Ethanol and hexane are mixed in equal mole fractions. Compared to what Raoult's law predicts, the observed total vapor pressure is higher. What best explains this positive deviation?
AMixing generates heat, which raises the temperature and therefore the vapor pressure above the predicted value
BThe ethanol-hexane interactions are weaker than the ethanol-ethanol and hexane-hexane self-interactions, so molecules escape the liquid more easily
CEthanol has a higher pure vapor pressure than hexane, pulling the mixture above Raoult's prediction
DPositive deviations occur whenever both components are polar, because polarity increases volatility
Positive deviations arise when cross-interactions (solute-solvent) are WEAKER than self-interactions (solute-solute and solvent-solvent). Pure ethanol is held by hydrogen bonds; in hexane, ethanol loses those partners and the weaker London-dispersion interactions with hexane don't compensate — molecules escape more readily than predicted. Option A (heat of mixing) is related but doesn't directly explain the direction of the vapor pressure deviation. Option C misunderstands the law — Raoult's law already accounts for pure vapor pressures via mole fraction weighting. Option D is wrong; positive deviation is about weakened interactions, not polarity per se.
Question 2 Multiple Choice
A solution contains benzene (mole fraction 0.4, P° = 75 mmHg) and toluene (mole fraction 0.6, P° = 25 mmHg). Assuming ideal behavior, what is the total vapor pressure above this solution?
A50 mmHg — the simple average of the two pure vapor pressures
B45 mmHg — calculated as (0.4 × 75) + (0.6 × 25)
C75 mmHg — dominated by the more volatile component
D100 mmHg — the sum of the two pure vapor pressures
Raoult's law: P_total = χ_A × P°_A + χ_B × P°_B = (0.4)(75) + (0.6)(25) = 30 + 15 = 45 mmHg. Each component contributes its partial pressure proportional to its mole fraction in the liquid. Option A (simple average) ignores the mole fraction weighting. Option C confuses the total pressure with the pure vapor pressure of the more volatile component. Option D would require the two pure vapor pressures to simply add, which would only be true if both components had mole fraction 1 simultaneously — impossible.
Question 3 True / False
A non-volatile solute that dissociates into two ions will lower the vapor pressure of a solvent approximately twice as much as a non-dissociating solute at the same molality.
TTrue
FFalse
Answer: True
Vapor pressure lowering is a colligative property — it depends on the number of solute particles, not their chemical identity. An ionic solute like NaCl dissociates into Na⁺ and Cl⁻, doubling the number of particles compared to a molecular solute at the same molality. Since more particles occupy surface sites and reduce the solvent's mole fraction, twice as many particles produce roughly twice the vapor pressure lowering. This same principle gives ionic solutes larger boiling point elevation and freezing point depression effects than molecular solutes at equal molality.
Question 4 True / False
Raoult's law is accurate for any mixture of two liquids, as long as the mole fractions are correctly calculated.
TTrue
FFalse
Answer: False
Raoult's law applies strictly to ideal solutions, where solute-solvent interactions are essentially identical to the self-interactions of each pure component. In practice, this requires chemically similar components (e.g., benzene and toluene, or hexane and heptane). Most real solutions deviate — positively when cross-interactions are weaker, negatively when they are stronger. The law is an approximation, not a universal rule; whether a mixture is 'ideal' always depends on the similarity of intermolecular forces between the components.
Question 5 Short Answer
Acetone and chloroform form a solution with lower-than-expected vapor pressure (negative deviation from Raoult's law). Using intermolecular force reasoning, explain why.
Think about your answer, then reveal below.
Model answer: Acetone's carbonyl oxygen (C=O) accepts a hydrogen bond from chloroform's acidic C-H (made acidic by the three adjacent electronegative Cl atoms). This C-H···O=C cross-interaction is stronger than either pure component's self-interactions (acetone has only dipole-dipole; chloroform has only weak C-H interactions with itself). Stronger cross-interactions hold molecules in solution more tightly, reducing their tendency to escape into the vapor phase — lowering vapor pressure below Raoult's prediction.
The sign of the deviation is a direct readout of relative interaction strengths: stronger cross-interactions → negative deviation (lower vapor pressure than ideal); weaker cross-interactions → positive deviation (higher vapor pressure than ideal). Acetone-chloroform is the textbook example of a hydrogen-bond-driven negative deviation. This reasoning — compare cross-interactions to self-interactions — is the general rule for predicting deviation direction from intermolecular forces.