ap chemistry 2차 - unit 3

Correct Answer: (D) H₂O (water)
Explanation: Water can form multiple hydrogen bonds between its molecules, resulting in a much higher boiling point than substances held together only by weaker forces (dipole-dipole in NH₃, or London dispersion in CH₄ and CO₂).

Correct Answer: (B) London dispersion forces
Explanation: London dispersion forces (induced dipole-induced dipole attractions) occur between all molecules. Even nonpolar molecules and single atoms experience these weak forces due to temporary fluctuations in their electron clouds.

Correct Answer: (B) HCl (hydrogen chloride)
Explanation: HCl is a polar molecule (H–Cl bond is polar and the molecule is linear with a dipole moment). Polar molecules like HCl attract each other via dipole-dipole forces. CO₂ and CH₄ are nonpolar (so only LDF), and C₂H₄ is also nonpolar, so none of those have dipole-dipole attractions.

Correct Answer: (C) CH₃OH (methanol)
Explanation: Hydrogen bonding requires a hydrogen atom bonded to a highly electronegative atom (N, O, or F) and an unshared electron pair on a N/O/F in a neighboring molecule. Methanol (CH₃OH) has an O–H group, so it can hydrogen bond. H₂S and CH₃Cl do not have H bonded to N/O/F, and CH₄ has no polar bonds.

Correct Answer: (B) H₂O and O₂
Explanation: Dipole–induced dipole forces occur when a polar molecule induces a temporary dipole in a nonpolar molecule. Here H₂O is polar and can induce a dipole in O₂ (nonpolar). In (A) both are nonpolar (only LDF), (C) both are polar (primarily dipole-dipole or ion-dipole in solution), (D) both are polar (dipole-dipole).

Correct Answer: (A) Hydrogen bonding
Explanation: Water (H₂O) molecules can form multiple hydrogen bonds with each other. These strong intermolecular hydrogen bonds give water a high surface tension (and high boiling point) relative to molecules of similar size. London forces and dipole interactions exist in water but are much weaker; metallic bonding is not applicable to molecular liquids.

Correct Answer: (B) The strength of London forces increases as the size and electron count of a molecule increase.
Explanation: London dispersion forces are induced dipole attractions that become stronger with greater polarizability. Larger atoms/molecules with more electrons (and more diffuse electron clouds) are more polarizable, leading to stronger LDF. LDF act between all molecules (polar or not), are much weaker than covalent bonds, and do not specifically involve hydrogen.

Correct Answer: (D) Xe (xenon)
Explanation: Xenon has the largest atomic size and the most electrons among the choices, making its electron cloud very polarizable. This results in the strongest dispersion forces. Down a group (or with increasing molar mass), London dispersion forces increase: Ne < Ar < Kr < Xe.

Correct Answer: (A) London dispersion forces
Explanation: London dispersion forces are the weakest of the intermolecular forces listed when comparing molecules of similar size. Dipole-dipole forces are moderate, hydrogen bonds are stronger, and ion-dipole interactions (between ions and polar molecules) are typically even stronger. (Note: Very large nonpolar molecules can have substantial LDF, but in general LDF are weakest.)

Correct Answer: (C) London dispersion < dipole-dipole < hydrogen bonding
Explanation: For molecules of comparable size, London dispersion forces are weakest, dipole-dipole forces are stronger, and hydrogen bonds are the strongest of these three. (Ionic and ion-dipole forces can be even stronger but are not in this list.)

Correct Answer: (D) Covalent bond
Explanation: A covalent bond is an intramolecular force (it holds atoms together within a molecule). The others (dipole-dipole, hydrogen bonding, and London dispersion) are intermolecular forces, which occur between separate molecules.

Correct Answer: (D) I₂
Explanation: Iodine (I₂) is the largest and heaviest of the halogen molecules listed and thus has the strongest London dispersion forces. As a result, it has the highest boiling point. (At room temperature, I₂ is a solid, Br₂ is a liquid, and Cl₂ and F₂ are gases, consistent with their boiling points increasing down the group.)

Correct Answer: (C) Atomic mass
Explanation: Boiling point and melting point are higher when intermolecular forces are stronger, and vapor pressure is lower when intermolecular forces are stronger (since fewer molecules can escape to the vapor phase). Atomic mass is a property of an atom (the sum of protons and neutrons) and is not influenced by intermolecular forces.

Correct Answer: (C) It increases.
Explanation: For a fixed amount of gas at constant volume, pressure is directly proportional to absolute temperature (Gay-Lussac’s law). Heating the gas gives molecules more kinetic energy, so they collide with the container walls more frequently and forcefully, increasing the pressure.

Correct Answer: (C) The pressure doubles.
Explanation: By Boyle’s law (P ∝ 1/V at constant T and moles), reducing volume to 1/2 causes pressure to increase by a factor of 2. The gas molecules are in a smaller space, so they hit the container walls twice as often, doubling the pressure.

Correct Answer: (A) It will increase.
Explanation: At constant pressure, gas volume is directly proportional to temperature (Charles’s law). Heating the gas causes it to expand in order to maintain the same pressure (e.g., a balloon will expand when warmed).

Correct Answer: (B) 22.4 L
Explanation: At STP (standard temperature and pressure), 1 mole of an ideal gas occupies about 22.4 liters. This is a standard molar volume for ideal gases (derived from PV = nRT with P = 1 atm, T = 273 K).

Correct Answer: (C) Mole fraction in the mixture
Explanation: Dalton’s law of partial pressures states that Pᵢ = Xᵢ * P_total, where Xᵢ is the mole fraction of gas i. This means each gas’s partial pressure is proportional to the number of moles (and thus mole fraction) of that gas in the mixture. (It does not depend on the gas’s molecular weight or density; all gases in the same container have the same volume.)

Correct Answer: (C) 0.5 Ptotal
Explanation: O₂ makes up half of the moles in the container (0.50 mol out of 1.00 mol total). Therefore, its mole fraction X₍O₂₎ = 0.5, and its partial pressure is 0.5 * P_total (half of the total pressure). Similarly, N₂ would also have partial pressure 0.5 P_total in this case.

Correct Answer: (C) 600 torr
Explanation: Helium constitutes 2.0 out of 3.0 total moles of gas in the mixture, which is a mole fraction of 2/3. Its partial pressure is therefore (2/3) * 900 torr = 600 torr. (Argon would be the remaining 300 torr.)

Correct Answer: (C) Gas particles lose energy with each inelastic collision they undergo.
Explanation: The kinetic molecular theory (KMT) assumes that collisions between gas molecules (and with container walls) are perfectly elastic, meaning no net energy is lost. The other statements are all assumptions of KMT for ideal gases: gas particles have essentially no volume, they move randomly, and they do not attract or repel each other.

Correct Answer: (B) Molecules colliding with the walls of the container.
Explanation: In KMT, pressure is explained by the force of gas molecules colliding with container walls. The more frequent and forceful the collisions, the higher the pressure. Gas molecules are assumed not to exert significant repulsive or attractive forces on each other in an ideal gas. Gravity’s effect is negligible in a container, and chemical reactions are not related to pressure in an inert gas sample.

Correct Answer: (A) All molecular motion would cease.
Explanation: Absolute zero (0 K) is the temperature at which particles would have zero kinetic energy. According to KMT, as temperature approaches 0 K, molecular motion diminishes and at 0 K it would theoretically stop entirely. (In reality, gases condense to liquids/solids before reaching 0 K.)

Correct Answer: (A) Gas molecules are in constant random motion and travel in straight lines until they collide.
Explanation: Gases spread out to fill any container because their molecules move rapidly in all directions with no significant attractions holding them together. They bounce off each other and the walls, eventually distributing uniformly. There are no strong repulsive forces needed for this (B is incorrect), and gravity (C) does not cause uniform filling (it would concentrate gas at the bottom if anything). Gas particles do not expand in size (D).

Correct Answer: (D) Very high pressure
Explanation: At very high pressures, gas particles are forced much closer together. Under these conditions, the volume of the particles is no longer negligible (violating one KMT assumption) and intermolecular forces can become significant if distance is small. Low pressure and high temperature favor ideal behavior. Helium, being small and nonpolar, is closer to ideal (so (C) is not where KMT fails; it actually helps KMT hold true).

Correct Answer: (D) All gases have the same average kinetic energy per molecule.
Explanation: Temperature (in Kelvin) is directly proportional to the average kinetic energy of gas particles. Thus, at a given temperature, any ideal gas—regardless of its molar mass—has the same average kinetic energy per molecule. However, lighter molecules will move faster on average than heavier molecules (since KE = ½ m v²).

Correct Answer: (C) He
Explanation: At a given temperature, lighter gas particles move faster on average than heavier ones. Helium (M ≈ 4 g/mol) has a much lower molar mass than O₂ (~32), N₂ (~28), or CO₂ (~44), so He atoms have the highest average speed at 25°C.

Correct Answer: (A) Gas A molecules have a lower average speed than gas B molecules.
Explanation: If both gases are at the same temperature, they have the same average kinetic energy. The heavier gas (A) must therefore move more slowly on average than the lighter gas (B) to have the same kinetic energy. (B) is incorrect because their average kinetic energies are equal at the same T. (C) is not necessarily true unless moles/volume differ. (D) is opposite of the truth.)

Correct Answer: (A) It doubles.
Explanation: Average kinetic energy of gas particles (for an ideal gas) is directly proportional to the Kelvin temperature. If T is doubled, the average kinetic energy is doubled. (Note: Molecular speeds would increase by a factor of √2, since KE ∝ v².)

Correct Answer: (B) The Ne container
Explanation: 10.0 g of Ne (atomic mass ≈ 20) is about 0.50 mol, whereas 10.0 g of CO₂ (molar mass ≈ 44) is about 0.23 mol. At equal volume and temperature, the container with more moles of gas (Ne) will have the higher pressure (by ideal gas law, P = (nRT)/V). Thus, the neon container has roughly twice the pressure of the CO₂ container in this scenario.

Correct Answer: (A) 2
Explanation: The rms speed of gas molecules is proportional to the square root of the absolute temperature (v_rms ∝ √T). If the temperature is increased by a factor of 4 (from 300 K to 1200 K), then v_rms increases by a factor of √4 = 2 (doubles). (Note: average kinetic energy increases by factor of 4, but speed increases by factor of 2.)

Correct Answer: (A) Low pressure and high temperature
Explanation: Ideal gas behavior is approached when intermolecular forces and molecular volumes are negligible. Low pressure (molecules far apart) and high temperature (molecules moving fast, overcoming attractions) minimize interactions and make gas particles’ own volume relatively negligible. In contrast, high pressure and/or low temperature increase deviations from ideality.

Correct Answer: (C) He (helium)
Explanation: Helium is a small, light, nonpolar atom with very weak intermolecular forces (just very weak London forces). It comes closest to ideal gas behavior under standard conditions. In contrast, NH₃ and H₂O are polar and can hydrogen-bond (strong IMF), and CO₂, while nonpolar, has more electrons (larger van der Waals forces) and a larger size than He.

Correct Answer: (D) NH₃ (ammonia)
Explanation: Ammonia is polar and can hydrogen-bond, indicating strong intermolecular attractions. It will deviate more from ideal behavior (particularly at moderate pressures or low temperatures) than the other gases listed. Helium is the most ideal. N₂ and CH₄ are nonpolar and relatively small, so they are also fairly ideal compared to NH₃.

Correct Answer: (A) Intermolecular attractive forces between gas molecules
Explanation: The van der Waals “a” term is a correction for the attractive forces between gas molecules. These attractions cause the observed pressure to be lower than it would be for an ideal gas (molecules hit the walls less forcefully). The “b” term accounts for finite molecular volume, not “a”.

Correct Answer: (B) The volume occupied by the gas molecules themselves
Explanation: The “b” constant in the van der Waals equation adjusts for the finite volume of gas particles. The term (V – nb) effectively subtracts the volume excluded by the particles, since gas molecules are not point masses. (A) is addressed by the “a” term, (C) is simply n, and (D) high temperature deviations are minor and not specifically corrected by b.)

Correct Answer: (C) The finite volume of molecules effectively reduces the free space, so collisions occur more frequently than ideal.
Explanation: At extremely high pressures, gas molecules are crowded together and their own volume is no longer negligible compared to the container volume. This means the “free” volume available is less than V, causing more frequent collisions (and thus higher observed pressure) than predicted by PV = nRT. (A) is incorrect: attractive forces would lower pressure, not raise it. (B) Molecules’ speed is set by temperature, not pressure. (D) KMT assumes elastic collisions; inelastic collisions aren’t a major factor unless chemical reactions occur.)

Correct Answer: (A) Gas molecules attract each other, so they hit the walls less forcefully.
Explanation: When a gas is cooled (and moving slower), intermolecular attractive forces have a greater effect. Molecules tug on each other and don’t strike the container walls as hard or as often, reducing the observed pressure compared to an ideal gas with no attractions. The other statements are incorrect: molecules have less kinetic energy at low T, occupy slightly less space (if anything), and do not dissociate simply from cooling.

Correct Answer: (D) Real gases can condense into liquids when sufficiently cooled or compressed.
Explanation: Unlike an ideal gas, real gases do exhibit intermolecular forces, which means that at low temperatures or high pressures, the gas molecules can stick together and condense into a liquid. Ideal gases are an approximation that assumes no condensation ever occurs. Real gases do deviate from PV = nRT, especially under non-ideal conditions, and their molecules do occupy volume.

Correct Answer: (C) The peak of the distribution shifts to higher speeds and the curve broadens (flattens out).
Explanation: As temperature increases, the most probable speed (peak of the curve) moves to a higher value, and the distribution spreads out (more speeds become populated). The curve becomes lower and broader because a greater fraction of molecules have very high speeds. At lower temperature, the curve is taller and narrower with a lower most-probable speed.

Correct Answer: (D) H₂
Explanation: Hydrogen (H₂) has the lowest molar mass of the gases listed. At a given temperature, lighter gas molecules have a higher most probable speed. H₂ molecules on average move faster than Ne, N₂, or O₂ at the same temperature, so H₂’s speed distribution peak is at the highest speed.

Correct Answer: (A) The lighter gas has a broader (more spread-out) distribution of speeds.
Explanation: At a given temperature, a lighter gas (lower molar mass) will have a wider range of molecular speeds – its Maxwell-Boltzmann distribution is more spread out and shifted to higher speeds. A heavier gas has a narrower distribution centered at lower speeds. Thus, the lighter gas also has a greater fraction of molecules at very high speed compared to the heavier gas (contrary to option B). They are not identical (C), and the heavier gas’s peak (most probable speed) is actually lower, not higher (D).

Correct Answer: (A) v_mp < v_avg < v_rms
Explanation: For any gas at a given temperature: the most probable speed is the speed at the peak of the distribution (mode) and is the lowest of these three measures. The average speed is a bit higher than v_mp. The root-mean-square (rms) speed is slightly higher than the average speed. In general, v_mp < v_avg < v_rms for Maxwell-Boltzmann distributions.

Correct Answer: (B) It increases.
Explanation: At higher temperature, the Maxwell-Boltzmann distribution spreads out, and a larger proportion of molecules achieve higher speeds. So the fraction of molecules with very high speeds increases when going from 300 K to 600 K. (Not all molecules double their speed—speeds distribute statistically. The average speed increases by factor √2, and more molecules will have speeds above any given high threshold.)

Correct Answer: (C) The lighter gas has a larger fraction of molecules with very high speeds.
Explanation: At the same temperature, lighter gas molecules (lower molar mass) have a more spread-out speed distribution and a higher tail, meaning a greater proportion of molecules move at very high speeds compared to a heavier gas. The heavier gas’s distribution is narrower and centered at a lower speed (so A and B are incorrect). Therefore, the gases do not have the same high-speed fraction (D is wrong).

Correct Answer: (B) The volume remains unchanged.
Explanation: Using the combined gas law (P₁V₁/T₁ = P₂V₂/T₂) with P₂ = 2P₁ and T₂ = 2T₁, we get:
2P₁ * V₂ / (2T₁) = P₁ * V₁ / T₁ ⇒ V₂ = V₁. 
The effects of doubling temperature and doubling pressure cancel out, leaving the volume the same as initially.

Correct Answer: (C) Move in straight lines at constant velocity.
Explanation: KMT assumes gas molecules are free of forces between collisions, so they move in straight lines with uniform motion (constant speed and direction) until they collide with another molecule or the wall. There are no significant attractions or repulsions altering their paths (A is false), gravity is negligible on their microscopic motion (B is false), and without collisions or forces they don’t slow down (D is false).

Correct Answer: (C) (C₂H₅)₂O (diethyl ether)
Explanation: Diethyl ether has the weakest intermolecular forces among the choices, resulting in the highest volatility. It has a boiling point of about 35°C and evaporates readily at 25°C (high vapor pressure). Water and ethanol can hydrogen bond (strong IMF) and thus have much lower vapor pressures (water BP 100°C, ethanol 78°C). Hexane (nonpolar) has only LDF but is a larger molecule (BP ~69°C), so it’s less volatile than ether.

Correct Answer: (A) ~2 g/mol
Explanation: By Graham’s law, Rate₁/Rate₂ = √(M₂/M₁). If the unknown effuses 4 times faster than O₂, then:
4 = √(M_O2 / M_unknown). Square both sides: 16 = M_O2 / M_unknown. For O₂, M ≈ 32 g/mol. So 16 = 32 / M_unknown ⇒ M_unknown = 32/16 ≈ 2 g/mol. This is close to the molar mass of H₂ (~2 g/mol), consistent with hydrogen being extremely light and effusing very quickly.

Correct Answer: (D) CO₂
Explanation: At STP, 1 mole of any ideal gas occupies 22.4 L. Thus, gas density at STP is directly proportional to molar mass. Among the choices, CO₂ (≈44 g/mol) has the largest molar mass, so 1 mole (22.4 L) of CO₂ has the greatest mass and thus highest density. (He ~4 g/mol; N₂ ~28 g/mol; O₂ ~32 g/mol.)