Part II. Multiple choice, continued (3 points each). Name:__________________ Please circle your answer. There is only one correct answer to each question. 52. A mixture of He and Ne at a total pressure of 0.95 atm is found to contain 0.32 mol of He and 0.56 mol of Ne. The partial pressure of Ne is __________ atm. A) 0.60 B) 1.5 C) 1.0 D) 0.35 53. Of the following gases, _____________ will have the greatest rate of effusion at a given temperature. A) NH3 B) CH4 C) Ar D) HCl 54. Which noble gas is expected to show the largest deviations from the ideal gas behavior? A) helium B) neon C) argon D) krypton 55. Together, liquids and solids constitute _______ phases of matter. A) the compressible B) the fluid C) the condensed D) the disordered 56. What is the predominant intermolecular force in liquid CBr4? A) dispersion B) dipole-dipole C) ion-dipole D) hydrogen bonds 57. Of the following substances, ____________ has the highest boiling point. A) H2O B) CO2 C) CH4 D) Kr 58. Of the following substances, ____________ has the lowest boiling point. A) Kr B) Ar C) Ne D) Xe 59. What is the predominant intermolecular force in liquid HBr? A) dispersion B) dipole-dipole C) ion-dipole D) hydrogen bonds 60. The total pressure of a mixture of O2 and N2 is 730 torr. The partial pressure of N2 is 0.78 atm. The partial pressure of O2 is A) 0.18 atm B) 0.22 atm C) 0.26 atm D) 0.78 atm

Here's a breakdown of the concepts behind the multiple-choice questions, along with the answers:

52. A) 0.60 atm

    • Concept: Dalton's Law of Partial Pressures. P{total} = P1 + P2 + … + Pn. The partial pressure of a gas is also proportional to its mole fraction in the mixture: Pi = Xi * P{total}, where Xi is the mole fraction of gas i.
    • Solution:
      • Total moles = 0.32 mol (He) + 0.56 mol (Ne) = 0.88 mol
      • Mole fraction of Ne, X_{Ne} = 0.56 mol / 0.88 mol ≈ 0.636
      • P_{Ne} = 0.636 * 0.95 atm ≈ 0.60 atm

53. B) CH4

    • Concept: Graham's Law of Effusion. The rate of effusion of a gas is inversely proportional to the square root of its molar mass. Lighter gases effuse faster.
    • Solution:
      • NH3 (17 g/mol), CH4 (16 g/mol), Ar (40 g/mol), HCl (36.5 g/mol)
      • CH4 has the lowest molar mass, so it will effuse the fastest.

54. D) krypton

    • Concept: Ideal Gas Law Deviations. Real gases deviate from ideal behavior at high pressures and low temperatures. Larger, more complex molecules with greater intermolecular forces deviate more.
    • Solution:
      • Larger noble gases have stronger London dispersion forces, leading to greater deviations. Krypton is the largest among the options.

55. C) the condensed

    • Concept: Phases of Matter. Liquids and solids are condensed phases because their particles are close together compared to gases.
    • Solution:
      • Liquids and solids have high densities and low compressibility due to close intermolecular distances.

56. A) dispersion

    • Concept: Intermolecular Forces. CBr4 is a nonpolar molecule, so the predominant force is London dispersion forces.
    • Solution:
      • CBr4 is symmetrical, so the individual bond dipoles cancel out, making the molecule nonpolar.

57. A) H2O

    • Concept: Boiling Point and Intermolecular Forces. Higher intermolecular forces lead to higher boiling points.
    • Solution:
      • H2O can form hydrogen bonds, which are stronger than the dipole-dipole forces in CO2 and Kr and the dispersion forces in CH4.

58. C) Ne

    • Concept: Boiling Point and Intermolecular Forces. For noble gases, boiling point increases with increasing atomic size due to stronger London dispersion forces.
    • Solution:
      • Ne < Ar < Kr < Xe in terms of size and thus boiling point.

59. B) dipole-dipole

    • Concept: Intermolecular Forces. HBr is a polar molecule, so the predominant force is dipole-dipole.
    • Solution:
      • HBr has a significant difference in electronegativity between H and Br, creating a dipole moment.

60. A) 0.18 atm

    • Concept: Dalton's Law and Unit Conversions. Convert torr to atm, then subtract the partial pressure of N2 to find the partial pressure of O2.
    • Solution:
      • Total pressure in atm: 730 torr / 760 torr/atm ≈ 0.96 atm
      • P{O2} = P{total} - P_{N2} = 0.96 atm - 0.78 atm = 0.18 atm

Dalton's Law of Partial Pressures in Depth:

Dalton's Law of Partial Pressures states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases. The partial pressure of a gas is the pressure that the gas would exert if it occupied the same volume alone.

Mathematically, this is expressed as:
P{total} = P1 + P2 + P3 + … + P_n
Where:

• P_{total} is the total pressure of the gas mixture.
• P1, P2, P3, …, Pn are the partial pressures of the individual gases in the mixture.

Key Points:

1. Ideal Gas Behavior: Dalton's Law assumes that the gases in the mixture behave ideally, meaning there are no significant intermolecular interactions between the gas molecules.
2. Non-reacting Gases: The gases in the mixture should not chemically react with each other. If a reaction occurs, the law is not applicable without adjustments.
3. Mole Fraction: The partial pressure of a gas in a mixture is related to its mole fraction. The mole fraction Xi of a gas i in a mixture is given by: Xi = \frac{ni}{n{total}}
   Where:

• n_i is the number of moles of gas i.
• n_{total} is the total number of moles of gas in the mixture.

The partial pressure of gas i can then be calculated as:
Pi = Xi * P_{total}

Applications:

• Calculating Total Pressure: If you know the partial pressures of all gases in a mixture,