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Chapter 15 1. Explain the difference/relationship between temperature, thermal energy, and kinetic energy. Answer: Temperature is a measure of the average kinetic energy of particles in a substance. Thermal energy is the total energy possessed by the particles in a substance due to their motion, including both kinetic and potential energy. Kinetic energy specifically refers to the energy of particles due to their motion and is directly related to their mass and velocity. 2. Explain why particles in gases move with a range of different velocities at a given temperature. Identify the Boltzmann distributions of particles at different temperatures, or for particles of different molecular weights at a specified temperature. Answer: Particles in gases move with a range of different velocities at a given temperature due to the distribution of kinetic energies. This distribution is described by the Boltzmann distribution, which shows that particles at a given temperature have a spread of kinetic energies. Higher temperatures correspond to higher average kinetic energies and a broader range of velocities. Additionally, at a specified temperature, particles of different molecular weights will have different kinetic energies, with lighter molecules generally having higher velocities. 3. Explain how temperature and kinetic energy are related, including the energy associated with vibration, rotation, and translation in different phases. Answer: Temperature and kinetic energy are directly related. As temperature increases, the average kinetic energy of particles increases as well. This increase in kinetic energy can manifest in different forms of molecular motion in different phases of matter. In gases, the kinetic energy is primarily associated with translation (linear motion) of particles. In liquids, there is also rotational motion, where molecules spin around their own axes. In solids, in addition to translation and rotation, there is vibrational motion, where atoms oscillate around fixed positions. 4. Explain the causes of water’s anomalous properties (high melting point, high boiling point, lower density of ice relative to liquid water, specific heat). Answer: Water's anomalous properties can be attributed to its hydrogen bonding and unique molecular structure. The high melting and boiling points of water are a result of the strong hydrogen bonds between water molecules, which require more thermal energy to break and transition to the liquid and gaseous states. The lower density of ice compared to liquid water is due to the arrangement of water molecules in an open hexagonal lattice structure in ice, creating larger spaces between the molecules. Water's specific heat capacity is relatively high because breaking and forming hydrogen bonds requires significant thermal energy.
5. Explain why the specific heat capacity of a substance is affected by the molecular-level structure. Answer: The specific heat capacity of a substance is affected by its molecular-level structure because different structures have varying abilities to store and release thermal energy. Substances with complex molecular structures or strong intermolecular forces tend to have higher specific heat capacities. This is because more energy is required to raise the temperature of these substances since the energy is used to overcome stronger intermolecular attractions. 6. Draw heating or cooling curves showing how the temperature changes when thermal energy is added to a substance (including a phase change). Explain why the temperature changes except during the phase change. Answer: During a heating or cooling process, the temperature changes as thermal energy is added or removed from a substance. The temperature increases as thermal energy is added until it reaches the boiling or melting point of the substance. At the boiling or melting point, the temperature remains constant despite the continuous addition of thermal energy because the energy is used to overcome intermolecular forces and facilitate the phase change. Once the phase change is complete, the temperature resumes changing as thermal energy is further added or removed. 7. Define and give examples of open, closed, and isolated systems. Answer: An open system allows both energy and matter to be exchanged with the surroundings. Examples include a pot of boiling water with water vapor escaping or an organism exchanging heat and nutrients with its environment. A closed system allows the exchange of energy but not matter with the surroundings. An example is a sealed container with a fixed amount of gas that can undergo temperature changes. An isolated system does not allow the exchange of energy or matter with the surroundings. The universe is often considered an isolated system since it is not influenced by external factors. 8. Explain the difference between state and path functions and give examples. Answer: State functions are properties that depend only on the current state of the system and are independent of the path taken to reach that state. Examples include temperature, pressure, volume, and internal energy. Path functions, on the other hand, depend on the path taken to reach a particular state. Examples include heat (q) and work (w), as they are determined by the specific process or path taken during a system's transformation.
9. Identify the direction of the thermal energy change and the sign of q or ΔH for exothermic and endothermic processes. Answer: In an exothermic process, thermal energy flows out of the system into the surroundings. Therefore, q and ΔH have negative signs since the system loses energy. In an endothermic process, thermal energy flows into the system from the surroundings. Consequently, q and ΔH have positive signs since the system gains energy. 10. Identify the direction of the thermal energy change and the sign of q or ΔH for a phase change. Answer: During a phase change, the direction of the thermal energy change depends on the specific phase change. For example, during melting or vaporization, the process is endothermic, so q and ΔH will be positive since the system absorbs energy from the surroundings. On the other hand, during freezing or condensation, the process is exothermic, so q and ΔH will be negative since the system releases energy to the surroundings. 11. For exothermic and endothermic processes, identify the direction of the thermal energy change and the sign of q or ΔH. Answer: In exothermic processes, the thermal energy change is from the system to the surroundings, resulting in a negative sign for q and ΔH. In endothermic processes, the thermal energy change is from the surroundings to the system, leading to a positive sign for q and ΔH. 12. For phase changes, identify the direction of the thermal energy change and the sign of q or ΔH. Answer: During phase changes, the thermal energy change depends on the specific phase change. If it involves a transition from a less-ordered phase to a more-ordered phase (e.g., gas to liquid or liquid to solid), the process is exothermic, and q and ΔH will be negative since the system releases energy. Conversely, if it involves a transition from a more-ordered phase to a less-ordered phase (e.g., liquid to gas or solid to liquid), the process is endothermic, and q and ΔH will be positive since the system absorbs energy. 13. Explain the role of probability in entropy changes. Answer: Probability plays a fundamental role in entropy changes. Entropy is a measure of the system's disorder or the number of microstates (possible arrangements of particles) that correspond to a particular macrostate (observed state). As the number of microstates increases, the system becomes more disordered, and the probability of finding the system in a particular macrostate increases, resulting in higher entropy.
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