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1. Which of the following is NOT a characteristic property of gases?
ⓐ. High compressibility
ⓑ. Indefinite shape and volume
ⓒ. Strong intermolecular forces
ⓓ. Ability to expand to fill the container
Correct Answer: Strong intermolecular forces
Explanation: Gases are characterized by weak intermolecular forces, allowing them to expand, flow, and be highly compressible. Solids and liquids exhibit strong intermolecular forces. Hence, option C is incorrect, while A, B, and D are true for gases.
2. Which state of matter has definite volume but no definite shape?
ⓐ. Solid
ⓑ. Liquid
ⓒ. Gas
ⓓ. Plasma
Correct Answer: Liquid
Explanation: Liquids have definite volume due to intermolecular attractions but take the shape of their container. Solids have both definite shape and volume, while gases have neither. Plasma is ionized gas-like matter without definite shape or volume.
3. Which postulate of the kinetic molecular theory explains the pressure exerted by gases?
ⓐ. Gas molecules are in continuous random motion.
ⓑ. Collisions between molecules and container walls are elastic.
ⓒ. The volume occupied by molecules is negligible.
ⓓ. Intermolecular forces are negligible.
Correct Answer: Collisions between molecules and container walls are elastic.
Explanation: Pressure arises from molecules colliding elastically with container walls, transferring momentum. Other postulates describe random motion, negligible molecular volume, and weak forces, but pressure is directly explained by elastic collisions.
4. At STP, 1 mole of an ideal gas occupies:
ⓐ. 11.2 L
ⓑ. 21.4 L
ⓒ. 44.8 L
ⓓ. 22.4 L
Correct Answer: 22.4 L
Explanation: According to Avogadro’s law, 1 mole of any ideal gas at standard temperature (273 K) and pressure (1 atm) occupies 22.4 liters. Option A corresponds to half a mole, C corresponds to 2 moles, and B is incorrect.
5. Which gas law relates pressure and volume at constant temperature?
ⓐ. Boyle’s law
ⓑ. Charles’ law
ⓒ. Gay-Lussac’s law
ⓓ. Avogadro’s law
Correct Answer: Boyle’s law
Explanation: Boyle’s law states $P \propto \frac{1}{V}$ at constant temperature. Charles’ law relates volume and temperature, Gay-Lussac’s law relates pressure and temperature, while Avogadro’s law relates volume and moles.
6. Which equation represents the ideal gas law?
ⓐ. $PV = k$
ⓑ. $PV = nRT$
ⓒ. $V/T = k$
ⓓ. $P/T = k$
Correct Answer: $PV = nRT$
Explanation: The ideal gas equation combines Boyle’s, Charles’, and Avogadro’s laws: $PV = nRT$. Here, $P$ is pressure, $V$ is volume, $n$ is moles, $R$ is the gas constant, and $T$ is absolute temperature.
7. Which of the following explains why gases are more compressible than liquids?
ⓐ. Gases have higher density.
ⓑ. Gas molecules occupy negligible volume compared to container volume.
ⓒ. Gas molecules move slower.
ⓓ. Gas molecules have strong intermolecular forces.
Correct Answer: Gas molecules occupy negligible volume compared to container volume.
Explanation: In gases, molecules are far apart, leaving empty spaces, so they can be compressed easily. Liquids and solids have closely packed molecules, hence less compressibility.
8. Which of the following correctly represents SI units of pressure?
ⓐ. atm
ⓑ. bar
ⓒ. N/m²
ⓓ. mmHg
Correct Answer: N/m²
Explanation: The SI unit of pressure is Pascal (Pa), which equals 1 N/m². atm, bar, and mmHg are common units but not SI units.
9. The temperature at which a gas theoretically occupies zero volume is called:
ⓐ. Absolute zero
ⓑ. Triple point
ⓒ. Critical temperature
ⓓ. Normal boiling point
Correct Answer: Absolute zero
Explanation: Absolute zero (0 K or –273.15 °C) is the temperature at which molecular motion ceases, and gas volume extrapolates to zero. Triple point is where solid, liquid, and gas coexist, while critical temperature relates to liquefaction.
10. Which of the following properties is common to both liquids and gases?
ⓐ. Fixed shape
ⓑ. Definite volume
ⓒ. Ability to flow
ⓓ. High density
Correct Answer: Ability to flow
Explanation: Liquids and gases are called fluids because their molecules can move freely, allowing them to flow. Solids have fixed shape and high density. Liquids have definite volume, gases don’t. Hence, flowing is the shared property.
11. Which set of properties best distinguishes solids from liquids and gases?
ⓐ. High compressibility, no definite shape, no definite volume
ⓑ. Definite shape, definite volume, very low compressibility
ⓒ. No definite shape, definite volume, high compressibility
ⓓ. No definite shape, no definite volume, highest kinetic energy
Correct Answer: Definite shape, definite volume, very low compressibility
Explanation: Solids have tightly packed particles with strong intermolecular forces, giving them definite shape and volume and making them least compressible. Liquids have definite volume but no fixed shape, while gases have neither shape nor volume and are highly compressible. Option D describes gases (highest kinetic energy).
12. Which property enables gases to be stored in large amounts in small steel cylinders?
ⓐ. High density
ⓑ. High compressibility
ⓒ. Definite volume
ⓓ. Strong intermolecular forces
Correct Answer: High compressibility
Explanation: Gas particles are far apart, so external pressure can substantially reduce volume, allowing large amounts to be stored. High density (A) actually opposes storage for the same volume, gases do not have definite volume (C), and they have weak—not strong—intermolecular forces (D).
13. The correct order of diffusion rate at room temperature is:
ⓐ. Solids > Liquids > Gases
ⓑ. Solids > Gases > Liquids
ⓒ. Liquids > Gases > Solids
ⓓ. Gases > Liquids > Solids
Correct Answer: Gases > Liquids > Solids
Explanation: Diffusion depends on particle mobility and intermolecular spacing. Gas molecules move fastest and are far apart, so they diffuse most rapidly. Liquids have intermediate mobility; solids have only very slow diffusion via defects. Hence B is correct.
14. Which row correctly matches “shape–volume” for the three states?
ⓐ. Solid: no shape, no volume; Liquid: shape, volume; Gas: shape, no volume
ⓑ. Solid: shape, volume; Liquid: no shape, volume; Gas: no shape, no volume
ⓒ. Solid: shape, no volume; Liquid: shape, volume; Gas: no shape, volume
ⓓ. Solid: no shape, volume; Liquid: shape, no volume; Gas: shape, no volume
Correct Answer: Solid: shape, volume; Liquid: no shape, volume; Gas: no shape, no volume
Explanation: Solids possess both definite shape and volume. Liquids take the container’s shape but keep volume. Gases neither retain shape nor volume, expanding to fill any container.
15. Which statement about densities is correct?
ⓐ. All solids are denser than their liquids without exception.
ⓑ. Gases have densities comparable to liquids at room temperature.
ⓒ. Water has maximum density at $4^\circ\text{C}$; ice floats because it is less dense than liquid water.
ⓓ. Liquids are always less dense than gases.
Correct Answer: Water has maximum density at $4^\circ\text{C}$; ice floats because it is less dense than liquid water.
Explanation: Generally, $\rho_{\text{solid}} > \rho_{\text{liquid}} \gg \rho_{\text{gas}}$. Water is an important exception: due to hydrogen bonding, ice is less dense than liquid water; hence it floats. Gases have far smaller densities than liquids (B false). Statements A and D are absolute claims and incorrect.
16. A fixed mass of gas expands from $1.0\,\text{L}$ to $3.0\,\text{L}$ at the same temperature. The ratio of final to initial density is:
ⓐ. $3:1$
ⓑ. $1:3$
ⓒ. $1:2$
ⓓ. $2:1$
Correct Answer: $1:3$
Explanation: Density $\rho = \dfrac{m}{V}$. For a fixed mass, $\rho \propto \dfrac{1}{V}$. Increasing volume threefold reduces density to one-third. So $\rho_f/\rho_i = 1/3$. Options A and D invert the relation; C corresponds to doubling volume, not tripling.
17. Which state has the largest bulk modulus $B$ (thus the least compressible)?
ⓐ. Gas
ⓑ. Liquid
ⓒ. Solid
ⓓ. All equal
Correct Answer: Solid
Explanation: Bulk modulus $B = -V\,\dfrac{dP}{dV}$ quantifies resistance to compression. Strongest cohesion and closest packing in solids give the highest $B$. Liquids have moderate $B$; gases have the smallest $B$ (high compressibility). Hence solids are least compressible.
18. The coefficient of volume expansion $(\beta)$ is generally in the order:
ⓐ. Solids > Liquids > Gases
ⓑ. Gases > Liquids > Solids
ⓒ. Liquids > Gases > Solids
ⓓ. Solids > Gases > Liquids
Correct Answer: Gases > Liquids > Solids
Explanation: Expansion reflects how strongly volume changes with temperature. Weak intermolecular forces and large free space in gases cause large expansion on heating. Liquids expand less; solids expand the least because of rigid structures and strong cohesive forces.
19. Which statement best describes the internal arrangement (“order”) of particles?
ⓐ. Solids: short-range order; Liquids: long-range order; Gases: long-range order
ⓑ. Solids: long-range order; Liquids: short-range order; Gases: no regular order
ⓒ. Solids: no order; Liquids: long-range order; Gases: short-range order
ⓓ. All three have long-range order
Correct Answer: Solids: long-range order; Liquids: short-range order; Gases: no regular order
Explanation: Crystalline solids show long-range periodicity. Liquids possess only short-range order (neighbors are transient). Gases have random distribution with negligible ordering due to large separations and continuous motion.
20. In which state is surface tension a characteristic, macroscopically measurable property of the bulk phase?
ⓐ. Nothing
ⓑ. Solid
ⓒ. Gases
ⓓ. Liquids
Correct Answer: Liquids
Explanation: Surface tension arises from cohesive forces at a liquid–air interface and manifests in phenomena like droplets, capillarity, and insects walking on water. Gases have negligible cohesive forces in bulk; solids possess surface energy but not fluid surface tension in the same sense. Thus liquids uniquely exhibit measurable surface tension as a bulk property.
21. Which of the following terms refers to the direct conversion of a solid into vapor without passing through the liquid state?
ⓐ. Condensation
ⓑ. Sublimation
ⓒ. Vaporization
ⓓ. Fusion
Correct Answer: Sublimation
Explanation: Sublimation is the process in which a solid directly changes to vapor, e.g., dry ice ($CO_2$) and iodine crystals. Condensation is gas to liquid, vaporization is liquid to gas, and fusion is solid to liquid.
22. Which process explains the formation of dew drops on grass in the early morning?
ⓐ. Evaporation
ⓑ. Sublimation
ⓒ. Freezing
ⓓ. Condensation
Correct Answer: Condensation
Explanation: Water vapor in the air loses heat at low temperatures overnight and condenses into liquid droplets on cool grass. Evaporation is liquid to vapor, freezing is liquid to solid, and sublimation is solid to vapor.
23. The heat absorbed during the conversion of 1 mole of ice at $0^\circ C$ to water at $0^\circ C$ is called:
ⓐ. Heat of vaporization
ⓑ. Heat of fusion
ⓒ. Heat of sublimation
ⓓ. Heat capacity
Correct Answer: Heat of fusion
Explanation: Heat of fusion is the energy required to overcome intermolecular forces during melting at constant temperature. Heat of vaporization is for liquid to gas, sublimation is solid to gas, and heat capacity measures heat needed per degree rise.
24. Which factor mainly affects the rate of evaporation of a liquid?
ⓐ. Intermolecular forces, surface area, and temperature
ⓑ. Only molecular mass of liquid
ⓒ. Density of liquid only
ⓓ. Color of the container
Correct Answer: Intermolecular forces, surface area, and temperature
Explanation: Weak intermolecular forces, larger surface area, and higher temperature all increase evaporation rate. Molecular mass and density indirectly affect but are not sole determinants. Container color is irrelevant.
25. When pressure on a solid increases, its melting point usually:
ⓐ. Increases
ⓑ. Decreases
ⓒ. Remains constant
ⓓ. Becomes zero
Correct Answer: Increases
Explanation: Higher pressure compresses solid molecules, making it harder for them to move apart, so higher temperature is needed to melt. An exception is ice, where hydrogen bonding causes the melting point to decrease under pressure.
26. The direct conversion of water vapor into ice without becoming liquid is known as:
ⓐ. Fusion
ⓑ. Deposition
ⓒ. Vaporization
ⓓ. Sublimation
Correct Answer: Deposition
Explanation: Deposition is the reverse of sublimation: gas directly becomes solid. Frost formation on windows in winter is a common example. Fusion is solid to liquid, vaporization is liquid to gas, and sublimation is solid to gas.
27. Which statement about boiling and evaporation is correct?
ⓐ. Both occur only at a fixed temperature.
ⓑ. Boiling occurs throughout the liquid, while evaporation occurs only at the surface.
ⓒ. Evaporation requires external heating, but boiling does not.
ⓓ. Evaporation always occurs at higher temperature than boiling.
Correct Answer: Boiling occurs throughout the liquid, while evaporation occurs only at the surface.
Explanation: Boiling is bulk vaporization at boiling temperature, requiring vapor pressure to equal external pressure. Evaporation is a surface phenomenon occurring at all temperatures.
28. Which of the following processes is exothermic?
ⓐ. Fusion
ⓑ. Vaporization
ⓒ. Condensation
ⓓ. Sublimation
Correct Answer: Condensation
Explanation: Condensation releases latent heat as gas molecules lose energy and form liquid. Fusion, vaporization, and sublimation are endothermic as they require heat absorption.
29. Which of the following correctly represents enthalpy of sublimation ($\Delta H_{sub}$)?
ⓐ. $\Delta H_{sub} = \Delta H_{fusion} + \Delta H_{vaporization}$
ⓑ. $\Delta H_{sub} = \Delta H_{fusion} – \Delta H_{vaporization}$
ⓒ. $\Delta H_{sub} = \Delta H_{fusion} \times \Delta H_{vaporization}$
ⓓ. $\Delta H_{sub} = \dfrac{\Delta H_{fusion}}{\Delta H_{vaporization}}$
Correct Answer: $\Delta H_{sub} = \Delta H_{fusion} + \Delta H_{vaporization}$
Explanation: Sublimation involves two steps: solid to liquid (fusion) and liquid to vapor (vaporization). Thus, enthalpy of sublimation is the sum of enthalpies of fusion and vaporization.
30. Which change of state explains the working of solid air fresheners (naphthalene balls)?
ⓐ. Deposition
ⓑ. Sublimation
ⓒ. Fusion
ⓓ. Vaporization
Correct Answer: Sublimation
Explanation: Naphthalene balls directly convert from solid to gas at room temperature, slowly releasing fragrance. This sublimation process is responsible for their disappearance without forming liquid.
31. Which of the following molecules exhibits dipole–dipole interactions?
ⓐ. $HCl$
ⓑ. $Cl_2$
ⓒ. $O_2$
ⓓ. $CH_4$
Correct Answer: $HCl$
Explanation: Dipole–dipole forces arise between polar molecules with permanent dipole moments. $HCl$ is polar due to electronegativity difference between H and Cl. $Cl_2, O_2,$ and $CH_4$ are nonpolar and show only London dispersion forces.
32. Which intermolecular force plays the major role in dissolving NaCl in water?
ⓐ. Dipole–dipole
ⓑ. Ion–dipole
ⓒ. London dispersion forces
ⓓ. Covalent bonding
Correct Answer: Ion–dipole
Explanation: In NaCl solution, water (polar) molecules surround $Na^+$ and $Cl^-$ ions. The attraction between ion and dipole (polar water) stabilizes ions in solution. Dipole–dipole forces act between neutral molecules, not ions. Dispersion forces are weak compared to ion–dipole interactions.
33. Which of the following molecules is expected to exhibit only London dispersion forces?
ⓐ. $H_2O$
ⓑ. $HCl$
ⓒ. $CO_2$
ⓓ. $NH_3$
Correct Answer: $CO_2$
Explanation: $CO_2$ is a linear, nonpolar molecule. It lacks permanent dipole and hydrogen bonding, so intermolecular attraction is only due to temporary fluctuations in electron density—London dispersion forces. $H_2O$ and $NH_3$ exhibit hydrogen bonding; $HCl$ has dipole–dipole interactions.
34. London dispersion forces arise due to:
ⓐ. Attraction between permanent dipoles
ⓑ. Attraction between ions and polar molecules
ⓒ. Temporary fluctuations in electron density creating instantaneous dipoles
ⓓ. Sharing of electrons between atoms
Correct Answer: Temporary fluctuations in electron density creating instantaneous dipoles
Explanation: Dispersion forces occur when electron clouds momentarily shift, inducing dipoles in neighboring molecules. These weak forces are universal in all molecules, stronger in heavier atoms with more electrons.
35. Which intermolecular force explains the higher boiling point of HF compared to HCl?
ⓐ. Dipole–dipole interaction
ⓑ. Ion–dipole interaction
ⓒ. London dispersion force
ⓓ. Hydrogen bonding
Correct Answer: Hydrogen bonding
Explanation: HF forms strong hydrogen bonds due to high electronegativity of fluorine, leading to higher boiling point. HCl lacks significant hydrogen bonding; its attractions are mainly dipole–dipole. Dispersion forces exist but are weaker than H-bonds.
36. Which force is dominant in the interaction between $Na^+$ ion and $H_2O$ molecule?
ⓐ. Dipole–dipole
ⓑ. Ion–dipole
ⓒ. London dispersion
ⓓ. Ion–ion
Correct Answer: Ion–dipole
Explanation: $Na^+$ attracts the negative pole (oxygen) of $H_2O$. This is ion–dipole interaction, crucial in solvation. Ion–ion exists in ionic solids like NaCl, dipole–dipole acts between neutral polar molecules, and dispersion is much weaker.
37. Which molecule will show stronger London dispersion forces?
ⓐ. $He$
ⓑ. $Ne$
ⓒ. $Ar$
ⓓ. $Kr$
Correct Answer: $Kr$
Explanation: Dispersion forces increase with number of electrons and molecular size. Krypton has the largest electron cloud among the given noble gases, making its dispersion forces the strongest. Helium has the weakest.
38. What causes dipole–dipole interaction strength to increase?
ⓐ. Higher temperature
ⓑ. Greater molecular mass
ⓒ. Larger dipole moment
ⓓ. Smaller number of electrons
Correct Answer: Larger dipole moment
Explanation: Stronger permanent dipoles create stronger attractions between molecules. Molecular mass increases dispersion forces, not dipole–dipole directly. Higher temperature disrupts interactions. Smaller electron count reduces polarizability, weakening dispersion forces.
39. Which intermolecular force is primarily responsible for the liquefaction of nonpolar gases like $O_2$ and $N_2$?
ⓐ. Dipole–dipole
ⓑ. Ion–dipole
ⓒ. London dispersion
ⓓ. Hydrogen bonding
Correct Answer: London dispersion
Explanation: Nonpolar gases like $O_2$ and $N_2$ rely solely on dispersion forces to condense at low temperatures. They have no permanent dipoles or hydrogen bonding. These temporary induced dipole interactions enable liquefaction.
40. Which of the following order of strength is correct for intermolecular forces?
ⓐ. London dispersion < Dipole–dipole < Ion–dipole
ⓑ. Ion–dipole < Dipole–dipole < London dispersion
ⓒ. Dipole–dipole < London dispersion < Ion–dipole
ⓓ. Dipole–dipole < Ion–dipole < London dispersion
Correct Answer: London dispersion < Dipole–dipole < Ion–dipole
Explanation: Ion–dipole interactions are strongest because ions have full charges. Dipole–dipole are weaker than ion–dipole but stronger than temporary dispersion forces. Dispersion forces are the weakest, though significant in large atoms/molecules.
41. Which type of intermolecular force is most significant in liquid $NH_3$?
ⓐ. London dispersion forces
ⓑ. Dipole–dipole interactions
ⓒ. Ion–dipole interactions
ⓓ. Hydrogen bonding
Correct Answer: Hydrogen bonding
Explanation: In $NH_3$, strong hydrogen bonding arises between highly electronegative nitrogen and hydrogen atoms. While dispersion and dipole–dipole forces are also present, hydrogen bonding dominates and explains higher boiling point than expected for its molar mass.
42. Which type of intermolecular force exists in all substances, regardless of polarity?
ⓐ. Dipole–dipole forces
ⓑ. Ion–dipole forces
ⓒ. London dispersion forces
ⓓ. Hydrogen bonding
Correct Answer: London dispersion forces
Explanation: Dispersion forces originate from temporary instantaneous dipoles due to electron cloud movement, so they exist in both polar and nonpolar molecules. Dipole–dipole and ion–dipole require polarity or ions, and hydrogen bonding requires specific atoms (N, O, F).
43. Which interaction is responsible for the solubility of ionic compounds in polar solvents?
ⓐ. London dispersion
ⓑ. Ion–dipole
ⓒ. Dipole–dipole
ⓓ. Covalent bonding
Correct Answer: Ion–dipole
Explanation: Polar solvents like water stabilize ions through strong ion–dipole interactions, aligning the solvent’s partial charges around ions. Dipole–dipole alone cannot overcome ionic lattice energy. Dispersion forces are too weak to cause ionic dissolution.
44. Which intermolecular force is the weakest?
ⓐ. Dipole–dipole interaction
ⓑ. Ion–dipole interaction
ⓒ. London dispersion force
ⓓ. Hydrogen bonding
Correct Answer: London dispersion force
Explanation: Dispersion forces are the weakest because they arise from temporary fluctuations. Dipole–dipole and hydrogen bonds are stronger due to permanent dipoles and specific N–H, O–H, F–H bonds. Ion–dipole is stronger still, involving full ionic charges.
45. Which compound has the highest boiling point due to intermolecular forces?
ⓐ. $CH_4$
ⓑ. $H_2S$
ⓒ. $H_2O$
ⓓ. $CO_2$
Correct Answer: $H_2O$
Explanation: Water has strong hydrogen bonds, requiring significant energy to break, resulting in a much higher boiling point than predicted by molecular mass. $CH_4$ and $CO_2$ are nonpolar (only dispersion). $H_2S$ is polar but lacks strong hydrogen bonding.
46. Which type of force is mainly responsible for the high solubility of $Na^+$ and $Cl^-$ in water?
ⓐ. Dipole–dipole interaction
ⓑ. Ion–dipole interaction
ⓒ. London dispersion force
ⓓ. Induced dipole interaction
Correct Answer: Ion–dipole interaction
Explanation: Water molecules surround ions, with oxygen facing $Na^+$ and hydrogen facing $Cl^-$. These strong ion–dipole attractions overcome lattice forces, making ionic salts soluble. Dispersion and induced dipoles are too weak in comparison.
47. Which of the following molecules will have the strongest dipole–dipole interactions?
ⓐ. $CH_3Cl$
ⓑ. $CH_4$
ⓒ. $Cl_2$
ⓓ. $CO_2$
Correct Answer: $CH_3Cl$
Explanation: $CH_3Cl$ is polar due to electronegativity difference and asymmetric structure, resulting in dipole–dipole attractions. $CH_4$ and $Cl_2$ are nonpolar. $CO_2$ is linear and nonpolar overall, despite polar bonds.
48. Which statement is true regarding London dispersion forces?
ⓐ. They are independent of molecular mass.
ⓑ. They increase with greater electron cloud size.
ⓒ. They only exist in ionic compounds.
ⓓ. They are stronger than ion–dipole forces.
Correct Answer: They increase with greater electron cloud size.
Explanation: Larger molecules with more electrons are more polarizable, leading to stronger dispersion forces. These forces are universal but weakest compared to dipole or ion–dipole. Options A, C, and D are false.
49. What type of intermolecular force dominates in liquid $HCl$?
ⓐ. Hydrogen bonding
ⓑ. Dipole–dipole
ⓒ. Ion–dipole
ⓓ. London dispersion
Correct Answer: Dipole–dipole
Explanation: $HCl$ is polar due to electronegativity difference between H and Cl. Molecules align via permanent dipole–dipole attractions. It lacks hydrogen bonding (since H is not bonded to N, O, or F). Dispersion also acts but is weaker.
50. Which gas can be liquefied more easily due to stronger London dispersion forces?
ⓐ. Helium
ⓑ. Neon
ⓒ. Argon
ⓓ. Xenon
Correct Answer: Xenon
Explanation: Xenon has the largest atomic size and electron count among noble gases, making it highly polarizable. Stronger dispersion forces allow it to liquefy more readily than lighter noble gases like He or Ne.
51. Which of the following correctly arranges intermolecular forces in increasing order of strength?
ⓐ. Ion–dipole < Dipole–dipole < Hydrogen bond < London dispersion
ⓑ. London dispersion < Dipole–dipole < Hydrogen bond < Ion–dipole
ⓒ. Hydrogen bond < Dipole–dipole < London dispersion < Ion–dipole
ⓓ. London dispersion < Ion–dipole < Dipole–dipole < Hydrogen bond
Correct Answer: London dispersion < Dipole–dipole < Hydrogen bond < Ion–dipole
Explanation: Dispersion forces are weakest, dipole–dipole are stronger due to permanent polarity, hydrogen bonds are special strong dipole interactions, and ion–dipole is strongest because ions have full charges interacting with polar molecules.
52. Why is the boiling point of HF higher than HCl, though HCl has greater molar mass?
ⓐ. Stronger ion–dipole forces in HF
ⓑ. Stronger hydrogen bonding in HF
ⓒ. Stronger London dispersion in HF
ⓓ. Stronger dipole–dipole in HF
Correct Answer: Stronger hydrogen bonding in HF
Explanation: HF molecules form strong hydrogen bonds due to high electronegativity of fluorine. These bonds require extra energy to break, raising the boiling point. HCl has larger molar mass but lacks significant hydrogen bonding.
53. Which intermolecular force is strongest among the following interactions?
ⓐ. Between two nonpolar $O_2$ molecules
ⓑ. Between HCl molecules
ⓒ. Between water molecules
ⓓ. Between $Na^+$ and $H_2O$
Correct Answer: Between $Na^+$ and $H_2O$
Explanation: The $Na^+$ and $H_2O$ interaction is ion–dipole, strongest among given choices. Water molecules have hydrogen bonding (strong but weaker than ion–dipole). HCl shows dipole–dipole, and $O_2$ relies only on weak dispersion forces.
54. The reason noble gases show an increasing trend in boiling point down the group is:
ⓐ. Increase in dipole–dipole interactions
ⓑ. Increase in ion–dipole forces
ⓒ. Increase in hydrogen bonding
ⓓ. Increase in London dispersion forces
Correct Answer: Increase in London dispersion forces
Explanation: Larger noble gases like Xe have more electrons and greater polarizability, so dispersion forces strengthen, raising boiling points. Noble gases are nonpolar, so dipole, ion–dipole, or hydrogen bonding do not apply.
55. Which of the following liquids shows the strongest intermolecular forces?
ⓐ. $CCl_4$
ⓑ. $CH_3Cl$
ⓒ. $NH_3$
ⓓ. $Ar$
Correct Answer: $NH_3$
Explanation: $NH_3$ exhibits hydrogen bonding, stronger than the dispersion forces in $CCl_4$ and $Ar$, or dipole–dipole in $CH_3Cl$. This results in higher boiling point and stronger cohesive properties.
56. Order the intermolecular forces from weakest to strongest:
ⓐ. Hydrogen bond < London dispersion < Dipole–dipole < Ion–dipole
ⓑ. London dispersion < Dipole–dipole < Ion–dipole < Hydrogen bond
ⓒ. London dispersion < Dipole–dipole < Hydrogen bond < Ion–dipole
ⓓ. Ion–dipole < Hydrogen bond < Dipole–dipole < London dispersion
Correct Answer: London dispersion < Dipole–dipole < Hydrogen bond < Ion–dipole
Explanation: Dispersion is weakest, dipole–dipole stronger, hydrogen bond is special strong dipole–dipole, and ion–dipole is the strongest due to full ionic charges interacting with permanent dipoles.
57. Which interaction explains the high solubility of salts in water?
ⓐ. London dispersion
ⓑ. Dipole–dipole
ⓒ. Ion–dipole
ⓓ. Hydrogen bonding
Correct Answer: Ion–dipole
Explanation: Water molecules surround ions with strong ion–dipole forces, stabilizing ions in solution. Hydrogen bonding is important in pure water, dipole–dipole acts in neutral polar molecules, and dispersion is weaker.
58. Which has stronger intermolecular forces: $H_2O$ or $H_2S$?
ⓐ. $H_2O$, due to hydrogen bonding
ⓑ. $H_2S$, due to greater molar mass
ⓒ. Both equal
ⓓ. Cannot be determined
Correct Answer: $H_2O$, due to hydrogen bonding
Explanation: Although $H_2S$ is heavier, water has strong hydrogen bonds, giving it much higher boiling point and stronger intermolecular forces than $H_2S$.
59. Which intermolecular force determines the liquefaction ease of $CO_2$?
ⓐ. Dipole–dipole
ⓑ. London dispersion
ⓒ. Ion–dipole
ⓓ. Hydrogen bonding
Correct Answer: London dispersion
Explanation: $CO_2$ is a linear nonpolar molecule, so its liquefaction depends only on dispersion forces, which are relatively strong for a molecule of its size. It lacks dipole or hydrogen bonding.
60. Why is ion–dipole force stronger than dipole–dipole force?
ⓐ. Because ions carry complete charges
ⓑ. Because ions are lighter than molecules
ⓒ. Because dipole–dipole acts only in solids
ⓓ. Because dispersion forces cancel them
Correct Answer: Because ions carry complete charges
Explanation: In ion–dipole interaction, ions with full charges interact strongly with permanent dipoles, much stronger than the partial charges in dipole–dipole. Molecular weight or phase does not change this principle.
61. Thermal energy of a system is defined as:
ⓐ. The total potential energy of molecules only
ⓑ. The total kinetic energy of random motion of molecules
ⓒ. The ordered mechanical energy of the system
ⓓ. The energy stored in chemical bonds only
Correct Answer: The total kinetic energy of random motion of molecules
Explanation: Thermal energy is the collective kinetic energy of molecules in random translational, rotational, and vibrational motion. It is not the ordered energy (C) or potential energy alone (A). Chemical bond energy (D) is chemical potential energy, not thermal energy.
62. Which of the following factors directly determines the thermal energy of a substance?
ⓐ. Volume of the container
ⓑ. Temperature of the substance
ⓒ. Color of the container
ⓓ. Pressure applied externally
Correct Answer: Temperature of the substance
Explanation: Thermal energy depends on the average kinetic energy of molecules, which is proportional to temperature. Volume and pressure affect macroscopic properties, while container color has no role in microscopic molecular energy.
63. At absolute zero ($0 \, K$), the thermal energy of an ideal gas is:
ⓐ. Maximum
ⓑ. Zero
ⓒ. Infinite
ⓓ. Constant but nonzero
Correct Answer: Zero
Explanation: At $0 \, K$, all molecular motion ceases, so average kinetic energy (thermal energy) becomes zero. This is the theoretical limit of temperature, where thermal energy vanishes.
64. Which of the following best describes thermal energy compared to heat?
ⓐ. Thermal energy is the same as heat.
ⓑ. Thermal energy is energy in transit, heat is stored energy.
ⓒ. Thermal energy is internal kinetic energy of molecules, while heat is energy transfer due to temperature difference.
ⓓ. Thermal energy and heat are unrelated.
Correct Answer: Thermal energy is internal kinetic energy of molecules, while heat is energy transfer due to temperature difference.
Explanation: Thermal energy is a state property, while heat is a process of energy transfer between systems at different temperatures. This distinction is important in thermodynamics.
65. Which of the following will have greater thermal energy at the same temperature?
ⓐ. 1 mole of oxygen gas at 300 K
ⓑ. 2 moles of oxygen gas at 300 K
ⓒ. Both have equal thermal energy
ⓓ. Cannot be compared
Correct Answer: 2 moles of oxygen gas at 300 K
Explanation: Thermal energy depends on both temperature and number of molecules (moles). Since both are at the same temperature, the system with more moles (2 moles) has greater total thermal energy.
66. The average translational kinetic energy of one molecule of an ideal gas is given by:
ⓐ. $\dfrac{1}{2} k_B T$
ⓑ. $\dfrac{3}{2} k_B T$
ⓒ. $RT$
ⓓ. $\dfrac{3}{2} RT$
Correct Answer: $\dfrac{3}{2} k_B T$
Explanation: For an ideal gas molecule, thermal energy comes from translational motion in 3 degrees of freedom. Each degree contributes $\dfrac{1}{2} k_B T$, so total translational KE = $\dfrac{3}{2} k_B T$.
67. Which of the following is an example of conversion of thermal energy to mechanical energy?
ⓐ. Melting of ice
ⓑ. Steam engine
ⓒ. Evaporation of water
ⓓ. Diffusion of gas
Correct Answer: Steam engine
Explanation: In a steam engine, heat energy (thermal) from burning fuel produces steam pressure, which does mechanical work on pistons. Melting and evaporation are phase changes involving absorption of heat but not mechanical work.
68. Which is true regarding thermal energy in liquids compared to solids?
ⓐ. Molecules in solids have higher thermal energy because they vibrate more
ⓑ. Molecules in liquids have higher thermal energy due to weaker intermolecular forces and greater freedom of motion
ⓒ. Both have equal thermal energy at the same mass
ⓓ. Liquids have zero thermal energy
Correct Answer: Molecules in liquids have higher thermal energy due to weaker intermolecular forces and greater freedom of motion
Explanation: Liquids have more freedom to move than solids, so at the same temperature, the average molecular kinetic energy is greater in liquids. Solids are more restricted to vibrations about fixed positions.
69. Which of the following is measured by a thermometer?
ⓐ. Heat content
ⓑ. Thermal energy directly
ⓒ. Temperature (related to average kinetic energy per particle)
ⓓ. Potential energy of molecules
Correct Answer: Temperature (related to average kinetic energy per particle)
Explanation: A thermometer measures temperature, which reflects the average kinetic energy of molecules. Thermal energy is proportional to temperature times number of molecules but is not measured directly.
70. Which condition leads to higher total thermal energy?
ⓐ. A small hot cup of tea at 80°C
ⓑ. A large bucket of water at 40°C
ⓒ. Both equal since they are hot
ⓓ. Cannot be compared without mass information
Correct Answer: A large bucket of water at 40°C
Explanation: Though the tea has higher temperature, the bucket has a much larger number of molecules. Total thermal energy depends on both temperature and amount of substance. Hence, a bucket at lower temperature may contain more total thermal energy.
71. As temperature of a gas increases at constant volume, what happens to the molecular motion?
ⓐ. Molecules slow down
ⓑ. Molecules stop moving
ⓒ. Molecules move faster with higher average kinetic energy
ⓓ. Molecular motion remains unchanged
Correct Answer: Molecules move faster with higher average kinetic energy
Explanation: Temperature is directly proportional to the average kinetic energy of gas molecules. Increasing temperature increases molecular speed and motion. Molecules never stop moving unless at absolute zero.
72. What happens to the motion of particles when a solid melts into a liquid?
ⓐ. Motion becomes completely random and unbounded
ⓑ. Vibrational motion decreases to zero
ⓒ. Vibrational motion becomes more vigorous and particles gain freedom to slide past each other
ⓓ. Particles stop moving
Correct Answer: Vibrational motion becomes more vigorous and particles gain freedom to slide past each other
Explanation: In solids, molecules vibrate in fixed positions. On melting, thermal energy increases vibrations until particles overcome some intermolecular forces and slide over one another, becoming a liquid.
73. In liquids, compared to solids, molecules exhibit:
ⓐ. Only vibrational motion
ⓑ. Translational and rotational motions in addition to vibration
ⓒ. No motion at all
ⓓ. Purely ordered motion
Correct Answer: Translational and rotational motions in addition to vibration
Explanation: Liquids have higher energy than solids. Molecules not only vibrate but also rotate and translate, giving them fluidity while still experiencing intermolecular forces.
74. Which of the following states of matter has the highest degree of random molecular motion?
ⓐ. Solid
ⓑ. Liquid
ⓒ. Gas
ⓓ. Plasma
Correct Answer: Plasma
Explanation: Plasma consists of highly energetic ions and electrons moving randomly with very high speeds. Among common states, gases already have the most random motion, but plasma surpasses them due to ionization and higher energy.
75. As pressure on a gas increases at constant temperature, what happens to molecular motion?
ⓐ. Average kinetic energy increases
ⓑ. Molecules move slower due to reduced volume
ⓒ. Frequency of collisions increases though average kinetic energy remains constant
ⓓ. Motion ceases entirely
Correct Answer: Frequency of collisions increases though average kinetic energy remains constant
Explanation: At constant temperature, kinetic energy and molecular speeds stay constant. Increasing pressure compresses the gas, decreasing intermolecular distances, thus increasing collision frequency.
76. What is the effect of decreasing temperature on molecular motion in a liquid?
ⓐ. Molecules vibrate more vigorously
ⓑ. Molecular motion decreases, leading to reduced fluidity
ⓒ. Molecular motion remains unchanged
ⓓ. Molecules stop moving instantly
Correct Answer: Molecular motion decreases, leading to reduced fluidity
Explanation: Cooling lowers average kinetic energy. Molecules move less freely, viscosity increases, and eventually, at freezing point, motion becomes restricted to vibrations in a solid lattice.
77. At absolute zero, molecular motion is:
ⓐ. Random but minimal
ⓑ. Completely stopped theoretically
ⓒ. Maximum due to compression
ⓓ. Unchanged compared to room temperature
Correct Answer: Completely stopped theoretically
Explanation: Absolute zero ($0\,K$) is the theoretical point where molecules possess no kinetic energy and motion ceases. Practically unattainable, but serves as a limit in thermodynamics.
78. Which of the following best explains Brownian motion?
ⓐ. Random motion of molecules in solids due to strong bonds
ⓑ. Continuous zig-zag motion of suspended particles caused by collisions with fluid molecules
ⓒ. Systematic vibration of liquid molecules
ⓓ. Ordered motion of gas molecules
Correct Answer: Continuous zig-zag motion of suspended particles caused by collisions with fluid molecules
Explanation: Brownian motion demonstrates the kinetic molecular theory: invisible molecules in fluids collide with suspended particles, making them move randomly.
79. How does increased molecular motion affect vapor pressure of a liquid?
ⓐ. Vapor pressure decreases
ⓑ. Vapor pressure increases
ⓒ. Vapor pressure remains constant
ⓓ. Vapor pressure becomes zero
Correct Answer: Vapor pressure increases
Explanation: As molecular motion increases with temperature, more molecules escape the liquid surface into vapor phase, raising vapor pressure.
80. Which molecular motion dominates in solids?
ⓐ. Translational
ⓑ. Rotational
ⓒ. Vibrational
ⓓ. Random free motion
Correct Answer: Vibrational
Explanation: In solids, strong intermolecular forces keep particles fixed in a lattice. They cannot translate or rotate freely, so only vibrational motion about mean positions is possible.
81. The physical state of a substance at a given temperature depends mainly on the balance between:
ⓐ. Atomic mass and density
ⓑ. Intermolecular forces and thermal energy
ⓒ. Pressure and molar mass
ⓓ. Color and crystalline structure
Correct Answer: Intermolecular forces and thermal energy
Explanation: The state of matter arises from competition between cohesive intermolecular forces (which hold particles together) and disruptive thermal energy (which makes them move apart). If intermolecular forces dominate, a solid forms; if thermal energy dominates, a gas forms; if both are comparable, the liquid state results. Options A, C, and D may influence properties but do not directly decide the state.
82. Why do solids have fixed shape compared to liquids and gases?
ⓐ. Because solids have no intermolecular forces
ⓑ. Because solids have dominant intermolecular forces over thermal motion
ⓒ. Because solids have higher thermal energy
ⓓ. Because solids contain gaseous molecules
Correct Answer: Because solids have dominant intermolecular forces over thermal motion
Explanation: In solids, strong intermolecular forces keep particles in a fixed lattice. Vibrations are allowed, but translational freedom is absent, leading to a definite shape and volume. Liquids and gases have weaker intermolecular forces relative to thermal energy, allowing particles to move freely.
83. Why does water exist as a liquid at room temperature while $H_2S$ exists as a gas?
ⓐ. Water is lighter, so it condenses more easily
ⓑ. Water forms hydrogen bonds, making intermolecular forces much stronger than thermal energy
ⓒ. $H_2S$ has stronger dipole–dipole interactions than water
ⓓ. $H_2S$ has higher molar mass, so it must be gaseous
Correct Answer: Water forms hydrogen bonds, making intermolecular forces much stronger than thermal energy
Explanation: Despite lower molar mass, water remains liquid because hydrogen bonding between molecules overcomes thermal motion. In contrast, $H_2S$ molecules are held only by weak dipole–dipole and dispersion forces, which are easily disrupted by thermal energy at room temperature, so it stays gaseous.
84. Which of the following best explains why gases are compressible but solids are not?
ⓐ. Solids have higher molar masses
ⓑ. Intermolecular forces dominate in solids, while in gases thermal energy dominates, leaving large empty spaces
ⓒ. Solids are colorless, gases are not
ⓓ. Solids contain no kinetic energy
Correct Answer: Intermolecular forces dominate in solids, while in gases thermal energy dominates, leaving large empty spaces
Explanation: In solids, molecules are tightly bound by strong cohesive forces, leaving no significant empty space, making compression nearly impossible. In gases, high thermal motion keeps molecules far apart, leaving large compressible spaces. Thus, option B correctly describes the interplay.
85. What happens to a liquid when intermolecular forces and thermal energy are nearly equal?
ⓐ. It instantly freezes into a solid
ⓑ. It instantly vaporizes into gas
ⓒ. It maintains fluidity but with definite volume
ⓓ. It becomes unobservable
Correct Answer: It maintains fluidity but with definite volume
Explanation: Liquids represent the intermediate state where intermolecular forces are strong enough to give definite volume, but thermal energy allows molecules to flow past each other. This balance makes liquids fluid yet incompressible. Neither freezing (A) nor vaporization (B) happens unless the balance shifts further.
86. Which factor explains the variation in boiling points among different substances?
ⓐ. Density of the liquid
ⓑ. Balance of intermolecular forces and thermal energy
ⓒ. Size of the container
ⓓ. Pressure of light falling on molecules
Correct Answer: Balance of intermolecular forces and thermal energy
Explanation: The boiling point is reached when thermal energy is sufficient to overcome intermolecular forces completely, letting molecules escape into vapor. Substances with strong intermolecular forces (like ionic compounds or hydrogen-bonded molecules) require more thermal energy, giving higher boiling points. Container size and light pressure have negligible effect.
87. Which of the following is an example of thermal energy dominating over intermolecular forces?
ⓐ. Ice at 0°C
ⓑ. Steam at 100°C
ⓒ. Water at 25°C
ⓓ. Diamond structure
Correct Answer: Steam at 100°C
Explanation: In steam, molecules have sufficient thermal energy to overcome cohesive hydrogen bonding and move independently in random motion. In ice and diamond, intermolecular or covalent forces dominate, giving solid form. Liquid water at 25°C is intermediate, with comparable thermal energy and intermolecular forces.
88. Why do liquids evaporate even at temperatures below their boiling point?
ⓐ. Because intermolecular forces are absent
ⓑ. Because some surface molecules have enough thermal energy to overcome intermolecular attractions
ⓒ. Because pressure inside liquid becomes infinite
ⓓ. Because molecules at the bottom push surface molecules out
Correct Answer: Because some surface molecules have enough thermal energy to overcome intermolecular attractions
Explanation: In a liquid, molecules have a distribution of energies. Even at low temperatures, a fraction at the surface acquires enough energy to escape intermolecular forces, causing evaporation. This shows the dynamic competition between intermolecular cohesion and molecular motion.
89. Which condition favors the dominance of intermolecular forces over thermal energy?
ⓐ. Low temperature and high pressure
ⓑ. High temperature and low pressure
ⓒ. High temperature and high volume
ⓓ. Low pressure and high temperature
Correct Answer: Low temperature and high pressure
Explanation: At low temperature, molecular motion is reduced, and at high pressure, molecules are forced closer together, enhancing intermolecular forces. This condition promotes condensed states (liquid/solid). High temperature and low pressure promote gaseous state.
90. Which concept explains why liquefaction of gases is possible at low temperature and high pressure?
ⓐ. Gases have no intermolecular forces
ⓑ. Intermolecular forces become significant when thermal energy is lowered and molecules are compressed
ⓒ. Gases are colorless, so they condense
ⓓ. Thermal energy increases with high pressure
Correct Answer: Intermolecular forces become significant when thermal energy is lowered and molecules are compressed
Explanation: Cooling reduces thermal energy, while high pressure reduces intermolecular spacing. Both favor attractive forces, allowing molecules to bind and form a liquid. This principle is exploited in liquefying gases like $O_2, N_2, CO_2$.
91. Which state of matter results when intermolecular forces greatly exceed thermal energy?
ⓐ. Gas
ⓑ. Liquid
ⓒ. Solid
ⓓ. Plasma
Correct Answer: Solid
Explanation: In solids, strong intermolecular forces lock particles into a fixed lattice. Thermal energy is insufficient to break these bonds, so molecules only vibrate about equilibrium positions. This dominance of cohesive forces explains solids’ rigidity and definite shape.
92. Why does butter remain solid in a refrigerator but melt at room temperature?
ⓐ. Butter has no intermolecular forces at low temperature
ⓑ. At higher temperature, thermal energy overcomes the weak van der Waals forces in butter
ⓒ. Butter becomes ionic at room temperature
ⓓ. Butter molecules form hydrogen bonds only at low temperature
Correct Answer: At higher temperature, thermal energy overcomes the weak van der Waals forces in butter
Explanation: Butter is composed mainly of fat molecules held by weak dispersion forces. At low temperature, intermolecular forces dominate, keeping it solid. At room temperature, thermal energy disrupts these forces, producing a soft or liquid form.
93. Which factor explains why solid $CO_2$ (dry ice) sublimes at room temperature?
ⓐ. Strong dipole–dipole interactions
ⓑ. Weak dispersion forces easily overcome by thermal energy
ⓒ. Presence of hydrogen bonding
ⓓ. High ion–dipole attraction
Correct Answer: Weak dispersion forces easily overcome by thermal energy
Explanation: Dry ice molecules are held together only by London dispersion forces. At room temperature, thermal energy is sufficient to overcome these weak attractions, causing direct sublimation to gas without forming liquid under normal pressure.
94. Why does boiling point increase as we move from helium to xenon among noble gases?
ⓐ. Because thermal energy decreases down the group
ⓑ. Because dispersion forces become stronger with increasing electron cloud size
ⓒ. Because noble gases form dipole–dipole interactions
ⓓ. Because xenon forms hydrogen bonds
Correct Answer: Because dispersion forces become stronger with increasing electron cloud size
Explanation: Down the group, atoms are larger and more polarizable, enhancing dispersion forces. Stronger cohesive forces require more thermal energy to separate atoms, raising boiling points. Noble gases are nonpolar, so hydrogen bonding and dipole–dipole forces do not apply.
95. Which situation illustrates intermolecular forces dominating over thermal energy?
ⓐ. Ice cube floating in water
ⓑ. Steam rising from boiling water
ⓒ. Evaporation of alcohol at room temperature
ⓓ. Balloon filled with helium gas
Correct Answer: Ice cube floating in water
Explanation: In ice, strong hydrogen bonding holds water molecules rigidly despite low thermal energy, giving solid structure. Steam (B) and evaporation (C) show thermal energy dominance, while helium in a balloon (D) reflects negligible intermolecular forces.
96. Why does ethanol have a lower boiling point than water, even though both exhibit hydrogen bonding?
ⓐ. Ethanol has weaker intermolecular hydrogen bonds compared to water
ⓑ. Ethanol has more hydrogen bonds per molecule than water
ⓒ. Water molecules are nonpolar
ⓓ. Ethanol is heavier in molar mass
Correct Answer: Ethanol has weaker intermolecular hydrogen bonds compared to water
Explanation: Water forms extensive hydrogen bonding networks (each molecule can bond up to four neighbors), requiring more energy to break. Ethanol forms fewer hydrogen bonds due to only one –OH group and bulky alkyl chain, so intermolecular forces are weaker relative to thermal energy at boiling point.
97. Which condition favors gaseous state most strongly?
ⓐ. Strong intermolecular forces, low temperature
ⓑ. Weak intermolecular forces, high temperature
ⓒ. Strong intermolecular forces, high pressure
ⓓ. Strong intermolecular forces, high density
Correct Answer: Weak intermolecular forces, high temperature
Explanation: Weak forces reduce cohesion between particles, while high thermal energy increases random motion, favoring the gas phase. Conversely, strong forces with low temperature lead to solids or liquids.
98. What happens to viscosity of a liquid as thermal energy increases significantly?
ⓐ. Viscosity increases because molecules stick together more
ⓑ. Viscosity decreases because molecular motion overcomes cohesive forces
ⓒ. Viscosity remains unchanged
ⓓ. Viscosity becomes infinite
Correct Answer: Viscosity decreases because molecular motion overcomes cohesive forces
Explanation: With rising temperature, molecules move faster and escape intermolecular attractions more easily. This reduces internal resistance to flow (viscosity). For example, honey flows more readily when heated.
99. Why does high pressure promote liquefaction of gases?
ⓐ. Because pressure increases thermal energy of molecules
ⓑ. Because pressure reduces intermolecular distances, enhancing attractive forces
ⓒ. Because pressure eliminates intermolecular forces
ⓓ. Because pressure changes gases into solids directly
Correct Answer: Because pressure reduces intermolecular distances, enhancing attractive forces
Explanation: Compression brings molecules closer together, so attractive forces can dominate over thermal motion, allowing condensation into liquid. This effect is critical in industrial gas liquefaction processes.
100. Which example demonstrates thermal energy being reduced until intermolecular forces dominate?
ⓐ. Freezing of water into ice
ⓑ. Sublimation of iodine crystals
ⓒ. Vaporization of petrol
ⓓ. Expansion of a hot air balloon
Correct Answer: Freezing of water into ice
Explanation: Cooling reduces molecular kinetic energy. When thermal energy falls below the stabilizing hydrogen bonding forces, molecules lock into a crystalline lattice, forming solid ice. Sublimation (B) and vaporization (C) show the opposite, with thermal energy dominance, while balloon expansion (D) reflects increased molecular motion.
Welcome to Class 11 Chemistry MCQs – Chapter 5: States of Matter (Part 1). This chapter from the NCERT/CBSE Class 11 Chemistry syllabus explains the fundamental nature of matter in its three forms—solids, liquids, and gases. It builds the conceptual bridge between microscopic particle interactions and macroscopic observable properties. Understanding this chapter is crucial for board exams, JEE Main, NEET, and state-level entrance tests, since it forms the basis of Physical Chemistry. Practicing MCQs here will sharpen your grip on gas laws, intermolecular forces, thermodynamics of states, and kinetic theory, helping you tackle numerical and reasoning questions with speed and accuracy.
States of Matter is one of the most scoring chapters in Physical Chemistry—its principles are universal and applied in both academic exams and real-life scientific understanding.
Navigation & pages: The complete chapter contains 494 MCQs, organized into 5 structured parts (100 + 100 + 100 + 100 + 94). Part 1 features the first 100 MCQs, arranged into 10 pages with 10 questions per page. Use the page numbers above to view questions, and the Part buttons to continue with the next sets.
What you will learn & practice
- Intermolecular forces vs thermal energy
- Conditions for liquefaction of gases
- Gas laws:
- Boyle’s law (Pressure–Volume relation)
- Charles’ law (Temperature–Volume relation)
- Gay-Lussac’s law (Temperature–Pressure relation)
- Avogadro’s law
- Dalton’s law of partial pressure
- Ideal gas equation and gas constant (R)
- Kinetic theory of gases — assumptions, derivations, and limitations
- Deviation from ideal behavior; real gases and van der Waals equation
- Compressibility factor (Z) and its significance
- Liquefaction of gases (critical temperature, pressure, volume)
- Surface tension and viscosity of liquids
- Numerical problems on gas laws, density, molar mass, partial pressure, and kinetic energy
How this practice works
- Click an option to check instantly: green dot = correct, red icon = incorrect. The Correct Answer with explanation appears immediately.
- Use the 👁️ Eye icon to reveal the solution with explanation directly.
- Use the 📝 Notebook icon as a temporary scratchpad (notes are not saved permanently).
- Use the ⚠️ Alert icon to report mistakes or corrections instantly.
- Use the 💬 Message icon to ask questions, share insights, or start discussions.
Real value: These MCQs are strictly aligned with NCERT/CBSE syllabus, framed with previous-year question trends in mind, and explained with concise, exam-focused answers. Perfect for concept clarity, board exam one-mark practice, competitive tests, and last-minute revision.
The full chapter contains 494 solved multiple-choice questions in 5 systematic parts. This page (Part 1) presents the first 100 MCQs with detailed answers and explanations. Explore more with: Class 11 Chemistry – Chapter Index, Class 11 – Subject Index, All Classes – MCQ Library.
- 👉 Total MCQs in this chapter: 494 (100 + 100 + 100 + 100 + 94)
- 👉 This page: First 100 MCQs with answers & explanations (in 10 pages)
- 👉 Best for: Boards • JEE/NEET • states of matter practice • concept revision • Chemistry quizzes
- 👉 Next: Use the Part buttons and page numbers above to continue
FAQs on States of Matter ▼
▸ What are States of Matter MCQs in Class 11 Chemistry?
These are multiple-choice questions from Chapter 5 of NCERT Class 11 Chemistry – States of Matter. They test your knowledge of intermolecular forces, gas laws, liquefaction, and properties of solids, liquids, and gases.
▸ How many MCQs are available in this chapter?
There are a total of 494 MCQs from States of Matter. They are divided into 5 structured sets – four sets of 100 questions each and one set of 94 questions.
▸ Are these Chemistry MCQs useful for NCERT and CBSE board exams?
Yes, these MCQs are designed strictly from the NCERT/CBSE Class 11 syllabus and are highly useful for board exam preparation, quick revision, and conceptual clarity.
▸ Are States of Matter MCQs important for JEE and NEET?
Yes, this chapter is very important for JEE, NEET, and other entrance exams. Topics like ideal gas equation, deviations from ideal behavior, liquefaction of gases, and critical temperature are frequently asked in competitive tests.
▸ Do these MCQs include correct answers and explanations?
Yes, every MCQ is provided with the correct answer along with explanations wherever required. This ensures students understand the reasoning behind the solutions, not just memorize facts.
▸ Who should practice States of Matter MCQs?
These MCQs are useful for Class 11 students, CBSE and state board aspirants, and candidates preparing for JEE, NEET, NDA, UPSC, and other competitive entrance exams.
▸ Can I practice these Chemistry MCQs online for free?
Yes, all States of Matter MCQs on GK Aim are available online for free and can be practiced anytime using mobile, tablet, or desktop.
▸ Are these MCQs helpful for quick revision?
Yes, practicing these MCQs regularly helps in quick revision, strengthens memory recall, and improves exam performance by enhancing speed and accuracy in Chemistry problem-solving.
▸ Do these MCQs cover both basics and advanced concepts?
Yes, the MCQs cover everything from basics like properties of gases and liquids to advanced concepts such as Van der Waals equation, compressibility factor, and critical constants.
▸ Why are the 494 MCQs divided into 5 parts?
The MCQs are divided into 5 smaller sets to make practice more structured and manageable. This step-by-step division allows students to prepare without feeling overloaded.
▸ Which subtopics are covered in these States of Matter MCQs?
These MCQs cover important subtopics such as intermolecular forces, gas laws (Boyle’s law, Charles’s law, Gay-Lussac’s law, Avogadro’s law), ideal gas equation, Dalton’s law, Graham’s law, liquefaction of gases, and liquid crystals.
▸ Are there MCQs on real-life applications of gas laws?
Yes, the MCQs also include real-life applications such as scuba diving (Dalton’s law), hot air balloons (Charles’s law), and respiration (Boyle’s law), which are frequently highlighted in exams like JEE and NEET.
▸ Can teachers and institutes use these MCQs?
Yes, teachers and coaching institutes can use these MCQs as ready-made practice material, assignments, and quizzes for students preparing for board and competitive exams.
▸ Are these States of Matter MCQs mobile-friendly?
Yes, the States of Matter MCQs pages are fully optimized for smartphones and tablets so that students can practice anytime, anywhere.
▸ Can I download or save these MCQs for offline study?
Yes, you can download these States of Matter MCQs in PDF format for offline study. Please visit our website shop.gkaim.com