a solid and liquid, and can be hammered into thin sheets. However, when dissolved in water, the resulting solution does not conduct electricity. What type of solid is this?

  1. Ionic solid that decomposes upon dissolution
  2. Metallic solid that forms covalent compounds in solution
  3. Covalent network solid with some metallic character
  4. Molecular solid with unusually strong intermolecular forces
  5. Metallic solid that is insoluble in water (correct answer)
Explanation: This question tests your ability to classify solids based on their physical and chemical properties. When analyzing an unknown solid, you need to consider multiple characteristics together to determine its bonding type and structure. The key properties here paint a clear picture: the extremely high melting point (1545°C), electrical conductivity in both solid and liquid states, and malleability (ability to be hammered into sheets) are all classic characteristics of metallic solids. Metals have delocalized electrons that allow conductivity and metallic bonding that enables malleability while maintaining strong attractions between atoms, resulting in high melting points. The crucial clue is that the solution doesn't conduct electricity when dissolved in water. This indicates the metal doesn't ionize significantly in solution, meaning it forms covalent compounds with water or simply doesn't dissolve appreciably. Many transition metals behave this way, forming neutral complexes or remaining largely undissolved. Choice A is incorrect because ionic solids are typically brittle (not malleable) and don't conduct electricity as solids. Choice C is wrong because covalent network solids, while having high melting points, are typically electrical insulators as solids and aren't malleable. Choice D fails because molecular solids have much lower melting points and don't conduct electricity. Remember that metallic properties cluster together: high melting points, conductivity, and malleability. When you see these three properties combined, especially with non-conducting aqueous solutions, think of metals that don't readily ionize in water, like many transition metals.

Question 2

A metallic solid has a face-centered cubic (FCC) structure with an edge length of 4.08 × 10⁻¹⁰ m. How many atoms are contained in one unit cell of this structure?

  1. 1 atom
  2. 2 atoms
  3. 3 atoms
  4. 4 atoms (correct answer)
  5. 8 atoms
Explanation: When analyzing crystal structures, you need to understand how atoms are arranged and shared between adjacent unit cells. The key is counting the effective number of atoms that belong to each unit cell. In a face-centered cubic (FCC) structure, atoms are located at two positions: the eight corners of the cube and the centers of all six faces. To find the total atoms per unit cell, you must account for how these atoms are shared with neighboring unit cells. Corner atoms are each shared by eight adjacent unit cells, so each corner contributes 18\frac{1}{8} atom to the unit cell. With 8 corners, this gives: 8×18=18 \times \frac{1}{8} = 1 atom. Face-centered atoms are shared between only two unit cells (the cell in question and its neighbor), so each face contributes 12\frac{1}{2} atom. With 6 faces, this gives: 6×12=36 \times \frac{1}{2} = 3 atoms. Total atoms per unit cell: 1+3=41 + 3 = 4 atoms. Option A (1 atom) only counts the corner contribution while ignoring face-centered atoms entirely. Option B (2 atoms) might result from incorrectly assuming face atoms contribute 13\frac{1}{3} each rather than 12\frac{1}{2}. Option C (3 atoms) counts only the face-centered contribution while neglecting the corners. Option D correctly accounts for both corner and face contributions. Remember this pattern: simple cubic has 1 atom per unit cell, body-centered cubic has 2, and face-centered cubic has 4. Always consider atom sharing when counting atoms in crystal structures.

Question 3

Silicon carbide (SiC) has a very high melting point (2830°C), is extremely hard, and does not conduct electricity. It is insoluble in water and most solvents. These properties are most consistent with which type of crystalline solid?

  1. Ionic solid with strong electrostatic interactions
  2. Metallic solid with delocalized electron sea
  3. Covalent network solid with extended three-dimensional bonding (correct answer)
  4. Molecular solid with strong hydrogen bonding
  5. Molecular solid with strong dipole-dipole interactions
Explanation: When you encounter a question about crystalline solids, focus on matching the given properties to the characteristic features of each solid type. The key is recognizing how atomic-level bonding translates to macroscopic properties. Silicon carbide's properties point directly to a covalent network solid. The extremely high melting point (2830°C) indicates very strong bonds that require enormous energy to break. The hardness suggests atoms are locked in a rigid three-dimensional framework where every atom is covalently bonded to multiple neighbors. The lack of electrical conductivity tells you there are no mobile electrons or ions. These characteristics perfectly match covalent network solids like diamond, quartz, and SiC, where strong covalent bonds extend throughout the entire crystal structure. Let's examine why the other options don't fit. Choice A (ionic solid) fails because while ionic solids can have high melting points, SiC's extreme hardness and the specific combination of Si and C (both nonmetals with similar electronegativity) don't support significant ionic character. Choice B (metallic solid) is wrong because metals conduct electricity due to their delocalized electron sea—exactly the opposite of SiC's insulating behavior. Choice D (molecular solid) doesn't work because molecular solids typically have much lower melting points since you're only breaking relatively weak intermolecular forces, not covalent bonds. Study tip: Remember the property patterns—covalent network solids are the "extreme" solids with the highest melting points, greatest hardness, and complete lack of conductivity among the main solid types.

Question 4

The density of a simple cubic crystal is 5.96 g/cm³. The atomic mass of the element is 87.6 g/mol, and each unit cell contains 1 atom. What is the edge length of the unit cell?

  1. 2.85 × 10⁻⁸ cm
  2. 3.42 × 10⁻⁸ cm (correct answer)
  3. 4.18 × 10⁻⁸ cm
  4. 5.67 × 10⁻⁸ cm
  5. 6.24 × 10⁻⁸ cm
Explanation: When you encounter crystal structure problems, you're working with the fundamental relationship between atomic-scale dimensions and bulk properties. The key is connecting density (a macroscopic property) to unit cell dimensions (atomic-scale measurements) through the crystal's atomic arrangement. For this simple cubic crystal, start with the density equation: density=mass of unit cellvolume of unit cell\text{density} = \frac{\text{mass of unit cell}}{\text{volume of unit cell}}. Since each unit cell contains 1 atom, the mass equals 87.6 g/mol6.022×1023 atoms/mol=1.455×10−22 g/atom\frac{87.6 \text{ g/mol}}{6.022 \times 10^{23} \text{ atoms/mol}} = 1.455 \times 10^{-22} \text{ g/atom}. The unit cell volume is a3a^3 (where aa is the edge length), so: 5.96=1.455×10−22a35.96 = \frac{1.455 \times 10^{-22}}{a^3} Solving for a3a^3: a3=1.455×10−225.96=2.441×10−23 cm3a^3 = \frac{1.455 \times 10^{-22}}{5.96} = 2.441 \times 10^{-23} \text{ cm}^3 Therefore: a=2.441×10−233=3.42×10−8 cma = \sqrt[3]{2.441 \times 10^{-23}} = 3.42 \times 10^{-8} \text{ cm}, confirming answer B. Answer A (2.85 × 10⁻⁸ cm) would result from incorrectly using 2 atoms per unit cell instead of 1. Answer C (4.18 × 10⁻⁸ cm) suggests an error in the cube root calculation or unit conversion. Answer D (5.67 × 10⁻⁸ cm) likely stems from forgetting to convert atomic mass units properly or miscalculating Avogadro's number. Remember: always verify that your crystal structure matches the given atoms per unit cell—simple cubic has 1, body-centered cubic has 2, and face-centered cubic has 4.

Question 5

The coordination number of an atom in a crystal structure refers to the number of nearest neighboring atoms. What is the coordination number for atoms in a face-centered cubic (FCC) structure?

  1. 4
  2. 6
  3. 8
  4. 10
  5. 12 (correct answer)
Explanation: When analyzing crystal structures, coordination number tells you how many atoms are in direct contact with any given atom. To find this for face-centered cubic (FCC) structures, you need to visualize the atomic arrangement and identify which atoms actually touch. In an FCC structure, atoms are positioned at each corner of the cube and at the center of each face. The key insight is determining which atoms are truly "nearest neighbors" - those that physically touch. If you imagine an atom at a corner position, it doesn't actually touch the corner atoms of the same unit cell. Instead, it touches the face-centered atoms from adjacent unit cells. Each corner atom touches atoms on the faces of surrounding cubes. When you account for all directions (up, down, and the four horizontal directions), plus the periodic nature of the crystal extending in three dimensions, each atom contacts exactly 12 nearest neighbors. Looking at the given choices: A) 4 represents the coordination number for diamond or zinc blende structures. B) 6 corresponds to simple cubic structures where atoms touch only along the face directions. C) 8 matches body-centered cubic (BCC) structures. D) 10 doesn't correspond to any common crystal structure. However, there appears to be an error in the question setup - the correct answer for FCC should be 12, but this isn't listed among the choices A-D, and option E isn't provided in the question. Study tip: Memorize the coordination numbers for common structures: simple cubic (6), BCC (8), FCC (12), and practice visualizing 3D atomic arrangements to avoid confusion.

Question 6

Solid iodine (I₂) sublimes readily at room temperature, has a low melting point (113.7°C), and is a poor conductor of electricity. These properties indicate that solid iodine is which type of solid?

  1. An ionic solid with weak electrostatic forces between ions
  2. A metallic solid with a low density of free electrons
  3. A covalent network solid with some defects in the structure
  4. A molecular solid held together by weak intermolecular forces (correct answer)
  5. An amorphous solid with no regular crystal structure
Explanation: When you encounter properties like sublimation, low melting point, and poor electrical conductivity, you're looking at clues about the type of bonding and intermolecular forces in a solid. These properties tell you how strongly the constituent particles are held together. Iodine's behavior points to a molecular solid. As I₂, iodine exists as discrete diatomic molecules held together by weak van der Waals forces (specifically London dispersion forces). These weak intermolecular attractions explain all three observed properties: molecules can easily escape to the gas phase (sublimation), require little energy to disrupt the crystal structure (low melting point), and provide no mobile electrons for conduction. Choice A is incorrect because ionic solids, even with weak electrostatic forces, would still conduct electricity when molten and typically have much higher melting points than 113.7°C. Choice B fails because metallic solids always conduct electricity due to their "sea" of delocalized electrons, regardless of electron density. Choice C is wrong because covalent network solids like diamond or quartz have very high melting points due to strong covalent bonds throughout the structure, and defects wouldn't change this fundamental characteristic. When analyzing solid types, remember this pattern: molecular solids show low melting/boiling points, poor electrical conductivity, and easy sublimation due to weak intermolecular forces. If you see a nonmetal element or simple compound with these "weak" properties, think molecular solid first. The key is recognizing that weak intermolecular forces, not strong chemical bonds, dominate the solid's behavior.

Question 7

Graphite can conduct electricity along its planes but not perpendicular to them, while diamond is an electrical insulator in all directions. This difference is best explained by which structural feature?

  1. Graphite has ionic character while diamond is purely covalent
  2. Graphite has delocalized π electrons while diamond has only localized σ bonds (correct answer)
  3. Graphite has weaker C-C bonds than diamond in all directions
  4. Graphite has metallic bonding while diamond has covalent bonding
  5. Graphite has fewer electrons per carbon atom than diamond
Explanation: When you encounter questions about electrical conductivity in different materials, focus on the electronic structure and bonding types present. The key is understanding how electrons can move through the material's structure. Graphite and diamond are both carbon allotropes, but their dramatically different electrical properties stem from their distinct bonding arrangements. In graphite, carbon atoms form hexagonal sheets where each carbon is sp2sp^2 hybridized and bonded to three other carbons. This leaves one unhybridized p orbital per carbon atom, creating a system of delocalized π electrons that can move freely within each plane. These mobile electrons act like an "electron sea" along the graphite planes, enabling electrical conductivity in those directions. However, the planes are held together only by weak van der Waals forces with no mobile electrons between them, so conductivity perpendicular to the planes is negligible. Diamond has a completely different structure where each carbon is sp3sp^3 hybridized and forms four strong covalent σ bonds in a tetrahedral arrangement. All electrons are localized in these bonds, leaving no mobile charge carriers in any direction. Choice A is wrong because both materials are covalently bonded. Choice C incorrectly suggests graphite has universally weaker bonds—actually, the C-C bonds within graphite planes are quite strong. Choice D is incorrect because graphite doesn't have true metallic bonding; it has covalent bonding with delocalized π electrons. Remember: electrical conductivity requires mobile electrons. Always look for delocalized electron systems (π bonds, metallic bonding) when explaining conductivity in materials.

Question 8

The packing efficiency of a crystal structure is the percentage of space occupied by atoms. Which crystal structure has the highest packing efficiency?

  1. Simple cubic (atoms at corners only)
  2. Body-centered cubic (BCC)
  3. Face-centered cubic (FCC)
  4. Hexagonal close-packed (HCP)
  5. Both FCC and HCP have equally high packing efficiency (correct answer)
Explanation: When you encounter questions about crystal structure packing efficiency, you're dealing with how efficiently spherical atoms can be arranged in three-dimensional space. This is a fundamental concept in solid-state chemistry that determines many material properties. The correct answer is actually both C (Face-centered cubic) and D (Hexagonal close-packed), as they share the highest packing efficiency at 74%. Both structures achieve this maximum efficiency through close-packed arrangements where each atom is surrounded by 12 nearest neighbors. In FCC, atoms occupy corner and face-center positions, creating layers that stack in an ABCABC pattern. HCP achieves the same efficiency through ABAB stacking of close-packed layers. Choice A (simple cubic) has the lowest packing efficiency at only 52% because atoms touch only along cube edges, leaving large gaps between spheres. This creates a very open, inefficient structure. Choice B (body-centered cubic) achieves 68% packing efficiency. While better than simple cubic, it's still less efficient than close-packed structures because the central atom doesn't create the optimal close-packed geometry. The question appears to have an error since the correct answer is listed as "E" but only four options are provided. Both FCC and HCP represent the theoretical maximum for packing identical spheres. Study tip: Remember the packing efficiency hierarchy: Simple cubic (52%) < BCC (68%) < FCC = HCP (74%). Close-packed structures (FCC and HCP) always achieve maximum efficiency because they eliminate wasted space through optimal sphere arrangements.

Question 9

Zinc sulfide (ZnS) crystallizes in a structure where each Zn²⁺ ion is surrounded by four S²⁻ ions in a tetrahedral arrangement, and each S²⁻ ion is surrounded by four Zn²⁺ ions. What is the coordination number for both ions in this structure?

  1. Zn²⁺: 4, S²⁻: 6
  2. Zn²⁺: 6, S²⁻: 4
  3. Zn²⁺: 4, S²⁻: 4 (correct answer)
  4. Zn²⁺: 6, S²⁻: 6
  5. Zn²⁺: 8, S²⁻: 8
Explanation: When you encounter crystal structure problems, focus on coordination number—the number of nearest neighbor ions that directly surround a central ion. This is determined by counting actual contacts, not by charge considerations or geometric assumptions. The problem directly states the key information: each Zn²⁺ ion is surrounded by four S²⁻ ions in tetrahedral arrangement, and each S²⁻ ion is surrounded by four Zn²⁺ ions. This means both ions have a coordination number of 4, making answer C correct. Let's examine why the other options are wrong. Answer A suggests Zn²⁺ has coordination number 4 (correct) but S²⁻ has coordination number 6 (incorrect). This contradicts the given information that each sulfide ion is surrounded by four zinc ions. Answer B reverses the relationship, giving Zn²⁺ a coordination number of 6 and S²⁻ a coordination number of 4. This ignores the stated tetrahedral arrangement around zinc, which specifically involves four neighbors. Answer D assigns coordination number 6 to both ions, which would describe an octahedral arrangement—completely different from the tetrahedral geometry described. The zinc blende (ZnS) structure described here is a classic example where both cation and anion have the same coordination number despite different charges. This differs from structures like fluorite or rutile where cations and anions have different coordination numbers. Study tip: Always read crystal structure problems carefully for explicit geometric information. When the problem states specific arrangements (tetrahedral, octahedral), use that to determine coordination numbers rather than making assumptions based on ionic charges or ratios.

Question 10

Silicon and germanium are both semiconductors with similar crystal structures, but silicon has a larger band gap than germanium. How does this difference affect their electrical conductivity at room temperature?

  1. Silicon has higher conductivity because its larger band gap creates more charge carriers
  2. Germanium has higher conductivity because its smaller band gap allows easier electron promotion (correct answer)
  3. Both have identical conductivity because they have the same crystal structure
  4. Silicon has higher conductivity because it has smaller atoms
  5. Germanium has higher conductivity because it has more valence electrons
Explanation: When you encounter semiconductor questions, focus on the relationship between band gap energy and electron mobility. The band gap is the energy difference between the valence band (where electrons are bound) and the conduction band (where electrons can move freely and conduct electricity). A smaller band gap means electrons need less thermal energy to jump from the valence band to the conduction band. At room temperature, the available thermal energy is fixed, so materials with smaller band gaps will have more electrons promoted to the conduction band, creating more charge carriers and higher conductivity. Since germanium has a smaller band gap than silicon, it more easily promotes electrons at room temperature, making it more conductive. Answer A incorrectly suggests that larger band gaps create more charge carriers—this is backwards. Larger band gaps actually make it harder for electrons to reach the conduction band, reducing the number of charge carriers. Answer C ignores the crucial role of band gap energy. While crystal structure affects how electrons move once they're in the conduction band, the band gap determines how many electrons get there in the first place. Answer D confuses atomic size with electrical properties. Although silicon atoms are indeed smaller, this doesn't directly determine conductivity—the band gap energy does. Remember this pattern: smaller band gap = easier electron promotion = higher conductivity at a given temperature. This principle explains why germanium was used in early transistors despite silicon's other advantages, and why different semiconductors are chosen for specific temperature ranges in modern electronics.

Question 11

A material scientist measures the hardness of four different solids using the same testing method. The results show: Sample A (hardness 9), Sample B (hardness 2), Sample C (hardness 6), Sample D (hardness 1). Based on hardness alone, which sample is most likely a covalent network solid?

  1. Sample A (correct answer)
  2. Sample B
  3. Sample C
  4. Sample D
  5. Cannot be determined from hardness alone
Explanation: When you encounter questions about material properties, think about how different types of bonding affect physical characteristics like hardness. The key insight is that stronger bonding leads to harder materials. Covalent network solids have atoms connected by strong covalent bonds extending throughout the entire crystal structure. This creates extremely hard materials because breaking the solid requires breaking many strong covalent bonds. Classic examples include diamond (hardness 10 on Mohs scale) and silicon carbide (hardness 9-10). These materials rank among the hardest substances known. Sample A, with a hardness of 9, fits perfectly with the expected properties of a covalent network solid. This high hardness indicates the presence of strong, directional covalent bonds throughout the structure that resist deformation. Sample B (hardness 2) is too soft to be a covalent network solid. This hardness suggests a material like gypsum or graphite, where weak intermolecular forces or layered structures allow easy deformation. Sample C (hardness 6) could be a moderately hard ionic solid like feldspar, but lacks the extreme hardness characteristic of covalent networks. Sample D (hardness 1) represents a very soft material like talc, indicating weak bonding forces such as van der Waals interactions between layers or molecules. Remember this pattern: when predicting material properties from bonding type, covalent network solids will always be among the hardest materials due to their extended three-dimensional network of strong covalent bonds. Look for hardness values of 7 or higher when identifying potential covalent network solids.

Question 12

Tungsten carbide (WC) is used in cutting tools because of its exceptional hardness, while tungsten metal is used in light bulb filaments because of its high melting point and electrical conductivity. Which statement best explains these different applications?

  1. WC is a molecular solid while tungsten is a covalent network solid
  2. WC is a covalent network solid while tungsten is a metallic solid (correct answer)
  3. Both are metallic solids but WC has stronger metallic bonding
  4. WC is an ionic solid while tungsten is a metallic solid
  5. Both are covalent network solids but with different crystal structures
Explanation: When you encounter questions about material properties and their applications, focus on identifying the type of solid structure, as this directly determines the material's characteristics. Tungsten carbide (WC) forms a covalent network solid where tungsten and carbon atoms are connected by strong covalent bonds in a three-dimensional network. This extensive covalent bonding throughout the structure creates exceptional hardness, making WC ideal for cutting tools that must resist deformation and wear. Tungsten metal, however, is a metallic solid with a "sea of electrons" that can move freely throughout the structure. This electron mobility gives tungsten its excellent electrical conductivity for filaments, while the strong metallic bonding between tungsten atoms (due to many valence electrons) creates an extremely high melting point. Choice A reverses the solid types—WC isn't molecular (which would involve discrete molecules held by weak intermolecular forces), and tungsten isn't a covalent network. Choice C incorrectly classifies WC as metallic; while both materials are hard, their bonding types are fundamentally different. WC's hardness comes from covalent bonds, not metallic bonding. Choice D misidentifies WC as ionic, but tungsten and carbon don't form ionic compounds due to their similar electronegativities and the way they actually bond covalently. The correct answer is B because it properly identifies WC as a covalent network solid (explaining its hardness) and tungsten as a metallic solid (explaining its conductivity and high melting point). Study tip: Remember that material properties directly reflect bonding type—covalent networks give hardness, while metallic bonding provides conductivity and often high melting points.

Question 13

The bulk modulus of a solid measures its resistance to uniform compression. Which type of bonding would be expected to produce the highest bulk modulus?

  1. Ionic bonding with large, singly charged ions
  2. Metallic bonding with loosely held valence electrons
  3. Covalent bonding in a three-dimensional network (correct answer)
  4. Hydrogen bonding between polar molecules
  5. Van der Waals forces between nonpolar molecules
Explanation: When you encounter questions about mechanical properties like bulk modulus, think about how different types of chemical bonding affect a material's ability to resist deformation. Bulk modulus specifically measures resistance to uniform compression - essentially how hard it is to squeeze a material and reduce its volume. Covalent bonding in a three-dimensional network produces the highest bulk modulus because it creates the strongest, most rigid structural framework. In materials like diamond or silicon carbide, every atom is connected to its neighbors through strong covalent bonds extending in all three spatial dimensions. When you try to compress such a material, you're directly working against these strong directional bonds throughout the entire structure. The electrons are localized between specific atom pairs, creating a rigid network that strongly resists any change in interatomic distances. Option A is incorrect because ionic bonding with large, singly charged ions actually produces relatively soft materials. Large ions have their charge spread over a bigger volume, creating weaker electrostatic attractions, and the single charges provide less binding energy than multiply charged systems. Option B is wrong because metallic bonding involves delocalized electrons that can move freely. This electron mobility allows metal atoms to slide past each other more easily, making metals generally more compressible than covalent networks. Option D fails because hydrogen bonds are much weaker than covalent bonds (typically 10-50 times weaker), so materials held together primarily by hydrogen bonding compress much more readily. Remember: when comparing mechanical strength properties, covalent network solids consistently outperform other bonding types due to their strong, directional bonds extending throughout three-dimensional space.

Question 14

A crystalline solid has a high melting point (1850°C), conducts electricity when molten but not when solid, and is soluble in water. The aqueous solution conducts electricity. Which type of solid is this most likely to be?

  1. Ionic solid (correct answer)
  2. Metallic solid
  3. Covalent network solid
  4. Molecular solid
  5. Amorphous solid
Explanation: When you encounter a question about crystalline solids, focus on how their properties relate to their bonding and structure. Each type of solid has a distinctive combination of characteristics that can help you identify it. The key clues here point to an ionic solid. The high melting point (1850°C) indicates strong electrostatic forces between oppositely charged ions. The fact that it conducts electricity when molten but not when solid is the most telling characteristic—ionic solids have ions that are locked in fixed positions in the crystal lattice, preventing electrical conduction. However, when melted, these ions become mobile and can carry electric current. The water solubility and electrical conductivity of the aqueous solution further confirm this, as ionic compounds typically dissociate into ions when dissolved. Let's examine why the other options don't fit. Option B (metallic solid) is wrong because metals conduct electricity in the solid state due to their "sea" of delocalized electrons, which contradicts the given properties. Option C (covalent network solid) is incorrect because these materials are typically insoluble in water and don't conduct electricity in any state, since electrons are localized in covalent bonds. Option D (molecular solid) fails because these solids generally have low melting points due to weak intermolecular forces, not the high melting point described. Remember this pattern: if a solid doesn't conduct electricity when solid but does when molten or dissolved, think ionic. The "on/off" conductivity switch is ionic solids' signature characteristic.

Question 15

A body-centered cubic (BCC) metal has an atomic radius of 1.24 × 10⁻¹⁰ m. In this structure, atoms touch along the body diagonal. What is the edge length of the unit cell?

  1. 2.48 × 10⁻¹⁰ m
  2. 2.86 × 10⁻¹⁰ m (correct answer)
  3. 3.51 × 10⁻¹⁰ m
  4. 4.30 × 10⁻¹⁰ m
  5. 4.96 × 10⁻¹⁰ m
Explanation: When you encounter crystal structure problems, you need to visualize how atoms are arranged and use geometric relationships to connect atomic radius with unit cell dimensions. In a body-centered cubic (BCC) structure, atoms are positioned at each corner of the cube plus one atom at the center. The key insight is that atoms touch along the body diagonal—the line connecting opposite corners through the center of the cube. This body diagonal contains exactly 4 atomic radii: you travel from the surface of a corner atom, through its center, across the central atom (2 radii), through the center of the opposite corner atom, to its surface. The body diagonal of a cube with edge length aa equals a3a\sqrt{3}. Since this diagonal contains 4 atomic radii: a3=4ra\sqrt{3} = 4r Solving for the edge length: a=4r3=4×1.24×10−103=4.96×10−101.732=2.86×10−10a = \frac{4r}{\sqrt{3}} = \frac{4 \times 1.24 \times 10^{-10}}{\sqrt{3}} = \frac{4.96 \times 10^{-10}}{1.732} = 2.86 \times 10^{-10} m Answer B (2.86 × 10⁻¹⁰ m) is correct. Answer A (2.48 × 10⁻¹⁰ m) represents 2r, which would be the edge length if atoms touched along the face diagonal. Answer C (3.51 × 10⁻¹⁰ m) comes from incorrectly using 2r32r\sqrt{3}. Answer D (4.30 × 10⁻¹⁰ m) would result from using a=2r2a = 2r\sqrt{2}, mixing up BCC with other crystal geometries. Remember: In BCC problems, always identify where atoms touch (body diagonal) and count how many radii fit along that distance—this geometric approach works for all crystal systems.

Question 16

The thermal expansion coefficient of a crystal depends on the nature of the bonding. Which type of solid would be expected to have the smallest thermal expansion coefficient?

  1. Ionic solids with highly charged ions
  2. Metallic solids with delocalized bonding
  3. Covalent network solids with directional bonding (correct answer)
  4. Molecular solids with hydrogen bonding
  5. Molecular solids with van der Waals forces
Explanation: When analyzing thermal expansion coefficients, you need to consider how strongly atoms are held in their positions and how rigid the crystal structure is. Stronger, more directional bonding leads to greater resistance to structural changes with temperature. Covalent network solids like diamond, silicon carbide, and quartz have the smallest thermal expansion coefficients because their atoms are connected by strong, directional covalent bonds throughout the entire crystal structure. These bonds create a rigid three-dimensional network where atoms are locked into specific geometric arrangements. When temperature increases, the strong covalent bonds resist stretching, and the directional nature means atoms cannot easily move out of their optimal bonding positions. Option A is incorrect because while ionic solids with highly charged ions do have strong electrostatic attractions, the bonding is non-directional. This allows for more flexibility in the crystal structure compared to covalent networks. Option B is wrong because metallic bonding, though strong, involves delocalized electrons that create a "sea" of bonding. This delocalization actually allows metal atoms more freedom to move and expand with temperature changes. Option D is incorrect because molecular solids are held together by relatively weak intermolecular forces (even hydrogen bonding is much weaker than covalent bonds), making them quite susceptible to thermal expansion. Remember this hierarchy: covalent network solids have the most rigid structures due to strong, directional bonding throughout, making them the most resistant to thermal expansion. When you see thermal expansion questions, always consider both bond strength and directionality.

Question 17

Quartz (SiO₂) and sodium chloride (NaCl) both form crystalline solids, but quartz is much harder and has a higher melting point than NaCl. Which difference in bonding best explains this observation?

  1. Quartz has ionic bonding while NaCl has covalent bonding throughout the structure
  2. Quartz has metallic bonding while NaCl has ionic bonding throughout the structure
  3. Quartz has covalent network bonding while NaCl has ionic bonding throughout the structure (correct answer)
  4. Both have ionic bonding, but quartz has stronger electrostatic interactions
  5. Both have covalent bonding, but quartz has longer bond lengths
Explanation: When you encounter questions comparing the physical properties of different crystalline solids, focus on identifying the types of bonding present in each structure, as this directly determines properties like hardness and melting point. Quartz (SiO₂) forms a covalent network structure where each silicon atom is covalently bonded to four oxygen atoms in a continuous three-dimensional network. These covalent bonds extend throughout the entire crystal, creating an extremely strong and rigid structure. To melt quartz or break it apart, you must break these strong covalent bonds, which requires significant energy. Sodium chloride (NaCl) has ionic bonding, where Na⁺ and Cl⁻ ions are held together by electrostatic attractions. While these ionic interactions are substantial, they're generally weaker than the covalent bonds in network solids and can be disrupted more easily. Choice A incorrectly reverses the bonding types – NaCl is the ionic compound, not quartz. Choice B wrongly suggests quartz has metallic bonding; metallic bonding occurs in metals and alloys, not in compounds like SiO₂. Choice D assumes both compounds are ionic, missing the key distinction that quartz is a covalent network solid. Choice C correctly identifies that quartz has covalent network bonding while NaCl has ionic bonding, explaining why quartz is harder and has a higher melting point. Study tip: Remember the hierarchy of bond strength in solids: covalent network bonds > ionic bonds > metallic bonds > intermolecular forces. When comparing physical properties, always identify the bonding type first.

Question 18

A certain alloy exhibits electrical conductivity that increases with temperature, unlike typical metals. This unusual behavior suggests that the alloy contains which type of structural feature?

  1. Increased grain boundaries that scatter electrons at low temperature
  2. Semiconductor impurities that create additional charge carriers when heated (correct answer)
  3. Ionic regions that become more mobile at higher temperatures
  4. Covalent network regions that break bonds to release electrons when heated
  5. Amorphous regions that crystallize and improve conductivity when heated
Explanation: When you encounter unusual electrical conductivity behavior in materials, think about what controls electron flow and how temperature affects different types of charge carriers. Normal metals have high conductivity that decreases with temperature because thermal vibrations scatter electrons more as temperature rises. However, this alloy shows the opposite pattern - conductivity increases with temperature. This is characteristic of semiconductor behavior, where thermal energy promotes electrons from the valence band to the conduction band, creating more charge carriers. Answer B correctly identifies semiconductor impurities as the cause. These impurities introduce energy levels within the band gap of the host material. As temperature increases, more electrons gain enough thermal energy to jump from these impurity levels into the conduction band, dramatically increasing the number of mobile charge carriers and thus the conductivity. Answer A is incorrect because grain boundaries cause electron scattering at all temperatures, not just low ones, and wouldn't explain increasing conductivity with temperature. Answer C describes ionic conduction, which while temperature-dependent, doesn't typically occur in metal alloys under normal conditions. Answer D suggests covalent bond breaking, but this would require extremely high temperatures and would likely damage the material permanently rather than create reversible temperature-dependent conductivity. Remember: When you see conductivity that increases with temperature in an otherwise metallic material, immediately think "semiconductor doping." This is a classic signature of semiconductor impurities creating thermally activated charge carriers.

Question 19

Cesium chloride (CsCl) has a different crystal structure than sodium chloride (NaCl), even though both are ionic compounds with 1:1 stoichiometry. The structural difference is primarily due to which factor?

  1. CsCl has stronger ionic bonding than NaCl
  2. CsCl has more covalent character than NaCl
  3. The size ratio of Cs⁺/Cl⁻ differs significantly from Na⁺/Cl⁻ (correct answer)
  4. CsCl forms hydrogen bonds while NaCl does not
  5. CsCl has higher charge density than NaCl
Explanation: When you encounter questions about crystal structures of ionic compounds, the key factor governing structure is the size ratio of the cation to anion, which determines how the ions can most efficiently pack together in three-dimensional space. The size ratio of Cs+/Cl−Cs^+/Cl^- is significantly larger than Na+/Cl−Na^+/Cl^- because cesium ion is much larger than sodium ion (both have the same +1 charge, but Cs is in period 6 while Na is in period 3). This larger size ratio means that in CsCl, each cesium ion can accommodate 8 chloride ions around it in a cubic arrangement (called the cesium chloride structure), while in NaCl, each sodium ion can only accommodate 6 chloride ions in an octahedral arrangement (the rock salt structure). Looking at the incorrect options: (A) is wrong because ionic bond strength actually decreases as ion size increases, so CsCl has weaker, not stronger, ionic bonding than NaCl. (B) is incorrect because both compounds are highly ionic with minimal covalent character - the electronegativity differences are similar. (D) is false because neither compound forms hydrogen bonds; both are purely ionic crystals held together by electrostatic attractions. When studying crystal structures, remember that geometry follows from ion size ratios. The radius ratio rule predicts coordination numbers: ratios of 0.732-1.000 favor 8-coordinate structures (like CsCl), while ratios of 0.414-0.732 favor 6-coordinate structures (like NaCl). This principle applies broadly to ionic compound structures.

Question 20

A crystalline solid exhibits the following properties: high electrical conductivity, malleability, and metallic luster. When heated, its electrical conductivity decreases. Which type of bonding model best explains these observations?

  1. Ionic bonding with mobile ions providing conductivity
  2. Covalent bonding with delocalized π electrons
  3. Metallic bonding with a sea of delocalized electrons (correct answer)
  4. Hydrogen bonding between polar molecules
  5. Van der Waals forces between nonpolar molecules
Explanation: When you encounter questions about crystalline solids and their properties, focus on connecting the observed characteristics to the underlying bonding model. The key is recognizing which combination of properties uniquely identifies each bonding type. The described solid shows high electrical conductivity, malleability, metallic luster, and decreasing conductivity when heated. This property combination points directly to metallic bonding with a sea of delocalized electrons (C). In metals, valence electrons are not bound to specific atoms but form an "electron sea" that moves freely throughout the crystal lattice. This explains the high electrical conductivity (mobile electrons carry current), malleability (atoms can slide past each other while maintaining bonding), and metallic luster (free electrons interact with light). The decreased conductivity upon heating occurs because thermal vibrations of metal atoms interfere with electron movement. Option A is incorrect because ionic solids are typically brittle, not malleable, and their conductivity comes from mobile ions, not electrons. Pure ionic crystals don't conduct electricity in solid form. Option B fails because while covalent compounds with delocalized π electrons (like graphite) can conduct electricity, they lack the malleability and metallic luster described. Option D is wrong since hydrogen-bonded substances are generally poor electrical conductors and don't exhibit metallic properties. Remember this pattern: when you see high electrical conductivity combined with malleability and metallic luster in a solid, think metallic bonding. The temperature-dependent conductivity behavior (decreasing with heat for metals, increasing for semiconductors) is often a key distinguishing feature on chemistry exams.