In an ideal gas mixture, which statement is true about partial pressures?

• A. It equals the total pressure.
• B. It is independent of the amount of gas present.
• C. It equals the mole fraction times the total pressure.
• D. It is equal to the concentration in mol/L.

Answer: (C)

Partial pressures in an ideal gas mixture come from Dalton's law: the total pressure is the sum of the individual partial pressures, and each gas contributes a portion of the total pressure proportional to how much of that gas is present. The partial pressure of a gas is P_i = X_i P_total, where X_i is the mole fraction of that gas in the mixture (X_i = n_i / n_total). Since X_i equals the number of moles of that gas divided by the total moles, you get P_i = (n_i / n_total) P_total. This is why the statement that a gas’s partial pressure equals its mole fraction times the total pressure is the correct description.

This makes sense because if there were only one gas, its partial pressure would equal the total pressure, and if you add more of a particular gas, its partial pressure increases accordingly. It’s not simply equal to a concentration in mol/L, although you can relate pressure to amount via PV = nRT; partial pressure itself is a pressure, not a concentration.