Properties of Substances and Mixtures

How this unit is tested

What you have to know

14 practice questions

1. Which of the following intermolecular forces is present in all molecules and atoms, regardless of polarity?
   • London dispersion forces
   • Dipole-dipole forces
   • Hydrogen bonding
   • Ion-dipole forces
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   Answer. London dispersion forces

   Dispersion forces arise from temporary, instantaneous shifts in electron density that create momentary dipoles, so they exist in every atom or molecule. Dipole-dipole and hydrogen bonding require permanent polarity or specific N-H, O-H, or F-H bonds.
2. NH3 has a significantly higher boiling point than PH3, even though PH3 has a larger molar mass. Explain why.
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   Answer. NH3 molecules form hydrogen bonds because H is bonded directly to N, while PH3 cannot hydrogen bond since P-H bonds aren't polar enough and P isn't N, O, or F.

   Hydrogen bonding is a much stronger intermolecular force than the dispersion forces that dominate in PH3, so despite PH3's greater mass and stronger dispersion forces, NH3 requires more energy to separate its molecules and boils at a higher temperature.
3. Which of the following is NOT a correct assumption of the kinetic molecular theory of gases?
   • Gas particles have negligible volume compared to their container
   • Collisions between gas particles are perfectly elastic
   • Gas particles exert no attractive or repulsive forces on each other
   • Average kinetic energy is directly proportional to the volume of the gas
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   Answer. Average kinetic energy is directly proportional to the volume of the gas

   KMT states that average kinetic energy is proportional to Kelvin temperature, not volume. The other three statements (negligible particle volume, elastic collisions, no intermolecular forces) are core KMT assumptions.
4. A 2.00 L rigid container holds 0.500 mol of an ideal gas at 25°C. What is the pressure inside the container, in atm?
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   Answer. Approximately 6.12 atm

   Using PV=nRT: P = nRT/V = (0.500 mol)(0.0821 L·atm/mol·K)(298 K)/(2.00 L) ≈ 6.12 atm. Remember to convert 25°C to 298 K before substituting.
5. Under which of the following conditions would a real gas be expected to deviate most from ideal gas behavior?
   • High temperature and low pressure
   • Low temperature and high pressure
   • High temperature and high pressure
   • Low temperature and low pressure
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   Answer. Low temperature and high pressure

   At low temperature, particles move slowly enough for intermolecular attractions to matter; at high pressure, particles are forced close together so their finite volume becomes significant. Both conditions violate the KMT assumptions of negligible volume and no interactions.
6. In the van der Waals equation (P + an²/V²)(V − nb) = nRT, what physical property does the constant b account for?
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   Answer. The finite volume occupied by the gas particles themselves

   The term nb is subtracted from the container volume V because real gas particles are not point masses; they take up actual space, reducing the volume available for particle motion compared to the ideal case.
7. A solid is hard, brittle, has a very high melting point, and does not conduct electricity as a solid but does conduct when melted. This solid is best classified as:
   • Metallic solid
   • Ionic solid
   • Molecular solid
   • Covalent network solid
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   Answer. An ionic solid

   Ionic solids are held together by strong electrostatic attractions between fixed ions, making them hard and high-melting; they can't conduct as solids because ions are locked in place, but once molten the ions become mobile and can carry charge.
8. Gas A effuses 2.0 times faster than gas B. If gas B has a molar mass of 64.0 g/mol, what is the molar mass of gas A?
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   Answer. 16.0 g/mol

   Using Graham's law, rateA/rateB = √(MB/MA), so 2.0 = √(64.0/MA). Squaring gives 4.0 = 64.0/MA, so MA = 16.0 g/mol. Lighter gases effuse faster, consistent with A being lighter than B.
9. How many mL of a 6.00 M HCl stock solution are needed to prepare 250. mL of 1.50 M HCl?
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   Answer. 62.5 mL

   Using M1V1 = M2V2: (6.00 M)(V1) = (1.50 M)(250. mL), so V1 = (1.50 × 250.)/6.00 = 62.5 mL of stock, then diluted with water to a total of 250. mL.
10. As temperature increases, what happens to the solubility of a gas dissolved in a liquid, and why?
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    Answer. Solubility decreases, because dissolved gas molecules gain kinetic energy and escape the liquid more easily as temperature rises

    Unlike most solids, gas solubility in liquids falls with increasing temperature. Higher temperature increases the average kinetic energy of gas molecules, allowing more of them to overcome the intermolecular attractions holding them in solution and escape as vapor.
11. In a paper chromatography experiment, a pigment travels 4.2 cm from the origin while the solvent front travels 6.0 cm. What is the Rf value of the pigment?
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    Answer. 0.70

    Rf = distance traveled by solute ÷ distance traveled by solvent front = 4.2 cm / 6.0 cm = 0.70. A valid Rf value must fall between 0 and 1.
12. A solution has an absorbance of 0.450 at a wavelength where ε = 3.00×10³ M⁻¹cm⁻¹, measured in a 1.00 cm cuvette. What is the molar concentration of the solution?
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    Answer. 1.50×10⁻⁴ M

    Rearranging the Beer-Lambert law A = εbc gives c = A/(εb) = 0.450/(3.00×10³ × 1.00) = 1.50×10⁻⁴ M.
13. A solution appears blue to the eye. In which region of the visible spectrum is the solution primarily absorbing light?
    • Blue light
    • Green light
    • Orange/red light
    • Violet light
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    Answer. Orange/red light

    The color observed is the complement of the color absorbed. A solution that absorbs light in the orange-red region (roughly 600-650 nm) transmits the remaining blue light, which is what reaches the observer's eye.
14. Which separation technique would be most effective for separating a mixture of ethanol and water based on their different boiling points?
    • Chromatography
    • Distillation
    • Filtration
    • Decantation
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    Answer. Distillation

    Distillation separates components of a liquid mixture by heating until the lower-boiling-point substance (ethanol, ~78°C) vaporizes and is collected first, leaving the higher-boiling water behind. Chromatography separates based on differing affinities for stationary/mobile phases, not boiling point.

What people get wrong

1. Labeling any polar bond as 'hydrogen bonding' — hydrogen bonding only occurs when H is directly bonded to N, O, or F. Check for these specific bonds before invoking hydrogen bonding as the explanation.
2. Assuming a heavier molecule always has stronger intermolecular forces — dispersion forces do scale with molar mass, but a lighter polar or hydrogen-bonding molecule can still have a higher boiling point than a heavier nonpolar one. Compare the strongest force present, not just mass.
3. Using R with mismatched pressure units (plugging kPa into an equation set up for atm, or forgetting to convert Celsius to Kelvin) in PV=nRT. Always convert temperature to Kelvin and match R's units to the given pressure unit.
4. Believing gas solubility increases with temperature the way solid solubility usually does. Gas solubility in liquids decreases as temperature rises because dissolved gas molecules gain enough kinetic energy to escape the solution.
5. Inverting the Rf formula by dividing solvent-front distance by solute distance. Rf is always distance traveled by the substance divided by distance traveled by the solvent front, and must be between 0 and 1.
6. Treating 'high pressure' and 'low temperature' as independent conditions rather than both pushing toward non-ideal behavior — students often say a gas is 'ideal' at high pressure alone. Both conditions bring molecules close enough that volume and attractive forces matter.