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Consider the equilibrium system shown - HSC - SSCE Chemistry - Question 36 - 2022 - Paper 1
Question 36
Consider the equilibrium system shown. H2O(l) ⇌ H2O(g) In a laboratory at 23°C, a 100 mL sample of water is held in a beaker and another 100 mL sample is held in a... show full transcript
Worked Solution & Example Answer:Consider the equilibrium system shown - HSC - SSCE Chemistry - Question 36 - 2022 - Paper 1
Step 1
Explain the differences in evaporation for the two samples
Answer
At room temperature, the 100 mL sample of water in the beaker will evaporate more readily than the sealed bottle. This is due to the fact that the water in the beaker is open to the atmosphere, allowing water molecules to escape freely into the air, whereas the water in the sealed bottle is restricted from evaporating as it reaches an equilibrium state with no additional space for the water vapor.
Thus, the beaker will continuously lose water to the gas phase, while the sealed system will maintain a constant level of liquid water and vapor due to its closed nature.
Step 2
Consider changes in enthalpy and entropy
Answer
The evaporation process is endothermic, meaning it requires energy (heat) from the surroundings. For both samples, the enthalpy change ( ( \Delta H )) during vaporization is positive, as heat is absorbed.
However, the entropy ( ( S )) change differs between the two systems. In the beaker, as water evaporates, the entropy increases significantly due to the release of water molecules into the gas phase, resulting in a greater disorder in the system. In contrast, for the sealed bottle, the entropy change is less pronounced since the closed environment limits the movement of water vapor molecules, maintaining some degree of order compared to the open system.
Therefore, evaporation leads to a larger increase in entropy in the beaker than in the sealed bottle.