AP Intermolecular Forces and Properties 2 — Questions and Answers

Question 1: Which of the following pairs of substances would be expected to be miscible (mix completely) in each other?

Ethanol (C₂H₅OH) and water (H₂O) (Correct answer)
Hexane (C₆H₁₄) and water (H₂O)
Octane (C₈H₁₈) and water (H₂O)
Carbon tetrachloride (CCl₄) and water (H₂O)

Correct answer: Ethanol (C₂H₅OH) and water (H₂O)

Ethanol and water are both polar and can form hydrogen bonds with each other ('like dissolves like'). Hexane, octane, and CCl₄ are nonpolar and immiscible with polar water.

'Like dissolves like' is the fundamental principle governing miscibility and solubility, rooted in intermolecular forces. Water is highly polar and an excellent hydrogen bond donor and acceptor. For a substance to dissolve in water, it must be able to form favorable interactions with water molecules that replace the water-water and solute-solute interactions being broken. Ethanol (CH₃CH₂OH): has a polar –OH group that forms hydrogen bonds with water AND the hydrophobic ethyl group is small enough that overall dissolution is favorable. Ethanol is completely miscible with water. Hexane (C₆H₁₄): nonpolar, only London dispersion forces. It cannot form H-bonds with water. The energy cost of disrupting water's H-bond network far exceeds the energy gained from weak water-hexane LDF interactions → immiscible (two layers form). CCl₄: nonpolar despite polar C–Cl bonds (tetrahedral geometry, bond dipoles cancel). Immiscible with water but miscible with other nonpolar solvents. This principle is critical for pharmaceutical applications (drug solubility), environmental chemistry (oil spills), and chemical separations (extraction).

Question 2: Arrange the following substances in order of increasing boiling point: CH₄, NH₃, Ne, HF.

Ne < CH₄ < NH₃ < HF (Correct answer)
CH₄ < Ne < HF < NH₃
Ne < CH₄ < HF < NH₃
CH₄ < Ne < NH₃ < HF

Correct answer: Ne < CH₄ < NH₃ < HF

Ne (bp −246°C): only tiny LDF. CH₄ (bp −161°C): small LDF, no dipole. NH₃ (bp −33°C): strong H-bonding and dipole-dipole. HF (bp +19.5°C): very strong H-bonding (F is most electronegative). Increasing bp: Ne < CH₄ < NH₃ < HF.

Boiling points reflect the energy needed to overcome intermolecular forces and convert liquid to vapor. Ne (MM = 20 g/mol): noble gas with only very weak London dispersion forces. bp = −246°C. CH₄ (MM = 16 g/mol): slightly heavier than Ne, nonpolar, only LDF. bp = −161°C. NH₃ (MM = 17 g/mol): polar molecule with hydrogen bonding capability (N–H···N bonds, though N is less electronegative than O or F). bp = −33°C. HF (MM = 20 g/mol): though similar in molar mass to Ne and NH₃, the extremely electronegative F creates very strong F–H···F hydrogen bonds. bp = +19.5°C. Ordering: Ne (−246°C) < CH₄ (−161°C) < NH₃ (−33°C) < HF (+19.5°C). Interestingly, within the hydrogen halides, HF has the highest boiling point despite being the lightest, while HI has higher bp than HBr and HCl due to larger LDF from more electrons. This anomaly across groups highlights H-bonding exceptions.

Question 3: Why does ice float on liquid water, which is unusual for solids compared to their liquid phases?

Ice has a less dense crystal structure due to directional H-bonding that holds molecules in an open hexagonal lattice, making solid water less dense than liquid water (Correct answer)
Ice has higher kinetic energy than liquid water, causing expansion
Water molecules are smaller in the solid phase due to compressed H-bonds
Ice floats due to surface tension, not density differences

Correct answer: Ice has a less dense crystal structure due to directional H-bonding that holds molecules in an open hexagonal lattice, making solid water less dense than liquid water

In ice, each water molecule forms 4 directional hydrogen bonds in a hexagonal lattice with large open spaces. This arrangement (density ~0.917 g/mL) is less dense than liquid water (~1.00 g/mL) where H-bond networks are more disordered and compact.

Water's anomalous behavior — solid less dense than liquid — has profound consequences for aquatic ecosystems and Earth's climate. In liquid water at 0–4°C, molecules form a dynamic, partially disordered H-bond network. On average, each water molecule has 3.5–4 H-bonds, but they are constantly breaking and reforming, allowing molecules to pack more closely (density ≈ 0.998–1.000 g/mL). Maximum density is at 4°C. In ice (hexagonal ice Ih), every water molecule forms exactly 4 directional hydrogen bonds (tetrahedral geometry) in a rigid, open hexagonal crystal lattice. This perfect tetrahedral arrangement creates large empty spaces (hexagonal channels), making ice less dense (≈ 0.917 g/mL) than liquid water. When water freezes, it expands ~9% in volume. This has critical consequences: 1. Ice floats on water, insulating aquatic ecosystems and preventing complete freezing of lakes. 2. Freezing water in rock crevices can cause weathering (frost wedging). 3. Water pipes burst when water freezes inside them. Most substances are denser in the solid phase than liquid (molecules pack more tightly when stationary), making water's behavior truly anomalous.

Question 4: Which of the following has the highest viscosity at room temperature?

Glycerol (HOCH₂CHOHCH₂OH) (Correct answer)
Ethanol (C₂H₅OH)
Water (H₂O)
Hexane (C₆H₁₄)

Correct answer: Glycerol (HOCH₂CHOHCH₂OH)

Glycerol has three –OH groups that form extensive hydrogen bonds both within and between molecules, creating a highly viscous liquid. The other compounds have fewer H-bonding sites or only weak LDF, resulting in lower viscosity.

Viscosity measures a liquid's resistance to flow — its 'thickness' or 'stickiness.' High viscosity results from strong and/or extensive intermolecular forces that resist the movement of molecules past each other. Glycerol (propane-1,2,3-triol) has three –OH groups, each capable of both donating and accepting hydrogen bonds. This creates an extensive H-bond network between molecules, requiring significant energy to overcome — very high viscosity (≈ 1412 mPa·s at 20°C). Water: two H-bond donor/acceptor sites. Moderate H-bonding but low molar mass → relatively low viscosity (1.002 mPa·s at 20°C). Ethanol: one –OH group, moderate H-bonding. Viscosity ≈ 1.074 mPa·s — slightly higher than water due to larger molecular size but fewer H-bonds per unit volume. Hexane: nonpolar, only weak LDF. Very low viscosity (0.294 mPa·s). Flows freely. Order of viscosity: Glycerol >> Ethanol ≈ Water >> Hexane. Glycerol's high viscosity makes it useful as a lubricant, in pharmaceuticals (cough syrup base), and in cosmetics. Honey and motor oil are other examples of highly viscous liquids with extensive molecular interactions.

Question 5: What is the primary intermolecular force responsible for DNA's double helix structure holding together the two strands?

Hydrogen bonding between complementary base pairs (A-T and G-C) (Correct answer)
Covalent bonding between bases on opposite strands
Ionic interactions between positively and negatively charged bases
London dispersion forces between stacked bases

Correct answer: Hydrogen bonding between complementary base pairs (A-T and G-C)

The two DNA strands are held together by hydrogen bonds between complementary base pairs: adenine (A) forms 2 H-bonds with thymine (T), and guanine (G) forms 3 H-bonds with cytosine (C). These H-bonds allow separation during replication and transcription.

DNA's double helix structure is stabilized by two types of intermolecular interactions, but H-bonding between base pairs is the primary force for strand complementarity: 1. Hydrogen bonding (base pairing): A–T base pairs share 2 H-bonds; G–C base pairs share 3 H-bonds. These H-bonds hold the two anti-parallel strands together via specific Watson-Crick base pairing. G-C pairs are stronger (3 H-bonds vs. 2), which is why GC-rich regions of DNA have higher melting temperatures. 2. Base stacking (π-π stacking / London dispersion forces): The aromatic base rings stack on top of each other along the helix axis, providing additional stabilization through hydrophobic effects and dispersion forces. This contributes significantly to overall stability. The H-bond nature of base pairing is critical for biology: (1) H-bonds are strong enough to maintain double-strand stability at physiological temperatures; (2) they are weak enough to be broken by helicases during replication and transcription, with minimal energy expenditure; (3) they are highly specific (A only pairs with T; G only with C) due to geometric complementarity, ensuring faithful genetic information transmission.

Question 6: A substance has a high melting point, conducts electricity only when molten or dissolved in water, and is hard but brittle. What type of solid is it?

Ionic solid (Correct answer)
Metallic solid
Covalent network solid
Molecular solid

Correct answer: Ionic solid

Ionic solids have high melting points (strong electrostatic forces), conduct electricity only when ions are free to move (molten or dissolved), and are hard (rigid lattice) but brittle (layers crack when shifted, aligning like charges).

This question tests the ability to identify solid types from physical properties. Ionic solids (e.g., NaCl, MgO, CaCO₃) are characterized by: 1. High melting points: Strong electrostatic attractions between oppositely charged ions require significant energy to overcome. MgO (mp 2852°C) > NaCl (mp 801°C). 2. Electrical conductivity: In the solid state, ions are locked in the lattice and cannot move → non-conducting. When melted or dissolved, ions become mobile → conducts electricity. This is a diagnostic property. 3. Hardness: The rigid, ordered lattice structure resists deformation. 4. Brittleness: When force is applied, ions shift so like charges become adjacent → repulsion causes fracture (cleavage planes). Compare with: - Metallic solids: Conduct electricity as solids (delocalized e⁻), malleable/ductile (not brittle). - Covalent network: Very high mp, hard, non-conducting (usually), not brittle in the same way as ionic. - Molecular solids: Low mp, soft, non-conducting (no free ions or electrons). The key distinguishing feature for ionic solids is conductivity only when molten or dissolved.

AP Intermolecular Forces and Properties 2

Which of the following pairs of substances would be expected to be miscible (mix completely) in each other?