Which of the following statements about the partial pressures of gases is true?

• A. The total pressure of a gas mixture is the sum of the partial pressures of each individual gas.
• B. Partial pressure is the pressure a gas would exert if it occupied the entire volume of the mixture at the same temperature.
• C. The partial pressure of a gas is independent of its mole fraction in a mixture.
• D. Partial pressures are only applicable to ideal gases and not real gases.

Answer: (B)

Verified step by step guidance

1. Understand the concept of partial pressure: Partial pressure is the pressure that a single gas in a mixture of gases would exert if it occupied the entire volume by itself at the same temperature. This is a fundamental concept in gas mixtures.
2. Apply Dalton's Law of Partial Pressures: According to Dalton's Law, the total pressure of a gas mixture is the sum of the partial pressures of each individual gas. This can be expressed mathematically as: $P_{\mathrm{total}}=P_1+P_2+\dots$, where $P_i$ represents the partial pressure of each gas.
3. Consider the relationship between partial pressure and mole fraction: The partial pressure of a gas in a mixture is directly proportional to its mole fraction. This can be expressed as: $P_i=X_i{P}_{\mathrm{total}}$, where $X_i$ is the mole fraction of the gas.
4. Evaluate the applicability to real gases: While partial pressures are most accurately applied to ideal gases, they can also be used for real gases under certain conditions, typically at low pressures where the behavior of real gases approximates that of ideal gases.
5. Review the statements: Based on the explanations above, identify which statements are true regarding partial pressures. The correct statements are: 'The total pressure of a gas mixture is the sum of the partial pressures of each individual gas' and 'Partial pressure is the pressure a gas would exert if it occupied the entire volume of the mixture at the same temperature.'