States of matter practice · Mixed difficulty

Option A ($10\%$) is correct because, according to Charles's Law, at constant pressure, volume is directly proportional to temperature in Kelvin.

Option A (4.0 bar) is correct because it applies Boyle's Law to calculate the pressure required to compress the gas.

Option D is correct because Statement I incorrectly defines surface tension as force per unit area, while Statement II correctly states that liquids tend to minimize their surface area.

Option D (Increase but to different extent) is correct because the volume increase depends on the initial volume and the thermal expansion coefficient, which varies for different gases.

Option C $\left( {p + \frac{a}{{{v^2}}}} \right)$ is correct because the term $\frac{a}{v^2}$ in the van der Waals equation accounts for the intermolecular forces of attraction.

The correct answer is $\frac{{2{p_1}{T_1}}}{{{T_1} + {T_2}}}$. The final pressure is derived by considering the conservation of moles and applying the ideal gas law.

Option C (Mixture of gases) is correct because Van der Waals' forces are intermolecular forces that exist between all types of molecules, including mixtures of gases.

Option B (6:7) is correct because the average kinetic energy is directly proportional to the absolute temperature.

Option C (270 ml) is correct because the final volume is calculated using Charles's Law, which relates volume and temperature at constant pressure.

Option D ($\sqrt 3 \,\,times$) is correct because the root mean square speed is proportional to the square root of the absolute temperature.

Option C is correct. The assertion is correct because oxygen has a higher molar mass and thus a smaller effusion rate. The reason is incorrect because nitrogen is smaller than oxygen.

Option B (Less volume than that of an ideal gas) is correct because at high pressure, the intermolecular forces of attraction become significant, reducing the volume compared to an ideal gas.

Option B ($7.3 \times {10^4}\mathrm{erg}$) is correct because work done is the product of surface tension and the change in surface area.

Option A ($> 1$) is correct because the compressibility factor for $H_2$ is usually greater than 1 at moderate pressures due to repulsive forces dominating.

Option D ($\frac{{283}}{{273}}{V_0}$) is correct because it applies Charles's Law to calculate the volume at 10°C based on the volume at 273 K.

The correct answer is 1/4. The fraction of air that goes out is due to the expansion of air when heated at constant pressure, and can be calculated using Charles's Law.

Option C is correct because, using the ideal gas law and the given density, temperature, and pressure, the calculated molar mass is approximately 44 g/mol, which corresponds to CO2.

Option C (twice that of $O_2$) is correct because at the same conditions (STP), equal volumes contain equal moles, and the weight is proportional to the molar mass. $SO_2$ has a molar mass twice that of $O_2$.

Option B (Both Assertion and Reason are true, but Reason is not the correct explanation for Assertion) is correct because both Dalton's law and Raoult's law deal with partial pressures, but they apply to different systems, and the reason doesn't explain why Dalton's law is a special case of Raoult's law.

Option C is correct because the kinetic theory of gases relates the average kinetic energy of molecules to temperature, not critical temperature.

Option B (Dipole forces and van der Waals’ forces) is correct because hydrogen is a nonpolar molecule, and the forces responsible for its liquefaction are dipole-induced dipole interactions and London dispersion forces (van der Waals forces).

Option B ($P/3$) is correct because, according to Boyle's Law, pressure and volume are inversely proportional at constant temperature.

Option B ($28\,g\,mo{l^{ - 1}}$) is correct because it is calculated using the ideal gas law and the given density, pressure, and temperature.

Option D (Equal to atmospheric pressure minus aqueous tension at that temperature) is correct because the total pressure inside the jar equals atmospheric pressure, and it's the sum of oxygen pressure and water vapor pressure.

Option A (five times as heavy as O2) is correct because the mass of a gas is directly proportional to its molar mass at the same temperature and pressure, and CCl4's molar mass is approximately five times that of O2.

Option C is correct because high relative humidity reduces the rate of evaporation, leading to discomfort on a hot day.

Option B ($\mathrm{N}{\mathrm{H}_3} + \mathrm{HCl} + \mathrm{HBr}$) is correct because Dalton's law of partial pressures is not applicable when gases react chemically with each other, and NH3 and HCl react to form NH4Cl.

Option A is correct because, using the ideal gas law and the given concentrations, the volumes of O2 and CO2 are calculated to be approximately 1.511 L and 4.396 L, respectively.

Option B ($9.33\,\,atm$) is correct because it is calculated using the van der Waals equation of state, accounting for intermolecular forces and the finite volume of gas molecules.

Option D (Van der Waals’ forces) is correct because Van der Waals' forces are the primary attractive forces between non-polar molecules, arising from temporary fluctuations in electron distribution.