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1. Who is credited with developing the kinetic theory of gases in its modern form?
ⓐ. Albert Einstein
ⓑ. James Clerk Maxwell
ⓒ. Isaac Newton
ⓓ. Robert Boyle
Correct Answer: James Clerk Maxwell
Explanation: James Clerk Maxwell and Ludwig Boltzmann developed the modern kinetic theory of gases. Newton contributed to mechanics, Boyle studied gas laws experimentally, and Einstein later applied kinetic theory to Brownian motion.
2. What is the main postulate of kinetic theory of gases?
ⓐ. Gases are made up of heavy particles that move slowly.
ⓑ. Gas molecules are in constant random motion and collisions are perfectly elastic.
ⓒ. Gas particles attract each other strongly.
ⓓ. Gas particles have negligible kinetic energy at all temperatures.
Correct Answer: Gas molecules are in constant random motion and collisions are perfectly elastic.
Explanation: The kinetic theory assumes gas molecules are point masses, moving randomly with elastic collisions. This explains macroscopic properties like pressure and temperature. Strong attractions (C) contradict the ideal gas assumption, and (A, D) are incorrect.
3. Which property of a gas is directly proportional to the average kinetic energy of its molecules?
ⓐ. Pressure
ⓑ. Volume
ⓒ. Temperature
ⓓ. Density
Correct Answer: Temperature
Explanation: The average kinetic energy of a molecule is $\frac{3}{2}k_B T$, where $k_B$ is Boltzmann’s constant and $T$ is temperature. Hence, temperature is a direct measure of average kinetic energy. Pressure and volume depend on both kinetic energy and molecular number.
4. In kinetic theory, what is the assumption about the size of gas molecules compared to the distance between them?
ⓐ. They are much larger than the distance.
ⓑ. They are equal to the distance.
ⓒ. They are negligible compared to the distance.
ⓓ. They are infinitely large.
Correct Answer: They are negligible compared to the distance.
Explanation: One key assumption is that gas molecules are point-like and their actual size is negligible compared to the intermolecular distance. This allows us to ignore molecular volume in the ideal gas model.
5. Which equation is derived from kinetic theory of gases?
ⓐ. $PV = nRT$
ⓑ. $E = mc^2$
ⓒ. $F = ma$
ⓓ. $V = IR$
Correct Answer: $PV = nRT$
Explanation: The ideal gas equation $PV = nRT$ can be derived using the assumptions of kinetic theory by relating pressure to the average kinetic energy of molecules. The other equations belong to different fields: relativity, mechanics, and electricity.
6. According to kinetic theory, pressure of a gas arises due to:
ⓐ. Intermolecular attraction
ⓑ. Weight of molecules
ⓒ. Collisions of molecules with container walls
ⓓ. Vibrations of molecules
Correct Answer: Collisions of molecules with container walls
Explanation: Gas pressure results from continuous bombardment of container walls by gas molecules. Each elastic collision imparts momentum, and collectively this produces measurable pressure.
7. The mean kinetic energy of one mole of an ideal gas at temperature $T$ is:
ⓐ. $\frac{1}{2}k_B T$
ⓑ. $\frac{3}{2}RT$
ⓒ. $\frac{3}{2}k_B T$
ⓓ. $RT$
Correct Answer: $\frac{3}{2}RT$
Explanation: For one molecule, mean kinetic energy is $\frac{3}{2}k_B T$. For one mole, multiply by Avogadro’s number, giving $\frac{3}{2}RT$. $k_B$ and $R$ are related as $R = N_A k_B$.
8. What is the root mean square (rms) speed of gas molecules?
ⓐ. $\sqrt{\frac{2RT}{M}}$
ⓑ. $\sqrt{\frac{3RT}{M}}$
ⓒ. $\sqrt{\frac{RT}{M}}$
ⓓ. $\sqrt{\frac{4RT}{M}}$
Correct Answer: $\sqrt{\frac{3RT}{M}}$
Explanation: The rms speed is defined as $v_{rms} = \sqrt{\frac{3RT}{M}}$, where $R$ is the gas constant, $T$ the absolute temperature, and $M$ the molar mass. This comes from equating pressure with kinetic energy per mole.
9. If temperature of a gas is doubled (in Kelvin), its rms speed will:
ⓐ. Remain constant
ⓑ. Double
ⓒ. Increase by factor $\sqrt{2}$
ⓓ. Decrease by half
Correct Answer: Increase by factor $\sqrt{2}$
Explanation: Since $v_{rms} \propto \sqrt{T}$, doubling $T$ gives $v_{rms,new} = \sqrt{2} v_{rms,old}$. Hence speed increases by $\sqrt{2}$, not by factor of 2.
10. Which law can be derived from the kinetic theory assumption that average kinetic energy is proportional to absolute temperature?
ⓐ. Boyle’s law
ⓑ. Charles’ law
ⓒ. Newton’s law
ⓓ. Hooke’s law
Correct Answer: Charles’ law
Explanation: Charles’ law states that at constant pressure, the volume of a gas is directly proportional to absolute temperature. Kinetic theory shows average molecular kinetic energy depends on $T$, hence supporting Charles’ law.
11. Who was the first scientist to propose that matter is made up of indivisible particles called “atoms”?
ⓐ. John Dalton
ⓑ. Democritus
ⓒ. James Clerk Maxwell
ⓓ. Albert Einstein
Correct Answer: Democritus
Explanation: In the 5th century BCE, Democritus introduced the idea that all matter consists of indivisible particles called atoms. Dalton (1803) revived this idea with his atomic theory. Maxwell and Einstein contributed later to kinetic theory and statistical mechanics, respectively.
12. Which scientist formulated the Atomic Theory in the early 19th century that laid the foundation for kinetic theory?
ⓐ. Robert Boyle
ⓑ. John Dalton
ⓒ. Ludwig Boltzmann
ⓓ. Galileo Galilei
Correct Answer: John Dalton
Explanation: In 1803, John Dalton proposed the Atomic Theory stating that matter is composed of tiny indivisible atoms, which combine in simple ratios to form compounds. This became a cornerstone for later development of kinetic theory.
13. Boyle’s Law, one of the early gas laws, relates pressure and volume as:
ⓐ. $P \propto V$
ⓑ. $PV = \text{constant}$
ⓒ. $P \propto T$
ⓓ. $V \propto T$
Correct Answer: $PV = \text{constant}$
Explanation: In 1662, Robert Boyle found experimentally that at constant temperature, the product of pressure and volume of a gas remains constant. This was an early empirical step towards kinetic theory.
14. Which scientist provided the first clear statement that gas pressure is due to molecular collisions?
ⓐ. Newton
ⓑ. Bernoulli
ⓒ. Maxwell
ⓓ. Avogadro
Correct Answer: Bernoulli
Explanation: In 1738, Daniel Bernoulli explained that pressure arises from collisions of molecules with container walls. His work, Hydrodynamica, is regarded as the foundation of kinetic theory.
15. Avogadro’s Hypothesis (1811) states:
ⓐ. Equal volumes of gases contain equal number of molecules at same temperature and pressure.
ⓑ. Pressure is inversely proportional to volume.
ⓒ. Energy is equally distributed among molecules.
ⓓ. Molecules are in continuous random motion.
Correct Answer: Equal volumes of gases contain equal number of molecules at same temperature and pressure.
Explanation: Amedeo Avogadro proposed that equal volumes of gases under identical conditions of temperature and pressure contain the same number of molecules. This hypothesis clarified molecular mass and paved the way for kinetic theory.
16. Who introduced the statistical interpretation of entropy and molecular distribution of speeds?
ⓐ. Maxwell
ⓑ. Boltzmann
ⓒ. Dalton
ⓓ. Joule
Correct Answer: Boltzmann
Explanation: Ludwig Boltzmann, in the late 19th century, introduced statistical mechanics, the Maxwell-Boltzmann distribution, and the entropy formula $S = k_B \ln W$. This was crucial for modern kinetic theory.
17. Maxwell’s contribution to kinetic theory includes:
ⓐ. Pressure–volume law
ⓑ. Statistical distribution of molecular speeds
ⓒ. Discovery of Brownian motion
ⓓ. Law of conservation of energy
Correct Answer: Statistical distribution of molecular speeds
Explanation: James Clerk Maxwell (1859) derived the distribution law of molecular speeds in a gas, now called the Maxwell-Boltzmann distribution. This connected microscopic motion with macroscopic properties of gases.
18. Who explained Brownian motion as evidence of molecular kinetic theory?
ⓐ. Robert Brown
ⓑ. Einstein
ⓒ. Avogadro
ⓓ. Newton
Correct Answer: Einstein
Explanation: Although Robert Brown observed the motion of pollen grains in 1827, Einstein in 1905 explained Brownian motion quantitatively using kinetic theory, proving the molecular nature of matter.
19. Joule’s contribution to kinetic theory was primarily through:
ⓐ. The kinetic equation of gases
ⓑ. The law of conservation of energy
ⓒ. Avogadro’s hypothesis
ⓓ. Distribution of speeds
Correct Answer: The law of conservation of energy
Explanation: James Prescott Joule established the mechanical equivalent of heat, showing heat is a form of energy. This connected thermodynamics with kinetic theory by relating molecular motion to temperature and energy.
20. Which scientist’s work finally unified the microscopic kinetic theory with macroscopic thermodynamics?
ⓐ. Boltzmann
ⓑ. Maxwell
ⓒ. Einstein
ⓓ. Carnot
Correct Answer: Boltzmann
Explanation: Ludwig Boltzmann developed statistical mechanics, providing a bridge between microscopic particle dynamics and macroscopic thermodynamic laws, thus completing the historical development of kinetic theory.
21. Which fundamental quantity connects kinetic theory with thermodynamics?
ⓐ. Heat capacity
ⓑ. Temperature
ⓒ. Volume
ⓓ. Pressure
Correct Answer: Temperature
Explanation: In kinetic theory, the average kinetic energy of gas molecules is proportional to absolute temperature: $\langle E_k \rangle = \tfrac{3}{2} k_B T$. In thermodynamics, temperature is a macroscopic variable that determines heat flow direction. Thus, temperature bridges microscopic particle motion with macroscopic thermodynamic behavior.
22. In statistical mechanics, entropy is expressed as:
ⓐ. $S = \frac{Q}{T}$
ⓑ. $S = k_B \ln W$
ⓒ. $S = mc\Delta T$
ⓓ. $S = PV$
Correct Answer: $S = k_B \ln W$
Explanation: Boltzmann’s entropy formula $S = k_B \ln W$ relates entropy $S$ to the number of microstates $W$. This connects thermodynamic entropy, defined via $\Delta S = \frac{Q_{rev}}{T}$, with the microscopic probability distribution of particles, thus linking kinetic theory to statistical mechanics.
23. The internal energy of an ideal monatomic gas is given by:
ⓐ. $U = \frac{1}{2}nRT$
ⓑ. $U = \frac{3}{2}nRT$
ⓒ. $U = nRT$
ⓓ. $U = \frac{5}{2}nRT$
Correct Answer: $U = \frac{3}{2}nRT$
Explanation: From kinetic theory, average kinetic energy per molecule is $\tfrac{3}{2}k_B T$. For one mole, $U = N_A \cdot \tfrac{3}{2}k_B T = \tfrac{3}{2}RT$. For $n$ moles, $U = \tfrac{3}{2}nRT$. Thermodynamically, this matches the energy calculated from specific heats, thus confirming consistency between the two theories.
24. Which law of thermodynamics is consistent with the conservation of molecular kinetic energy in elastic collisions?
ⓐ. Zeroth Law
ⓑ. First Law
ⓒ. Second Law
ⓓ. Third Law
Correct Answer: First Law
Explanation: The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed. In kinetic theory, during elastic molecular collisions, total kinetic energy remains conserved, directly aligning with the principle of energy conservation in thermodynamics.
25. What does the Maxwell-Boltzmann distribution describe in relation to thermodynamics?
ⓐ. Variation of pressure with temperature
ⓑ. Distribution of molecular speeds at a given temperature
ⓒ. Change of entropy with volume
ⓓ. Energy conversion between heat and work
Correct Answer: Distribution of molecular speeds at a given temperature
Explanation: Maxwell-Boltzmann distribution shows the probability of molecules having certain speeds in a gas at temperature $T$. This statistical distribution connects microscopic speed variations with macroscopic thermodynamic variables like pressure, temperature, and entropy.
26. In statistical mechanics, the partition function $Z$ is important because:
ⓐ. It determines the type of collisions.
ⓑ. It acts as a bridge between microscopic states and macroscopic properties.
ⓒ. It gives the heat capacity directly.
ⓓ. It is unrelated to thermodynamics.
Correct Answer: It acts as a bridge between microscopic states and macroscopic properties.
Explanation: The partition function $Z = \sum e^{-E_i/k_BT}$ sums over all possible energy states. It allows calculation of thermodynamic quantities like free energy, entropy, and internal energy. Thus, it connects statistical mechanics with thermodynamics by providing a link between microstates and observable quantities.
27. Which thermodynamic property is most directly associated with disorder in statistical mechanics?
ⓐ. Enthalpy
ⓑ. Entropy
ⓒ. Internal Energy
ⓓ. Pressure
Correct Answer: Entropy
Explanation: In thermodynamics, entropy measures disorder and spontaneity of processes. In statistical mechanics, entropy is related to the number of microstates accessible to a system. The greater the number of microstates, the higher the entropy, which quantitatively links microscopic randomness with macroscopic thermodynamic behavior.
28. The Second Law of Thermodynamics states that entropy of an isolated system never decreases. How is this interpreted statistically?
ⓐ. Molecules stop moving at equilibrium.
ⓑ. Probability favors the most disordered microstate distribution.
ⓒ. Heat capacity becomes zero.
ⓓ. Internal energy becomes constant.
Correct Answer: Probability favors the most disordered microstate distribution.
Explanation: In statistical mechanics, systems evolve toward macrostates with the largest number of microstates, which corresponds to higher entropy. Thus, the Second Law reflects the natural statistical tendency of systems to move toward equilibrium with maximum disorder.
29. The equipartition theorem relates directly to which thermodynamic property?
ⓐ. Heat capacity
ⓑ. Enthalpy
ⓒ. Pressure
ⓓ. Entropy
Correct Answer: Heat capacity
Explanation: Equipartition theorem states each quadratic degree of freedom contributes $\tfrac{1}{2}k_B T$ per molecule to energy. For a monatomic ideal gas with 3 translational degrees, this leads to $C_v = \tfrac{3}{2}R$. Thus, kinetic theory predictions of heat capacity match thermodynamic values, showing the link between micro- and macroscopic descriptions.
30. In thermodynamics, free energy determines spontaneity. Which statistical mechanics quantity helps calculate free energy?
ⓐ. Partition function
ⓑ. Maxwell’s equations
ⓒ. Pressure-volume work
ⓓ. Specific heat
Correct Answer: Partition function
Explanation: The Helmholtz free energy is given by $F = -k_B T \ln Z$, where $Z$ is the partition function. Thus, free energy—central to thermodynamics—can be calculated statistically, bridging molecular-level energy states with macroscopic spontaneity criteria.
31. How does kinetic theory help explain the states of matter?
ⓐ. By describing atomic masses
ⓑ. By analyzing nuclear structure
ⓒ. By relating molecular motion to macroscopic states (solid, liquid, gas)
ⓓ. By assuming matter has no internal motion
Correct Answer: By relating molecular motion to macroscopic states (solid, liquid, gas)
Explanation: In solids, molecules vibrate about fixed positions due to strong intermolecular forces. In liquids, molecules move but remain close together. In gases, molecules move freely with negligible forces. Kinetic theory links these behaviors of molecular motion to the macroscopic states of matter.
32. Why are gases highly compressible compared to solids and liquids?
ⓐ. Gas molecules are very heavy.
ⓑ. Gas molecules occupy negligible volume relative to empty space.
ⓒ. Gas molecules are strongly bound to each other.
ⓓ. Gas molecules lose energy on collision.
Correct Answer: Gas molecules occupy negligible volume relative to empty space.
Explanation: Kinetic theory assumes molecules are point masses with large intermolecular distances compared to their size. Thus, gases contain mostly empty space, making them compressible. Solids and liquids have closely packed molecules, so they are nearly incompressible.
33. Which molecular-level property explains diffusion in gases?
ⓐ. Vibrational motion of molecules
ⓑ. Collisions with container walls
ⓒ. Random motion of molecules due to kinetic energy
ⓓ. Attraction between gas molecules
Correct Answer: Random motion of molecules due to kinetic energy
Explanation: Diffusion is the spreading of molecules from high to low concentration. Kinetic theory states that gas molecules move randomly with kinetic energy, causing them to mix spontaneously. Intermolecular attractions play almost no role in ideal gases.
34. Why does pressure of a gas increase with temperature at constant volume?
ⓐ. Molecules become heavier.
ⓑ. Molecules move faster and collide more frequently with container walls.
ⓒ. Number of molecules increases.
ⓓ. Container shrinks at high temperature.
Correct Answer: Molecules move faster and collide more frequently with container walls.
Explanation: According to kinetic theory, pressure arises from collisions with walls. As temperature rises, average kinetic energy $\langle E_k \rangle = \tfrac{3}{2}k_B T$ increases, leading to stronger and more frequent collisions, and thus higher pressure.
35. Why do liquids evaporate even at room temperature?
ⓐ. All molecules have the same kinetic energy.
ⓑ. Some molecules always have enough kinetic energy to overcome intermolecular forces.
ⓒ. Molecules stop colliding at the surface.
ⓓ. External pressure forces molecules out.
Correct Answer: Some molecules always have enough kinetic energy to overcome intermolecular forces.
Explanation: Maxwell-Boltzmann distribution shows molecules have varying speeds. Even at moderate temperatures, some molecules gain enough energy to escape liquid’s surface tension, leading to evaporation. This is an application of kinetic theory to explain phase change at the molecular level.
36. Why does ice float on water, considering kinetic theory and molecular structure?
ⓐ. Ice molecules move faster than water molecules.
ⓑ. Ice is denser due to strong hydrogen bonds.
ⓒ. Open hexagonal structure in ice traps empty spaces, lowering density.
ⓓ. Ice molecules are heavier than water molecules.
Correct Answer: Open hexagonal structure in ice traps empty spaces, lowering density.
Explanation: Kinetic theory combined with molecular bonding explains that water molecules form hydrogen-bonded hexagonal structures in ice. These structures have more empty space, reducing density. Hence, ice floats on water despite being solid.
37. How does kinetic theory explain Brownian motion?
ⓐ. Molecules move in fixed paths.
ⓑ. Random collisions of molecules transfer momentum to suspended particles.
ⓒ. Molecules exert long-range forces on particles.
ⓓ. Particles move due to magnetic effects.
Correct Answer: Random collisions of molecules transfer momentum to suspended particles.
Explanation: Brownian motion is the zig-zag movement of small particles suspended in a fluid. Kinetic theory explains this as unequal molecular bombardment on different sides of the particle, giving evidence of molecular motion at the microscopic level.
38. Why do real gases deviate from ideal behavior at high pressure?
ⓐ. Molecular collisions stop.
ⓑ. Intermolecular forces and finite molecular volume become significant.
ⓒ. Molecules lose kinetic energy permanently.
ⓓ. Gas molecules disappear.
Correct Answer: Intermolecular forces and finite molecular volume become significant.
Explanation: Ideal kinetic theory ignores molecular size and forces. At high pressure, molecules are closer, so intermolecular attractions and excluded volume cannot be neglected, causing deviations described by Van der Waals equation.
39. Why does heat conduction occur in gases according to kinetic theory?
ⓐ. Molecules carry energy from hotter to cooler regions via random motion and collisions.
ⓑ. Molecules radiate energy spontaneously without collisions.
ⓒ. Heat moves only by convection currents.
ⓓ. Heat is stored permanently in molecules.
Correct Answer: Molecules carry energy from hotter to cooler regions via random motion and collisions.
Explanation: Kinetic theory states molecules transfer kinetic energy during collisions. Molecules in hotter regions move faster, collide with slower ones in cooler regions, and transfer energy, resulting in thermal conduction.
40. How does kinetic theory explain the pressure exerted by liquids in a container?
ⓐ. Molecules in liquids are at rest.
ⓑ. Liquid molecules vibrate but do not collide.
ⓒ. Random motion and collisions of molecules with each other and container walls create hydrostatic pressure.
ⓓ. Pressure arises only due to gravity.
Correct Answer: Random motion and collisions of molecules with each other and container walls create hydrostatic pressure.
Explanation: Although intermolecular forces are stronger in liquids than gases, molecules still move randomly. Their collisions with container walls and transfer of momentum contribute to pressure in liquids, consistent with kinetic theory principles.
41. Who is credited with proposing the “plum pudding” model of the atom?
ⓐ. Ernest Rutherford
ⓑ. J.J. Thomson
ⓒ. Niels Bohr
ⓓ. James Chadwick
Correct Answer: J.J. Thomson
Explanation: In 1897, after discovering the electron, J.J. Thomson proposed that atoms are made of a positively charged sphere with negatively charged electrons embedded in it, like plums in a pudding. This model was later replaced by Rutherford’s nuclear model.
42. Rutherford’s alpha scattering experiment led to the conclusion that:
ⓐ. Electrons orbit randomly in the atom.
ⓑ. Atoms have a nucleus with most of the mass concentrated in it.
ⓒ. Electrons are spread uniformly in the atom.
ⓓ. Atoms are indivisible.
Correct Answer: Atoms have a nucleus with most of the mass concentrated in it.
Explanation: Rutherford observed that most alpha particles passed undeflected, but some were deflected at large angles. This showed atoms are mostly empty space, with mass and positive charge concentrated in a small nucleus.
43. According to Bohr’s model, the angular momentum of an electron in orbit is quantized as:
ⓐ. $L = mvr$
ⓑ. $L = nh/2\pi$
ⓒ. $L = nh$
ⓓ. $L = \frac{mv^2}{r}$
Correct Answer: $L = nh/2\pi$
Explanation: Bohr postulated that only those electron orbits are stable for which the angular momentum $L$ is an integer multiple of $h/2\pi$. This quantization explained the stability of atoms and spectral lines of hydrogen.
44. The atomic number of an element represents:
ⓐ. Number of protons in the nucleus
ⓑ. Number of neutrons in the nucleus
ⓒ. Number of nucleons in the nucleus
ⓓ. Mass of the atom in atomic units
Correct Answer: Number of protons in the nucleus
Explanation: Atomic number (Z) is the number of protons in the nucleus. It determines the chemical identity of an element. Mass number (A) is the sum of protons and neutrons. Neutrons (B) and nucleons (C) are related but not equal to Z.
45. Isotopes of an element have:
ⓐ. Same number of protons, different number of neutrons
ⓑ. Different number of protons, same number of neutrons
ⓒ. Same mass number, different atomic number
ⓓ. Completely different electronic configuration
Correct Answer: Same number of protons, different number of neutrons
Explanation: Isotopes have the same atomic number (same protons and electrons, hence same chemical properties) but different mass number due to different neutrons. For example, $^{12}C$ and $^{14}C$.
46. Which scientist discovered the neutron?
ⓐ. Rutherford
ⓑ. Bohr
ⓒ. Chadwick
ⓓ. Einstein
Correct Answer: Chadwick
Explanation: James Chadwick discovered the neutron in 1932. He showed that the nucleus contained electrically neutral particles in addition to protons, explaining previously unexplained atomic masses.
47. Which of the following correctly describes atomic orbitals?
ⓐ. Fixed circular paths of electrons
ⓑ. Definite tracks of electrons around nucleus
ⓒ. Regions around nucleus with high probability of finding electrons
ⓓ. Solid shells where electrons reside
Correct Answer: Regions around nucleus with high probability of finding electrons
Explanation: Modern quantum mechanics replaces Bohr’s fixed orbits with orbitals—probability distributions described by Schrödinger’s wave equations. Orbitals show where electrons are most likely to be found.
48. What is the principal quantum number $n$ associated with?
ⓐ. Shape of orbital
ⓑ. Orientation of orbital
ⓒ. Energy level and size of orbital
ⓓ. Spin of electron
Correct Answer: Energy level and size of orbital
Explanation: The principal quantum number $n$ determines the shell number, energy, and approximate distance of electrons from the nucleus. Larger $n$ means higher energy and larger orbital size. Shape (l), orientation (m), and spin (s) are other quantum numbers.
49. The Pauli Exclusion Principle states that:
ⓐ. No two electrons can have the same set of four quantum numbers.
ⓑ. Electrons occupy orbitals of lowest energy first.
ⓒ. Orbitals are filled singly before pairing.
ⓓ. Electrons revolve in circular paths.
Correct Answer: No two electrons can have the same set of four quantum numbers.
Explanation: Pauli’s principle ensures the uniqueness of electron arrangements in atoms. While Hund’s rule (C) and Aufbau principle (B) also govern filling, Pauli’s rule states no orbital can hold more than two electrons with opposite spins.
50. Which experiment provided evidence for the wave nature of electrons?
ⓐ. Rutherford scattering experiment
ⓑ. Thomson’s cathode ray experiment
ⓒ. Davisson–Germer electron diffraction experiment
ⓓ. Millikan oil drop experiment
Correct Answer: Davisson–Germer electron diffraction experiment
Explanation: In 1927, Davisson and Germer observed diffraction patterns of electrons scattered from a nickel crystal. This confirmed de Broglie’s hypothesis of matter waves and wave-particle duality, a cornerstone of atomic and molecular structure.
51. Which of the following is an example of an ionic bond?
ⓐ. $\text{H}_2$
ⓑ. $\text{NaCl}$
ⓒ. $\text{CH}_4$
ⓓ. $\text{O}_2$
Correct Answer: $\text{NaCl}$
Explanation: Ionic bonds are formed by transfer of electrons from one atom to another, usually between metals and non-metals. Sodium loses an electron to form $\text{Na}^+$, and chlorine gains an electron to form $\text{Cl}^-$. The electrostatic attraction forms an ionic bond. Covalent bonds occur in $\text{H}_2, \text{CH}_4, \text{O}_2$.
52. Which of the following is a property of covalent compounds?
ⓐ. High melting and boiling points
ⓑ. Conduct electricity in molten state
ⓒ. Poor electrical conductivity
ⓓ. Always soluble in water
Correct Answer: Poor electrical conductivity
Explanation: Covalent compounds consist of shared electron pairs with no free ions. This makes them poor conductors of electricity. Ionic compounds conduct in molten or aqueous states, and metals conduct due to free electrons. Solubility in water (D) varies depending on polarity.
53. What type of bond holds the atoms together in a diamond crystal?
ⓐ. Ionic bond
ⓑ. Covalent bond
ⓒ. Metallic bond
ⓓ. Hydrogen bond
Correct Answer: Covalent bond
Explanation: Diamond is made of carbon atoms, each covalently bonded to four others in a tetrahedral network. This strong, extensive covalent bonding makes diamond very hard with a high melting point. Metallic and ionic bonds are absent here.
54. Metallic bonds are best explained by which model?
ⓐ. Localized electron pairs
ⓑ. Electron cloud or sea model
ⓒ. Proton exchange model
ⓓ. Magnetic attraction
Correct Answer: Electron cloud or sea model
Explanation: In metals, valence electrons are delocalized, forming a “sea of electrons” around positively charged metal ions. This explains high electrical and thermal conductivity, malleability, and ductility. Other models do not account for these metallic properties.
55. Which compound is held together by covalent bonding?
ⓐ. $\text{MgCl}_2$
ⓑ. $\text{NaBr}$
ⓒ. $\text{CO}_2$
ⓓ. $\text{KCl}$
Correct Answer: $\text{CO}_2$
Explanation: In carbon dioxide, carbon shares electrons with oxygen atoms, forming double covalent bonds. Ionic bonds exist in salts like $\text{MgCl}_2, \text{NaBr}, \text{KCl}$. Thus, $\text{CO}_2$ is a covalently bonded molecule.
56. What is the primary reason ionic compounds conduct electricity in molten state but not in solid state?
ⓐ. Ions are fixed in lattice in solid but free to move in molten state.
ⓑ. Electrons are free in solid but bound in molten state.
ⓒ. Covalent bonds break only in molten state.
ⓓ. Metallic bonds form only in molten state.
Correct Answer: Ions are fixed in lattice in solid but free to move in molten state.
Explanation: In solid ionic crystals, ions are held rigidly in a lattice and cannot move. On melting, the lattice breaks down, and free ions move, allowing conduction of electricity. Metallic conduction does not apply here.
57. Which type of bond explains the malleability and ductility of metals?
ⓐ. Ionic bond
ⓑ. Covalent bond
ⓒ. Metallic bond
ⓓ. Hydrogen bond
Correct Answer: Metallic bond
Explanation: In metallic bonding, delocalized electrons form a flexible “glue” that holds metal ions together. When hammered or stretched, layers of ions can slide without breaking the bond, giving metals malleability and ductility.
58. Which of the following has both ionic and covalent bonds?
ⓐ. $\text{NaCl}$
ⓑ. $\text{CaCO}_3$
ⓒ. $\text{O}_2$
ⓓ. $\text{Fe}$
Correct Answer: $\text{CaCO}_3$
Explanation: In calcium carbonate, the $\text{Ca}^{2+}$ and $\text{CO}_3^{2-}$ ions are held together by ionic bonds. Within the carbonate ion, carbon and oxygen are linked by covalent bonds. Hence, it exhibits both types.
59. Which of the following best describes a polar covalent bond?
ⓐ. Equal sharing of electrons between atoms
ⓑ. Complete transfer of electrons from one atom to another
ⓒ. Unequal sharing of electrons leading to partial charges
ⓓ. No sharing of electrons at all
Correct Answer: Unequal sharing of electrons leading to partial charges
Explanation: In polar covalent bonds, the more electronegative atom attracts electrons more strongly, creating partial charges. For example, in $\text{H}_2\text{O}$, oxygen pulls electrons closer, giving oxygen a partial negative charge and hydrogen a partial positive charge.
60. Which of the following has metallic bonding?
ⓐ. Sodium metal
ⓑ. $\text{NaCl}$
ⓒ. Water
ⓓ. Ammonia
Correct Answer: Sodium metal
Explanation: In sodium, atoms release valence electrons into a delocalized sea, forming metallic bonds. This explains sodium’s electrical conductivity and metallic properties. $\text{NaCl}$ has ionic bonds, while water and ammonia have covalent bonds.
61. Which of the following statements about ionic bonds is correct?
ⓐ. They are formed by equal sharing of electrons.
ⓑ. They involve complete transfer of electrons between atoms.
ⓒ. They exist only in metals.
ⓓ. They are weaker than van der Waals forces.
Correct Answer: They involve complete transfer of electrons between atoms.
Explanation: Ionic bonds result from the transfer of electrons from a metal to a non-metal, forming oppositely charged ions held by electrostatic attraction. For example, NaCl forms when sodium donates an electron to chlorine. Covalent bonds involve sharing, not transfer, and metallic bonds involve delocalized electrons.
62. Which of the following is a characteristic property of metallic bonds?
ⓐ. High brittleness
ⓑ. Good electrical conductivity
ⓒ. Low melting points
ⓓ. Lack of luster
Correct Answer: Good electrical conductivity
Explanation: In metallic bonds, valence electrons are delocalized into an “electron sea,” allowing free movement. This explains metals’ ability to conduct electricity and heat efficiently. Metals are usually not brittle, and they have luster and high melting points.
63. Which of the following compounds has a purely covalent bond?
ⓐ. $\text{KCl}$
ⓑ. $\text{CaF}_2$
ⓒ. $\text{CH}_4$
ⓓ. $\text{MgO}$
Correct Answer: $\text{CH}_4$
Explanation: In methane ($\text{CH}_4$), carbon shares its valence electrons equally with hydrogen atoms, forming covalent bonds. $\text{KCl}, \text{CaF}_2, \text{MgO}$ are ionic compounds formed by electron transfer.
64. Which factor most strongly determines whether a bond is ionic or covalent?
ⓐ. Atomic radius
ⓑ. Electronegativity difference between atoms
ⓒ. Ionization energy only
ⓓ. Atomic mass difference
Correct Answer: Electronegativity difference between atoms
Explanation: A large electronegativity difference (greater than \~1.7) favors ionic bonding due to electron transfer. Smaller differences lead to covalent bonding, with equal or unequal sharing of electrons. Atomic radius and mass play secondary roles.
65. Which of the following best explains metallic luster?
ⓐ. Strong ionic forces
ⓑ. Vibration of nuclei
ⓒ. Free electrons reflecting incident light
ⓓ. Presence of covalent bonds
Correct Answer: Free electrons reflecting incident light
Explanation: In metals, delocalized electrons oscillate in response to incident light and re-emit radiation, giving metals their shiny surface. Ionic and covalent bonds cannot explain this property.
66. Why does sodium chloride (NaCl) have a high melting point?
ⓐ. Presence of covalent bonds
ⓑ. Strong electrostatic forces between $\text{Na}^+$ and $\text{Cl}^-$ ions
ⓒ. Free electron sea
ⓓ. Weak van der Waals forces
Correct Answer: Strong electrostatic forces between $\text{Na}^+$ and $\text{Cl}^-$ ions
Explanation: Ionic compounds like NaCl are held together by strong Coulombic forces between oppositely charged ions, requiring large energy to break, leading to high melting and boiling points.
67. Which type of bonding explains why aluminum is both strong and a good conductor of electricity?
ⓐ. Covalent bonding
ⓑ. Ionic bonding
ⓒ. Metallic bonding
ⓓ. Hydrogen bonding
Correct Answer: Metallic bonding
Explanation: Aluminum atoms contribute electrons to a delocalized “sea,” giving high conductivity. At the same time, strong metallic bonds between positive ions and electron sea give structural strength, making aluminum both strong and conductive.
68. Which of the following compounds contains polar covalent bonds?
ⓐ. $\text{Cl}_2$
ⓑ. $\text{NaBr}$
ⓒ. $\text{H}_2\text{O}$
ⓓ. $\text{MgO}$
Correct Answer: $\text{H}_2\text{O}$
Explanation: Water has O–H bonds where oxygen is more electronegative, leading to unequal electron sharing. This creates partial charges, making the bonds polar covalent. $\text{Cl}_2$ is non-polar covalent, $\text{NaBr}$ and $\text{MgO}$ are ionic.
69. Which of the following is NOT true for covalent bonds?
ⓐ. They involve electron sharing.
ⓑ. They usually occur between non-metal atoms.
ⓒ. They form ions in solid state.
ⓓ. They can be single, double, or triple bonds.
Correct Answer: They form ions in solid state.
Explanation: Covalent compounds consist of molecules with shared electrons, not ions. They generally exist as discrete molecules like $\text{O}_2, \text{H}_2\text{O}$. Ionic bonding, not covalent, forms crystalline lattices of ions.
70. Which of the following is an example of a metallic bond with a practical application?
ⓐ. Ductility of copper wires used in electricity
ⓑ. Solubility of sugar in water
ⓒ. High melting point of sodium chloride
ⓓ. Covalent bonding in graphite
Correct Answer: Ductility of copper wires used in electricity
Explanation: Copper’s metallic bonding allows free electron mobility (conductivity) and sliding of ions without breaking the metallic bond (ductility). This makes copper ideal for electrical wiring. The other options describe covalent or ionic bond properties.
71. Which type of intermolecular force is primarily responsible for the high boiling point of water?
ⓐ. Van der Waals (London dispersion) forces
ⓑ. Dipole-dipole interactions
ⓒ. Hydrogen bonding
ⓓ. Ionic interactions
Correct Answer: Hydrogen bonding
Explanation: Water molecules form strong hydrogen bonds due to attraction between the partially positive hydrogen of one molecule and the partially negative oxygen of another. These bonds require significant energy to break, giving water a higher boiling point compared to other molecules of similar size.
72. What type of intermolecular force exists between non-polar molecules like $\text{O}_2$ or $\text{N}_2$?
ⓐ. Hydrogen bonds
ⓑ. Dipole-dipole forces
ⓒ. London dispersion forces
ⓓ. Ionic bonds
Correct Answer: London dispersion forces
Explanation: Non-polar molecules lack permanent dipoles. Instead, temporary fluctuations in electron distribution create instantaneous dipoles, inducing weak attractions called London dispersion forces. These are the only forces present in noble gases and diatomic non-polar molecules.
73. Why does ethanol ($\text{C}_2\text{H}_5\text{OH}$) dissolve in water?
ⓐ. Only because of van der Waals forces
ⓑ. Due to ionic bonding
ⓒ. Due to hydrogen bonding between $-OH$ groups and water
ⓓ. Due to metallic bonding
Correct Answer: Due to hydrogen bonding between $-OH$ groups and water
Explanation: The hydroxyl group in ethanol forms hydrogen bonds with water molecules, allowing ethanol to mix completely with water. This intermolecular force explains miscibility. Van der Waals forces alone would not account for high solubility.
74. Which of the following has the strongest intermolecular forces?
ⓐ. $\text{H}_2$
ⓑ. $\text{CO}_2$
ⓒ. $\text{HF}$
ⓓ. $\text{CH}_4$
Correct Answer: $\text{HF}$
Explanation: Hydrogen fluoride forms strong hydrogen bonds due to the high electronegativity of fluorine. $\text{H}_2$ and $\text{CH}_4$ have only weak dispersion forces, and $\text{CO}_2$ is non-polar overall, so its intermolecular forces are weaker.
75. Which property of liquids is most influenced by intermolecular forces?
ⓐ. Atomic number
ⓑ. Viscosity
ⓒ. Nuclear stability
ⓓ. Ionization energy
Correct Answer: Viscosity
Explanation: Viscosity is the resistance of a liquid to flow. Strong intermolecular forces (e.g., hydrogen bonding in glycerol) increase viscosity. Liquids with weak intermolecular forces (like hexane) flow more easily. Atomic number and ionization energy are atomic properties, not molecular.
76. Why does table salt ($\text{NaCl}$) dissolve easily in water?
ⓐ. Water molecules form hydrogen bonds with each other only.
ⓑ. Water molecules orient such that their dipoles stabilize $\text{Na}^+$ and $\text{Cl}^-$ ions.
ⓒ. Water molecules lose polarity in salt solution.
ⓓ. Salt breaks into covalent molecules.
Correct Answer: Water molecules orient such that their dipoles stabilize $\text{Na}^+$ and $\text{Cl}^-$ ions.
Explanation: The polar nature of water allows its partially negative oxygen to surround $\text{Na}^+$ and partially positive hydrogen to surround $\text{Cl}^-$, stabilizing ions and dissolving salt. This is an ion-dipole interaction, a strong type of intermolecular force.
77. Which of the following explains why iodine ($\text{I}_2$) is solid at room temperature while chlorine ($\text{Cl}_2$) is a gas?
ⓐ. Iodine has stronger London dispersion forces due to more electrons.
ⓑ. Chlorine forms hydrogen bonds.
ⓒ. Chlorine has stronger metallic bonds.
ⓓ. Iodine has covalent bonds between molecules.
Correct Answer: Iodine has stronger London dispersion forces due to more electrons.
Explanation: Larger atoms like iodine have more electrons, leading to stronger instantaneous dipoles and stronger London dispersion forces. This makes iodine a solid, while chlorine, with weaker forces, remains gaseous at room temperature.
78. Which of the following best describes dipole–dipole forces?
ⓐ. Attraction between instantaneous dipoles
ⓑ. Attraction between permanent partial charges on polar molecules
ⓒ. Electrostatic attraction between ions
ⓓ. Sharing of electrons
Correct Answer: Attraction between permanent partial charges on polar molecules
Explanation: Dipole–dipole forces arise when polar molecules with permanent dipoles attract each other. For example, HCl molecules align so that positive H atoms attract negative Cl atoms. Dispersion forces involve temporary dipoles, while ionic bonds involve full charges.
79. Which property of water is NOT directly explained by hydrogen bonding?
ⓐ. High surface tension
ⓑ. High boiling point
ⓒ. Expansion on freezing
ⓓ. Electrical conductivity in liquid state
Correct Answer: Electrical conductivity in liquid state
Explanation: Hydrogen bonding explains high boiling point, surface tension, and expansion upon freezing (open lattice structure in ice). Conductivity, however, is due to ionization of dissolved salts or impurities, not hydrogen bonding itself.
80. Which of the following liquids has the highest surface tension?
ⓐ. Ether
ⓑ. Benzene
ⓒ. Water
ⓓ. Alcohol
Correct Answer: Water
Explanation: Water’s extensive hydrogen bonding between molecules results in high cohesion and strong surface tension. Ether, benzene, and alcohol have weaker intermolecular forces, so their surface tensions are lower than water’s.
81. Which of the following assumptions is true for an ideal gas but not for a real gas?
ⓐ. Gas molecules are in constant random motion.
ⓑ. Gas molecules occupy negligible volume compared to container volume.
ⓒ. Collisions between molecules are elastic.
ⓓ. Pressure arises from molecular collisions with walls.
Correct Answer: Gas molecules occupy negligible volume compared to container volume.
Explanation: Ideal gases are assumed to have point particles with negligible volume. Real gases, however, have finite molecular size, so at high pressure or low volume this assumption fails. The other assumptions are approximately true for both.
82. Which equation represents the behavior of an ideal gas?
ⓐ. $PV = nRT$
ⓑ. $PV = k$
ⓒ. $P + \frac{a}{V^2}$($V – b$) = $nRT$
ⓓ. $P \propto T$
Correct Answer: $PV = nRT$
Explanation: The ideal gas law combines Boyle’s, Charles’s, and Avogadro’s laws into one relation, where $R$ is the universal gas constant. Option C is Van der Waals equation (real gas correction), while B and D are partial forms.
83. Why do real gases deviate from ideal gas behavior at high pressure?
ⓐ. Molecular collisions stop.
ⓑ. Molecular volume and intermolecular forces become significant.
ⓒ. Temperature becomes zero.
ⓓ. Pressure no longer depends on collisions.
Correct Answer: Molecular volume and intermolecular forces become significant.
Explanation: At high pressures, molecules are close together, so their finite volume and mutual attractions affect behavior. Ideal gas assumptions break down, requiring corrections like the Van der Waals equation.
84. At what conditions do real gases behave most like ideal gases?
ⓐ. Low pressure and high temperature
ⓑ. High pressure and low temperature
ⓒ. Low pressure and low temperature
ⓓ. High pressure and high temperature
Correct Answer: Low pressure and high temperature
Explanation: At low pressure, intermolecular distances are large, reducing volume effects. At high temperature, kinetic energy dominates over attractive forces. Thus, real gases approximate ideal gas behavior under these conditions.
85. Which scientist proposed corrections to the ideal gas law to account for real gas behavior?
ⓐ. Avogadro
ⓑ. Van der Waals
ⓒ. Maxwell
ⓓ. Dalton
Correct Answer: Van der Waals
Explanation: Johannes Diderik Van der Waals introduced the equation $(P + \tfrac{a}{V^2})(V – b) = nRT$, correcting for intermolecular attractions (a) and finite molecular volume (b). This explained deviations of real gases from ideal behavior.
86. Which term in the Van der Waals equation accounts for molecular attractions?
ⓐ. $V – b$
ⓑ. $P + \frac{a}{V^2}$
ⓒ. $nRT$
ⓓ. $V^2$
Correct Answer: $P + \frac{a}{V^2}$
Explanation: The factor $\tfrac{a}{V^2}$ is added to pressure to account for intermolecular attractive forces, which reduce observed pressure in real gases compared to ideal gases. The term $V – b$ corrects for finite molecular size.
87. Which of the following gases is closest to ideal gas behavior under normal conditions?
ⓐ. Helium
ⓑ. Water vapor
ⓒ. Ammonia
ⓓ. Carbon dioxide
Correct Answer: Helium
Explanation: Helium has very small atomic size and extremely weak intermolecular forces, making it behave nearly ideally. Water vapor and ammonia have strong hydrogen bonding, while $\text{CO}_2$ shows significant deviations at high pressure.
88. Which of the following is NOT an assumption of kinetic theory for an ideal gas?
ⓐ. Gas molecules move in random directions with various speeds.
ⓑ. Intermolecular forces are negligible.
ⓒ. Collisions between molecules are inelastic.
ⓓ. Average kinetic energy is proportional to temperature.
Correct Answer: Collisions between molecules are inelastic.
Explanation: In ideal gas theory, collisions are assumed perfectly elastic. Inelastic collisions would result in energy loss and contradict the principle of conservation of kinetic energy, which is fundamental in kinetic theory.
89. Why does compressibility factor $Z = \frac{PV}{nRT}$ deviate from 1 for real gases?
ⓐ. Due to high molecular velocities
ⓑ. Due to intermolecular attractions and finite volume
ⓒ. Due to errors in measurement only
ⓓ. Because Avogadro’s law is invalid
Correct Answer: Due to intermolecular attractions and finite volume
Explanation: For ideal gases, $Z = 1$. For real gases, $Z < 1$ at moderate pressure (dominant attractions) and $Z > 1$ at very high pressure (dominant repulsions due to finite volume). This deviation quantifies non-ideal behavior.
90. Why does liquefaction of gases occur in real gases but not in ideal gases?
ⓐ. Ideal gases have infinite energy.
ⓑ. Real gases have intermolecular attractions that allow condensation.
ⓒ. Ideal gases contain no particles.
ⓓ. Liquefaction is unrelated to molecular interactions.
Correct Answer: Real gases have intermolecular attractions that allow condensation.
Explanation: In real gases, attractive forces between molecules cause them to come together at low temperatures and high pressures, resulting in liquefaction. Ideal gas assumptions exclude intermolecular forces, so liquefaction cannot be explained in the ideal model.
91. In the Van der Waals equation $(P + \frac{a}{V^2})(V – b) = nRT$, what does the constant $a$ represent?
ⓐ. Correction for molecular volume
ⓑ. Correction for molecular attraction
ⓒ. Gas constant value
ⓓ. Boltzmann’s constant
Correct Answer: Correction for molecular attraction
Explanation: The term $\frac{a}{V^2}$ corrects for intermolecular attractions that reduce observed pressure in real gases. Without this correction, the pressure calculated by the ideal gas law would be higher than actual observed pressure.
92. In the same Van der Waals equation, what does the constant $b$ represent?
ⓐ. Correction for intermolecular attraction
ⓑ. Correction for finite molecular volume
ⓒ. Correction for compressibility factor
ⓓ. Correction for entropy
Correct Answer: Correction for finite molecular volume
Explanation: The constant $b$ is subtracted from the volume to account for the finite size of gas molecules. This “excluded volume” represents the space unavailable for molecular motion, which is ignored in the ideal gas law.
93. At very high pressure, which effect dominates the deviation from ideal gas behavior?
ⓐ. Attractive forces between molecules
ⓑ. Repulsive forces due to finite molecular size
ⓒ. Zero kinetic energy of molecules
ⓓ. Constant entropy of the system
Correct Answer: Repulsive forces due to finite molecular size
Explanation: At very high pressures, molecules are packed closely together, so the finite volume correction (b) becomes significant. Repulsive interactions dominate, causing observed pressure to be greater than predicted by the ideal gas law.
94. At intermediate pressures, real gases often show compressibility factor $Z < 1$. What does this indicate?
ⓐ. Repulsive forces dominate.
ⓑ. Attractive forces dominate.
ⓒ. Molecules have zero volume.
ⓓ. The gas is ideal.
Correct Answer: Attractive forces dominate.
Explanation: When attractions between molecules reduce the effective pressure, $Z < 1$. This means the gas is more compressible than an ideal gas. At high pressures, repulsions make $Z > 1$.
95. The critical temperature of a gas is defined as:
ⓐ. The lowest temperature at which gas can be liquefied by pressure.
ⓑ. The highest temperature at which solid melts.
ⓒ. The temperature where intermolecular forces vanish.
ⓓ. The boiling point at standard pressure.
Correct Answer: The lowest temperature at which gas can be liquefied by pressure.
Explanation: Above the critical temperature, no amount of pressure can liquefy the gas, because molecular kinetic energy exceeds intermolecular attraction. Van der Waals equation helps explain this concept quantitatively.
96. Which of the following gases has the largest Van der Waals constant $a$?
ⓐ. Helium
ⓑ. Hydrogen
ⓒ. Ammonia
ⓓ. Neon
Correct Answer: Ammonia
Explanation: The constant $a$ depends on the strength of intermolecular attractions. Ammonia has strong hydrogen bonding, so its $a$ value is high. Helium and neon are noble gases with very weak attractions, so their $a$ values are small.
97. Why does Van der Waals equation reduce to the ideal gas equation at low pressure and high temperature?
ⓐ. $a$ and $b$ both become zero.
ⓑ. Corrections for attraction and volume become negligible compared to $P$ and $V$.
ⓒ. Molecular mass becomes infinite.
ⓓ. Kinetic energy becomes zero.
Correct Answer: Corrections for attraction and volume become negligible compared to $P$ and $V$.
Explanation: At low pressure, molecules are far apart, so volume and attractive corrections are insignificant. At high temperature, kinetic energy dominates over attractions. Hence, the gas approaches ideal behavior.
98. Which of the following explains why carbon dioxide liquefies more easily than oxygen?
ⓐ. Lower critical pressure
ⓑ. Larger intermolecular attractions (higher $a$)
ⓒ. Zero molecular volume
ⓓ. Greater compressibility factor
Correct Answer: Larger intermolecular attractions (higher $a$)
Explanation: Carbon dioxide molecules have stronger intermolecular forces than oxygen, reflected in a higher Van der Waals constant $a$. Stronger attractions facilitate liquefaction under moderate conditions.
99. The isotherms of real gases on a $P$-$V$ graph deviate from ideal gas isotherms because:
ⓐ. Ideal gas molecules have higher molecular weight.
ⓑ. Real gases undergo phase transitions due to attractions.
ⓒ. Ideal gases expand indefinitely.
ⓓ. Real gases have zero entropy.
Correct Answer: Real gases undergo phase transitions due to attractions.
Explanation: Van der Waals isotherms show “loops” in $P$-$V$ diagrams, corresponding to condensation of gas into liquid. Ideal gas law predicts smooth curves without accounting for phase transitions.
100. The Van der Waals constants $a$ and $b$ are determined experimentally for gases. Which of the following units are correct for $a$ and $b$?
ⓐ. $a = \text{L}^2\text{atm mol}^{-2}, \, b = \text{L mol}^{-1}$
ⓑ. $a = \text{atm L}^2 \text{mol}^{-2}, \, b = \text{L mol}^{-1}$
ⓒ. $a = \text{atm L mol}^{-2}, \, b = \text{mol L}^{-1}$
ⓓ. $a = \text{atm L mol}^{-1}, \, b = \text{mol}^{-1}$
Correct Answer: $a = \text{atm L}^2 \text{mol}^{-2}, \, b = \text{L mol}^{-1}$
Explanation: From Van der Waals equation, $a$ corrects pressure and has units of pressure × volume$^2$/mole$^2$. $b$ corrects volume and has units of volume/mole. These units ensure dimensional consistency in the equation.
Welcome to Class 11 Physics MCQs – Chapter 13: Kinetic Theory (Part 1). This page is a chapter-wise question bank for the NCERT/CBSE Class 11 Physics syllabus—built for quick revision and exam speed. Practice MCQs / objective questions / Physics quiz items with solutions and explanations, ideal for CBSE Boards, JEE Main, NEET, competitive exams, and Board exams.
These MCQs are suitable for international competitive exams—physics concepts are universal.
Navigation & pages: The full chapter has 498 MCQs in 5 parts (100 + 100 + 100 + 100 + 98). Part 1 contains 100 MCQs split across 10 pages—you’ll see 10 questions per page. Use the page numbers above to view the remaining questions.
What you will learn & practice
- Introduction to Kinetic Theory and the molecular nature of matter
- Behaviour of gases and the kinetic theory of an ideal gas
- Law of equipartition of energy; specific heat capacity (molar & specific)
- Mean free path, collision theory, and transport phenomena (diffusion, viscosity, thermal conductivity)
- Brownian motion and microscopic interpretation of temperature
- RMS/average/most probable speeds and speed distributions (overview)
- Avogadro’s law, ideal gas equation, and pressure of a gas from molecular impacts
- Statistical mechanics (intro), phase transitions (qualitative)
- Thermal conductivity & viscosity of gases (dependence and trends)
How this practice works
- Click an option to check instantly: green dot = correct, red icon = incorrect. The Correct Answer and brief Explanation then appear.
- Use the 👁️ Eye icon to reveal the answer with explanation without guessing.
- Use the 📝 Notebook icon as a temporary workspace while reading (notes are not saved).
- Use the ⚠️ Alert icon to report a question if you find any mistake—your message reaches us instantly.
- Use the 💬 Message icon to leave a comment or start a discussion for that question.
Real value: Strictly aligned to NCERT/CBSE topics, informed by previous-year paper trends, and written with concise, exam-oriented explanations—perfect for one-mark questions, quick concept checks, and last-minute revision.
The full chapter includes 498 solved multiple-choice questions in 5 parts. This page (Part 1) gives the first 100 MCQs with answers and explanations to build your foundations. Explore more with: Class 11 Physics – Chapter Index, Class 11 – Subject Index, All Classes – MCQ Library.
- 👉 Total MCQs in this chapter: 498 (100 + 100 + 100 + 100 + 98)
- 👉 This page: first 100 multiple-choice questions with answers & brief explanations (in 10 pages)
- 👉 Best for: Boards • JEE/NEET • chapter-wise test • one-mark revision • quick Physics quiz
- 👉 Next: use the Part buttons and the page numbers above to continue
FAQs On Kinetic Theory ▼
▸ What are Kinetic Theory MCQs in Class 11 Physics?
These are multiple-choice questions from Chapter 13 of NCERT Class 11 Physics – Kinetic Theory. They test understanding of gases, molecular motion, mean free path, pressure of gases, and assumptions of kinetic theory.
▸ How many MCQs are available in this chapter?
There are a total of 498 MCQs from Kinetic Theory. They are divided into 5 sets of 100 questions each, with the last set containing 98 questions.
▸ Are these MCQs useful for NCERT, CBSE, and state board exams?
Yes, these MCQs are directly based on NCERT and state board Class 11 Physics syllabus, making them highly useful for CBSE and state-level board exams.
▸ Are Kinetic Theory MCQs important for JEE and NEET?
Yes, Kinetic Theory is an important chapter for JEE and NEET. Questions on gas laws, kinetic energy of molecules, RMS speed, and molecular collisions are frequently asked in these exams.
▸ Do these MCQs include answers and explanations?
Yes, every MCQ comes with the correct answer and explanations wherever necessary. This ensures conceptual clarity and helps students understand the logic behind answers.
▸ Who should practice Kinetic Theory MCQs?
These MCQs are helpful for Class 11 students, CBSE/State board aspirants, and students preparing for competitive exams like JEE, NEET, NDA, UPSC, and other entrance tests requiring strong physics fundamentals.
▸ Can I practice these Kinetic Theory MCQs online for free?
Yes, all Kinetic Theory MCQs on GK Aim are available online for free. Students can practice anytime using mobile, tablet, or desktop.
▸ Are these MCQs helpful for quick revision before exams?
Yes, solving these MCQs regularly helps in quick revision, improves memory recall, and boosts exam performance by enhancing accuracy and speed in numerical problem-solving.
▸ Do these MCQs cover both basics and advanced concepts?
Yes, the MCQs cover both fundamental topics like Boyle’s Law, Charles’s Law, and Avogadro’s hypothesis as well as advanced concepts like mean free path, pressure due to molecular collisions, and transport phenomena.
▸ Which subtopics are included in Kinetic Theory MCQs?
These MCQs include subtopics such as assumptions of kinetic theory, equation of state of an ideal gas, pressure of a gas from molecular motion, mean free path, and specific heat capacities of gases.
▸ Why are the 498 MCQs divided into 5 parts?
The MCQs are divided into 5 sets to make practice more structured and manageable, allowing students to study step by step without feeling overloaded.
▸ Can teachers and coaching institutes use these MCQs?
Yes, teachers and coaching institutes can use these MCQs as ready-made assignments, quizzes, and practice material for classroom teaching and test preparation.
▸ Are these MCQs mobile-friendly?
Yes, the Kinetic Theory MCQs pages are fully optimized for smartphones and tablets, so students can practice anytime, anywhere.
▸ Can I download or save Kinetic Theory MCQs for offline study?
Yes, you can download these Kinetic Theory MCQs in PDF format for offline study. Please visit our website shop.gkaim.com