The Correct Answer and Explanation is:
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The correct answer is D. Neither Technician A nor B.
Here is an explanation of why both technicians are incorrect.
Technician A’s statement is incorrect because it misrepresents the fundamental properties of a vapor, which is the gaseous state of a substance. A key characteristic of gases is that they do not have a fixed volume or a fixed shape. The molecules in a gas are spread far apart and move randomly, allowing the gas to expand and fill the entire volume of any container it is placed in. While it is possible to hold a vapor within a rigid container of a fixed volume, the vapor itself does not inherently possess a fixed volume that remains constant regardless of conditions. For instance, if the container were flexible like a balloon, increasing the temperature would cause the gas to expand, thus increasing its volume. Technician A’s assertion that vapor has a fixed volume is fundamentally flawed.
Technician B’s statement is also incorrect as it contradicts the established relationship between the temperature and pressure of a gas, described by Gay-Lussac’s Law. This law, a component of the Ideal Gas Law (PV=nRT), states that for a given amount of gas held at a constant volume, the pressure is directly proportional to its absolute temperature. This means that as the temperature of a gas increases, its molecules gain kinetic energy, move faster, and collide with the container walls more frequently and with greater force, resulting in higher pressure. Conversely, decreasing the temperature causes the molecules to slow down, leading to less frequent and weaker collisions, which in turn decreases the pressure. Technician B claims the opposite, that decreasing the temperature increases the pressure, which is physically incorrect.
Since both technicians have made statements that defy the basic principles of gas behavior, neither of them is correct.
