When a substance is at its melting point, any added heat will go into the phase change (an endothermic process) before causing a temperature increase. As long as both solid and liquid phases are present in equilibrium, the temperature of the system will remain constant at the melting point (0°C for water).

During a phase change, the temperature of a pure substance remains constant. The added energy, known as latent heat of vaporization, is used to overcome intermolecular forces and convert the liquid to a gas, rather than to increase the average kinetic energy (temperature) of the molecules.

Freezing is an exothermic process where liquid turns into a more ordered solid. Energy is released as stronger, more stable intermolecular forces (hydrogen bonds in an ice lattice) are formed. This means the potential energy of the system decreases. Because the phase change occurs at a constant temperature (0°C), the average kinetic energy of the molecules does not change.

First, calculate the moles of benzene that can be frozen. Since heat is removed, $$q = -4.975 \text{ kJ}$$. Using $$q = n \Delta H_{\text{solidification}}$$, we find $$n = \frac{q}{\Delta H_{\text{solidification}}} = \frac{-4.975 \text{ kJ}}{-9.95 \text{ kJ/mol}} = 0.500 \text{ mol}$$. Then, convert moles to grams: $$mass = 0.500 \text{ mol} \times 78.11 \text{ g/mol} = 39.1 \text{ g}$$.

During a phase change, the added or removed energy alters the potential energy of the molecules by changing the distance between them and the strength of their intermolecular interactions. The temperature, which is a measure of the average kinetic energy of the molecules, remains constant until the phase change is complete.

The energy required for melting is determined by the strength of the forces holding the particles in the solid lattice. NaCl is an ionic solid with strong electrostatic attractions (ionic bonds) between ions. Methane is a molecular solid with only weak London dispersion forces between molecules. Overcoming the strong ionic bonds in NaCl requires far more energy than overcoming the weak dispersion forces in solid methane. Intramolecular covalent bonds are not broken during phase changes.

Fusion (melting) involves weakening the intermolecular forces enough for molecules to move past each other, but significant attractions remain. Vaporization involves completely overcoming the intermolecular forces to separate molecules into the gas phase. This requires much more energy, hence $$\Delta H_{\text{vap}}$$ is much larger than $$\Delta H_{\text{fus}}$$. Phase changes involve changes in potential energy, not kinetic energy, so C is incorrect.

This question tests the calculation of energy released during freezing (liquid to solid transition). When benzene freezes, it releases energy equal to its enthalpy of fusion multiplied by the number of moles. The energy released = ΔHfus × moles = 9.95 kJ/mol × 0.80 mol = 7.96 kJ. A common error is to use the enthalpy value directly without accounting for the partial mole (choice C: 9.95 kJ), treating the molar enthalpy as if it were the total energy for any amount. For phase transitions, always multiply the per-mole enthalpy value by the actual number of moles present.

This question tests the ability to calculate the energy absorbed during melting using the enthalpy of fusion and the mass of the substance. Convert 27.0 g of ice to moles with 18.0 g/mol, yielding 1.50 moles. Multiply by 6.01 kJ/mol to get 9.02 kJ absorbed, reflecting the endothermic nature of breaking hydrogen bonds in ice. This matches choice B, the energy for complete melting at 0°C. Choice C, 6.01 kJ, tempts those who omit the mole conversion, confusing the per-mole value with the total energy. A transferable strategy is to identify whether the phase change is endothermic or exothermic by considering if it's increasing or decreasing molecular disorder.

This question tests the ability to calculate the energy absorbed during sublimation using the enthalpy of sublimation and the mass of the substance. Convert 10.0 g of iodine to moles using 254 g/mol, resulting in about 0.0394 moles. Multiply by 62.4 kJ/mol to find approximately 2.46 kJ absorbed, as sublimation is endothermic, directly transitioning solid to gas and requiring energy to break bonds. This corresponds to choice B, the energy input for the phase change. Choice D, 62.4 kJ, attracts those who skip the mole calculation and use the enthalpy value alone, mistakenly treating it as per gram instead of per mole. A transferable strategy is to remember that enthalpies of phase changes are molar values, so always scale by the number of moles involved.