- In a closed system, if the temperature of a gas increases, which of the following will occur if the volume is held constant?
A) The pressure decreases.
B) The pressure increases.
C) The number of moles decreases.
D) The temperature does not change.
Answer: B) The pressure increases.
Explanation: If the temperature increases at constant volume, the pressure of the gas must increase according to Gay-Lussac’s Law.
- The distribution of molecular speeds in a gas is influenced by which factor?
A) Type of gas only.
B) Temperature only.
C) Both type of gas and temperature.
D) Volume only.
Answer: C) Both type of gas and temperature.
Explanation: The type of gas (molar mass) and the temperature affect the distribution of molecular speeds, altering their average speeds and energy.
- In the kinetic theory of gases, the term “ideal gas” refers to:
A) A gas that does not exist in reality.
B) A gas that perfectly follows the ideal gas law under all conditions.
C) A gas that behaves according to the kinetic theory assumptions.
D) A gas with strong intermolecular forces.
Answer: C) A gas that behaves according to the kinetic theory assumptions.
Explanation: An ideal gas is one that follows the assumptions of the kinetic theory closely, meaning negligible intermolecular forces and volume.
- In a mixture of gases, the total pressure is equal to:
A) The pressure of the heaviest gas only.
B) The pressure of the lightest gas only.
C) The sum of the partial pressures of all gases present.
D) The average of the partial pressures.
Answer: C) The sum of the partial pressures of all gases present.
Explanation: Dalton’s Law states that the total pressure of a mixture is the sum of the partial pressures exerted by each individual gas.
- When gases are mixed, the pressure exerted by the mixture is the sum of the partial pressures of each gas. This principle is known as:
A) Dalton’s Law of Partial Pressures.
B) Avogadro’s Law.
C) Graham’s Law.
D) Boyle’s Law.
Answer: A) Dalton’s Law of Partial Pressures.
Explanation: Dalton’s Law states that the total pressure of a gas mixture is equal to the sum of the partial pressures of its individual components.
- Which of the following is true for the relationship between temperature and kinetic energy?
A) Kinetic energy increases as temperature decreases.
B) Kinetic energy is independent of temperature.
C) Kinetic energy is proportional to the square of the temperature.
D) Kinetic energy increases as temperature increases.
Answer: D) Kinetic energy increases as temperature increases.
Explanation: The kinetic energy of gas molecules is directly proportional to the absolute temperature of the gas.
- Which of the following factors would NOT affect the rate of diffusion of a gas?
A) Molar mass.
B) Temperature.
C) Pressure.
D) The color of the gas.
Answer: D) The color of the gas.
Explanation: The rate of diffusion is affected by molar mass, temperature, and pressure, but not by the color of the gas.
- According to kinetic theory, the average velocity of gas molecules is:
A) Constant for all temperatures.
B) Dependent on pressure only.
C) Dependent on temperature and molar mass.
D) Always equal to zero.
Answer: C) Dependent on temperature and molar mass.
Explanation: The average velocity is influenced by both the temperature (higher temperature means higher velocity) and the molar mass (lighter gases have higher velocities).
- If the temperature of a gas in a rigid container is doubled, what will happen to the pressure?
A) It will remain the same.
B) It will double.
C) It will triple.
D) It will quadruple.
Answer: B) It will double.
Explanation: If the temperature is doubled in a rigid container (constant volume), the pressure also doubles due to the direct relationship established by Gay-Lussac’s Law.
- In a real gas, which of the following can occur at high pressures?
A) Increased volume.
B) Decreased temperature.
C) Significant deviations from ideal behavior.
D) Constant pressure.
Answer: C) Significant deviations from ideal behavior.
Explanation: At high pressures, real gases experience interactions between molecules and volume effects, leading to deviations from the ideal gas behavior.
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